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Transcript
Table of Contents Unit 1- Matter and Change ............................................................................1 Pages - 12 Unit 2- Atomic Theory and Structure ............................................................13 - 33 Unit 3- Electrons and Periodicity...................................................................34 - 52 Unit 4- Bonding .............................................................................................53 - 67 Unit 5- Nomenclature ....................................................................................68 - 87 Unit 6- Chemical Reactions ...........................................................................88 - 101 Unit 7- Stoichiometry ....................................................................................102 - 114 Unit 8- Gas Laws ...........................................................................................115 - 132 Unit 9- Solids, Liquids and Phase Changes ...................................................133 - 144 Unit 10- Solutions and Solubility ..................................................................145 - 156 Unit 11- Acids and Bases ...............................................................................157 - 174 Unit 12- Kinetics and Thermochemistry........................................................175 - 193 Unit 13- Oxidation and Reduction .................................................................194 - 203 Unit 1- Matter and Change CHEMISTRY: THE STUDY OF MATTER Chemistry is the science that investigates and explains the structure and properties of ________________. This includes its composition, properties and the changes it undergoes. SCIENTIFIC METHOD A scientific method is a systematic approach to answer a ______________________ or study a situation. It starts with _____________________ - noting and recording facts. A _________________________ is a possible explanation for what has been observed. It is an educated _________________ as to the cause of the problem or answer to the question. An experiment is a set of controlled observations that _____________ a hypothesis. The variable that is changed in an experiment is called the ________________________ variable. The variable that you watch to see how it _________________ as a result of your changes to the independent variable is called the dependent variable. The cycle (hypothesis followed by experimentation) repeats many times, and the hypothesis gets more and more certain. The hypothesis becomes a _______________________, which is a thoroughly tested model that explains why things behave a certain way. Theories can never be ____________________; they are always subject to additional research. Another outcome is that certain behavior is repeated many times. A scientific ____________ describes a relationship in nature that is supported by many experiments and for which no exception has been found. Identify the dependent variable and the independent variable in the following experiments. a) A student tests the ability of a given chemical to dissolve in water at three different temperatures. independent variable: _________________________________ dependent variable: _________________________________ b) A farmer compares how his crops grow with and without phosphorous fertilizers. independent variable: _________________________________ dependent variable: _________________________________ 1 Unit 1- Matter and Change MATTER Matter is anything that takes up __________________ and has mass. ______________ is the measure of the amount of matter that an object contains. Virtually all of the matter around us consists of mixtures. A mixture can be defined as something that has _____________________ composition. Soda is a mixture (carbon dioxide is dissolved in it), and ____________________ is a mixture (it can be strong, weak or bitter). If matter is not uniform throughout, then it is a _______________________ mixture. If matter is uniform throughout, it is homogeneous. Homogeneous mixtures are called ___________________. A heterogeneous mixture contains regions that have ____________________ properties from those of other regions. When we pour sand into water, the resulting mixture contains two distinct regions. ___________________ pavement, which has small rocks mixed with tarry goo, is a simple example of a heterogeneous mixture. Oil-and-vinegar salad dressing, which has a layer of oil floating on a layer of vinegar, is another example. Homogeneous mixtures (also known as solutions) are mixtures in which the composition is _______________________, there are no chunks or layers. Salt water, ___________________ ___________________ and dust free air (mixture of nitrogen, oxygen, argon, carbon dioxide, water vapor and other gases) are examples of homogeneous mixtures. Brass (solid mixture of copper and ______________) is also a homogeneous mixture. Brass is a(n) _________________, which is a mixture of metals. Since heterogeneous mixtures contain chunks or layers, they are often easier to separate than homogeneous mixtures. A mixture of solid particles in a liquid can be separated by pouring the mixture through a ___________________ that traps the solid particles while the liquid passes through in a process called filtering. Some simple methods also exist for separating homogeneous mixtures. A solid dissolved in a liquid solution can be separated by letting it dry out in the process of ___________________. Mixtures are separated into pure _____________________. A pure substance always has the same composition. Pure substances are either elements or _________________________. Elements are substances that cannot be broken down into other substances chemically or _______________________. Examples include sodium, carbon and aluminum. Compounds are substances made of two or more ______________________ combined chemically. Compounds have properties ___________________________ from those of the original elements. Examples of compounds include water (hydrogen and oxygen) and table salt (sodium and chlorine). 2 Unit 1- Matter and Change Classify each of the following as a pure substance, a homogeneous mixture or a heterogeneous mixture. A. gasoline ____________________ B. copper metal ____________________ C. a stream with gravel at the bottom ____________________ D. maple syrup _________________ E. chunky peanut butter _____________ F. common salt ____________________ G. margarine ____________________ H. a Spanish omelet ________________ I. a multivitamin tablet ________________ J. oxygen gas ____________________ K. carbon dioxide gas _________________ PROPERTIES The properties of matter describe the characteristics and behavior of matter, including the changes that matter undergoes. _____________________ properties are characteristics that a sample of matter exhibits without any change in its identity. This property can be observed and measured without _____________________ the substance. Examples of the physical properties of a chunk of matter include its: 1. __________________________________ 2. _________________________________ 3. __________________________________ 4. _________________________________ 5. __________________________________ 6. _________________________________ 7. __________________________________ Chemical properties are those that can be observed only when there is a change in the ___________________________ of the substance. Rusting is a chemical reaction in which iron combines with __________________ to form a new substance, iron (III) oxide. Classify each of the following as a chemical or physical property. density ___________________________ reactivity ___________________________ color _____________________________ melting point ________________________ Using the Chemistry Reference Tables, which substance has a A. density = 19.31 g/cm3 _____________________ B. melting point = -119°C _____________________ C. boiling point = 65°C _____________________ D. melting point = -73°C _____________________ 3 Unit 1- Matter and Change Using the Chemistry Reference Tables, are the following substances soluble or insoluble in water? A. zinc nitrate _____________________ B. sodium sulfate _____________________ C. calcium carbonate _____________________ D. potassium oxide _____________________ E. lead (II) fluoride _____________________ F. barium hydroxide _____________________ G. copper (II) sulfide _____________________ H. silver chloride _____________________ CHANGES A physical change is a change in matter that does not involve a change in the chemical identity of individual substances. The matter only changes in appearance. Examples: ______________, _________________, __________________, _________________, ___________________, and _____________________. A chemical property always relates to a chemical change, the change of one or more substances _____________ other substances. Another term for chemical change is chemical ___________________. Indications of a chemical reaction: __________________ absorbed or released, _________________ change, formation of a precipitate - ______________ that separates from solution, and formation of a ___________. All matter is made of atoms, and any chemical change involves only a rearrangement of the atoms. Atoms do not just appear. Atoms do not just disappear. This is an example of the law of conservation of mass (or matter), which says that in a chemical change, matter is neither ________________ nor destroyed. All chemical changes also involve some sort of energy change. Energy is either taken in or __________________ ____________ as the chemical change takes place. Energy is the capacity to do _________________. Work is done whenever something is moved. Chemical reactions that give off heat energy are called ____________________ reactions. Chemical reactions that _________________ heat energy are called endothermic reactions. Freezing, condensation and ___________________ are exothermic. Melting, _______________________ and sublimation are endothermic. 4 Unit 1- Matter and Change State whether each of the following is an endothermic or exothermic process. 1. melting of ice __________________________ 2. combustion of gasoline __________________________ 3. Natural gas is burned in a furnace. __________________________ 4. When solid potassium bromide is dissolved in water, the solution gets colder. __________________________ 5. When concentrated sulfuric acid is added to water, the solution gets very hot. __________________________ In an endothermic graph, the ____________________ have greater energy than the reactants. The change in energy is a ______________________ value. In a(n) ____________________ graph, the reactants have greater energy than the products. The change in ________________ is a negative value. Sketch endothermic and exothermic graphs below. Label the reactants, products and activation energy. Endothermic Exothermic E n e r g y E n e r g y Reaction Progress Reaction Progress Conservation of Energy: Energy can be neither created nor destroyed in ordinary changes (not nuclear); it can only change _______________. 5 Unit 1- Matter and Change DENSITY Density is the amount of matter (mass) contained in a unit of ___________________. Styrofoam has a low density or small mass per unit of volume. density mass volume D m V Solve the following density problems. 1. The density of sugar is 1.59 g/cm3. Calculate the mass of sugar in 15.0 ml. (1 mL = 1 cm3). 2. The density of helium is 0.178 g/L. Calculate the volume of helium that has a mass of 23.5 g. 3. A 14.95 g sample of gold has a volume of 0.774 cm3. Calculate the density of gold. 4. Balsa wood has a density of 0.12 g/cm3. What is the mass of a sample of balsa wood if its volume is 134 cm3? 5. The density of ice at 0°C is 0.917 g/cm3. Calculate the volume of 145 g of ice. 6. The density of a sample of lead is found by the process of water displacement. A graduated cylinder is filled with water to the 30.0 mL mark. The cylinder with the water is placed on an electronic balance and weighs 106.82 g. A piece of lead is added to the cylinder. The cylinder is reweighed with the water and the lead and the scale reads 155.83 g. The volume of all the material in the cylinder is 34.5 mL. Calculate the density of the lead. 7. The density of an unknown solid was found by the process of water displacement. The object was massed on an electronic balance. The balance reads 125 g. 50.0 cm3 of water was poured into a 100.0 mL graduated cylinder. The unknown sample was then gently placed into the graduated cylinder. The volume in the cylinder rose to 60.7 cm3. Calculate the density of the unknown solid. Practice / Homework Reference Packet Study Density: Identify the substance based on the density value given D = m / V 1. D = 0.66g/cm3 3. M = 20 g, V = 4.44 cm3 2. D = 2.702g/cm3 4. M = 3 g, V = 2.1 L 6 Unit 1- Matter and Change Melting and Boiling points: Identify the substance based on the given temperature value. 5. Melting point = 801oC 7. Melting point = 1455oC 6. Boiling point = 79oC 8. Boiling point = 1413oC Solubility: Identify if the substance is soluble or insoluble. 9. Lithium sulfate 11. Lead (IV) bromide 10. Strontium oxide 12. Ammonium carbonate Identify each of the following as an element, a compound, a homogeneous mixture or a heterogeneous mixture. 13. Water 16. Silver 14. Cheerios in milk 17. Salsa 15. Apple juice 18. A bag of nuts and bolts Identify each of the following as a chemical or physical property 19. Combustible 21. Volume 20. Mass 22. Ability to rust Identify each of the following as a chemical or physical change 23. Melts 26. Rips 24. Burns 27. Tarnishes 25. Dissolves 28. Shatters 7 Unit 1- Matter and Change Density Practice: Solve each problem below, writing the equation and showing the substitution. Provide a unit for each answer. 1. A block of aluminum occupies a volume of 15.0 mL and weighs 40.5 g. What is its density? 2. Find the mass of gold that occupies 965 cm3 of space. 3. Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the cylinder weighs 306.0 g. From this information, calculate the density of mercury. 4. Find the volume occupied by 250.0 g of O2. 5. A cube of metal has a side length of 1.55 cm. If the sample is found to have a mass of 26.7 g, find the density and identity of the metal. 6. An irregularly-shaped sample of aluminum (Al) is put on a balance and found to have a mass of 43.6 g. The student decides to use the water-displacement method to find the volume. The initial volume reading is 25.5 mL and, after the Al sample is added, the water level has risen to 41.7 mL. Find the density of the Al sample in g/cm3. (Remember: 1 mL = 1 cm3.) 7. If you are sure that a sample of material is aluminum but have no measuring instruments AND are not allowed to handle the sample, how would you determine the sample’s density? 8. A gas has a mass of 7914 g and takes up enough space to fill a room that is 2.00 m X 2.00 m X 2.50 m. Determine the density of the gas in g/m3. 9. A flask that weighs 345.8 g is filled with 225 mL of carbon tetrachloride. The weight of the flask and carbon tetrachloride is found to be 703.55 g. From this information, calculate the density of carbon tetrachloride. 8 Unit 1- Matter and Change The Study of Matter Practice Test Directions: Define and/or describe the following terms relating to the scientific method. 1. Dependent Variable _________________________________________________ 2. Hypothesis __________________________________________________________ Directions: Indicate if the process listed is a physical or chemical change. 3. Food digests _________________________________________________________ 4. Bending a piece of copper wire. _______________________________________ 5. Two clear liquids react to form a yellow clumps ___________________________ Directions: Solve the following problems. Show all work! Be sure to include the correct unit with your final answer. 6. What is the density of a substance that has a volume of 2.8cm3 and mass of 25grams? 7. What is the density of a solid that has a volume of 4 cm3 and a mass of 6 grams? Directions: For each sample of matter below, correctly classify it as a pure substance or a mixture. 8. Trail Mix ____________________ 9. Helium ____________________ Directions: For each, correctly classify as homogeneous or heterogeneous mixture. 10. Vegetable soup . ______________________ 11. Gatorade . ______________________ 12. Orange juice, with pulp . ______________________ Multiple Choice Practice 13. The amount of mass per unit volume refers to the a. Density b. Specific weight c. Volume d. Weight 14. A substance that can be further simplified may be either a. An element or a compound b. An element or a mixture c. A mixture or a compound d. A mixture or an atom 9 Unit 1- Matter and Change 15. A substance composed of two or more elements chemically united is called a. An isotope c. An element b. A compound d. A mixture 16. An example of a chemical change is the a. Breaking of a glass bottle b. Sawing of a piece of wood c. Rusting of iron d. Melting of an ice cube 17. A substance that cannot be further decomposed by ordinary chemical means is a. Water c. Sugar b. Air d. Silver 18. An example of a physical change is a. The fermenting of sugar to alcohol b. The rusting of iron c. The burning of paper d. A solution of sugar in water 19. The property of matter that is independent of its surroundings and position is a. Volume c. Mass e. State b. Density d. Weight 20. What Kelvin temperature is equal to 25°C? a. 248 K b. 298 K c. 100 K d. 200 K 21. Which substance cannot be decomposed into simpler substances? a. ammonia c. methane b. aluminum d. methanol 22. A compound differs from a mixture in that a compound always has a a. homogeneous composition b. maximum of two components c. minimum of three components d. heterogeneous composition 23. Which statement describes a chemical property? a. Its crystals are a metallic gray. b. It dissolves in alcohol. c. It forms a violet-colored gas. d. It reacts with hydrogen to form a gas. 10 Unit 1- Matter and Change 24. To determine the density of an irregularly shaped object, a student immersed the object in 21.2 milliliters of H2O in a graduated cylinder, causing the level of the H2O to rise to 27.8 milliliters. If the object had a mass of 22.4 grams, what was the density of the object. a. 27.8 g / mL c. 3.0 g / mL b. 6.6 g / mL d. 3.4 g/ mL 25. Sarah designed an experiment to find out which mouthwash was most effective against some bacteria. She cut out four different circles from a paper towel and soaked each circle in a different mouthwash. She put the circles on a nutrient agar-coated Petri dish that was covered with bacteria commonly found in the mouth. She then incubated the plate for 24 hours. The picture shows the results of this test. Which of the following should Sarah do to improve her experiment? a. Use different kinds of bacteria b. Use the same type of mouthwash on each paper circle. c. Use the same size paper circles for all mouthwashes. d. Use a smaller Petri dish. 26. A student decided to set up an experiment to see if cats prefered skim milk or 2%milk. She put out a cup of milk for 5 kittens and then measured how much the kittens drank over the course of a day. The same kittens were used and the milk was served at the same temperature. The student discovered that the cats liked the 2% more than the skim milk. What is wrong with the above experiment? a. Repeated experimentation is needed. b. The milk should be served at different temperatures. c. There are no constants present. d. There is no independent variable. 27. An experiment for a new asthma medication was set up into two groups. Group one was given the new drug for asthma, while group 2 was given a sugar pill. The sugar pill serves as a. Control c. experimental variable b. Constant d. dependent variable 28. A scientist plants two rows of corn for experimentation. She puts fertilizer on row 1 but does not put fertilizer on row 2. Both rows receive the same amount of water and light intensity. She checks the growth of the corn over the course of 5 months. What is a constant in this experiment? a. Plant height d. Amount of water b. Corn without fertilizer c. Corn with fertilizer 11 Unit 1- Matter and Change 29. Which sentence best states the importance of using control groups? a. Control groups eliminate the need for large sample sizes, reducing the number of measurements needed. b. Control groups eliminate the need for statistical tests and simplify calculations. c. Control groups provide a method by which statistical variability can be reduced. d. Control groups allow comparison between subjects receiving a treatment and those receiving no treatment 30. The measurable factor in an experiment is known as the: a. Control b. independent variable c. constant d. dependent variable 31. A student decides to set up an experiment to see if detergent affects the growth of seeds. He sets up 10 seed pots. 5 of the seed pots will receive a small amount of detergent in the soil and will be placed in the sun. The other 5 seed pots will not receive detergent and will be placed in the shade. All 10 seed pots will receive the same amount of water, the same number of seeds, and the same type of seeds. He grows the seeds for two months and charts the growth every 2 days. What is wrong with his experiment? a. More than one variable is being tested. b. The student should have a larger number of pots. c. There is no way of measuring the outcome. d. There is no control set-up. 32. A scientific study showed that the depth at which algae were found in a lake varied from day to day. On clear days, the algae were found as much as 6 meters below the surface of the water but were only 1 meter below the surface on cloudy days. Which hypothesis best explains these observations? a. Nitrogen concentration affects the growth of algae. b. Precipitation affects the growth of algae. c. Light intensity affects the growth of algae. d. Wind currents affect the growth of algae. 12 Unit 2- Atomic Theory and Structure ATOMIC THEORY HISTORY OF THE ATOM The original idea (400 B.C.) came from ______________________, a Greek philosopher. He expressed the belief that all matter is composed of very small, indivisible particles, which he named atomos. John Dalton (1766-1844), an English school teacher and chemist, proposed his atomic theory of matter in 1803. Dalton’s Atomic Theory states that: 1. All matter is made of tiny __________________________ particles called atoms. 2. Atoms of the ____________ element are identical; those of different atoms are different. 3. Atoms of different elements combine in whole number ________________ to form compounds 4. Chemical reactions involve the rearrangement of atoms. No _______ atoms are created or destroyed. PARTS OF THE ATOM Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts. In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate. Thomson’s experiments used a __________________ ray tube. It is a vacuum tube - all the air has been pumped out. Because these rays originate at the ____________________, they are called cathode rays. Thomson concluded that cathode rays are made up of invisible, _________________________ charged particles referred to as electrons. From Thomson’s experiments, scientists had to conclude that atoms were not just neutral _________________, but somehow were composed of electrically charged particles. Matter is not negatively charged, so atoms can’t be negatively charged either. If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles — probably with a much greater _____________ than electrons. J.J. Thomson said the atom was like ______________ pudding, a popular English dessert. In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important experiments that revealed an arrangement far different from the plum pudding model of the atom. The experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole. 13 Unit 2- Atomic Theory and Structure The sheet of ________________ foil was surrounded by a screen coated with zinc sulfide, which glows when struck by the positively charged particles of the beam. The ________________ particles were expected to pass through without changing direction very much because Rutherford thought the mass was evenly distributed in the atom. Because most of the particles passed through the foil, they concluded that the atom is nearly all _______________ ______________. Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a ____________________. Alpha particles are deflected by it if they get close enough to the nucleus. R.A. Millikan found the charge of an electron to be -1.60 x 10-19 Coulombs in his famous oil drop experiment. In 1910, J.J. Thomson discovered that neon consisted of atoms of two different masses. Atoms of an element that are chemically alike but differ in mass are called ______________________ of the element. Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass. Calculations showed that such a particle should have a mass ____________________ to that of a proton but no electrical _______________. The existence of this neutral particle, called a neutron, was confirmed in the early 1930s. James _________________ is given credit for discovering the neutron. 14 Unit 2- Atomic Theory and Structure NAME SYMBOL CHARGE RELATIVE MASS 1/2000 proton no MODERN VIEW OF THE ATOM The atom has two regions and is ___-dimensional. The nucleus is at the ___________________ and contains the protons and _____________________. The electron cloud is the region where you might find an electron and most of the volume of an atom. The atomic _________________ of an element is the number of protons in the nucleus of an atom of that element. The number of protons determines ____________________ of an element, as well as many of its chemical and physical properties. Because atoms have no overall electrical charge, an atom must have as many ____________________ as there are protons in its nucleus. Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element. The mass of a neutron is almost the same as the mass of a ________________. The sum of the protons and neutrons in the nucleus is the ________________ number of that particular atom. _____________________ of an element have different mass numbers because they have different numbers of _______________, but they all have the same atomic number. AVERAGE ATOMIC MASS The atomic mass is the weighted average mass of all the naturally occurring isotopes of that element. To determine the average atomic mass, first calculate the contribution of each isotope to the average atomic mass, being sure to convert each ___ a fractional abundance. The average atomic mass of the element is the sum of the mass contributions of each isotope. Elements can be represented by using the symbol of the element, the mass number and the atomic number. The mass number is the __________________ mass rounded to a whole number. 15 Unit 2- Atomic Theory and Structure 1. Determine the following for the fluorine-19 atom. a) number of protons b) number of neutrons d) atomic number e) mass number c) number of electrons 2. Repeat #1 for bromine-80. 3. If an element has an atomic number of 34 and a mass number of 78, what is the a) number of protons b) number of neutrons c) number of electrons d) complete symbol 4. If an element has 91 protons and 140 neutrons, what is the a) atomic number b) mass number c) number of electrons d) complete symbol 5. If an element has 78 electrons and 117 neutrons what is the a) atomic number b) mass number c) number of protons d) complete symbol Practice/Homework Isotopes and Subatomic Particles Complete the chart Isotope symbol 207 82 Atomic # Mass Protons Neutrons Electrons 8 9 77 54 Pb 38 12 50 13 106 11 7 157 12 14 11 126 85 238 92 U 75 33 As 32 16 65 29 S Cu 202 80 Hg 16 Unit 2- Atomic Theory and Structure Isotope symbol 261 104 Atomic # Mass Protons Neutrons 47 40 24 61 51 28 Electrons Rf MOLES We measure ________________ in grams. We measure volume in __________________. We count pieces in _________________. The number of moles is defined as the number of __________________ atoms in exactly ____ grams of carbon-12. ____ mole is 6.022 x 1023 particles. 6.022 x 1023 is called __________________ number. Representative particles are the smallest pieces of a substance. For a molecular compound it is a(n) ______________________. For an ionic compound it is a ______________________ ______________. For an element it is a(n) ________________. How many oxygen atoms are in the following? a) CaCO3 b) Al2(SO4)3 How many total ions are in the following? a) CaCl2 b) NaF c) Al2S3 MOLE CONVERSIONS 1. How many atoms of carbon are there in 1.23 moles of carbon? 2. How many molecules of CO2 are in 4.56 moles of CO2? 3. How many atoms of iron are in 0.600 moles of iron? 4. How many moles are in 7.78 x 1024 formula units of MgCl2? 5. How many moles of water are 5.87 x 1022 molecules of water? 6. How many moles of aluminum are 1.2 x 1024 atoms of aluminum? 17 Unit 2- Atomic Theory and Structure Representative particles Volume Use 22.4 L Use 6.02x1023 moles Remember DIMO: Divide In Multiply Out Use molar mass mass Calculate the number of particles (atoms, ions or molecules) in each of the following. a) 3.4 moles Na2S b) 0.0020 moles Zn c) 1.77 x 10-11 moles C d) 92.35 moles O2 Calculate the number of moles in each of the following. a) 3.4 x 1024 molecules HCl b) 8.7 x 1021 atoms Zn c) 1.77 x 1018 ions Al+3 d) 2.66 x 1026 atoms Cu MOLAR MASS Molar mass is the generic term for the mass of one _____________. It may also be referred to as gram molecular mass, gram formula mass, and gram atomic mass. The unit is ______________. To determine the molar mass of an element, find the element’s symbol on the periodic table and round the mass so there is __________ digit beyond the decimal. 18 Unit 2- Atomic Theory and Structure Determine the molar mass of the each of the following elements. a) sulfur (S) b) chromium (Cr) c) bromine (Br) To determine the molar mass of a compound, find the mass of all elements in the compound. If necessary, ___________________ an element’s mass by the subscript appearing beside that element in the compound’s formula (or ________________ of the subscripts). Calculate the molar mass of each of the following compounds. a) Na2S b) N2O4 c) C6H12O6 d) Ca(NO3)2 MASS-PARTICLE/MOLE CONVERSIONS 1. How many atoms of lithium are in 1.00 g of Li? 2. How many molecules of sodium oxide are in 42.0 g of Na2O? 3. How much would 3.45 x 1022 atoms of uranium (U) weigh? 4. How many moles of magnesium are in 56.3 g of Mg? 5. How many moles is 5.69 g of NaOH? 6. How many grams of sodium chloride are in 3.45 moles of NaCl? 7. How many moles is 4.8 g of CO2? 8. How many grams is 9.87 moles of H2O? 9. How many molecules are in 6.8 g of CH4? 10. What is the mass of 49.0 molecules of C6H12O6? GASES Many of the chemicals we deal with are gases. They are difficult to weigh, and we need to know how many moles of gas we have. Two things affect the volume of a gas: temperature and pressure. Standard temperature is ______ ºC, and standard pressure is ______ atm. Standard temperature and pressure is abbreviated STP. At STP 1 mole of gas occupies ______ L. 22.4 L is called the _____________ volume. Avogadro’s Hypothesis - At the same temperature and pressure equal volumes of gas have the same number of _______________________. 19 Unit 2- Atomic Theory and Structure GAS CONVERSIONS 1. What is the volume of 4.59 mole of CO2 gas at STP? 2. How many moles is 5.67 L of O2 at STP? 3. What is the volume of 8.8 g of CH4 gas at STP? 4. How many grams is 16.2 L of O2 at STP? Calculate the number of liters in each of the following. a) 3.10 x 1024 molecules Cl2 b) 8.7 moles Ne c) 2.77 x 1018 atoms He d) 266 grams SO2 Homework/Practice Part 1--Convert between particles and moles 1. 24 atoms of sodium = _____ moles of sodium atoms 2. 5 molecules of chlorine gas = _____ moles of chlorine molecules 3. 900 atoms of silver = _____ moles of silver atoms 4. 2.89 x 1023 molecules of ammonia = _____ moles of ammonia molecules 5. 15 moles of arsenic atoms = ______ atoms of arsenic 6. 4.00 x 103 moles of barium atoms = __________ atoms of barium Part 2--Convert between mass and moles 7. Calculate the mass of 1.000 mole of CaCl2 8. Calculate grams in 3.0000 moles of CO2 9. Calculate number of moles in 32.0 g of CH4 10. Calculate moles in 168.0 g of HgS 11. Calculate moles in 510.0 g of Al2S3 12. How many moles are in 27.00 g of H2O 13. What is the mass of 2.55 moles Cu2CrO4 20 Unit 2- Atomic Theory and Structure Part 3- Multiple steps 14. Arrange the following in order of increasing weight. a. 10.4 g of sulfur c. 6.33 x 1025 atoms of hydrogen b. 0.179 moles of iron d. 0.77 moles of N2 15. How many grams would 8.1 1021 molecules of sucrose (C12H22O11) weigh? 16. How many atoms are in a 2.0 kg ingot of gold? (Note mass units.) 17. What is the mass of 2.3x1024 molecules of KCl? 18. Calculate the number of molecules in 50.0 grams of H2SO4 19. Calculate the number of molecules in 100. grams of KClO4 20. Calculate the number of molecules in 8.76 grams of NaOH 21. Calculate the mass of 1.2x1022 molecules of Fe3(PO4)2 22. Calculate mass of 7.2x1024 molecules of Na2CO3 NUCLEAR CHEMISTRY Nuclear chemistry is the study of the structure of _________________ nuclei and the changes they undergo. Marie Curie named the process by which materials such as uranium give off rays radioactivity; the rays and particles emitted by a radioactive source are called __________________. As you may recall, isotopes are atoms of the same element that have different numbers of _________________. Isotopes of atoms with unstable nuclei are called ______________________. These unstable nuclei emit radiation to attain more stable atomic configurations in a process called radioactive ________________. During radioactive decay, unstable atoms lose _________________ by emitting one of several types of radiation. TYPES OF RADIATION The three most common types of radiation are alpha (α), ____________ (β), and gamma (γ). An alpha particle (α) has the same composition as a __________________ nucleus - two protons and ________ neutrons - and is therefore given the symbol _________. The charge of an alpha particle is 2+ due to the presence of the two ___________________. Because of their mass and charge, alpha particles are relatively slow-moving compared with other types of radiation. Thus, alpha particles are not very ________________________ - a single sheet of paper stops alpha 21 Unit 2- Atomic Theory and Structure particles. A beta particle is a very-fast moving ______________________ that has been emitted from a neutron of an unstable nucleus. Beta particles are represented by the symbol _________. The zero superscript indicates the insignificant mass of an electron in comparison with the mass of a ____________________. The –1 subscript denotes the _____________________ charge of the particle. Beta radiation consists of a stream of fast-moving electrons. Because beta particles are both lightweight and fast moving, they have _____________________ penetrating power than alpha particles. A thin metal foil is required to stop beta particles. Gamma rays are highenergy (_________________ wavelength) electromagnetic radiation. They are denoted by the symbol __________. As you can see from the symbol, both the subscript and superscript are zero. Thus, the emission of gamma rays does not change the __________________ number or mass number of a nucleus. Gamma rays almost always accompany alpha and beta radiation, as they account for most of the energy loss that occurs as a nucleus decays. NAME SYMBOL Alpha FORMULA 4 2 MASS CHARGE DESCRIPTION He β -1 0 High energy radiation NUCLEAR STABILITY and DECAY Radioactive nuclei undergo decay in order to gain _____________________. All elements with atomic numbers greater than 83 are radioactive. Nuclear equations are used to show nuclear transformations. Balanced nuclear equations require that both the ____________________ number and the mass number must be balanced. When beryllium-9 is bombarded with alpha particles (helium nuclei), a neutron is produced. The balanced nuclear reaction is given as: ________________________________________________ 22 Unit 2- Atomic Theory and Structure The atomic number (the number on the bottom) determines the identity of the element. When nitrogen-14 is bombarded with a neutron, a proton is produced. The balanced nuclear equation can be written as: _________________________________________________________ Polonium-230 undergoes alpha decay: ________________________________________________ Uranium-234 undergoes alpha decay: ________________________________________________ Cobalt-50 undergoes beta decay: ____________________________________________________ Provide symbols for each of the following: neutron ___________, proton ___________ or ___________, and the positron ___________. What element is formed when iron-60 undergoes beta decay? Give the atomic number and mass number of the element. ____________ Write a balanced nuclear equation for the alpha decay of the following radioisotope, uranium-235. ____________________________________________________________ Nitrogen-12 decays into a positron and another element. Write the balanced nuclear equation. ____________________________________________________________ Uranium-238 is bombarded with a neutron. One product forms along with gamma radiation. Write the balanced nuclear equation. ____________________________________________________________ Nitrogen-14 is bombarded with deuterium (hydrogen-2). One product forms along with an alpha particle. Write the balanced nuclear equation. ____________________________________________________________ 23 Unit 2- Atomic Theory and Structure RADIOACTIVE DECAY RATES Radioactive decay rates are measured in half-lives. A half-life is the time required for one-half of a radioisotope’s nuclei to ________________ into its products. For example, the half-life of the radioisotope strontium-90 is 29 years. If you had 10.0 g of strontium-90 today, 29 years from now you would have 5.0 g left. The decay continues until negligible strontium-90 remains. Iron-59 is used in medicine to diagnose blood circulation disorders. The half-life of iron-59 is 44.5 days. How much of a 2.000-mg sample will remain after 133.5 days? Cobalt-60 has a half-life of 5.27 years. How much of a 10.0 g sample will remain after 21.08 years? Carbon-14 has a half-life of 5730 years. How much of a 250. g sample will remain after 5730 years? FISSION and FUSION Heavy atoms (mass number > 60) tend to break into smaller atoms, thereby increasing their ________________________. Using a neutron to split a nucleus into fragments is called nuclear _______________________. Nuclear fission releases a large amount of energy and several neutrons. Since neutrons are products, one fission reaction can lead to more fission reactions, a process called a ________________ reaction. A chain reaction can occur only if the starting material has enough mass to sustain a chain reaction; this amount is called __________________ mass. The _____________________ of atomic nuclei is called nuclear fusion. For example, nuclear fusion occurs within the Sun, where hydrogen atoms fuse to form __________________ atoms. Fusion reactions can release very large amounts of energy but require extremely high temperatures. For this reason, they are also called _____________________________ reactions. APPLICATIONS OF NUCLEAR REACTIONS Geiger counters, scintillation counters, and film badges are devices used to detect and measure radiation. Geiger counters use ________________________ radiation, which produces an electric current in the counter, to rate the strength of the radiation on a scale. Film badges are often used to monitor the approximate radiation ______________________ of people working with radioactive materials. Scintillation counters measure ionizing radiation. With proper safety 24 Unit 2- Atomic Theory and Structure procedures, radiation can be useful in industry, in scientific experiments, and in medical procedures. Nuclear power plants use the process of nuclear fission to produce heat in nuclear reactors. The heat is used to generate steam, which is then used to drive __________________ that produce electricity. A radiotracer is a radioisotope that emits non-ionizing radiation and is used to signal the presence of an element or of a specific substance. Radiotracers are used to detect ______________________ and to analyze complex chemical reactions. Ionizing radiation has many uses. A(n) _______________ is ionizing radiation, and ionizing radiation can be used in medicine to kill cancerous cells. Most medical devices require sterilization after they are packaged, and another trend has been the move to sterilization by __________________ radiation as opposed to other methods such as ethylene oxide gas. Advantages of gamma irradiation include ______________, cost-effectiveness, and the elimination of the need for special packaging. Chemical reaction rates are greatly affected by changes in temperature, _____________________, and concentration, and by the presence of a catalyst. In contrast, nuclear reaction rates remain ____________________ regardless of such changes. In fact, the half-life of any particular radioisotope is constant. Because of this, radioisotopes, especially carbon-14, can be used to determine the ____________ of an object. The process of determining the age of an object by measuring the amount of a certain radioisotope remaining in that object is called radiochemical dating. EFFECTS OF NUCLEAR REACTIONS Any exposure to radiation can damage living ____________. Gamma rays are very dangerous because they penetrate ______________________ and produce unstable and reactive molecules, which can then disrupt the normal functioning of cells. The amount of radiation the body absorbs (a dose) is measured in units called rads and ____________. Everyone is exposed to radiation, on average 100–300 millirems per year. A dose exceeding ____________ rem can be fatal. 25 Unit 2- Atomic Theory and Structure Atomic Theory, The Mole, and Nuclear Chemistry- Practice Test 1. Given the work of Dalton, please check the box for the postulate(s) that have since been proven to be incorrect. Explain what we now know to be the true case. [] All atoms of a specific element are identical. [] Compounds consist of atoms of different elements combined together. [] Atoms of different elements have different masses. Directions: For the scientist listed below, explain what was done in the experiment, what knowledge was developed as a result, 2. Rutherford 3. Thomson Directions: Fill in the table for the following isotopes. 4. 5. 6. 7. Isotope H-1 Cu-65 Atomic # Mass # Protons 18 40 19 9 Neutrons Electrons 8. What is the charge of a beta particle? Directions: Solve the following problems be sure to include the correct unit with your final answer. 9. Given the equation: X 24 He + 220 84 Po The nucleus represented by X is 10. How many moles of sodium are 6.02 x 10 23 atoms of sodium? 11. What is the mass of 6 moles of Carbon? 12. How many atoms are in 45g of Neon? Multiple Choice Practice 13. What is the approximate formula mass of Ca(NO3)2 a. 70 c. 102 b. 82 d. 150 e. 164 14. How many molecules are in 1 mole of water? a. 3 c. 6.02x1023 b. 54 d. 2(6.02x1023) e. 3(6.02x1023) 26 Unit 2- Atomic Theory and Structure 15. How many atoms are represented in the formula Ca3(PO4)2 a. 5 c. 9 b. 8 d. 12 16. Four grams of hydrogen gas at STP contain a. 6.02x1023 atoms b. 12.04x1023atoms c. 12.04x1046atoms e. 13 d. 1.2x1022molecules e. 12.04x1023molecules 17. What is the mass in grams of 1 mole of KAl(SO4)2*12H2O a. 132 c. 394 b. 180 d. 474 e. 516 18. Compared to the charge and mass of a proton, an electron has a. the same charge and a smaller mass b. the same charge and the same mass c. an opposite charge and a smaller mass d. an opposite charge and the same mass 19. When alpha particles are used to bombard gold foil, most of the alpha particles pass through undeflected. This result indicates that most of the volume of a gold atom consists of ____. a. deuterons c. protons b. neutrons d. unoccupied space 20. A proton has approximately the same mass as a. a neutron b. an alpha particle c. a beta particle d. an electron 21. A neutron has approximately the same mass as a a. an alpha particle b. a beta particle c. an electron d. a proton 22. Which symbols represent atoms that are isotopes? a. C-14 and N-14 b. O-16 and O-18 c. I-131 and I-131 d. Rn-222 and Ra-222 23. Which atom contains exactly 15 protons? a. P-32 b. S-32 c. O-15 d. N-15 24. An ion with 5 protons, 6 neutrons, and a charge of 3+ has an atomic number of a. 5 b. 6 c. 8 27 d. 11 Unit 2- Atomic Theory and Structure 25. What is the mass number of an atom which contains 28 protons, 28 electrons, and 34 neutrons? a. 56 b. 62 c. 90 d. 28 26. What is the total number of atoms represented in the formula CuSO4 . 5H2O? a. 8 b. 13 c. 21 27. What is the gram formula mass of K2CO3? a. 138 g b. 106 g d. 27 c. 99 g d. 67 g 28. What is the total number of atoms contained in 2.00 moles of nickel? a. 58.9 c. 6.02 x 1023 b. 118 d. 1.2 x 1024 29. What is the total number of moles of hydrogen gas contained in 9.03 x 1023molecules a. 1.5 moles c. 6.02 moles b. 2.00 moles d. 9.03 moles 30. What is the mass in grams of 3.0 x 1023 molecules of CO2? a. 22 g c. 66 g b. 44 g d. 88 g 31. The amount of substance having 6.022 x 1023 of any kind of chemical unit is called a(n): a. formula c. mole b. mass number d. atomic weight 32. The total number of atoms in the formula of aluminum dichromate, Al2(Cr2O7)3 is: a. 5 b. 29 c. 17 d. 11 33. The total number of OXYGEN atoms in the formula of aluminum dichromate, Al2(Cr2O7)3 is: a. 21 b. 10 c. 29 34. The formula mass of calcium hydroxide, Ca(OH)2 is: a. 57.05 grams b. 74.10 grams 28 c. 128 grams d. 97.07 grams d. 7 Unit 2- Atomic Theory and Structure 35. The formula mass of ammonium chlorate, NH4ClO3 is: a. 101.5 g c. 211.43 g b. 78.06 g d. 172.40 g 36. What is the molar mass of the gas butane, C4H10? a. 13.02 grams b. 485.2 grams c. 68 24 grams d. 58.14 grams 37. The molar mass of sodium chloride, NaCl is: a. 69.71 grams b. 2 grams c. 6.022 x 1023 grams d. 58.44 grams 38. The formula mass of magnesium hydroxide, Mg(OH)2 is: a. 42.33 grams c. 41.32 grams b. 58.33 grams d. 5 grams 39. What is the mass in grams of 3 moles of water molecules, H2O? a. 54.06 grams c. 0.166 grams b. 21.02 grams d. 6.01 grams 40. What is the mass of 4 moles of hydrogen molecules (H2)? a. 4.04 grams c. 3.96 grams b. 8.08 grams d. 1.98 grams 41. What is the mass in grams of 10 moles of ammonia, NH3? a. 170.4 grams c. 1.704 grams b. 0.587 grams d. 27.04 grams 42. How many moles of water molecules, H2O, are present in a 42 gram sample of water? a. 23.98 moles c. 2.33 moles b. 0.429 moles d. 757 moles 43. How many moles of methane molecules, CH4, are in 80 grams of methane? a. 0.201 moles c. 6.022 x 1080 moles b. 4.98 moles d. 1284 moles 29 Unit 2- Atomic Theory and Structure 44. How many moles of calcium hydroxide, Ca(OH)2 are in 150 grams of the compound? a. 2.02 moles c. 0.494 moles b. 224.1 moles d. 11115 moles 45. About how many atoms of helium would be found in 2 grams of helium? a. 4.00260 c. 2 b. 6.02 x 1023 d. 3.01 x 1023 46. By knowing the number of electrons in a neutral atom, you should also be able to determine a. the number of neutrons in the neutral atom b. the number of protons in the neutral atom c. the atomic number of the neutral atom d. the mass of the neutral atom e. two of these 47. How many oxygen atoms are there in one formula unit of Al2(SO4)3? a. 3 b. 4 c. 7 d. 12 48. The average mass of a magnesium atom is 24.31. If you were able to select and measure a single atom of magnesium, the chance that you would select an atom of mass 24.31 is about: a. 0% d. greater than 50% b. 0.31% e. 100% c. 24.31% 49. Which of the following arrangements represent different isotopes of the same element? i. ii. iii. iv. v. 12 protons, 11 neutrons, 12 electrons 11 protons, 12 neutrons, 11 electrons 10 protons, 12 neutrons, 12 electrons 11 protons, 12 neutrons, 10 electrons 12 protons, 12 neutrons, 12 electrons a. 1 and 5 b. 2 and 4 c. 2, 3, 4 and 5 d. all of these qualify e. None of these qualify 30 Unit 2- Atomic Theory and Structure 50. If the abundance of 6Li (6.015121 amu) is 7.500% and the abundance of 7Li (7.016003 amu) is 92.500%, what is the average atomic mass? a. 6.0750 amu c. 6.9250 amu b. 6.0902 amu d. 6.9409 amu 51. Which of following is not true of the carbon-14 atom? a. It has six protons b. It has an average mass of 12.011 amu. c. It has six electrons d. It has eight neutrons e. It is the less common than carbon-12 52. A particle with 15 protons and 18 electrons would be symbolized as: a. Ar c. P3b. Ar3d. P3+ 53. Which of these is the correct number of particles in this nuclide? a. 34 protons, 79 neutrons, 2 electrons b. 34 protons, 45 neutrons, 32 electrons c. 34 protons, 45 neutrons, 2 electrons d. 34 protons, 45 neutrons, 36 electrons e. 34 protons, 113 neutrons, 36 electrons 79 34 e. Ar3+ Se 2 54. Which of the following have equal numbers of neutrons? a. I, II and III b. II and III c. I and V d. I and IV e. II, III and IV 55. The element hafnium (Hf) has five stable isotopes. The correct number of nuclear particles in an atom of hafnium-178 is: a. 72 protons, 178 neutrons d. 72 protons, 106 neutrons b. 72 protons, 72 electrons e. 72 protons, 106 neutrons, c. 106 protons, 72 neutrons 72 electrons 31 Unit 2- Atomic Theory and Structure 56. J.J. Thomson's model of the atom can be summarized with the visual image of: a. planets orbiting the sun d. a small central nucleus and b. plum pudding an electron cloud c. bees around a hive e. none of the above 57. The number of errors contained in the table below is: 58. 10 5 Nuclide protons neutrons electrons mass Sodium-23 11 12 11 23 Cobalt-59 27 22 27 59 Tungsten-184 74 110 184 184 Fluorine-19 a. one b. two 10 9 c. three d. four 10 19 e. five B + _____ 13 H + 2 24 He In the equation shown here, the missing particle is: a. b. 1 1 1 0 H n c. d. o 1 1 1 e p e. 4 2 37 K → _____ + 59. Identify the missing particle in the following nuclear reaction: 19 37 38 36 Ar Ar Ar a. 18 b. 18 c. 18 He 0 1 e d. 37 20 60. For the most common types of radioactive decay, the order of least penetrating to human tissue, to most penetrating to human tissue is: a. gamma, beta, alpha c. beta, gamma, alpha b. alpha, beta, gamma d. gamma, alpha, beta 61. Very large nuclei tend to be unstable because of the: a. repulsive forces between protons b. attraction of protons for neutrons c. repulsive forces between neutrons d. attraction of electrons for the positively charged nucleus e. repulsive forces between electrons 32 Ca Unit 2- Atomic Theory and Structure 62. An alpha () particle is essentially a ____________________ nucleus. a. plutonium c. hydrogen b. helium d. uranium e. carbon-12 63. Phosphorus-15 has a half-life of 14 days. What proportion of the original phosphorus-15 remains after 8 weeks? a. 1/2 c. 1/4 e. 1/8 b. 1/16 d. 1/32 64. The nuclide radium-226 is the daughter nuclide resulting from the decay of what parent nuclide? a. radon-222 d. thorium-228 b. polonium-214 e. radium-225 c. thorium-230 65. An electron emitted from the nucleus during some kinds of radioactive decay is known as: a. A gamma ray c. A beta () particle b. A positron d. An alpha () particle 66. A process in which a very heavy nucleus splits into more-stable nuclei of intermediate mass is called: a. nuclear fission d. nuclear fusion b. radiocarbon dating e. radioactive decay c. a chain reaction 33 Unit 3- Electrons and Periodicity ELECTRONS IN ATOMS LIGHT Light is a kind of electromagnetic _____________________. All forms of electromagnetic radiation move at 3.00 x 108 m/s. The ______________ is the baseline of a wave. The crest is the high point on a wave, and the trough is low point on a wave. The amplitude of a wave is the wave’s _____________ from the origin to a crest, or from the origin to a trough. Wavelength (represented by λ, the Greek letter lambda) is the ___________________ distance between equivalent points on a continuous wave. Wavelength is the distance from crest to crest or trough to trough and is usually expressed in meters (m). _____________________ (represented by f ) is the number of “waves” that pass a given point per second, and the units are cycles/sec or hertz (Hz) c = fλ c = the speed of light Frequency and wavelength are __________________ related, which means that as one goes up the other goes _____________. Different frequencies of light correspond to different colors of light. In 1900, the German physicist Max Planck began searching for an explanation as he studied the light emitted from ___________________ objects. Matter can gain or lose energy only in small, specific amounts called _______________. That is, a quantum is the minimum amount of energy that can be gained or lost by a(n) ____________. That is, while a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called ________________. Thus, a photon is a particle of electromagnetic radiation with no _____________ that carries a quantum of energy. Planck went further and demonstrated mathematically that the energy of a quantum is ___________________ related to the frequency of the emitted radiation. E=hf E = energy of the photon (J – Joules); f = frequency (Hz); h = Planck’s constant (J.s) The energy of radiation increases as the radiation’s frequency, f, __________________. Scientists knew that the wave model of light could not explain a phenomenon called the ____________________ effect. In the photoelectric effect, electrons, called __________________________, are emitted from a metal’s surface when light of a certain _______________________ shines on the surface. Einstein proposed that for the photoelectric 34 Unit 3- Electrons and Periodicity effect to occur, a photon must possess, at a minimum, the energy required to _______________ an electron from an atom of the metal. THE BOHR MODEL OF THE ATOM Niels Bohr, a young Danish physicist working in Rutherford’s laboratory in 1913, suggested that the single electron in a ___________________ atom moves around the nucleus in only certain allowed circular orbits. The atom looked like a miniature _________________ system. The nucleus is represented by the sun, and the electrons act like the planets. The orbits are circular and are at different levels. Amounts of ___________________ separate one level from another. (Modern View: The atom has two regions and is 3-dimensional. The nucleus is at the _________________ and contains the protons and neutrons. The electron _________________ is the region where you might find an electron and most of the volume of an atom.) Bohr proposed that electrons must have enough energy to keep them in constant motion around the ___________________. Electrons have energy of motion that enables them to overcome the attraction of the _________________ nucleus. Further away from the nucleus means more energy. Electrons reside in ________________ levels. THE QUANTUM MECHANICAL MODEL Building on Planck’s and Einstein’s concepts of ____________________ energy (quantized means that only certain values are allowed), Bohr proposed that the hydrogen atom has only certain allowable energy ______________. The lowest allowable energy state of an atom is called its _______________ state. When an atom gains energy, it is said to be in a(n) __________________ state. When the atom is in an excited state, the electron can drop from the higher-energy orbit to a _______________-energy orbit. As a result of this transition, the atom emits a ____________________ corresponding to the difference between the energy levels associated with the two orbits. ATOMIC AND EMISSION SPECTRA By heating a gas of a given element with electricity, we can get it to give off _______________. Each element gives off its own characteristic colors. The spectrum can be used to __________________ the atom. These are called line _______________. Each is unique to an 35 Unit 3- Electrons and Periodicity element. The spectrum of light released from excited atoms of an element is called the _________________ spectrum of that element. As the electrons fall from the excited state, they __________________ energy in the form of light. The further they fall, the ________________ the energy. This results in a higher frequency. Use the Chemistry Reference Tables to answer the following: (a) An electron falls from energy level 5 to 3. What is the wavelength of the light emitted? (b) An electron falls from energy level 6 to 2. What is the wavelength of the light emitted? (c) An electron falls from energy level 3 to 1. What type of electromagnetic radiation is emitted (infrared, visible or ultraviolet)? (d) An electron falls from energy level 4 to 2. What type of electromagnetic radiation is emitted (infrared, visible or ultraviolet)? (e) An electron falls from energy level 5 to 2. What color of visible light is emitted? (f) An electron falls from energy level 3 to 2. What color of visible light is emitted? MORE QUANTUM MECHANICAL MODEL Like Bohr’s model, the quantum mechanical model limits an electron’s energy to certain values. The space around the nucleus of an atom where the atom’s electrons are found is called the electron ________________. A three-dimensional region around the nucleus called an atomic __________________ describes the electron’s probable location. In general, electrons reside in principal ________________ levels. As the energy level number increases, the orbital becomes _______________, the electron spends more time ___________________ from the nucleus, and the atom’s energy level increases. Principal energy levels contain energy ___________________. Principal energy level 1 consists of a single sublevel, principal energy level 2 consists of __________ sublevels, principal energy level 3 consists of three sublevels, and so on. Sublevels are labeled s, p, d, or f. The s sublevel can hold 2 electrons, the p sublevel can hold _____ electrons, the d sublevel can hold 10 electrons, and the f sublevel can hold 14 electrons. Sublevels contain __________________. Each orbital may contain at most ________ electrons. There is one s orbital for every energy level, and the s orbital is ____________________ shaped. They are called the 1s, 2s, 3s, etc… orbitals. The p orbitals start at the second energy level, reside along ______ different directions and have 3 different 36 Unit 3- Electrons and Periodicity ________________ shapes. The d orbitals start at the ________________ energy level and have ____ different shapes. The f orbitals start at the fourth energy level and have ______ different shapes. ELECTRON CONFIGURATIONS Electron configurations represent the way electrons are arranged in atoms. The Aufbau principle states that electrons enter the __________________ energy first. This causes difficulties because of the ________________ of orbitals of different energies. At most there can be only 2 electrons per orbital, and they must have __________________ “spins.” Hund’s rule states that when electrons occupy orbitals of equal energy, they don’t _________ up with an electron of opposite spin until they have to. Let’s determine the electron configuration for phosphorus. ______________________________ Let’s determine the electron configuration for chromium. _______________________________ Write the electron configuration for aluminum (Al). ________________________________ Write the electron configuration for neon (Ne). ____________________________________ Write the electron configuration for calcium (Ca). __________________________________ Write the electron configuration for iron (Fe). _____________________________________ Write the electron configuration for bromine (Br). _________________________________ To identify an element with a given electron configuration, add the _________________ numbers together and find the element with that atomic number. Directions: Identify the element with the following electron configuration: a. 1s2 2s2 2p6 3s2 3p4 _________________________________ b. 1s2 2s2 2p6 3s2 3p6 4s2 3d9 _________________________________ c. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 _________________________________ Electron Configuration Using a Noble Gas Abbreviation - In order to write this type of configuration, find the _______________ gas (from Group 8A) that comes before the element in question. Put the symbol for the noble gas in _____________________ and then write the part of the configuration that follows to reach the desired element. 37 Unit 3- Electrons and Periodicity Write the electron configuration using a noble gas abbreviation for: magnesium (Mg) _________________________ • nickel (Ni) ___________________ fluorine (F) _________________________ • silicon (Si) ___________________ zirconium (Zr) _________________________ VALENCE ELECTRONS The electrons in the ______________________ energy level are called valence electrons. You can also use the periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18. All atoms in Group 1, like hydrogen, have __________ valence electron. All atoms in Group 2 have two, in Group 13 have _______, in Group 14 have four, in Group 15 have five, in Group 16 have six, and in Group 17 have ________ valence electrons. All atoms in Group 18 have eight valence electrons, except helium which only has two. All atoms in sublevels d and f have _________ valence electrons. How many valence electrons does each of the following elements have? a) carbon (C) c) iron (Fe) b) bromine (Br) d) potassium (K) e) aluminum (Al) LEWIS DOT DIAGRAMS Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols. A Lewis dot diagram illustrates ___________________ electrons as dots (or other small symbols) around the chemical symbol of an element. Each dot represents _____________ valence electron. In the dot diagram, the element’s symbol represents the core of the atom - the nucleus plus all the _______________ electrons. Write a Lewis dot diagram for a) chlorine 38 b) calcium c) potassium Unit 3- Electrons and Periodicity PERIODIC TABLE HISTORY The Russian chemist, Dmitri ______________________ was studying the properties of the elements and realized that the chemical and physical properties of the elements repeated in an orderly way when he organized the elements according to increasing atomic ___________. Mendeleev later developed an improved version of his table with the elements arranged in horizontal ___________. This arrangement was the forerunner of today’s periodic table. Patterns of changing properties repeated for the elements across the horizontal rows. Elements in vertical ___________________ showed similar properties. Mendeleev grouped elements in columns by similar properties in order of increasing atomic mass. He found some inconsistencies and felt that the properties were more important than the mass, so he switched order. Mendeleev left some _____________ in his periodic table, deciding there must be undiscovered elements. He predicted their properties before they were found. Mendeleev is considered to be the _________________ of the Periodic table. This repeated pattern (when Mendeleev grouped elements in columns by similar properties) is an example of __________________ in the properties of elements. Periodicity is the tendency to recur at regular intervals. By 1860, scientists had already discovered _________ elements and determined their atomic masses. THE MODERN PERIODIC TABLE Fifty years after Mendeleev, the British scientist Henry ________________ discovered that the number of protons in the nucleus of a particular type of atom was always the same. When atoms were arranged according to increasing atomic ___________________, the few problems with Mendeleev's periodic table disappeared. Because of Moseley's work, the modern periodic table is based on the atomic numbers of the elements. The statement that the physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number is known as the periodic _____________. On the periodic table a _________________, sometimes also called a series, consists of the elements in a horizontal row. A __________________, sometimes also called a family, consists of the elements in a vertical column. Elements are placed in columns by similar properties. 39 Unit 3- Electrons and Periodicity The elements in the A groups are called the __________________ elements. The B groups are called the ____________________ elements. The two rows at the bottom of the table are called the inner transition elements. Group 1A elements are the _________________ metals. Group 1A elements have ______ valence electron and form _______ ions after losing the one valence electron. Group 2A elements are the alkaline earth metals. Group 2A elements have ______ valence electrons and form 2+ ions after losing the two __________________ electrons. Group 3A is called the _________________ group. Group 3A elements have ________ valence electrons and form 3+ ions after losing the three valence electrons. Group _______ is called the carbon group. Group 4A elements have four valence electrons and form 4+ ions after ___________________ the four valence electrons or 4- ions after ___________________ four additional electrons. Group 5A is called the _____________________ group. Group 5A elements have five valence electrons and form ________ ions after gaining three more electrons. Group 6A is called the oxygen group. Group 6A elements have _______ valence electrons and form 2- ions after ____________________ two more electrons. Group 7A is called the ____________________. Group 7A elements have seven valence electrons and form 1- ions after gaining one more electron. The word halogen is from the Greek words for “______________ former” so named because the compounds that halogens form with metals are salt-like. Group 8A elements are the ________________ gases. Group 8A elements have eight valence electrons except for helium which only has ________. The noble gases, with a full complement of valence electrons, are generally unreactive. All transition elements have _______ valence electrons. How many valence electrons are in an atom of each of the following elements? a) Magnesium (Mg) ______ b) Selenium (Se) ______ c) Tin (Sn) _____ METALS, NONMETALS AND METALLOIDS Metals are elements that have ________________, conduct ____________ and electricity, and usually bend without breaking. Most metals have one, two, or three valence electrons. All metals except _________________ are solids at room temperature; in fact, most have extremely _____________ melting points. A metal’s ___________________ is its ability to react with another substance. 40 Unit 3- Electrons and Periodicity Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the more active metal. a) cobalt (Co) or manganese (Mn) ________ b) barium (Ba) or sodium (Na) ________ Although the majority of the elements in the periodic table are _________________, many nonmetals are abundant in nature. Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are __________________ when solid. Many are ______________ at room temperature; those that are solids lack the luster of metals. Their _____________________ points tend to be lower than those of metals. With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons. A nonmetal’s reactivity is its ability to react with another substance. Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the less active nonmetal. a) fluorine (F2) or chlorine (Cl2) _________ b) chlorine (Cl2) or iodine (I2) _________ __________________ have some chemical and physical properties of metals and other properties of nonmetals. In the periodic table, the metalloids lie along the border between metals and nonmetals. Some metalloids such as silicon, germanium (Ge), and arsenic (As) are _____________________. A semiconductor is an element that does not conduct electricity as well as a ________________, but does conduct slightly better than a nonmetal. Match each element in Column A with the best matching description in Column B. Each Column A element may match more than one description from Column B. Column A Column B 1. strontium a. halogen 2. chromium b. alkaline earth metal 3. iodine c. representative element d. transition element PERIODIC TRENDS Because the periodic table relates group and period numbers to valence electrons, it’s useful in predicting atomic structure and, therefore, ______________________ properties. 41 Unit 3- Electrons and Periodicity Atomic Radius Atomic radius is half the distance between two __________________ of a diatomic molecule. Atomic size is influenced by two factors: (1) energy level – A _________________ energy level is further away. (2) charge on nucleus - More charge (_________________) pulls electrons in closer. As you go down a ___________________, each atom has another energy level so the atoms get bigger. As you go across a period, the radius gets ____________________. Atoms are in the same energy level, but as you move across the chart, atoms have a greater ___________________ charge (more protons). Therefore, the outermost electrons are closer. Choose the element from the pair with the larger atomic radius. a) lithium (Li) or beryllium (Be) _________ b) silicon (Si) or tin (Sn) _________ Choose the element from the pair with the smaller atomic radius. a) silver (Ag) or gold (Au) _________ b) cesium (Cs) or barium (Ba) _______ Ionic Radius When an atom gains or loses one or more electrons, it becomes a(n) ______________. Because an electron has a negative charge, gaining electrons produces a _______________________ charged ion, an anion, whereas losing electrons produces a positively charged ion, a ________________. As you might expect, the loss of electrons produces a positive ion with a radius that is ___________________ than that of the parent atom. Conversely, when an atom gains electrons, the resulting negative ion is larger than the parent atom. Practically all of the elements to the _____________ of group 4A of the periodic table commonly form positive ions. As with neutral atoms, ___________________ ions become smaller moving across a period and become larger moving down through a group. As you go down a group, you are adding a(n) _________________ level. Ions get bigger as you go down. Most elements to the right of group 4A (with the exception of the noble gases in group 8A) form negative ions. These ions, although considerably larger than the positive ions to the left, also decrease in ______________ moving across a period. Like the positive ions, the negative ions increase in size moving down through a group. Across the period, nuclear charge __________________ so they get smaller. Energy level changes between anions and cations. 42 Unit 3- Electrons and Periodicity Choose the element from the pair with the smaller radius. a) silver (Ag) or the silver ion (Ag1+) _________ b) oxygen (O) or the oxygen ion (O2-) _________ For each of the following pairs, predict which atom is larger. a) Mg, Sr _________ b) Sr, Sn _________ d) Ge, Br _________ e) Cr, W _________ c) Ge, Sn _________ For each of the following pairs, predict which atom or ion is larger. a) Mg, Mg2+ _________ b) S, S2– _________ d) Cl–, I– _________ e) Na+, Al3+ _________ c) Ca2+, Ba2+ ______ Ionization Energy Ionization energy (IE) is the amount of energy required to completely _____________________ an electron from a gaseous atom. Removing one electron makes a ________ ion. The energy required to do this is called the first ionization energy. The _____________________ the nuclear charge (# of protons), the greater IE. The distance from the ____________________ increases IE. As you go down a group, first IE decreases because the electron is further away, thus there is more shielding by the _______________ electrons from the pull of the positive nucleus. All the atoms in the same period have the same energy level. They have the same shielding, but as you move across the chart there is a(n) _____________________ nuclear charge. Therefore, IE generally increases from left to right. Choose the element from the pair with the greater ionization energy. a) silver (Ag) or iodine (I) _________ b) oxygen (O) or selenium (Se) ________ Choose the element from the pair with the smaller ionization energy. a) chromium (Cr) or tungsten (W) ______ b) sodium (Na) or magnesium (Mg) _______ Electronegativity Electronegativity is the tendency for an atom to ___________________ electrons to itself when it is chemically combined with another element. Large electronegativity means it _______________ the electron toward it. The further you go down a group, the farther the electron is away from the nucleus and the _____________ electrons an atom has. It is harder to attract extra electrons if the available energy level is far from the nucleus, so the 43 Unit 3- Electrons and Periodicity electronegativity _____________________. As you go across a row, electronegativity increases as the ________________________ character of the elements decreases. Choose the element from the pair with the greater electronegativity. a) sodium (Na) or rubidium (Rb) _______ b) selenium (Se) or bromine (Br) _______ Choose the element from the pair with the smaller electronegativity. a) magnesium (Mg) or calcium (Ca) _______ b) nitrogen (N) or oxygen (O) _______ Homework / Practice Write the configuration notation for each of the following elements: 1) sodium 3) bromine 2) iron 4) barium Write the noble gas notation for each of the following elements: 5) cobalt 7) tellurium 6) silver 8) radium Determine what elements are denoted by the following electron configurations: 9) 1s22s22p63s23p4 11) [Kr] 5s24d105p3 10) 1s22s22p63s23p64s23d104p65s1 12) [Rn] 7s25f11 Write the orbital notation for the following: 13) carbon 14) neon 15) sulfur 16) P 17) B 18) Na Write configuration notation for atoms containing the following number of electrons: 19) 3 20) 6 21) 8 22) 13 Draw the Lewis Dot Notation for the following elements 23) Sodium 25) Silver 27) Antimony 24) Sulfur 26) Aluminum 28) Argon 44 Unit 3- Electrons and Periodicity Reference Labeling- Label the following on the blank periodic table on the next page-29) Metals 30) Nonmetals 31) Metalloids 32) Transition metals 33) Actinides 34) Lanthanides 35) Alkali metals 36) Alkaline earth metals 37) Halogens 38) Noble gases (inert gases) 39) Group numbers: 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A 40) Label the most likely charge of an ion from groups 1A, 2A, 5A, 6A, 7A, and Noble gases 41) Show which way are periods (left to right or up and down) 42) Show which way are groups/families (left to right or up and down) 43) Using color, identify the s, p, d and f blocks on the table 44) Label the rows with energy level numbers 45) Using arrows, indicate the trends of increasing atomic radius, ionization energy, and electronegativity (from left to right and top to bottom) 46) Put a star in the box for each diatomic element 47) Show the charge of the ions that form for silver (+1), zinc (+2), and aluminum (+3) 45 Unit 3- Electrons and Periodicity 46 Unit 3- Electrons and Periodicity Electrons and Periodicity Practice Test Directions: Match each of the following terms with a number or chemical symbol from the periodic table below. 1. Alkaline earth metals: 3. Noble gases 2. 4. The transition metals Halogens: 5. Draw the orbital notation for Sodium 6. Given the electron configuration, identify the element 1s2 2s2 2p6 3s2 2p6 3s2 3p6 4s2 3d7 7. Write the complete configuration notation for Silver 8. Write the shorthand method (Noble Gas notation) for Antimony: 9. Give the energy level for the valence electrons in Helium 10. Determine the color of light emitted when an electron jumps from the following quantum levels n=4 to n=2 11. How many valence electrons does Carbon have? 12. Draw the Lewis Dot notation for Sodium 13. Describe why the atomic radius of elements increases as you go down a group. 47 Unit 3- Electrons and Periodicity Multiple Choice Practice 14. The two main parts of an atom are the a. b. c. d. Principle energy levels and energy sublevels Nucleus and kernel Nucleus and energy levels Planetary electrons and energy levels 15. The sublevel that has only one orbital is identified by the letter a. s b. p c. d d. f 16. The sublevel that can be occupied by a maximum of ten electrons is identified by the letter a. s b. p c. d d. f 17. An orbital may never be occupied by a. 1 electron b. 2 electrons c. 3 electrons d. 0 electrons 18. An atom of beryllium consists of 4 protons, 5 neutrons, 4 electrons. The mass number of this atom is a. 13 b. 9 c. 8 d. 5 19. Which of the following is the correct electron configuration for the bromide ion, Br-? a. [Ar] 4s24p5 b. [Ar] 4s23d104p5 c. [Ar] 4s23d104p6 d. [Ar] 4s23d104p65s1 e. [Ar] 4s23d103p6 20. Which is the first element to have 4d electrons in its electron configuration? a. Ca b. Sc c. Rb d. Y e. La 21. When electrons in an atom in an excited state fall to lower energy levels, energy is a. absorbed, only b. released, only c. neither released nor absorbed d. both released and absorbed 22. Which of the following elements has the greatest electronegativity? a. Mg b. K c. S d. F 23. Which of the following elements would react with chlorine in a one to one ratio a. Mg b. K c. S 48 d. F Unit 3- Electrons and Periodicity 24. Which of the following elements would have the smallest radius a. Mg b. K c. S d. F 25. Which of following elements has the lowest first ionization energy a. Mg b. K c. S d. F c. S d. F 26. Which of the following elements is an alkali metal? a. Mg b. K 27. Which element's ionic radius is smaller than its atomic radius? a. neon b. nitrogen c. sodium d. sulfur 28. Which three groups of the Periodic Table contain the most elements classified as metalloids (semimetals)? a. 1, 2, and 13 b. 2, 13, and 14 c. 14, 15, and 16 d. 16, 17, and 18 29. Which element has the highest first ionization energy? a. sodium b. aluminum c. calcium d. phosphorus 30. Which of the following elements has the smallest atomic radius? a. nickel b. cobalt c. calcium d. potassium 31. Which set of elements contains a metalloid? a. K, Mn, As, Ar b. Li, Mg, Ca, Kr c. Ba, Ag, Sn, Xe d. Fr, F, O, Rn 32. Atoms of elements in a group on the Periodic Table have similar chemical properties. This similarity is most closely related to the atoms' a. number of principal energy levels b. number of valence electrons c. atomic numbers d. atomic masses 33. As atoms of elements in Group 16 are considered in order from top to bottom, the electronegativity of each successive element a. decreases b. increases 49 c. remains the same Unit 3- Electrons and Periodicity 34. An atom of which of the following elements has the greatest ability to attract electrons? a. silicon b. sulfur c. nitrogen d. chlorine 35. At STP, which substance is the best conductor of electricity? a. nitrogen b. neon c. sulfur d. silver 36. A strontium atom differs from a strontium ion in that the atom has a greater a. number of electrons b. number of protons c. atomic number d. mass number 37. Which gas is monatomic at STP? a. chlorine b. fluorine c. neon d. nitrogen 38. How many valence electrons does an oxygen atom have? a. 2 b. 6 c. 8 d. 16 39. The identity of an element is determined by... a. the number of its protons. b. the number of its neutrons. c. the number of its electrons d. its atomic mass. 40. Which of the following electron configurations represents the electron configuration for a magnesium cation (Mg2+)? a. 1s22s22p63s2 b. 1s22s22p63s23p2 c. 1s22s22p6 d. 1s22s22p4 41. Which of the following atoms has the largest diameter? a. F b. Cl c. Br d. I 42. Which of the following elements has the greatest electronegativity? a. Si b. P c. N d. O 43. Which scientist noted a definite pattern in valence numbers and arranged an early periodic table in order of the elements atomic mass? a. Enrico Fermi b. Dmitri Mendeleev c. Albert Einstein d. Madame Curie 44. The periodic law states that the properties of elements are periodic functions of their a. Mass b. Symbol c. atomic number d. valence 50 Unit 3- Electrons and Periodicity 45. Which of the following is a noble gas? a. Sodium b. Gold c. Chlorine d. Neon 46. A gas is called "noble" because a. b. c. d. it is normally unreactive it is normally inert it has a complete outer energy level of electrons all of the above 47. Of the following elements, the one that forms cations with varying positive charges is: a. iron, Fe b. sodium, Na c. aluminum, Al d. strontium, Sr e. nitrogen, N 48. Which of the following statements is incorrect?: a. b. c. d. e. metals generally form cations nonmetals are generally poor conductors of electricity metals are malleable nonmetals are generally brittle metalloids are metals with some nonmetallic characteristics 49. Which of the following statements are true? a. It is easier to remove an electron from Na than from Na+. b. As the atomic number increases within a group of the representative elements, the tendency is for first ionization energy to increase. c. All particles with the electron configuration [Ar]4s2 have the same ionization energy. d. The first ionization energy of fluorine is greater than the first ionization energy of oxygen. e. All are false 50. An element having the configuration [Xe]6s1 belongs to the Group: a. alkali metals b. halogens c. alkaline earth metals d. None of these e. noble gases 51 Unit 3- Electrons and Periodicity 51. How many unpaired electrons are there in an atom of tin in its ground state? a. 4 b. 0 c. 3 d. 2 e. 1 52. Of the following elements, which one is most likely to form an ion through the loss of two electrons? a. strontium b. chlorine c. aluminum d. sodium e. sulfur 53. Which of the following particles has the greatest atomic radius? a. Al b. Si c. S d. Al3+ e. P 54. Which of the following forms of electromagnetic radiation has the shortest wavelength? a. ultraviolet b. radio waves c. infrared d. visible light e. microwaves 55. For which of the following transitions does the light emitted have the shortest wavelength? a. n = 4 to n = 2 b. n = 2 to n = 1 c. n = 5 to n = 3 d. n = 4 to n = 3 e. n = 3 to n = 2 56. Researchers at Lawrence Berkeley National Lab have recently formed a new synthetic element with atomic number 118 and mass number 293. Which of the following elements would have chemical properties most similar to this new element? a. Ir c. Ta e. S b. Xe d. Pb BONDING As atoms bond with each other, they _____________________ their potential energy, thus creating more stable arrangements of matter. The force that holds two ________________ together is called a chemical bond. There are 3 types of bonding: ionic, ___________________, and metallic. The number of valence electrons are easily found by looking up the group number on the periodic table. 52 Unit 4 – Types of Bonding Group 1A (Li, Na, K, etc.): 1 valence electron Group 2A (Be, Mg, Ca, etc.): ______ valence electrons Group 3A (B, Al, Ga, etc.): 3 valence electrons Group 4A (C, Si, Ge, etc.): 4 valence electrons Group 5A (N, P, As, etc.): _______ valence electrons Group 6A (O, S, Se, etc.): 6 valence electrons Group 7A (F, Cl, Br, etc.): 7 valence electrons Group 8A (He, Ne, Ar, etc.): 8 valence electrons (except He has _______ valence electrons) Recall the rules for drawing Lewis dot diagrams: Write the _____________________ of the element. Put one dot for each valence electron. Don’t _____________ electrons up until you have to. Electron Configurations and Electron Dot Diagrams for Cations Metals lose electrons to attain noble gas configuration. They make positive ions, ____________. If we look at an electron configuration, it makes sense. Example: Sodium (Na), 1s22s22p63s1, has _________ valence electron(s). The electron that is removed comes from the ____________ energy level. As a result of the loss of the electron, the sodium ion (Na+) has the following electron configuration: 1s22s22p6 Calcium has 2 valence electrons. These will come off, forming a positive ion. 53 Unit 4 – Types of Bonding Electron Configurations and Electron Dot Diagrams for Anions Nonmetals gain electrons to attain noble gas configuration. This means they want a(n) ________________ of electrons, 8 electrons. They make negative ions, ___________________. If we look at an electron configuration, it makes sense. Example: Sulfur (S), 1s22s22p63s23p4, has _______ valence electrons and needs to gain 2 more to have an octet. The sulfur ion (S-2) has the same electron configuration as a noble gas: 1s22s22p63s23p6 Phosphorous has 5 valence electrons. It will gain _________ electrons to fill the outer shell. Stable Electron Configurations All atoms react to achieve __________________ gas configuration. Noble gases, except He, have 2 s electrons and 6 p electrons, totaling 8 valence electrons. They obey the ____________________ rule. IONIC BONDING Anions and cations are involved in ionic bonding and are held together by __________________ charges, electrostatic attraction. The bond is formed through the ______________________ of electrons. Electrons are transferred to achieve noble gas configuration. Ionic bonds occur between _________________ and nonmetals. All the electrons must be accounted for! A compound that is composed of _______________ is called an ionic compound. Note that only the arrangement of electrons has changed. Nothing about the atom’s nucleus has changed. Ionic compounds have a _______________________ structure, a regular repeating arrangement of ions in the solid. Even though the ions are ___________________ bonded to one another, ionic compounds are __________________. Strong repulsion breaks crystal apart. The structure is rigid. They have _______________ melting points because of strong forces between ions. They also conduct electricity in the _________________ and dissolved states. Any compound that conducts electricity when melted or dissolved in water is a(n) ___________________________. How many valence electrons must an atom have in its outer energy level in order to be considered stable? The energy required to separate one mole of the ions of an ionic compound is called ____________________ energy, which is expressed as a negative quantity. The greater (that is, the more negative) the lattice energy is, the ______________________ the force of attraction 54 Unit 4 – Types of Bonding between the ions. Lattice energy tends to be __________________________ for more-highlycharged ions (those atoms that have more electrons to give or those atoms that can take more electrons). Lattice energy also tends to be greater for __________________ ions. Between the following ionic compounds, which would be expected to have the higher (more negative) lattice energy? LiF or KBr Between the following ionic compounds, which would be expected to have the higher (more negative) lattice energy? NaCl or MgS The electronegativity difference for two elements in an ionic compound is greater than or equal to _______________. COVALENT BONDING A _______________________ is an uncharged group of two or more atoms held together by covalent bonds. Covalent compounds occur between two ___________________ or a nonmetal and hydrogen. The attraction of two atoms for a shared _______________ of electrons is called a covalent bond. In a covalent bond, atoms share electrons and neither atom has an ionic ______________________. Covalent bonds occur between 2 ___________________________ because nonmetals hold onto their valence electrons. They can’t give away electrons to bond, yet, they still want _______________ gas configuration. They get it by sharing valence electrons with each other. By sharing both atoms get to count the electrons toward noble gas configuration. A ____________________ bond is formed from the sharing of two valence electrons. The electronegativity difference for two elements in a covalent compound is between _________ and 1.7. Do atoms that share a covalent bond have an ionic charge? Sometimes atoms share more than one pair of valence electrons. A ____________________ bond is when atoms share two pair of electrons, 4 electrons. A triple bond is when atoms share three pair of electrons, _____ electrons. Triple bonds are ________________________ and shorter than double bonds. Double bonds are stronger and shorter than ______________________ bonds. 55 Unit 4 – Types of Bonding THE WETTER WAY You can easily determine the number of bonds in a compound by performing the “Wetter Way.” # bonds = (Σe- after – Σe- before ) / 2 The number of electrons before bonding is equal to the __________________________ number. To get the number of electrons after bonding, double the number of electrons before bonding, BUT DO NOT EXCEED ___________! Example: CO2 C is in column 4A and therefore has 4 valence electrons before bonding. O is in column 6A and therefore has 6 valence electrons before bonding. Σ electrons before bonding Carbon = 4 Two oxygens = 6 x 2 = 12 (The formula CO2 implies there are 2 oxygen atoms.) Σ electrons before bonding = 4 + 12 = 16 Σ electrons after bonding C has 4 valence e- before bonding, so it has (4x2) = 8 electrons after bonding O has 6 valence e- before bonding, so it has (6x2) = 12 electrons after bonding Remember 8 is the maximum electrons after bonding, so oxygen can only have 8 electrons even though we calculated 12. Carbon = 8 Two oxygens = 8 x 2 = 16 Σ electrons after bonding = 8 + 16 = 24 # bonds = (Σe- after – Σe- before) / 2 = (24 - 16) / 2 = 4 bonds The element you have only 1 of goes in the center. The other elements surround it. Connect the elements with a single line (a single bond). You have only used 2 of your calculated 4 bonds, so you need to double up. A line represents 2 electrons. Count your lines for each element to determine if extra electrons need to be added. Carbon has 4 lines attached which represent 8 electrons. No extra electrons are needed around carbon. Each oxygen has 2 lines attached which represent 4 electrons. Oxygen needs 8 electrons after bonding, so each oxygen needs 4 electrons. Determine the number of bonds and draw the dot-dash diagram for HBr. Determine the number of bonds and draw the dot-dash diagram for N2. 56 Unit 4 – Types of Bonding METALLIC BONDING The bonding in metals is explained by the _______________________ ____________ model, which proposes that the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic __________________. These delocalized electrons are not held by any specific atom and can ________________ easily throughout the solid. Metal atoms release their valence electrons into a sea of electrons shared by all of the metal atoms. The bond that results from this _____________________ pool of valence electrons is called a metallic bond. Metals hold onto their valence electrons very _______________________. Think of them as positive ions floating in a sea of electrons. Because electrons are free to move through the solid, metals conduct _______________________. Metals generally have extremely ______________ melting points because it is difficult to pull metal atoms completely away from the group of cations and attracting electrons. Metals are ________________________ (able to be hammered into sheets). Metals are also ________________________ (able to be drawn into wire) because of the mobility of the particles. Electrons allow atoms to slide by. A mixture of elements that has metallic properties is called a(n) _____________________. Homework / Practice Complete the table by identifying the charge of each of the elements listed and then indicating the formula for ionic compounds formed between the two substances Charge Charge Na O N P S Cl 21+ Na2O Mg Ca Al Li Zn What type of boding is present in the following compounds: 1) SbBr3 2) Ag 3) MgBr2 4) ClO2 5) KCl 6) Fe 57 7) PbO 8) FeCl2 F Unit 4 – Types of Bonding What type of boding is present in the following compounds: 9) NI3 10) CO2 11) Ni 12) Au 13) LiF 14) Al2O3 15) N2O3 16) Mg3P2 17) CCl4 18) H2O POLARITY and VSEPR How each atom fares in a tug-of-war for shared electrons is determined by comparing the _________________________________ of the two bonded atoms. Recall that electronegativity is the measure of the ability of an atom in a bond to ________________________ electrons. Atoms with large electronegativity values, such as fluorine, attract shared valence electrons more __________________ than atoms such as sodium that have small electronegativities. Electronegativity is a periodic property. With only a few exceptions, electronegativity values_____________________ as you move from left to right in any period of the periodic table. Within any group, electronegativity values decrease as you go ___________________ the group. Fluorine has the highest value of ____________. The greater the difference between the electronegativities of the bonding atoms, the more _____________________________ the electrons are shared and the more polar the bond. If the electronegativity difference between the two elements in question is: between 0.0 – 1.7, the bond is ______________________ greater than 1.7, the bond is _____________________ When the electronegativity difference in a bond is 1.7 or greater, the sharing of electrons is so unequal that you can assume that the electron on the less electronegative atom is ________________________ to the more electronegative atom. For example, ∆EN for cesium and fluorine is 4.0 − 0.7 = 3.3. Therefore the bond is _________________. COVALENT BONDS AND POLARITY When the atoms in a bond are the same, the electrons are shared ________________________. This results in a _______________________ covalent bond. _________________________ elements (H2, O2, N2, Cl2, Br2, I2, F2 and At2) have pure nonpolar covalent bonds. All other covalent bonds are polar. The electron sharing is not equal, but it is not so unequal that a complete _____________________ of electrons takes place. 58 Unit 4 – Types of Bonding Consider hydrogen and chlorine. Hydrogen has an electronegativity of 2.20, and chlorine has an electronegativity of 3.16. The ________ pulls harder on the electrons because its electronegativity is greater. The electrons spend more time near the Cl. These symbols, __________________ plus (δ+) and delta minus (δ-), represent a partial positive charge and a partial negative charge. Polar molecules are molecules with a positive and a negative ______________. This requires two things to be true: The molecule must contain _______________ bonds. (This can be determined from differences in electronegativity.) Symmetry cannot ______________________ out the effects of the polar bonds. (Must determine geometry first.) In the following compounds, determine whether the molecule is polar or nonpolar a. hydrogen fluoride (HF) d. ammonia (NH3) b. water (H2O) e. carbon dioxide (CO2) c. carbon tetrachloride (CCl4) VSEPR VSEPR stands for Valence Shell ______________________ ________________ Repulsion. It predicts three-dimensional geometry of molecules. The valence shell includes the ______________________ electrons. The electron pairs try to get as far away as possible to _______________________ repulsion. You can determine the angles of the bonds. VSEPR is based on the number of pairs of valence electrons, both bonded and unbonded. An unbonded pair of electrons is referred to as a _______________pair. Use the Wetter Way to calculate the number of bonds and then draw the dot-dash diagram. The shape of the molecule and bond angle can be determined from this diagram. LINEAR Consider the simplest molecule that exists—hydrogen, H2. Determine the number of bonds using the Wetter Way. H is in column 1A and therefore has 1 valence electron before bonding. Σ electrons before bonding Two hydrogens = 1 x 2 = 2 H – 1 valence e- before bonding, so it has (1x2) = 2 electrons after bonding Σ electrons after bonding 59 Unit 4 – Types of Bonding Two hydrogens = 2 x 2 = 4 (4 – 2) / 2 = 1, so H2 has 1 bond! Connect the elements with a single bond. You have used all of the calculated bonds. Each hydrogen has 1 line attached which represents _______________ electrons. No extra electrons are needed around hydrogen to have the 2 electrons needed after bonding. A hydrogen molecule is linear. The electrons attempt to maximize their distance from one another by having bond angle of ____________°. Linear compounds are NOT ____________________. TETRAHEDRAL Consider CH4. The Wetter Way shows that CH4 has __________ bonds! The element you have only one of goes in the _______________________. The other elements surround it. Connect the elements with a single _________________ (a single bond). You have used all 4 of your calculated bonds. Remember a line represents ___________ electrons. Count your lines for each element to determine if extra electrons need to be added. Carbon has 4 lines attached which represent _________ electrons. No extra electrons are needed around carbon. Each hydrogen atom has one line attached which represents 2 electrons. No extra electrons are needed around hydrogen. Single bonds fill all atoms. There are _________ bond pairs of electrons pushing away. The electrons can _________________ their distance from one another by forming a 3-D shape. The furthest they can get away is ___________º. This basic shape is a tetrahedral, a pyramid with a triangular base. The tetrahedral is the shape for everything with 4 bond pairs and ____________ lone pairs around the central atom. TRIGONAL PYRAMIDAL Perform the Wetter Way for phosphorous trichloride (PCl3). How many bonds are in this molecule? ________ Sketch the dot-dash diagram for phosphorous trichloride. Please include all electrons. Only the electrons around the ______________________ atom affect the shape. The shape is a basic _______________________________ but you can’t see the lone pair. The shape is called trigonal pyramidal. The bond angle is ____________° between the chlorines because the electron pair forces the chlorines closer to each other. 60 Unit 4 – Types of Bonding BENT Perform the Wetter Way for water (H2O). How many bonds are in this molecule? _____ Sketch the dot-dash diagram for water. Please include all electrons. Only the electrons around the central atom affect the shape. The shape is still basic tetrahedral, but you can’t see the _________ lone pairs. The shape is called bent. The bond angle between hydrogens is ____________°. TRIGONAL PLANAR Perform the Wetter Way for H2CO. How many bonds are in this molecule? _____ Sketch the dot-dash diagram for H2CO. Please include all electrons. (Carbon is the central atom.) The farthest you can get the elements apart is __________º. The shape is flat and called trigonal planar. Determine the number of bonds, draw the dot-dash diagram, state the VSEPR shape and provide the bond angle for the following compounds a. CO2 b. BCl3 c. SCl2 d. SiF4 HYBRID ORBITALS Hybrid orbitals combine bonding with geometry. ______ hybridization has a linear shape. sp2 hybridization has a trigonal ____________________ shape. ____________ hybridization has tetrahedral, trigonal pyramidal and bent shapes. Homework/Practice Draw the Lewis structure for each of the following compounds, identify the shape of the molecule, and identify the polarity of the molecule. 1. CCl4 4. H2CS 2. BF3 5. SiO2 3. NF3 6. H2S 61 Unit 4 – Types of Bonding INTERMOLECULAR FORCES Intermolecular forces are forces of _______________________. They are what make solid and liquid molecular compounds possible. The three intermolecular forces are _________________ bonds, dipole–dipole forces and London ____________________________ forces. Hydrogen Bonding A hydrogen bond is a _________________________________________ attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least ____________ lone electron pair. For a hydrogen bond to form, hydrogen must be bonded to a fluorine, __________________________, or nitrogen atom. F, O, and N are very electronegative so it is a very _______________________ dipole. Hydrogen bonding is the _________________________ of the intermolecular forces. Examples include H2O, NH3, and HF. Dipole-dipole Forces Polar molecules contain ___________________________ dipoles; that is, some regions of a polar molecule are always ___________________________ negative and some regions of the molecule are always partially positive. Attractions between _____________________________ charged regions of polar molecules are called dipole–dipole forces. Neighboring polar molecules orient themselves so that oppositely charged regions _______________ up. Opposites attract but are not completely hooked as in ionic solids. Dipole-dipole forces depend on the number of _______________________. Bigger molecules result in more electrons, and more electrons mean ________________________ forces. Dipole–dipole forces are stronger than dispersion forces as long as the molecules being compared have approximately the same mass. Examples of compounds that exhibit dipole-dipole forces include CO, HCl, and PH3. London Dispersion Forces Dispersion forces are ____________________ forces that result from temporary shifts in the ______________________ of electrons in electron clouds. Remember that the electrons in an electron cloud are in constant _____________________. When two nonpolar molecules are in 62 Unit 4 – Types of Bonding close contact, especially when they collide, the electron cloud of one molecule _______________________ the electron cloud of the other molecule. The electron density around each nucleus is, for a moment, greater in one region of each cloud. Each molecule forms a __________________________ dipole. When temporary dipoles are close together, a weak dispersion force exists between oppositely charged regions of the dipoles. Due to the temporary nature of the dipoles, dispersion forces are the __________________________ intermolecular force. Dispersion forces exist between ____________ gases and compounds that are nonpolar. Examples include Ar, Cl2, Br2, CH4, and CO2. Dispersion forces ______________________ as the mass of the molecule increases. C2H6 (MW = 30.0 g/mol) has stronger dispersion forces than CH4 (MW = 16.0 g/mol). This difference in dispersion forces explains why fluorine and chlorine are gases, bromine is a __________________________, and iodine is a solid at room temperature. The molecular mass of iodine is greater than that of bromine, and bromine has a greater mass than chlorine. Intermolecular Forces To determine what type of intermolecular force a compound has, ask yourself the following questions. Does the compound contain hydrogen attached to N, O, or F? o If yes, the force is hydrogen bonding. Determine the number of bonds from the Wetter Way and draw the dash-dot diagram. Does the central element of the compound contain any lone pairs of electrons? o If yes, the force is dipole-dipole. Does the central element of the compound contain ZERO lone pairs of electrons? o If yes, the force is dispersion. Determine the type of intermolecular force in each of the following compounds 1) BCl3 _____________________________ 2) Xe _____________________________ 3) NH3 _____________________________ 4) CH4 _____________________________ 5) SO2 _____________________________ 6) H2 _____________________________ 7) SO3 _____________________________ 8) CH3Cl ___________________________ 9) HF _____________________________ 10) HBr ____________________________ 63 Unit 4 – Types of Bonding Types of Bonding Practice Test 1. In a complete sentence, compare and contrast metallic bonds and ionic bonds. Directions- For each of the following pairs of elements, write the formula for the ionic compound that would form between them 2. K and Cl 5. Calcium and Chlorine 3. Na and N 6. Zinc and Sulfur 4. Al and O 7. Lithium and Phosphorous Directions- Draw the Lewis structure , Identify the shape of the molecule, Identify the polarity of the bonds, Identify the polarity of the molecule, Identify the IMF that would be exhibited 8. CCl4 9. SF2 10. SiO2 11. BI3 12. PCl3 13. N2 14. What does IMF stand for? Which of the three IMF’s is the weakest? Multiple Choice Practice 15. What type of bond exists between atoms of potassium and chloride in a crystal of potassium chloride? a. Hydrogen bond d. Nonpolar covalent bond b. Ionic bond e. Metallic bond c. Polar covalent bond 16. What type of bond exists between atoms in a nitrogen molecule? a. Hydrogen bond d. Nonpolar covalent bond b. Ionic bond e. Metallic bond c. Polar covalent bond 17. What type of bond exists between atoms of calcium in a crystal of calcium? a. Hydrogen bond d. Nonpolar covalent bond b. Ionic bond e. Metallic bond c. Polar covalent bond 18. All of the following have covalent bonds except a. HCl c. H2O b. CCl4 d. CsF 19. Which of the following atoms normally forms monatomic molecules? a. Cl c. O b. H d. N 64 e. CO2 e. He Unit 4 – Types of Bonding 20. The complete loss of an electron of one atom to another atom with the consequent formation of electrostatic charges is said to be a. Covalent bonding c. Ionic bonding b. Polar covalent bonding d. Coordinate covalent bonding 21. When a metal atom combines with a nonmetal atom, the nonmetal atom will a. lose electrons and decrease in size b. lose electrons and increase in size c. gain electrons and decrease in size d. gain electrons and increase in size 22. Which formula represents a molecular substance? a. CaO b. CO c. Li2O d. Al2O3 23. Which combination of atoms can form a polar covalent bond? a. H and H c. N and N b. H and Br d. Na and Br 24. Fluorine atoms tend to.______.when they form chemical compounds with metals. a. lose electrons b. gain electrons c. neither lose nor gain electrons...they usually share electrons equally with metals. d. Fluorine atoms do not form compounds with other atoms...fluorine is an inert gas. 25. What is a compound composed of? a. two or more different elements that are physically combined in a fixed proportion b. two or more different mixtures that are physically combined in a fixed proportion c. two or more different elements that are chemically combined in a fixed proportion d. two or more different elements that are chemically combined in a variable proportion 26. How can a chemical compound be broken? a. can be broken down by physical means b. can be broken down by chemical means c. cannot be broken down d. can be broken down by physical or chemical means 27. Which of the following compounds is most likely to be ionic? a. CO2 c. FeCl3 b. CCl4 d. MgCl2 65 e. HBr Unit 4 – Types of Bonding 28. How many unshared electron pairs must be included in the Lewis structure for water, H2O? a. 3 c. 1 e. 0 b. 2 d. 4 29. Nitrogen triiodide, NI3, is an unstable molecule that is used as a contact explosive. Its molecular structure is: a. none of these d. tetrahedral b. octahedral e. pyramidal c. square planar 30. In which of the following compounds does the bond between the central atom and chlorine have the greatest ionic character? a. BCl3 c. CCl4 e. CaCl2 b. FeCl2 d. HCl 31. Which of the following molecules must contain at least one double bond a. H2O d. CH3I b. CCl4 e. CH3COOH c. H2O2 32. The Lewis structure for hydrogen cyanide is: c. a. d. b. e. 33. In the Lewis structure for CH2Cl2, the number of unshared electron pairs is: a. 10 c. 2 b. 8 d. 4 e. 6 34. The only intermolecular forces existing between oxygen molecules are: a. ion-ion attractive forces d. nuclear forces b. hydrogen bonding forces e. London dispersion forces c. permanent dipole forces 66 Unit 4 – Types of Bonding 35. Reactions between alkali metals and phosphorous result in compounds with the formula: a. M3P d. M2P3 b. None of these e. MP3 c. M2P 36. Which of the following arrangements represent ions? i. ii. iii. iv. v. 12 protons, 12 neutrons, 11 electrons 12 protons, 11 neutrons, 12 electrons 11 protons, 12 neutrons, 12 electrons 11 protons, 12 neutrons, 11 electrons 12 protons, 12 neutrons, 12 electrons a. 1 and 2 b. 1, 3, and 4 c. 1 and 3 d. 2 and 3 e. all of these are ions 37. A particle X contains 10 electrons, seven neutrons and has a net charge of 3-. The particle is: a. a nitride ion d. a neon ion b. obviously polyatomic e. none of these are correct c. an oxide ion 38. Ions are formed in chemical reactions by: i. Gaining electrons ii. Losing electrons iii. Gaining protons iv. Losing protons v. All of these a. 1 and 2 are correct b. 3 and 4 are correct c. 5 is correct d. 1 and 3 are correct e. 2 and 4 are correct 67 Unit 5 – Nomenclature NAMING COMPOUNDS AND WRITING FORMULAS Recall that the periodic table is more than a list of elements. Elements are put in columns because of _____________________ properties. Each column is called a group. A compound is made of two or more ______________________. The name should tell us how many and what type of atoms. There are two types of compounds: ___________________ compounds and molecular compounds. The simplest ratio of the ions represented in an ionic compound is called a ______________________ unit. The overall charge of any formula unit is ________________. In order to write a correct formula unit, one must know the charge of each ion. Atoms are electrically _____________________. They have the same number of protons and electrons. ________________ are atoms, or groups of atoms, with a charge. Ions have a different numbers of electrons. An anion is a _____________________ ion. An anion has gained electrons. Nonmetals can ________________ electrons. The charge is written as a superscript on the right. F1- has gained _________ electron. O2- has gained __________ electrons. A ___________________ is a positive ion. It is formed by __________________ electrons. There are more _____________________ than electrons. ______________________ form cations. K1+ has lost one electron. Ca2+ has lost __________ electrons. The charges of monatomic ions, or ions containing only one atom, can often be determined by referring to the periodic table or table of common ions based on group number. The charge of a monatomic ion is equal to its _________________________ number. For most of the Group ________ elements, the Periodic Table can tell what kind of ion they will form from their location. Elements in the same group have similar properties, including the charge when they are ions. NAMING CATIONS We will use the systematic way. For cations, if the charge is always the same (Group A) just write the _________________ of the metal. Transition metals (as well as tin and lead) can have more than one type of charge. The charge is indicated with ___________________ numerals in parenthesis. Zinc (Zn2+) and silver (Ag1+), although transition metals, only have __________ possible charge. Roman numerals ARE NOT used for zinc and silver. Li1+ is called the ____________________ ion. __________ is called the Strontium ion. Fe2+ is called the Iron (II) ion. Iron is a transition metal, so the charge is not always the same. The name of the metal 68 Unit 5 – Nomenclature is written, and the charge is denoted with Roman numerals in parenthesis. Pb2+ is called the Lead __________ ion. Name the following cations. a) Ca2+ _________________________ b) Al3+ ___________________________ c) Sn4+ _________________________ d) Na+ _________________________ e) Fe3+ _________________________ f) Cu+ _________________________ WRITING FORMULAS FOR CATIONS Write the formula for the metal. If a Roman numeral is in parenthesis use that number for the _____________________. Indicate the charge with a superscript. If no Roman numeral is given, find the Group A metal on the periodic table and determine the charge from the _____________________ number. The formula for the nickel (II) ion is Ni2+. The formula for the gallium ion is ____________. Write the formulas for the following cations. a) magnesium ion ________________ b) copper (II) ion ___________________ c) potassium ion ________________ d) silver ion _________________ e) chromium (VI) ion ________________ f) mercury (II) ion ________________ NAMING ANIONS Naming monatomic anions is always the same. Change the element ending to – ___________. F is the symbol for fluorine, F1- is fluoride. Cl1- is called the chloride ion. _______ is called the oxide ion. Name the following anions. a) S2- _________________________ b) Br1- ___________________________ c) N3- _________________________ d) As3- _________________________ e) Te2- _________________________ WRITING FORMULAS FOR ANIONS Write the formula for the nonmetal. Find the Group A nonmetal on the periodic table and determine the charge from the column number. 69 Unit 5 – Nomenclature Write the formulas for the following anions. a) iodide ion ________________ b) phosphide ion ___________________ c) selenide ion ________________ d) carbide ion _________________ IONIC COMPOUNDS Oxidation numbers can be used to determine the chemical formulas for ionic compounds. If the oxidation number of each ion is _________________________ by the number of that ion present in a formula unit, and then the results are added, the sum must be _______________. In the formula for an ionic compound, the symbol of the _________________ is written before that of the anion. Subscripts, or small numbers written to the lower ______________________ of the chemical symbols, show the numbers of ions of each type present in a formula unit. BINARY IONIC COMPOUNDS Binary ionic compounds are composed of a metal bonded with a ________________________. Name the metal ion using a Roman numeral in parenthesis if necessary. Follow this name with the name of the nonmetal ion. Name the following binary ionic compounds. a) NaCl __________________ b) Ca3P2 __________________ c) CuO __________________ d) SnBr2 __________________ e) Fe2S3 __________________ f) AlF3 __________________ g) KCl __________________ h) Na3N __________________ i) CrN __________________ j) PbO2 __________________ Write the symbol for the metal. Determine the oxidation number from either the column number or the Roman numeral and write it as a superscript to the right of the metal’s symbol. To the right of the metal’s symbol, write the symbol for the nonmetal. Determine the oxidation number from the column number and write it as a superscript to the right of the nonmetal’s symbol. Example: potassium fluoride - K1+ F1- If the two oxidation numbers add together to get zero, the formula is a one-to-one ratio of the elements. Answer = KF Example: aluminum sulfide - Al3+ S2- If the two oxidation numbers DO NOT add together to get zero, you will need to “criss-cross” the superscripts. These numbers now 70 Unit 5 – Nomenclature become subscripts. Omit all positive and negative signs and omit all 1’s. Answer = Al2S3 Write the formulas for the following binary ionic compounds. a) lithium selenide __________________ b) tin (II) oxide __________________ c) tin (IV) oxide __________________ d) magnesium fluoride ________________ e) copper (II) sulfide __________________ f) iron (II) phosphide _________________ g) gallium nitride __________________ h) iron (III) sulfide __________________ TERNARY IONIC COMPOUNDS Ternary ionic compounds are composed of at least _________ elements. Name the metal ion, using a Roman numeral in parenthesis if necessary. Follow this name with the name of the polyatomic ion. Polyatomic ions are groups of atoms that stay together and have a __________________. Examples are provided on page 7 of the NCDPI Reference Tables for Chemistry. There is one polyatomic ion with a positive oxidation number (NH4+) that may come first in a compound. Name the ion. Follow this name with the name of the anion or second polyatomic ion. Certain polyatomic ions, called ________________________, contain oxygen and another element. Name the following ternary ionic compounds. a) LiCN __________________ b) Fe(OH)3 ___________________ c) (NH4)2CO3 __________________ d) NiPO4 __________________ e) NaNO3 __________________ f) CaSO4 __________________ g) (NH4)2O __________________ h) CuSO3 __________________ Write the symbol for the metal or ammonium ion. Write the oxidation number as a superscript to the right of the metal’s/ammonium ion’s symbol. To the right of the metal’s symbol, write the symbol for the nonmetal or polyatomic ion. Write the oxidation number as a superscript to the right of the nonmetal’s/polyatomic ion’s symbol. Example: potassium nitrate - K1+ NO31- If the two oxidation numbers add together to get zero, the formula is a one-to-one ratio of the elements. Answer = KNO3 Example: aluminum hydrogen sulfate – Al3+ HSO4 1- If the two oxidation numbers DO NOT add together to get zero, you will need to “criss-cross” the superscripts. These numbers now 71 Unit 5 – Nomenclature become subscripts. Parentheses are to be placed around polyatomic ions before crisscrossing. Omit all positive and negative signs and omit all 1’s. Answer = Al(HSO4)3 Write the formulas for the following ternary ionic compounds. a) ammonium chloride __________________ b) ammonium sulfide _________________ c) barium nitrate __________________ d) zinc iodate __________________ e) sodium hypochlorite __________________ f) chromium (III) acetate ______________ g) iron (II) dichromate __________________ h) mercury (I) bromate ________________ MOLECULAR COMPOUNDS Molecular compounds are made of molecules. They are made by joining _______________________ atoms together into molecules. A molecular compound’s name tells you the number of atoms through the use of ____________________________. 1 mono- 4 tetra- 7 hepta- 2 di- 5 penta- 8 octa- 3 tri- 6 hexa- 9 nona- The name will consist of two words. Prefix name prefix name –ide 10 deca- One exception is we don’t write mono- if there is only one of the first element. The following double vowels cannot be used when writing names: (oa) and (oo). Example: NO2 There is one nitrogen. Mononitrogen But, you cannot use mono- on the first element, so drop the prefix. There are two oxygens. dioxygen You need the suffix – ide. dioxide Answer: nitrogen dioxide. Example: N2O There are two nitrogens. Dinitrogen There is one oxygen. monooxygen You cannot run (oo) together, so monoxygen. You need the suffix –ide. monoxide Answer: dinitrogen monoxide. Name the following molecular compounds. a) Cl2O7 ____________________________ b) CBr4 ____________________________ c) CO2 ________________________ d) BCl3 ___________________________ 72 Unit 5 – Nomenclature When writing a formula of a molecular compound from the name, you will not need to crisscross oxidation numbers. Molecular compounds name tells you the number of atoms through the use of prefixes. Example: diphosphorus pentoxide The name implies there are 2 phosphorous atoms and 5 oxygens. Answer: P2O5 Example: sulfur hexafluoride The name implies there is 1 sulfur atom and 6 fluorines. Answer: SF6 Write the formulas for the following molecules. a) tetraiodide nonoxide __________________ b) nitrogen trioxide __________________ c) carbon tetrahydride __________________ d) phosphorus trifluoride ______________ IONIC MOLECULAR Smallest Piece Molecule Types of metal and nonmetal Elements State of Matter solid Melting Point Low <300°C ACIDS Acids are compounds that give off hydrogen ions (H+) when dissolved in water. Acids will always contain one or more hydrogen ions next to an anion. The anion determines the name of the acid. Binary Acids Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, put the prefix hydro- and change -ide to -ic acid. Example: HCl The acid contains the hydrogen ion and chloride ion. Begin with the prefix hydro-, name the nonmetallic ion and change -ide to -ic acid. Answer: hydrochloric acid 73 Unit 5 – Nomenclature Example: H2S The next step is change -ide to -ic acid, but for sulfur the “ur” is added before -ic. Answer: hydrosulfuric acid Name the following binary acids. a) HF ____________________________ b) H3P ____________________________ The prefix hydro- lets you know the acid is binary. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the nonmetal. Example: hydrobromic acid The acid contains the hydrogen ion and the bromide ion. H1+ Br1- The two oxidation numbers add together to get zero. Answer: HBr Example: hydrotelluric acid The acid contains the hydrogen ion and the telluride ion. H1+ Te2- The two oxidation numbers do NOT add together to get zero, so you must criss-cross. Answer: H2Te Write the formulas for the following binary acids. a) hydrocyanic acid __________________ b) hydroselenic acid __________________ Ternary Acids The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to -ic acid. Change the suffix -ite to -ous acid. The hydro- prefix is NOT used! Example: HNO2 The acid contains the hydrogen ion and nitrite ion. Name the polyatomic ion and change -ite to -ous acid. Answer: nitrous acid Example: H3PO4 The acid contains the hydrogen ion and phosphate ion. Name the polyatomic ion and change -ate to -ic acid. Answer: phosphoric acid Name the following ternary acids. a) H2CO3 ____________________________ b) H2SO4 __________________________ c) H2CrO4 ________________________ d) HClO2 __________________________ The lack of the prefix hydro- from the name implies the acid is ternary, made of the hydrogen ion and a polyatomic ion. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the polyatomic ion. 74 Unit 5 – Nomenclature Example: acetic acid The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of H+ and the acetate ion. H1+ C2H3O21- The two charges when added equal zero. Answer: HC2H3O2 Example: sulfurous acid Again the lack of the prefix hydro- implies the acid is ternary, made of the hydrogen ion and a polyatomic ion. The polyatomic ion must end in –ite since the acid ends in -ous. The acid is made of H+ and the sulfite ion. H1+ SO32- The two charges when added do not equal zero, so you must crisscross the oxidation numbers. Ignore the negative sign and ones are understood. Answer: H2SO3 Write the formulas for the following binary acids. a) perchloric acid __________________ b) iodic acid __________________ c) dichromic acid __________________ d) hypochlorous acid ________________ Homework / Practice Name each of the following compounds 1. CuS 18. K2SO4 35. N2O3 2. FeI2 19. NH4NO3 36. Mg3P2 3. Cu2SO4 20. FePO4 37. CCl4 4. CuCl2 21. SbBr3 38. H2O 5. Ni(C2H3O2)2 22. Ag 39. NH4Cl 6. MnO 23. 7. CoO 24. ClO2 41. TiBr3 8. Mn2O3 25. KCl 42. Cu3P 9. Co2S3 26. Fe 43. SnSe2 10. AuF 27. PbO 44. GaAs 11. CrBr2 28. FeCl2 45. Pb(SO4)2 12. AlPO4 29. NI3 46. Be(HCO3)2 13. KNO2 30. CO2 47. Mn2(SO3)3 14. CaCO3 31. Ni 48.Al(CN)3 15. Mg(OH)2 32. Au 49. P4O6 16. Na2CrO4 33. LiF 50. N2O3 17. Ba(CN)2 34. Al2O3 51. N2O MgBr2 75 40. Fe(NO3)3 Unit 5 – Nomenclature 52. BrO3 57. PF5 62. As4O6 53. SiF4 58. P2O5 63. N2O4 54. P4S10 59. Cl2O 64. NI3 55. Cl2O3 60. CCl4 65. ClF3 56. PCl3 61. N2O5 66. I2O5 Write the formula for each of the following compounds 67. iron (II) chloride 90. calcium sulfide 68. copper (I) sulfide 91. silicon diselenide 69. mercury (I) bromide 92. diphosporus trioxide 70. chromium (III) oxide 93. iron (III) chloride 71. gold (I) iodide 94. ,magnesium fluoride 72. manganese (II) nitride 95. zinc oxide 73. cobalt (III) phosphide 96. arsenic tribromide 74. iron (III) chloride 97. carbon dioxide 75. copper (II) sulfide 98. gold (III) chlorite 76. manganese (III) chloride 99. silver dichromate 77. potassium perchlorate 100. barium Iodide 78. aluminum sulfate 101. carbon tetrafluoride 79. iron (II) carbonate 102. sodium sulfide 80. barium iodate 103. chromium (VI) phosphate 81. magnesium phosphate 104. vanadium (IV) carbonate 82. silver phosphate 105. tin (II) nitrite 83. cobalt (III) nitrite 106. cobalt (III) oxide 84. ammonium sulfite 107. titanium (II) acetate 85. dihydrogen oxide 108. vanadium (V) sulfide 86. lithium Iodide 109. chromium (III) hydroxide 87. silver chloride 110. lithium iodide 88. carbon monoxide 111. lead (II) nitride 89. selenium difluoride 112. silver bromide 76 Unit 5 – Nomenclature Write the formula for the compound formed between the two ions listed carbonate chromate chlorate sulfide chloride sulfite sodium Na2CO3 Na2CrO4 · · · · silver · · · · · · ammonium · · · · · · Tin (II) · · · · · · magnesium · · · · · · lead (II) · · · · · · aluminum · · · · · · manganese (III) · · · · · · potassium · · · · · · barium · · · · · · APPLICATIONS OF THE MOLE Molar Mass The molar mass of a compound is the mass of a mole of the _________________________________ particles of the compound. Because each representative particle is composed of two or more atoms, the molar mass of the compound is found by adding the molar masses of all of the ________________ in the representative particle. To determine the molar mass of an element, find the element’s symbol on the periodic table and round the mass so there is __________ digit beyond the decimal. For example, the molar mass of carbon (C) is ____________ g/mol, of chlorine (Cl) is ___________ g/mol and of iron (Fe) is _____________ g/mol. In the case of NH3, the molar mass equals the mass of one mole of nitrogen atoms plus the mass of ___________ moles of hydrogen atoms. Molar mass of NH3 = molar mass of N + 3 (molar mass of H) Molar mass of NH3 = 14.0 + 3 (1.0) = 17.0 g/mol Mole Conversions You can use the molar mass of a compound to convert between mass and moles, just as you used the molar mass of elements to make these conversions. How many moles of magnesium in 56.3 g of Mg? 77 Unit 5 – Nomenclature How many moles are in 146 grams of NH3? How many moles are in 295 grams of Cr(OH)3? How many moles are in 22.5 grams of HCl? How many grams of sodium chloride in 3.45 moles of NaCl? How many grams are in 0.120 moles of AlF3? How many grams are in 13.0 moles of H2SO4? How many grams are in 1.6 moles of K2CrO4? Percent Composition Recall that every chemical compound has a definite composition - a composition that is always the same wherever that compound is found. The composition of a compound is usually stated as the percent by mass of each element in the compound. The percent of an element (X) in a compound can be found in the following way. %X molarmassX # X ' s MolarMassC ompound Determine the percent composition of chlorine in calcium chloride (CaCl2). First, analyze the information available from the formula. A mole of calcium chloride consists of one mole of calcium ions and ___________ moles of chloride ions. Next, gather molar mass information from the atomic masses on the periodic table. To the mass of one mole of CaCl2, a mole of calcium ions contributes ______________ g, and two moles of chloride ions contribute 2 x 35.5 g = 71.0 g for a total molar mass of 111.1 g/mol for CaCl2. Finally, use the data to set up a calculation to determine the percent by mass of an element in the compound. Determine the percent composition of carbon in sodium acetate (NaC2H3O2). Calculate the percent composition aluminum of aluminum oxide (Al2O3). Determine the percent composition of oxygen in magnesium nitrate, which has the formula Mg(NO3)2. Determine the percent composition of sulfur in aluminum sulfate, which has the formula Al2(SO4)3. 78 Unit 5 – Nomenclature Determine the percent composition of oxygen in zinc nitrite, which has the formula Zn(NO2)2. Percent Water in a Hydrate Hydrates are compounds that incorporate ________________________ molecules into their fundamental solid structure. In a hydrate (which usually has a specific crystalline form), a defined number of water molecules are associated with each formula unit of the primary material. Gypsum is a hydrate with __________ water molecules present for every formula unit of CaSO4. The chemical formula for gypsum is CaSO4 • 2 H2O and the chemical name is calcium sulfate __________________. Note that the dot in the formula (or multiplication sign) indicates that the waters are there. Other examples of hydrates are: lithium perchlorate trihydrate - LiClO4 • 3 H2O; magnesium carbonate pentahydrate - MgCO3 • 5 H2O; and copper (II) sulfate pentahydrate - CuSO4 • 5 H2O. The water in the hydrate (referred to as "water of hydration") can be removed by heating the hydrate. When all hydrating water is removed, the material is said to be __________________________ and is referred to as a(n) ___________________________. Experimentally measuring the ________________________ water in a hydrate involves first heating a known mass of the hydrate to remove the waters of hydration and then measuring the mass of the anhydrate remaining. The difference between the two masses is the mass of water _______________. Dividing the mass of the water lost by the original mass of hydrate used is equal to the fraction of water in the compound. Multiplying this fraction by ___________ gives the percent water. Determine the percent water in CuSO4 • 5 H2O (s). Determine the percent water in MgCO3 •5 H2O (s). Determine the percent water in LiClO4 • 3 H2O (s). Empirical Formula You can use percent composition data to help identify an unknown compound by determining its empirical formula. The empirical formula is the ________________________ whole-number ratio of atoms of elements in the compound. In many cases, the empirical formula is the actual 79 Unit 5 – Nomenclature formula for the compound. For example, the simplest ratio of atoms of sodium to atoms of chlorine in sodium chloride is 1 atom Na : 1 atom Cl. So, the empirical formula of sodium chloride is Na1Cl1, or NaCl, which is the true formula for the compound. The formula for glucose is C6H12O6. The coefficients in glucose are all divisible by 6. The empirical formula of glucose is CH2O. Determine the empirical formula for Tl2C4H4O6. Determine the empirical formula for N2O4. The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O. Determine the compound’s empirical formula. Because percent means “parts per hundred parts,” assume that you have ___________ g of the compound. Then calculate the number of moles of each element in the 100 g of compound. To obtain the simplest wholenumber ratio of moles, _________________ each number of moles by the smallest number of moles. Find the whole number mole ratio for the compound. These numbers become the____________________________ in the empirical formula. Determine the empirical formula of the following compound: 31.9 g Mg, 27.1 g P The composition of an unknown acid is 40.00% carbon, 6.71% hydrogen, and 53.29% oxygen. Calculate the empirical formula for the acid. The composition of an unknown ionic compound is 60.7% nickel and 39.3% fluorine. Calculate the empirical formula for the ionic compound. The composition of a compound is 6.27 g calcium and 1.46 g nitrogen. Calculate the empirical formula for the compound. Find the empirical formula for a compound consisting of 63.0% Mn and 37.0% O. Molecular Formula For many compounds, the empirical formula is not the true formula. A molecular formula tells the ___________________ number of atoms of each element in a molecule or formula unit of a compound. The molecular formula for a compound is either the same as the empirical formula or a whole-number _______________________ of the empirical formula. In order to determine the molecular formula for an unknown compound, you must know the molar mass of the compound in addition to its empirical formula. Then you can compare the molar mass of the compound with the molar mass represented by the empirical formula. 80 Unit 5 – Nomenclature The molecular mass of benzene is 78 g/mol and its empirical formula is CH. What is the molecular formula for benzene? HINT: Calculate the molar mass represented by the formula CH. Calculate the whole number multiple, n, and apply it to its empirical formula. The simplest formula for butane is C2H5 and its molecular mass is about 60.0 g/mol. What is the molecular formula of butane? What is its molecular formula of cyanuric chloride, if the empirical formula is CClN and the molecular mass is 184.5 g/mol? The simplest formula for vitamin C is C3H4O3. Experimental data indicates that the molecular mass of vitamin C is about 180. What is the molecular formula of vitamin C? Maleic acid is a compound that is widely used in the plastics and textiles industries. The composition of maleic acid is 41.39% carbon, 3.47% hydrogen, and 55.14% oxygen. HINT: Start by determining the empirical formula for the compound. The composition of silver oxalate is 71.02% silver, 7.91% carbon, and 21.07% oxygen. If the molar mass of silver oxalate is 303.8 g/mol, what is its molecular formula? The composition of silver oxalate is 71.02% silver, 7.91% carbon, and 21.07% oxygen. If the molar mass of silver oxalate is 303.8 g/mol, what is its molecular formula? Homework / Practice 1) A compound is found to have (by mass) 48.38% carbon, 8.12% hydrogen and the rest oxygen. What is its empirical formula? 2) A compound is found to have 46.67% nitrogen, 6.70% hydrogen, 19.98% carbon and 26.65% oxygen. What is its empirical formula? 3) A compound is known to have an empirical formula of CH and a molar mass of 78.11 g/mol. What is its molecular formula? 4) Another compound, also with an empirical formula if CH is found to have a molar mass of 26.04 g/mol. What is its molecular formula? 5) A compound is found to have 1.121 g nitrogen, 0.161 g hydrogen, 0.480 g carbon and 0.640 g oxygen. What is its empirical formula? (Note that masses are given, NOT percentages.) 6) A compound containing only carbon, hydrogen and oxygen is found to be 48.38% carbon and 8.12% hydrogen by mass. What is its empirical formula? 7) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula? 8) Calculate the empirical formula of the compound that contains 1.0 g S for each 1.5 g O. 9) Calculate the empirical formula of the compound containing 75.0% C and 25.0% H. 10) Calculate the empirical formula of the compound containing 81.8% C and 18.2% H. 81 Unit 5 – Nomenclature 11) The active ingredient in chocolate is theobromine; a sample was analyzed and determined to be composed of: 147.0 g C 14.0 g H 56.0 g O 98.0 g N a. Determine the % composition for each element. b. Determine the empirical formula for theobromine. c. The molecular weight of theobromine is known to be 180.0 g/mole. What is the molecular formula? 12) Determine the empirical formula of the compound containing 37.5% C, 12.5% H, and 50.0% O by weight. 13) Determine the empirical formula of the compound containing 26.1% C, 4.3% H, and 69.6% O by weight. 14) Determine the empirical formula of the compound containing 38.7% C, 16.1% H, and 45.2% N by weight. 15) What is the empirical formula of a compound if a 50.0 g sample of it contains 9.1 g Na, 20.6 g Cr, and 22.2 g O? 16) A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H. What is the empirical formula of the compound? 17) NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet and find the molecular formula. (The molar mass of NutraSweet is 294.30 g/mol) 18) A compound consists of 85% silver and 15% florine by mass. What is the empirical formula? 19) A compound consists of 40% calcium, 12% carbon, and 48% oxygen by mass. What is the empirical formula by mass? 20) A compound consists of 75% Magnesium, and 25% oxygen by mass. What is the empirical formula by mass? 21) A compound contains 50% Magnesium, 24% carbon, 16% oxygen, and 10% hydrogen. What is the empirical formula? 22) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula? 23) Benzoic acid contains 68.8% Carbon, 4.95% Hydrogen and 26.2% Oxygen. Find the empirical formula? 24) Freons are gaseous compounds used in refrigeration. A particular freon contains 9.93% carbon, 56% chlorine, and 31.4% fluorine. What is the empirical formula? Naming and Formula Math Practice Test Directions: Write the formula for the compound 1. 2. 3. 4. 5. sodium phosphide iron (II) perchlorate vanadium (V) nitrite nickel (I) oxide magnesium hydroxide 6. 7. 8. 9. 82 cesium nitride nitrogen trichloride hydroxic acid carbon tetrahydride Unit 5 – Nomenclature Directions: Name the compound 10. KCl 11. FeSO4 12. Li2O 13. Cr2S3 14. Ca3N2 15. Fe2S3 16. CuI2 17. PBr3 18. CO2 19. HNO3 20. What is the percent nitrogen in potassium nitrate? 21. What is the ratio of barium ions to Nitrogen ions in a formula unit of barium nitrate? 22. A compound is found to have 46.67% nitrogen, 6.70% hydrogen, 19.98% carbon and 26.65% oxygen. What is its empirical formula? 23. A compound is found to have (by mass) 48.38% carbon, 8.12% hydrogen and the rest oxygen. What is its empirical formula? 24. Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula? 25. Determine the empirical formula of N3O6 26. The empirical formula of a compound is CH2. Its molecular mass is 70g/mol. What is its molecular formula? (Show all work.) 27. Determine the empirical formula of a compound that is 82% carbon, 18% hydrogen. 28. Determine the molecular formula of a compound that is composed of 40.0% carbon, 6.7% hydrogen and 53.5% oxygen. The molecular mass is 120.0g/mol. Multiple Choice Practice 29. Which of the following is a binary compound? a. hydrogen sulfide b. hydrogen sulfate c. ammonium sulfide d. ammonium sulfate 30. What is the formula for sodium oxalate? a. NaClO b. Na2ClO c. Na2C2O4 d. NaC2H3O2 31. Which of the following is a binary compound? a. potassium chloride b. ammonium chloride c. potassium chlorate d. ammonium chlorate 83 Unit 5 – Nomenclature 32. Which is the correct formula for nitrogen monoxide? a. NO b. N2O c. NO2 d. N2O3 33. Which of the following represents the correct formula for aluminum oxide? a. AlO c. AlO2 b. Al2O3 d. Al2O 34. Which of the following is the correct name for NaHCO3? a. sodium hydrogen carbonate b. sodium acetate c. nitrogen hydrogen carbonate d. sodium hydrogen carbon trioxide 35. In which of the following compounds does ionic bonding predominate? a. NH4Cl c. CH4 b. CO2 d. LiBr 36. Which of the following molecules contains only one non-bonding pair of valence electrons? a. NH4+ c. C2H4 b. HCN d. N2 37. What is the name of CaCl2? a. calcium dichloride b. calcium (II) chloride c. monocalcium dichloride d. calcium chloride 38. What is the name of Mg(NO3)2 a. Magnesium nitrate b. Magnesium (II) nitrate c. Magnesium dinitrate d. Magnesium nitrogen oxide 39. What is the name of P2O5? a. phosphorus oxide b. phosphorus pentaoxide c. diphosphorus pentaoxide d. phosphorus (III) oxide 40. What is the formula for sulfur hexachloride? a. S5Cl b. SHCl c. SCl5 d. SCl6 41. What is the name of the formula Fe(NO3)2? a. iron nitrate b. iron (II) nitrate c. iron dinitrate d. iron (III) nitrate 42. What is the formula for the compound nitrogen (II) oxide. a. N2O3 c. NO2 b. NO d. N2O 84 Unit 5 – Nomenclature 43. Which of the following is not a type of chemical formula? a. Empirical c. Structural b. Molecular d. Parabola 44. What is the approximate percentage oxygen in the formula mass of Ca(NO3)2? a. 28 c. 58 e. 164 b. 42 d. 96 45. Which formulas could represent the empirical formula and the molecular formula of a given compound? a. CH2O and C4H6O4 d. CH2 and C3H6 b. CHO and C6H12O6 e. CO and CO2 c. CH4 and C5H12 46. When combining with nonmetallic atoms, metallic atoms generally will a. lose electrons and form negative ions b. lose electrons and form positive ions c. gain electrons and from negative ions d. gain electrons and form positive ions 47. What is the empirical formula of the compound whose molecular formula is P4O10? a. PO c. P2O5 b. PO2 d. P8O20 48. What is the percent by mass of oxygen in magnesium oxide, MgO? a. 20% c. 50% b. 40% d. 60% 49. A compound is 86% carbon and 14% hydrogen by mass. What is the empirical formula for this compound? a. CH c. CH3 b. CH2 d. CH4 50. What is the percent by mass of water in the hydrate Na2CO3 * 10H2O (formula mass = 286)? a. 6.89% c. 26.1% b. 14.5% d. 62.9% 51. What is the gram formula mass of (NH4)3PO4? a. 113 g b. 121 g 85 c. 149 g d. 404 g Unit 5 – Nomenclature 52. What is the empirical formula of a compound that contains 85% Ag and 15% F by mass? a. AgF c. AgF2 b. Ag2F d. Ag2F2 53. What is the percent water in hydrated calcium chloride... CaCl2·2H2O? a. 66.67% c. 24.51% b. 32.47% d. 12.26% 54. What is the empirical formula for a compound that contains 17.34% hydrogen and 82.66% carbon? a. C5H c. CH3 b. C2H5 d. CH2 55. What is the molecular formula for a compound that is 46.16% carbon, 5.16% hydrogen and 48.68% fluorine if the molar mass of this compound is 156.12 g? a. C3H4F2 c. C6H8F4 b. C5H10F5 d. C6H6F3 56. One mole of (NH4)2HPO4 contains _?_ moles of hydrogen atoms. a. 1 b. 5 c. 6 d. 9 57. Which of the following is named incorrectly? a. H2CO2 : carbonous acid b. HClO2 : chlorous acid c. H2SO4 : sulfuric acid d. HClO : hydrochlorous acid e. H3PO3 : phosphorous acid 58. The name for the compound NaHSO4 is: a. sodium hydrogen sulfate b. sodium persulfate c. sodium bisulfate d. two of these are correct e. none of these is correct 59. A sample of an alcohol is tested and found to contain 52% carbon, 35% oxygen, and 13% hydrogen by mass. Tests indicate that the molecular weight of the molecule is between 30 and 80. What is the molecular formula of the alcohol? a. C2H5OH b. C3H7OH c. C5H11OH d. C4H9OH e. CH3OH 86 Unit 5 – Nomenclature 60. A 5.15 gram sample of a hydrocarbon is burned in oxygen, producing 15.6 grams of carbon dioxide and 8.45 grams of water. Assuming an excess of oxygen, what is the empirical formula of the hydrocarbon? a. CH4 b. C7H16 c. C9H20 d. C5H12 e. C3H8 61. For which of the following compounds does 0.400 mol have a mass of 12.8 grams? a. CH3OH b. CH4 c. CCl4 d. CO2 e. C4H10 62. What is the molar mass of glucose, C6H12O6? a. 18 g b. 220.17 g c. 12.01 g d. 180.18 g e. 160.00 g 63. Methane, CH4, and ethane, C2H6 are both hydrocarbons that exist as gases at room conditions. How many grams of ethane contain the same number of molecules as 6.00 g of methane? a. 0.374 grams b. 0.09 grams c. 2896 grams d. 11.2 grams e. 80.5 grams 87 Unit 6- Chemical Reactions CHEMICAL REACTIONS All chemical reactions have two parts: (1) A substance that undergoes a reaction is called a __________________________. In other words, reactants are the substances you start with. (2) When reactants undergo a chemical change, each new substance formed is called a ___________________________. In other words, the products are the substances you end up with. The reactants turn into the products. Reactants → Products. In a chemical reaction, the way atoms are joined is changed. Atoms aren’t __________________________ or destroyed. Chemical reactions can be described several ways. In a sentence: Copper reacts with chlorine to form copper (II) chloride. In a word equation: Copper + chlorine → copper (II) chloride The arrow separates the reactants from the products. The arrow reads “reacts to ________________.” The plus sign reads “_____________.” (s) after the formula implies the substance is a ___________________. (g) after the formula implies the substance is a gas. (l) after the formula implies the substance is a ______________________. (aq) after the formula implies the substance is aqueous, a solid dissolved in _____________________. __________ used after a product indicates a gas, same as (g). ↓ used after a product indicates a ________________, same as (s). _____________ indicates a reversible reaction. ________________ or ________________ shows that heat is supplied to the reaction. ___________________ is used to indicate a catalyst used supplied, in this case, platinum. A catalyst is a substance that ____________________ ____________ a reaction without being changed by the reaction. Enzymes are biological or ______________________ catalysts. There are seven elements that never want to be alone. They form ________________________ molecules.H2 , N2 , O2 , F2 , __________ , Br2 , I2. (1 + 7 pattern on the periodic table) The following are indications that a chemical reaction has occurred: formation of a ____________________________, evolution of a gas, _____________________ change, and absorption or release of ________________. A ________________________ formula uses formulas and symbols to describe a reaction. All chemical equations are sentences that describe reactions. 88 Unit 6- Chemical Reactions Convert the following sentences to chemical equations. a) Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form solid iron (II) chloride and hydrogen sulfide gas. b) Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water. Convert the following chemical equations to sentences. a) Fe (s) + O2 (g) → Fe2O3 (s) b) Cu (s) + AgNO3 (aq) → Ag (s) + Cu(NO3)2 (aq) Balancing Equations Atoms can’t be ______________________ or destroyed. All the atoms we start with we must end up with. A balanced equation has the same number of each element on both _________________ of the equation. Example: C + O2 → CO This equation is NOT balanced. There is one carbon atom on the left and ________ on the right. There are two oxygen atoms on the left and only one on the right. We need one more oxygen atom in the products. We can’t change the formula, because it describes what it is. In order to have two oxygen atoms, another CO must be produced. But where did the other carbon come from? We must have started with two carbon atoms. The balanced chemical equation is 2 C + O2 → 2 CO Rules for Balancing Write the correct formulas for all the reactants and products. Count the number of atoms of each type appearing on both sides. Balance the elements one at a time by adding coefficients (the numbers in front). Check to make sure it is balanced. Never change a _________________________ to balance an equation. If you change the formula you are describing a different reaction. Never put a coefficient in the middle of a formula. 2 NaCl is okay; Na2Cl is not. Balance the following reaction. H2 + O2 → H2O Balance elements in the following order: (1) metals; (2) nonmetals; (3) hydrogen; and (4) oxygen 89 Unit 6- Chemical Reactions If an atom appears more than once on a side, balance it last. If you fix everything except one element, and it is even on one side and odd on the other, double the first number, then move on from there. Balance the following equations. 1) _____ CH4 + _____ O2 → _____ CO2 + _____ H2O 2) _____ AgNO3 + _____ Cu → _____ Cu(NO3)2 + ______ Ag 3) _____ Mg + _____ N2 → _____ Mg3N2 4) _____ P + _____ O2 → ______ P4O10 5) _____ Na + _____ H2O → _____ H2 + _____ NaOH 6) _____ Pb(NO3)2 + _____ K2CrO4 ______ PbCrO4 + ______ KNO3 7) _____ MnO2 + _____ HCl _____ MnCl2 + ______ H2O + _____ Cl2 8) _____ Ba(CN)2 + _____ H2SO4 _____ BaSO4 + _____ HCN 9) _____ Zn(OH)2 + _____ H3PO4 _____ Zn3(PO4)2 + _____ H2O TYPES OF REACTIONS Reactions fall into 5 categories. We will recognize the type by the reactants. We will be able to predict the products. For some we will be able to predict whether they will happen at all. Synthesis Reactions Synthesize means to put together. Whenever two or more substances combine to form one single product, the reaction is called a synthesis reaction. Examples: Ca + O2 → CaO and P2O5 + 3 H2O → 2 H3PO4 We can predict the products if they are two elements. All you need to do is combine the elements, metals first, and criss-cross oxidation numbers if necessary. After predicting the product, the reaction must be balanced. Mg + N2 → CaO + H2O → On page 6 of the Chemistry Reference Packet, this reaction is an example of “Metal oxide water reactions.” The product listed in the packet is “base.” A base is a metallic hydroxide. SO2 + H2O → 90 Unit 6- Chemical Reactions On page 6 of the Chemistry Reference Packet, this reaction is an example of “Nonmetal oxide water reactions.” The product listed in the packet is “________________.” The acid is a ternary acid. Ternary acids start with _________ and end in O. The other element goes in the center. This is the only compound for which you can add the number of elements and use these numbers as subscripts. Write and balance the following synthesis reactions. a) Ca + Cl2 → b) Fe + O2 → HINT: Use iron (II). c) K2O + H2O → d) Al + O2 → e) SO3 + H2O → f) N2O5 + H2O → Decompositions Reactions The word decompose implies the compound will “fall apart.” In a decomposition reaction, one compound breaks down into _____________ or more simple substances. NaCl → Na + Cl2 CaCO3 → CaO + CO2 We can easily predict the products if it is a binary compound. A binary compound is made up of only two elements. The compound merely falls apart into its elements. H2O → HgO → If the compound has more than two elements, you must consult the Reference Tables, page 6. NiCO3 → NiCO3 is called nickel (II) carbonate. The packet states that a metallic carbonate decomposes to form a MO (metallic oxide) and CO2. The metallic oxide is nickel (II) oxide. Use the Chemistry Reference Tables to write and balance the following decomposition reactions. a) KClO3 → b) CaBr2 → c) Li2CO3 → 91 Unit 6- Chemical Reactions d) Cr(OH)2 → e) NaHCO3 → f) HNO3 → (Dinitrogen pentoxide one of the products.) Single Replacement In a single-displacement reaction, one element takes the place of another in a compound. One reactant must be an element, and the one reactant must be a _______________________. The products will be a different element and a different compound. F2 + LiCl → LiF + Cl2 Remember zinc, Zn, always forms a ___________ ion doesn’t need parenthesis. ZnCl2 is zinc chloride. In addition, silver, Ag, always forms a ___________ ion. AgCl is silver chloride. Some single replacement reactions do not occur because some elements are not as ________________ as others. A more active element _________________________ a less active element. There is a list referred to as the Activity Series on page 7 of your Chemistry Reference Packet. A higher element on the list replaces lower element. If the element by itself is lower on the list, the reaction will ___________ occur. Metals replace metals (and hydrogen) K + NaCl → Potassium wants to replace ________________________. You must check the activity series on page 7 of your Chemistry Reference Packet to see if this is possible. Because K is higher, potassium can replace sodium. The potassium will bond with the _____________________ and the sodium will be alone. You must always check to see if the compound formed needs crisscrossing. Check for balancing. Sn + FeCl3 → Because Sn is NOT higher, tin cannot replace iron. No reaction occurs. Write and balance the following single replacement reaction. a) Rb + AlN → c) Ag + CoBr2 → b) Zn + HCl → Metals replace hydrogen Na + H2O (cold) → 92 Unit 6- Chemical Reactions Think of water as HOH. Metals high enough on the activity series replace the first ______ and combine with the OH1- (hydroxide) according to page 6 of the Reference Tables. Is sodium above hydrogen and higher than the line marked “Replace hydrogen from cold water” on the activity series? Since the answer is yes, sodium replaces the first H, bonding with hydroxide. Mg + HCl → Metals higher on the activity series replace the H and combine with the nonmetal according to page 6 of the Reference Tables. Hydrogen gas is a second product. Is magnesium above hydrogen on the activity series? Write and balance the following single replacement reactions. a) Ag + H2O (steam) → c) Cr + H3PO4 → (HINT: Use Cr3+ ) b) Cu + H2SO4 → d) Ca + H2O (steam) → Nonmetals can replace other ________________________. This is limited to F2 , Cl2 , Br2 and I2 The order of activity is listed in the Chemistry Reference Packet, page 7. Higher replaces _____________. F2 + HCl → Is fluorine above chlorine in the activity series of halogens? Since the answer is yes, fluorine replaces the chlorine, bonding with hydrogen. Write and balance the following single replacement reactions. a) Br2 + KCl → b) Cl2 + KI → Double Replacement In double-displacement reactions, the positive portions of two ___________________ compounds are interchanged. The reactants must be two ionic compounds or ______________. Double replacement reactions usually take place in ________________________ solution. NaOH + FeCl3 → The positive ions change place. You must check to see if you need to criss-cross the products. Now balance. A double replacement reaction will only happen if one of the products: (1) doesn’t dissolve in water and forms a __________________, (2) is a _____________ that bubbles out, or (3) is a _________________________ compound usually water. 3NaOH + FeCl3 → Fe(OH)3 + 3NaCl 93 Unit 6- Chemical Reactions None of the products are familiar gases. Both products are ionic (not covalent) because they start with metals. We must consult the Solubility Rules on page 6 of the Chemistry Reference Tables to determine if a solid (a ________________________) is formed. The “Soluble” side of the Solubility Rules states that Group 1 (IA) salts are soluble; therefore, NaCl is soluble and is NOT the precipitate. The “Insoluble” side of the Solubility Rules states that all hydroxides except Group 1, Sr, Ba and NH41+ are INSOLUBLE. Therefore, Fe(OH)3 is the precipitate (solid). In molecular equations, the formulas of the compounds are written as though all species existed as molecules or whole units. An ionic equation shows dissolved ionic compounds in terms of their free ions. Ions that are not involved in the overall reaction are called spectator ions. The net ionic equation indicates only the species that actually take part in the reaction. The following steps are useful for writing ionic and net ionic equations: 1) Write a balanced molecular equation for the reaction. 2) Rewrite the equation to indicate which substances are in ionic form in solution. Remember that all soluble salts (and other strong electrolytes), are completely dissociated into cations and anions. This procedure gives us the ionic equation. 3) Lastly, identify and cancel spectator ions on both sides of the equation to arrive at the net ionic equation. Example: sodium hydroxide + iron (III) chloride yields iron (III) hydroxide + sodium chloride Balanced Molecular Equation: 3 NaOH + FeCl3 Fe(OH)3 + 3 NaCl Complete Ionic Equation: 3Na1+ + 3OH1- + Fe3+ + 3Cl1- Fe(OH)3 + 3Na1+ + 3Cl1Net Ionic Equation: 3OH1- + Fe3+ Fe(OH)3 Write and balance the following double replacement reaction. Assume the reaction takes place. In addition, identify the precipitate and write the net ionic equation. a) CaCl2 + NaOH → c) KOH + Fe(NO3)3 → b) CuCl2 + K2S → d) (NH4)2SO4 + BaF2 → 94 Unit 6- Chemical Reactions Combustion A combustion reaction is one in which a substance rapidly combines with ____________________ to form one or more oxides. Combustion reactions involve a compound composed of only C and H (and maybe O) that is reacted with oxygen gas. If the combustion is complete, the products will be CO2 and __________________. Combustion reactions produce heat, and are therefore considered exothermic reactions. Complete and balance the following combustion reactions. a) C4H10 + O2 → c) C8H8 + O2 → b) C6H12O6 + O2 → d) C3H8O3 + O2 → To determine which type a reaction is, look at the reactants. (E = element and C = compound) E+E Synthesis C Decomposition E+C Single replacement C+C Double replacement CH cpd + O2 Combustion Note: Two other common synthesis reactions include: nonmetallic oxide + water and metallic oxide + water. Identify whether the reaction is synthesis, decomposition, single replacement, double replacement or combustion. a) H2 + O2 → e) KBr + Cl2 → b) H2O → f) Zn + H2SO4 → c) Mg(OH)2 + H2SO3 → g) AgNO3 + NaCl → d) HgO → h) C6H6 + O2 → Homework / Practice Directions: balance the following equations. SYNTHESIS 1. S + O2 ---> SO2 2. S + O2 ---> SO3 3. P + O2 ---> P2O3 4. Mg + N2 ---> Mg3N2 5. N2 + O2 --> NO2 6. Na + O2 ---> Na2O 95 Unit 6- Chemical Reactions 7. Cu + S ---> Cu2S 8. Al + N2 ---> AlN 9. Hg + I2 ---> HgI2 10. Fe + O2 ---> Fe2O3 DECOMPOSITION 11. HgO ---> Hg + O2 12. MgSO4.7H2O ---> MgSO4 + H2O 13. KClO3 ---> KCl + O2 14. NH4NO3 ---> N2O + H2O 15. NaNO3 ---> NaNO2 + O2 16. BaO2 ---> BaO + O2 17. H2O2 ---> H2O + O2 18. NO2 ---> N2 + O2 19. CaCO3 ---> CaO + CO2 20. H2O ---> H2 + O2 SINGLE REPLACEMENT (SINGLE DISPLACEMENT) 21. AlI3 + Cl2 ---> AlCl3 + I2 26. ZnS + O2 ---> ZnO + SO2 22. CH4 + Cl2 ---> CHCl3 + HCl 27. Na + H2O ---> NaOH + H2 23. Al + CuSO4 ---> Al2(SO4)3 + Cu 28. Al + H2SO4 ---> Al2(SO4)3 + H2 24. Fe2O3 + Al ---> Al2O3 + Fe 29. Zn + NaOH ---> Na2ZnO2 + H2 25. Zn + HCl ---> ZnCl2 + H2 30. AgNO3 + Zn ---> Zn(NO3)2 + Ag DOUBLE REPLACEMENT 31. Fe(OH)3 + H2SO4 ---> Fe2(SO4)3 + H2O 32. AgNO3 + K2CrO4 ---> Ag2CrO4 + KNO3 33. AgNO3 + CuCl2 ---> AgCl + Cu(NO3)2 34. Pb(NO3)2 + HCl ---> PbCl2 + HNO3 35. MgCl2 + NaOH ---> Mg(OH)2 + NaCl 36. AgNO3 + H2S ---> Ag2S + HNO3 37. CaCO3 + HCl ---> CaCl2 + H2CO3 38. Hg2(NO3)2 + NaCl ---> Hg2Cl2 + NaNO3 39. BaCl2 + (NH4)2CO3 ---> BaCO3 + NH4Cl 40. Al(OH)3 + NaOH ---> NaAlO2 + H2O COMBUSTION 41. CH4 + O2 ---> CO2 + H2O 42. C4H10 + O2 ---> CO2 + H2O 43. C3H6 + O2 ---> CO2 + H2O 44. C5H8 + O2 ---> CO2 + H2O 45. CH3OH + O2 ---> CO2 + H2O 46. C6H12O6 + O2 ---> CO2 + H2O 96 Unit 6- Chemical Reactions Predict the products of the following reactions 47. MgCl2 + O2 ---> 48. Na + O2 ---> 49. P2O3 + H2O ---> 50. K2O + H2O ---> 51. BaO + CO2 ---> 52. BeO + CO2 ---> 53. Al2O3 + H2O ---> 54. N2O5 + H2O ---> 55. NaCl + O2 ---> 56. Ra + Cl2 ---> 57. Ni(ClO3)2 ---> 58. Ag2O ---> 59. HNO2 ---> 60. Fe(OH)3 ---> 61. ZnCO3 ---> 62. Cs2CO3 ---> 63. Al(OH)3 ---> 64. H2SO4 ---> 65. RbClO3 ---> 66. RaCl2 ---> 67. ZnS + O2 ---> 68. K + H2O ---> 69. Fe + HCl ---> 70. NaI + Br2 ---> 71. Al + Pb(NO3)2 ---> 72. Cl2 + NaI ---> 73. Fe + AgC2H3O2 ---> 74. Al + CuCl2 ---> 75. Br2 + CaI2 ---> 76. Al + HCl ---> 77. Mg + HCl ---> 78. Zn + H2SO4 ---> 79. Fe + CuSO4 ---> 80. Cl2 + MgI2 ---> 81. Ca(OH)2 + H3PO4 ---> 82. K2CO3 + BaCl2 ---> 83. Cd3(PO4)2 + (NH4)2S ---> 84. Co(OH)3 + HNO3 ---> 85. AgNO3 + KCl ---> 86. Na2CO3 + H2SO4 ---> 87. Al(OH)3 + HC2H3O2 ---> 88. Al2(SO4)3 + Ca3(PO4)2 ---> 89. Cr2(SO3)3 + H2SO4 ---> 90. AgC2H3O2 + K2CrO4 ---> Write the complete ionic equation, identify the spectator ions, write the net ionic equation 91. K2CO3 + BaCl2 ---> BaCO3 + KCl 92. Cd3(PO4)2 + (NH4)2S ---> CdS + (NH4)3PO4 93. AgNO3 + KCl ---> AgCl2 + K2NO3 94. Na2CO3 + H2SO4 ---> H2CO3 + Na2SO4 95. Al(OH)3 + HC2H3O2 ---> Al(C2H3O2)3 + HOH 96. Al2(SO4)3 + Ca3(PO4)2 ---> AlPO4 + CaSO4 97. Cr2(SO3)3 + H2SO4 ---> Cr2(SO4)3 + H2SO3 98. AgC2H3O2 + K2CrO4 ---> Ag2CrO4 + KC2H3O2 Write the balanced equation (including states) and identify the type of reaction: 99. Aqueous solutions of ammonium chloride and lead (II) nitrate produce lead (II) chloride precipitate and aqueous ammonium nitrate. 100. Solid carbon disulfide burns in oxygen to yield carbon dioxide and sulfur dioxide gases. 101. Iron metal reacts with aqueous silver nitrate to produce aqueous iron (III) nitrate and silver metal. 102. Solid potassium nitrate yields solid potassium nitrite and oxygen gas. 103. Calcium metal reacts with chlorine gas to produce solid calcium chloride. 97 Unit 6- Chemical Reactions Chemical Reactions and Balancing Practice Test 1. Write the balanced equation of the reaction that occurs when calcium carbonate decomposes to form calcium oxide and carbon dioxide. 2. 3. 4. 5. 6. Directions: Balance and identify the type of reaction ____K3PO4 + ____Al(NO3)3 ____KNO3 + ____AlPO4 ____Fe2O3 + ____Al ____Fe + ____Al2O3 ____NaOH ____Na2O + ____H2O ____HCl + ____Mg ____MgCl2 + ____H2 ____C2H4 + ____O2 ____CO2 + ____H2O 7. Write the balanced equation of the synthesis reaction that occurs when iron metal and oxygen gas react to form iron (III) oxide. 8. Write the balanced equation of the combustion reaction that occurs when ethane (C2H6) reacts with oxygen to form carbon dioxide and water. 9. Write the balanced equation of the reaction that occurs when calcium carbonate decomposes to form calcium oxide and carbon dioxide. Directions: Predict the products and balance the reaction 10. _____Na + ______O2 11. _____Al + _____Pb(NO3)2 12. _____NaF + _____Br2 Multiple Choice Practice 13. In the following unbalanced reaction, what is the coefficient of HOH once the reaction is balanced? H2CO3 + KOH --> HOH + K2CO3 a. 1 b. 2 c. 3 d. 4 14. What are the different types of chemical reactions? a. synthesis, fusion, combustion, fission, and decomposition b. single replacement, combustion, and double replacement c. synthesis, fission, single replacement, combustion, and fusion d. synthesis, decomposition, combustion, single and double replacement 15. What type of reaction is represented by 2 or more elements forming a compound? a. Decomposition c. Combustion b. Synthesis d. Single replacement 16. Decomposition is the burning of hydrocarbons in the presence of oxygen. a. True b. False 98 Unit 6- Chemical Reactions 17. Which equation represents a double replacement reaction? a. CaCO3 -> CaO + CO2 c. LiOH + HCl -> LiCl + H2O b. CH4 + 2O2 -> CO2 + 2H2O d. C3H8 + 5O2 -> 3CO2 + 8H2O 18. MgSO4 + BaCl2 --> MgCl2 + BaSO4, is an example of what type of chemical reaction? a. Single replacement c. Combustion b. Synthesis d. Double replacement 19. Zn + 2 AgNO3 --> 2 Ag + Zn(NO3)2, is an example of what type of chemical reaction? a. Synthesis c. Decomposition b. Single replacement d. Double replacement 20. The cation of one compound replaces the cation in another compound in a double replacement reaction. a. True b. False 21. Which statement best describes the conservation of atoms in all balanced chemical equations? a. There is a conservation of mass, number of protons, and charge. b. There is a conservation of mass, electronegativity, and charge. c. There is a conservation of only energy, and charge. d. There is a conservation of mass, energy, and charge. 22. Given the unbalanced equation ; CuS + O2 -> CuO + SO2 When it is balanced, what is the sum of the coefficients? a. 8 b. 9 c. 10 d. 11 23. Which one of these chemical reactions is balanced? a. Na + Cl2 -> NaCl c. CuCO3 -> CuO + CO2 b. H2 + O2 -> H2O d. KClO3-> KCl + O2 24. Which is the correct way of setting up a word equation for this balanced chemical equation, 2Na + Cl2 -> 2NaCl? a. Sodium react with chlorine gas to produce sodium chloride. b. 2 moles of sodium react with 1 mole of chlorine to yield 1 mole of sodium chloride. c. 2 moles of sodium react with 1 mole of chlorine gas to yield 2 mole of sodium chloride. d. 2 moles of sodium added with 1 mole of chlorine gas to yield 1 mole of sodium chloride. 25. 2 moles of copper react with 1 mole of oxygen gas to yield 2 moles of copper (ll) oxide. How would you express this word equation into a balanced chemical equation? a. Cu + O -> CuO c. 2Cu + O -> CuO b. Cu + O2 -> CuO d. 2Cu + O2 -> 2CuO 99 Unit 6- Chemical Reactions 26. When the following reaction is balanced, the sum of all of the coefficients in the equation is: NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2 a. 11 b. 10 c. 6 d. 16 e. 14 27. When two ionic compounds are dissolved in water, a double replacement reaction can...never occur since all ions in water are "spectator ions".occur if two of the ions form an insoluble ionic compound, which precipitates out of solutionoccur if the ions react to form a gas, which bubbles out of the solution.occur only if the ions form covalent bonds with each other. 28. Which equation represents combustion? a. 4 Fe + 3 O2 2 Fe2O3 b. 2H2O - 2H2 + O2 c. CH4 + 2O2 CO2 + 2 H2O d. Cu + 2 AgNO3 Cu(NO3)2 + 2 Ag 29. Given the unbalanced equation: Al + O2 Al2O3 When this equation is completely balanced using the smallest whole numbers, what is the sum of the coefficients? a. 9 b. 7 c. 5 d. 4 30. Given the reaction CH4 + 2O2 --> CO2 + 2H2O, What amount of oxygen is needed to completely react with 1 mole of CH4? a. 2 moles c. 2 grams b. 2 atoms d. 2 molecules 31. The balanced equation for the complete combustion of benzene, C6H6, is a. b. c. d. e. C6H6 + 12 H2O 2 + 15 H2 C6H6 6 C + 3 H2 2 C6H6 + 9 O2 12 CO + 6 H2O C6H6 + O2 CO2 + H2O 2 C6H6 + 15 O2 12 CO2 + 6 H2O 32. Which equation shows conservation of atoms? a. H2 + O2 2O b. H2 + O2 2O 100 c. 2H2 + O2 d. 2H2 + 2O2 2O 2O Unit 6- Chemical Reactions 33. The reaction of magnesium with elemental iodine, I2, yields magnesium iodide. Write a balanced chemical equation for this reaction. Which equation shows conservation of mass and charge? a. b. c. d. NH4Br 3 + Br2 3+ 2+ 2Mg + Fe + 3Fe H2SO4 + LiOH 2SO4 + H2O 2+ Cu + 2Ag + 2Ag 34. When ethanol undergoes complete combustion, the products are carbon dioxide and water. __ C2H5OH + __ O2 __ CO2+ __ H2O What are the respective coefficients when the equation is balanced with the smallest whole numbers? a. 2, 7, 4, 6 b. 1, 3, 2, 3 c. 2, 2, 1, 4 d. 1, 2, 3, 2 e. 2, 4, 6, 4 35. What is the net ionic equation? a. b. c. d. e. AgNO3(aq) + KBr(aq) AgBr(s) + KNO3(aq) K+(aq) + NO3-(aq) KNO3(s) AgNO3(aq) + KBr(aq) AgBr(s) K+(aq) + NO3-(aq) KNO3(aq) AgNO3(aq) + KBr(aq) AgBr(s) + KNO3(aq) Ag+(aq) + Br-(aq) AgBr(s) 36. Which equation shows conservation of atoms? a. H2 + O2 2O b. H2 + O2 2O c. 2H2 + O2 d. 2H2 + 2O2 2O 2O STOICHIOMETRY The word stoichiometry is Greek for “________________________ elements.” The calculations of quantities in chemical reactions are based on a ______________________ equation. We can interpret balanced chemical equations several ways. Using the methods of stoichiometry, we can measure the amounts of substances involved in chemical reactions and relate them to one 101 Unit 7- Stoichiometry another. The group or unit of measure used to count numbers of atoms, molecules, or formula units of substances is the ______________ (abbreviated mol). Moles in Chemical Reactions The coefficients tell us how many moles of each kind of element or compound we have. 2 Al2O3 → 4 Al + 3 O2 2 moles of aluminum oxide form 4 moles of aluminum and 3 moles of oxygen gas. 2 H2 + O2 → 2 H2O ___ mole(s) of hydrogen gas and ___ mole of oxygen form ___ mole(s) of water. 2 Na + 2 H2O → 2 NaOH + H2 ___ moles of sodium and ___ moles of water form ___ moles of sodium hydroxide and ___ mole of hydrogen gas. 2 Al2O3 → 4 Al + 3 O2 Every time we use 2 moles of Al2O3 we make 3 moles of O2. Every time we use 2 moles of Al2O3 we make 4 moles of Al. Using the balanced equation above, how many moles of O2 are produced when 3.34 moles of Al2O3 decompose? 2 C2H2 + 5 O2 → 4 CO2 + 2 H2O a) If 3.84 moles of C2H2 are burned, how many moles of O2 are needed? b) How many moles of C2H2 are needed to produce 8.95 moles of H2O? c) If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed? SiCl4 + 2 Mg → 2 MgCl2 + Si 3.74 mol of Mg would make how many moles of Si? Mass in Chemical Reactions 2 Al2O3 → 4 Al + 3 O2 2 x (102.0) grams of aluminum oxide form 4 x (27.0) grams of aluminum and 3 x (32.0) grams of oxygen. The mass of Al2O3 was found by adding the masses of 2 aluminums & 3 oxygens. (2 x 27.0 + 3 x 16.0 = 102.0) 2 H2 + O2 → 2 H2O 102 Unit 7- Stoichiometry 2 x (_________) grams of hydrogen and ___ x (16.0) of oxygen form ___ x (_______) grams of water. 2 Na + 2 H2O → 2 NaOH + H2 ___ x (23.0) grams of sodium and 2 x (________) grams of water form ___ x (_________) grams of sodium hydroxide and ___ x (________) grams of hydrogen gas. The law of conservation of _____________ applies in chemical reactions. The mass of the reactants equals the mass of the ________________________. Show that the following equation follows the Law of Conservation of Mass. 2 Al2O3 → 4 Al + 3 O2 Mass – Mole Stoichiometry The mass of 1 mole of a pure substance is called its _________________ mass. To convert the mass of an element or compound to the number of moles, use the mass of 1 mol as a conversion factor. We can convert ___________________ to moles using the periodic table. Then we must apply the mole to mole conversion to change chemicals using the balanced equation. Finally we will turn the moles back to grams using the periodic table. 2 C2H2 + 5 O2 → 4 CO2 + 2 H2O a) How many moles of C2H2 are needed to produce 8.95 g of H2O? b) If 2.47 moles of C2H2 are burned, how many grams of CO2 are formed? SiCl4 + 2 Mg → 2 MgCl2 + Si How many moles of Mg are needed to make 9.3 g of Si? 3 Al + 3 NH4ClO4 → Al2O3 + AlCl3 + 3 NO + 6 H2O How many moles of water are produced when 32 grams of aluminum are used? CO2 + 2 LiOH → Li2CO3 + H2O What mass of water can be produced from 3.66 moles of lithium hydroxide (LiOH)? 2 Al + 3 I2 → 2 AlI3 Calculate the mass of AlI3 (Aluminum Iodide) that can be produced from 3.00 mol of Al. Mass – Mass Stoichiometry 2 Fe + 3 CuSO4 → Fe2(SO4)3 + 3 Cu 103 Unit 7- Stoichiometry If 10.1 g of Fe are added to a solution of copper (II) sulfate, how much solid copper would form? 2 Al + 3 I2 → 2 AlI3 Calculate the mass of I2 needed just to react with 35.0 g of Al. SiCl4 + 2 Mg → 2 MgCl2 + Si How many grams of MgCl2 are produced along with 9.3 g of silicon? 3 Al + 3 NH4ClO4 → Al2O3 + AlCl3 + 3 NO + 6 H2O a) How many grams of Al must be used to react with 652 g of NH4ClO4? b) How many grams of NO are produced if 150.0 grams of AlCl3 are also produced? Particles in Chemical Reactions The number of things in one mole is 6.022 x 1023. This big number has a short name: the Avogadro constant. Atom - ______________________ Molecule - Molecular compound (non-metals) or ______________________ (O2 etc.) Formula unit - _________________ Compounds (Metal and non-metal or metal and a polyatomic ion) 2 Al2O3 → 4 Al + 3 O2 2 x (6.022 x 1023) formula units of aluminum oxide form 4 x (6.022 x 1023) atoms of aluminum and 3 x (6.022 x 1023) molecules of oxygen. 2 H2 + O2 → 2 H2O 2 x (___________) molecules of hydrogen and ___ x (6.022 x 1023) molecules of oxygen form ___ x (___________) molecules of water. 104 Unit 7- Stoichiometry 2 Na + 2 H2O → 2 NaOH + H2 ___ x (6.022 x 1023) atoms of sodium and ___ x (___________________) molecules of water form ___ x (__________________) formula units of sodium hydroxide and ___ x (6.022 x 1023) molecules of hydrogen gas. SiCl4 + 2 Mg → 2 MgCl2 + Si How many moles of MgCl2 are produced along with 8.76 x 1024 atoms of silicon? Gases and Reactions In gas conversions, liters of a gas are converted to moles and vice-versa. ____________ stands for standard temperature and pressure. 0ºC is standard _________________________, and 1 atmosphere is standard pressure. At STP, ____________ L of a gas = 1 mole 2 H2O → 2 H2 + O2 If 6.45 grams of water are decomposed, how many liters of oxygen will be produced at STP? CH4 + 2 O2 → CO2 + 2 H2O How many liters of CH4 at STP are required to completely react with 17.5 L of O2? 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O Octane, C8H18, is one of the hydrocarbons in gasoline. How many liters of oxygen are required, at STP, to burn 1.00 g of octane? 2 C4H10 + 13 O2 → 8 CO2 + 10 H2O How many liters of CO2 at STP will be produced from the complete combustion of 23.2 g C4H10? 2 NiS + 3 O2 → 2 NiO + 2 SO2 What volume of sulfur dioxide is produced from 123 grams of nickel (II) sulfide at STP? According to Avogadro, equal volumes of gas, at the _____________ temperature and pressure, contain the same number of particles. _______________ are numbers of particles. We can also change between particles and liters at STP. 2 C4H10 + 13 O2 → 8 CO2 + 10 H2O a) How many molecules of CO2 at STP will be produced from the complete combustion of 18.2 L C4H10 ? 105 Unit 7- Stoichiometry b) How many molecules of O2 at STP are needed to produce 18.2 L of steam? c) How many liters of CO2 at STP are produced from 3.2 x 1024 molecules of butane, C4H10? 4 NH3 + 6 NO → 5 N2 + 6 H2O Nitrogen monoxide is a pollutant found in smokestack emissions. How many liters of ammonia, NH3, at STP are needed to produce 1.4 x 1023 molecules of H2O? Homework / Practice Solve the following problems. The reactions may not be balanced. 1. If 20.0 g of magnesium react with excess hydrochloric acid, how many grams of magnesium chloride are produced? Mg + HCl ---> MgCl2 + H2 2. How many grams of chlorine gas must be reacted with excess sodium iodide if 10.0 g of sodium chloride are needed? NaI + Cl2 ---> NaCl + I2 3. How many moles of oxygen gas are produced in the decomposition of 5.00 g of potassium chlorate? KClO3 ---> KCl + O2 4. What mass of copper is required to replace silver from 4.00 g of silver nitrate dissolved in water? Cu + AgNO3 ---> Cu(NO3)2 + Ag 5. If excess ammonium sulfate reacts with 20.0 g of calcium hydroxide, how many grams of ammonia (NH3) are produced? (NH4)2SO4 + Ca(OH)2 ---> CaSO4 + NH3 + H2O 6. If excess sulfuric acid reacts with 0.2564 moles of sodium chloride, how many grams of hydrochloric acid are produced? H2SO4 + NaCl ---> HCl + Na2SO4 7. How many grams of silver phosphate are produced if 10.0 g of silver acetate react with excess sodium phosphate? AgC2H3O2 + Na3PO4 ---> Ag3PO4 + NaC2H3O2 8. How many moles of sodium hydroxide are needed to completely neutralize 25.0 g of sulfuric acid? NaOH + H2SO4 ---> Na2SO4 + H2O 9. When calcium carbonate is heated strongly, carbon dioxide gas is released. CaCO3 ---> CaO + CO2 What volume of carbon dioxide, measured at STP, is produced if 15.2 g of calcium carbonate is heated? 106 Unit 7- Stoichiometry 10. What volume of oxygen gas at STP is needed for complete combustion of 5.63 g of propane? C3H8 + O2 ---> CO2 + H2O 11. What volume of chlorine gas, measured at STP, is needed to produce 10.0 g of potassium permanganate (KMnO4)? K2MnO4 + Cl2 ---> KMnO4 + KCl 12. Suppose that you could decompose 0.250 mol of Ag2S into its elements. a. How many moles of silver would form? b. How many moles of sulfur would form from 38.8 g of silver sulfide? 13. Ammonia (NH3) is made industrially by reacting nitrogen gas and hydrogen gas under pressure, at high temperature and in the presence of a catalyst. If 4.0 mol of hydrogen react, how many moles of ammonia will be produced? 14. How many grams of sodium hydroxide can be produced from 500. g of calcium hydroxide according to the equation: Ca(OH)2 + Na2CO3 ---> 2 NaOH + CaCO3? 15. How many liters of Cl2 can be produced from 5.60 mole HCl at STP? 4 HCl + O2 ---> 2 Cl2 + 2 H2O 16. Given the equation Al4C3 + 12 H2O ---> 4 Al(OH)3 + 3 CH4 of water are needed to react with 100. g Al4C3? How many moles 17. How many grams of zinc phosphate are formed when 10.0 g of Zn are reacted with the phosphoric acid? The other product is hydrogen gas. 18. Given the equation 4 FeS2 + 11 O2 ---> 2 Fe2O3 + 8 SO2 are required to react with 4.50 mol of FeS2 at STP? How many liters of O2 19. Given the equation C2H4 + 3 O2 ---> 2 CO2 + 2 H2O a. If 6.0 mol of CO2 are produced, how many moles of O2 were reacted? b. How many liters of O2 are required for the complete reaction of 45 g of C2H4 at STP? c. If 18.0 g of CO2 are produced, how many grams of H2O are produced? 20. According to his pre-laboratory theoretical yield calculations, a student’s experiment should have produced 1.44 g of magnesium oxide. When he weighed his product after reaction, only 1.23 g of magnesium oxide was present. What was the student’s percent yield? 107 Unit 7- Stoichiometry Balance the following equations to use with questions 25 – 32: 21. ____ Al + ____ O2 ____ Al2O3 22. ____ Cu + ____ AgNO3 ____ Ag + ____ Cu(NO3)2 23. ____ Zn + ____ HCl ____ ZnCl2 + ____ H2 24. ____ Fe + ____ Cl2 ____ FeCl3 Perform the following calculations: 25. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen. How many moles of HCl are required to produce 7.50 moles of ZnCl2? 26. Copper metal reacts with silver nitrate to form silver and copper (II) nitrate. How many grams of copper are required to form 250 g of silver? 27. When aluminum is burned in excess oxygen, aluminum oxide is produced. How many grams of oxygen are required to produce 0.75 moles of Al2O3? 28. How many grams of iron (III) chloride are produced when 15.3 g of iron react with excess chlorine gas? 29. Copper metal reacts with silver nitrate to form silver and copper (II) nitrate. How many moles of silver will be produced from 3.65 moles of silver nitrate? 30. When 9.34 g of zinc react with excess hydrochloric acid how many grams of zinc chloride will be produced? 31. How many grams of iron are needed to react with 31.0 L of chlorine gas at STP to produce iron (III) chloride? 32. How many liters of oxygen gas at STP are required to react with 65.3 g of aluminum in the production of aluminum oxide? 33. How many liters of ammonia gas are formed when 25 L of hydrogen at STP? ____ N2 + ____ H2 ____ NH3 34. Given the following equation: Na2O + H2O ---> 2 NaOH How many grams of Na2O are required to produce 1.60 x 102 grams of NaOH? 108 Unit 7- Stoichiometry 35. Given the following equation: 8 Fe + S8 ---> 8 FeS a. What mass of iron is needed to react with 16.0 grams of sulfur? b. How many grams of FeS are produced? 36. Given the following equation: 2 NaClO3 ---> 2 NaCl + 3 O2 a. 12.00 moles of NaClO3 will produce how many grams of O2? b. How many grams of NaCl are produced when 80.0 grams of O2 are produced? 37. Given the following equation: Cu + 2 AgNO3 ---> Cu(NO3)2 + 2 Ag a. How many moles of Cu are needed to react with 3.50 moles of AgNO3? b. If 89.5 grams of Ag were produced, how many grams of Cu reacted? 38. Molten iron and carbon monoxide are produced in a blast furnace by the reaction of iron (III) oxide and coke (pure carbon). If 25.0 kilograms of pure Fe2O3 is used, how many kilograms of iron can be produced? The reaction is: Fe2O3 + 3 C ---> 2 Fe + 3 CO 39. The average human requires 120.0 grams of glucose (C6H12O6) per day. How many grams of CO2 (in the photosynthesis reaction) are required for this amount of glucose? The photosynthetic reaction is: 6 CO2 + 6 H2O ---> C6H12O6 + 6 O2 40. Given the reaction: 4 NH3 (g) + 5 O2 (g) ---> 4 NO (g) + 6 H2O (l) When 1.20 mole of ammonia reacts, the total number of moles of products formed is: Stoichiometry Practice Test Directions: Solve the following problems, showing all work. 1. Balance the equation ______NaOH ______Na2O + _____H2O 2. How many moles of water are produced from 4 moles of sodium hydroxide? 3. Balance the equation ____KCl + _____O2 _____KClO3 4. How many moles of potassium chlorate are produced from 9 moles of oxygen? 5. Fe2O3 + 2Al 2Fe + Al2O3 What mass of aluminum oxide is produced when 4 moles of aluminum react? 6. P4O10 + 6H2O 4H3PO4 How many grams of phosphoric acid are produced by the reaction of 12.5g of water? 7. Using the reaction in question 6, how many grams of P4H10 must react if the reaction produces 25 moles H3PO4? 109 Unit 7- Stoichiometry 8. Balance the equation _____HNO3 + ______Cu _______Cu(NO3)2 + ______H2 9. How many moles of nitric acid must react in order to form 83g of copper nitrate? 10. Balance the equation ______NH3 + ________O2 _______NO + ______H2O 11. What mass of nitrogen monoxide will be formed when 7.2g nitrogen trihydride react? 12. C2H4 + 3O2 2CO2 + 2H2O Determine the mass of water produced if 50g C2H4 and 50g O2 react. Multiple Choice Practice 13. Consider the balanced equation Zn + 2HCl -> ZnCl2 + H2 How many moles of ZnCl2 will be produced if 7 moles of HCl are used? a. 2 moles c. 3.5 moles b. 2.5 moles d. 4 moles 14. Given : C2H2(g) + 5O2(g) -> 4CO2(g) + 2H2O(g) Is this chemical equation balanced? a. True b. False 15. In the reaction below, how many moles of oxygen gas is produced by the decomposition of 4moles of mercury (II) oxide? 2HgO -> 2Hg + O2 a. 1 mole c. 3 moles b. 2 moles d. 4 moles 16. True or False, 6 moles of H2 is needed to completely react with 2 moles of N2 in the balanced chemical reaction N2 + 3H2 -> 2NH3 a. True b. False 17. If 18.0 grams of carbon are burned in 55.0 grams of oxygen, how many grams of carbon dioxide are formed? a. 44.01 grams CO2 b. 75.6 grams CO2 c. 151 grams CO2 d. 66.0 grams CO2 18. A mass of 21.5 grams of calcium hydroxide reacts with an excess of phosphoric acid. What mass of calcium phosphate could be recovered from solution? a. 284 grams b. 186 grams c. 94.7 grams d. 31.6 grams e. 326 grams 110 Unit 7- Stoichiometry 19. If 3.64 g of calcium hydroxide react with excess sodium sulfate in aqueous solution to produce solid sulfate and aqueous sodium hydroxide, how many moles of calcium atoms are reacting here? Ca(OH)2(aq) + Na2SO4(aq) CaSO4(s) + 2NaOH(aq) a. 0.00982 mol d. 0.0491 mol b. 0.0246 mol e. 0.0909 mol c. 0.0266 mol 20. If one mole of the rocket fuel ammonium perchlorate, NH4ClO4 (s) is allowed to react with excess Al so all of the NH4ClO4 is consumed, how many molecules of water will be produced? 3NH4ClO4 (s) + 3Al (s) Al2O3 (s) + AlCl3 (s) + 3NO (g) + 6H2O (g) a. 3.61 x 1023 d. 1.20 x 1024 23 b. 1.0 x 10 e. 3.01 x 1024 c. 6.02 x 1023 21. How many grams of potassium cyanide, PCl3, is produced from 93.0 grams of P4 (s) and 213 g of Cl2 (g), assuming the reaction goes to completion? The balanced equation for the reaction is: P4 (s) + 6Cl2 (g) 4PCl3 (g) a. 277 g d. 104 g b. 416 g e. 69.3 g c. 213 g 22. How many moles of Al2O3 are formed when a mixture of 0.36 moles Al and 0.36 moles O2 is ignited? a. 0.12 c. 0.28 e. 0.72 b. 0.18 d. 0.46 23. In the oxidation of ethane: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of O2 are required to react with 1 mole of ethane? a. 7 moles b. 2 moles c. 3.5 moles 24. In the reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of CO2 are formed when 1mole of O2 is consumed? a. 7 moles b. 1.75 moles c. 0.57 moles 25. In the reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of CO2 are formed when 5moles of ethane are consumed? a. 10 moles b. 4 moles 111 c. 2 moles Unit 7- Stoichiometry 26. How many mL of 0.1 M HCl are required to react with 0.01 mole of Na2CO3? HCl + Na2CO3 NaCl + H2CO3 a. 100 mL b. 200 mL c. 50 mL 27. How many liters of H2 at STP are required to react with 2.3 g of Fe3O4? H2 + Fe3O4 FeO + H2O a. 0.22 L b. 0.44 L c. 0.56 L 28. When 0.05 mole H2 is mixed with 0.05 mole CO, what is the maximum number of moles of methanol (CH3OH) that can be obtained? H2 + CO CH3OH a. 0.10 mole b. 0.05 mole 112 c. 0.025 mole Unit 8- Gas Laws THE GAS LAWS The gas laws describe how gases behave. They can be predicted by theory and the amount of change can be calculated with mathematical equations. One ____________________________ is equal to 760 mm Hg, 760 torr, or _______________ kPa (kilopascals). Perform the following pressure conversions. a) 144 kPa = ______________ atm b) 795 mm Hg = ______________ atm c) 669 torr = ______________ kPa d) 1.05 atm = ______________ mm Hg Air pressure at higher altitudes, such as on a mountaintop, is slightly ______________________ than air pressure at sea level. Air pressure is measured using a ________________________. More molecules mean more ____________________ between the gas molecules themselves and more collisions between the gas molecules and the walls of the container. Number of molecules is ________________________ proportional to pressure. Doubling the number of gas particles in a basketball _____________________ the pressure. Gases naturally move from areas of high pressure to ____________ pressure because there is empty space to move in. If you double the number of molecules, you _____________________ the pressure. As you remove molecules from a container, the pressure ________________________ until the pressure inside equals the pressure outside. In a smaller container, molecules have less room to move. The molecules hit the sides of the container _________________ often, striking a smaller area with the same force. As volume decreases, pressure increases. Volume and pressure are ______________________ proportional. As the pressure on a gas increases, the volume decreases. Raising the temperature of a gas increases the _______________________ if the volume is held constant. At higher temperatures, the particles in a gas have greater ________________________ energy. They move faster and collide with the walls of the container more often and with greater ___________________, so the pressure rises. If you start with 1 liter of gas at 1 atm pressure and 300 K and heat it to 600 K, one of 2 things happens. Either the volume will increase to 2 liters at ______ atm, or the pressure will increase to ______ atm while the volume remains constant. 113 Unit 8- Gas Laws Ideal Gases and the Kinetic Molecular Theory In this unit we will assume the gases behave ideally. _____________________ gases do not really exist, but this makes the math easier and is a close approximation. Gas particles are much smaller than the spaces between them. The particles have negligible _____________________. There are no attractive or repulsive ___________________ between gas molecules. Gas particles are in constant, _________________________ motion. Until they bump into something (another particle or the side of a container), particles move in a straight line. No kinetic energy is ____________________ when gas particles collide with each other or with the walls of their container. All gases have the same ______________________ energy at a given temperature. Temperature is a measure of the average kinetic energy of the particles in a sample of matter. There are no gases for which this is true. Real gases behave more ideally at ________________ temperature and _________________ pressure. At low temperature, the gas molecules move more _______________________, so attractive forces are no longer negligible. As the pressure on a gas increases, the molecules are forced closer together and _________________________ forces are no longer negligible. Therefore, real gases behave more ideally at high temperature and low pressure. Avogadro’s Law Avogadro’s law states that equal volumes of different gases (at the same temperature and pressure) contain equal numbers of ________________ or molecules. 2 liters of helium has the same number of particles as ______ liters of oxygen. The molar volume for a gas is the volume that one mole occupies at 0.00ºC and 1.00 atm. 1 mole = 22.4 L at STP (standard temperature and pressure). As a result, the volume of gaseous reactants and products can be expressed as small whole numbers in reactions. How many moles are in 45.0 L of a gas at STP? How many liters are in 0.636 moles of a gas at STP? The volume of a gas is directly proportional to the number of moles. V1 V2 n1 n 2 114 Unit 8- Gas Laws Consider two samples of nitrogen gas. Sample 1 contains 1.5 mol and has a volume of 36.7 L. Sample 2 has a volume of 16.5 L at the same temperature and pressure. Calculate the number of moles of nitrogen in sample 2. If 0.214 mol of argon gas occupies a volume of 652 mL at a particular temperature and pressure, what volume would 0.375 mol of argon occupy under the same conditions? If 46.2 g of oxygen gas occupies a volume of 100. L at a particular temperature and pressure, what volume would 5.00 g of oxygen gas occupy under the same conditions? Boyle’s Law At Boyle’s law states that the pressure and volume of a gas at constant temperature are inversely proportional. Inversely proportional means as one goes up the other goes ________________. P1 V1 = P2 V2 Sketch the PV graph that represents Boyle’s law. The P-V graph for Boyle’s law results in a _____________________________ because pressure and volume are inversely proportional. A balloon is filled with 25 L of air at 1.0 atm pressure. If the pressure is changed to 1.5 atm, what is the new volume? (Make sure the pressure and volume units in the question match.) A balloon is filled with 73 L of air at 1.3 atm pressure. What pressure is needed to change the volume to 43 L? A gas is collected in a 242 cm3 container. The pressure of the gas in the container is measured and determined to be 87.6 kPa. What is the volume of this gas at standard pressure? A gas is collected in a 24.2 L container. The pressure of the gas in the container is determined to be 756 mm Hg. What is the pressure of this gas if the volume increases to 30.0 L? Charles’ Law The volume of a gas is directly proportional to the Kelvin temperature if the pressure is held constant. 115 Unit 8- Gas Laws V1 V2 = T1 T2 K = °C + 273 Sketch the PV graph that represents Charles’ law. The V-T graph for Charles’ law results in a _____________________________ _________________ because pressure and volume are directly proportional. What is the temperature of a gas that is expanded from 2.5 L at 25 ºC to 4.1 L at constant pressure? (Make sure the volume units in the question match and make sure to convert degrees Celsius to Kelvin.) What is the final volume of a gas that starts at 8.3 L and 17 ºC and is heated to 96 ºC? A 225 cm3 volume of gas is collected at 57 ºC. What volume would this sample of gas occupy at standard temperature? A 225 cm3 volume of gas is collected at 42 ºC. If the volume is decreased to 115 cm3, what is the new temperature? Gay-Lussac’s Law The temperature and the pressure of a gas are directly related at constant volume. P1 P2 = T1 T2 Sketch the PT graph that represents Gay-Lussac’s law. What is the pressure inside a 0.250 L can of deodorant that starts at 25 ºC and 1.2 atm if the temperature is raised to 100 ºC? Volume remains constant. (Make sure the pressure units in the question match and make sure to convert degrees Celsius to Kelvin.) A can of deodorant starts at 43 ºC and 1.2 atm. If the volume remains constant, at what temperature will the can have a pressure of 2.2 atm? A can of shaving cream starts at 25 ºC and 1.30 atm. If the temperature increases to 37 ºC and the volume stays constant, what is the pressure of the can? A 12 ounce can of a soft drink starts at STP. If the volume remains constant, at what temperature will the can have a pressure of 2.20 atm? 116 Unit 8- Gas Laws The Combined Gas Law The gas laws may be combined into a single law, called the combined gas law, which relates two sets of conditions of pressure, volume, and temperature by the following equation. P1 V1 P2 V2 = T1 T2 A 15 L cylinder of gas at 4.8 atm pressure at 25 ºC is heated to 75 ºC and compressed to 17 atm. What is the new volume? If 6.2 L of gas at 723 mm Hg at 21 ºC is compressed to 2.2 L at 4117 mm Hg, what is the temperature of the gas? A sample of nitrogen monoxide has a volume of 72.6 mL at a temperature of 16 °C and a pressure of 104.1 kPa. What volume will the sample occupy at 24 °C and 99.3 kPa? A hot air balloon rises to an altitude of 7000 m. At that height the atmospheric pressure drops to 300. mm Hg and the temperature cools to -33 °C. Suppose on the hot air balloon there was a small balloon filled to 1.00 L at sea level and a temperature of 27 °C. What would its volume ultimately be when it reached the height of 7000 m? Dalton’s Law of Partial Pressures Dalton’s law of partial pressures states that the _________________ pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture, as shown below. Pt = P1 + P2 + P3 + … Pt = total pressure The partial pressure is the contribution by that gas. What is the total pressure in a balloon filled with air if the pressure of the oxygen is 170 mm Hg and the pressure of nitrogen is 620 mm Hg? In a second balloon the total pressure is 1.30 atm. What is the pressure of oxygen (in mm Hg) if the pressure of nitrogen is 720. mm Hg? A container has a total pressure of 846 torr and contains carbon dioxide gas and nitrogen gas. What is the pressure of carbon dioxide (in kPa) if the pressure of nitrogen is 50. kPa? When a container is filled with 3 moles of H2, 2 moles of O2 and 4 moles of N2, the pressure in the container is 8.7 atm. The partial pressure of H2 is _____. 117 Unit 8- Gas Laws It is common to synthesize gases and collect them by displacing a volume of ________________. Hydrogen was collected over water at 21°C on a day when the atmospheric pressure is 748 torr. The volume of the gas sample collected was 300. mL. The vapor pressure of water at 21°C is 18.65 torr. Determine the partial pressure of the dry gas. A sample of oxygen gas is saturated with water vapor at 27ºC. The total pressure of the mixture is 772 mm Hg and the vapor pressure of water is 26.7 mm Hg at 27ºC. What is the partial pressure of the oxygen gas? The Ideal Gas Law Remember ideal gases do not exist. Molecules do take up ______________________. There are _________________________ forces; otherwise, there would be no liquids. PV = nRT Pressure times volume equals the number of ___________________ (n) times the ideal gas constant (R) times the temperature in Kelvin. R = 0.0821 (L atm)/(mol K) or R = 8.314 (L kPa)/(mol K) or R = 62.4 (L mm Hg)/(mol K) The one you choose depends on the unit for pressure! How many moles of air are there in a 2.0 L bottle at 19 ºC and 747 mm Hg? What is the pressure in atm exerted by 1.8 g of H2 gas exerted in a 4.3 L balloon at 27 ºC? Sulfur hexafluoride (SF6) is a colorless, odorless and very unreactive gas. Calculate the pressure (in atm) exerted by 1.82 moles of the gas in a steel vessel of volume 5.43 L at 69.5 ºC. Calculate the volume (in liters) occupied by 7.40 g of CO2 at STP. A sample of nitrogen gas kept in a container of volume 2.30 L and at a temperature of 32 ºC exerts a pressure of 476 kPa. Calculate the number of moles of gas present. A 1.30 L sample of a gas has a mass of 1.82 g at STP. What is the molar mass of the gas? Calculate the mass of nitrogen gas that can occupy 1.00 L at STP. 118 Unit 8- Gas Laws Homework / Practice 1. Identify whether the descriptions below describe an ideal gas or a real gas. a) Gas particles move in straight lines until they collide with other particles or the walls of their container. b) Individual gas particles have a measurable volume. c) The gas will not condense even when compressed or cooled. d) Collisions between molecules are perfectly elastic. e) Gas particles passing close to one another exert an attraction on each other. 2. The formula for kinetic energy is KE = 1/2mv2. a) What happens to KE if the mass is tripled (at constant speed)? b) What happens to KE if the speed is halved (at constant mass)? c) If two gases at the same temperature share the same KE, it follows that the molecules of greater mass have the _____ speed. (faster or slower) 3. Explain the following using the kinetic-molecular theory: a) As a gas is heated, its rate of effusion through a small hole increases if all other factors remain constant. b) A strong-smelling gas released from a container in the middle of a room is soon detected in all areas of the room. 4. Pressure = force / area. For a constant force, when the surface area is quadrupled, the pressure 5. Rank the following in increasing order. a) 50 kPa b) 2 atm c) 76 torr 6. Does atmospheric pressure increase or decrease as altitude above sea level increases? 7. Convert the following: a. 0.200 atm = _____ mm Hg b. 790 mm Hg = _____ Pa c. 123 kPa = _____ atm d. 0.935 atm = ______ torr 8. The mercury level in an open manometer is 25 mm Hg higher in the arm connected to the atmosphere. If atmospheric pressure is 765 mm Hg, calculate the pressure of the enclosed gas. 9. The mercury level in an open manometer is 32 mm Hg lower in the arm connected to the atmosphere. If atmospheric pressure is 765 mm Hg, calculate the pressure of the enclosed gas. 10. A 24 L sample of a gas (at fixed mass and constant temperature) exerts a pressure of 3.0 atm. What pressure will the gas exert if the volume is changed to 16 L? 119 Unit 8- Gas Laws 11. An air bubble with a volume of 2.0 mL forms at the bottom of a lake, where the pressure is 3.0atm. As it rises, the pressure on the bubble decreases. Assume the temperature remains constant. a. Will the bubble expand or contract as it rises? b. Predict the volume of the bubble just as it reaches the surface, where the pressure is 1.0atm. 12. A common laboratory system to study Boyle’s law uses a gas trapped in a syringe. The pressure in the system is changed by adding or removing identical weights on the plunger. The original gas volume is 50.0 mL when two weights are present. Predict the new gas volume when 4 more weights are added. 13. Helium gas in a balloon occupies 2.40 L at 400. K. What volume will it occupy at 300 K? 14. A bicycle tire is inflated to 55 lb/in2 at 15 °C. Assume that the volume of the tires does not change appreciably once it is inflated. a. The tire and the air inside it are heated to 30 °C by road wear, does the pressure in the tire increase or decrease? b. Because the temperature has doubled, does the pressure double to 110 psi? Why or why not? 15. If 0.864 mol of neon gas occupies a volume of 452 mL at a particular temperature and pressure, what volume would 0.575 mol of neon occupy under the same conditions? 16. If 26.5 g of oxygen gas occupies a volume of 100. L at a particular temperature and pressure, how many moles of oxygen gas will there be in 350. L under the same conditions? 17. At one point in the cycle of a piston in an automobile engine, the volume of the trapped fuel mixture is 400 cm3 at a pressure of 1.0 atm and a temperature of 27 °C. In the compression of the piston, the temperature reaches 77 °C and the volume decreases to 50.0 cm3. What is the new pressure? 18. On a cold winter morning when the temperature is - 13 °C, the air pressure in an automobile tire is 1.5 atm. If the volume does not change, what will the pressure be after the tire has warmed to 13 °C? 19. A gas storage tank has a volume of 3.5 x1 5 m3 when the temperature is 27 °C and the pressure is 1.0 atm. What is the new volume of the tank if the temperature drops to - 10.°C and the pressure drops to 0.95 atm? 20. Explain how to correct for the partial pressure of water vapor when calculating the partial pressure of a dry gas that is collected over water. 120 Unit 8- Gas Laws 21. A sample of chlorine gas is collected by water displacement at 23 °C. If the atmospheric pressure is 751 torr, what is the partial pressure due to the chlorine? The vapor pressure of water at 23 °C is 21.1 mm Hg. 22. When an explosive like TNT is detonated, a mixture of gases at high temperature is created. Suppose that gas X has a pressure of 50 atm, gas Y has a pressure of 20 atm, and gas Z has a pressure of 10 atm. a. What is the total pressure of the system? b. Once the gas mixture combines with air, Ptotal soon drops to 2 atm. By what factor does the volume of the mixture increases? (Assume mass and temperature are constant.) 23. A gas occupies a volume of 180 mL at 35.0 °C and 740 mm Hg. What is the volume of the gas at STP? 24. Perform the following calculations a. How many moles of methane, CH4, are present in 5.6 L of the gas at STP? b. How many moles of gas are present in 5.6 L of any ideal gas at STP? c. What is the mass of the 5.6 L sample of methane gas? 25. What is the pressure exerted by 32 g of oxygen gas in a 20. L container at 30.0 °C? 26. How many grams of nitrogen gas are in a flask with a volume of 250 mL at a pressure of 3.0 atm and a temperature of 300. K? 27. A container holds three gases: oxygen, carbon dioxide, and helium. The partial pressures of the three gases are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container? 28. A gas occupies 12.3 liters at a pressure of 40.0 mm Hg. What is the volume when the pressure is increased to 60.0 mm Hg? 29. If a gas at 25.0 °C occupies 3.60 liters at a pressure of 1.00 atm, what will be its volume at a pressure of 2.50 atm? 30. A gas occupies 1.56 L at 1.00 atm. What will be the volume of this gas if the pressure becomes 3.00 atm? 31. A gas occupies 11.2 liters at 0.860 atm. What is the pressure if the volume becomes 15.0 L? 32. How much will the volume of 75.0 mL of neon change if the pressure is lowered from 50.0 torr to 8.00 torr? 33. Calculate the decrease in temperature when 2.00 L at 20.0 °C is compressed to 1.00 L. 121 Unit 8- Gas Laws 34. 600.0 mL of air is at 20.0 °C. What is the volume at 60.0 °C? 35. A gas occupies 900.0 mL at a temperature of 27.0 °C. What is the volume at 132.0 °C? 36. What change in volume results if 60.0 mL of gas is cooled from 33.0 °C to 5.00 °C? 37. A gas occupies 1.00 L at standard temperature. What is the volume at 333.0 °C? 38. Determine the pressure change when a constant volume of gas at 1.00 atm is heated from 20.0°C to 30.0 °C. 39. A gas has a pressure of 0.370 atm at 50.0 °C. What is the pressure at standard temperature? 40. A gas has a pressure of 699.0 mm Hg at 40.0 °C. What is the temperature at standard pressure? 41. If a gas is cooled from 323.0 K to 273.15 K and the volume is kept constant what final pressure would result if the original pressure was 750.0 mm Hg? 42. If a gas in a closed container is pressurized from 15.0 atmospheres to 16.0 atmospheres and its original temperature was 25.0 °C, what would the final temperature of the gas be? 43. A gas has a volume of 800.0 mL at negative 23.00 °C and 300.0 torr. What would the volume of the gas be at 227.0 °C and 600.0 torr of pressure? 44. 500.0 liters of a gas are prepared at 700.0 mm Hg and 200.0 °C. The gas is placed into a tank under high pressure. When the tank cools to 20.0 °C, the pressure of the gas is 30.0 atm. What is the volume of the gas? 45. What is the volume of gas at 2.00 atm and 200.0 K if its original volume was 300.0 L at 0.250 atm and 400.0 K. 46. At conditions of 785.0 torr of pressure and 15.0 °C temperature, a gas occupies a volume of 45.5 mL. What will be the volume of the same gas at 745.0 torr and 30.0 °C? 47. A gas occupies a volume of 34.2 mL at a temperature of 15.0 °C and a pressure of 800.0 torr. What will be the volume of this gas at standard conditions? 48. The volume of a gas originally at standard temperature and pressure was recorded as 488.8 mL. What volume would the same gas occupy when subjected to a pressure of 100.0 atm and temperature of minus 245.0 °C? 49. At a pressure of 780.0 mm Hg and 24.2 °C, a certain gas has a volume of 350.0 mL. What will be the volume of this gas under STP 122 Unit 8- Gas Laws 50. A gas sample occupies 3.25 liters at 24.5 °C and 1825 mm Hg. Determine the temperature at which the gas will occupy 4250 mL at 1.50 atm. 51. If 10.0 liters of oxygen at STP are heated to 512 °C, what will be the new volume of gas if the pressure is also increased to 1520.0 mm of mercury? 52. If 2.00 liters of hydrogen, originally at 25.0 °C and 750.0 mm of mercury, are heated until a volume of 20.0 liters and a pressure of 3.50 atmospheres is reached. What is the new temperature? 53. How many moles of gas are contained in 890.0 mL at 21.0 °C and 750.0 mm Hg pressure? 54. Calculate the volume 3.00 moles of a gas will occupy at 24.0 °C and 762.4 mm Hg. 55. What volume will 20.0 g of Argon occupy at STP? 56. How many moles of gas would be present in a gas trapped within a 100.0 mL vessel at 25.0 °C at a pressure of 2.50 atmospheres? 57. How many moles of a gas would be present in a gas trapped within a 37.0 liter vessel at 80.00°C at a pressure of 2.50 atm? 58. At what pressure would 0.150 mole of nitrogen gas at 23.0 °C occupy 8.90 L? 59. What volume would 32.0 g of NO2 gas occupy at 3.12 atm and 18.0 °C? 60. Find the volume of 2.40 mol of gas whose temperature is 50.0 °C and whose pressure is 2.00atm. 61. How many moles of gas are contained in a 50.0 L cylinder at a pressure of 100.0 atm and a temperature of 35.0 °C? 62. Determine the number of moles of Krypton contained in a 3.25 liter gas tank at 5.80 atm and 25.5 °C. If the gas is Oxygen instead of Krypton, will the answer be the same? Why or why not? 63. A 500.0 mL sample of a gas is collected at 745.0 mm Hg. What will the volume be at standard pressure? 64. Convert 350.0 mL at 740.0 mm of Hg to its new volume at standard pressure. 65. The temperature of a sample of gas in a steel container at 30.0 kPa is increased from 100.0 °C to 1.00 x 103 °C. What is the final pressure inside the tank? 123 Unit 8- Gas Laws Gas Laws Practice Test 1. Explain how the temperature is related to the kinetic energy and motion of gas particles. 2. If the volume of a gas contained within balloon were to be tripled, what would be the impact upon the pressure if Kelvin temperature is maintained as constant? Directions: Solve the following problems. Show all your work, including units 3. In a mixture of carbon dioxide, oxygen gas, sulfur dioxide and carbon monoxide, the pressure of the carbon dioxide is 0.3 atm, oxygen gas is 0.5 atm, sulfur dioxide is 0.6 atm, and the pressure of the carbon monoxide is 0.1 atm. What is the total pressure in the container? 4. A high-altitude balloon contains 4 Liters of helium gas at 1.35atm. What is the volume when the balloon rises to an altitude where the pressure is only 1.20atm? (Assume that the temperature remains constant.) 5. If a sample of gas occupies 27 Liters at 12 Celsius, what will be its volume at 112 Celsius if the pressure does not change? 6. A gas has a pressure of 122kPa at -6 Celsius (negative 6). What will be the pressure at 85 Celsius if the volume does not change? 7. A gas at 10 kPa and 45 Celsius occupies a container with an initial volume of 4 Liters. By changing the volume, the pressure of the gas increases to 25 kPa as the temperature is raised to 190 Celsius. What is the new volume? 8. You fill a rigid steel cylinder that has a volume of 840 milliliters with oxygen gas to a final pressure of 1.1 atmospheres at 145 Celsius. How many moles of nitrogen gas does the cylinder contain? 9. What is the temperature when 4moles of carbon dioxide occupies a 2 L container and exerts a pressure of 745torr? 10. What pressure, in atm, will be exerted by 1.25 moles of a gas at 39 Kelvin if it is contained in a 5 Liter vessel? 11. What volume will 29 grams of nitrogen gas occupy at 10 Celsius and a pressure of 620 torr? 12. A 35mL sample of hydrogen gas is collected over water at a temperature of 24oC, the vapor pressure of the water at that temperature is 2.99kPa, and the atmospheric pressure is 765.5 torr. What is the pressure of the dry hydrogen gas? 124 Unit 8- Gas Laws Multiple Choice Practice 13. As the pressure of a gas at 2 atm is changed to 1 atm at constant temperature, the volume of the gas a. decreases c. remains the same b. increases 14. According to the kinetic molecular theory, molecules increase in kinetic energy when they a. Are mixed with other molecules at lower temperature b. Are frozen into a solid c. Are condensed into a liquid d. Are heated to a higher temperature 15. Collide with each other in a container at lower temperature At STP, 32.0 liters of O2 contain the same number of molecules as a. 22.4 L Ar c. 32. 0 L of H2 b. 28.0 L of N2 d. 44.8 L of He 16. What is the total pressure, in atmospheres, of a 10.0 L container that contains 10 moles of nitrogen gas and 10 moles of oxygen gas at 300 K? a. 24.6 L c. 2460 L b. 49.3 L d. 4930 L 17. An 8.25 L sample of oxygen is collected at 25°C and 1.022 atm pressure. What volume will the gas occupy 0.940 atm and -15°C? a. 1.78 L c. 10.4 L e. 7.77 L b. 5.00 L d. 8.76 L 18. A motorist fills his car tires to 32 lb/in2 pressure at a temperature of 30°C. Assuming no change in volume, what will be the pressure in the tires when the motorist drives across Death Valley, with a pavement temperature of 78°C? a. 12 lb/in2 c. 37 lb/in2 e. 83 lb/in2 b. 28 lb/in2 d. 4.8 lb/in2 19. The mass of 2.37 liters of a gas is 8.91 grams. What is the density of the gas? a. 3.76 g/L c. None of these e. 21.1 g/L b. 6.54 g/L d. 0.266 g/L 20. If temperature is constant, the relationship between pressure and volume is a. Direct b. inverse 125 Unit 8- Gas Laws 21. A 268 cm3 sample of an ideal gas at 18°C and 748 torr pressure is placed in an evacuated container of volume 648cm3. To what centigrade temperature must the assembly be heated so that the gas will fill the whole chamber at 748 torr? a. 431°C c. 704°C e. 324°C b. 120°C d. 597°C 22. How big a volume of dry oxygen gas at STP would you need to take to get the same number of oxygen molecules as there are hydrogen molecules in 25.0 liters at 0.850 atm and 35°C a. 18.8 L c. 0.656 L e. 32.3 L b. 0.068 L d. 4.2 L 23. Nitrogen has a molar mass of 28.02 g/mol. What is the density of nitrogen at 1.05 atm and 37°C? a. None of these c. 0.89 g/L e. 4.72 g/cm3 b. 2.82 g/L d. 1.25 g/L 24. How many moles of gas would it take to fill an average man's lungs, total capacity of which is about 4.5 liters? Assume 1.00 atm pressure and 37.0°C. a. 37.0 mol c. 0.75 mol e. 11.2 mol b. 1.24 mol d. 0.18 mole 25. Which flask contains the greatest number of molecules? a. Flask 3 (O2) b. Flask 1 (NH3) c. Flask 2 (CH4) d. Flasks 2 and 3 e. All are the same 26. You have two samples of the same gas in the same size container, at the same pressure. The temperature in the first container is -23°C and the temperature in the second container is 227°C. What is the ratio of the number of moles of gas in the first container compared to that in the second container? a. 2:1 c. 1:2 e. 4:1 b. 1:4 d. 1:1 126 Unit 8- Gas Laws 27. If pressure is constant, the relationship between temperature and volume is a. Direct b. Inverse 28. If pressure of a gas is increased and its volume remains constant, what will happen to its temperature? a. Increase b. Decrease c. Stay the same 29. One way to increase pressure on a gas is to a. decrease temperature b. increase volume c. increase the number of gas particles d. lower the kinetic energy of the gas molecules 30. How do gas particles respond to an increase in volume? a. increase in kinetic energy and decrease in temperature b. decrease in kinetic energy and decrease in pressure c. increase in temperature and increase in pressure d. increase in kinetic energy and increase in temperature 31. If a gases volume is decreased and pressure is constant, its temperature will a. Increase b. Decrease c. Stay the same 32. If the temperature of a gas remains constant but pressure is decreased, the volume will a. Increase b. Decrease c. Stay the same 33. Convert 2.3 atm into mmHg a. 2300 mmHg b. 1750 mmHg c. 2.3 mmHg d. 0.0030 mmHg 34. The pressure of a gas is 750.0 torr when its volume is 400.0 mL. Calculate the pressure (in atm) if the gas is allowed to expand to 600.0 mL at constant temperature. a. 0.660 atm c. 500.0 atm b. 1.48 atm d. 1125 atm 127 Unit 8- Gas Laws 35. The volume of a gas is increased from 150.0 mL to 350.0 mL by heating it. If the original temperature of the gas was 25.0 °C, what will its final temperature be (in °C)? a. - 146°C c. 58.3°C e. 695°C b. 10.7°C d. 422°C 36. Standard temperature and pressure (STP) refers to which conditions? a. 0 oC and 1 kPa d. 0 K and 1 atm b. 0 oC and 1 mm Hg e. 273 K and 1 atm c. 0 K and 1 kPa 37. If 4 moles of a gas are added to a container that already holds 1 mole of gas, how will the pressure change within the container? (Assume volume and temperature are constant.) a. The pressure will be 5 times as great. b. The pressure will be 2 times as great. c. The pressure will be 4 times as great. d. The pressure will not change. e. None of the above are correct. 38. A 4.0 L sample of hydrogen gas at 700 mm Hg would occupy what volume at 250 mm Hg? (Assume temperature and number of particles stays constant.) a. 1.4 x 10 -7 L c. 11.2 L e. 7.0 x 10 5 L b. 1.4 L d. 2.4 L 39. A 25 L tank of oxygen under a pressure of 80. atm would require what pressure to decrease the volume to 1.0 L? (Assume temperature and number of particles stays constant.) a. 0.31 atm b. 3.2 atm c. 2000 atm d. There is not enough information to answer the question. e. None of these is correct. 40. A balloon containing 2.50 L of gas at 1 atm would be what volume at a pressure of 300 KPa? (Assume temperature and number of particles stays constant.) a. 6.33 L c. 0.844 L e. 000833 L b. 8.11 L d. 120. L 41. A syringe containing 75.0 mL of air is at 298 K. What will the volume of the syringe be if it is placed in a boiling water bath (373 K). Assume pressure and the number of particles are held constant. a. 59.9 mL b. 188 mL c. 300. mL d. 8.34 x 106 mL e. None of the above are correct. 42. A gas occupies 40.0 mL at 127 oC. What volume will it occupy at -73 oC? (Assume pressure and number of particles is constant.) a. 182 mL b. 8.80 mL c. 80.0 mL d. 20.0 mL e. None of these is correct 128 Unit 8- Gas Laws 43. If 88.0 grams of solid carbon dioxide evaporates, how many liters of CO2 gas will be formed at a temperature of 300 K and 2.00 atmospheres of pressure? a. 98.5 liters c. 24.6 liters b. 2170 liters d. 1080 liters 44. Which of the following equations correctly combines Boyle's and Charles' Laws? a. d. b. e. c. 45. A 50.0 mL sample of a gas is at 3.00 atm of pressure and a temperature of 298 K . What volume would the gas occupy at STP? a. 0.00728 mL d. 137 mL b. 15.3 mL e. None of these is correct. c. 18.2 mL 46. A syringe contains 60.0 mL of air at 740 mm Hg pressure and 20 oC. What would be the temperature at which the syringe would contain 30.0 mL at a pressure of 370 mm Hg? (Assume no gas could leak in or out of the syringe.) a. -200 oC d. 73.3 oC o b. 0.0137 C e. None of these is correct o c. 5 C 47. A sealed container contains 1.0 mol of hydrogen and 2.0 moles of nitrogen gas. If the total pressure in the container is 1.5 atm, what is the amount of pressure exerted by each gas? a. H2 = 1.0 atm and N2 = .50 atm b. H2 = 0.50 atm and N2 = 1.0 atm c. H2 = 1.0 atm and N2 = 2.0 atm d. H2 = 2.0 atm and N2 = 1.0 atm e. There is not enough information given to answer the question. 48. A sample of gas is collected by water displacement. The atmospheric pressure in the room is 757mm Hg and the vapor pressure of water is 17 mm Hg. What is the partial pressure of hydrogen under these conditions? a. 17 mm Hg c. 757 mm Hg b. 740 mm Hg d. 774 mm Hg e. You cannot answer this question because you do not know the temperature. 129 Unit 9- Solids, Liquids and Phase Changes SOLIDS AND LIQUIDS States of Matter There are ______ states of matter. A solid is a form of matter that has its own definite _____________ and volume. A solid cannot _________________. The particles can vibrate but cannot move around. The particles of matter in a solid are very tightly ____________________; when heated, a solid expands, but only slightly. A liquid is a form of matter that flows, has ____________________ (definite) volume, and takes the _________________ of its container. The particles in a liquid are not rigidly held in place and are _______________ closely packed than are the particles in a solid; liquid particles are able to move past each other. A liquid is not very __________________________. Like solids, liquids tend to expand when heated. A gas is a form of matter that flows to conform to the ____________________ of its container and fills the entire _______________________ of its container. Compared to solids and liquids, the particles of gases are very far apart. Because of the significant amount of space between particles, gases are easily compressed. _____________________ is composed of electrons and positive ions at temperatures greater than ____________ °C. The sun and other stars are examples of plasma. Identify the following as a property of a solid, liquid or gas. The answer may include more that one state of matter. 1. flows and takes the shape of a container 2. compressible 3. made of particles held in a specific arrangement 4. has definite volume 5. always occupies the entire space of its container 6. has a definite volume but flows The word_____________________ refers to the gaseous state of a substance that is a solid or a liquid at room temperature. For example, steam is a vapor because at room temperature water exists as a liquid. Some substances are described as _______________________, which means that they change to a gas easily at room temperature. Alcohol and gasoline are ______________ volatile than water. Kinetic-molecular theory predicts the constant motion of the liquid particles. Individual liquid molecules do not have fixed positions in the liquid. However, forces of 130 Unit 9- Solids, Liquids and Phase Changes ________________________ between liquid particles limit their range of motion so that the particles remain closely packed in a fixed volume. These attractive forces are called ___________________________ forces. Inter = between. Molecular = molecules. A liquid diffuses more _______________________ than a gas at the same temperature, however, because intermolecular attractions interfere with the flow. __________________________ is a measure of the resistance of a liquid to flow. Viscosity decreases with ________________________ temperature. Particles in the middle of the liquid can be attracted to particles above them, below them, and to either side. For particles at the surface of the liquid, there are no attractions from above to balance the attractions from _______________. Thus, there is a net attractive force pulling down on particles at the surface. _____________________ ____________________ is a measure of the inward pull by particles in the interior. Soaps and detergents decrease the surface tension of water by disrupting the _______________________ bonds between water molecules. For a substance to be a solid rather than a liquid at a given temperature, there must be strong attractive forces acting between particles in the solid. These forces limit the motion of the particles to __________________________ around fixed locations in the solid. Thus, there is more order in a solid than in a liquid. The particles can only vibrate and revolve in place. Because of this order, solids are much less _________________ than liquids and gases. In fact, solids are not classified as fluids. Most solids are more _________________ than most liquids. A crystalline solid is a solid whose atoms, ions, or molecules are arranged in an orderly, geometric, three-dimensional structure. Most solids are _____________________. Amorphous solids lack an orderly internal structure. Think of them as __________________________ liquids. Examples of amorphous solids include ____________________ and glass. Phase Changes If a substance is usually a liquid at room temperature (as water is), the gas phase is called a _________________. Vaporization is the process by which a liquid changes into a gas or vapor. Vaporization is an endothermic process - it requires _______________. When vaporization occurs only at the _____________________ of an uncontained liquid (no lid on the container), the process is called evaporation. Molecules at the surface break away and become gas. Only those with enough _____________________ energy (KE) escape. Evaporation is a _______________________ process. It requires heat, which is endothermic. 131 Unit 9- Solids, Liquids and Phase Changes __________________ pressure is the pressure exerted by a vapor over a liquid. As temperature increases, water molecules gain kinetic energy and vapor pressure ______________________. When the vapor pressure of a liquid equals atmospheric pressure, the liquid has reached its boiling point, which is 100°C for water at sea level. Recall that standard atmospheric pressure equals ______ atm. At this point, molecules throughout the liquid have the energy to enter the gas or vapor phase. The temperature of a liquid can never ______________ above its boiling point. Boiling is an __________________________ process. It requires the addition of heat. As you go up into the mountains (increase in elevation), atmospheric pressure ______________. Lower external pressure requires ______________________ vapor pressure. Lower vapor pressure means lower ______________________ point. As a result, spaghetti cooks slower in the mountains than at the beach. When you use a pressure cooker to can vegetables, the external pressure around the mason jars rises. This raises the vapor pressure needed in order for water to boil. In turn, the boiling point is raised so the food cooks ______________________. Some phase changes release energy into their surroundings. For example, when a vapor loses energy, it may change into a __________________. Condensation is the process by which a gas or vapor becomes a liquid. It is the ___________________ of vaporization. In a closed system, the rate of vaporization can equal the rate of condensation. When first sealed, the molecules gradually _________________ the surface of the liquid. As the molecules build up above the liquid, some condense back to a liquid. Equilibrium is reached when the rate of vaporization __________________ the rate of condensation. Molecules are constantly changing phase. The total amount of liquid and vapor remains _______________________. The melting point of a solid is the temperature at which the ____________________ holding the particles together are broken and the solid becomes a liquid. When heated the particles vibrate more _____________________ until they shake themselves free of each other. The freezing point is the temperature at which a liquid becomes a _________________________ solid. The freezing point is the _______________ as the melting point. The process by which a solid changes directly into a gas without first becoming a liquid is called _______________________. Solid air fresheners and dry ice are examples of solids that sublime. When a substance changes from a gas or vapor directly into a solid without first becoming a liquid, the process is called 132 Unit 9- Solids, Liquids and Phase Changes _________________________. Deposition is the reverse of sublimation. _______________ is an example of water deposition. Classify the following phase changes. 1. dry ice (solid carbon dioxide) to carbon dioxide gas ____________________________ 2. ice to liquid water ________________________________ 3. liquid water to ice ________________________________ 4. water vapor to liquid water ________________________________ Phase Diagrams Temperature and _____________________ control the phase of a substance. A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure. A phase diagram typically has ______ regions, each representing a different phase and three curves that ________________________ each phase. 0.0098 Temperature (°C) The points on the curves (lines) indicate conditions under which two phases coexist. The critical point indicates the critical pressure and the critical temperature above which a substance cannot exist as a ____________________. The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can __________________________. The __________________________ slope of the solid-liquid line in the phase diagram for water indicates that the solid floats on its liquid. 133 Unit 9- Solids, Liquids and Phase Changes What happens to solid CO2 at -100 ºC and 1 atm pressure as it is heated to room temperature? What happens to water at 1 atm as the temperature rises from -15°C to 60°C? What state of matter is water at 50°C and 20 atm? At what temperature does the triple point occur for water? At what temperature does the critical point occur for carbon dioxide? At standard pressure and -78°C, what phase change occurs for carbon dioxide? What state of matter is carbon dioxide at -80°C and 2 atm? Solids and Liquids Practice Test Directions: Identify the proper sections by indicating the interval between 2 letters. Heating Curve 1. Which section represents the gas being warmed? _____ to _______ 2. Which section represents a phase change from solid to liquid? _____ to H A F 200º C G 60º C B D E C Energy 3. Define viscosity4. The temperature at which the vapor pressure of a liquid equals the external or atmospheric pressure is known as the________________________ 134 _______ Unit 9- Solids, Liquids and Phase Changes Directions: Using the phase diagram below, answer questions 5-7: Pressure (atm) 5. What does letter T represent? _______________ C 0.75 6. At 0.75 atm, what is the T 0.25 MP_____and BP______ 7. At 0.25 atm, what is the freezing 125 point? ________ 175 250 o Temperature ( C) Multiple Choice 8. Under the same conditions of temperature and pressure, a liquid differs from a gas because the particles of the liquid a. are in constant straight-line motion b. take the shape of the container they occupy c. have no regular arrangement d. have stronger forces of attraction between them 9. The phase change represented by the equation I2 (s) ----> I2 (g) is called a. sublimation c. melting b. condensation d. boiling 10. Which of the following terms represents the temperature and pressure at which three states of a compound can coexist a. Law of definite composition d. Triple point b. Van der Waals forces e. Critical point c. Graham’s Law of Diffusion 11. What is the smallest portion of a crystal lattice that reveals the 3-dimensional pattern? a. unit cell c. coordinate system b. crystal structure d. crystalline symmetry 135 Unit 9- Solids, Liquids and Phase Changes 12. What forces hold nonpolar particles together? a. magic b. hydrogen bonding c. London dispersion d. dipole-dipole 13. Compared with the particles in a solid, the particles in a liquid usually are a. higher in energy c. more massive b. closer together d. less fluid 14. What is the process of a substance changing from a vapor to a solid without passing through the liquid phase? a. condensation c. sublimation b. deposition d. evaporation 15. A liquid forms when the average energy of a solid substance's particles a. increases c. creates an orderly b. changes form arrangement d. decreases 16. Which of the following is an NOT an amorphous solid? a. silly putty c. ice b. play dough d. glass 17. Which term best describes the process by which particles escape from both the surface of a liquid and from within the liquid itself and enter the gas phase? a. boiling c. aeration b. evaporation d. surface tension 18. The attractive forces in a solid are a. too weak to prevent the particles from changing positions b. strong enough to hold the particles in fixed positions c. less effective than those in a liquid d. weaker than those of a liquid particles 19. When electrons in a covalent bond spend more time around on nucleus of the compound than the other, the molecule is considered a. weak c. ionic b. polar d. nonpolar 136 Unit 9- Solids, Liquids and Phase Changes 20. Which of the following phase changes results in an overall increase in randomness of particles over the course of the change? a. deposition c. melting b. condensation d. freezing 21. What type of crystals are like giant molecules? a. covalent network b. covalent molecular c. metallic d. ionic 22. The difference between crystalline and amorphous solids is determined by a. temperature changes b. pressure when the substances are formed c. strength of molecular forces d. the particle arrangement 23. Which of the following statements is false? a. Condensed states have much higher densities than gases. b. Molecules are very far apart in gases and closer together in liquids and solids. c. Gases completely fill any container they occupy and are easily compressed. d. Vapor refers to a gas formed by evaporation of a liquid or sublimation of a solid. e. Solid water (ice), unlike most substances, is denser than its liquid form (water). 24. Which physical state/ property is incorrectly matched? a. liquids and solids - rigid d. solids - higher density than shape gases b. gases - easily compressed e. liquids – incompressible c. gases and liquids – flow 25. Which one of the following statements does not describe the general properties of liquids accurately? a. Liquids have characteristic volumes that do not change greatly with changes in temperature. (Assuming that the liquid is not vaporized.) b. Liquids have characteristic volumes that do not change greatly with changes in pressure c. Liquids diffuse only very slowly when compared to solids. d. The liquid state is highly disordered compared to the solid state. e. Liquids have high densities compared to gases. 137 Unit 9- Solids, Liquids and Phase Changes 26. For which of the following would permanent dipole-dipole interactions play an important role in determining physical properties in the liquid state? a. BF3 c. BeCl2 e. CCl4 b. ClF d. F2 27. Identify which property liquids do not have in common with solids. a. rigid shape b. volumes do not change significantly with pressure c. hydrogen bonding forces can be significant d. practically incompressible e. volumes do not change significantly with temperature 28. Which one of the following statements does not describe the general properties of solids accurately? a. Solids have characteristic volumes that do not change greatly with changes in temperature. b. Solids have characteristic volumes that do not change greatly with changes in pressure. c. Solids diffuse only very slowly when compared to liquids and gases. d. Solids are not fluid. e. Most solids have high vapor pressures at room temperature. 29. Which of the following interactions are the strongest? a. hydrogen bonding force c. dipole- dipole force b. ion-ion interactions d. London-dispersion force 30. For which of the following would dispersion forces be the most important factor in determining physical properties in the liquid state? a. H2O c. F2 e. NH4Cl b. NaCl d. HF 31. For which of the following would hydrogen bonding not be an important factor in determining physical properties in the liquid state? a. HI c. HF e. H2O2 b. H2O d. NH3 138 Unit 9- Solids, Liquids and Phase Changes 32. Which one of the following statements does not describe the general properties of liquids accurately? a. In the liquid state the close spacing of molecules leads to large intermolecular forces that are strongly dependent on the nature of the molecules involved. b. Liquids are practically incompressible. c. As the temperature of a liquid is increased, the vapor pressure of the liquid decreases. d. The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid becomes equal to exactly 760 torr. e. Vapor pressures of liquids at a given temperature differ greatly, and these differences in vapor pressure are due to the nature of the molecules in different liquids. 33. Which technique listed below separates a mixture of liquids on the basis of their boiling points? a. Distillation d. Reflux b. Extraction e. None of the above c. Filtration 34. The melting point of a solid is the same as the ____ of its liquid. a. Boiling point d. Condensation point b. Freezing point e. Critical point c. Sublimation point 35. Some solids can be converted directly to the vapor phase by heating. The process is called ____. a. Fusion d. Condensation b. Sublimation e. Distillation c. Vaporization 36. Which of the images shown here depicts a phase that has definite volume but not definite shape? a. The one on b. The one in the c. The one on the left middle the right 139 Unit 9- Solids, Liquids and Phase Changes 37. 38. Which phase depicted here has both a definite shape and a definite volume? a. The one in the middle b. The one in the middle and the one on the right c. The one on the right d. The one on the left Which of the phases depicted here can be easily compressed? a. The one in the middle d. The one in the middle and the b. The one on the right one on the right c. The one on the left 39. a. Liquid b. Gas Which phase of matter is depicted here? c. Plasma d. Solid 40. Ice floats in water because: a. Water is denser than ice b. Ice is colder than water c. Water has a substantial surface tension d. Ice is denser than water 140 Unit 9- Solids, Liquids and Phase Changes 41. Which phase(s) depicted here have the ability to flow? a. The one on the right b. The one on the left c. The ones on the right and the left d. The one in the middle and the one on the right e. The one in the middle 42. During the phase change from liquid to solid: a. energy must be removed b. energy must be absorbed c. there is no change in energy 43. Definite shape, definite volume, and a low rate of diffusion are characteristics of: a. Fluids c. Gases b. Liquids d. Solids 44. a. Solid b. Gas Which phase of matter is depicted here? c. Liquid d. Plasma 141 Unit 10- Solutions and Solubility SOLUTIONS A solution is made up of a solute and a _______________________________. The solvent does the ________________________________. The solute is the substance that is dissolved. If a solution is made of two liquids, the one in ______________________ quantity is the solute. _________________________ is the universal solvent. Water is a versatile solvent because of its attraction to other molecules and its ___________________________. Most of the water on the Earth is not pure, but rather is present in solutions. Table salt (NaCl), like a great many ionic compounds, is _________________________ in water. The salt solution is also an excellent ___________________________ of electricity. This high level of electrical conductivity is always observed when ionic compounds dissolve to a significant extent in water. The process by which the charged particles in an ionic solid separate from one another is called _____________________________. You can represent the process of dissolving and dissociation in shorthand fashion by the following equation. ________________________________________ Water is not only good at dissolving ionic substances. It also is a good solvent for many _________________________________ compounds. Consider the covalent substance sucrose, commonly known as table sugar, as an example. Although water dissolves an enormous variety of substances, both ionic and covalent, it does not dissolve everything. The phrase that scientists often use when predicting solubility is “________________ dissolves like.” The expression means that dissolving occurs when similarities exist between the solvent and the solute. A salt dissolves faster if it is _________________________ or shaken, if the particles are made ___________________________ and if the temperature is ___________________________. In order to dissolve the solvent molecules must come in ______________________________ with the solute. Stirring moves fresh _____________________________ next to the solute. The solvent touches the surface of the solute. __________________________________ pieces increase the amount of surface of the solute. For solids in liquids, as the temperature goes up the solubility goes ______________________. A higher temperature makes the molecules of the solvent move around ______________________________ and contact the solute harder and ________________________ often. It speeds up dissolving. Higher temperature usually increases the _______________________________ that will dissolve. 142 Unit 10- Solutions and Solubility Figure 1 How many grams of sodium hydroxide (NaOH) will dissolve in 100 g of water at 15ºC? How many grams of sodium hydroxide will dissolve in 100 g of water at 40ºC? At what temperature will 90 grams of Pb(NO3)2 dissolve in 100 g of water? At what temperature will 30 grams of KNO3 dissolve in 100 g of water? For gases in a liquid, as the temperature goes up the solubility goes _______________________. For gases in a liquid, as the pressure goes up the solubility goes ______________________. Solubility is the ________________________________ amount of substance that will dissolve at that temperature (usually measured in grams/liter). If the amount of solute dissolved is less than the maximum that could be dissolved, the solution is called a(n) ___________________________ solution. A solution which holds the maximum amount of solute per amount of the solution under the given conditions is called a(n) _____________________________ solution. A(n) _________________________________ solution contains more solute than the usual maximum amount and are unstable. They cannot permanently hold the excess solute in solution and may release it suddenly. A(n) __________________ crystal will make the extra come out. Generally, a supersaturated solution 143 Unit 10- Solutions and Solubility is formed by dissolving a solute in the solution at an elevated temperature, at which solubility is _______________________ than at room temperature, and then slowly cooling the solution. Figure 2 If 45 g of KCl is dissolved in 100 g of water at 60ºC, is the solution unsaturated, saturated or supersaturated? If 90 g of Pb(NO3)2 is dissolved in 100 g of water at 40ºC, is the solution unsaturated, saturated or supersaturated? If 30 g of KNO3 is dissolved in 100 g of water at 20ºC, is the solution unsaturated, saturated or supersaturated? If 10 g of KClO3 is dissolved in 100 g of water at 50ºC, is the solution unsaturated, saturated or supersaturated? ___________________________ means that two liquids can dissolve in each other. ________________________________ means they cannot. Oil and ______________________ are immiscible. 144 Unit 10- Solutions and Solubility Measuring Solutions Chemists never apply the terms strong and weak to solution concentrations. Instead, use the terms concentrated and _________________________. Concentration is a measure of the amount of solute dissolved in a certain amount of solvent. A concentrated solution has a _________________________ amount of solute. A dilute solution has a __________________________ amount of solute. For chemistry applications, the concentration term molarity is generally the most useful. Molarity is the number of moles of _______________________ in 1 Liter of the solution. moles Molarity = ------------------------Liters of solution Note that the volume is the total solution volume that results, not the volume of solvent alone. Suppose you need 1.0 Liter of a 1 M copper (II) sulfate solution. STEP 1: Measure a mole of copper (II) sulfate. STEP 2: Place the CuSO4 in a volumetric flask. STEP 3: Add some water to dissolve the CuSO4 and then add enough additional water to bring the total volume of the solution to 1.0 L. What is the molarity of a solution with 2.0 moles of NaCl in 4.0 Liters of solution? What is the molarity of a solution with 3.0 moles dissolved in 250 mL of solution? How many moles of NaCl are needed to make 6.0 L of a 0.75 M NaCl solution? 0.200 moles of NaOH are dissolved in a small amount of water then diluted to 500. mL. What is the concentration? 1.25 moles of NaCl are dissolved in a small amount of water then diluted to 625 mL. What is the concentration? How many moles are in 2.00 L of a 3.00 M solution of sulfuric acid (H2SO4)? How many moles are in 1500 mL of a 3.2 M solution of nitric acid (HNO3)? 10.3 g of NaCl are dissolved in a small amount of water then diluted to 250 mL. What is the concentration? 145 Unit 10- Solutions and Solubility 20.3 g of NaOH are dissolved in a small amount of water then diluted to 500. mL. What is the concentration? 80.6 g of KCl are dissolved in a small amount of water then diluted to 500. mL. What is the concentration? 125 g of NaC2H3O2 are dissolved in a small amount of water then diluted to 750. mL. What is the concentration? How many grams of CaCl2 are needed to make 625 mL of a 2.00 M solution? How many grams of sugar are needed to make 125 mL of a 0.500 M C6H12O6 solution? How many grams of sodium hydroxide are needed to make 500. mL of a 0.750 M NaOH solution? How many grams of aluminum nitrate are needed to make 600. mL of a 0.500 M Al(NO3)2 solution? Refer to Figure 1 to answer the following questions: What is the molarity of a KNO3 solution at 10ºC? (100 g of water = 100 mL of water) What is the molarity of a Pb(NO3)2 solution at 61ºC? What is the molarity of a KNO3 solution at 71ºC? Dilution The number of moles of solute doesn’t change if you add more solvent. M1 x V1 = M2 x V2 M1 and V1 represent the starting concentration and volume. M2 and V2 represent the ______________ concentration and volume. 2.0 L of a 0.88 M solution are diluted to 3.8 L. What is the new molarity? 6.0 L of a 0.55 M solution are diluted to 8.8 L. What is the new molarity? You have 150 mL of 6.0 M HCl. What volume of 1.3 M HCl can you make? 6.0 liters of a 0.55 M solution are diluted to a 0.35 M solution. What is the final volume? You need 450 mL of 0.15 M NaOH. All you have available is a 2.0 M stock solution of NaOH. How do you make the required solution? 146 Unit 10- Solutions and Solubility Compounds in Aqueous Solution and Double Replacement Reactions The _________________________________ of ions when an ionic compound dissolves in water is called dissociation. Although no compound is completely insoluble, compounds of very low solubility can be considered insoluble. Using the solubility rules printed on page 6 of the NCDPI Reference Tables for Chemistry, determine whether the following salts are soluble in water. a) sodium chloride _____________________ b) mercury (I) acetate _____________________ c) potassium nitrate _____________________ d) nickel carbonate _____________________ e) barium sulfate _____________________ f) ammonium bromide ____________________ g) calcium sulfide _____________________ In a double-replacement reaction, two compounds exchange partners with each other to produce two different compounds. The general form of the equation is AB + CD ---> AD + CB Signs that a double-replacement reaction has taken place include a color change, the release or absorption of energy, evolution of a gas, and formation of a _______________________________. Write the net ionic equation for each of the following precipitation reactions. a) barium chloride + silver nitrate ___________________________________________ b) lead (II) nitrate + potassium iodide _______________________________________ c) ammonium sulfate + barium nitrate _______________________________________ d) potassium sulfide + zinc acetate __________________________________________ e) strontium phosphate + aluminum hydroxide __________________________________ 147 Unit 10- Solutions and Solubility Homework / Practice 1. Suppose you had 2.00 moles of solute dissolved into 1.00 L of solution. What's the molarity? 2. Calculate the molarity of 25.0 grams of KBr dissolved in 750.0 mL. 3. 80.0 grams of glucose (C6H12O6, mol. wt = 180. g/mol) is dissolved in enough water to make 1.00 L of solution. What is its molarity? 4. What is the molarity when 0.75 mol is dissolved in 2.50 L of solution 5. What is the molarity of 245.0 g of H2SO4 dissolved in 1.00 L of solution? 6. What is the molarity of 5.00 g of NaOH in 750.0 mL of solution? 7. How many moles of Na2CO3 are there in 10.0 L of 2.0 M soluton? 8. How many moles of Na2CO3 are in 10.0 mL of a 2.0 M solution? 9. How many grams of Ca(OH)2 are needed to make 100.0 mL of 0.250 M solution? 10. What is the molarity of a solution made by dissolving 20.0 g of H3PO4 in 50.0 mL of solution? 11. What weight (in grams) of KCl is there in 2.50 liters of 0.50 M KCl solution? 12. What is the molarity of a solution containing 12.0 g of NaOH in 250.0 mL of solution? 13. Determine the number of moles of solute to prepare these solutions: a. 2.35 liters of a 2.00 M Cu(NO3)2 solution. b. 16.00 mL of a 0.415-molar Pb(NO3)2 solution. c. 3.00 L of a 0.500 M MgCO3 solution. d. 6.20 L of a 3.76-molar Na2O solution. 14. Determine the final volume of these solutions: e. 4.67 moles of Li2SO3 dissolved to make a 3.89 M solution. f. 4.907 moles of Al2O3 to make a 0.500 M solution. g. 0.783 grams of Na2CO3 to make a 0.348 M solution. h. 8.97 grams of (NH4)2CO3 to make a 0.250-molar solution. i. 48.00 grams of PbCl2 to form a 5.0-molar solution. 148 Unit 10- Solutions and Solubility Solutions Practice Test Directions: For credit, show all steps in your calculations and include units. 1. What is the molarity of a solution of NaOH if 12 liters of the solution contains 3 moles of NaOH? 2. You have a 3.5 L solution that contains 20 grams of NaCl. What is the molarity of the solution? For these common substances, identify what is the solute and solvent. 3. KoolAid (sugar, water) 4. Vinegar (acetic acid, water) Solute Solvent Determine whether, according to the solubility rules, the mixing of these substances will make a solution. 5. Water and Mg(OH)2 6. Water and Na2CO3 Yes No Faster Slower Determine how the following conditions can affect the rate of dissolving KCl in water. 7. Decrease the temperature of the water 8. Agitate the mixture Directions: Using the solubility curve below, answer the following questions. 9. Which is most soluble at 40ºC? _______ 10. How many grams of KClO3 can be dissolved in 100g H2O at 90ºC? _____ 11. At 40ºC, how much HCl can be dissolved in 300 g. H2O? _________ 149 Unit 10- Solutions and Solubility Multiple Choice 12. Which of these compounds are soluble in water? f. CaBr2 g. PbCl2 h. SrS i. CaCO3 13. Which of these compounds is insoluble in water? a. CuI b. CaCl2 c. MgS d. NaBr 14. Iron (III) sulfide is soluble in water. a. True b. False 15. LiBr is a. Soluble b. Insoluble c. can't tell the solubility d. a covalent compound 16. NH4OH is insoluble. a. True b. False 17. Which of these compounds is soluble? a. Pb(OH)4 b. NaHCO3 c. BaCrO4 d. Mg3(PO4)2 18. Powdered NaCl will dissolve slower then NaCl crystals because there is less surface area for the reaction to take place. a. True b. False 19. Which term indicates that there is a large quantity of solute, compared to the amount of solvent in a solution a. Dilute c. Unsaturated b. Concentrated d. Saturated 20. Ten grams of sodium hydroxide is dissolved in enough water to make 1L of solution. What is the molarity of the solution? a. 0.25 M c. 1 M b. 0.5 M d. 1.5 M 21. Which solution is the most concentrated? a. 1 mole of solute dissolved in 1 liter of solution? b. 2 moles of solute dissolved in 3 liters of solution? c. 6 moles of solute dissolved in 4 liters of solution? d. 4 moles of solute dissolved in 8 liters of solution? 150 22. What is the total number of moles of H2SO4 needed to prepare 5.0 liters of a 2.0 M solution of H2SO4? a. 2. b. 5.0 d. 20 5 c. 10 23. What is the molarity of a KF (aq) solution containing 116 grams of KF in 1.00 liter of solution? a. 1.00 M c. 3.00 M b. 2.00 M d. 4.00 M 24. The solubility of a gas will _?_ when a solution containing the gas is heated and the solubility of a gas in a solution will _?_ when the pressure over the solution is decreased. a. decrease...decrease c. increase...decrease b. decrease...increase d. increase...increase 25. How many grams of potassium nitrate are required to prepare 3.00 x 102 mL of 0.750 M solution? a. 2.28 x 104 g d. 2.4 g b. 84.5 g e. 0.00223 g c. 22.8 g 26. How many grams of sodium chloride are dissolved in 50.0 mL of 1.50 M solution? a. 0.00324 g d. 23.4 g b. 117 g e. 4.38 g 3 c. 4.38 x 10 g 27. A 500 mL sample of a 0.350 M solution is left open on a lab counter for two weeks, after which the concentration of the solution is 0.955 M. What is the new volume of the solution? a. 183 L d. 1.83 mL b. 223 mL e. 0.605 L c. 0.183 L 28. A chemist makes a stock solution of potassium chromate solution by dissolving 97.1 grams of the compound in 1.00 liter of solution. What volume of the solution must be diluted with water in order to prepare 200. mL of 0.200 M solution? a. 80.0 mL d. 0.0800 mL b. 0.150 L e. 120. mL c. 750. mL 29. A 25.0-g sample of sodium hydroxide is dissolved in 400. mL of water. What is the concentration of the solution? a. 0.10 M d. 100. M b. 62.5 M e. 1.56 x 10-3 M c. 1.56 M 30. How many milliliters of 6.0 M HNO3 are needed to prepare 500 mL of 0.50 M HNO3? a. 0.25 mL c. 15.76 mL b. 300 mL d. 40 mL e. None of these are correct 31. How many grams of calcium chloride are needed to prepare 300 mL of a 0.250 M solution? a. 832 g c. 566 g e. 112 g b. 5.66 g d. 8.32 g 32. In a solution of sugar and water, the solvent is the: a. sugar b. water 33. In a solution of sugar and water, the solute is the: a. sugar b. water 34. Gases dissolve best in liquids when: a. the pressure is high and the temperature is low b. the pressure is low and the temperature is low c. the pressure is low and the temperature is high d. the pressure is high and the temperature is high 35. The solubility of potassium nitrate in water at 35 °C is about 60 grams KNO3 per 100 grams of water. How many grams of KNO3 should dissolve in 300 grams of water at 35 °C? a. 180 grams b. 335 grams c. 20 grams 36. Breaking up a solid speeds dissolving in a liquid by: a. decreasing the pressure b. slowing hydration c. raising the temperature d. increasing surface area 37. Most salts become more soluble in water as the: a. temperature is decreased b. pressure is decreased c. pressure is increased d. temperature is increased 38. Calculate the concentration of the following solution in units of molarity, M, moles per liter: 80 grams of NaOH is dissolved in 2 L of solution a. 40 M NaOH b. 82 M NaOH c. 1 M NaOH d. 160 M NaOH 39. How many grams of NaOH are needed to make 100 milliliters of a 0.2 molar solution of NaOH? a. 0.002grams0.8grams20grams required to make 1 liter of 0.4 M d. 800gramsHow many NaCl solution? milliliters of 2 M NaCl solution are a. 5,000 mL800 mL200 mL0.2 42. Under which conditions are gases mLWhich procedure will most soluble in water? increases the solubility of KCl in a. high pressure and high water? temperature b. high pressure and low temperature a. stirring the solute and solvent c. low pressure and high mixture temperature b. increasing the surface area of d. low pressure and low the solute temperature c. raising the temperature of the solvent d. increasing the pressure on the surface of the solvent Unit 11- Acids and Bases Acids and Bases Properties of Acids and Bases Acids taste _________________. Lemon juice and _____________________________, for example, are both aqueous solutions of acids. Acids conduct electricity; they are ________________________. Some are strong electrolytes, while others are _________________ electrolytes. An acetic acid solution, which is a weak electrolyte, contains only a few ions and does not conduct as much current as a strong electrolyte. The bulb is only _____________________ lit. Acids cause certain colored dyes (_________________________) to change color. (Litmus paper turns _______________.) Acids react with metals to form ______________________________ gas. This property explains why acids corrode most metals. Acids react with hydroxides (bases) to form water and a ___________________. Bases taste _________________________ and feel _______________________________. Bases can be strong or weak electrolytes. Naming Acids Acids are compounds that give off _________________________ ions (H+) when dissolved in water. Acids will always contain one or more hydrogen ions next to an __________________________. The anion determines the name of the acid. Naming Binary Acids Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, put the prefix ______________________- and change -ide to -ic acid. Example: HCl The acid contains the hydrogen ion and chloride ion. Begin with the prefix hydro-, name the nonmetallic ion and change -ide to -ic acid. Example: H2S The acid contains the hydrogen ion and sulfide ion. Begin with the prefix hydro- and name the nonmetallic ion. The next step is change -ide to -ic acid, but for sulfur the “ur” is added before -ic. Name the following binary acids. a) HF ___________________________________________ b) H3P __________________________________________ Writing the Formulas for Binary Acids The prefix hydro- lets you know the acid is binary. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the nonmetal. 154 Unit 11- Acids and Bases Example: Hydrobromic acid The acid contains the hydrogen ion and the bromide ion. The two oxidation numbers add together to get zero. The prefix hydro- lets you know the acid is binary. Example: Hydrotelluric acid The acid contains the hydrogen ion and the telluride ion. The two oxidation numbers do NOT add together to get zero, so you must criss-cross. Write the formulas for the following binary acids. a) Hydrocyanic acid _______________ b) Hydroselenic acid _______________ Naming Ternary Acids The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to -_______ acid. Change the suffix -ite to -ous acid The hydro- prefix is NOT used! Example: HNO3 The acid contains the hydrogen ion and nitrate ion. Name the polyatomic ion and change -ate to -ic acid. Example: HNO2 The acid contains the hydrogen ion and nitrite ion. Name the polyatomic ion and change -ite to -ous acid. Example: H3PO4 The acid contains the hydrogen ion and phosphate ion. Name the polyatomic ion and change -ate to -ic acid. Name the following ternary acids. a) H2CO3 ___________________________________________________ b) H2SO4 ___________________________________________________ c) H2CrO4 ___________________________________________________ d) HClO2 ___________________________________________________ Writing the Formulas for Ternary Acids The lack of the prefix hydro- from the name implies the acid is ternary, made of the hydrogen ion and a polyatomic ion. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the polyatomic ion. Example: Acetic acid The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of H+ and the acetate ion. The two charges when added equal zero. Example: Sulfurous acid Again the lack of the prefix hydro- implies the acid is ternary, made of the hydrogen ion and a polyatomic ion. The polyatomic ion must end in –ite since the acid ends in -ous. The acid is made of H+ and the sulfite ion. The two charges when added do not equal zero, so you must crisscross the oxidation numbers. Write the formulas for the following ternary acids. a) perchloric acid ______________________ b) iodic acid _____________________ c) nitrous acid _______________________ 155 d) bromic acid ___________________ Unit 11- Acids and Bases Types of Acids and Bases Arrhenius Definitions - The simplest definition is that an acid is a substance that produces _____________________________ ions when it dissolves in water. A hydronium ion, H3O+, consists of a hydrogen ion attached to a __________________ molecule. A hydronium ion, H3O+, is equivalent to H+. HCl and H3PO4 are acids according to Arrhenius. A base is a substance that produces ________________________ ions, OH–, when it dissolves in water. Ca(OH)2 and NaOH are Arrhenius bases. NH3, ammonia, could not be an Arrhenius ___________________. Monoprotic acids have only ____________ ionizable hydrogen. Some acids have more than one ionizable hydrogen and are called ______________________________ acids. Bronsted-Lowry Definitions - An Bronsted-Lowry acid is a ________________________ (H+) donor. HBr and H2SO4 are Bronsted-Lowry acids. When a Bronsted-Lowry acid dissolves in water it gives its proton to water. HCl (g) + H2O (l) ↔ H3O+ + Cl- A Bronsted-Lowry base is a proton acceptor. B + H2O ↔ BH+ + OH- A Brønsted-Lowry base does not need to contain OH-. Consider HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) HCl donates a proton to water. Therefore, HCl is an _______________. H2O accepts a proton from HCl. Therefore, H2O is a ______________. Identify the acid and base in the following reactions. a) H2SO3 + H2O ↔ HSO3- + H3O+ Acid _____________________________ base _________________________ b) NH3 + H2SO4 ↔ NH4+ + HSO4Acid _____________________________ base _________________________ Molarity and Dilution The concentration of a solution is the amount of solute present in a given quantity of solution. _________________________ is the number of moles of solute in 1 liter of solution. moles solute Molarity = -------------------------liters of solution The procedure for preparing a less concentrated solution from a more concentrated one is called a ___________________________. M1 V1 = M2 V2 156 Unit 11- Acids and Bases PRACTICE: What is the molarity of an acetic acid (HC2H3O2) solution with 4.0 moles dissolved in 250 mL of solution? How many moles of hydrochloric acid (HCl) are needed to make 3.0 L of a 0.55 M HCl solution? 0.600 moles of the base sodium hydroxide (NaOH) are dissolved in a small amount of water then diluted to 500. mL. What is the concentration? 3.25 moles of the base potassium hydroxide (KOH) are dissolved in a small amount of water then diluted to 725 mL. What is the concentration? How many moles are in 2.00 L of a 6.00 M solution of sulfuric acid (H2SO4)? How many moles are in 1250 mL of a 3.60 M solution of nitric acid (HNO3)? 6.0 L of a 1.55 M LiOH solution are diluted to 8.8 L. What is the new molarity of the lithium hydroxide solution? You have 250 mL of 6.0 M HCl. How many milliliters of 1.2 M HCl can you make? 4.0 liters of a 0.75 M solution of sulfuric acid (H2SO4) are diluted to a 0.30 M solution. What is the final volume? You need 350 mL of 0.25 M NaOH. All you have available is a 2.0 M stock solution of NaOH. How do you make the required solution? Strength of Acids and Bases The strength of a base is based on the percent of units ___________________________________, not the number of OH– ions produced. The strength of a base does NOT depend on the _____________________________. 1A and _______ hydroxides, excluding __________, are strong bases. Some bases, such as Mg(OH)2, are not very soluble in water, and they don’t produce a large number of OH– ions. However, they are still considered to be strong bases because all of the base that does dissolve completely dissociates. The strength of an acid is based on the percent of units dissociated, not the number of ____________ ions produced. The strength of an acid does NOT depend on the _______________________________. There are 6 strong acids: HCl, HBr, HI, HClO4, HNO3, and H2SO4. Strong acids and bases are strong __________________________________ because they dissociate completely. Electrolytes conduct ______________________________. 157 Unit 11- Acids and Bases Weak acids and bases don’t completely ionize, so they are weak electrolytes. Although the terms weak and strong are used to compare the _____________________________ of acids and bases, dilute and concentrated are terms used to describe the _____________________________ of solutions. pH Scale Water ionizes; it falls apart into _________________. H2O H+ + OH- The preceding reaction is called the _____________________________________ of water. [H+ ] = [OH-] = 1 x 10-7 M When [H+ ] = [OH-], the solution is _________________________. At 25°C, Kw = [H+] [OH-] = 1 x 10-14 Kw is called the ion-product constant. If [H+] > 10-7 then [OH-] < 10-7. The solution is ______________________ when [H+] > [OH-]. If [H+] < 10-7 then [OH-] > 10-7. The solution is __________________________ when [OH-] > [H+]. In most applications, the observed range of possible hydronium or hydroxide ion concentrations spans 10–14 M to ______M. To make this range of possible concentrations easier to work with, the pH scale was developed. pH is a mathematical scale in which the concentration of hydronium ions in a solution is expressed as a number from _________ to __________. pH meters are instruments that measure the exact pH of a solution. Indicators register different colors at different pH’s. In neutral solution, pH = 7. In an acidic solution, pH < 7. In a basic solution, pH > 7. As the pH drops from 7, the solution becomes more acidic. As pH increases from 7, the solution becomes more basic. The pH of a solution equals the negative logarithm of the hydrogen ion concentration. pH = - log [H+] Chemists have also defined a pOH scale to express the basicity of a solution. pOH = - log [OH-] If either pH or pOH is known, the other may be determined by using the following relationship. pH + pOH = 14.00 Find the pH of the following solutions. a) The hydronium ion concentration equals: 10–2 M. pH = _________________ b) The hydronium ion concentration equals: 10–11 M. pH = _________________ c) The hydronium ion concentration equals: 1 x 10–6 M. pH = _________________ d) The hydroxide ion concentration equals: 10–8 M. pH = _________________ e) The hydroxide ion concentration equals: 10–5 M. pH = _________________ f) The hydroxide ion concentration equals: 10–3 M. pH = _________________ If a certain carbonated soft drink has a hydrogen ion concentration of 1.0 x 10–4 M, what are the pH and pOH of the soft drink? 158 Unit 11- Acids and Bases Calculating Ion Concentrations From pH If either pH or pOH is known, the hydrogen ion or hydroxide ion can be found. [H+] =10-pH [OH-] =10-pOH On the calculator, hit 2nd log (-) and then the number. Find the [H+] of a solution that has a pH equal to 6. Find the [H+] of a solution that has a pH equal to 12. Find the [H+] of a solution that has a pH equal to 5. Find the [H+] of a solution that has a pOH equal to 6. Find the [H+] of a solution that has a pOH equal to 6. Find the [H+] of a solution that has a pOH equal to 2. Find the [H+] of a solution that has a pOH equal to 4. Find the [OH-] of a solution that has a pH equal to 10. Calculating Ion Concentration From Ion Concentration If either [H+] or [OH-] is known, the hydrogen ion or hydroxide ion can be found. [H+] [OH-] = 1 x 10-14 Find the hydrogen ion concentration if the hydroxide ion concentration equals: 1 x 10–8 M. Find the hydrogen ion concentration if the hydroxide ion concentration equals: 1 x 10–2 M. Find the hydroxide ion concentration if the hydrogen ion concentration equals: 1 x 10–4 M. Find the hydroxide ion concentration if the hydrogen ion concentration equals: 1 x 10–9 M. Indicators Chemical _____________________ whose colors are affected by acidic and basic solutions are called indicators. Many indicators do not have a sharp color change as a function of ____________. Most indicators tend to be __________________ in more acidic solutions. 159 Unit 11- Acids and Bases Which indicator is best to show an equivalence point pH of 4? Which indicator is best to show an equivalence point pH of 11? Which indicator is best to show an equivalence point pH of 2? Neutralization Reactions The reaction of an acid and a base is called a neutralization reaction. Acid + base salt + water A salt is an ___________________ compound. Predict the products of and balance the following neutralization reactions. (Remember to check the oxidation numbers of the ions in the salt produced.) a) HNO3 + KOH The salt is composed of the ________________ ion and the _________________ ion. b) HCl + Mg(OH)2 c) H2SO4 + NaOH Neutralization How many moles of HNO3 are needed to neutralize 0.86 moles of KOH? How many moles of HCl are needed to neutralize 3.5 moles of Mg(OH)2? How many moles of H3PO4 are needed to neutralize 3.5 moles of Mg(OH)2? How many moles of HC2H3O2 are needed to neutralize 3.5 moles of Cr(OH)3? 160 Unit 11- Acids and Bases If it takes 87 mL of an HCl solution to neutralize 0.67 moles of Mg(OH)2 what is the concentration of the HCl solution? If it takes 58 mL of an H2SO4 solution to neutralize 0.34 moles of NaOH what is the concentration of the H2SO4 solution? If it takes 85 mL of an HNO3 solution to neutralize 0.54 moles of Mg(OH)2 what is the concentration of the HNO3 solution? If it takes 150. mL of an Ca(OH)2 solution to neutralize 0.800 moles of HCl what is the concentration of the Ca(OH)2 solution? Acid Rain Acid Rain is any rain with a pH less than __________. Pure rain is naturally acidic because of dissolved _______________. It is caused by the man-made oxides of _______________________ and nitrogen. SO3 + H2O H2SO4 Research shows acid rain is associated with parts of a country where heavy ______________________________ are situated and also down-wind from such sites. Analysis of acid rain indicates that especially sulfur oxides, SOx and nitrogen oxides, NOx are mostly responsible from rain acidity. Snow _________________, sleet, hail and drizzle all become contaminated with acids when SOx and NOx are present as pollutants. Titration The known reactant molarity is used to find the unknown _________________________ of the other solution. Solutions of known molarity that are used in this fashion are called _________________________ solutions. In a titration, the molarity of one of the reactants, acid or base, is known, but the other is unknown. A 15.0-mL sample of a solution of H2SO4 with an unknown molarity is titrated with 32.4 mL of 0.145M NaOH to the bromothymol blue endpoint. Based upon this titration, what is the molarity of the sulfuric acid solution? First find the number of moles of the solution for which you know the molarity and volume. Next, use the mole-mole ratio to determine the moles of the unknown. Finally, determine the molarity of the unknown solution. If it takes 45 mL of a 1.0 M NaOH solution to neutralize 57 mL of HCl, what is the concentration of the HCl ? If it takes 67.0 mL of 0.500 M H2SO4 to neutralize 15.0 mL of Al(OH)3 what was the concentration of the Al(OH)3 ? 161 Unit 11- Acids and Bases How many moles of 0.275 M HCl will be needed to neutralize 25.0 mL of 0.154 M NaOH? Titration Curves A plot of ___________ versus volume of acid (or base) added is called a titration curve. Strong Base-Strong Acid Titration Curve Consider adding a strong base (e.g. NaOH) to a solution of a strong acid (e.g. HCl). Before any base is added, the pH is given by the strong _________________ solution. Therefore, pH ____ 7. When base is added, before the equivalence point, the pH is given by the amount of strong acid in _________________________. Therefore, pH < 7. At ________________________________ point, the amount of base added is stoichiometrically equivalent to the amount of acid originally present. Therefore, pH =_________. To detect the equivalent point, we use an indicator that changes ____________________somewhere near 7.00. Past the equivalence point all acid has been consumed. Thus one need only account for excess __________________. Therefore, pH ______ 7. 162 Unit 11- Acids and Bases Homework / Practice Show all work for full credit. 1. Name the following compounds as acids. a. H2SO4 c. H2S b. H2SO3 d. HClO4 2. Write formulas for the following acids. a. nitrous acid c. phosphoric acid b. hydrobromic acid 3. Use an activity series to identify two metals that will not generate hydrogen gas when treated with an acid. 4. Write balanced molecular equations for the reactions of acids and bases. a. aluminum metal with dilute nitric acid b. calcium hydroxide solution with acetic acid 5. Write the equation that represents the following reaction: the ionization of HClO3 in water 6. Explain how strong acid solutions conduct an electric current. 7. CaCO3 (s) + HCl (aq) CaCl2 (aq) + H2O (l) + CO2 (g) a. Balance the above equation. b. How many liters of CO2 form at STP if 5.0 g of calcium carbonate are treated with excess hydrochloric acid? 8. Consider the following reaction: NH4+ (aq) + CO3-2 (aq) ↔ NH3 (aq) + HCO3-1 (aq) a. What reactant serves as the base? b. What reactant serves as the acid? 9. Given the following reaction: HCO3-1 (aq) + OH-1 (aq) ↔ CO3-2 (aq) + H2O (l) a. What reactant serves as the base? 163 Unit 11- Acids and Bases b. What reactant serves as the acid? 10. Write the formula for the salt formed in each of the following neutralization reactions. a. hydrobromic acid combined with barium hydroxide b. lithium hydroxide combined with sulfuric acid 11. : H2SO4 (aq) + NaOH (aq) Na2SO4 (aq) + H2O (l) a. Balance the above neutralization equation b. In order to completely consume all reactants, what should be the mole ratio of acid to base? 12. Consider the reaction represented by the following incomplete equation: Ba(OH)2 (aq) + H2SO4 (aq) a. Predict the products of this reaction, and write the balanced equation. b. Use the solubility rules to determine the solubility of the salt produced in the reaction. c. If 0.030 mol of Ba(OH)2 is consumed, how many grams of water are produced? 13. Name the acid that is present in vinegar. 14. Perform the following calculations. a. If the hydronium concentration is 1 x 10-6 M for a solution, calculate the hydroxide concentration. b. If the hydroxide concentration is 1 x 10-12 M for a solution, calculate the hydronium concentration. c. If the pOH = 4.00 for a solution, calculate the pH. Is the solution acidic or basic? d. If the hydronium concentration is 1.00 x 10-3M, calculate the pOH. e. If the pOH = 5.0 for a solution, calculate the hydroxide concentration. f. If the pH = 12.0 for a solution, calculate the hydronium concentration. g. If the pH = 3.00 for a solution, calculate the hydroxide concentration. h. If the hydronium concentration = 1.0 x 10-8 M for a solution, calculate the hydroxide concentration. 164 Unit 11- Acids and Bases 15. Summarize the two main acid-base theories in the table below. ACID BASE Arrhenius Brønsted-Lowry 16. Label the acid (A), base (B), conjugate acid (CA), and conjugate base (CB) in each of the following reactions. a. H2SO4 + NH3 HSO4 + NH4 b. HC2H3O2 + H2O H3O+ + C2H3O2 c. NaHCO3 + HCl NaCl + H2CO3 17. Give the conjugate base for each of the following Brønsted-Lowry acids. a. HI c. H2CO3 b. NH4+ d. HNO3 18. Give the conjugate acid for each of the following Brønsted-Lowry bases. a. CN– c. CH3COO– b. O2– d. NH3 19. Find [OH ] for 1.0 × 10-12M HClO4. 20. What is the pH of 1.0 × 10-4M HCl? 21. What is the pH of 1.5 × 10-3M NaOH? 22. A solution of HNO3 has a pH of 4.0. What is the molarity of HNO3? 23. What is the molarity of KOH in a solution that has a pH of 10.0? 165 Unit 11- Acids and Bases Acids and Bases Practice Test Directions: Give the names or formulas for the following acids, bases, and salts: 1. KOH________________________ 3. Sulfuric Acid__________________ 2. HNO3 _______________________ 4. Magnesium hydroxide ___________ Directions:. 5. In complete sentences, define an acid according to the Arrhenius theory. Directions: Label (according to Bronsted-Lowry) the Bronsted-Lowry acid, Bronsted-Lowry base, conjugate acid, and conjugate base in each of the equations below: 6. H₂O + HC2H3O2 ⇆ H₃O⁺ + C2H3O2⁻ 7. CN-- + H₃O⁺ ⇆ H₂O + HCN Directions Identify the following as an acid or a base, strong or weak. a.Acid or base b.Strong or weak 8. 2 M KOH __________ _____________ 9. 7 M H₂SO₄ __________ _____________ 10. 0.12 M H2S _________ _____________ Directions: Complete and balance the following neutralization reactions. 11. NaOH + HCl → _______________ + __________________ 12. H₂SO₄ + KOH → _____________ + __________________ 13. Determine the pOH for a solution of HNO3 that has a concentration of 0.01M. 14. Determine the pH for a solution of CuOH that has an [OH-] of 0.000001M. Directions: Complete the following chart. pH [H₃O⁺]or [H+] 15. 16. [OH⁻] pOH 1 x 10⁻6 2 17. 4 166 Acidic or basic Unit 11- Acids and Bases Multiple Choice Practice 18. According to the Arrhenius theory, a base yields a. H+ as the only positive ion in an aqueous solution b. OH+ as the only positive ion in an aqueous solution c. OH- as the only negative ion in an aqueous solution d. H- as the only negative ion in an aqueous solution 19. In the reaction H2SO4(aq) --> 2H+(aq) + SO4-2(aq) H2SO4 is a(n) a. Arrhenius acid c. Salt b. Arrhenius base 20. Arrhenius acids yield a. OH- as the only negative ion in an aqueous solution b. H- as the only negative ion in an aqueous solution c. H3O+ as the only positive ion in an aqueous solution d. OH+ as the only positive ion in an aqueous solution 21. Which of the following is an Arrhenius base? a. HCl b. H2SO4 c. NaCl d. NH3 22. Which of the following is a salt? a. HOH b. NH4NO3 c. HCl d. H2CO3 23. A substance that conducts an electrical current when dissolved in water is called a. an acid c. an ionic compound b. an electrolyte d. a nonelectrolyte 24. Which of the following can conduct an electric current? a. Mg(OH)2(s) c. NaOH(aq) b. H2O(s) d. NH4Cl(s) 25. Which electrolyte is best at conducting electricity when dissolved in an aqueous solution? a. KCl(s) c. CaCl2(s) b. Na2SO4(s) d. H3PO4(s) 26. A student dissolved NaCl(s) in water, and tested with a battery, wire, and a light blub to see if it conducted an electric current. The solution conducted an electric current. This is because NaCl(s) is a. a salt and an electrolyte b. a salt and a nonelectrolyte c. a Arrhenius acid and an electrolyte d. a Arrhenius base and an electrolyte 167 Unit 11- Acids and Bases 27. In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form which of the following? a. Water only c. Water and Carbon dioxide b. Salt and Carbon dioxide d. Water and Salt 28. What type of chemical reaction is neutralization? a. single replacement b. double replacement c. synthesis d. decomposition 29. In the process of neutralization a salt and a base react to yield water and an acid. a. True b. False 30. Which of the following can be a product of neutralization? a. LiOH(s) c. MgI2(l) b. HCl(l) d. NaCl(aq) 31. Which of the following reactants will represent a neutralization? a. BaCl2 + CaSO4 c. Ca(OH)2 + H2SO4 b. HCl + F d. NaCl + H2O 32. When NaOH and HCl react, what will be on the product side? a. only NaCl c. NaCl and HOH b. only HOH d. NaCl and Cl2 33. A titration reaction involves a complete neutralization reaction where the moles of H+ equal the moles of OH-. a. True b. False 34. Titration is a process in which a. which a volume of solution of unknown concentration is used to determine the concentration of another solution. b. which a volume of solution of known concentration is used to determine the volume of another solution. c. which a volume of solution of known concentration is used to determine the concentration of another solution. d. which a volume of solution of known concentration is used to determine the curve of another solution. 35. In MaVa=MbVb, what is Mb? a. molarity of OHb. molarity of H+ c. molarity of Hd. molarity of O2 36. What is the molarity of HCl(aq) if 25 mL of 8.0M NaOH(aq) neutralizes exactly 20.0 mL of HCl(aq)? a. 5 b. 10M d. 20M M c. 15M 168 Unit 11- Acids and Bases 37. At the end point of titration, what is the relationship between moles of H+ and OH-? a. the moles of H+ are greater than OHb. the moles of OH- are greater than H+ c. the moles of H+ are equal to moles of OHd. there is no relationship between moles of H+ and OH38. The molarity of HBr(aq) is 2M when 10 milliliters of 8M NaOH(aq) neutralizes exactly 20 milliliters of HBr. a. True b. False 39. What is the molarity of NaOH if 5. milliliters of 4M HCl(aq) neutralizes exactly 10. mL of NaOH(aq)? a. .5M c. 1.5M b. 1M d. 2M 40. If 10.milliliters of a 0.40M HBr solution is required to neutralize exactly 0.2M of NaOH, what is the volume of the base? a. 10ml c. 30ml b. 20ml d. 40ml 41. The molarity of an acid can be calculated if a base of known concentration (standard base) is added, drop by drop, to a specific volume of the acid until the indicator changes color. a. True b. False 42. One acid-base theory states that an acid is an H+ a. Acceptor b. Eliminator c. Dissolver d. Donor 43. According to the Bronsted-Lowry acid-base theory, a base is a substance that can a. donate an electron c. donate a proton b. accept a proton d. accept a electron 44. In the following reaction, NH3 + HCl --> NH4+ + Cl- NH3 acts as a(n) a. base in the reverse reaction. c. base in the forward reaction. b. acid in the forward reaction. d. acid in the reverse reaction. 45. The acidity or alkalinity of a solution can be measured by its pH value. a. True b. False 46. The relative level of acidity or alkalinity of a solution can be shown by using their pH values. a. True b. False 169 Unit 11- Acids and Bases 47. Which of these 1 M solutions will have the highest pH? a. H3PO4 c. NaCl b. HCl d. NaOH 48. Which pH indicates an acidic solution? a. 1 b. 7 c. 9 49. Which of these pH numbers indicates the lowest level of acidity? a. 1 b. 3 c. 8 50. Which formula represents a salt? a. KOH b. KCl d. 12 d. 12 c. CH3OH d. CH3COOH 51. Which substance can be classified as an Arrhenius acid? a. HCl c. LiOH b. NaCl d. KOH 52. Which solution will change red litmus to blue? a. HCl(aq) b. NaCl(aq) c. CH3OH(aq) d. NaOH(aq) 53. An acidic solution could have a pH of a. 7 b. 10 c. 3 d. 14 54. What is the pH of a 0.00001 molar HCl solution? a. 1 b. 9 c. 5 d. 4 55. What is the pH of a solution with a hydronium ion concentration of 0.01 moles per liter? a. 1 b. 2 c. 10 d. 14 56. Which 0.1 M solution will turn phenolphthalein pink? a. HBr(aq) c. LiOH(aq) b. CO2(aq) d. CH3OH(aq) 57. Given the equation: H+ + OH- <-> H2O Which type of reaction does the equation represent? a. esterification b. decomposition c. hydrolysis d. neutralization 58. As the hydrogen ion concentration of an aqueous solution increases, the hydroxide ion concentration of this solution will a. decrease c. remain the same b. increase 170 Unit 11- Acids and Bases 59. A student wishes to prepare approximately 100 milliliters of an aqueous solution of 6M HCl using 12 M HCl. Which procedure is correct? a. adding 50 mL of 12 M HCl to 50 mL of water while stirring the mixture steadily. b. adding 50 mL of 12 M HCl to 50 mL of water and then stirring the mixture steadily. c. adding 50 mL of water to 50 mL of 12 M HCl while stirring the mixture steadily. d. adding 50 mL of water to 50 mL of 12 M HCl and then stirring the mixture steadily. 60. The reaction of an acid like HCl and a base like NaOH always a. Forms a precipitate c. Forms a salt and water b. Forms a volatile product d. Forms a sulfate and water 61. What is the pH of an acetic acid solution if the [H3O+] = 1x10-4 mol/L? a. 1 c. 3 b. 2 d. 4 171 e. 5 Unit 12- Kinetics and Thermochemistry REACTION KINETICS Energy Diagrams Reactants always start a reaction so they are on the _________________ side of the diagram. Products are on the right. The exothermic reaction gives off ___________________ because the products are at a lower energy level than the reactants. In an exothermic graph, the reactants have _____________________ energy than the products. The change in energy is a _________________________ value. The endothermic reaction absorbs heat because the products are at a _________________________ energy level than the reactants. In an endothermic graph, the products have _____________________ energy than the reactants. The change in energy is a _______________________ value. Scientists have observed that the energy released in the formation of a compound from its elements is always identical to the energy required to ______________________ that compound into its elements. _____________________________ energy is the minimum amount of energy that reacting particles must have to form the activated complex. The activated complex is a short-lived, _________________ arrangement of atoms that may break apart and re-form the reactants or 172 Unit 12- Kinetics and Thermochemistry may form products. To calculate the activation energy, ______________________ the energy of the reactants from the energy at the top of the peak. The enthalpy or heat of reaction (ΔH) is the amount of ___________________ released or absorbed in the reaction. To determine ΔH, take the energy of the products and _______________________ the energy of the reactants. The heat content of the reactants of the forward reaction is about ________ kilojoules. The heat content of the products of the forward reaction is about ________ kilojoules. The heat content of the activated complex of the forward reaction is about _______ kilojoules. The activation energy of the forward reaction is about _______ kilojoules. The heat of reaction (ΔH) of the forward reaction is about _______ kilojoules. The forward reaction is (endothermic or exothermic). The heat content of the reactants of the reverse reaction is about ________ kilojoules. The heat content of the products of the reverse reaction is about ________ kilojoules. The heat content of the activated complex of the reverse reaction is about _______ kilojoules. The activation energy of the reverse reaction is about _______ kilojoules. The heat of reaction (ΔH) of the reverse reaction is about _______ kilojoules. The reverse reaction is (endothermic or exothermic). 173 Unit 12- Kinetics and Thermochemistry The activation energy can be lowered by adding a __________________________. The catalyst _________________ the activation energy by providing an alternate pathway for the reaction to occur. Expressing Reaction Rates We generally define the average _____________________ of an action or process to be the change in a given quantity during a specific period of time. Reaction rates cannot be calculated from balanced equations as stoichiometric amounts can. Reaction rates are determined experimentally by measuring the _____________________________ of reactants and/or products in an actual chemical reaction. Collision Theory According to the collision theory, atoms, ions, and molecules must collide with each other in order to react. The following three statements summarize the collision theory. 1. Particles must ________________________ in order to react. 2. The particles must collide with the correct _______________________________. 3. The particles must collide with enough ________________________ to form an unstable activated complex, also called a _____________________________ state, which is an intermediate particle made up of the joined reactants. The _____________________________ amount of energy that colliding particles must have in order to form an activated complex is called the activation energy of the reaction. Particles that collide with energy less than the activation energy ____________________________ form an activated complex. In an exothermic reaction, molecules collide with enough energy to overcome the activation energy barrier, form an activated complex, then __________________________ energy and form products at a lower energy level. In the reverse endothermic reaction, the reactant molecules lying at a _______________ energy level must absorb energy to overcome the activation energy barrier and form high-energy products. Factors Affecting Reaction Rates The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction. 174 Unit 12- Kinetics and Thermochemistry 1) An important factor that affects the rate of a chemical reaction is the reactive nature of the reactants. As you know, some substances react more readily than others. The more reactive a substance is, the _____________________________ the reaction rate. 2) Another important factor that affects the rate of a chemical reaction is the concentration of the reactants. Reactions ______________________ _________________ when the concentrations of reacting particles are increased. Increasing the number of reactants increases probability of collisions. The rate of gaseous reactions can be ___________ by pumping more gas into the reaction container. 3) Surface area of the reactants affects the rate of a chemical reaction. ______________________ the surface area of reactants provides more opportunity for collisions with other reactants, thereby increasing the reaction rate. 4) Temperature affects the rate of a chemical reaction. Generally, increasing the temperature at which a reaction occurs _________________________ the reaction rate. Raising the temperature raises both the collision frequency and the collision energy. 5) Adding a catalyst affects the rate of a chemical reaction. A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed in the reaction. In fact, catalysts are not included in the chemical equation. 6) Compressing gases affects the rate of a chemical reaction. When two gases react, compressing the gases generally ___________ the rate of reaction. REACTION ENERGY Energy is the ability to do___________________ or produce heat. It exists in two basic forms, potential energy and ______________________ energy. Potential energy is energy due to the ______________________ or position of an object. Kinetic energy is energy of ________________. The potential energy of the dammed water is converted to kinetic energy as the dam gates are opened and the water flows out. Chemical systems contain _________________ kinetic energy and potential energy. As temperature increases, the motion 175 Unit 12- Kinetics and Thermochemistry of submicroscopic particles ______________________, so its kinetic energy __________________________. The potential energy of a substance depends upon its composition: the type of atoms in the substance, the number and type of chemical bonds joining the atoms, and the particular way the atoms are arranged. Law of Conservation of Energy and Heat The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor ________________________. Heat, which is represented by the symbol ____, is energy that is in the process of flowing from a _____________________ object to a cooler object. The SI unit of heat and energy is the joule (J). Heat involves a transfer of energy between 2 objects due to a ________________________ difference. When the warmer object loses heat, its temperature decreases and q is _________________________. When the cooler object absorbs heat, its temperature ________________ and q is positive. The specific heat of any substance is the amount of heat required to raise the temperature of ____ gram of that substance by one degree Celsius. Because different substances have different compositions, each substance has its own specific heat. q = m Cp ∆T q = heat (J); m = mass (g); Cp = specific heat (J/(g.°C); ∆T = change in temperature = Tf – Ti (°C) Exothermic: Heat flows _________ of the system (to the surroundings). The value of ‘q’ is negative. Endothermic: Heat flows _________ the system (from the surroundings). The value of ‘q’ is positive. The temperature of a sample of iron with a mass of 10.0 g changed from 50.4°C to 25.0°C with the release of 114 J heat. What is the specific heat of iron? A piece of metal absorbs 256 J of heat when its temperature increases by 182°C. If the specific heat of the metal is 0.301 J/g.°C, determine the mass of the metal. If 335 g water at 65.5°C loses 9750 J of heat, what is the final temperature of the water? The specific heat of water is 4.18 J/g.°C. As 335 g of aluminum at 65.5°C gains heat, its final temperature is 300.°C. The specific heat of aluminum is 0.897 J/g.°C. Determine the energy gained by the aluminum. 176 Unit 12- Kinetics and Thermochemistry Heat changes that occur during chemical and physical processes can be measured accurately and precisely using a ___________________________. A calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process. A coffee-cup calorimeter made of ________ Styrofoam cups. Suppose you put 125 g of water into a foam-cup calorimeter and find that its initial temperature is 25.6°C. Then, you heat a 50.0 g sample of the unknown metal to a temperature of 115.0°C and put the metal sample into the water. Both water and metal have attained a final temperature of 29.3°C. Heat flows from the hot metal to the cooler water and the temperature of the water rises. The flow of heat stops only when the temperature of the metal and the water is equal. Assuming no heat is lost to the surroundings, the heat gained by the water is equal to the heat lost by the metal. Determine the specific heat of the metal. You put 352 g of water into a foam-cup calorimeter and find that its initial temperature is 22.0°C. What mass of 134°C lead, Clead = 0.129 J/g°C, can be placed in the water so that the equilibrium temperature is 26.5°C? You put water into a foam-cup calorimeter and find that its initial temperature is 25.0°C. What is the mass of the water if 14.0 grams of 125°C nickel, CNi = 0.444 J/g°C, can be placed in the water so that the equilibrium temperature is 27.5°C? Phase Changes Review Solid → liquid ________________________ Liquid → solid ___________________ Liquid → gas ________________________ Gas → liquid _____________________ Solid → gas ________________________ Gas → solid ______________________ Energy and Phase Changes q = m Hf q = m Hv Hf = latent heat of fusion (J/g) ; Hv = latent heat of vaporization (J/g) Heating Curve for Water 120 The heating curve has _____ distinct Steam Water and Steam 100 regions. The ______________________ 80 lines are where phase changes occur. 60 Water 40 Temperature is ___________________ 20 0 Ice during a phase change! Water and Ice -20 0 40 120 220 760 800 177 Unit 12- Kinetics and Thermochemistry Heat of vaporization (Hv) is the energy required to change one gram of a substance from ________________ to gas. Heat of fusion (Hf) is the energy required to change one gram of a substance from __________________ to liquid. How much heat does it take to melt 12.0 g of ice at 0 °C? Hf for water is 334 J/g. How much heat must be removed to condense 5.00 g of steam at 100 °C? Hv = 2260 J/g. Three equations can be used in calculating energy. q = m Cp ΔT q = m Hf q = m Hv Solving Problems The total heat equals the sum of all the heats you have to use. Go in the following order when energy is being added to the system. 1) Heat ice q = m Cice ∆T 2) Melt ice q = m Hf 3) Heat water q = m Cwater ∆T 4) Boil water q = m Hv 5) Heat steam q = m Csteam ∆T Numbers Needed For Energy Problems Involving Water (look up in reference tables) For ice, specific heat = __________ For steam, specific heat = _____________ Heat of fusion = _____________ For water, specific heat = _________ Heat of vaporization = ___________ How much heat does it take to heat 12 g of ice at – 6 °C to 25 °C water? Round to a whole number. How much heat does it take to heat 35 g of ice at 0 °C to steam at 150 °C? Round to a whole number. How much heat does it take to convert 16.0 g of ice to water at 0 °C? How much heat does it take to heat 21.0 g of water at 12.0°C to water at 75.0°C? How much heat does it take to heat 14.0 g of water at 12.0°C to steam at 122.0°C? Entropy Entropy (S) is a measure of the ______________________ or randomness of the particles that make up a system. Spontaneous processes always result in a(n) _____________________ in the entropy of the universe. Entropy of a solid < Entropy of a liquid << Entropy of a gas A solid has an orderly arrangement. A liquid has the molecules next to each other. A gas has molecules moving all over the place. 178 Unit 12- Kinetics and Thermochemistry Several factors affect the change in entropy of a system. 1. Changes of state. Entropy ___________________ when a solid changes to a liquid and when a liquid changes to a gas because these changes of state result in freer movement of the particles. 2. Dissolving of a gas in a solvent. When a gas is dissolved in a liquid or solid solvent, the motion and randomness of the particles are limited and the entropy of the gas _____________. 3. Change in the number of gaseous particles. When the number of gaseous particles increases, the entropy of the system usually ___________________ because more random arrangements are possible. 4. Dissolving of a solid or liquid to form a solution. When solute particles become dispersed in a solvent, the disorder of the particles and the entropy of the system usually ________________. 5. Change in temperature. A temperature increase results in increased disorder of the particles and a(n) ____________________ in entropy. Predict the sign of ∆Ssystem for: O2 (g) O2 (aq) Predict the sign of ∆Ssystem for: C6H6 (s) C6H6 (l) Predict the sign of ∆Ssystem for: C (s) + CO2 (g) 2 CO (g) Homework / Practice 1. If 200. g of water at 20.0 °C absorbs 41840 J of heat, what will its final temperature be? 2. Aluminum has a specific heat of 0.900 J/g.°C. How much energy is needed to raise the temperature of a 625 g block of aluminum from 30.7 °C to 82.1 °C? 3. The products in a reaction have a total heat content of 458 kJ and the reactants have a total heat content of 656 kJ. What is the value of ∆H? 179 Unit 12- Kinetics and Thermochemistry 4. If a reaction is exothermic, are the products or the reactants more stable? 5. The heat of combustion of ethene gas, C2H2, is - 1301.1 kJ/mol of ethene. (a) If 0.250 mol of ethene react, how much energy is released? (b) How many grams of ethene are needed to react to release 3900. kJ of heat? (c) Write a balanced reaction for the combustion of ethene. 6. For each of the following examples, state whether the change in entropy is positive, negative or remains the same. (a) HCl (l) → HCl (g) (b) C6H12O6 (aq) → C6H12O6 (s) (c) 2 NH3 (g) → N2 (g) + 3 H2 (g) (d) 3 C2H4 (g) → C6H12 (l) 7. List 4 factors that can speed up a chemical reaction. 8. Sketch on the graph above the changes when a catalyst is added. 9. Consider the following equilibrium equation: H2O (g) + C (s) → H2 (g) + CO (g) + heat energy. Will the reaction rate increase or decrease when each of the following occurs? (a) extra CO gas is introduced (b) a catalyst is introduced (c) the temperature of the system is lowered (d) the pressure of the system is increased 180 Unit 12- Kinetics and Thermochemistry 10.Convert from one unit to the other: (a) (b) (c) (d) (e) (f) (g) (h) 1.69 Joules to calories 0.3587 J to cal 820.1 J to kilocalories 68 calories to kilocalories 423 calories to kilocalories 20.0 calories to Joules 252 cal to J 2.45 kilocalories to calories (i) 556 kilocalories to cal (j) 6.78 kilocalories to kilojoules (k) 59.6 calories to kcal (l) 449.6 joules to kilojoules (m) 9.806 kJ to J (n) 5.567 cal to J (o) 5467.9 kcal to J 11. How much energy must be absorbed by 20.0 g of water to increase its temperature from 283.0 °C to 303.0 °C? 12. When 15.0 g of steam drops in temperature from 275.0 °C to 250.0 °C, how much heat energy is released? 13. How much energy is required to heat 120.0 g of water from 2.0 °C to 24.0 °C? 14. If 720.0 g of steam at 400.0 °C absorbs 800.0 kJ of heat energy, what will be its increase in temperature? 15. How much heat (in kJ) is given out when 85.0 g of lead cools from 200.0 °C to 10.0 °C? (Cp of lead = 0.129 J/g °C) 16. If it takes 41.72 joules to heat a piece of gold weighing 18.69 g from 10.0 °C to 27.0 °C, what is the specific heat of the gold? 17. A certain mass of water was heated with 41,840 Joules, raising its temperature from 22.0 °C to 28.5 °C. Find the mass of water. 181 Unit 12- Kinetics and Thermochemistry 18. Determine the energy required (in Joules) when the temperature of 3.21 grams of liquid water increases by 4.0 °C. 19. Determine the temperature change when: (a) 20.0 g of water is heated from 16.8 °C to 39.2 °C. (b) 35.0 g of water is cooled from 56.5 °C to 5.9 °C. (c) 50.0 g liquid water is heated from 0.0 °C to 100.0 °C. (d) 25.0 g of ice is warmed from -25.0 °C to 0.0 °C, but does not melt. (e) 30.0 g of steam heats from 373.2 K to 405.0 K. 20. Determine the energy needed (in Joules) when 55.6 grams of water at 43.2 °C is heated to 78.1°C. 21. Determine the energy required (in kilojoules) when cooling 456.2 grams of water at 89.2 °C to a final temperature of 5.9 °C. 22. Determine the energy required to: (a) melt 5.62 moles of ice at 0 °C. (b) melt 74.5 grams of ice at 0 °C. (c) boil 0.345 moles of water at 100.0 °C. (d) boil 43.89 grams of water at 100.0 °C. (e) Convert 16.2 grams of ice to liquid water. (f) Convert 5.8 grams of water to steam (g) Convert 98.2 grams of water to ice. (h) Convert 52.6 grams of steam to water (i) Convert 34.0 grams of water at 20.0 °C to steam at 100.0 °C. (j) Convert 125.0 grams of ice at 0.0 °C to steam at 100.0 °C. (k) Convert 25.9 grams of steam at 100.0 °C.to ice at 0.0 °C. 23. Determine the specific heat of a 150.0 gram object that requires 62.0 cal of energy to raise its temperature 12.0 °C. 24. Determine the energy required to raise the temperature of 46.2 grams of aluminum from 35.8 °C to 78.1 °C. Specific heat capacity of aluminum is 0.089 J/g °C. 182 Unit 12- Kinetics and Thermochemistry 25. How many degrees of temperature rise will occur when a 25.0-g block of aluminum absorbs 10.0 kJ of heat? The specific heat of aluminum is 0.897 J/g·°C. Kinetics and Thermochemistry Practice Test Directions: Match the terms below with their correct definitions. (1-4) a. calorimeter c. enthalpy b. thermochemistry d. system 1. The study of heat changes that accompany chemical reactions and phase changes 2. The specific part of the universe that contains the reaction or process you wish to study 3. The heat content of a system at constant pressure 4. An insulated device used to measure the amount of heat absorbed or released during a chemical or physical process 5. Predict the change in entropy (ΔS) for the following reaction (will it increase, decrease, or stay the same).CH4(g) + 2O2(g) 2H2O(l) + CO2(g) 6. Which of the following species has the highest entropy at 25°C? Explain your answer. a) CH3OH(l) d) H2O(l) b) CO(g) e) Ni(s) c) MgCO3(s) Directions: Solve the following problems, show all of your work 7. If the temperature of a 25-g sample of liquid water is raised 40°C, how much heat is absorbed by the water? 8. Copper metal has a specific heat of 0.385 J/g·°C , calculate the amount of heat required to raise the temperature of 38 g of copper from 20.0°C to 575°C. 183 Unit 12- Kinetics and Thermochemistry 9. When a 50.0-g nugget of pure gold is heated from 35.0°C to 50.0°C, it absorbed 5200.0 J of energy. Find the specific heat of gold. 10. When 80.0 grams of a certain metal at 90.0 °C was mixed with 100.0 grams of water at 30.0 °C, the final equilibrium temperature of the mixture was 36.0 °C. What is the specific heat of the metal? Directions: Use the energy diagram for the rearrangement reaction of methyl isonitrile to acetonitrile to answer the following questions. (1113) 11. What kind of reaction is represented by this diagram, endothermic or exothermic? 12. What does the symbol E represent? 13. How does a catalyst speed a reaction? Multiple Choice Practice 14. As ice cools from 273 K to 263 K, the average kinetic energy of its molecules will a) decrease c) remain the same b) increase 15. The heat of fusion is defined as the energy required at constant temperature to change 1 unit mass of a a) gas to a liquid c) solid to a gas b) gas to a solid d) solid to a liquid 16. What is the total number of joules of heat energy absorbed by 15 grams of water when it is heated from 30°C to 40°C? a) 10 c) 150 b) 63 d) 630 184 Unit 12- Kinetics and Thermochemistry 17. How many joules of heat are absorbed when 70.0 grams of water is completely vaporized at its boiling point? a) 23, 352 c) 15, 813 b) 7, 000 d) 158, 130 18. What occurs as potassium nitrate is dissolved in a beaker of water, indicating that the process is endothermic? a) The temperature of the solution decreases. b) The temperature of the solution increases. c) The solution changes color. d) The solution gives off a gas. 19. A solid is dissolved in a beaker of water. Which observation suggests that the process is endothermic? a) The solution gives off a gas. b) The solution changes color. c) The temperature of the solution decreases. d) The temperature of the solution increases. 20. When a catalyst is added to a system at equilibrium, a decrease occurs in the a) activation energy b) heat of reaction c) potential energy of the reactants d) potential energy of the products 21. Which statement explains why the speed of some chemical reactions is increased when the surface area of the reactant is increased? a) This change increases the density of the reactant particles. b) This change increases the concentration of the reactant. c) This change exposes more reactant particles to a possible collision. d) This change alters the electrical conductivity of the reactant particles. 22. Which conditions will increase the rate of chemical reaction? a) decreased temperature and decreased concentration of reactants? b) decreased temperature and increased concentration of reactants? c) increased temperature and decreased concentration of reactants? d) increased temperature and increased concentration of reactants? 185 Unit 12- Kinetics and Thermochemistry 23. In a chemical reaction, a catalyst changes the a) potential energy of the products b) potential energy of the reactants c) heat of reaction d) activation energy 24. Given the change of phase: CO2(g) changes to CO2(s), the entropy of the system a) decreases c) remains the same b) increases 25. The following reaction coordinate diagram represents... a) an endothermic reaction b) an exothermic reaction c) a reaction that is neither endothermic nor exothermic d) a reaction in which a catalyst is used 26. If 1.45 J of heat are added to a 2.00 g sample of aluminum metal and the temperature of the metal increases by 0.798 oC, what is the specific heat of aluminum? a) 0.579 J/g deg c) 1.68 J/g deg b) 0.909 J/g deg d) 3.63 J/g deg 27. Water has a specific heat of 4.184 J/g deg while glass (Pyrex) has a specific heat of 0.780 J/g deg. If 10.0 J of heat is added to 1.00 g of each of these, which will experience the larger increase of temperature? a) glass b) water c) They both will experience the same change in temperature since only the amount of a substance relates to the increase in temperature. 28. How many PopTarts® are needed to convert 1,000.0 g of water at 20.0 oC to steam at 100.0 oC? One PopTart is equal to 800,000 J of energy...wow! a) 0.419 PopTart c) 3.25 PopTarts b) 2.83 PopTarts 29. The collision theory states that a reaction is most likely to occur if reactant particles collide with the proper a) energy and concentration b) energy and orientation c) concentration and orientation d) pressure and orientation 186 Unit 12- Kinetics and Thermochemistry 30. What will happen to the rate of reaction when temperature increases? a) Increase d) increase then b) Decrease decrease c) remains the same 31. Given the following reaction H+(aq) + OH-(aq) --> H2O(l) if the concentration of the reactants is increased, the rate of reaction will a) Increase d) increase then b) Decrease decrease c) remains the same 32. What does a catalyst decrease when introduced to a reaction? a) the rate of reaction c) the activation b) the energy released energy during the reaction d) the kinetic energy 33. Catalysts increase the rate of reaction by being consumed. a) True b) False 34. Entropy is a measure of the randomness or disorder of a system. a) True b) False 35. Systems in nature tend to undergo changes toward what kind of energy and entropy? a) lower energy and lower entropy b) lower energy and higher entropy c) higher energy and higher entropy d) higher energy and lower entropy 36. Which process is accompanied by a increase in entropy? a) melting of ice b) freezing of water c) condensing of water vapor What phase change represents a decrease in entropy? a) solid to liquid b) liquid to gas c) gas to liquid d) solid to gas 37. 38. 39. H2O(g)--> H2O(l) The entropy in this equation a) Increases b) Decreases Which sample has the lowest entropy? a) 1 mole of KNO3(l) b) 1 mole of KNO3(g) 187 c) remains the same c) 1 mole of H2O(g) d) 1 mole of KNO3(s) Unit 12- Kinetics and Thermochemistry 40. In a chemical reaction, if the products have more entropy than the reactants, the change in entropy is negative. a) True b) False 41. What is the change in entropy in the following reaction C + O2 -> CO2 a) Increases b) Decreases c) remains the same d) not enough information to tell 42. What is the change in entropy in the following reaction 4Al(s) + 3O2(g) -> 2Al2O3(s) a) Increase b) Decrease c) remains the same d) not enough information to tell 43. What is the change in entropy in the following reaction N2(g) + O2(g) -> 2NO(g) a) Increase b) Decrease c) remains the same d) not enough information to tell 44. A 47.5 gram sample of a metal at a temperature of 425°C is placed in 1.00 liters (1000 g) of water which had an initial temperature of 18°C. What is the specific heat capacity of the metal if the final temperature of the metal and water at equilibrium is 21°C? (The specific heat capacity of water is 4.18 J/°C·g. a) 0.03 J/°C·g d) 0.65 J/°C·g b) 1.47 J/°C·g e) 12.54 J/°C·g c) -0.75 J/°C·g 45. A bomb calorimeter has a heat capacity of 3.18 kJ/K. When 0.0038 mol of a gas is burned in the calorimeter, the temperature increased from 25.0°C to 27.3°C. Calculate the energy released by the combustion of one mole of the gas. a) 2.8 x 10-2 kJ d) -2.8 x 10-2 kJ b) 7.3 kJ e) -3.6 x 102 kJ c) -1.9 x 103 kJ 46. A sample of wood has a heat of combustion of 3.29 kJ/g. What quantity of the wood must be burned to heat 250. g of water from 18°C to 85°C? Once again, the specific heat capacity of water is 4.18 J/°C·g. a) 85.1 g d) 2.13 x 104 g b) 0.45 g e) 21.3 g c) 12.4 g 188 Unit 12- Kinetics and Thermochemistry 47. 20.0 mL of pure water at 285 K is mixed with 48 mL of water at 315 K. What is the final temperature of the mixture in kelvins? a) 306 K d) None of these b) 290 K e) 275 K c) 318 K 48. The specific heat of liquid water is 4.18 J/°C·g and the specific heat of carbon is 0.71J/°C·g. A 10 gram sample of water and a 10 gram sample of carbon are each subjected to 155 J of heat. If both samples started at 25°C, which substance will have the higher final temperature, and by what magnitude? a) Neither. They will have the same final temperature because the started at the same temperature and were exposed to the same quantity of heat. b) carbon, by 18.1°C c) water, by 3.47°C d) carbon, by 215°C e) water, by 28.7°C 49. As a result of an exothermic reaction, a) the energy of the system is increased and the energy of the surroundings are decreased. b) the energy of the system and the energy of the surroundings are decreased. c) the energy of the system is decreased and the energy of the surroundings are increased. d) the energy of the system and the energy of the surroundings are increased. e) None of these are accurate 50. The specific heat of iron is 0.450 J/(g·°C). How much heat is required to raise the temperature of a 5.00 gram sample of iron from 22°C to 53°C? a) -43 J d) 18 J b) 155 J e) 69.8 J c) 344 J 51. How much energy is needed to convert 50.0 g of ice at -5.00°C to water at 25°C? a) 22.4 kJ d) 18.0 J b) 37.3 kJ e) 175 J c) 21.9 kJ 189 Unit 12- Kinetics and Thermochemistry 52. Which of the following is not an endothermic process? a) combustion b) melting c) crystallization d) vaporization e) sublimation 53. Increasing the temperature at which a reaction occurs speeds up the reaction by: a) improving the orientation of collisions b) increasing the energy of collisions c) activating catalysts d) increasing the frequency of collisions e) two of these 54. For which of the following processes would ΔS° be expected to be most positive? a) CO2(g) → CO2(s) b) C6H12O6(s) + 6 O2(g) → 6 CO2(g) + 6 H20(g) c) H+(aq) + OH-(aq) → H2O(l) d) Cl2(g) + H2(g) → 2 HCl(g) e) Na+(aq) + Cl-(aq) → NaCl(s) 190 Unit 13- Oxidation and Reduction OXIDATION-REDUCTION One of the defining characteristics of single-replacement and combustion reactions is that they always involve the transfer of ____________________________________ from one atom to another. So do many, but not all, synthesis and decomposition reactions. The ________________________________ number of an atom in an ionic compound is the number of electrons lost or gained by the atom when it forms ions. Oxidation numbers are tools that scientists use in written chemical equations to help them keep track of the movement of electrons in a ________________________ reaction. Like some of the other tools you have learned about in chemistry, oxidation numbers have a specific notation. Oxidation numbers are written with the positive or negative sign __________________________ the number (+3, +2), whereas ionic charge is written with the sign after the number (3+, 2+). Oxidation number: +3. Ionic charge: 3+. In order to understand all kinds of redox reactions, you must have a way to determine the oxidation number of the atoms involved in the reaction. Chemists use a set of rules to make this determination easier. 1. The oxidation number of an uncombined atom is _______________________. This is true for elements that exist as polyatomic molecules such as O2, Cl2, H2, N2, S8. 2. The oxidation number of a _____________________ ion is equal to the charge on the ion. For example, the oxidation number of a Ca2+ ion is +2, and the oxidation number of a Br– ion is –1. 3. The oxidation number of the _______________________ electronegative atom in a molecule or a complex ion is the same as the charge it would have if it were an ion. • In ammonia (NH3), for example, nitrogen is more electronegative than hydrogen, meaning that it attracts electrons more strongly than does hydrogen. So nitrogen is assigned an oxidation number of –3, as if it had gained three electrons to complete an octet. In the compound silicon tetrachloride (SiCl4), chlorine is more electronegative than silicon, so each chlorine has an oxidation number of –1 as if it had taken an electron from silicon. The silicon atom is given an oxidation number of +4 as if it had lost electrons to the four chlorine atoms. 191 Unit 13- Oxidation and Reduction 4. The most electronegative element, _________________________________, always has an oxidation number of –1 when it is bonded to another element. 5. The oxidation number of oxygen in compounds is always ___________, except in peroxides, such as hydrogen peroxide (H2O2), where it is –1. When it is bonded to fluorine, the only element more electronegative than oxygen, the oxidation number of oxygen is +2. 6. The oxidation number of hydrogen in most of its compounds is _____________. The exception to this rule occurs when hydrogen is bonded to less electronegative metals to form _______________________ such as LiH, NaH, CaH2, and AlH3. In these compounds, hydrogen’s oxidation number is –1 because it attracts electrons more strongly than does the metal atom. 7. The metals of groups 1A and 2A and aluminum in group 3A form compounds in which the metal atom always has a ___________________________ oxidation number equal to the number of its valence electrons (+1, +2, and +3, respectively). 8. The sum of the oxidation numbers in a neutral compound is ______________________. 9. The sum of the oxidation numbers of the atoms in a ________________________________ ion is equal to the charge on the ion. Determining Oxidation Numbers Determine the oxidation number of each element in the following compounds or ions. Assign the known oxidation numbers to their elements, set the sum of all oxidation numbers to zero or to the ion charge, and solve for the unknown oxidation number. (Let nelement = oxidation number of the element in question.) • KClO3 (potassium chlorate) Potassium chlorate is a _________________________ salt, so oxidation numbers must add up to zero. K’s oxid. # + Cl’s oxid. # + 3 (O’s oxid. #) = 0 Rule _________ states that Group 1 metals have a +1 oxidation number in compounds. According to rule ___________, the oxidation number of oxygen in compounds is –2. (+1) + (nCl) + 3 (-2) = 0. Chlorine’s oxidation number must equal _______. • SO32– (sulfite ion) Sulfite ion has a charge of 2–, so oxidation numbers must add up to ______. S’s oxid. # + 3 (O’s oxid. #) = –2. According to rule 5, the oxidation number 192 Unit 13- Oxidation and Reduction of oxygen in compounds is –2. (nsulfur) + 3 (-2) = –2. Sulfur’s oxidation number must equal _______. • Determine the oxidation number of each element in the following compound or ion. a) Na2CO3 ___________________________________________ b) H2SO4 ___________________________________________ c) Mg3P2 ___________________________________________ d) SrH2 ___________________________________________ e) PO4-3 ___________________________________________ f) ClO4-1 ___________________________________________ g) C4H10 ___________________________________________ Electron Transfer and Redox Reactions Complete Chemical Equation: 2 Na (s) + Cl2 (g) 2 NaCl (s) Net Ionic Equation: 2 Na (s) + Cl2 (g) 2 Na+ + 2 Cl• Solid sodium, on the reactant side, has an oxidation number of _______. The sodium ion product has an oxidation number of _______. Oxidation is defined as the ____________________ of electrons from atoms of a substance. When an element’s oxidation number ________________________, the element is said to be oxidized. • Sodium’s oxidation number changes from 0 to +1. Sodium is oxidized because it loses ______ electron. To state this reaction more clearly, Na → Na+ + e-. For oxidation to take place, the electrons lost by the substance that is oxidized must be accepted by atoms or ions of another substance. In other words, there must be an accompanying process that involves the ________________ of electrons. _____________________________ is defined as the gain of electrons by atoms of a substance. Net Ionic Equation: 2 Na (s) + Cl2 (g) 2 Na+ + 2 Cl• Chlorine gas, on the reactant side, has an oxidation number of _______. The chlorine ion product has an oxidation number of _______. Chlorine’s oxidation number changes from 0 to -1. Chlorine is reduced because it _________________ one electron. Following our sodium chloride example further, the reduction reaction that accompanies the oxidation of sodium is the reduction of chlorine. Cl2 + 2 e- → 2 Cl-. In this reaction, 193 Unit 13- Oxidation and Reduction an electron from each of two sodium atoms is transferred to the Cl2 molecule to form two Cl– ions. Complete Chemical Equation: 2 Mg (s) + O2 (g) 2 MgO (s) Net Ionic Equation: 2 Mg (s) + O2 (g) 2 Mg2+ + 2 O2• Solid magnesium, on the reactant side, has an oxidation number of ____. The magnesium ion product has an oxidation number of ____. Magnesium’s oxidation number changes from 0 to +2. Magnesium is oxidized because it loses _________ electrons. To state this reaction more clearly, _______________________________________________. Oxygen gas, on the reactant side, has an oxidation number of _______. The oxygen ion product has an oxidation number of _______. Oxygen’s oxidation number changes from 0 to -2. Oxygen is reduced because it gains _________ electrons. Following our magnesium oxide example further, the reduction reaction that accompanies the oxidation of magnesium is the reduction of oxygen. ______________________________________________________ When magnesium reacts with oxygen, each magnesium atom transfers two electrons to each oxygen atom. The two magnesium atoms become Mg2+ ions and the two oxygen atoms become O2– ions (oxide ions). A reaction in which electrons are ______________________________ from one atom to another is called an oxidation–reduction reaction. For simplicity, chemists often refer to oxidation– reduction reactions as ____________________ reactions. Now consider the single-replacement reaction in which chlorine in an aqueous solution replaces bromine from an aqueous solution of potassium bromide. Complete Chemical Equation: 2 KBr (aq) + Cl2 (aq) 2 KCl (aq) + Br2 (aq) Net Ionic Equation: 2 Br- (aq) + Cl2 (aq) 2 Cl- (aq) + Br2 (aq) • Note that chlorine “steals” electrons from bromide ions to become _______________________ ions. When the bromide ions lose their extra electrons, the two bromine atoms form a covalent bond with each other to produce Br2 molecules. The 194 Unit 13- Oxidation and Reduction result of this reaction, the characteristic color of elemental bromine in solution. The formation of the covalent bond by ________________________ of electrons also is an oxidation–reduction reaction. • Note that there is no change in the oxidation number of potassium. The potassium ion takes no part in the reaction and as a result DOES NOT appear in the net ionic equation. It is called a _________________________________ ion. Can oxidation occur without reduction? By our definitions, oxidation and reduction are complementary processes; oxidation cannot occur unless reduction also occurs. • Determine which element is oxidized and which element is reduced in the following reaction. (Provide the oxidation numbers.) a) 3Cu + 2NO3- + 8H+ 3Cu2+ + 2NO + H2O b) 2KNO3 2KNO2 + O2 c) H2 + CuO Cu + H2O d) H2S + NO3- SO42- + NO2 How DO Oxidation and Reduction Differ? LEO the lion says GER or, for short, LEO GER This phrase will help you remember that Loss of Electrons is Oxidation, and Gain of Electrons is Reduction. • Determine how many electrons are lost and how many electrons are gained in the following reaction. (Provide the oxidation numbers.) a) Fe2+ + MnO4- Fe3+ + Mn2+ b) Na2SnO2 + Bi(OH)3 Bi + Na2SnO3 + H2O c) 2Fe + O2 + 2H2O 2Fe(OH)2 195 Unit 13- Oxidation and Reduction Redox and Electronegativity The chemistry of oxidation–reduction reactions is not limited to atoms of an element changing to ions or the reverse. Example: N2 (g) + 3H2 (g) → 2NH3 (g) • Nitrogen’s oxidation number changes from _______ to _______; therefore nitrogen is ____________________________. For the purpose of studying oxidation–reduction reactions, the more electronegative atom (nitrogen) is treated as if it had been reduced by gaining electrons from the other atom. Hydrogen’s oxidation number changes from _______ to _______; therefore hydrogen is _______________________. The less electronegative atom (hydrogen) is treated as if it had been oxidized by losing electrons. Half-Reactions The oxidation process and the reduction process of a redox reaction can each be expressed as a half-reaction. • For example, consider the unbalanced equation for the formation of aluminum bromide. Al + Br2 → AlBr3 • The oxidation half-reaction shows the loss of electrons by aluminum. Al → Al3+ + 3eThe reduction half-reaction shows the gain of electrons by bromine. Br2 + 2e- →2Br- Write balanced half-reactions for each of the following molecular equations. a) K + S → K2S b) Ca + Cl2 → CaCl2 c) AlBr3 → Al + Br2 d) Zn3P2 → Zn + P Practical Applications of Redox A common application of redox chemistry is to remove ____________________________ from metal objects, such as a silver cup. When you add chlorine ________________________ to your laundry to whiten clothes, you are using an aqueous solution of sodium hypochlorite (NaClO). Chlorine is ________________________ in the process. Hydrogen peroxide (H2O2) 196 Unit 13- Oxidation and Reduction can be used as an antiseptic because it _____________________________ some of the vital biomolecules of germs, or to lighten hair because it oxidizes the dark pigment of the hair. When you put a battery into a flashlight, radio, or CD player, you complete the electrical circuit of a galvanic cell(s), providing a path for the electrons to flow through as they move from the site of ____________________________ to the site of the reduction. Electroplating is important in protecting objects from ______________________________ and uses the process of redox. Reduction of silver ions onto cheaper metals forms _______________________________. Homework / Practice What is the oxidation number of . 1) N in NO3¯ 6) Pb in PbOH+ 2) C in CO32¯ 7) V in VO2+ 3) Cr in CrO42¯ 8) V in VO2+ 4) Cr in Cr2O72¯ 9) Mn in MnO4¯ 5) Fe in Fe2O3 10) Mn in MnO42¯ Balance each half-reaction for atoms and charge: 11) Cl2 ---> Cl¯ 13) Fe2+ ---> Fe3+ 12) Sn ---> Sn2+ 14) I3¯ ---> I¯ Separate each of these redox reactions into their two half-reactions (but do not balance): 15) Sn + NO3¯ ---> SnO2 + NO2 16) HClO + Co ---> Cl2 + Co2+ 17) NO2 ---> NO3¯ + NO In each of the following, identify what is being oxidized and what is being reduced. 18) Cu+2 + Fe ---> Fe+2 + Cu 19) Co + Sn+2 ---> Co+2 + Sn 20) 2 Cr + 3 Sn+2 ---> 2 Cr+3 + 3 Sn 197 Unit 13- Oxidation and Reduction 21) CH4 + 2 O2 ---> CO2 + 2H2O 22) Mg + CO2 ---> 2MgO + C 23) 4H+1 + NO3-1 + 3 Fe+2 ---> 2 H2O + NO + 3 Fe+3 24) S-2 + 2 NO3-1 + 4H+1 ---> SO2 + NO2 + 2H2O 25) 2 NH4+1 + 2 NO3-1 ---> 2 N2 + 4 H2O + O2 Oxidation and Reduction Practice Test 1. Explain what takes place in an oxidation-reduction reaction. Directions: Identify the oxidation number for the listed element within each compound or ion 2. Cr in Cr2O72¯ __________ 5. Mn in Mn2O7 __________ 3. Pb in PbOH+ __________ 6. S in H2SO4 __________ 4. V in VO2+ 7. Br in BrO3- __________ __________ Directions: Label the oxidation states on all of the elements and trace the paths of oxidation and reduction. 8. HNO3 + 3Cu2O 6Cu(NO3)2 + 2NO + 7H2O 9. H2SO4 + HI H2S + I2 + H2O Directions: Balance each half-reaction for atoms and charge: 10. Sn ---> Sn2+ 11. Cr5+ Cr3+ 12. Ag Ag1+ 198 Unit 13- Oxidation and Reduction Directions: Separate each of these redox reactions into their two half-reactions: 13. Sn + NO3¯ ---> SnO2 + NO2 14. C2H4 + 3O2 2CO2 + 2H2O 15. 2Ca + O2 2CaO Multiple Choice Practice 16. The oxidation number of sulfur in H2SO4 is a. +2 c. +4 e. +8 b. +3 d. +6 17. What is the oxidation number of chromium in the ionic compound Na2Cr2O7? a. +2+6 18. What is the oxidation number of nitrogen in the molecule N2? a. 0357 19. In the following reaction, the oxidation numbers on sulfur are: NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2 a. +6, +6 and +4 respectively b. +6, +6 and +6 respectively c. +4, +6 and +4 respectively d. +4, +4 and +4 respectively e. +4, +6 and +6 respectively 20. NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2 Which of the reactants contains an element that is oxidized? a. H+ c. Na+ b. Mn4+ d. Cl- e. O2- 21. In the balanced molecular equation for the neutralization reaction between phosphoric acid and potassium hydroxide, the products are: a. KPO4 + H3OH b. 3H2O c. None of these d. K+(aq) + PO43-(aq) + 3H+ + 3OHe. K3PO4 + 3H2O 22. In the reaction 4Al + 3O2 → 2Al2O3 the substance oxidized is: a. O2b. O2 199 Unit 13- Oxidation and Reduction c. Al3+ d. Al e. None of these, because this is not a redox reaction 23. In which of the following compounds does sulfur have an oxidation state of +4? a. H2SO4 b. H2S c. H2SO3 d. MgSO4 e. SO3 200
