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9/6/12 What is chemistry? A chemical is any substance that has a definite composition. A chemical reaction is the process by which one or more substances change to produce one or more new substances. Physical states of matter -The states of matter are the physical forms of matter which are solid, liquid, gas, and plasma. Macroscopic refers to what you can see with the unaided eye Microscopic refers to what you would see if you could see individual atoms. 9/7/12 Solids have a fixed volume and shape that result from the way their particles are arranged. Liquids have a fixed volume but not a fixed shape Gases have neither a fixed volume nor shape Physical Changes are changes in which the identity of a substance doesn’t change - Changes of state are physical changes Chemical Changes occur when the identities of substances change and new substances form. 9/10/12 Chemical Changes Mercury(II) oxide mercury + oxygen Reactants are the substances on the left-hand side of the arrow (they are used up in the reaction Products are the substances on the right-hand side of the arrow (they are produced in the reaction) Evidence of a chemical change - the evolution of a gas - the formation of a precipitate (solid substance forms and typically sinks to the bottom) - the release or absorption of energy - a change in temperature or the giving off of light - a color change in the reaction system Matter Matter- Peanut butter, water, fish, garbage, human brain, carbon dioxide, air, yourself, tree Not Matter- light, time, motion, an idea, energy Distinguish between different characteristics of matter, including mass, volume, and weight. Identity and use SI units in measurements and calculations Set up conversion factors, and use them in calculations Identify and describe physical properties, including density Identify chemical properties Matter had Mass and Volume - Matter is anything that has mass and volume - Volume is the space an object occupies - Mass is the quantity of matter in an object o Devices used for measuring mass in a laboratory are called balances or scales - Weight is the force produced by gravity acting on a mass 9/11/12 Units of Measurements - when working with numbers, be careful to distinguish a quantity and its units o Quantity describes something that has magnitude, size, or amount o Unit is a quantity adopted as a standard of measurement Meter kilogram second ampere Unit symbol m kg s A kelvin K candela cd mole mol Unit name Name Symbol Factor Prefix exa peta tera giga mega kilo hecto deca Symbol E P T G M K h da deci centi milli micro nano pico d c m μ n p 100 1000m 10006 10005 10004 10003 10002 10001 10002/3 10001/3 10000 1000-1/3 1000-2/3 1000-1 1000-2 1000-3 1000-4 decada 101 Quantity name length mass time electric current thermodynamic temperature Dimension symbol L M T I Quantity symbol l (a lowercase L), x, r m t I (an uppercase i) T Θ luminous intensity Iv (an uppercase i with lowercase non-italicized v subscript) J amount of substance n N hectoh 102 kilo- megak M 103 106 giga- teraG T 109 1012 10n Decimal 1018 1000000000000000000 1015 1000000000000000 1012 1000000000000 109 1000000000 106 1000000 103 1000 102 100 101 10 100 1 10-1 0. .1 10-2 0. .01 10-3 0. .001 10-6 0. .000001 10-9 0. .000000001 10-12 0 .000000000001 peta- exa- zettaP E Z 1015 1018 1021 Short scale quintillion quadrillion trillion billion billionth trillionth Long scale trillion billiard billion milliard million thousand hundred ten one tenth hundredth thousandth millionth milliardth billionth A conversion factor is a simple ration that relates two units that express a measurement of the same quantity you can construct conversion factors between kg and grams as follows: 1kg/1000g or 1000g/1kg Given: 4.5kg ?g 4.5kg 1000g = 4500g 1kg - Given: 208lb ? kg 208lb .45kg = 93.6kg 1lb 9/13/12 Given: 0.851L ?mL .851L 1000mL = 851.0mL 1L Derived Units - Many quantities you can measure need units other than the seven basic SI units - These units are derived by multiplying or dividing the base unit o Speed is distance divided by time. A derived unit of speed is meters/second (m/s) o A rectangle’s area is found by multiplying its length (in meters) by its width (also in meters) Its unit is square meters (m2) - Volume is another commonly used derived unit o The volume of a book can be found by multiplying its length, width, and height. o The unit of volume is the cubic meters (m3) o This unit is too large and inconvenient in most labs. Chemists usually use the liter (L) 1L = 1000mL = 1000cm3 1mL = 1cm3 Properties of Matter Physical Properties - A physical property of a substance is a characteristic that does not involve a chemical change. - Physical properties of a substance can be determined without changing the nature of a substance. - Physical properties include malleability, color, texture, state, density, melting point, and boiling point. 9/14/12 Properties of Matter Density is the Ratio of Mass to Volume - The density of an object is the mass if the object divided by volume of the object Densities are expressed in derived unties such as g/cm3 or g/mL Density is calculated as follows: o Density = mass/volume OR D = m/v The density of a substance is the same no matter what the size of the sample is. Because the density of a substance is the same for all samples, you can use this property to help identify substances. 9/18/12 Do Now Kilo = 1000 .01 = centi 45g = .045 kg 100mL = .1L 4.8cg = 48mg Properties of Matter Chemical Properties - A chemical property is a property of matter that describes a substance’s ability to participate in chemical reactions – i.e.- Reactivity - A chemical property of many substances is their reactivity with oxygen. o Rusting, corrosion - Some substances break down into new substances when heated Classifying Matter - An atom is the smallest unit of an element that maintains the properties of that element. - Matter exists in many different forms but there are only 118+ types of atoms. - Atoms are joined together to make up all the different kinds of matter. Pure Substance - A pure substance is a sample of matter, either a single element or a single compound, which has definite chemical and physical properties. - Elements are pure substances that only contain one kind of matter (atom). They cannot be separated or broken down into simpler substances by ordinary chemical means. (exception- Higgs Boson particle) o Each element has its own unique set of physical and chemical properties - A molecule is the smallest unit of a substance that keeps all of the physical and chemical properties of that substance. - A molecule usually consists of two or more atoms combined in a definite ratio o Except diatomic elements - Diatomic elements exist as two atoms of the same element joined together. o H2 O2 N2 F2 Cl2 Br2 I2 o Acronym- HOFBrINCl - Some elements, such as oxygen, phosporus, sulfur, and carbon, have many molecular forms - An allotrope is one of a number of molecular forms of an element - The properties of allotropes vary widely o Carbon- graphite, diamond, buckyballs 9/20/12 Pure Substances Some Elements Exist in More than One Form - Oxygen exists as allotropes - Oxygen gas (O2) is colorless and odorless - Ozone (O3) is toxic and pale blue Compounds are Pure Substances - Pure substances that are not elements are compounds. Compounds are composed of more than one kind of atom. o Example: carbon dioxide - There may be easier ways of preparing them, but compounds can be made from their elements. - Compounds can be broken down into their elements, often with difficulty. Compounds are Represented by Formulas - Compounds have characteristic in properties and composition - Compounds can be represented by an abbreviation or Formula o A formula has subscripts, which represent the number of different atoms in the compound. o Example: H2O has 2 hydrogen atoms and 1 oxygen atoms - Molecular formulas give information only about what makes up a compound. o Example: the molecular formula for aspirin is C9H8O4 - A structural formula shows how the atoms are connected o Two-dimensional models do not show the molecule’s true shape 9/21/12 How is matter classified? - A call and stick model shows the distances between atoms and the angles between them in three dimensions. - A space filling model attempts to represent the actual size of the atoms nd not just their relative position - These models convey different information Mixtures - A mixture is a combination of two or more substances that are not chemically combined - Air is a mixture of mostly nitrogen and oxygen - Water is not a mixture o The H and O atoms are chemically bonded o The ratio of H to O atoms is always 2 to 1 - The properties of the mixture may vary - An alloy is a solid mixture o Example: an alloy of gold and other metal atoms is stronger than pure gold 18-karot gold contains 18 grams of gold per 24 grams of alloy 14-karot gold contains 14 grams of gold per 24 grams of alloy Homogenous mixture - A homogenous mixture describes something that has a uniform structure or composition throughout. o Example: gasoline, syrup, and air - Homogenous mixtures have the same properties throughout Heterogeneous Mixtures - A heterogeneous mixture describes something that is composed of dissimilar components o Example: a mixture of sand and water - Any two samples of a heterogeneous mixture will have the different proportions of ingredients o Heterogeneous mixtures have different properties throughout - Mixtures can be either homogenous or heterogeneous 9/24/12 Do Now 37 gigabytes = 37,000,000 kilobytes 763 mL = .763L 45 kilograms = 45,000,000 mg Understanding Concepts Tea is best classified as a homogenous mixture He is an element CHAPTER 2 Objectives - Explain that physical and chemical changes in matter involve transfers of energy - Apply the law of conservation of energy to analyze changes in matter - Distinguish between heat and temperature - Convert between Celsius and Kelvin Energy and Change - Energy is the capacity to do work, such as moving an object, forming a new compound, or generating light (Electromagnetic Energy) - Energy is always involved when there is a change of matter Changes in Matter Can Be Physical or Chemical - Ice melting and water boiling are examples of physical changes - A Physical Change is a change of matter that affects only the physical properties of matter. - In contrast, the reaction of hydrogen and oxygen to produce water is an example of a chemical change - A Chemical Change is a change that occurs when one or more substances change into entirely new substances with different properties - A chemical change occurs whenever a new substance is made - All physical and chemical changes involve a change in energy - Sometimes energy much be supplied for the change in matter to occur - Evaporation is the change of a substance form a liquid to a gas due to the absorption of Energy 9/25/12 Energy and Change Endothermic and Exothermic Processes - Any change in matter in which energy is absorbed from the surrounding in an Endothermic process o The melting of ice and boiling of water are examples of physical changes that are Endothermic o As the chemicals react, energy is absorbed. Energy is a Reactant - Any change in matter in which energy is released is an Exothermic process o The freezing of water and condensation of water vapor are two examples of physical changes that are exothermic processes. o Energy is released. Energy is a Product. - Energy can be absorbed (Endothermic) by the surroundings or released (exothermic) to the surroundings, but it cannot be created or destroyed. - The Law of Conservation of Energy states that during any physical or chemical change, the total quantity of energy remains constant. - In other words, energy cannot be created nor destroyed. Energy is often transferred - A system consists of all the components that are being studied at any given time. o The chemicals are the system, not the beaker they are in. - The surroundings include everything around outside the system. - Energy exists in different forms, including: o Heat, Light, Potential-Chemical, kinetic-Motion, Sound, Electrical, Pressure waves - The transfer of energy between a system and its surroundings can involve any one of these forms of energy Heat - Heat is the energy transferred between objects that are different temperatures. - Heat energy is always transferred from a warmer object to a cooler object - For example, when ice cubes are placed in water, heat energy is transferred from the water to the ice. - Energy is also transferred as heat during chemical changes. 9/27/12 Energy Can Be Absorbed as Heat - In an endothermic reaction, energy is absorbed by the chemicals that are reacting. Heat is Different from Temperature - Scientists define Temperature as a measurement of the average kinetic energy of the random motion of the particles in a substance - The transfer of energy as Heat can be measured by calculated changes in Temperature Temperature is Expressed Using Different Scales - Thermometers are usually marked with the Fahrenheit or Celsius temperature scales - A third temperature scale uses the unit Kelvin, K The zero point of the Celsius scale is designed as the freezing point of water The zero point on the Kelvin scale is designated as absolute zero, the temperature at which the minimum average kinetic energies of all particles occur. - There are no negative temperatures in Kelvin Transfer of Heat May Not Affect Temperature - The transfer of energy as heat does not always result in a change of temperature o For example, consider that happens when energy is transferred to a mixture of ice and water o As energy is transferred as heat to the ice-water mixture, the ice cubes start to melt. o The temperature of the mixture remains at 0C until all the ice has melted. - Transfer of Heat Affects Substances Differently - The specific heat of a substance is the quantity of energy as heat that mist be transferred to raise the temperature of 1g of a substance 1K - The SI (metric) Unit for energy is Joule (L) - Specific Heat is expressed in joules per gram Kelvin (J/gK) - Metals tend to have low specific heats - Water has a high specific heat 9/28/12 Scientific Method - The Scientific Method is a series of steps followed to solve problems, including: o Collecting data o Formulating a hypothesis o Testing the hypothesis o Stating conclusions - An experiment is the process by which scientific ideas are tested. - A Hypothesis is a reasonable and testable explanation for observations. - A Variable is a factor that could affect the results of an experiment. - When a variable is kept constant form one experiment to the next, the variable is called the control. - The procedure is called a Controlled Experiment - Any hypothesis that withstands repeated testing may become part of a theory - In science, a Theory is a well-tested explanation of observations Theories and Laws Have Different Purposes - Some facts in science always hold true. These facts are called laws. - A Law is a statement or mathematical expression that reliably describes a behavior of the natural world. - A theory is an attempt to explain the cause of certain events in the natural world. - For example, the Law of Conservation of Mass states that the products of a chemical reaction have the same mass as the reactants have. - Keep in mind that a hypothesis predicts an event, a theory explains it, and a law describes it. Models Can Illustrate the Microscopic World of Chemistry - A Model represents an object, a system, a process, or an idea 10/3/12 Measurements and Calculations in Chemistry Accuracy and Precision - The Accuracy of a measurement is how close the measurement is to the true or actual value - Precision is the exactness of a measurement - It refers to how closely several measurements of the same quantity made in the same way can agree. Significant Figures - Scientists report values using significant figures - The Significant Figures of a measurement or a calculation consists of all the digits known with certainty as well as one estimated, or uncertain, digit. - The last digit or significant figure reported after a measurement is uncertain and estimated Rules for Determining Significant Figures - Nonzero digits are always significant - Zeros between nonzero digits are significant - Zeros in front of nonzero digits are not significant - Zeros both at the end of a number and to the right of a decimal point are significant - Zeros both at the end of a number but to the left of a decimal point may not be significant. If a zero has not been measured or established, it is not significant. A decimal point placed after zeros indicated that the zeros are significant 10/8/21 Do Now Define: Density- mass divided my volume Formula- D=m/v Units- g/mL or g/cm3 Significant Figures - To determine the number of significant figures in the answer, you must first find the number of significant figures in the values used to calculate the answer. When multiplying and dividing, the answer cannot have more sigfigs than the measurement with the smallest number of sigfigs. - With addition and subtraction, the result can be no more certain than the least certain number in the calculation. - Conversion factors and countable values have unlimited numbers of sigfigs OR it does not limit my calculation. 10/9/12 Significant Figures Example: A student heats 23.62g of a solid and observes that the temperature increases from 21.6C to 36.79C. Calculate the temperature increase per gram of solid. 36.79C - 21.6C = 15.19C = 15.2C 15.2C/23.62g = 0.644C/g Specific Heat Depends on Various Factors - Specific heat depends on the nature of the material that is changing temperature, the mass of the material, and the size of the temperature change. - Recall that the specific heat is the quantity of energy that must be transferred as heat to raise the temperature of 1g of a substance by 1K. - The specific heat (cp) of a substance at a given pressure (p) is calculated by the following formula o Specific heat = cp = q / mxT (q-heat in joules, m-mass in grams, T- change in temperature in K) Example: A 4.0g sample of glass was heated from 274K to 314K and was found to absorb 32J of energy as heat. Calculate the specific heat of this glass. cp = q = 32J = .20J/gK mxT (4.0g)(40K) 10/10/12 Scientific Notation - Very large and very small numbers are often written in scientific notation. - A number in scientific notation has 2 parts - The first part is a number that is between 1 and 10 - The second part consists of a power of 10 - To write the first part of the number, move the decimal to the right of the left so that only one nonzero digit is to the left of the decimal. Scientific Notation with SigFigs - Use scientific notation to eliminate all place-holding zeros - Move the decimal in an answer so that only one digit is to the left, and change the exponent accordingly. The final value must contain the correct number of sigfigs. o 1001000000 = 1.001x109 o 0.0000456 = 4.56x10-5 o 0.0000036mm = 3.6x10-6mm o 1450000mg = 1.45x106mg - 10/11/12 SigFig Practice 1890.01m – 6 sigfigs 0.01702L – 4 sigfigs Specific heat- is the quantity of energy that must be transferred as heat to raise the temperature of 1g of a substance by 1K. Specific heat = cp = q / mxT (q-heat in joules, m-mass in grams, T- change in temperature in K) 10/15/12 CHAPTER 3 3.1 Atomic Theory - The idea of an atoms theory is more than 2000 years old. - Until recently, scientists had never seen evidence of atoms. - The law of definite proportions, the law of conservation of mass and the law of multiple proportions support the current atomic theory The Law of Definite Proportions - The Law of Definite Proportions states that a chemical compound always contains the same elements in exactly the same proportions by weight or mass. - The law of definite proportions also states that every molecule of a substance is made of the same number and types of elements. The Law of Conservation of Mass - The law of conservation of mass states that mass cannot be created nor destroyed in ordinary chemical and physical changes. - The mass of the reactants is equal to the mass of the products. The Law of Multiple Proportions - The law of multiple proportions states that when two elements combine to form two or more compounds, the mass of one element that combines with a given mass of the other is in the ratio of small whole numbers. Dalton’s Atomic Theory - In 1808, John Dalton developed an atomic theory. - Dalton believed that a few kinds of atoms made up all matter. - According to Dalton, elements are composed of only one kind of atom and compounds are made from two or more kinds of atoms. - Dalton’s Theory Contains Five Principles o All matter is composed of extremely small particles called atoms, which cannot be subdivided, created, or destroyed. o Atoms of a given element are identical in their physical and chemical properties. o Atoms of different elements differ in their physical and chemical properties. o Atoms of different elements combine in simple, whole number ratios to form compounds. o In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed. Data gathered since Dalton’s time shows that the first two principles are not true in all cases. HW- Read Section 3.1, do 1-9 (pg 78) 10/16/12 3.2 Structure of Atoms Subatomic Particles - Experiments by several scientists in the mid-1800’s led to the first change to Dalton’s atomic theory. - The smaller parts that make up atoms are called subatomic particles - The three subatomic particles that are most important for chemistry are the electron, proton, and neutron. Electrons were discovered using Cathode Rays - JJ Thompson studied currents - Thomson observed a glowing beam that came out of the cathode and struck the anode and the nearby glass walls of the tube. o He called these rays cathode rays o The glass tube Thomson used is known as a cathode-ray tube. An electron has a Negative Charge - Thomson’s experiments showed that a cathode ray consists of particles that have mass and a negative charge - There particles are called electrons - An electron is a subatomic particle that has a negative electric charge - Electrons are negatively charged, but atoms have no charge o Atoms contain some positive changes that are balanced by negative charges. - Electron: e, e-, or –10e ; charge of 1.602x10-19C ; common charge notation –1 ; mass is 9.109x10-31kg Rutherford Discovered the Nucleus - Thomson proposed that the electrons of an atom were embedded in a positively charged ball of matter. His model of an atom was named the plumpudding model. - Ernest Rutherford performed the gold foil experiment, which discovered the nucleus of an atom. o A beam of small, positively charged particles, called alpha particles, was directed at a thing gold foil. - Rutherford called the space around the atom, the nucleus. 10/17/12 - The nucleus is the dense, central portion of the atom. - The nucleus is made up of protons and neutrons - The nucleus has all the positive charge, and nearly all of the mass, but only a very small fraction of the volume of the atom. - Protons - Protons are the subatomic particles that have a positive charge and that is found in the nucleus of an atom. o The number of protons of the nucleus is the atomic number, which determines the identity of an element. o Because protons and electrons have equal but opposite charges, a neutral atom must contain equal numbers of protons and electrons. - Neutrons are the subatomic particles that have no charge and that are found in the nucleus of an atom Charge Particle Relative Charge Mass Relative mass Proton +1.60 x 10-19 C +1 1.672 x 10-24 g 1 amu Neutron neutral 0 1.674 x 10-24 g 1 amu Atomic Number - The number of protons that an atom has is known as the atomic number. o The atomic number is the same for all atoms of an element. o No two elements can have the same atomic number. - Atomic numbers are always whole numbers. - The atomic number also reveals the number of electrons in an element. o For atoms to be neutral, the number of negatively charged electrons must equal the number of positively charged protons. 10/18/12 Do Now 1) Who discovered the electron? JJ Thomson How? Cathode Ray Tube When? Early 1900’s 2) What did Rutherford determine? Atom is made up of mostly space and there is dense nucleus that is positively charged in the 1900’s 3) Who developed the first Atomic Theory? Dalton in the early 1800’s Mass Number is the Number of Particles in the Nucleus - The mass number is the sum of the number of protons and neutrons of an atom. o Mass# = p+ + n0 o Mass of e- = 0 (charge –1) o Mass of p+ = 1 (charge +1) o Mass of n0 = 1 (charge 0) - You can calculate the number of neutrons in an atom by subtracting the atomic number (the number of protons) from the mass number (the number of protons and neutrons) o Atomic# - p+ o Mass# - p+ + n0 o # of neutrons = mass# - atomic# - Unlike the atomic number, the mass number can vary depending on the number of neutrons. Example: a particular atom of neon has a mass number of 20. How many neutrons does is have? Mass# = 20 Atomic# = 10 (has 10 protons) Mass# - Atomic# = 20 – 10 = 10 Neutrons - Sample Problem A: How many protons, neutrons, and electrons are present in an atom of copper whose atomic number is 29 and whose mass number is 64. Copper (Cu) ?p+ = 29 ?e- = 29 ?n0 = mass – atomic = 64 – 29 = 35 - The atomic number always appears on the lower left side of the symbol. - Mass numbers are written on the upper left side of the symbol. 1 2 2 2 3 7 5 9 11 1H 1H 3He 4He 6Li 3Li 4Be 5B 5B + + + + + + + + 1p 2p 2p 2p 3p 7p 5p 9p 11p+ 0n0 1n0 0n0 0n0 0n0 4n0 1n0 4n0 6n0 1e- 1e- 3e- 4e- 6e- 3e- 4e- 5e- 5e10/22/12 Atomic Number and Mass Number Isotopes - All atom of an element have the same atomic number and the same number of protons. Atoms do not necessarily have the same number of neutrons. - Atoms of the same element that have different numbers of neutrons. - One standard method of identifying isotopes is to write the mass number with a hyphen after the name of the element. o helium-3 or helium-4 - The second method of identifying isotopes shoes the composition of a nucleus as the isotope’s nuclear symbol. o 32He or 42He 10/23/12 Do Now - Element Magnesium Magnesium Symbol Mg Mg Magnesium chloride Magnesium Potassium ClMg KMg Magnesium Phosphorus Magnesium Aluminum P Mg AlMg Magnesium U Mg Au Uranium Magnesium Gold Molybdenum Mo Cesium Cs p+ n0 e- Mass# 12 17 19 15 13 92Mg 79 42 55 12 18 21 16 14 146 118 54 78 12 17 19 15 13 92 79 42 55 24 35 40 31 27 238 197 96 133 Atomic# 12 17 19 15 13 92 79 42 55 10/24/12 3.3 Electron Configuration- Atomic Models Rutherford’s model proposed electron orbits - The experiments of Rutherford’s team led to the replacement of the plum pudding model of the atom with a nuclear model of the atom. o Rutherford suggested that electrons, like planets orbiting the sun, revolve around the nucleus in circular or elliptical orbits. o Rutherford’s model could not explain why electrons did not crash into the nucleus. - Neils Bohr replaced the Rutherford model of an atom 2 years later. Bohr’s Model Confines Electrons to Energy Levels - According to Bohr’s model, electrons can be only certain distances from the nucleus. Each distance corresponds to a certain quantity of energy that an electron can have. o An electron that is as close to the nucleus as it can be is in its lowest energy level. o The farther an electron is from the nucleus, the higher the energy level that the electron occupies. - The difference in energy between two energy levels is known as a Quantum of energy - Rutherford had a planetary model - Bohr had the quantum energy model Electrons Act like Both Particles and Waves - Thomson’s experiments demonstrated that electrons act like particles that have mass. - In 1924, Louis de Broglie pointed out that the behavior of electrons according to Bohr’s model was similar to the behavior of waves. - De Broglie suggested that electrons could be considered waves confined to the space around the nucleus. - The present day model of the atom takes into account both the particle and wave properties of electrons. - In this model, electrons are located in Orbitals, regions around a nucleus that correspond to specific energy levels o Orbitals are regions where electrons are likely to be found. 10/26/12 Spectroscopy- study of light absorption, separating light into its component colors. Electrons and light - By 1900, scientists knew that light could be though of as moving waves that have given frequencies, speeds, and wavelengths. - The wavelength is the distance between two consecutive peaks pr troughs of a wave. - The electromagnetic spectrum is all the frequencies or wavelengths of electromagnetic radiation. - In 1905, Albert Einstein proposed that light also has come properties of particles. His theory explained a phenomenon known as the Photoelectic Effect. o This effect happens when light strikes a metal and electrons are released. - Einstein proposed that light has the properties of both particles and waves. Light is an electromagnetic wave - Red light has a low frequency and a long wavelength - Violet light has a high frequency and short wavelength - 11/6/12 Wavelength and Frequency - The distance between any two corresponding points on one wave, such as from crest to crest, is the wavelength - We can measure the wave’s frequency, v, by observing how often the wave rises and falls at a given point. Light provides information about Electrons - An electron in a state of its lowest possible energy is in a Ground State. o The ground state is the lowest energy state of a quantized system. - If an electron gains energy, it moves to an Excited State. o An excited state is a state in which an atom has more energy than it does at its ground state. - An electron is an excited state will release a specific quantity of energy as it quickly “falls” back to its ground state. Quantum Numbers - The present-day model of the atom is known as the Quantum Model - A Quantum Number is a number that specified the properties of elements. 11/7/12 - The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron. o Values of n are positive integers, such as 1, 2, 3, and 4. o As n increases, the electron’s distance from the nucleus and the electron’s energy increases. - The main energy levels can be divided into sublevels. These sublevels are represented by the angular momentum quantum number, I. o This quantum number indicates shape or type of orbital that corresponds to a particular sublevel. o A letter code is used for this quantum number. I = 0 corresponds to an s orbital I = 1 to p orbital I = 2 to d orbital I = 3 to f orbital - The magnetic quantum number, symbolized by m, is a subset of the I quantum number. o It also indicates the numbers and orientation of orbitals around the nucleus. o The value of m takes whole-number values, depending on the value of I. o The number of orbitals includes One s orbital Three p orbitals ( dumbbells) Five d orbitals Seven f orbitals - The spin quantum number indicates the orientation of an electron’s magnetic field relative to an outside magnetic field o The spin quantum number is represented by + 1/2 or – 1/2 or or o A single orbital can hold a maximum of two electrons, which must have opposite spins. 11/9/12 Do Now What are the 4 orbitals? S p d f How many shapes do they each have? 1 3 5 7 How many (max) e- can each orbital hold? 2 6 10 14 Electron Configuration - Write the electron configuration for Nitrogen. 7 e1s22s22p3 14 Write the electron configuration for Calcium 20 e1s22s22p63s23p64s2 15 Write the electron configuration for Silicon 14 e 1s22s22p63s23p2 11/12/12 Do Now- Write the electron configuration for: Mg- 12 electrons valence electrons 2 6 2 1s22s 2p 3s outermost shell 6 2 5 (2 valence electrons) Cl- 1s22s22p 3s 3p C- 1s22s22p2 He- 1s2 Ne- 1s22s22p6 The max number of e-‘s in an outer energy level is 8 11/13/12 Aufbau Principle Lowest Energy Level Highest Energy Level 11/15/12 Patterns in Element Properties - The elements lithium, sodium, potassium, rubidium, and cesium can combine with chlorine in a 1:1 ratio o They form LiCl, NaCl, KCl, RbCl, and CsCl John Newlands - In 1865, the English chemist John Newlands arranged the known elements according to their properties and in order of increasing atomic mass. He placed the elements in a table. - Newlands noticed that all of the elements in a given row had similar chemical and physical properties - Because these properties seemed to repeat every eight elements, Newlands called this pattern the Law of Octaves. Dmitri Mendeleev- invented the first periodic table - In 1869, the Russian chemist Dmitri Mendeleev used Newlands’ observation and other information to produce the first orderly arrangement, or periodic table, of all 63 elements known at the time. - Mendeleev wrote the symbol for each element, along with the physical and chemical properties and the atomic mass and number on a card. - Mendeleev started a new row each time he noticed the chemical properties of the elements repeated. o He placed elements in the new row directly below elements of similar chemical properties in the preceding row. - Mendeleev made two interesting observations o Mendeleev’s table contains gaps that elements with particular properties should fill. o The elements do not always fit neatly in order of atomic mass. - Mendeleev predicted the properties of the missing elements. The Physical Benefits of the Periodic Table - About 40 years after Mendeleev published his periodic table, an English chemist named Henry Moseley found a different physical basis for the arrangement of elements. Moseley studied the lines in the X-ray spectra of 38 different elements; he found that the wavelengths of the lines in the spectra decreased in a regular manner as atomic mass increased. 11/16/12 - When elements were arranged by increasing atomic number, the discrepancies of Mendeleev’s table disappeared. Periodic Law Organization of the Periodic Table - When the elements are arranged by increasing atomic number, elements with similar properties appear at regular intervals. - Elements in each column have the same number of electrons in their outer shell. - The electrons in the outer shell are called Valence Electrons. o Valence Electrons are found in the outermost shell of an atom and that determines the atom’s chemical properties - Elements with the same number of valence electrons tend to react in similar ways. - A vertical column on the periodic table is called a group. Elements in a group share chemical properties. - A horizontal row on the periodic table is called a period. Elements in the same period have the same number of occupied energy levels. 11/19/12 Tour of the Periodic Table Main Group Elements - Elements in groups 1,2 and 13-18 are known as the main-group elements. Main- group elements are in s- and p-blocks of the periodic table. - Main- group elements are sometimes called representative elements because they have a wide range of properties o At room temperature and atmospheric pressure, many are solids, while others are liquids or gases. o About half of the main-group elements are metals o Many are extremely reactive, while several are nonreactive. - The main-group elements silicon and oxygen account for four of every five atoms found on or near the earth’s surface. - Four groups within the main-group elements have special names. These are: o Alkali metals (group 1) o Alkaline-earth metals (Group 2) o Halogens (Group 7 or 17) o Noble gases (Group 8 or 18) The Alkali Metals make up Group 1 - Elements in Group 1 are called alkali metals o Litium, sodium, potassium, rubidium, cesium, and francium - Alkali metals are so named because they are metals that react with water to make alkaline solutions. - Because the alkali metals have a single valence electron, they are very reactive. - Alkali metals are never found in nature as pure elements but are found as compounds. The Alkaline-Earth Metals Make up Group 2 - Group 2 elements are slightly less reactive than the alkali metals - The alkaline-earth metals are slightly less reactive than the alkali metals o They are usually found as compounds. - The alkaline-earth metals have two valence electrons and must lose both their valence electrons to get to a stable configuration. o It takes more energy to lose two electrons than it takes to lose just on electron The Halogens make up Group 7 (or 17) - They are the most reactive group of nonmetal elements. o When Halogens react, they often gain the one electron needed to have eight valence electrons, a filled outer energy level. - Because the alkali metals have one valence electron, they are ideally suited to react with halogens. - The halogens react with most metals to produce salts. The Noble Gases, Group 8 (or 18) are unreactive - Group 8 elements are called the noble gases - The noble gas atoms have a full set of electrons in their outermost energy level - The low reactivity of noble gases leads to some special uses. - The noble gases were once called inert gases because they were thought to be completely unreactive. o In 1962, chemists were able to get xenon to react, making the compound XePtF6 o In 1979, chemists were able to form the first xenon-carbon bonds. Hydrogen is in a class by itself - Hydrogen is the most common element in the universe. 11/27/12 Do Now Group 1- Alkali Metals Group 2- Alkaline Earth Metals Group 7/17- Halogens Group 8/18- Noble Gases Where are the metals located? Left of the staircase Where are the nonmetals located? Right of the staircase How are elements organized on the periodic table? atomic number (p+) Chapter 4 (cont) Most elements are Metals Transition Metals occupy the center of the periodic table - Generally, the transition metals are less reactive than other metals. Lanthanides and Actinides Fill f-orbitals - The elements in the first row are called Lanthanides because their atomic number follow the element lanthanum. - A lanthanide is a member of the rare-earth series of elements, whose atomic numbers range from 58-71. - The actinides are unique in that their nuclear structure Other Properties of Metals - An alloy is a solid or liquid mixture of two or more metals. - The properties of an alloy are different from the properties of the individual elements. - A common alloy is brass, a mixture of copper and zinc o Brass is harder than copper and more resistant to corrosion. Section 3- Trends in the Periodic Table Periodic Trends - The arrangement of the periodic table reveals trends in the properties of the elements. - A trend is a predictable change in a particular direction. - Understanding a trend among elements enables you to make predictions about the chemical behavior of the elements. - These trends in properties of the elements in a group or period can be explained in terms of electron configuration. Ionization Energy - The ionization energy required removing an electron from an atom or ion. A (neutral atom) + ionization energy A+ + e11/28/12 Do Now Define- Ductile- draw into a wire, stretch it Malleable- bendable, folded What “block” are the transition metals? D Block What group are the Halogens? 7/17 Alkaline Earth Metals? 2 Noble gases? 8/18 Alkali Metals? 1 Periodic Trends 1) Atomic Radius- distance from the nucleus to the edge of the electron cloud increases decreases - Electron cloud 2) 3) 4) 5) -Bond Radius – half the distance between bonding atoms (nucleii) Electronegativity- the measure of the ability of an atom in a compound to attract electrons decreases increases Ionic Size- the increase or decrease in the size of an atom when it gains or loses electrons (ion- an atom with a charge) Ionization Energy- the energy required to remove an electron from an atom or ion BP/MP(boiling point/melting point)- changes in state due to energy changes 11/29/12 Do Now - List the 4 evidences of a chemical change: color change, evolution of a gas, release or absorption of energy, precipitate. - Convert 3.70kg to grams: 3700g - Place 23700 into scientific notation: 2.37x104 - What is the metric system unit of volume: L 12/3/12 CHAPTER 7 The Mole - 1 mole = 6.02x1023 - Mole - defined in terms of carbon - the number of atoms in exactly 12 grams of the carbon isotope-12 - 1 mole = atomic mass on periodic table - Avogadro’s Number- the number of particles in 1 mole of a substance 12/4/12 Avogadro’s Number and the Mole - The mole is a counting unit - Amount in Moles can be converted to number of particles. - Counting units are used to make conversion factors - The conversion factor is 6.02x1023particles / 1mol 12/5/12 Converting Amount in Moles to Number of Particles - Law of Conservation of Mass- xg in reactants = xg in products - Example: Find the number if molecules in 2.5mol of sulfur dioxide? Given: 2.5mol sulfur dioxide ? # of molecules 2.5mol sulfur dioxide 6.02x1023 molecules = 1 mol sulfur dioxide 12/6/12 MUST KNOW THIS: 1mol or 6.02x1023 1mol molar mass or atomic mass or 6.02x1023 1mol molar mass or atomic mass 1 mol HW Review 6c) Given: 0.25molK+ ? molecules 0.25molK+ 6.02x1023moleculesK+ = 1.505x1023moleculesK+ 1molK+ 7b) Given: 3.00molNa4P2O7 ? # Na+ ions 3.00molNa4P2O7 6.02x1023Na4P2O7 4Na+ = 7.224x1024Na+ 1molNa4P2O7 1Na4P2O7 8a) Given: 3.01x1023moleculesH2O ? # moles 3.01x1023moleculesH2O 1molH2O = .5molH2O 6.02x1023moleculesH2O 12/7/12 Do Now 6.7molH2O 6.02x1023H2O = 4.0x1024moleculesH2O 1molH2O 2.68molCO2 6.02x1023CO2 = 1.61x1024moleculesCO2 1molCO2 0.005molH2O 6.02x1023H2O = 3.01x1021moleculesH2O 1molH2O Molar Mass - the mass, in grams, of 1 mole of an element - g/mol - the mass of 1 mole of an element = the element’s atomic mass (from P.T.) - ex: 1molCu = 63.55g - 197gAu and 63.5gCu are both 1mol which both have 6.02x1023particles - 1molH2O = 18g 12/10/12 Do Now - How many atoms in 2.75 moles of H2O? 2.75molH2O 6.02x1023atomsH2O = 1.65x1024atomsH2O 1molH2O - How many moles in 1.86x1024molecules of H2SO4? 1.86x1024moleculesH2SO4 1molH2O = 3.08molH2SO4 6.02x1023moleculesH2SO4 HW Review 13a) 2.0molH2 6.02x1023moleculesH2 = 1.2x1024moleculesH2 1molH2 b) 4.01gHF 1molHF 6.02x1023moleculesHF = 1.21x1023moleculesHF 20.01gHF 1molHF 9) 4.30x1016atomsHe 1molHe 4.00gHe = 2.85x10-7gHe 6.02x1023atomsHe 1molHe Molar Mass- The mass, in grams (g), of 1 mole of a substance, element, compound. Ex: H2O = 2H = 2x1 = 2 1O = 1x16 = 16 18g/mol 12/12/12 Percent Composition - percent, by mass, of each element in a compound molar mass of the part molar mass of the whole ex: XeF4 1Xe: 1x131= 131 %Xe: Xe = 131 100 = 63.3% 4F: 4x19 =+76 XeF4 207 207 %F: F4 = 76 100 = 36.7% XeF4 207 ex: FeO 1Fe: 1x56= 56 %Fe: Fe = 56 100 = 78% 1O: 1x16=+16 FeO 72 72 %O: O = 16 100 = 22% FeO 72 12/13/12 Do Now Determine the molar mass of Mg(ClO3)2 mass of each element from P.T. Mg(ClO3)2 1Mg: 1x24= 24 2Cl: 2x35.5= 71 6O: 6x16= +96 191g/mol molar mass of Mg(ClO3)2 Determine the %composition of NH4Br 1N: 1x14= 14 %N: N = 4H: 4x1= 4 NH4Br 1Br: 1x80= +80 98g/mol %H: H4 = NH4Br %Br: 14 100 = 14.3% 98 4 100 = 4.1% 98 Br = 80 100 = 81.6% NH4Br 98 Formula and Percentage Composition Empirical Formula - An empirical formula is chemical formula that shows the simplest ratio of the relative numbers and kinds of atoms in a compound. - An actual formula shoes the actual ratio of elements or ions in a single unit of a compound. - For example, the empirical formula for ammonium nitrate is NH2O, while the actual formula is NH4NO2. You can use percent composition for a compound to determine its empirical formula. o Convert the percentage of each element to grams o Convert from g to mol using the molar mass of each element as a conversion factor. Determine the empirical formula of the Do Now from today: (now divide by the smallest - mol) %N: N = NH4Br %H: H4 = NH4Br %Br: Br = NH4Br 14 100 = 14.3% 14.3gN 98 4 100 = 4.1% 4.1gH 98 80 100 = 81.6% 81.6gBr 98 1molN = 1molN = 1 14gN 1 1molH = 4.1molH = 4 1gH 1 1molBr = 1molBr =1 80gBr 1 NH4Br Ex: Chemical analysis of a liquid shoes that it is 60.0%C, 13.4%H, and 26.6%O by mass. Calculate the empirical formula of this substance. %C: 60.0% 60.0gC 1molC = 5molC = 2.9 = 3 12gC 1.7 %H: 13.4% 13.4gH 1molH = 13.4molH =7.9 = 8 C3H8O 1gH 1.7 %O: 26.6% 26.6gO 1molO = 1.7molO = 1 = 1 16gO 1.7 12/14/12 Do Now Determine the molar mass of Ca(NO3)2 Ca: 1x40 = 40 N: 2x14 = 28 O: 6x16 = + 96 164g/mol Determine the %composition of H2SO4 H: 2x1 = 2 %H: H = 2 100 = 2% S: 1x32 = 32 H2SO4 98 O: 4x16 = +64 %S: S = 32 100 = 32.7% (all % add to 100) H2SO4 98 %O: O = 64 100 = 65.3% H2SO4 98 HW Review 2) Molecular Formula: C8H18 , so Empirical Formula is C4H9 4b) ?Empirical Formula 50.1%S 50.1gS 1molS = 1.5molS = 1 32gS 1.5 SO2 49.9%O 49.9gO 1molO = 3.1molO = 2 16gO 1.5 5) Given an experimental mass of 64g/mol, what is the molecular formula for 4b? Since the empirical mass and experimental mass are the same, the molecular formula is the name as the empirical formula. 6) Calcium Sulfate CaSO4 , then determine the % composition. Molecular Formula (actual formula)- a whole number multiple of the Empirical Formula 12/19/12 CHAPTER 8 Chemical Reaction - A chemical reaction is the process by which one or more substances change into one or more new substances. - Reactants are the original substances in a chemical reaction. - Products are the substances that are created in a chemical reaction. Evidence of a Chemical Reaction - Release of energy (heat, light, sound, electricity) - Formation of gas - Formation of a precipitate (solid) - Change in color or odor Symbols: Reactants yields Products (s)(l)(g)- physical state (aq)- (aqueous) dissolved in water heat or add energy catalyst name or formula of catalyst reversible reaction Example Chemical Reaction: - Na2O(s) + H2O(l) 1NaOH(aq) 1/3/13 Do Now 1) 2H2 + O2 2H2O 2) N2 + 3H2 2NH3 3) 6CO2 + 6H2O C6H12O6 + 6O2 4) 2HgO 2Hg + O2 Coefficients = Moles 1/7/13 Types of Chemical Reactions Single Displacement Rxn- one element displaces (or replaces) another in a compound Ex: A + BC=B + AC Ex: 2Al(s) + 3CuCl2(aq)3Cu(s) + 1AlCl3(aq) Double Displacement Rxn- positive and negative portions of 2 compounds are interchanged. Ex: AB + CD AD + CB Ex: HCl + NaOH HOH + NaCl (or H2O) Decomposition Rxn- substances break up into simpler substances when energy is applied. Ex: ABA + B Ex: 2H2O electricity 2H2 + O2 Synthesis Rxn- Two or more substances combine to form a new substance. Ex: A + B AB Ex: C + O2 CO2 Combustion Rxn- Oxidation reaction of an organic compound in which heat/energy is released. Every combustion has O2, CO2, and H2O. Ex: + O2 CO2 + H2O + energy Ex: C3H8 + 5O23CO2 + 4H2O + energy 1/8/12 Do Now - What is a hypothesis? An educated guess on what you think will happen. - What is kinetic energy? The energy of motion. - A chemical change produces a new substance. A physical change changes the state of a substance. HW Review (pg 285 1-10) 1) 2 smaller compounds come together to form a larger one. 2) They are very reactive so they replace other elements. 6) For a double displacement reaction, 2 new compounds must be produced. 8) a) single displacement b) synthesis c) decomposition d) double displacement e) single displacement f) combustion Mole Ratio- uses coefficients -Molesknown (or given)mole ratiomolesunknown -Molesgiven Molesunknown Mole Molesgiven ratio C6H12O6 + 6O26CO2 + 6H2O + energy Coefficients are in green, tell us the number of moles Ex: C6H12O6 + 6O2 6H2O + 6CO2 + Energy 1mol C6H12O6 6molO2 6molH2O Ex: 2Al + Fe3N2 2AlN + 3Fe Given: 6molAl 6molAl Unknown: ?molFe Unknown: ?molAlN 6molAl 6molCO2 3molFe = 9molFe 2molAl 2molAlN = 6molAlN 2molAl 1/9/13 Do Now Write 2,350,000 in scientific notation: 2.35x106 Metric System unit for Mass: grams Volume: liters Length: meters What are the three subatomic particles? p+, no, enucleus electron cloud Examples: Br2 + Cl2 2BrCl - 2.74molCl2 ?molBrCl - - 239.7gCl2 779.87gBrCl ?gBrCl 2.74molCl2 239.7gCl2 1molCl2 71gCl2 2molBrCl = 5.48molBrCl 1molCl2 2molBrCl 115.5gBrCl= 1molCl2 1molBrCl 4.53x1025molecCl2 4.53x1025molecCl2 1molCl2 1molBr2 160gBr2= 1.20x104gBr2 ?gBr2 6.02x1023Cl2 1molCl2 1molBr2 1/10/13 Do Now - Mg-25 p+=12, no=13, e- =12 Write the electron configuration for phosphorus: 1s22s22p63s23p3 Name the 4 orbitals? S P D F e- = 2 6 10 14 shape =1 3 5 7 - What was Thompson’s Atomic Model? plum pudding model/chocolate chip cookie model - How many valence electrons in Be? 2eDiatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2 1/11/12 Do Now - How is the periodic table arranged? By atomic number/number of protons - Who developed the modern periodic table? Mendeleev - Name of horizontal rows? Periods - Name of vertical columns? Groups - Name of: o Group 1- Alkali Metals o Group 2- Alkaline Earth Metals o Group 7- Halogens o Group 8- Noble Gases o D block elements- transition metals - Name elements right of the zigzag? nonmetals - Name elements left of the zigzag? metals Examples: Br2 + 5F2 2BrF5 R P - Diatomic Elements - Type of reaction- synthesis - Balancing coefficients- moles - Molar mass- mass in grams, of 1 mole of a substance - Avogadro’s Number- number of particles in 1 mole of a substance Useful Ratios: 1mol/MM or MM/1mol from periodic table 1mol/6.02x1023 or 6.02x1023/1mol Avogadro’s number THINGS WE NEED TO COVER: Electron Dots- put one dot on each side of the element’s symbol first, then go back and start doubling. Mg: (2 valence electrons) .. ·S: (6 valence electrons) ˙ .. :Ar: (8 valence electrons) (octet rule- there can be, at most, 8 valence electrons) ˙˙ Ca: (2 valence electrons) Yield Theoretical yield- the maximum quantity of product that a reaction could make if everything works perfectly. Actual yield-the mass of the product actually formed. Why? -The actual yield is less than the theoretical yield. - Many reactants do not completely use up limiting reactants. - Purification - Unwanted “side reactions” Limiting Reactant- the substance that controls the quantity of the product that can be formed in a chemical reaction. The limiting reactant forms the least amount of product. (runs out) Excess Reactant- the substance that is not used up completely in a reaction. % yield- actual yield x100 theoretical yield Example: Determine the limiting reactant, theoretical yield and %yield if 14.0gN2 are mixed with 9gH2 and 16.1gNH3 form. N2+3H22NH3 Molar mass’s- 14g 9g 16.1g (actual yield) 14gN2 1molN2 2molNH3 17gNH3 = 17gNH3 theoretical yield 28gN2 1molN2 1molNH3 (17 is smaller then 51 so N2 is the limiting reactant) 9gH2 1molH2 2molNH3 17gNH3= 51gNH3 2gH2 3molH2 1molNH3 %yield- actual yield = 16.1g x100 = 94.7% theoretical yield 17g