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AP Chemistry Summer Assignment About this assignment: This assignment is meant to review the essential skills and knowledge that was covered in Regular/Honors Chemistry. There are four units to this assignment: •Unit 1: •Unit 2: •Unit 3: •Unit 4: Matter and measurement Atomic Structure Molecules and Compounds Chemical Equations and Stoichiometry Each unit is preceded by a list of learning targets. Pay attention to these targets! Not every learning target has problems that go with it. Read the textbook! Make sure that after you complete a unit that you feel confident that you have mastered each target. It benefits you to take this assignment seriously and master these skills/concepts. We will not have time to review them during AP Chemistry and if you do not understand them thoroughly then you will struggle all year. When This Assignment is due: This assignment is due the first day of school and you will have a test covering the entire assignment at the end of that week. Use the learning targets as your guide to prepare for the test! Resources at Your Disposal You have the following resources available to help you with this assignment •Your Textbook: You had a textbook assigned to you before summer started. It is an excellent resource and you should become accustomed to learning from it. We will use it a great deal throughout the year •Videos: I have posted videos on select topics on my teacher webpage in the Online Resources section. They go through several examples of the types of problems I have assigned. If you cannot find my webpage, email me and I will send you the link. •AP Chemistry Boot Camp: AP Chemistry Boot Camp will run July 14th-17th. You are highly encouraged to sign up. We will mostly be covering Units 3 and 4 during boot camp, but if time permits I can take questions from Units 1 and 2. •Email Me: I typically check my school email daily throughout the summer (although I will be away from the internet a few days here-and-there). If you have a specific question, please feel free to email me and I will respond as soon as I can. •Each Other: You are encouraged to study with a partner so you can help each other understand the material. While studying with a friend will help you complete the assignment quicker, remember that in the end it is you that is ultimately responsible for understanding the material. So make sure your friend helps you, but doesn’t do it for you. Unit 1 -Matter and Measurement MATH SKILLS LEARNING TARGETS Metric System You know the metric system. You know the meaning of the metric prefixes, kilo-, centi-, and milli-. You know that there are other metric prefixes and can look them up if needed (micro, mega, pico, etc.) You can convert one measurement into another (e.g., 0.532 cg = ______ mg). You can convert squared or cubed units (e.g., knowing that 2.54 cm = 1 inch, 38.5 in2 = _____ cm2). Dimensional Analysis & Showing Your Work When you convert one unit to another, you can show your work using dimensional analysis or unit analysis. You know that good examples of dimensional analysis are changing metric units, converting time units, or using density to convert mass to volume or volume to mass. You know that you should always show enough work so that if your answer is incorrect, I can tell where you went wrong. Scientific Notation You can translate regular numbers into scientific notation and numbers written in scientific notation into normal notation. You know the distinction between exponential notation and scientific notation. Making Measurements You can use a ruler or other measuring device to make a measurement to the correct number of significant figures, i.e. include all of the digits in the measurement that are a significant part of the measurement. You can correctly assign a value when making a given measurement. You always include a unit on a measurement. You know the distinction between a measurement and a defined number (e.g., 12 things in a dozen, pi). You can explain the difference between accuracy (how close a measurement is to a true or accepted value) and precision (how close a set of measurements are to each other). Significant Figures You can determine the number of significant figures in a given measurement (i.e., you know whether a “0” in a measurement is significant or not.) You can determine the precision in a calculation involving measurements when the measurements are written with the correct number of significant figures. You can determine the precision in a calculation involving measurements when the measurements are written with notation. Unit 1 -Matter and Measurement PROPERTIES OF MATTER Define the following: physical property – chemical property – physical change – chemical change – intensive property – extensive property element – compound – mixture – Differentiate between the three states of matter. State Gas Liquid Solid Picture Movement Shape Volume Compression List the commonly used metric prefixes and their meanings. Unit Symbol Giga Mega Kilo Deci Centi Milli Micro Nano Angstrom Pico Femto Meaning Perform the following calculations involving density. Example 1: A rectangular solid has the following dimensions: length: 2.00 cm width: 75.0 mm height: 4.00 cm mass: 436 grams Using the given data, is the solid most likely silicon, tin, or gold? Example 2: Assume you had a silver sphere with a mass of 1.50 kg. Calculate the diameter of the sphere (in cm). The density of silver is 10.5 g/cm3. The formula for the volume of a sphere is . . . V = 4/3 π r3. Convert temperatures between Celsius and Kelvin. Formula: K = oC + 273 Example 1: convert 195 Kelvin to Celsius. Example 2: convert 55C to Kelvin. DIMENSIONAL ANALYSIS PRACTICE You can rent a bucket of golf balls at the driving range to practice your swing. Make the following conversion factor cards and use them to set up the problems below. Bucket Facts: 1 bucket 918 grams of g.b. 918 grams of g.b. 1 bucket 1 bucket 20 golf balls 20 golf balls 1 bucket 1 bucket 2 Liters of g.b. 2 Liters of g.b. 1 bucket 1000 grams 1 kg 1 kg 1000 grams 1. Calculate the mass of 3.05 buckets of golf balls. Given: 3.05 buckets x _____________ = ___________________ 2. How many individual golf balls would be in 8.75 buckets? Given: 8.75 buckets x _____________ = _____________________ 3. Calculate the volume (in Liters) of 300. golf balls. Given: 300. golf balls x ______________ x ____________________ = ______________________ 4. Calculate the mass of a Liter of golf balls. Given: 1.00 Liter of golf balls x _________________ x _________________ = _________________ 5. Calculate the mass (in kg) of one golf ball. Given: 1.00 golf ball x ______________ x _______________ x _________________ = ____________ 6. How many golf balls would you need to have a mass of 575 grams? 7. What is the volume, in Liters, of 1500 g of golf balls? 8. How many buckets would you need to hold 375 kg of golf balls? 9. What is the volume (in Liters) of 155 golf balls? 10. Determine the volume of 13.2 kg of golf balls. 11. How many golf balls would have a mass of 19,9 kg? 12. What is the mass (in grams) of 50.0 L of golf balls? SCIENTIFIC NOTATION & UNIT ANALYSIS Change the following to Scientific Notation (maintain the number of significant figures): 1. 5.280 = _______________ 11. 2,560 = _______________ 2. 2,000 = _______________ 12. .0009 = _______________ 3. 15 = _______________ 13. 8,900,000 = _______________ 4. 6,589,000 = _______________ 14. .0920 = _______________ 5. 70,400,000,000 = _______________ 15. 6,300 = _______________ 6. .00263 = _______________ 16. .90 = _______________ 7. .00589 = _______________ 17. 250 = _______________ 8. .006 = _______________ 18. .006087 = _______________ 9. .400 = _______________ 19. 500,000 = 10. .08060 = _______________ 20. .0000000105 = _______________ _______________ Make the following Metric System conversions using “unit analysis” (you may use scientific notation): 1. 100 mg _______________ = _______________ g 2. 20 cm _______________ = _______________ m 3. 50 L _______________ = _______________ kL 4. 22 g _______________ = _______________ cg 5. 825 cm _______________ = _______________ km 6. 2,350 kg _______________ = _______________ g 7. 19 mL _______________ = _______________ cL 8. 52 km _______________ = _______________ m 9. 36 m _______________ = _______________ cm 10. 18 cm _______________ = _______________ mm 11. 6g _______________ = _______________ mg 12. 4,259 mg _______________ = _______________ g SIGNIFICANT FIGURES A. When taking measurements all certain digits plus the first uncertain number are significant. Example: Your bathroom scale weighs in 10 Newton increments and when you step onto it, the pointer stops between 550 and 560. Your look at the scale and determine your weight to 557 N. You are certain of the first two places, 55, but not the last place 7. The last place is a guess and if it is your best guess it also is significant. B. When given measurements, the numbers that are significant are the digits 1-9 and the 0 when it is not merely a place holder. 1. When 0’s are between sig. fig., 0’s are always significant. Example: 101 has 3 sig. fig. and 34055 has 5 sig. fig. 2. When the measurement is a whole number ending with 0’s, the 0’s are never significant. Example: 210 has 2 sig. fig. and 71,000,000 also has 2 sig. fig. 3. When the measurement is less than a whole number, the 0’s between the decimal and other significant numbers are never significant (they are place holders). Example: .0021 has 2 sig. fig. and .0000332 has 3 sig. fig. 4. When the measurement is less than a whole number and the 0’s fall after the other significant numbers, the 0’s are always significant. Example: .310 has 3 sig. fig. and .3400 has 4 sig. fig. 5. When the measurement is less than a whole and there is a 0 to the left of the decimal, the 0 is not significant. Example: 0.02 has only 1 sig. fig. and 0.110 has 3 sig. fig. 6. When the measurement is a whole number but ends with 0’s to the right of the decimal, the 0’s are significant. Example: 20.0 has 3 sig. fig., 18876.000 has 8 sig. fig. In case 4 and 6 the 0’s have no effect on the value (size) of the measurement. Therefore, these 0’s must have been included for another reason and that reason is to show precision of the measurement. Since these 0’s show precision they must therefore be significant. In cases 2 and 3 removal of the 0’s DO change the value (size) of the measurement, the 0’s are place holders and are thus not significant. In case 5 the 0 is completely unnecessary, it is neither a place holder nor adds to the accuracy of the measurement. UNCERTAINTIES IN CALCULATIONS 1. When adding or subtracting numbers written with the notation, always add the uncertainties and then round off the value to the largest significant digit. Round off the answer to match. Example: (22.4 .5) + (14.76 .25) = 37.16 .75 = 37.2 .8 The uncertainty begins in the tenths place… it is the last significant digit. 2. When adding or subtracting numbers written in significant figures, show the uncertainty by rounding the answer to match the largest place with uncertainty. Example: 267 + 11.8 = 278.8 = 279 The least accurate original measurement is only accurate to the ones place. 4. When multiplying or dividing measurements written in significant figures, show the uncertainty of your calculations by rounding off your answer to match the same number of significant figures as your least precise measurement (the measurement with the least number of significant figures). Example: 477.85 32.6 = 14.657975 = 14.7 32.6 is the least accurate measurement with only 3 significant figures. NOTE: There are two types of precision: “absolute precision” and “relative precision.” Example: 322.45 x 12.75 x 3.92 = 16116.051 = 16100 All the measurements are accurate to the hundredth place (absolute precision) but the answer is rounded to 3 significant figures because 3.92 has only 3 significant figures (relative precision). In Summary: Adding and Subtracting Multiplying and dividing #’s with notation Rule 1 Don’t Do This Case #’s with significant figures Rule 2 Rule 4 SIGNIFICANT FIGURES & ROUNDING Indicate the number of significant figures then round each to the number of significant figures indicated. For example: 1.234 has ______4___ significant figures and, rounded to 2 significant figures, is ___1.2____ 1. 0.6034 has __________ significant figures and, rounded to 2 significant figures, is __________ 2. 12,700 has __________ significant figures and, rounded to 2 significant figures, is __________ 3. 12,700.00 has __________ significant figures and, rounded to 1 significant figures, is __________ 4. 0.000983 has __________ significant figures and, rounded to 2 significant figures, is __________ 5. 123342.9 has __________ significant figures and, rounded to 5 significant figures, is __________ 6. 6.023 x 1023 significant figures, is __________ 7. .005600 has __________ significant figures and, rounded to 1 significant figures, is __________ 8. 10000.5006 has __________ significant figures and, rounded to 5 significant figures, is __________ 9. 2.0 x 10-3 has __________ significant figures and, rounded to 1 significant figures, is __________ 10. 3.456110 has __________ significant figures and, rounded to 3 significant figures, is __________ has __________ significant figures and, rounded to 2 Given calculations with the calculator answer, write the answers with the appropriate number of significant figures. Example: 6.00 x 3.00 = 18 The answer should be ______18.0_____ 1. 23 + 46 = 69 The answer should be _______________ 2. 23.0 + 46.0 = 69 The answer should be _______________ 3. 253 + 345.8 = 598.8 The answer should be _______________ 4. 56 – 35 = 21 The answer should be _______________ 5. 56.00 – 35.0 = 21 The answer should be _______________ 6. 46 x 12 = 552 The answer should be _______________ 7. 3.24 x 5.63 = 18.2412 The answer should be _______________ 8 (2.355 + 2.645) x 10.00 = 50 The answer should be _______________ 9 654 32 = 20.4375 The answer should be _______________ = 1.512 x 10-03 The answer should be _______________ 10. .024 x .063 Unit 2 -Atoms and Elements LEARNING TARGETS The Development of the Atomic Theory: I can: State the four “signpost scientists”, their experiments, what they added to the atomic theory, and the name of their model. Define the three theories that Dalton explained in terms of atoms: o Law of Conservation of Matter o Law of Definite/Constant Proportions o Law of Multiple Proportions Give examples and solve calculation problems related to each of the three theories. Sketch a cathode ray tube as demonstrated in class and state how J.J. Thomson’s experiments led to the idea that atoms have positive and negative parts, the negative parts are all the same, and the negative parts (called electrons) have a certain charge/mass ratio. Define cathode rays. State the factors that determine how much a moving charged particle will be deflected by an electric or magnetic field. Explain Millikan’s oil drop experiment & how it added to the atomic theory. Sketch the set-up used by Ernest Rutherford (the gold-foil experiment), show what he observed, and explain how these observations led to the idea that most of the mass of the atom is concentrated into a tiny, amazingly massive, positively-charged nucleus. Parts of the Atom: State the three particles that make up an atom, their symbol, their charge, their mass, and their location. State the number of protons, neutrons, and electrons in any atom or ion. Explain that isotopes are two atoms with the same atomic number (number of protons) but different mass numbers (number of nucleons—protons + neturons). Represent the nucleus with isotopic notation, 220 such as: 86 Rn Recognize when two nuclei are isotopes of each other. Molar Mass Calculations: Calculate the isotopic mass of an atom given the resting mass of protons and neutrons. Explain that a mole of any element is actually made up of various isotopes in constant percentage abundance. Calculate the average atomic mass of an element using the percent abundance and mass of each isotope. Calculate the percent abundance of isotopes given the average atomic mass and isotopic masses of an element. The Families of the Periodic Table: List the common families of the periodic table and recognize to which family any element belongs. Recognize metals, non-metals, and metalloids (semi-metals) on the periodic table. State and define the terms conductivity, malleability, ductility, and sectility. State some element facts such as which elements are too radioactive to exist, which is the largest non-radioactive element, which element has the greatest density, and which element has the highest melting point. Explain how Dmitri Mendeleev put together the periodic table and why we give him credit for the table even though others were working along the same lines. List the three elements that Mendeleev predicted and where they are located on the periodic table. A Little Nuclear Chemistry: State that Henri Becquerel discovered radioactivity and Marie Curie studied it. List the three “Becquerel rays” (alpha, beta, and gamma) and state why alpha particles were the perfect tool for Ernest Rutherford to study the structure of atoms. State that the alpha particle is the same as a helium nucleus, a beta particle is a high-speed electron, and a gamma ray is a high-energy form of light. Unit 2 -Atoms and Elements ATOMIC HISTORY Describe the works/experiments of the following: John Dalton (early 1800’s) 4 postulates of his atomic theory 1) 2) 3) 4) Henri Becquerel (1896) J.J. Thomson (1897) Robert Millikan (1909) Ernest Rutherford (1910) ATOMIC STRUCTURE Most people already know that the atom is made up of three main parts, the _______________ and ______________ in the nucleus and the ______________ somewhere outside of the nucleus. Let’s summarize: proton symbol charge location mass size (see below) neutron electron The atom is often represented as a miniature ___________ _____________. Draw it: How Many Particles in Each Atom? The particle that defines the identity of an atom is the _____________. (shown on the periodic table) Every hydrogen atom has ___ proton. Every magnesium atom has ___ protons. Any atom that has 23 protons is _________________. Any atom that has 92 protons is _________________. The mass of an atom is mostly from the ___________ and ____________. Find O on the periodic table. It’s mass is ______ amu. It has ___ protons. It must have ___ neutrons. Electrically neutral atoms (as opposed to ions) have one electron for every proton. Fill in this chart for these neutral atoms: Atom Mass protons neutrons electrons He Si Be The mass of the atom is due to the _____________________________ H Rn Ar F The size of the atom is due to the __________________ Pb If the mass is not close to a whole number, it is because the atom has several _____________. These are atoms with the same number of ___________ but different numbers of _____________. Chlorine has two isotopes: Cl-35 ( ___ p+ & ___ n) and Cl-37 ( ___ p+ & ___ n). THE NUCLEAR ATOM ALL of the answers to this worksheet can be logically figured out by looking at the Schematic Diagrams for Various Atoms , the Periodic Table, and discussing with your partners. All of the information you need is here somewhere. Determine each answer and be able to give convincing reasons for each answer. Good luck. C? 1. How many protons are found in 12 13 13 2. How many neutrons are found in 12 13 13 3. How many electrons are found in 12 13 13 4. Based on the model, a) what do all carbon atoms (and ions) have in common? b) C? C? C? C? C? C? C? C? what do all hydrogen atoms (and ions) have in common? 5. What is the significance of the atomic number, Z, above each atomic symbol in the periodic chart? 6. What do all nickel (Ni) atoms have in common? 7. How is the mass number, A, (left-hand superscript next to the atomic symbol) determined? 8. What structural feature is different in isotopes of a particular element? 9. a) What feature distinguishes a neutral atom from an ion? b) How is the charge on an ion determined? 10. Where is most of the mass of an atom, within the nucleus or outside of the nucleus? Explain your reasoning. 11. Complete the following table: Isotope 31 18 P Atomic Number Z O Ni2+ Number of electrons 15 8 19 58 Mass Number A 39 58 18 ISOTOPES 1. Give the mass number of each of the following atoms: (a) an iron atom with 30 neutrons (b) an americium atom with 148 neutrons (c) a tungsten atom with 110 neutrons 2. Give the complete symbol ( AZ X ) for each of the following atoms: (a) nitrogen with 8 neutrons (b) zinc with 34 neutrons (c) xenon with 75 neutrons 3. How many electrons, protons, and neutrons are there in an atom of: (a) carbon-13, 13 C (b) copper-63, 63 Cu (c) bismuth-205, 205 Bi 4. Fill in the blanks in the table (one column per element). 65 86 Symbol Cu Kr Number of protons Number of neutrons Number of electrons in the neutral atom Name of element 78 117 46 36 5. Radioactive americium-241 is used in household smoke detectors and in bone mineral analysis. Give the number of electrons, protons, and neutrons in an atom of americium-241. 6. Which of the following are isotopes of element X, with atomic number of 9: 7. Verify that the atomic mass of magnesium is 24.31 amu, given the following information: 24 Mg , mass = 23.985042 amu; percent abundance = 78.99% 9. 25 Mg , mass = 24.985837 amu; percent abundance = 10.00% 26 Mg , mass = 25.982593 amu; percent abundance = 11.01% Copper has two stable isotopes, 63 Cu and 65 19 9 X , 209 X , 189 X , and 21 9 X. Cu , with masses of 62.939598 amu and 64.927793 amu, respectively. Calculate the percent abundances of these isotopes of copper. 10. Strontium has four stable isotopes, Strontium-84 has a very low natural abundance, but 88 86 Sr are all reasonably abundant. Which of these more abundant isotopes predominates? Sr , 87 Sr , and STUDY QUESTIONS AND PROBLEMS: 1. Explain, at an atomic or molecular level, what happens when a. water freezes to form ice b. copper and tin combine to form bronze c. rainwater evaporates from the pavement 2. Which of the following atoms are isotopes of the same element? Identify the elements of these isotopes and describe the number of protons and neutrons in the nucleus of them all. 15 7 3. X 12 6 X 13 7 X 18 8 X 14 7 X 14 6 X 16 8 X 13 6 X 17 8 X There are three naturally occurring isotopes of neon: neon-20 neon-21 neon-22 mass 19.9924 amu mass 20.9940 amu mass 21.9914 amu abundance 90.84% abundance 0.260% abundance 8.90% a. Without calculation, what is the approximate atomic mass of neon? b. Calculate the actual atomic mass. 4. Uranium has an atomic mass equal to 238.0289. It consists of two isotopes: uranium-235 with an isotopic mass of 235.044 amu and uranium-238 with an isotopic mass of 238.051. Calculate the % abundance of the uranium-235 isotope. 5. From amongst the elements sodium, chlorine, nickel, argon, calcium, uranium, and oxygen, select the alkali metal, the alkaline earth metal, the transition metal, the actinide, the halogen, the noble gas, and the chalcogen (Group 6A). Unit 3 - Molecules & Compounds LEARNING TARGETS I can: Formulas Look at a formula and state how many elements and atoms are in that compound. Calculate the molecular mass or molar mass of any compound. State that the mass of a molecule is measured in amu’s and the mass of a mole is measured in grams. Give examples of empirical formulas, molecular formulas, and structural formulas. Identify a formula as empirical, molecular, or structural. Ionic Compounds I can: Name and determine the charge of a monoatimic ion from its place on the periodic table List the names the common polyatomic ions. State whether a compound is an ionic compound or a nonmetal compound. Write the formula of an ionic compound given the two ions or its name. Know when to use parentheses. Name an ionic compound given the Name a binary nonmetal compound (molecular compound) given its formula. Percent Composition Calculate the percent composition (by mass) for any compound. Calculate the empirical formula from percent composition data. Determine the molecular formula of a compound given its empirical formula and molar mass. Hydrates Give examples of hydrates and anhydrous compounds. Calculate the formula of a hydrate from dehydration data. The Mole State the significance of the mole. State the three mole facts for any substance (molar volume, molar mass, Avogadro’s number) 1 mole = 22.4 Liters @ STP (gases only) 1 mole = 6.02 x 1023 particles (particles = molecules or atoms) 1 mole = gram molecular mass of chemical formula. Determine the charge on an ion from information in an ionic formula. Nonmetal Compounds aka Molecular Compound Write the formula of a binary nonmetal compound (molecular compound) given its name. Use dimensional analysis to convert between moles, mass, volume, and number of particles for a chemical. Use density as a conversion factor in mole problems. Use gas density to calculate molar mass. Unit 3 -Molecules & Compounds QUICK NOTES THE PERIODIC TABLE Groups & families Periods Group IA Group IIA Group VIA Group VIIA Group VIIIA = vertical columns = horizontal rows Alkali metals Alkaline Earth metals Chalcogens Halogens Noble gases / rare gases / inert gases Metals - elements found on the left side of the “stair case” on the periodic table as well as the Lanthanoids and Actinides on the bottom, good conductors of heat & electricity, ductile, malleable, solids at room temperature (except Hg) Nonmetals - elements found on the right side of the staircase, gases, liquid, & solid; usually poor conductors and are brittle Metalloids - elements that lie along staircase which have properties of both metals and nonmetals (except Al, which is usually considered a metal) IONS Cations = ions with a positive charge formed by a metal atom losing one or more electrons Na (atom) 11 protons 11 electrons Na+ (ion) 11 protons 10 electrons Anions = ion with a negative charge formed by a nonmetallic atom gaining 1 or more electrons Charges of common monatomic ions Group 1A ions +1 charge Group 2A ions +2 charge Group 3A ions +3 charge (usually just aluminum) Group 5A ions -3 charge (usually just nitrogen, sometimes P) Group 6A ions -2 charge Group 7A ions -1 charge DIATOMICS Elements found as diatomic molecules in nature include: hydrogen, oxygen, fluorine, bromine, iodine, nitrogen, chlorine H2 O2 F2 Br2 I2 N2 Cl2 These were discovered by Prof. HOFBrINCl or was his name BrINClHOF? POLYATOMIC NAMING PRACTICES Polyatomic ions that don't appear on the above tables do NOT always follow these naming practices. If you can remember the formula of the ion whose name ends with ate, you can usually work out the formulas of the other family members as follows: modify stem name with: meaning examples -ate a common form, containing oxygen chlorate, ClO3nitrate, NO3sulfate, SO42- -ite one less oxygen than -ate form chlorite, ClO2sulfite, SO32nitrite, NO2- per-, -ate same charge, but contains one more oxygen than -ate form perchlorate, ClO4perbromate, BrO4- hypo-, -ite same charge, but contains one less oxygen hypochlorite, ClOhypobromite, BrOthan the -ite form thio- replace an O with an S thiosulfate, S2O32thiosulfite, S2O22- Some anions can capture hydrogen ions. For example, carbonate (CO32- can capture an H+ to produce hydrogen carbonate HCO3- (often called bicarbonate). Each captured hydrogen neutralizes one minus charge on the anion. modify stem name with: meaning examples hydrogen or bi- (1) captured H+ ions hydrogen carbonate, HCO3- (a.k.a. bicarbonate) hydrogen sulfate, HSO4- (a.k.a. bisulfate) dihydrogen (2) captured H+ ions dihydrogen phosphate, H2PO4- WRITING FORMULAS AND NAMING COMPOUNDS Introduction Writing formulas and naming compounds can be confusing because there are different types of compounds that follow different rules. Additionally, some compounds (H2O, NH3, CH4, etc.) simply have common names that must be memorized. The two types of compounds we will focus on first are ionic compounds (formed from positive and negative ions) and binary nonmetal compounds (molecular compounds). Later we will add acids. So… you must recognize the type of compound before you try to name it. [Note: + ion = “cation” and – ion = “anion”.] Ionic Formula + ion before – ion ex: NaCl (NH4)2SO4 Al2S3 Name of cation + name of anion sodium chloride ammonium sulfate Naming aluminum sulfide I. Binary Nonmetal usually the less electronegative atom is first ex: CO CO2 N2O Indicate the number (mono, di, tri, and kind of atoms. First element is simply name of element. Second element name ends with “ide” carbon monoxide carbon dioxide dinitrogen monoxide Writing Ionic Formulas Cl NO3 S2 CO32 N3 PO43 Na+ NH4+ Sn2+ Hg22+ Al3+ Sn4+ II. Naming Ionic Compounds Cation Anion Cu2+ OH Ba2+ SO42 NH4+ Cr2O72 Ag+ C2H3O2 Fe3+ S2 Formula Name OH mono III. di tri tetra hexa hepta octa Formula Name deca Formula nitrogen trifluoride phosphorus trichloride nitrogen monoxide phosphorus pentachloride nitrogen dioxide sulfur hexafluoride dinitrogen tetroxide disulfur decafluoride dinitrogen monoxide xenon tetrafluoride IV. nona Writing Formulas of Binary Nonmetal Compounds Name Naming Binary Nonmetal Compounds Name V. penta Formula Name Formula CCl4 HBr P4O10 N2F4 ClF3 XeF3 BCl3 PI3 SF4 SCl2 Practice for Both Types of Compounds Formula Name Formula Name HCl carbon dioxide PCl5 ammonium carbonate K2S sulfur dichloride NiSO4 calcium iodide ClF3 boron trifluoride OF2 phosphorus triiodide Al(OH)3 magnesium perchlorate NCl3 potassium permanganate (NH4)3PO4 aluminum phosphate S2Cl2 dioxygen difluoride M O L A R M A S S & % C O M P O S I T I O N I. Molar Masses Given a periodic table, you should be able to calculate the molecular mass (in u’s) or the molar mass (in grams) for any element or compound. Examples: (give answers to two decimal places) H2SO4 Cl2 CO2 N2O Ca(OH)2 HC2H3O2 NaOCl Al2S3 II. Fraction and Percent Composition It is useful to determine how much of a compound’s mass is made up of each element. Water, H2O, for example has a molar mass of 18.02 g. The H’s mass is 2(1.0079) = 2.02 g. The O’s mass is 16.00 g. 2.02 16.00 We can set up fractions for each element: H = = 0.112 = 11.2%. O= = 0.888 = 88.8%. 18.02 18.02 This is called the percent composition. The fraction composition is a good in-between step. Determine the fraction and percent composition of each element below (answer to one decimal place): 1. H2SO4 2. Ca(OH)2 3. HC2H3O2 4. CO2 5. N2O 6. NaOCl 7. Al2S3 HYDRATES & COMPOSITION PROBLEMS 1. Cupric chloride, CuCl2, when heated to 100C is dehydrated. If 0.235 g of CuCl2 · x H2O gives 0.185 g of CuCl2 on heating, what is the value of x? 2. The “alum” used in cooking is potassium aluminum sulfate hydrate, KAl(SO4)2 · x H2O . To find the value of x, you can heat a sample of the compound to drive off all of the water and leave only KAl(SO4)2. Assume you heat 4.74 g of the hydrated compound and that the sample loses 2.16 g of water. What is the value of x? 3. If “Epsom salt,” MgSO4 · x H2O is heated to 250C, all the water of hydration is lost. On heating a 1.687g sample of the hydrate, 0.824 g of MgSO4 remains. What is the formula of Epsom salt? 4. When CaSO4 · x H2O is heated, all of the water is driven off. If 34.0 g of CaSO4 (molar mass = 136) is formed from 43.0 g of CaSO4 · x H2O, what is the value of x? Mole Calculations - Difficulty Level 1 1 mole = 6.02 x 1023 molecules = 22.4 L (@ STP) 1. Calculate the mass of 1.58 moles CH4. [molar mass CH4 = 16.0 g/mol] G: 1.58 moles CH4 D: ? g CH4 1.58 moles CH4 = 2. What volume will 7.29 moles of CO2 gas occupy at STP? G: 7.29 moles CO2 D: ? L CO2 7.29 moles CO2 = 3. How many molecules are there in a 0.00583 mole sample of H2O? G: 0.00583 moles H2O D: ? molecules H2O 0.00583 moles H2O = Mole Calculations - Difficulty Level 2 4. What mass of CO2 gas occupies a volume of 395 Liters at STP? [molar mass CO2 = 44.0 g/mol] G: D: = 5. How many molecules are in a 0.250 gram sample of H2O? [molar mass H2O = 18.0 g/mol] G: D: = 6. What volume will 3.01 x 1022 molecules of CH4 occupy at STP? G: D: = Mole Calculations - Difficulty Level 3 1 mole = 6.02 x 1023 molecules = 22.4 L (@ STP) 1. Calculate the mass of 7.23 moles CH4. [molar mass CH4 = 16.0 g/mol] G: D: 2. What volume will 9.35 moles of CO2 gas occupy at STP? G: D: 3. How many molecules are there in a 0.0752 mole sample of H2O? G: D: 4. What mass of CO2 gas occupies a volume of 10.8 Liters at STP? [molar mass CO2 = 44.0 g/mol] G: D: 5. How many molecules are in a 1.44 gram sample of H2O? [molar mass H2O = 18.0 g/mol] G: D: 6. What volume will 1.21 x 1024 molecules of CH4 occupy at STP? G: D: Review Questions: 1. _______ is the state of matter which has a definite volume but no definite shape. 2. Another name for a homogeneous mixture is a(n) ________________. 3. Can the elements in a compound be separated by chemical means? by physical 4. List 2 intensive properties and 2 extensive properties. 5. The SI unit for mass is the _______________; the SI unit for length is the _________. 6. Which of the following is the longest distance? 1 km, 106 m, or 10-6 Mm? 7. 32°F = ________________°C = ___________________ K 8. 17.2 cm + 204.8 mm = ________________ mm 9. Which of John Dalton’s postulates was incorrect about atomic theory? Why? 10. 1 meter = _________________ cm 11. Name an element which is: a. a gas and also a halogen (assume room T) b. a metalloid c. an alkaline earth metal 12. Write the empirical formula for glucose. 13. Write the formula for: a. a polyatomic anion b. a monatomic cation c. any molecular compound d. magnesium cyanide e. nickel (II) nitride f. hyponitrous acid 14. Name each of the following: a. MnO4b. Cr(OH)3 c. NBr3 d. NH4Cl means? Unit 4 -Chemical Equations and Stoichiometry LEARNING TARGETS I can: Chemical Equations Limiting Reactant Problems Give examples of products and reactants in a chemical equation. State that Antoine Lavoisier introduced the law of conservation of matter. Combustion State that combustion is another name for burning. Write an equation for a combustion reaction given only the fuel that is burned. Correctly label substances in an equation as solid (s) , gas (g), liquid (l), or aqueous (aq) Balancing Equations Balance equations by adding coefficients. Recognize when an equation is balanced. State that the formulas of reactants and products should not be changed in order to balance equations. Stoichiometry Problems Use the stoichiometric factor (mole ratio) to convert from moles of one substance to moles of a different substance. (i.e. In the equation: N2 + 3H2 2NH3, 3 mol H2 2 mol NH3) Convert between the quantities of mass, volume, molecules and moles using dimensional analysis (i.e. use 1 mol = 22.4 L, 1 mol = 6.02 x 1023 molecules, and 1 mol = gram molecular mass) Show the units of molar mass as grams/mol or g·mol-1. Recognize that a problem with two “given values” is a limiting reactant problem. Determine the limiting reactant and excess reactant in a problem. Solve problems involving Limiting Reactants Calculate how much excess chemical is left over after a reaction. Percent Yield Problems Use stoichiometry to calculate the theoretical yield (mass of a product) in a problem. State that actual yields are usually given in a problem. Use the theoretical yield and actual yield to calculate the percent yield. Chemical Analysis Problems Calculate the mass of each element in a given compound given data such as the masses of CO2 and H2O formed in a combustion reaction. Use mass and mole information to calculate the empirical formula of an unknown substance. Use percent composition to equalize mass and mole information derived from different samples. Unit 4 -Chemical Equations and Stoichiometry COMBUSTION EQUATIONS For burning to occur, you need a fuel, an oxidizer, and heat. When hydrocarbons are the fuel and O2 in the air is the oxidizer, then CO2 and H2O are the products. Example: Write the balanced equation for the complete combustion of propane, C3H8, in air. Solution: First, set up the basic equation. You memorize the “+ O2 CO2 + H2O” part. C3H8 + O2 CO2 + H2O Next, balance. 3 C’s in C3H8 result in 3CO2’s; 8 H’s in C3H8 result in 4 H2O’s; C3H8 + __ O2 3 CO2 + 4 H2O Total O’s on the product side = 10 [(3 x 2) + (4 x 1)] = total O’s on the reactant side. This would mean that 5 O2’s were involved. Tip: If an UNEVEN number of O’s need to be represented, a fraction should be used. 7 O’s = 7/2 O2 Tip: Take into account fuels that contain oxygen. Subtract the O’s from that represented as O2’s Practice: Write the balanced combustion equations for the following substances. 1. CH4 2. C5H12 3. C9H20 4. C2H6 5. C8H18 6. C4H10 7. C2H5OH 8. C3H7OH 9. HC2H3O2 10. CH3COCH3 BALANCING EQUATIONS Balance the following chemical equations: 1. __ZnS + __HCl __ZnCl2 + __H2S 2. __HCl + __Cr __CrCl2 + __H2 3. __Al + __Fe3O4 __Al2O3 + __Fe 4. __H2 + __Br2 __HBr 5. __Na2S2O3 + __I2 __NaI + __Na2S4O6 6. __LaCl3 + __Na2CO3 __La2(CO3)3 + __NaCl 7. __NH4Cl + __Ba(OH)2 __BaCl2 + __NH3 + __H2O 8. __Ca(OH)2 + __H3PO4 __Ca3(PO4)2 + __H2O 9. __La2(CO3)3 + __H2SO4 __La2(SO4)3 + __H2O + __CO2 10. __Na2O + __(NH4)2SO4 __Na2SO4 + __H2O + __NH3 11. __C4H10 + __O2 __CO2 + __H2O 12. __C7H6O2 + __O2 __CO2 + __H2O 13. __P4O10 + __H2O __H3PO4 14. __FeS2 + __O2 __Fe2O3 + __SO2 15. __NH3 + __O2 __NO + __H2O 16. __Fe + __HCl __H2 + __FeCl2 17. __PbO2 + __HCl __H2O + __PbCl2 + __Cl2 18. __Fe2O3 + __H2SO4 __Fe2(SO4)3 + __H2O 19. __NO2 + __H2O __NO + __HNO3 20. __C2H6S + __O2 __CO2 + __H2O + __SO2 STOICHIOMETRY General Stoichiometry 1. Several brands of antacid tablets use aluminum hydroxide to neutralize excess acid. Al(OH)3(s) + 3 HCl(aq) AlCl3(aq) + 3 H2O(l) [Molar masses: 78.01 36.46 133.4 18.02] What quantity of HCl, in grams, can a tablet with 0.750 g of Al(OH)3 consume? What quantity of water is produced? 2. The equation for one of the reactions in the process of reducing iron ore to the metal is Fe2O3(s) + 3 CO(g) 2 Fe(s) + 3 CO2(g) [Molar masses: 159.7 28.01 55.85 44.01] (a) What is the maximum mass of iron, in grams, that can be obtained from 454 g (1.00 lb) of iron(III) oxide? (b) What mass of CO is required to reduce the iron(III) oxide to iron metal? Limiting Reactants 3. The reaction of methane and water is one way to prepare hydrogen: CH4(g) + H2O(g) CO(g) + 3 H2(g) [Molar masses: 16.04 18.02 28.01 2.02] If you begin with 995 g of CH4 and 2510 g of water, what is the maximum possible yield of H2? 4 Disulfur dichloride, S2Cl2, is used to vulcanize rubber. It can be made by treating molten sulfur with gaseous chlorine: S8(l) + 4 Cl2(g) 4 S2Cl2(l) [Molar masses: 256.6 70.91 135.0] Starting with a mixture of 32.0 g of sulfur and 71.0 g of Cl2, which is the limiting reactant? What mass of S2Cl2 (in grams) can be produced? What mass of the excess reactant remains when the limiting reactant is consumed? Percent Yield 5. Diborane, B2H6, is a valuable compound in the synthesis of new organic compounds. One of several ways this born compound can be made is by the reaction 2 NaBH4(s) + I2(s) B2H6(g) + 2 NaI(s) + H2(g) [Molar masses: 37.84 253.8 27.67 149.9 2.02] Suppose you use 1.203 g of NaBH4 with an excess of iodine and obtain 0.295 g of B2H6. What is the percent yield of B2H6? 6. Disulfur dichloride, which has a revolting smell, can be prepared by directly combining S8 and Cl2, but it can also be made by the following reaction: 3 SCl2(l) + 4 NaF(s) SF4(g) + S2Cl2(l) + 4 NaCl(s) [Molar masses: 103.0 41.99 108.1 135.0 58.46] Assume you begin with 5.23 g of SCl2 and excess NaF. What is the theoretical yield of S2Cl2? If only 1.19 g of S2Cl2 is obtained, what is the percent yield of the compound? CHEMICAL ANALYSIS Show all of your work (on a separate sheet of paper if necessary) Chemical Analysis 32. A mixture of CuSO4 and CuSO4 • 5H2O has a mass of 1.245 g, but, after heating to drive off all the water, the mass is only 0.832 g. What is the weight percent of CuSO4 • 5H2O in the mixture? 34. A 1.25-g sample contains some of the very reactive compound Al(C6H5)3. On treating the compound with aqueous HCl, 0.951 g of C6H6 is obtained. Al(C6H5)3(s) + 3HCl(aq) AlCl3(aq) + 3C6H6(l) Assuming that Al(C6H5)3 was converted completely to products, what is the weight percent of Al(C6H5)3 in original 1.25-g sample? Determination of Empirical Formulas 36. Styrene, the building block of polystyrene, is a hydrocarbon, a compound consisting only of C and H. If 0.438 g of styrene is burned in oxygen and produces 1.481 g of CO2 and 0.303 g of H2O, what is the empirical formula of styrene? 38. Menthol, from the oil of mint, has a characteristic cool taste. The compound contains only C, H, and O. If 95.6 mg of menthol burns completely in O2, and gives 269 mg of CO2 and 110 mg of H2O, what is the empirical formula of menthol? 40. Silicon and hydrogen form a series of compounds with the general formula SixHy. to find the formula of one of them, a 6.22-g sample of the compound is burned in oxygen. On doing so, all of the Si is converted to 11.64 g of SiO2 and all of the H to 6.980 g of H2O. What is the empirical formula of the silicon compound?