ࡱ> 35./012 Lbjbj DN{{{#^+^+88,9,9,9@9@9@98x9:@9]e(<(P<f<f<f<QNQNQNddddddd$[i le,9QN]MQNQNQNe88f<f<ePPPQN8Rf<,9f<dPQNdPPb 9"cf< Nbd-e0]eb6lUOrl,cc&l,9$dPQNQNQNeePQNQNQN]eQNQNQNQNlQNQNQNQNQNQNQNQNQN^+ ~7: AP Chemistry Unit I Targets Chapters 1-3 By the end of this unit you should be able to . . . chapter 1 1. define and provide examples for each of the following terms: physical property, chemical property, physical change, chemical change, intensive property, extensive property element, compound, mixture 2. differentiate between the three states of matter. 3. list the commonly used metric prefixes and their meanings. 4. determine the number of significant digits in a measured quantity and determine the appropriate number significant digits in a calculation. 5. convert temperatures between Celsius and Kelvin. 6. perform calculations involving density. 7. convert between units by using dimensional analysis. chapter 2 1. differentiate between protons, neutrons, and electrons in terms of charge, mass and location in an atom. 2. I can determine the number of protons, neutrons and electrons in isotopes and in ions. 3. describe the works of John Dalton, J.J. Thomson (cathode ray tube), Robert Millikan (Oil Drop Experiment) and Ernst Rutherford (Gold Foil Experiment). 4. use the periodic table to predict the charges of monatomic ions. 5. use the periodic table to predict whether an element is a metal, nonmetal or a metalloid. 6. write the names and formulas of ionic compounds, molecular/covalent compounds and acids. 7. calculate the atomic weight of an element given the abundances and masses of its isotopes. 8. distinguish between empirical formulas, molecular formulas and structural formulas. chapter 3 1. Predict the products for and write balanced equations for the following types of reactions: combustion, decomposition, sysnthesis (called combination reactions), single displacement and double displacement reactions. 2. interconvert between the number of moles and mass of a substance. I can also use Avogadros number and molar mass to calculate the number of particles (atoms, molecules or formula units) making up a substance. 3. calculate the percentage composition of a compound by mass. 4. calculate the empirical formula of a compound, having been given either: a) mass or % composition, or b) the mass of CO2 and H2O produced by combustion. 5. calculate the molecular formula, having been given the empirical formula and the molecular weight. 6. use stoichiometry to solve problems involving chemical reactions. 7. determine the limiting reactant in a reaction and determine the amount of excess reactant left over from a reaction. 8. calculate the theoretical and actual yields of a chemical reaction when given the appropriate data. Target 1. I can define and provide examples for each of the following terms: physical property, chemical property, physical change, chemical change, intensive property, extensive property element, compound, mixture physical property chemical property physical change chemical change intensive property extensive property - element compound mixture Describe the following mixture separation techniques: Decant - Filtration - Distillation  Chromatography Target 2: I can differentiate between the three states of matter. StatePictureMovementShapeVolumeCompression Gas  Liquid  Solid  Target 3: I can list the commonly used metric prefixes and their meanings. Memorize highlighted UnitGigaMegaKiloDeciCentiMilliMicroNanoAngstromPicoFemtoSymbol Meaning  Metric approximations - name 2 objects with a . . . a) mass of 1 gram d) length of a 1 mm b) mass of 1 kg e) volume of 1 mL c) length of 1000 cm f) volume of 1 liter List the best metric unit that would be used to measure each of the following: a) The distance from Atlanta to Greenville, SC. __________ b) The mass of your body. ________ c) The volume of coffee held in a cup of coffee. __________ d) The length of a soccer field. _________ e) The length of this sheet of paper. ________ Target 4: I can determine the number of significant digits in a measured quantity and determine the appropriate number significant digits in a calculation. Rules For Significant Digits 1) All non-zero numbers are significant. 2) Exact numbers are assumed to have an infinite amount of significant digits. Exact numbers are often seen in conversion factors, such as, 3 feet = 1 yard. 3) SANDWICH RULE - any zero between 2 significant digits is significant 4) RIGHT/RIGHT RULE - zeros to the right of a decimal point AND to the right of another significant digit is significant Note: Placing a number in scientific notation removes any ambiguity about whether or not a zero is significant. If a zero is used in scientific notation, then it must be significant. Examples: Underline all of the significant digits in the following measurements: a) 0.0090 mm d) 10.00 tons b) 102 L e) 2.50 X 10-3 grams c) 30,400 miles f) 50 angstroms Multiplication and Division Rule The result will have the same number of significant digits as the measurement with the fewest number of significant digits. Example: Calculate the density (in g/cm3)of a rectangular solid which has the following measurements: height: 2.00 cm width: 150 mm length: 0.0500 m mass: 300.0 grams Addition and Subtraction Rule The result cannot have more significant digits to the right of the decimal point than any of the original numbers. Example: Add the following numbers (first convert all of the measurements to the same unit!) 1.000 m + 175 cm + 0.0010 km = ________________ Note: When performing a multi-step calculation, always retain at least one additional significant digit for each of the intermediate answers. Why are the sig figs for pH different? i.e. Why is 12.3 only one sig fig? _______________ Target 5: I can convert temperatures between Celsius and Kelvin. Special note about units of temperature: Kelvin and Celsius are metric units of temperature and one degree of Kelvin is equal to one degree centigrade. The Kelvin scale is based upon absolute zero. -273(C = 0 Kelvin The Celsius scale is based upon the boiling and freezing point of water. 0(C = freezing point of water and 100(C = boiling point of water. Formula: K = oC + 273 Example 1: convert 195 Kelvin to Celsius. Example 2: convert 55(C to Kelvin. Target 6: I can perform calculations involving density. D = mass  volume Density is a property of matter that is widely used to characterize a substance. 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: A cylinder has Diameter = 25.0 mm Calculate its density in g/cm3. Height = 59.0 cm Mass = 38.0 g The volume of a cylinder is found by using the formula V = (r2h. Target 7: I can convert between units by using dimensional analysis. Conversion factor: __________________________________________________ Example 1: A typical rate of deposit of dust from unpolluted air is 10.0 tons per square mile per month (30 days). 1 ton = 907.2 kg 1 mile = 1.609 km What is this dust fall, in milligrams per square kmeter per hour? 30 days x 24 hours = 720 hrs 1 day 10 tons x 907.2 kg x 1000g x 1000 mg x 1 mile x 1 mile = 3.50 x 10 9mg/km2 Mile2 1 ton 1 kg 1 g 1.609 km 1.609 km 3.50 x 10 9mg/km2 / 720 hrs = 4.87 x 106 mg/km2/hr If dust has an average density of 2.0 g/cm3, how many hours would it take to accumulate a layer of dust 1.0 mm thick? Problems Solve the following using dimensional analysis. What is 50.0 mL in liters? How many moles of Ca(C2H3O2)2 are present in a 3980 mg of the compound? (a) A solution of AlCl3 has a concentration of 0.4405 mol/L, how many grams of aluminum chloride are dissolved in 150 mL of the solution? (b) How many atoms of chlorine are in the same 150 mL AlCl3 solution? A bug travels at the rate of 3.0 miles/hour. How fast is this in meters per second? Hint: 2.54 cm = 1 inch and 1 mile = 5,280 feet 5. The concentration of CO in a room is 48 (g/m3. What mass in grams is present in a room which measures 8.0 x 12.0 x 22 feet? 6. Ben Franklin showed that 1 teaspoon of oil would cover about 0.50 acre of still water. If you know that 1.0 x 104 m2 = 2.47 acres, and that there are 5.0 cm3 in a teaspoon, what is the thickness (in cm) of a layer of oil? 7. An empty 3.00 L bottle weighs 1.70 kg. Filled with wine it weighs 4.55 kg. The wine contains 11.0% ethyl alcohol by mass. How many ounces of ethyl alcohol are present in a 400.0 mL glass of wine? (1 lb = 16 oz = 453.6 g) AP Chem Unit I Quick Check For Understanding Directions: Label each of the following statements as either true or false. If false, change the statement to make it true. 1. Air, gaseous carbon dioxide, and light are all examples of matter. 2. Gases can be compressed quite considerably, liquids can be compressed fairly easily, and solids are nearly incompressible. 3. The particles in a solid can move. 4. Pure substances are either elements, compounds, or homogeneous mixtures. 5. Melting point, boiling point, and density . . . all three of these properties could be correctly classified as being both intensive and physical properties of matter. 6. The gram, kilometer, and hour are the SI units for mass, length, and time respectively. 7. 150 femtometers is a longer distance than 1.5 picometers. 8. All of the following represent absolute zero: -273oC, 0 Kelvin, and -419.4oF 9. Your AP chem teacher gave you a choice in the lab . . . your results could be either accurate or precise, but not both. The best choice would be for results to be precise. 10. If you were to round 62,421 km to 3 significant digits, it would be 624 km. 11. [23.05 - (14.0000 + 6.050)] X 1159 = ________ ; The correct answer to this problem is 1.802 X 104. AP Chem Notes Chapter 2 Target 1: I can differentiate between protons, neutrons, and electrons in terms of charge, mass and location in an atom.  The nucleus of the atom is centrally located and is comprised of protons and neutrons. The electrons are found in regions of space outside the nucleus called orbitals. Protons and neutrons have a relative mass of 1 amu (1.67 x 10-27 kg). The proton has a charge of +1 and the neutron has no charge. The electron has a relative mass of zero (9.11 x 10-31 kg) and a charge of -1. Target 2: I can determine the number of protons, neutrons and electrons in isotopes and in ions. Isotopes - atoms of a the same element (so they have the same atomic number) which have different masses, mass numbers, number of neutrons, and physical properties # neutrons = mass # - atomic # The number of protons determines the identity of the element, but its an atoms electrons (valence) which determine the chemical properties of the element. Isotopic Notation  Type I Protons Neutrons Electrons Carbon-14 6 8 6 Carbon-13 6 7 6 Oxygen-15 8 7 8 Type II C-4 6 8 10  O-2 8 7 10 Target 3: I can describe the works of John Dalton, J.J. Thomson (cathode ray tube), Robert Millikan (Oil Drop Experiment) and Ernst Rutherford (Gold Foil Experiment). Describe the contributionsof the following scientists: Democritus (400 B.C.) -vs- Aristotle B- John Dalton (early 1800s) 4 postulates of his atomic theory 1) 2) 3) 4) C- Henri Becquerel (1896) D- J.J. Thomson (1897) E- Robert Millikan (1909) F- Ernest Rutherford (1910) G- James Chadwick (1932) Target 4: I can use the periodic table to predict the charges of monatomic ions. Cations = ions with a positive charge formed by a metal atom losing one or more electrons Na (atom) Na+ (ion) 11 protons 11 protons 11 electrons 10 electrons Anions = ion with a negative charge formed by a nonmetallic atom gaining 1 or more electrons Charges of common monatomic ions Group # Group name: 1 ions +1 charge Alkali metals 2 ions +2 charge Alkaline earth metals 13 ions +3 charge (usually just aluminum) 14 ions +/-4 charge 15 ions -3 charge (Usually just N and P) 16 ions -2 charge Chalcogens (Oxygen group) 17 ions -1 charge Halogens 18 - (no charge) Noble gases / rare gases / inert gases Target 5: I can use the periodic table to predict whether an element is a metal, nonmetal or a metalloid. Periodic Table Notes Groups & families = vertical columns Periods = horizontal rows Describe and give location on the chart: Metals - Nonmetals - Metalloids - --------------------------------------------------------------------------------------------------------------------- Elements found as diatomic molecules in nature include: hydrogen, oxygen, fluorine, bromine, iodine, nitrogen, chlorine H2 O2 F2 Br2 I2 N2 Cl2 Mnemonics: These were discovered by Prof. HOFBrINCl or was his name BrINClHOF? I bring clay from our new house. I have no bright or clever friends I bring clams from oceans near here. --------------------------------------------------------------------------------------------------------------------- Target 6: I can write the names and formulas of ionic compounds, molecular/covalent compounds and acids. Ionic Compounds - compound made up of cations and anions Naming ionic compounds: 1) Name cations first - cation has the same name as metal; transition metals often need roman numerals to indicate their charge 2) Name monatomic anions by dropping their ending and adding suffix of -ide. Polyatomic anions most commonly end in -ate but can end in -ite; do not change the ending of a polyatomic ion. The prefix per- indicates the anion contains one more oxygen than the common ion. The prefix hypo- indicates that the oxyanion contains one less oxygen atom than the anion which ends in -ite. NaClO4 = NaClO2 = NaClO3 = NaClO = Naming Acids prefix suffix example Contains 1 more oxygen HClO4 than the common ion Contains the common ion HClO3 Contains 1 less oxygen HClO2 than the common ion Contains 2 less oxygen HClO than the common ion Contains no oxygen HCl (called binary acids) Naming Binary Molecular Compounds 1) Name the element farthest to the left in the periodic table first. If both elements are in the same group, name the lower one first. 2) Use suffix -ide on the second element. 3) Use the Greek prefixes in front of both elements (except if you have just one of the first element). Greek prefixes 1-10: Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca Name the following binary molecular compounds: N2O3 ____________________________ SO2 ________________________________ Name or write the formulas for each of the following: _________ 1. phosphoric acid ________________________ 11. K2SO4 _________ 2. dinitrogen heptoxide ________________________ 12. Cu2CO3 _________ 3. calcium acetate ________________________ 13. (NH4)2S _________ 4. iron (II) chromate ________________________ 14. P2O5 _________ 5. magnesium hydride ________________________ 15. H2CO2 _________ 6. hydrofluoric acid ________________________ 16. Na3AsO4 _________ 7. potassium dichromate ________________________ 17. N2O _________ 8. ammonium sulfite ________________________ 18. LiMnO4 _________ 9. hyposulfurous acid ________________________ 19. KOH _________ 10. tin (IV) phosphate ________________________ 20. CoCO3 Name each of the following: NO3- __________________________ NO3 _________________________ Target 7: I can calculate the atomic weight of an element given the abundances and masses of its isotopes. Atomic mass unit (amu or u) = a unit used to express small masses; 1 amu = 1.66054 X 10-24 g. The amu is defined by assigning a mass of exactly 12 amu to the C-12 isotope. ATOMIC MASSES: The atomic mass of an element is the weighted average of the masses of the isotopes of that element. A weighted average reflects both the mass and the relative abundance of the isotopes as they occur in nature. Most elements occur as two or three isotopes in nature. For example, chlorine has two isotopes, both of which have 17 protons in their atomic nuclei. One isotope has 18 neutrons (Cl-35) and the other isotope has 20 neutrons (Cl-37). In order to calculate the atomic mass (also known as the average atomic mass) you use the following formula: (mass)(% abundance) + (mass)(% abundance) . . . . = atomic mass 100 example: isotope mass (amu) % ab Cl-35 34.969 75.53 Cl-37 36.966 24.47 atomic mass = (34.969 amu)(0.7553) + (36.966 amu)(0.2447) = 35.46 amu Example 1: Magnesium has three naturally occurring isotopes. The data for these isotopes are given below. Calculate the atomic mass of magnesium. Isotope Percent abundanceMass (amu)Mg-2478.7023.985Mg-2510.1324.985Mg-2611.1725.983 Example 2: Gallium has two isotopes: Ga-69 (mass of 68.9255 amu) and Ga-70 (mass of 70.9247). If the average atomic mass of gallium is 69.723 amu, what is the relative abundance (in %) for each of the two isotopes of gallium? Practice: What is an isotope? What is the atomic weight of Lithium if Lithium has 2 isotopes: Li-6 (7.5%) and Li-7 (92.5%)? Why doesnt the periodic table use a normal average (mean)? (In other words why doesnt C-12, C-13 and C-14 just average out to 13?) What is the difference between a normal average (mean) and a weighted average? Chlorine has 2 major isotopes, Cl-35 and Cl-37. What is the %abundance of each isotope? Silver has 2 major isotopes, Ag-107 and Ag- 109. What is the % abundance of each isotope? Copper has 2 major isotopes, Cu-63 and Cu-65. Without doing any major calculations, which isotope is more abundant and why? Target 8: I can distinguish between empirical formulas, molecular formulas and structural formulas. Empirical Formula - the simplest whole number ratio of atoms of each element present in a compound. Lets look at ethene, C2H4 Ethene can undergo polymerization and become a long strand of linked ethene molecules. The empirical formula would be C2H4 but the polymer may have a molecular formula several hundred times larger than the building block of the substance. Molecular Formula - indicates the number of atoms of each element found in each discrete molecule of that compound. An example of polyethene: C2H4 x 150 = C300H600 Structural Formula - a graphical representation of the molecular structure, showing how the atoms are arranged. Used extensivily in organic chemistry. Ethene would look like this:  INCLUDEPICTURE "http://upload.wikimedia.org/wikipedia/commons/thumb/b/b5/Ethylene.svg/82px-Ethylene.svg.png" \* MERGEFORMATINET  Polyethene would look like this  INCLUDEPICTURE "http://pslc.ws/macrog/kidsmac/images/pe02.gif" \* MERGEFORMATINET  AP Chemistry Name ________________________ Review Chapter 1 and Chapter 2 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 means? 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. 32F = ________________C = ___________________ K 8. 17.2 cm + 204.8 mm = ________________ mm 9. Which of John Daltons 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 Temp) 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. MnO4- b. Cr(OH)3 c. NBr3 d. NH4ClTarget 1: I can predict the products for and write balanced equations for the following types of reactions: combustion, decomposition, sysnthesis (called combination reactions), single displacement and double displacement reactions. 1- ALKALI METALS WITH WATER M (s) + H2O (l) ----------> MOH (aq) + H2(g) Na(s) + H2O (l) ----------> NaOH (aq) + H2(g) K(s) + H2O (l) ----------> KOH (aq) + H2(g) 2- COMBUSTION IN AIR Hydrocarbon (fuel) + oxygen -------> CO2 + H2O CH4(g) + O2(g) -----------> CO2(g) + H2O(l) C3H8(g) + O2(g) -----------> CO2(g) + H2O(l) CH3OH(l) + O2(g) -----------> CO2(g) + H2O(l) 3- COMBINATION REACTIONS - 2 or more reactants form one product A + B -------> C (generic equation) 2 Mg(s) + O2(g) --------------> 2 MgO(s) N2(g) + 3 H2(g) -----------> 2 NH3(g) 4- DECOMPOSITION REACTIONS - one reactant forms 2 or more products A ----------> B + C (generic equation) 2 KClO3(s) -----------> 2 KCl(s) + 3 O2(g) ; chlorates decompose to form chloride salts and oxygen gas 2 NaN3(s) ----------> 2 Na(s) + 3 N2(g) CaCO3(s) ---------> CaO(s) + CO2(g) ; metal carbonates decompose to form metal oxides and carbon dioxide gas Target 2: I can interconvert between the number of moles and mass of a substance. I can also use Avogadros number and molar mass to calculate the number of particles (atoms, molecules or formula units) making up a substance. Avogadros Number: A number equal to the number of atoms in exactly 12 grams of C-12. Experimentally we have found this number to be 6.022 X 1023. Molar mass Avogadros # GRAMS <--------------------> MOLES <--------------------> PARTICLES (atoms, ions, molecules) Example: How many moles of CO2 are in 10.0 grams of CO2? Practice 1: How many molecules of H2O are contained in 150. grams of water? SYW! Practice 2: How many hydrogen atoms are there in 36.0 grams of sulfuric acid? SYW! Formula weight = sum of atomic weights of the atoms in a formula; we can use the term molecular weight if the substance is a molecular formula or molar mass for either. Example: Find the formula weight of calcium nitrate. calcium nitrate is Ca(NO3)2 1 Ca atom = 1(40.1 amu) = 40.1 amu 2 N atoms = 2(14.0 amu) = 28.0 amu 6 O atoms = 6(16.0 amu) = 96.0 amu ___________ 164.1 amu Practice: Find the formula weight of ammonium phosphate. Target 3: I can calculate the percentage composition of a compound by mass. Percent Composition: % composition is the % by mass contributed by each element in the substance (# of atoms of element)(atomic weight) X 100 molar mass of the compound Example: Find the percentage of nitrogen in calcium nitrate. (Answer is 17.1% N) Practice: Find the percentage of water in copper sulfate pentahydrate. Target 4: I can calculate the empirical formula of a compound, having been given either: a) mass or % composition, or b) the mass of CO2 and H2O produced by combustion. mass % ------------> grams of each ---------------> moles of each --------------> empirical element element formula Example: A compound is 62.58% C, 9.63 % H, and 27.79 % O by mass. Calculate its empirical formula. Practice: What is the empirical formula of a compound which contains 25.9% nitrogen and 74.1% oxygen? (N2O5) Determining the empirical formula of a compound from combustion analysis: Example: Combustion of a 0.2000 gram sample of vitamin C (which contains C, H, and O) yields 0.2998 g of CO2 and 0.0819 g of H2O. Calculate the empirical formula of vitamin C. Practice: Combustion analysis of toluene, a common organic solvent, gives 3.53mg of CO2 and 0.822mg of H2O. If the compound contains only carbon and hydrogen, what is its empirical formula? Target 5: I can calculate the molecular formula, having been given the empirical formula and the molecular weight. Example: A compound is found to contain 16.66 grams of carbon and 3.49 grams of hydrogen. a. Find the empirical formula. (C2H5) b. The molar mass is 87. Find the molecular formula: (C6H15) Practice: Determine the molecular formula for a compound that contains 22.5 % Na, 30.4 % P and 47.1 % O if it has a molar mass of 306 g/mole. Target 6: I can use stoichiometry to solve problems involving chemical reactions. Stoichiometry is a process which uses balanced equations to solve problems. Example: How many grams of O2 can be prepared from the decomposition of 4.50 grams of potassium chlorate? Target 7: I can determine the limiting reactant in a reaction and determine the amount of excess reactant left over from a reaction. Example: In order to bake 1 cake you need the following ingredients: 1 box of cake mix 1 egg 2 cups of flour 1/2 cup of sugar 3 tablespoons of water Assume that you had 1 dozen eggs, 6 boxes of cake mix, 10 cups of flour, 10 cups of sugar, one quart of oil, and an unlimited supply of water. How many cakes can you make? What is the limiting reagent? Example: A 2.00 gram piece of zinc is placed in 2.50 grams of aqueous silver nitrate. a) Write the balanced equation. b) Determine the limiting reagent. c) How many grams of silver will be formed? d) How many grams of the excess reactant will be left over? Target 8: I can calculate the theoretical and actual yields of a chemical reaction when given the appropriate data. Percent Yield Problems Definitions: theoretical yield = the quantity of product that should theoretically be produced (if all of the reactant is used up) actual yield = the quantity of product that is actually produced when the reaction is done in the lab percent yield = the relationship between the actual yield and the theoretical yield percent yield = actual yield X 100 theoretical yield Example: Fe2O3(s) + 3 CO(g) -------> 2 Fe(s) + 3 CO2(g) If 150.0 grams of Fe2O3 reacts with excess CO, 87.9 grams of Fe is produced. Calculate the % yield. Practice: Nitrogen gas reacts with hydrogen gas to make ammonia. 15.5 L of Nitrogen reacts at STP to make 30 L of ammonia. What is the % yield? Law of Definite and Multiple Proportions Experimental Demonstration Purpose To experiemntally determine the emperical formula of a copper iodide compound. The compound will either be copper (I) iodide, CuI, or copper (II) iodide, CuI2. Your instructor will demonstrate a reaction between copper and iodine in the hood. The procedure is as follows: Procedure 1. Mass an empty beaker. 2. Clean the surface of a Cu strip with steel wool. Place the clean Cu strip into the empty beaker and determine the mass of both. 3. Bend the strip into a concave shape, and sprinkle a small amount of iodine onto it. Under a fume hood and using tongs, heat gently to react the iodine with the copper. The unreacted iodine will leave as a dense, purple gas. The copper iodide product will remain on the slip as a creamy white solid. 4. Carefully place the strip into the beaker and find the mass. 5. Remove the strip from the flask and remove the compound from the strip by scraping it off. 6. Place the cleaned Cu strip in the beaker and find the mass. Data mass of copper before . . . . . . . . . . . _______________ g mass of copper & compound . . . . . . _______________ g mass of copper after . . . . . . . . . . . . . _______________ g Calculations Using the data above, show by calculations whether the formula of the compound is CuI or CuI2. Discussion Questions Using complete sentences, explain how the ratio of copper to iodine would be altered if: a) some of the copper iodide had been knocked off before being massed. b) some of the copper iodide had been knocked off after being massed. c) In some procedures, after the copper iodiide is scraped off, the copper strip is rinsed with water and then dried off. How would the copper to iodine ratio be altered if the strip were not completely dried? d) some scientists dry the strip by heating it over a bunsen burner. If heated too long some dark copper (II) oxide can form on the surface. Explain how this would affect the Cu:I ratio? AP Chemistry Ch. 3 Review Sheet Name_____________________ Directions: Answer the following questions on a separate sheet of paper. Show all of your work, label your final answers with correct units, and circle your final answers! 1. Predict the products for the following reactions and write a balanced equation for each: a) The combustion of glucose. b) The synthesis reaction between potassium and chlorine gas. c) The decomposition of magnesium carbonate. d) Reacting magnesium oxide and water. e) Reacting sulfur trioxide and water. f) The decomposition of sodium chlorate. 2. Below is a chart containing data for the three naturally occurring isotopes of Mg: Isotope abundance (%) mass (u) Mg-24 78.70 23.98504 Mg-25 10.13 24.98584 Mg-26 11.17 25.98259 Calculate the atomic mass of magnesium. 3. Calculate the percentage of oxygen (by mass) in nickel (II) acetate. 4. Assume you have 5.0 liters of water. Calculate the number of each of the following: a) the number of grams of water. b) the number of moles of water. c) the number of molecules of water. d) the number of hydrogen atoms in this sample of water. 5. Antifreeze is composed of 51.6 % oxygen, 9.70% hydrogen, and 38.7% carbon by mass. The molar mass of antifreeze is 62.1 g/mol. Calculate its empirical and molecular formulas. 6. Menthol, the substance we can smell in mentholated cough drops, is composed of C, H, and O. A 1.005-g sample of menthol is combusted, producing 0.2829 g CO2 and 0.1159 g of H2O. What is the empirical formula of menthol? If the compound has a molecular mass of 156 g/mol, what is its molecular formula? 7. When a mixture of 10.0 g of acetylene, C2H2, and 10.0 g of oxygen, O2, is ignited, the resultant combustion produces CO2 and H2O. a) Write the balanced equation for this reaction. b) Which reactant is the limiting reactant? c) How many grams of C2H2, O2, CO2, and H2O are present after the reaction is complete?      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