Final Exam Review - 271 Unit 1 Atomic Structure Terms: mass number neutron sublevels average atomic mass atomic number electrons orbitals valence electrons proton ion isotopes kernel electrons 1. What is the mass, charge and location of an electron, proton and neutron? Mass Charge Location proton 1 amu positive nucleus neutron 1 amu neutral nucleus electron 0 amu negative outside nucleus 2. Assume these are all neutral atoms. Element Atomic # Mass # Neutrons Electrons Ar 18 38 20 18 Rb 37 72 35 37 Si 14 29 15 14 Na 11 22 11 11 3. List the number of protons, neutrons and electrons for the following ions. Element Protons Neutrons Electrons 12 13 10 8 9 10 4. Neon has 2 isotopes. Neon-20 has a mass of 19.992 u and neon-22 has a mass of 21.991 u. In any sample of 100 neon atoms, 90 will be neon-20 and 10 will be neon-22. What is the average atomic mass of neon? Mass % Abundance Neon - 20 19.992 x 90 = 17.9928 Neon – 22 21.991 x 10 = + 2.1991 20.1919 u 5. What is the average atomic mass of silicon if 92.21 % of its atoms have mass 27.97 u, 4.70% have mass 28.976 u, and 3.09 % have mass 29.974 u? Mass % Abundance Si - 27.97 x 92.21 = 25.79 Si - 28.976 x 4.70 = 1.362 Si - 29.974 x 3.09 = + .9262 28.08 u 6. What were some problems with Dalton’s Atomic Theory? Solid mass, no isotopes 7. What did Rutherford’s experiment prove? The existence of a positively charged nucleus. 8. How many electrons are in an s orbital? 2 9. Which sublevels are in the 4th energy level? s, p, d, f 1s 2s 2p 3s 3p 4s 10. Al 11. Write electron configurations for the following elements using the Diagonal Rule. Identify the energy level of the valence shell, the number of valence electrons and the number of kernel electrons. a. Zn 1s22s22p63s23p64s23d10 c. I 1s22s22p63s23p64s23d104p65s24d105p5 VS = 4 VE = 2 KE = 28 VS = 5 VE = 7 KE = 46 b. Ba 1s22s22p63s23p64s23d104p65s24d105p66s2 d. Cd 1s22s22p63s23p64s23d104p65s24d10 VS = 6 VE = 2 KE = 54 VS = 5 VE = 2 KE = 46 Unit 2 Periodic Table Terms: metals/nonmetals transition elements groups/periods ionization energy atomic radius Periodic Law 1. What were Mendeleev’s contributions to the Periodic Table? Initial organization/Predicted undiscovered elements. 2. Who arranged the periodic table to its current format? Henry Moseley 3. Write the symbol of 2: a. metals K, Fe d. noble gases He. Ne g. alkaline earth metals Mg, Ca b. nonmetals S, O e. transition metals Ag, Mo h. inner transition c. halogens F, Br f. alkali metals Rb, Li metals U, Ce 4. Write complete electron configurations and noble gas notation for the following elements using the Periodic Table. Identify the energy level of the valence shell, the number of valence electrons and the number of kernel electrons. a. Li [He]2s1 VS=2 VE=1 KE= 2 c. Fe [Ar]4s23d6 VS=4 VE=2 KE= 24 b. P [Ne]3s23p3 VS=3 VE=5 KE= 10 d. Ba [Xe]6s2 VS=6 VE=2 KE= 54 5. What are the trends for atomic radius and ionization energy? Atomic radius increases top to bottom and decreases left to right. Ionization energy is the opposite. 6. Which is larger? a. S or O b. Na or K c. Ca or Ca2+ Unit 3 Chemical Bonding Terms: double bond cation covalent bond molecular geometry anion ionic bond 1. Draw dot diagrams for Ca, S, and K1+ Ca S [ K ]1+ 2. Describe how the ionic bond forms between an atom of potassium and an atom of chlorine. Potassium transfers 1 electron to the chlorine in order to bond. Potassium becomes a 1+ ion and chlorine (chloride) becomes a 1- ion. 3. Describe how a covalent bond forms between an atom of hydrogen and an atom of chlorine. Hydrogen and chlorine share electrons to form a bond. There is an unequal sharing (polar covalent bond) since chlorine has a higher electronegativity than hydrogen but they are still sharing. No ions are formed. 4. Why do atoms form bonds? [Hint: Octet Rule] Atoms form bonds in order to become more stable in their electron configuration like the nearest noble gas. Metals and nonmetals achieve this through ionic bonding while nonmetals bond by covalent bonding. 12. Draw a structure and identify the molecular geometry of the following: a. NH3 c. CH4 b. H2O d. CCl4 Unit 4 Study Terms: chemical formula empirical formula ternary compound binary compound traditional naming system molecular formula stock naming system diatomic molecule binary ionic binary acid 13. Classify the compound. Write the name when given the formula or vice-versa. Identify the name and number of cations and anions for ionic compounds. Identify the number of atoms of each element for ternary ionic compounds. Write the alternate name for molecular compounds. a. Hydrochloric acid HCl h. Fe2(CO3)3 iron(III)carbonate o. MgCl2 magnesium chloride b. MgSO3 magnesium sulfite i. acetic acid HC2H3O2 p. copper(I)oxide Cu2O c. ammonium oxalate (NH4)2C2O4 j. NH4Cl ammonium chloride q. CaC2O4 calcium oxalate d. sodium chloride NaCl k. AgOH silver hydroxide r. strontium bromide SrBr2 e. iron(II)oxide FeO l. aluminum sulfide Al2S3 s. AlCl3 aluminum chloride f. Sr(NO3)2 strontium nitrate m. H2SO4 sulfuric acid t. Titanium(IV)oxide TiO2 g. nitric acid HNO3 n. Rubidium oxide Rb2O 14. What is the molecular and empirical formula for the following structural formulas? C6H6 C2H4O2 CH2O CH Unit 5 Measurement Terms: scientific notation rules for calculating qualitative/quantitative accuracy/precision 1. What is the difference between quantitative and qualitative? Quantitative uses numbers to describe a measurement, while qualitative is a general description without numbers. 2. What is the difference between accuracy and precision? Accuracy is a measure of how close the answer is to the accepted value, while precision is being able to get a consistent result with each trial. 3. Express these in scientific notation: a. 400 4.00 x 102 b. .0404 4.04 x 10-2 c. 600 x 103 6.00 x 105 d. 25.3 x 10-4 2.53 x 10-3 4. Calculate: a. (6.0 X 104)(2.0 X 105)= 1.20 X 1010 b. (6.0 x 106)/(2.0 x 104)= 3.00 x 102 Unit 6 The Mole Terms: Avogadro’s number percentage composition formula mass molar mass Molarity atomic mass empirical/molecular formula 1. Calculate the formula mass of the following compounds: a. N2O5 108g b. CO2 44 g c. H3PO4 98 g 2. Do the following mole conversion problems: a. A sample of neon has a volume of 11.2 L. How many moles does this represent? 11.2 L Ne x 1 mol Ne = 0.5 mol = 5.00 x 10-1 mol Ne 22.4 L Ne b. How many atoms are there in 2.75 moles of copper? 2.75 mol Cu x 6.02 x 1023 atoms Cu = 1.66 x 1024 atoms Cu 1 mol Cu c. What is the mass of 1.51 x 1024 atoms of carbon? 1.51 x 1024 atoms C x 1 mol C x 12 g C = 30.09 g C = 3.01 x 101 g C 23 6.02 x 10 atoms C 1 mol C d. Determine the mass of 213 L of NO2? 213 L NO2 x 1 mol NO2 x 46 g NO2 = 437.4 g NO2 22.4 L NO2 1 mol NO2 e. Calculate the mass of 2.26 x 1024 molecules of PCl3 at STP? 24 2 2.26 x 10 molecules PCl3 x 1 mol PCl3 x 137.5 g PCl3 = 516.2 g = 5.16 x 10 g PCl3 23 6.02 x 10 molecules 1 mol PCl3 3. Determine the percentage composition by mass for the following compounds. a. CO 43% C b. NaCl 39.3% Na 2.04% H c. H2SO4 57% O 60.7% Cl 32.7% S 65.2% O 4. What is the empirical formula of a compound, given that a 212.1 g sample of the compound contains 42.4 g of hydrogen and 169.7 g of carbon? What is the molecular formula of the compound, given it has a molecular mass of 30.0 u. 42.4 g H x 1 mol H = 42.4 mol H/14.08 = 3 mol H 1gH empirical = CH3 30.0 u/15.0 = 2 169.7 g C x 1 mol C = 14.08 mol C/14.08 = 1 mol C f.m. = 15 2(CH3) = C2H6 = molecular 12 g C 5. Calculate the empirical and molecular formula of a substance that is 38.7% C, 9.6% H and 51.6% O. The molecular mass of the compound is 62 u. 38.7 g C x 1 mol C = 3.225 mol C/3.225 = 1 mol C 12 g C empirical = CH3O 62 u/31 = 2 9.6 g H x 1 mol H = 9.6 mol H/3.225 = 3 mol H f.m. = 31 2(CH3O) = C2H6O2 = molecular 1gH 51.6 g O x 1 mol O = 3.225 mol O/3.225 = 1 mol O 16 g O 6. The molecular mass of a compound is 118.0 u and its empirical formula is C2H3O2. What is the molecular formula of this compound? 118.0 u / 59 = 2 2(C2H3O2) = C4H6O4 7. What is the molarity of 2.3 moles of sodium chloride, NaCl in 0.45 liters of water? M= mol 2.3 mol L 0.45 L = 5.11 M solution 8. What is the molarity of 98 grams of sodium hydroxide, NaOH in 2.2 liters of water? 98 g NaOH x 1 mol NaOH = 1.11 M solution 2.2 L 40 g NaOH 9. How many grams of sodium sulfate are needed to make 0.75 L of a 0.25 M sodium sulfate, Na2SO4. M = mol mol = M x L so… 0.75 L x 0.25 mol Na2SO4 x 142.1 g Na2SO4 = 26.64 g L 1L 1 mol Na2SO4 10. What is the molarity of a solution that contains 2.5 L of solvent and 660 grams of calcium phosphate, Ca3(PO4)2? 660 g Ca3(PO4)2 x 1 mol Ca3(PO4)2 = 0.851 M solution 2.5 L 310.3 g Ca3(PO4)2 Unit 7 Stoichiometry Terms: reactants yields endothermic products exothermic coefficient perfect mixing ratio limiting reactant types of reactions 1. Solid sodium reacts with oxygen gas to produce solid sodium oxide. Write a word equation and a formula equation, using the correct chemical formulas and phase symbols. Balance the equation. sodium(s) + oxygen (g) sodium oxide(s) 4Na(s) + O2 (g) 2Na2O (s) 2. Write the word and the formula equation for the decomposition of solid calcium carbonate into solid calcium oxide and carbon dioxide gas. Balance the equation. calcium carbonate(s) calcium oxide(s) + carbon dioxide(g) CaCO3 (s) CaO(s) + CO2(g) 3. Balance the following equations: a. 5O2 + Sb2S3 Sb2O4 + 3SO2 b. Cu + Cl2 CuCl2 Balanced c. CuO + H2 Cu + H2O Balanced d. 2Sb + 3H2O Sb2O3 + 3H2 e. 2Re + 3Br2 2ReBr3 4. Which type of reactions are these? a. 2K(s) + Cl2(g) 2KCl(s) synthesis b. Fe2O3(s) + 2Al(s) Al2O3(s) + 2Fe(s) single replacement c. HNO3 (aq) + NaOH (aq) H2O (l) + NaNO3 (aq) double replacement d. PbO2 (s) Pb (s) + O2 (g) decomposition 5. How many liters of oxygen gas at STP can be produced by the decomposition of 36.0 grams of water? The balanced equation is as follows. 2H2O 2H2 + O2 36.0 g H2O x 1mol H2O x 1 mol O2 x 22.4 L O2 = 22.4 L O2 18.0 g H2O 2 mol H2O 1 mol O2 6. How many grams of water vapor will be produced when 56.5 L of hydrogen at STP reacts completely with oxygen according to the balanced chemical equation? 2H2 + O2 2H2O 56.5 L H2 x 1 mol H2 x 2 mol H2O x 18 g H2O = 45.4 g H2O 22.4 L H2 2 mol H2 1 mol H2O 7. What volume of chlorine gas, measured at STP, can be produced by the decomposition of 2.15 x 1024 molecules of hydrogen chloride gas given the following balanced chemical equation? 2HCl H2 + Cl2 24 2.15 x 10 molecules HCl x 1 mol HCl x 1 mol Cl2 x 22.4 L Cl2 = 40 L Cl2 23 6.02 x 10 molecules HCl 2 mol HCl 1 mol Cl2 8. Determine the number of molecules of sulfur trioxide that is produced when 64 L of oxygen gas at STP reacts with excess sulfur according to this balanced chemical equation. S + 3O2 2SO3 23 24 64 L O2 x 1 mol O2 x 2 mol SO3 x 6.02 x 10 molecules SO3 = 1.15 x 10 molecules S 22.4 L O2 3 mol O2 1 mol SO3 NOTE: Review all labs. At least 1 question related to each lab will appear on the final.