Final Exam Review Chem 101

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1 Final Exam Review Chem Know your nomenclature. a) Know how to go from the name to the formula. b) Know how to go from the formula to the name. 1. Ionic compounds (binary and ternary) a. Example: What is the formula for Cobalt(III) phosphate? (CoPO 4 ) b. Example: What is the formula for Calcium nitride? (Ca 3 N 2 ) c. Example: What is the name for ZnSO 4? (Zinc sulfate) d. Example: What is the name for MnCl 2? (Manganese (II) chloride) 2. Binary molecular compounds a. Example: What is the name for P 2 S 5? (Diphosphorous pentasulfide) b. Example: What is the formula for Sulfur trioxide? (SO 3 ) 3. Simple acids a. Example: What is the name for HF? (Hydrofluoric acid) b. Example: What is the formula for Hydroselenic acid? (H 2 Se) 4. Oxoacids a. Example: What is the name for HBrO 3? (Bromic acid) b. Example: What is the formula for hydrobromic acid? (HBr) 5. Hydrates a. Example: What is the name for CuSO 4 3H 2 O? (Copper (II) trihydrate) b. Example: What is the formula for chromium (III) chloride hexahydrate? (CrCl 3 6H 2 O) 2. Know what ions the main group elements tend to for, and the formulas for the ionic compounds they for. a) Example: Which elements would likely form an ionic compound with the formula CaX? (oxygen, sulfur, or selenium) 3. Be able to predict relative melting points based on relative lattice enthalpy. a) Example: Which would likely have the strongest lattice enthalpy, SrBr 2 or MgF 2? (MgF 2 ) b) Example: Which would likely have the highest melting point, SrBr 2 or MgF 2? (MgF 2 ) 4. Be able to characterize bonds in a compound as either ionic or covalent. a) Example: Which has bonds that are covalent, CS 2 or AgCl? (CS 2 ) 5. Be able to recognize isotope pairs. a) Example, Which of the following would be an isotope of, or? ( ) 6. Be able to find the isotopic mass of one isotope of a given element. a) Example: Thallium has two naturally occurring isotopes, one with a mass of amu and a relative abundance of 29.50%. What is the mass of the other isotope? (205.0 amu) 1 of 7

2 7. Know how many protons, neutrons, and electrons are in a given isotope. a) Example: How many protons, neutrons, and electrons are there in 197 Au 3+? (79 p +, 118 n 0, 76 e ) 8. Know which elements have similar chemical properties to each other. a) Example: Which of the following elements has similar chemical properties to potassium, Cs, N, or Kr? (Cs) 9. Given the mass and molar mass, know how to find moles of an ion are in a compound. a) Example: How many moles of sodium ions are there in 62.1 g of Na 2 Cr 2 O 7? The molar mass of Na 2 Cr 2 O 7 is g/mol. (0.474 mol Na + ) 10. Know how to find the mass of an elements in a given number of molecules of that compound. a) Example: What is the mass of carbon atoms in 3.14 x molecules of C 3 H 6 O? The molar mass of C is g/mol. (1.88 x 10 4 g) 11. Know how to calculate mass percent of an element in a compound. a) What is the mass percent of potassium in K 3 PO 4? The molar mass of K 3 PO 4 is g/mol. (55.2%) 12.Given the empirical formula of a compound be able to calculate the mass percent of an element in the compound. a) If there are 2 phosphorous atoms for every 5 oxygen atoms in a compound, what is the mass percent phosphorous in that compound? (43.6%) 13. Be able to calculate the energy of a photon given it s wavelength. a) Example: What is the energy of a photon with wavelength 659 nm?(3.01x10 19 J) 14. Know what information is given by the quantum numbers. a) Example: What information does each quantum number give? (n = distance from nucleus and energy, l = shape of orbitals in that sub level and energy, m l = which orbital within a sublevel the electron is in, m s = spin state of the electron.) 15. Know the rules for the quantum numbers. a) Example: Give an allowed set of quantum numbers for a 3p electron in the ground state of aluminum? (n = 3, l = 1, m l = 1, m s = +1/2) (There are others too). 16. Know how to get electron configurations. a) Example: What is the ground state electron configuration for the Fe 2+ ion? ([Ar]3d 6 ) 17. Know the periodic trends (atomic radius, electronegativity, first ionization energy) and the reason for them. a) Example: What happens to atomic radius as you go up and to the right in the periodic table? (It decreases) b) Why? (Effective nuclear charge increases) 2 of 7

3 18. Know how to predict relative sizes and first ionization energies of an isoelectronic series. a) Example: Which of the following has the smallest radius: K +, S 2, Cl, or Ca 2+? (Ca 2+ ) 19. Know about multiple ionization energies, and the octet rule. a) Example: Which group of elements will have a large jump in their second ionization energy, IE(2)? (The alkali metals, group 1A) b) Example: Which group of elements will have a large jump in their third ionization energy, IE(3)? (The alkaline earth metals, group 2A) c) Example: Which group of elements will have a large jump in their fourth ionization energy, IE(4)? (Group 3A) 20. Be able to draw Lewis structures and answer questions about the resulting structures. a) Example: Draw the Lewis structure for XeO 2. How many nonbonding (lone) pairs of electrons are on the central atom? (2) b) Example: How many double bonds are there in XeO 2? (2) 21. Know how to calculate formal charge. a) Example: What is the formal charge on carbon in carbon monosulfide, CS? ( 1) 22. Know how to determine the existence of resonance forms and what they mean. a) Example: Draw the Lewis structure(s) for sulfate, SO 4 2. How many resonance forms are there? (6) b) Example: What do these resonance forms tell you about the actual structure of sulfate? (There is only one structure that is an average of the 6 resonance forms) 23. Know the relationship between single, double, and triple bonds and their length and strength. a) Example: Which is longer, an oxygen oxygen single bond or and oxygen oxygen double bond? (The single bond) b) Example Which is stronger, an oxygen oxygen single bond or and oxygen oxygen double bond? (The double bond) 24. Know how to use VSEPR theory to predict the molecular geometry and approximate bond angles of a compound. a) Example: What is the molecular geometry of XeCl 4? (Square planar) b) Example: What are the approximate bond angles in XeCl 4? (90 o ) 3 of 7

4 25. Know the difference between a nonpolar covalent, polar covalent, and an ionic bond. a) Example: How are the electrons in a nonpolar covalent bond shared? (Equally between the two atoms) b) Example: How are the electrons in a polar covalent bond shared? (Unequally between the two atoms) c) Example: How are the electrons in an ionic bond shared? (They re not, the electron density is almost entirely on one atom (the anion).) 26. Know how to predict the relative polarity of bonds. a) Which of the following bonds is most polar: S O, C H, or O F? ( S O) 27. Know how to determine whether a molecule is polar (has a permanent dipole moment) or not. a) Which of the following has a permanent dipole moment, CF 4, AlCl 3, CO 2, NH 3? (NH 3 ) 28. Know how to count and bonds. a) Draw the Lewis structure for CH 2 CCO. How many and bonds are there? (5 and 3 ) 29. Know how to predict the hybridization of an atom in a compound. a) Example: Draw the Lewis structure for CH 2 O. What is the hybridization of the carbon atom? (sp 2 ). 30. Be able to predict relative boiling points. a) Example: Which of the following is predicted to have the highest boiling point? H 2, CF 4, CH 3 F, CH 3 OH, CBr 3 OH (CBr 3 OH) 31. Know the difference between strong electrolytes, weak electrolytes, and nonelectrolytes. a) Example: What species are present in an aqueous solution of the strong electrolyte, hydrobromic acid (other than water)? (H + and Br ) 32. Know how to determine oxidation numbers and which species is oxidezed and which is reduced in a redox reaction. a) Example: Is Ag + being oxidized or reduced in the following reaction? What happens to it s oxidation number? (Reduced, decreases) 2Ag + + Cu Cu Ag 33. Know how to calculate oxidation numbers. a) Example: What is the oxidation number of Cr in K 2 CrO 4? (+6) 34. Know how to balance chemical equations. a) Example: When the following equation is balanced with the smallest set of whole numbers, what is the coefficient for sodium nitrate? (6) Pb(NO 3 ) 2 (aq) + Na 3 PO 4 (aq) NaNO3(aq) + Pb 3 (PO 4 ) 2 (s) 4 of 7

5 35. Know how to find a net ionic equation. a) Example: What is the net ionic equation for the following reaction? (Ag + (aq) + Cl (aq) AgCl(s) ) AgNO 3 (aq) + NaCl(aq) AgCl(s) + NaNO3(aq) 36. Know how to calculate the molarity of a given ion in a solutiion of a soluble ionic compound. a) Example: What is the molarity of sodium ions in a 2.50 M solution of Na 2 SO 4? (5.00 M) 37. Know how to calculate the molarity of a solution given the volume of solution and mass of solute. a) Example: What is the molar concentration of a 275 ml solution that has 7.23 g of NaF? Take the molar mass of NaF as g/mol. (0.626 M) 38. Know how (and when) to use the dilution equation. a) Example: How many ml of a 9.79 M KOH solution is required to make 75.0 ml of a 1.50 M KOH solution? (11.5 ml) 39. Know your stoichiometry. a) Example: What is the concentration of 20.0 ml of a sulfuric acid solution that requires 27.4 ml of a M NaOH solution to reach the equivalence point? (0.290 M) b) Example: If 140 molecules of H 2 are mixed with 75 molecules of O 2 and react according to the following equation, how many molecules of each species will be in the reaction vessel once the reaction goes to completion? (0 molecules H 2, 5 molecules O 2, 140 molecules H 2 O) 2H 2 (g) + O 2 (g) 2H 2 O(g) c) Example: How many moles of Au can be produced by the complete reaction of g of Pb? Take the molar mass of Pb as g/mol. ( moles) 3Pb(s) + 2Au(NO 3 ) 3 (aq) 3Pb(NO3) 2 (aq) + 2Au(s) d) Example: What mass of aluminum sulfate is required to produce g of solid aluminum? Take the molar mass of Al as g/mol and that of aluminum sulfate as g/mol. (109.3 g) Al 2 (SO 4 ) 3 (aq) + 3Zn(s) 2Al(s) + 3ZnSO 4 (aq) e) A mixture originally contains moles of H 3 PO 4 and moles of NaOH. How many moles of the excess reactant are left after the reaction goes to completion? (0.025 moles of H 3 PO 4 ) H 3 PO 4 (aq) + 3NaOH(aq) Na 3 PO 4 (aq) + 3H 2 O(l) f) Lithium reacts with fluorine according to the following reaction. If this reaction has a % yield, what mass of F 2 is required to produce grams of LiF? Take the molar mass of F 2 as g/mol and that of LiF as g/mol. (2.036 g F 2 ) 2Li + F 2 2LiF 5 of 7

6 g) SO 2 reacts with H 2 S as follows: 2H 2 S + SO 2 3S + 2H 2 O When 7.50 g of H 2 S reacts with g of SO 2, 7.34 g of sulfur are formed. What is the percent yield? Take the molar mass of H 2 S as g/mol, that of SO 2 as g/mol, and that of S as g/mol. (69.2%) 40. Know the concepts contained in the kinetic theory of gases. a) Example: There are 2 balloons, one filled with a certain number of atoms of Kr and the other with the same number of atoms of H 2 at the same temperature and pressure. What can we say about these two balloons? (They have the volume) 41. Know the ideal gas law (it will not be given to you, but R will) and how to use it. a) Example: What is the volume of a balloon filled with moles of CO 2 gas at 23 o C and a pressure of 1.12 atm? Take R as L atm/k mol. (2.67 L) 42. Know Graham s law of effusion. a) Example: Which gas will have the highest rate of effusion, H 2 or CO 2? (H 2 ) 43. Know what specific heat capacity is and what it means. a) Example: Consider the following specific heats of metals. Metal Specific Heat Zinc J/(g C) Magnesium 1.02 J/(g C) Lithium 3.58 J/(g C) Silver J/(g C) Bismuth J/(g C) If the same amount of heat is added to 25.0 g of each of the metals, which are all at the same initial temperature, which metal will have the highest temperature? (Bismuth) 44. Know what exothermic and endothermic mean. a) Example: In an endothermic reaction what is the sign of H o? Are the products at a higher or lower energy than the reactants? (Positive, higher) b) Example: In an exothermic reaction what is the sign of H o? Are the products at a higher or lower energy than the reactants? (Negative, lower) 45. Know the definition of H o f and the equation that describes it. a) Example: What is the chemical equation whose enthalpy change is the H o f of NO 2 (g)? (½ N 2 (g) + O 2 (g) NO 2 (g)) 46. Know how to use a thermochemical equation. a) Example: CH 4 (g) + 4Cl 2 (g) CCl 4 (g) + 4HCl(g), H = 434 kj Based on the above reaction, what energy change occurs when 12.5 grams of chlorine (Cl 2 ) reacts with excess CH 4? Take the molar mass of chlorine as g/mol. (19.1 kj of energy is released to the surrounding) 6 of 7

7 47. Know how to use Hess s Law. a) What is the H o for the following reaction given these heats of formation? ( kj/mol) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(g) H o f(ch 4 (g)) = 74.8 kj/mol H o f(o 2 (g)) = 0 kj/mol H o f(co 2 (g)) = kj/mol H o f(h 2 O(g)) = kj/mol b) Example: At 25 C, the following heats of reaction are known: H (kj/mol) 2ClF + O 2 Cl 2 O F 2 O 2ClF 3 + 2O 2 Cl 2 O F 2 O 2F 2 + O 2 2F 2 O 43.4 What is H o for the following reaction? ClF + F 2 ClF 3 ( kj/mol) 48. Know how to estimate H o for a reaction from bond dissociation energies. a) Example: Using the following bond energies estimate the heat of combustion for one mole of acetylene. ( 1228 kj) C 2 H 2 (g) + O 2 (g) 2 CO2(g) + H 2 O(g) Bond Bond Energy (kj/mol) C=C 839 C H 413 O=O 495 C=O 799 O H of 7

Chemistry 104 Final Exam Content Evaluation and Preparation for General Chemistry I Material

Chemistry 104 Final Exam Content Evaluation and Preparation for General Chemistry I Material What is 25 mph in mm s 1? Unit conversions What is 1025 K in o F? Which is larger 1 ft 3 or 0.1 m 3? What is