AP Chemistry
AP Chemistry Name _________________________
Practice Units 1-5 Period _____
Multiple Choice (No calculator)
Questions 1-3 refer to the following molecules.
(A) CO2 (B) CH4 (C) HF (D) PH3
1. Contains only two σ-bonds
| | |
2. Has the highest dipole moment
| | |
3. Has a molecular geometry that is trigonal pyramidal
| | |
Questions 4-6 refer to neutral atoms for which the atomic orbitals are represented below
(A) 1s(↑↓)
(B) 1s(↑↓) 2s(↑ ) 2p(↑ )( )( )
(C) 1s(↑↓) 2s(↑↓) 2p(↑ )(↑ )(↑ )
(D) [Ar] 4s(↑↓) 3d(↑↓)(↑↓)(↑↓)(↑ )(↑ )
4. contains the most valence electrons
| | |
5. is diamagnetic
| | |
6. is a transition metal
| | |
Questions 7-8 refer to a various points in time during an experiment conducted at 1.0 atm. Heat is added at a constant rate to a sample of a pure substance that is solid at time to. The graph below shows the temperature of the sample as a function of time.
(A) t1 (B) t2 (C) t3 (D) t5
7. Time when solid and liquid states are in equilibrium.
| | |
8. Time when the temperature of the substance is between its melting point and its boiling point
| | |
9. Three gases in the amounts shown in the table above are added to a previously evacuated rigid tank.
|Gas |Ar |CH4 |N2 |
|Amount |0.35 mol |0.90 mol |0.25 mol |
If the total pressure in the tank is 3.0 atm at 25oC, the partial pressure of CH4(g) in the tank is closest to
(A) 0.9 atm (B) 1.5 mol (C) 1.8 mol (D) 3.0 mol
| | |
10. Which of the following substances exhibits significant hydrogen bonding in the liquid state?
(A) CH2F2 (B) N2H4 (C) C2H4 (D) CH3OCH3
| | |
11. Which of the following best explains why the normal boiling point of CCI4(I) (350 K) is higher than the normal boiling point of CF4(I) (145 K)?
(A) The C-CI bonds in CCI4 are less polar than the C-F bonds in CF4.
(B) The C-CI bonds in CCI4 are weaker than the C-F bonds in CF4.
(C) The mass of the CCI4 molecule is greater than that of the CF4 molecule.
(D) The electron cloud of the CCI4 molecule is more polarizable than that of the CF4 molecule.
| | |
12. At which of the following temperatures and pressures would a real gas be most likely to deviate from ideal behavior?
Temperature (K) Pressure (atm)
(A) 100 50
(B) 200 5
(C) 300 0.01
(D) 500 0.01
| | |
13. After 195 days, a 10.0 g sample of pure 95Zr has decayed to the extent that only 2.50 g of the original 95Zr remains. The half-life of 95Zr is closest to
(A) 195 days (B) 97.5 days
(C) 65.0 days (D) 48.8 days
| | |
14. The first seven ionization energies of element X are shown in the table below.
| |1st |2nd |3rd |4th |5th |6th |7th |
|Ionization |787 |1,580 |3,200 |4,400 |16,000 |20,000 |24,000 |
|Energy | | | | | | | |
|(kJ•mol-1) | | | | | | | |
On the basis of these data, element X is most likely a member of which of the following groups of elements?
(A) Alkaline earth metals (B) Boron group
(C) Carbon group (D) Nitrogen group
| | |
15. Which of the following particles is emitted by an atom of 3920Ca when it decays to produce an atom of 3921Sc?
(A) 10n (B) 11H (C) β- (D) β+
| | |
16. At approximately what temperature will 40. g of argon gas at 1.0 atm occupy a volume of 11.2 L?
(A) 600 K (B) 550 K (C) 270 K (D) 140 K
| | |
17. Which of the following 0.25 M aqueous solutions has the highest boiling point at 1.0 atm?
(A) CaCl2 (B) Na2SO4 (C) NaCl (D) C6H12O6
| | |
18. An aqueous solution contains 11.0 % RbCl by mass. From the following list, what is needed to determine the molarity?
I. Mass of the sample
II. Volume of the sample
III. Temperature of the sample
(A) I only (B) II only (C) III only (D) I and II only
| | |
19. Heat energy is added slowly to a pure covalent network compound at its melting point. About half of the solid melts to become a liquid. Which of the following must be true about this process?
(A) Covalent bonds are broken as the solid melts.
(B) The temperature of the solid/liquid mixture remains the same while heat is being added.
(C) The volume of the compound increases as the solid melts to become a liquid.
(D) The average kinetic energy of the molecules become greater as the molecules leave the solid state and enter the liquid state.
| | |
20. Which of the following molecules contains bonds that have a bond order of 1.5?
(A) N2 (B) O3 (C) NH3 (D) CO2
| | |
21. In which of the following are the chemical species correctly ordered from smallest radius to largest radius?
(A) B < C < N (B) At < Xe < Kr
(C) CI < S < S2- (D) Na < Na+ < K
| | |
22. Of the following elements, which would be expected to have chemical properties most similar to those of sulfur, S?
(A) Br (B) CI (C) P (D) Se
| | |
23. Of the following gases, which has the greatest average molecular speed at 298 K?
(A) Cl2 (B) NO (C) H2S (D) HCN
| | |
24. Types of hybridization exhibited by carbon atoms in a molecule of propyne, CH3CHCH2, include which of the following?
I. sp II. sp2 III. sp3
(A) I only (B) II only (C) III only (D) I and II only
| | |
25. In which of the following processes are covalent bonds broken?
(A) Solid silver melts.
(B) Solid potassium chloride melts.
(C) Solid carbon (graphite) sublimes.
(D) Solid iodine sublimes.
| | |
Free Response (Calculator)
1. Use the information in the table below to respond to the statements and questions that follow. Your answers should be in terms of principles of molecular structure and intermolecular forces.
|Compound |Formula |Lewis Electron-Dot Diagram |
|Ethanethiol |CH3CH2SH |[pic] |
|Ethane |CH3CH3 |[pic] |
|Ethanol |CH3CH2OH |[pic] |
a. Draw the complete Lewis electron-dot diagram for ethyne, C2H2.
| |
b. Which of the four molecules contains the shortest carbon-to-carbon bond? Explain.
| |
c. A Lewis electron-dot diagram of a molecule of ethanoic acid is given below. The carbon atoms in the molecule are labeled x and y, respectively.
[pic]
Identify the geometry of the arrangement of atoms bonded to each of the following.
(1) Carbon x
| |
(2) Carbon y
| |
d. Energy is required to boil ethanol. Consider the statement “As ethanol boils, energy goes into breaking C−C bonds, C−H bonds, C−O bonds, and O−H bonds.” Is the statement true or false? Justify your answer.
| |
| |
e. Identify a compound from the table above that is nonpolar. Justify your answer.
| |
f. Ethanol is completely soluble in water, whereas ethanethiol has limited solubility in water. Account for the difference in solubilities between the two compounds in terms of intermolecular forces.
| |
| |
2. 2 Al(s) + 3 Zn2+ → 2 Al3+ + 3 Zn(s)
Respond to the following statements and questions that relate to the species and the reaction represented above.
a. Write the complete electron configuration
(e.g., 1s2 2s2 . . .) for Zn2+.
| |
b. Which species, Zn or Zn2+, has the greater ionization energy? Justify your answer.
| |
3. Answer the following using chemical concepts and principles of the behavior of gases.
a. A metal cylinder with a volume of 5.25 L contains 3.22 g of He(g) and 11.56 g of N2(g) at 15.0oC.
(1) Calculate the total pressure.
| |
| |
(2) Calculate the partial pressure of N2(g).
| |
b. A 1.50 L container holds a 9.62 g sample of an unknown gaseous saturated hydrocarbon at 30oC and 3.62 atm.
(1) Calculate the density of the gas.
| |
(2) Calculate the molar mass of the gas.
| |
(3) Write the formula of the hydrocarbon.
| |
(4) Calculate the root-mean-square speed of the gas molecules in the container at 30'C.
| |
4. Answer each of the following using principles of atomic or molecular structure, and/or intermolecular or intramolecular forces
a. Explain why the H-O-H bond angle in H2O is less than the H-N-H bond angles in NH3, as shown in the table below.
|H–O–H Bond Angle in H2O |H–N–H Bond Angle in NH3 |
|104.5o |107o |
| |
| |
b. Explain why the radius of the Br atom is less than the radius of the Br- ion, as shown in the table below.
|Radius of Br |Radius of Br- |
|0.111 nm |0.196 nm |
| |
| |
c. Explain why the dipole moment of HI is less than the dipole moment of HCI, as shown in the table below.
|Dipole Moment of HI |Dipole Moment of HCl |
|0.42 debye |1.08 debyes |
| |
| |
d. Explain why the normal boiling point of Ne is less than the normal boiling point of Kr, as shown in the table below.
|Normal Boiling Point of Ne |Normal Boiling Point of Kr |
|27 K |121 K |
| |
| |
5. Answer the following questions, which pertain to binary compounds.
a. In the box provided, draw a complete Lewis electron-dot diagram for the IF3 molecule.
| |
| |
b. On the basis of the Lewis electron-dot diagram that you drew in part a, predict the molecular geometry of the IF3 molecule.
| |
c. In the SO2 molecule, both of the bonds between sulfur and oxygen have the same length. Explain this observation, supporting your explanation by drawing in the box below a Lewis electron-dot diagram (or diagrams) for the SO2 molecule
| |
| |
| |
| |
| |
| |
d. On the basis of your Lewis electron-dot diagram(s) in part c, identify the hybridization of the sulfur atom in the SO2 molecule.
| |
6. Answer the following questions about the structures of ions that contain only sulfur and fluorine.
a. The compounds SF4 and BF3 react to form an ionic compound according to the following equation.
SF4 + BF3 → SF3BF4
(1) Draw a complete Lewis structure for the SF3+ cation in SF3BF4.
| |
| |
| |
(2) Identify the type of hybridization exhibited by sulfur in the SF3+ cation.
| |
(3) Identify the geometry of the SF3+ cation that is consistent with the Lewis structure drawn in part a(1).
| |
(4) Predict whether the F-S-F bond angle in the SF3+ cation is larger than, equal to, or smaller than 109.5°. Justify your answer.
| |
| |
| |
b. The compounds SF4 and CsF react to form an ionic compound according to the following equation.
SF4 + CsF → CsSF5
(1) Draw a complete Lewis structure for the SF5- anion in CsSF5.
| |
| |
(2) Identify the type of hybridization exhibited by sulfur in the SF5- anion.
| |
(3) Identify the geometry of the SF5- anion that is consistent with the Lewis structure drawn in part b(1).
| |
(4) Identify the oxidation number of sulfur in the compound CsSF5.
| |
AP Chemistry Name _________________________
Practice Units 6-9 Period _____
Multiple Choice (No calculator)
1. What mass of KBr (MM = 119 g mol-1) is required to make 500. mL of a 0.200 M KBr solution?
(A) 0.595 g (B) 1.19 g
(C) 2.50 g (D) 11.9 g
| | |
2. The standard enthalpy of formation, ΔHof, of HI(g) is +26 kJ mol-1. What is the approximate mass of HI(g) that is decompose into H2(g) and I2(s) to release 500 kJ of energy?
(A) 250 g (B) 650 g
(C) 1,300 g (D) 2,500 g
| | |
3. _CH3OCH3(g) + _O2(g) → _CO2(g) + H2O(g)
When the equation above is balanced using the lowest whole-number coefficients, the coefficient for CO2(g) is
(A) 6 (B) 4 (C) 3 (D) 2
| | |
4. What mass of Cu(s) would be produced if 0.50 mol of Cu2O(s) was reduced completely with excess H2(g)?
(A) 13 g (B) 25 g (C) 32 g (D) 64 g
| | |
5. Which of the following is a formula of a ketone?
(A) CH3–CO–CH3 (B) CH3–CH2–COOH
(C) CH3–CH2–CH2OH (D) CH3–CH2–O–CH3
| | |
6. Which of the following would produce the MOST mass of CO2 if completely burned in excess oxygen gas?
(A) 10.0 g CH4 (B) 10.0 g CH3OH
(C) 10.0 g C2H4 (D) 10.0 g C2H6
| | |
Questions 7-9 refer to a galvanic cell constructed using two half-cells and based on the two half-reactions below.
Zn2+ + 2 e- → Zn(s) Eo = -0.76 V
Fe3+ + e- → Fe2+ Eo = 0.77 V
7. As the cell operates, ionic species that are found in the half-cell containing the cathode include which of the following?
I. Zn2+ II. Fe2+ III. Fe3+
(A) I only (B) II only (C) III only (D) II and III
| | |
8. What is the standard cell potential for the galvanic cell?
(A) -0.01 V (B) 0.01 V
(C) 1.53 V (D) 2.30 V
| | |
9. What is the oxidizing agent for the forward reaction?
(A) Zn2+ (B) Zn (C) Fe2+ (D) Fe3+
| | |
10. C(diamond) → C(graphite)
For the reaction represented above, the standard Gibbs free energy change, ΔGo298, is negative. However the reaction does not occur at 298 K and 1.00 atm. This is probably because the following value is extremely large.
(A) ΔS (B) ΔH (C) ΔG (D) Ea
| | |
11. A reaction is spontaneous below 400 K but is not spontaneous above 400 K. If ΔSo for the reaction is –50.0 J•mol-1 K-1, then the value of ΔHo for the reaction is
(A) -50.0 kJ mol-1 (B) -20 kJ mol-1
(C) -0.050 kJ mol-1 (D) 2.0 x 104 kJ mol-1
| | |
12. _ Au3+ + _I- → _Au(s) + _I2(s)
When the equation is balanced using the lowest whole-number coefficients, the coefficient for I2(s) is
(A) 8 (B) 6 (C) 4 (D) 3
| | |
13. Z → X + Y
A pure substance Z decomposes into two products, X and Y, as shown by the equation. Which of the following graphs of the concentration of Z versus time is consistent with the rate of the reaction being second order with respect to Z?
(A) (B)
(C) (D)
| | |
14. CS2(l) + 3 O2(g) → CO2(g) + 2 SO2(g)
When 0.60 mol of CS2(l) reacts as completely as possible with 1.5 mol of O2(g) according to the equation above, the total number of moles of SO2(g) is
(A) 0.60 mol (B) 1.0 mol (C) 1.2 mol (D) 1.5 mol
| | |
Questions 15-16 refer to an experiment to determine the heat of fusion of ice. A student used a calorimeter consisting of a polystyrene cup and a thermometer. The cup was weighed, then filled halfway with warm water, then weighed again. The temperature of the water was measured, and some ice cubes from a 0oC ice bath were added to the cup. The mixture was gently stirred as the ice melted, and the lowest temperature reached by the water in the cup was recorded. The cup and its contents were weighed again.
15. The purpose of weighing the cup and its contents again at the end of the experiment was to
(A) determine the mass of ice that was added
(B) determine the mass of the thermometer
(C) determine the mass of water that evaporated
(D) verify the mass of water that was cooled
| | |
16. A student weighs out 0.0154 mol of pure, dry NaCI in order to prepare a 0.154 M NaCI solution. Of the following pieces of laboratory equipment, which would be most essential for preparing the solution?
(A) 50 mL volumetric pipet
(B) 100 mL Erlenmeyer flask
(C) 100 mL graduated beaker
(D) 100 mL volumetric flask
| | |
17. Suppose that during the experiment, a significant amount of water from the ice bath adhered to the ice cubes. How does this affect the calculated heat of fusion of ice?
(A) The calculated value is too large because less warm water had to be cooled
(B) The calculated value is too large because more cold water had to be heated
(C) The calculated value is too small because less ice was added than the student assumed
(D) The calculated value is too small because the total mass of the calorimeter contents was too large.
| | |
18. A steady electric current is passed through molten MgCl2 for exactly 1.00 hour, producing 243 g of Mg metal. If the same current is passed through AlCl3 for 1.00 hour, the mass of Al metal produced is closest to
(A) 27.0 g (B) 54.0 g (C) 120. g (D) 180. g
| | |
19. What is the empirical formula of a hydrocarbon that is 20 % hydrogen by mass?
(A) CH3 (B) C2H5 (C) C3H4 (D) C4H9
| | |
20. Pb(s) Δ Pb(l)
Which of the following is true for the process represented above at 327oC and 1 atm; the normal melting point for Pb(s)
(A) ΔH = 0 (B) ΔS = 0 (C) ΔG = 0 (D) ΔH/ΔS = 0
| | |
21. Equal volumes of 0.15 M HCl and 0.25 M HCl are mixed. The molarity of the resulting solution is
(A) 0.25 M HCl (B) 0.20 M HCl
(C) 0.15 M HCl (D) 0.10 M HCl
| | |
22. 2 HClO + 3 O2 → 2 HClO4
As the reaction represented above proceeds to the right, the oxidation number of chlorine changes from
(A) -1 to +5 (B) +1 to +5
(C) +1 to +7 (D) +3 to +7
| | |
23. A large piece of wood burns slowly, but sawdust combusts explosively. The primary reason for the difference is that compared with a large piece of wood sawdust
(A) has a greater surface area per kilogram
(B) has a greater carbon content per kilogram
(C) absorbs more atmospheric moisture per kilogram
(D) contains more compounds that act as catalysts for combustion
| | |
24. When a solution is formed by adding some methanol, CH3OH, to water, processes that are exothermic include which of the following?
I. Methanol molecules move water molecules apart as the methanol goes into solution.
II. Water molecules move methanol molecules apart as the methanol goes into solution.
III. Intermolecular attractions form between molecules of water and methanol as the methanol goes into solution.
(A) I only (B) II only (C) III only (D) I and II only
| | |
25. 8 H2(g) + S8(s) → 8 H2S(g)
When 0.1 mol of S8(s) reacts completely with excess H2(g) according to the equation above, the volume of H2S(g), measured at 0oC and 1.00 atm, produced is closest to
(A) 30 L (B) 20 L (C) 10 L (D) 5 L
| | |
Free Response (Calculator)
1.
5 Fe2+ + MnO4- + 8 H+ → 5 Fe3+ + Mn2+ + 4 H2O
A galvanic cell and the balanced equation for the spontaneous cell reaction are shown above. The two reduction half-reactions for the overall reaction that occurs in the cell are shown below.
Half-Reaction E° (V) at 298 K
Fe3+ + e- → Fe2+ + 0.77
MnO4- + 8 H+ + 5 e- → Mn2+ + 4 H2O +1.49
a. On the diagram, clearly label the cathode.
b. Calculate the value of the standard potential, E°, for the spontaneous cell reaction.
| |
c. How many moles of electrons are transferred when 1.0 mol of MnO4- is consumed in the overall reaction?
| |
d. Calculate the value of the equilibrium constant Keq for the cell reaction at 25°C. Explain what the magnitude of Keq tells you about the extent of the reaction.
| |
| |
| |
Three solutions, one containing Fe2+, one containing MnO4-, and one containing H+, are mixed in a beaker and allowed to react. The initial concentrations of the species in the mixture are 0.60 M Fe2+, 0.10 M MnO4-, and 1.0 M H+.
e. Which has the higher concentration, Mn2+ or MnO4-, when the reaction mixture has come to equilibrium? Explain.
| |
| |
f. When the reaction mixture has come to equilibrium, what are the molar concentrations of Fe2+ and Fe3+?
| |
2. A sample of ore containing the mineral tellurite (TeO2) was dissolved in acid. The resulting solution was then reacted with a solution of K2Cr2O7 to form telluric acid (H2TeO4). The unbalanced chemical equation for the reaction.
_TeO2(s) + _Cr2O72- + _H+ → _H2TeO4(aq) + _Cr3+ + _H2O
a. Identify the molecule or ion that is being oxidized.
| |
b. Give the oxidation number of Cr in the Cr2O72- ion.
| |
c. Balance the chemical equation given above.
| |
| |
| |
In the procedure described above, 46.00 mL of 0.03109 M K2Cr2O7 was added to the ore sample after it was dissolved in acid. When the chemical reaction had progressed as completely as possible, the amount of unreacted (excess) Cr2O72- was determined by titrating the solution with 0.110 M Fe(NO3)2. The reaction that occurred during the titration is represented by the balanced equation:
6 Fe2+ + Cr2O72- + 14 H+ → 2 Cr3+ + 6 Fe3+ + 7 H2O
A volume of 9.85 mL of 0.110 M Fe(NO3)2 was required to reach the equivalence point.
d. Calculate the moles of excess Cr2O72- titrated.
| |
e. Calculate the moles of Cr2O72- that reacted.
| |
f. Calculate the mass of tellurite that was in the ore.
| |
3. The mass of an aqueous solution of H2O2 is 6.951 g. The H2O2 in the solution decomposes completely according to the reaction represented below. The O2(g) produced is collected in an inverted graduated tube over water at 23.4oC and has a volume of 182.4 mL when the water levels inside and outside of the tube are the same. The atmospheric pressure in the lab is 762.6 torr, and the equilibrium vapor pressure of water at 23.4°C is 21.6 torr.
2 H2O2(aq) → 2 H2O(l) + O2(g)
a. Calculate the partial pressure of O2(g) in the tube.
| |
b. Calculate moles of O2(g) produced in the reaction.
| |
c. Calculate the mass of H2O2 that decomposed.
| |
d. Calculate the percent of H2O2 in the 6.951 g of solution.
| |
e. Write the oxidation number of the oxygen atoms in H2O2 and the oxidation number of the oxygen atoms in O2 in the appropriate cells in the table below.
|Substance |H2O2 |O2 |
|Oxidation Number of Oxygen Atoms | | |
f. Write the balanced oxidation half-reaction.
| |
4. Methane gas reacts with chlorine gas to form dichloro-methane and hydrogen chloride.
CH4(g) + 2 Cl2(g) → CH2Cl2(g) + 2 HCl(g)
a. A 25.0 g sample of methane gas is placed in a reaction vessel containing 2.58 mol of Cl2(g).
(1) Identify the limiting reactant. Show calculations.
| |
| |
| |
(2) Calculate the moles of CH2Cl2(g) after the limiting reactant has been totally consumed.
| |
5. A(g) + B(g) → C(g) + D(g)
For the gas-phase reaction represented above, the following experimental data were obtained.
|Exp. |Initial [A] |Initial [B] |Initial Reaction Rate |
| |(mol•L−1) |(mol•L−1) |(mol•L−1•s-1) |
|1 |0.033 |0.034 |6.67 x 10-4 |
|2 |0.034 |0.137 |1.08 x 10-2 |
|3 |0.136 |0.136 |1.07 x 10-2 |
|4 |0.202 |0.233 |? |
a. Determine the order of the reaction with respect to reactant A. Justify your answer.
| |
b. Determine the order of the reaction with respect to reactant B. Justify your answer.
| |
c. Write the rate law for the overall reaction.
| |
d. Determine the value of the rate constant, k, for the reaction. Include units with your answer.
| |
e. Calculate the initial reaction rate for experiment 4.
| |
f. The following mechanism has been proposed for the reaction. Provide two reasons why the mechanism is acceptable.
Step 1: B + B → E + D slow
Step 2: E + A Δ B + C fast equilibrium
| |
g. In the mechanism in part (f), is species E a catalyst, or is it an intermediate? Justify your answer.
| |
6. The reaction represented above is one that contributes significantly to the formation of photochemical smog.
a. Calculate the quantity of heat released when 73.1 g of NO(g) is converted to NO2(g).
| |
b. The standard free-energy ΔGo is -70.4 kJ at 25oC.
(1) Calculate the equilibrium constant Keq at 25°C.
| |
(2) Indicate whether ΔGo would become more negative, less negative, or remain unchanged as the temperature is increased. Justify your answer.
| |
c. Calculate the standard molar entropy So for O2(g) at 25°C.
|Substance |NO(g) |NO2 (g) |
|So (J•K-1•mol-1) |210.8 |240.1 |
| |
d. Calculate the bond energy, in kJ•mol-1, of the nitrogen-oxygen bond in NO2.
| |Bond Energy (kJ•mol-1) |
|Nitrogen-oxygen bond in NO |607 |
|Oxygen-oxygen bond in O2 |495 |
|Nitrogen-oxygen bond in NO2 |? |
| |
AP Chemistry Name _________________________
Practice Units 10-12 Period _____
Multiple Choice (No calculator)
Questions 1-3 refer to the chemical reactions represented below.
(A) HC2H3O2(aq) + NH3(aq) → C2H3O2- + NH4+
(B) Ba2+ + SO42- → BaSO4(s)
(C) Zn(OH)2(s) + 2 OH- → [Zn(OH)4]2-
(D) 2 K(s) + Br2(l) → 2 KBr(s)
1. A precipitation reaction
| | |
2. A reaction in which a coordination complex is formed
| | |
3. A Lewis acid-base reaction that is not a Brønsted-Lowry acid-base reaction
| | |
4. The acid-dissociation constant, Ka, for a weak monoprotic acid HA is 2.5 x 10-6. The pH of 0.40 M HA is closest to
(A) 2.0 (B) 3.0 (C) 4.0 (D) 6.0
| | |
5. Which of the systems in equilibrium represented below will exhibit a shift to the left (toward reactants) when the pressure on the system is increased by reducing the volume of the system? (Assume that the temperature is constant.)
(A) 2 Mg(s) + O2(g) Δ 2 MgO(s)
(B) SF4(g) + F2(g) Δ SF6(g)
(C) H2(g) + Br2(g) Δ 2 HBr(g)
(D) SO2Cl2(g) Δ SO2(g) + Cl2(g)
| | |
6. In an aqueous solution with a pH of 11.50 at 25oC, the molar concentration of OH-(aq) is approximately
(A) 3.2 x 10-12 M (B) 3.2 x 10-3 M
(C) 2.5 x 10-1 M (D) 2.5 M
| | |
7. Which of the following changes to a reaction system in equilibrium would affect the value of the equilibrium constant, Keq, for the reaction? (Assume in each case that all other conditions are held constant.)
(A) Adding more of the reactants to the system
(B) Adding a catalyst for the reaction to the system
(C) Increasing the temperature of the system
(D) Increasing the pressure on the system
| | |
8. N2(g) + O2(g) + Cl2(g) Δ 2 NOCl(g) ΔHo = +104 kJ•mol-1
The equilibrium system represented above is contained in a sealed, rigid vessel. Which of the following will increase if the temperature of the mixture is raised?
(A) [N2]
(B) The rate of the forward reaction only
(C) The rate of the reverse reaction only
(D) The rate of both the forward and reverse reactions
| | |
9. F- + H2O → HF(aq) + OH-
Which of the following species, if any, acts as a Brønsted-Lowry base in the reversible reaction represented above?
(A) HF(aq) (B) H2O (C) F- only (D) F- and OH-
| | |
10. At 25oC a saturated solution of a metal hydroxide, M(OH)2, has a pH of 9.0. What is the value of the solubility-product constant, Ksp, of M(OH)2(s) at 25oC?
(A) 1.0 x 10-27 (B) 5.0 x 10-19
(C) 5.0 x 10-16 (D) 1.0 x 10-15
| | |
11. XY2(aq) Δ X2+ + 2 Y-
A soluble compound XY2 dissociates in water according to the equation above. In a 0.050 m solution of the compound, the XY2(aq) species is 40.0 % dissociated. In the solution, the number of moles of particles of solute per 1.0 kg of water is closest to
(A) 0.15 (B) 0.090 (C) 1.040 (D) 0.020
| | |
12. Ascorbic acid, H2C6H6O6(s), is a diprotic acid with
K1 = 7.9 x 10-5 and K2 = 1.6 X 10-12. In a 0.005 M aqueous solution of ascorbic acid, which of the following species is present in the lowest concentration?
(A) H3O+ (B) H2C6H6O6
(C) HC6H6O6- (D) C6H6O62-
| | |
Questions 13-14 The graph below shows the titration curve that results when 100. mL of 0.0250 M acetic acid is titrated with 0.100 M NaOH.
[pic]
13. Which indicator is the best choice for this titration?
Indicator pH Range of Color Change
(A) Methyl orange 3.2 - 4.4
(B) Methyl red 4.8 - 6.0
(C) Bromothymol blue 6.1 - 7.6
(D) Phenolphthalein 8.2 - 10.0
| | |
14. What part of the curve corresponds to the optimum buffer action for the acetic acid/acetate ion pair?
(A) Point V (B) Point X
(C) Point Z (D) Along all of section YZ
| | |
15. HCO3- + OH- Δ H2O(l) + CO32- ΔH = -41.4 kJ
When the reaction represented by the equation above is at equilibrium at 1 atm and 25°C, the ratio [CO32-]/[ HCO3-] can be increased by doing which of the following?
(A) Decreasing the temperature
(B) Adding acid
(C) Adding a catalyst
(D) Diluting the solution with distilled water
| | |
Free Response (Calculator)
1. Answer the following questions about the solubility of the salts Li3PO4 and PbCl2. Assume that hydrolysis effects are negligible. The equation for the dissolution of Li3PO4(s) is
Li3PO4(s) Δ 3 Li+ + PO43- Ksp = 3.2 x 10-9
a. Write the equilibrium-constant expression for the dissolution of Li3PO4(s).
| |
b. Assuming that volume changes are negligible, calculate the maximum number of moles of Li3PO4(s) that can dissolve in
(1) 0.50 L of water
| |
| |
(2) 0.50 L of 0.20 M LiNO3
| |
| |
The equation for the dissolution of PbCl2, is
PbCl2(s) Δ Pb2+ + 2 Cl- Ksp = 1.6 x 10-5
c. Calculate the concentration of Cl- in a saturated solution of PbCl2.
| |
| |
d. An open container holds 1.000 L of 0.00400 M PbCl2, which is unsaturated. Calculate the minimum volume of water, in mL, that must evaporate from the container before solid PbCl2 can precipitate.
| |
| |
2. Butane, C4H10, occurs in two isomeric forms, n-butane and isobutane. Both exist as gases at 25°C and 1.0 atm. The two isomers exist in equilibrium by the equation:
n-butane(g) Δ isobutane(g) Kc = 2.5 at 25°C
A 0.010 mol sample of pure n-butane is placed in an evacuated 1.0 L rigid container at 25°C.
a. Write the expression for the equilibrium constant Kc.
| |
b. Calculate the initial pressure in the container when the n-butane is first introduced (before the reaction starts).
| |
c. The n-butane reacts until equilibrium has been established at 25°C.
(1) What is the total pressure at equilibrium? Justify your answer.
| |
(2) Calculate the molar concentration of each species at equilibrium.
| |
| |
| |
(3) The volume of the system is reduced to 0.50 L, what is the new concentration of n-butane after equilibrium has been reestablished at 25°C? Justify your answer.
| |
| |
| |
In another experiment 0.010 mol of pure isobutane is placed in an evacuated 1.0 L rigid container and allowed to come to equilibrium at 25°C.
d. What are the molar concentration of each species after equilibrium has been established? Justify your answer.
| |
3. A pure 14.85 g sample of the weak base ethylamine, C2H5NH2, is dissolved in enough distilled water to make 500. mL of solution.
a. Calculate the molar concentration of the C2H5NH2 in the solution.
| |
The aqueous ethylamine reacts with water according to the equation: C2H5NH2(aq) + H2O Δ C2H5NH3+ + OH-
b. Write the equilibrium-constant expression for the reaction between C2H5NH2(aq) and water.
| |
c. Which is present in the solution at the higher concentration at equilibrium, C2H5NH2(aq) or C2H5NH3+? Justify your answer.
| |
d. A different solution is made by mixing 500. mL of 0.500 M C2H5NH2 with 500. mL of 0.200 M HCl. The pH of the resulting solution is found to be 10.93.
(1) Calculate the concentration of OH- in the solution.
| |
(2) Write the net-ionic equation that represents the reaction that occurs when the C2H5NH2 solution is mixed with the HCl solution.
| |
(3) Calculate the molar concentration of the C2H5NH3+ that is formed in the reaction.
| |
| |
(4) Calculate the value of Kb for C2H5NH2.
| |
4. Answer the questions that relate to halogen oxyacids.
a. Use the information in the table to answer part a(1).
|Acid |Ka at 298 K |
|HOCl |2.9 x 10-8 |
|HOBr |2.4 x 10-9 |
(1) Which of the two acids is stronger, HOCl or HOBr? Justify your answer in terms of Ka.
| |
(2) Draw a complete Lewis electron-dot diagram for the stronger acid.
| |
(3) Hypoiodous acid has the formula HOI. Predict whether HOI is a stronger acid or a weaker acid than the acid that you identified in part a(1). Justify your answer in terms of chemical bonding.
| |
| |
| |
| |
b. Write the equation for the reaction that occurs between hypochlorous acid and water.
| |
c. A 1.2 M NaOCl solution is prepared by dissolving solid NaOCl in distilled water at 298 K according to the reaction: OCl– + H2O Δ HOCl(aq) + OH−
(1) Write the equilibrium-constant expression for the hydrolysis reaction between OCl- and H2O.
| |
(2) Calculate the value of the equilibrium constant at 298 K for the hydrolysis reaction.
| |
(3) Calculate the value of [OH-] in the 1.2 M NaOCl solution at 298 K.
| |
| |
d. A buffer solution is prepared by dissolving some solid NaOCl in a solution of HOCl at 298 K. The pH of the buffer solution is determined to be 6.48.
(1) Calculate the [H3O+] in the buffer solution.
| |
(2) Indicate which of HOCl(aq) or OCl− is present at the higher concentration in the buffer solution. Support your answer with a calculation.
| |
| |
5. Hydrofluoric acid, HF(aq), dissociates in water as represented by the equation:
HF(aq) Δ H+ + F- Ka = 7.2 x 10-4
a. Write the equilibrium-constant expression for the dissociation of HF(aq) in water.
| |
b. Calculate the molar concentration of H3O+ in a 0.40 M HF(aq) solution.
| |
| |
HF(aq) reacts with NaOH(aq) according to the reaction.
HF(aq) + OH- → H2O + F-
A volume of 15 mL of 0.40 M NaOH(aq) is added to 25 mL of 0.40 M HF(aq) solution.
c. Calculate the moles of HF(aq) remaining in the solution.
| |
| |
d. Calculate the molar concentration of F- in the solution.
| |
e. Calculate the pH of the solution
| |
| |
| |
6. Answer the questions that relate to salts of lead and barium.
a. A saturated solution is prepared by adding excess PbI2(s) to distilled water to form 1.0 L of solution at 25°C. The concentration of Pb2+ in the saturated solution is found to be 1.3 x 10-3 M. The equation is:
PbI2(s) Δ Pb2+ + 2 I-
(1) Write the equilibrium expression for the equation.
| |
(2) Calculate the molar concentration of I-.
| |
(3) Calculate the equilibrium constant Ksp.
| |
b. A saturated solution is prepared by adding PbI2(s) to distilled water to form 2.0 L of solution at 25°C. What are the molar concentrations of Pb2+ and I-? Justify your answer.
| |
c. Solid NaI is added to a saturated solution of PbI2 at 25°C. Assuming that the volume of the solution does not change, does the molar concentration of Pb2+ in the solution increase, decrease, or remain the same? Justify your answer.
| |
| |
d. The value of Ksp for the salt BaCrO4 is 1.2 x 10-10. When a 500. mL sample of 8.2 x 10-6 M Ba(NO3)2 is added to 500. mL of 8.2 x 10-6 M Na2CrO4, no precipitate is observed.
(1) Assuming that volumes are additive, calculate the molar concentrations of Ba2+ and CrO42- in the 1.00 L of solution.
| |
| |
| |
| |
(2) Use the molar concentrations of Ba2+ ions and CrO42- ions as determined above to show why a precipitate does not form. You must include a calculation as part of your answer.
| |
................
................
In order to avoid copyright disputes, this page is only a partial summary.
To fulfill the demand for quickly locating and searching documents.
It is intelligent file search solution for home and business.