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Chemistry 40S Supplementary Problems Electronic Structure, etc… page 1

Electronic Structure, Bonding & The Periodic Table 1. What are the definitions of “wavelength” and “frequency” of waves? 2. List the regions of the electromagnetic spectrum in order of increasing wavelength. 3. List the colours of the rainbow in order of increasing frequency. 4. What is the “speed of light”? What symbol is used to represent this quantity? 5. A recent news report stated that scientists had succeeded in slowing down the speed

of light. How is this possible if the speed of light is supposed to be “constant”? 6. CJOB radio emits a signal at a frequency of 680 kHz. Calculate the wavelength of

these radio waves. 7. A cordless telephone operates at 900 MHz. What is the wavelength of this signal? 8. Visible light falls between 400 nm and 700 nm. What is the range of frequencies for

visible light? 9. A sodium lamp emits yellow light at a frequency of 5.1 x 1014 s-1. Calculate the

wavelength of this light. 10. Which color of light is at the 400 nm end of the visible spectrum? 11. What is a “quantum”? What is a “photon”? 12. Calculate the minimum and maximum energies of light in the visible spectrum (see

question 6). 13. How is the energy of electromagnetic radiation related to its wavelength? How is it

related to the frequency of the radiation? 14. Calculate the energy and frequency of light where λ = 6.56 x 10-5 cm. 15. Microwave radiation may have ν = 3.2 x 1011 s-1. Calculate the energy and

wavelength of this radiation. 16. Green light has an energy of 3.84 x 10-19 J/photon. Calculate λ and ν for this green

light. 17. Calculate the energy of an electron in the 3rd energy level of a hydrogen atom. 18. Calculate the energy of an electron in the 2nd energy level of a lithium cation, Li2+. 19. An electron falls from n = 5 to n = 2 in a hydrogen atom.

a) Is energy released or absorbed? Explain.

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b) Calculate the wavelength of light emitted during this electron transition.


20. An electron falls from n = 6 to n = 3 in a helium cation, He+. Calculate the wavelength of the light emitted during this transition. Would this light be visible to the naked eye?

21. Calculate the frequency of light that would allow an electron to jump from the 2nd to

the 3rd energy level in a hydrogen atom. 22. What does it mean to say that an atom is in the “ground state”? What is the term

applied when the atom is not in the ground state? 23. When an atom is “ionized”, it loses an electron. The electron can be thought to jump

from to an energy level where n = ∞. Calculate the maximum wavelength of radiation that would ionize a ground-state hydrogen atom.

24. What does the “Heisenberg Uncertainty Principle” state? 25. How did Bohr’s model of the atom violate the Heisenberg Uncertainty Principle? 26. Is an electron a particle? or is it a wave? Explain. 27. What is an “orbital”? What is the probability of finding an electron in an orbital as it

is pictured in your text? 28. What are the four “quantum numbers” that are used to describe an electron within

an atom? What does each quantum number determine? 29. What does the Pauli Exclusion Principle state? 30. How many electrons can be held in any given orbital? This is a direct result of the

Pauli Exclusion Principle. 31. How many electrons (in total) can be held on the 3rd energy level of an atom? Show

how you calculated the answer. 32. Why is the electron configuration, 1s2 2s2 2p4 3s1, called an “excited state”? 33. Why is the electron configuration, 1s2 2s2 2p7 3s1, an impossible configuration? 34. Draw orbital diagrams for: a) silicon b) fluorine c) copper 35. Write electron configurations for: a) nitrogen b) argon c) chromium 36. Write abbreviated electron configurations for: a) calcium b) chlorine 37. Identify the element whose valence electron configuration is “5s2 5p2”. 38. Write the valence electron configurations for: a) bromine b) barium 39. Identify the element whose electron configuration is 1s2 2s2 2p6 3s2 3p6 4s2 3d2.


40. Identify the element whose abbreviated electron configuration is [Ar] 4s2 3d10 4p3.


41. Identify the element whose valence configuration is 5s2 5p5. 42. How many unpaired electrons are there in ground state atoms of…

a) carbon b) sulfur c) iron d) nickel 43. Which element(s) in period 4 have two unpaired electrons in their ground states? 44. Which two families of elements comprise the “s-block” on the periodic table? 45. Which family of elements comprises the “d-block” on the periodic table? 46. Which families of elements comprise the “f-block” on the periodic table? 47. Give the symbol of the atom whose ground state electron configuration is…

a) 1s2 2s2 2p3 b) 1s2 2s2 2p6 3s1 c) 1s2 2s2 2p6 3s2 3p6 4s2 d) [Xe] 6s2 4f14 5d10 6p2

48. Which family of elements has electron configurations ending in “p5”? 49. Describe the trend in atomic radii of elements within a period on the periodic table.

Explain why the trend exists. 50. Describe the trend in atomic radii of elements within a group on the periodic table.

Explain why the trend exists. 51. Lithium (Z = 3) has the following successive ionization energies:

IE 1: 520 kJ/mol IE 2: 7 300 kJ/mol IE 3: 11 810 kJ/mol Explain why there is such a large increase between IE 1 and IE 2. 52. An element in period 3 has the following successive ionization energies:

IE 1: 787 kJ/mol IE 2: 1577 IE 3: 3232 IE 4: 4356 IE 5: 21 270 How many valence electrons does the element have? Identify the element.

53. What is the trend in first ionization energies within a period on the periodic table?

Explain the trend. 54. Explain the following trend in first ionization energies:

Li: 520 kJ/mol Na: 496 K: 419 Rb: 403 Cs: 376

55. Arrange the following in order of increasing first I.E.: As, Sn, Br, Sr


56. Arrange the following in order of increasing atomic radius: Rb, As, Sb, Br, Sr


57. An element is considered to be paramagnetic if it has unpaired electrons. Which elements in the third period on the periodic table are paramagnetic?

58. Write the symbols of the four most electronegative elements, in order of increasing

electronegativity. 59. What is a “polar covalent bond”? Discuss “electronegativity” and “sharing of

electrons” in your answer. 60. Classify the following bonds as “non-polar covalent”, “polar covalent”, or “ionic”.

Include calculations as part of your answers. a) Li-F b) C-Br c) Zn-Cl d) Cu-S

61. Classify the following bonds as “non-polar covalent”, “polar covalent”, or “ionic”.

Include calculations as part of your answers. a) Na-O b) Al-I c) Si-O d) N-N

62. Arrange the following in expected order of increasing radius: K+, Cl-, S2-, Ca2+, Ar.

Explain your answer. 63. Arrange the following in expected order of increasing radius: F-, O2-, Na+, Ne. 64. Draw the Lewis electron dot diagrams for: a) Mg b) O c) Br 65. Draw the Lewis electron dot diagrams for: a) Na b) P c) Ar 66. Draw the Lewis electron dot diagrams for: a) Ca2+ b) O2- c) F- 67. Draw the Lewis electron dot diagrams for: a) NH4

+ b) H3O+ c) OH- 68. Draw the Lewis electron dot diagrams for: a) C2H6 b) C2H4 c) CO2 69. Draw the Lewis electron dot diagrams for: a) N2 b) O2 c) Br2 70. Draw the Lewis electron dot diagrams for: a) NaCl b) K2O c) NH4Br 71. Draw the Lewis electron dot diagrams for: a) NH3 b) CN- c) NaBr 72. What is a coordinate covalent bond? Draw the Lewis dot diagram for the

hydronium ion and label a coordinate covalent bond. 73. Use the periodic table and concepts in this unit to answer the following:

a) What is the electron dot formula for the element with Z = 51. b) Which of the following elements has the smallest radius? (Li, C, O, Br) c) What type of bonding would you expect between carbon and sulfur? d) Which element has the lowest first ionization energy – Ca, Sr, or Mg? e) Which of the following elements is most metallic – Si, S, N, or Cl?


f) Which is larger – P3- or P, Ca2+ or Ca?

AP Chemistry 42S Supplementary Problems Chemical Kinetics… page 5

Chemical Kinetics – Rates of Chemical Reactions 1. Describe the collision theory of reactions. What is a “successful” collision? What

factors determine if a collision is successful? 2. List the factors that affect the rates of chemical reactions. 3. Which of the following units would NOT be used to measure the rate of a reaction?

Explain why. a) cm/s b) mL/min c) mol/L 4. Refrigerated food stays fresh for long periods. The same food stored at room

temperature spoils quickly. Why? 5. Use collision theory to explain how “temperature” affects the rate of a reaction. 6. What is the “activation energy” of a reaction? How is this related to the rate of the

reaction? 7. When the gas to a Bunsen burner is turned on, the gas does not begin to burn (react

with the oxygen in the air) until a flame lights it. Once lit, however, it burns completely until the gas is turned off. Explain these observations.

8. A charcoal grill has air vents that can be adjusted. As the air vents are made

smaller, explain what happens to the rate of burning of the charcoal. Will the temperature become higher or lower?

9. In your body, sugar is oxidized at 37°C (body temperature). Outside the body, sugar

will burn only at a temperature of over 600°C. What accounts for this difference? 10. What is a “catalyst”? How does the catalyst affect the rate of a reaction? 11. What is an “inhibitor”? How does the inhibitor affect the rate of a reaction? 12. Explain the difference between a homogeneous and a heterogeneous catalyst. 13. What is the “mechanism” of a reaction? 14. What is an “elementary step”? What is the “rate determining step” of a mechanism? 15. Use collision theory to explain why most reactions probably occur through a multi-

step reaction mechanism. 16. Define the term “reaction intermediate”. Why are intermediates important for

studying reaction mechanisms?


17. For each reaction below, explain how the rate of disappearance of each reactant is related to the rate of appearance of each product: a) H2O2 → H2O + ½O2 b) CO + 2 H2 → CH3OH


18. What is an “endothermic” reaction? What is the sign of ∆H for this reaction? 19. Construct and label a potential energy diagram from the information below.

Calculate ∆H for the reaction and include it on your diagram.

P.E. of reactants = 300 kJ P.E. of products = 500 kJ Ea = 400 kJ 20. Construct and label a potential energy diagram for the two-step reaction, using the

information below. Calculate ∆H for step 1 and ∆H for the overall reaction.

P.E. of reactants = 400 kJ ∆H for step 2 = -100 kJ P.E. of intermediates = 250 kJ Ea for step 2 = 300 kJ Ea for step 1 = 200 kJ

21. Which step in the reaction in the previous question is more likely the “rate

determining step”? Explain your choice. 22. Use the P.E. diagram below to answer the questions that follow:

a) What is the P.E. of the reactants? b) What is the P.E. of the products? c) What is the P.E. of the intermediates? d) What is the activation energy of step 1? e) What is the activation energy of step 2? f) Which step is more likely the rate-determining step? Explain. g) What is ∆H for the overall reaction? h) What is ∆H for step 2? i) Is step 1 endothermic or exothermic? Explain. j) What is the P.E. of the second Transition State (activated complex)?


800 kJ

700 kJ

600 kJ

500 kJ

400 kJ

300 kJ

200 kJ

100 kJ

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l E

nerg

y

Reaction Coordinate


23. Define the term “activated complex”. Why is it a “complex”? Why is it “activated”? Why is it sometimes called the “Transition State”?

24. What is a “rate law”? 25. The rate law for the reaction between bromine and nitrogen monoxide is given

below: Br2 + 2 NO → 2 NOBr rate = k [Br2] [NO]

Indicate whether the following statements are true or false. Explain your reasoning.

a) The reaction is 1st order with respect to bromine. b) The reaction is 2nd order with respect to nitrogen monoxide. c) The reaction is 2nd order overall. d) The reaction rate will double if the concentration of bromine is doubled. e) The reaction rate will double if the concentration of nitrogen monoxide is doubled. f) The reaction rate will double if the concentrations of both reactants are doubled

at the same time. g) The reaction cannot involve a single-step reaction mechanism. (This is not an

elementary reaction) 26. Based on the rate law given in the previous question, propose two simple two-step

reaction mechanisms for the reaction. Identify any “intermediates” in your mechanisms. Identify the rate-determining step in each mechanism.

27. The rate law for the decomposition of ozone into oxygen gas is given below:

2 O3 → 3 O2 rate = k [O3]

Indicate whether the following statements are true or false. Explain your reasoning.

a) The reaction is 2nd order with respect to the ozone. b) The reaction involves a single collision between two ozone molecules. c) If the concentration of ozone is tripled, the rate of decomposition will increase 9x.

28. Based on the rate law given in the previous question, propose a simple two-step

reaction mechanism for the reaction. Identify any “intermediates” in your mechanism. Identify the rate-determining step in the mechanism.

29. A reaction has the experimental rate law, rate = k[A]2[B].

a) What is the overall order of the reaction? b) What would be the effect on the rate if [A] were tripled? c) What would be the effect on the rate if [B] were halved?


d) What would be the effect on the rate if [A] were doubled while [B] was tripled?


30. The hypochlorite ion, ClO-, undergoes a decomposition into the chlorate and chloride ions. A two-step reaction mechanism has been proposed for the decomposition:

step 1: (slow) ClO- + ClO- → ClO2- + Cl-

step 2: (fast) ClO2- + ClO- → ClO3

- + Cl- a) What is the net (overall) reaction for the decomposition reaction? b) What intermediate is formed in this reaction mechanism? c) Which step is the rate-determining step? Explain.

31. The reaction, 2 NO + 2 H2 → N2 + 2 H2O, has been studied at 904°C:

Reactant Concentrations (mol/L) [NO] [H2]

Rate of Appearance of N2 (mol/L·s)

0.420 0.122 0.136 0.210 0.122 0.0339 0.210 0.244 0.0678 0.105 0.488 ?

a) Determine the rate law for the reaction. b) What is the overall order of the reaction? c) Calculate the value of the rate constant for the reaction at 904°C. d) Calculate the rate of the reaction in trial 4 in the data table. e) Notice that the rates are measured for the “appearance of N2”. What would be

the rate of appearance of “H2O” in trial 1? Explain. 32. The reaction, 2 NO + O2 → 2 NO2, has been studied at 940°C:

Reactant Concentrations (mol/L) [NO] [O2]

Rate of Disappearance of NO (mol/L·s)

0.020 0.010 1.0 x 10-4

0.040 0.010 4.0 x 10-4

0.020 0.040 4.0 x 10-4

0.010 0.020 ? a) Determine the rate law for the reaction. b) What is the overall order of the reaction? c) Calculate the value of the rate constant for the reaction at 940°C. d) Calculate the rate of the reaction in trial 4 in the data table.


e) Notice that the rates are measured for the “disappearance of NO”. What would be the rate of appearance of “NO2” in trial 1? Explain.


33. The reaction, O3 + NO → O2 + 2 NO2, has been studied at 750°C:

Reactant Concentrations (mol/L) [O3] [NO]

Rate of Appearance of NO2 (mol/L·s)

0.21 0.21 0.016 0.42 0.21 0.032 0.63 0.63 0.144

a) Determine the rate law for the reaction. b) What is the overall order of the reaction? c) Calculate the value of the rate constant for the reaction at 904°C. d) Calculate the rate of the reaction if the concentrations in trial 1 were both

doubled. e) Notice that the rates are measured for the “appearance of NO2”. What would be

the rate of disappearance of ozone in trial 1? Explain. 34. A chemist collected data for the reaction: CH3Cl + H2O → CH3OH + HCl

Reactant Concentrations (mol/L) [CH3Cl] [H2O]

Rate of Appearance of CH3OH (mol/L·s)

0.50 0.50 22.700 0.75 0.50 34.050 0.50 0.75 51.075

? 0.125 2.128

a) Determine the rate law for this reaction. b) What is the overall order of the reaction? c) Calculate the rate constant for the reaction at this temperature. d) What was the concentration of chloromethane in trial 4?

35. Explain why the rate of a chemical reaction such as: 2 NO + O2 → 2 NO2 , is likely to be most rapid at the beginning of the reaction. In other words, why does the reaction slow down as it progresses?

36. In an experiment, a student reacts magnesium with excess hydrochloric acid:

trial 1: 1.0 g Magnesium Ribbon is added to 1.0-M HCl at 20°C trial 2: 1.0 g Magnesium Ribbon is added to 0.50-M HCl at 20°C trial 3: 1.0 g Magnesium powder is added to 1.0-M HCl at 20°C trial 4: 1.0 g Magnesium powder is added to 1.0-M HCl at 40°C

a) Rank the four trials in the expected order of increasing reaction rates. b) Which trial would require the least time for the magnesium to dissolve? c) If the student collected the hydrogen gas produced in the reactions in 4 identical

balloons, which balloon would have the largest volume. Explain your answer. (Assume the balloons have been tied and cooled to room temperature).


d) If the student performed the four trials simultaneously, collecting the hydrogen gas in balloons, what would an observer notice about the rate of filling of the balloons?

AP Chemistry 42S Supplementary Problems Chemical Equilibrium… page 10

37. Explain the effect of a catalyst on the activation energy of a reaction. Does the catalyst affect the heat of reaction? Sketch a P.E. diagram to help explain your answer.

38. Explain why a mixture of methane gas and oxygen can stand quietly for years if kept

away from sparks or flame, but will explode violently if a spark is passed through the container.

Chemical Equilibrium 1. What is a “reversible” reaction? Give an example of a reaction that is NOT

reversible. 2. What is the definition of chemical equilibrium? 3. Why is equilibrium sometimes referred to as “dynamic” equilibrium? 4. Water in an open glass will evaporate completely. If the glass is sealed, the water

appears to stop evaporating. Does evaporation still occur in the sealed glass? How is this an example of a physical equilibrium?

5. Dry ice, solid carbon dioxide, sublimes at room temperature. One piece of dry ice is

placed in a sealed container and another is placed in an open beaker. The next day, the dry ice in the sealed container is seen to have decreased in mass only slightly, whereas there is no more dry ice in the open beaker. Explain these observations.

6. Write equilibrium expressions for each reaction below:

a) S (s) + O2 (g) SO2 (g) b) C6H12O6 (s) + 6 O2 (g) 6 CO2 (g) + 6 H2O (l) c) 2 NO2 (g) 2 NO (g) + O2 (g) d) 2 BrCl (g) Br2 (g) + Cl2 (g) e) H2O (l) H2 (g) + ½ O2 (g)

7. Write equilibrium expressions for each reaction below:

a) 2 H2O (l) H3O+ (aq) + OH- (aq) b) Fe3+ (aq) + SCN- (aq) FeSCN2+ (aq) c) C (s) + H2O (g) CO (g) + H2 (g) d) CaO (s) + CO2 (g) CaCO3 (s) e) NH3 (g) + HCl (g) NH4Cl (s)

8. Consider the reaction, 2 HCl (g) H2 (g) + Cl2 (g)

a) Write the equilibrium expression for the reaction. b) Write the new equilibrium expression for the reaction if the coefficients are all

doubled: 4 HCl (g) 2 H2 (g) + 2 Cl2 (g). c) Write the new equilibrium expression for the reaction if the reaction is written in

reverse: H2 (g) + Cl2 (g) 2 HCl (g).


d) Write the new equilibrium expression for the reaction if the coefficients are all halved: HCl (g) ½ H2 (g) + ½ Cl2 (g).


9. The equilibrium constant (K) for the reaction, A 2 B, is “9” at 20°C. Calculate the value of the equilibrium constant at the same temperature for the reaction as balanced below:

a) 2 A 4 B b) 2 B A c) ½ A B

10. Consider the reaction, H2 (g) + I2 (g) 2 HI (g) K = 45.9

a) What is K for the reverse reaction? b) What is K for the reaction, ½ H2 (g) + ½ I2 (g) HI (g) ?

11. Describe the significance of the size of an equilibrium constant on the extent of a

chemical reaction. 12. What does it mean to say “reactants are favored” in a chemical reaction? What is

the size of the equilibrium constant for this type of reaction? 13. Silver chloride is almost completely ISOLUBLE in water. The equation for the

dissolving of silver chloride is : AgCl (s) Ag+ (aq) + Cl- (aq). Predict whether the equilibrium constant

14. State whether the following statements are true or false. Explain your answers.

a) Only gases are included in equilibrium expressions. b) When a reaction reaches equilibrium, it appears to “stop”. c) Equilibrium requires a “closed” system. d) The value of an equilibrium constant depends on how a reaction is balanced. e) The value of an equilibrium constant depends on temperature. f) Reactions with very small equilibrium constants are considered to be “complete”

reactions. g) Reactions with large equilibrium constants favor the formation of products. h) A large equilibrium constant means a reaction will be very fast.

15. How does a catalyst affect the value of an equilibrium constant? Explain. 16. For the reaction, 2 H2S (g) 2 H2 (g) + S2 (g), the equilibrium concentrations at

1130°C are: [H2S] = 0.15 M, [H2] = 0.010 M, and [S2] = 0.051 M. Write the equilibrium expression for the reaction and calculate K.


17. At a high temperature, the following reaction reaches equilibrium:

N2 (g) + O2 (g) 2 NO (g)

An analysis of the equilibrium mixture in a 2.0-L flask shows 0.80 mol NO, 0.64 mol N2 and 0.42 mol O2. What is the value of the equilibrium constant for the reaction at this temperature?


18. A chemist studying the reaction, H2 (g) + I2 (g) 2 HI (g) placed hydrogen and iodine into a 10.0-L vessel. At equilibrium, she found 0.222 mol H2, 0.222 mol I2 and 1.56 mol HI.

a) What is the value of the equilibrium constant for the reaction as written above? b) What is the value of the equilibrium constant for the reverse reaction?

19. While studying the Haber process for the production of ammonia gas from its

elements, a chemist performed three different experiments at the same temperature:

Initial Concentration Equilibrium ConcentrationN2 1.00 M 0.921 M H2 1.00 M 0.763 M Experiment 1

NH3 - 0.157 M N2 - 0.399 M H2 - 1.197 M Experiment 2

NH3 1.00 M 0.203 M N2 2.00 M 2.59 M H2 1.00 M 2.77 M Experiment 3

NH3 3.00 M 1.82 M

a) Write a balanced chemical equation for the Haber Process (all reactants & products are gases).

b) Write the equilibrium expression for the reaction. c) For each of the three experiments, calculate the value of the equilibrium

constant. d) Does the value of K for a reaction appear to depend on the initial concentrations?

20. A student placed 4.00 mol of HI into a 10.0-L vessel and allowed the HI to

decompose into H2 and I2. Use your answer to question 18 to calculate the number of moles of H2 and I2 that will be present when the system reaches equilibrium.

21. At 80°C, the gas BrNO will decompose slightly according to the equation:

2 BrNO (g) Br2 (g) + 2 NO (g)

When 2.0 mol of BrNO are placed into a 5.0-L vessel at 80°C, 5.0% of the molecules are found to decompose. Calculate the value of the equilibrium constant.


22. 4.00 mol of SO2 gas and 1.50 mol of O2 gas are placed in a 2.00-L flask and allowed to reach equilibrium in the reaction:

2 SO2 (g) + O2 (g) 2 SO3 (g)

At equilibrium, there is 0.80 mol of O2 remaining in the flask. Calculate K at this temperature.


23. A substance (CD) decomposes into C and D: CD (g) C (g) + D (g). At a certain temperature, 15% of the CD is decomposed when equilibrium is reached. If the initial concentration of the CD is 0.20 M, what is K for the reaction?

24. A reaction may be represented by: A (g) + B (g) AB (g). At a given

temperature, 1.0 mol of A and 1.0 mol of B are placed into a 2.0-L container and allowed to reach equilibrium. Analysis showed that the equilibrium concentration of AB was 0.40 M.

a) What percent of “A” was converted to AB? b) What is the value of K for the reaction at this temperature?

25. F2 reacts with H2 according to: H2 (g) + F2 (g) 2 HF (g). A student placed 0.50

mol each of the two reactants into a 1.0-L vessel and allowed the system to reach equilibrium. The equilibrium concentration of H2 was 0.10 M. What is the equilibrium constant for the reaction at this temperature?

26. Consider the equilibrium reaction, A (g) 2 B (g) + C (g). When 1.00 mol of A is

placed into a 4.00-L vessel at a given temperature, the concentration of C at equilibrium is found to be 0.050 M. What is the value of K for the reaction?

27. Students were studying the reaction: PCl3 (g) + Cl2 (g) PCl5 (g). An equilibrium

mixture of the three gases in a 4.0-L flask was found to contain 2.40 mol PCl3, 2.00 mol Cl2 and 1.20 mol PCl5. Calculate the value of the equilibrium constant for the reaction.

28. A chemist was studying the reaction: SO2 (g) + SCl2 (g) 2 SOCl K = 22.4.

An equilibrium mixture of the three gases was found to contain 0.20 M SO2 and 0.40 M SCl2. Calculate the concentration of SOCl in the mixture of gases.

29. A researcher placed 2.0 mol of HBr (g) into a 5.0-L tank and allowed it to

decompose: 2 HBr (g) H2 (g) + Br2 (g). When equilibrium was reached, 15% of the HBr had decomposed. What is the value of the equilibrium constant for the reaction?

30. If the same researcher repeated his experiment from the previous question, starting

with 5.0 mol of HBr, what percent of the HBr would decompose? Use the value of K to calculate the answer!

31. In a sealed container, nitrogen dioxide is in equilibrium with dinitrogen tetraoxide:

2 NO2 (g) N2O4 (g). The equilibrium constant for the reaction is 1.15 at 55°C. If the equilibrium concentration of NO2 at 55°C is found to be 0.050-M, what is the concentration of N2O4 in the vessel?


32. A student placed 0.20-mol of phosphorus pentachloride gas into a 20.0-L flask and allowed it to decompose into phosphorus trichloride gas and chlorine gas. At equilibrium, she found 0.040 mol of the PCl5 left in the flask. Calculate K for the reaction at this temperature: PCl5 (g) PCl3 (g) + Cl2 (g)


33. 5.00 g of HBr (g) was placed into a 15.0-L flask and allowed to decompose into H2 and Br2. At equilibrium, 2.21 g of Br2 was discovered. Calculate K for the reaction at this temperature: 2 HBr (g) H2 (g) + Br2 (g)

34. When studying the reaction, 2 CO2 (g) 2 CO (g) + O2 (g), a chemist placed 10.0 g

of carbon dioxide gas into a 2.00-L vessel. She found that 8.0% of the CO2 decomposed in reaching equilibrium. Calculate K for the reaction at this temperature.

35. A student placed 10.00 g of solid magnesium carbonate into a 5.00-L container and

heated the system to 80°C. The MgCO3 decomposed slightly and at equilibrium there was only 9.15 g of the MgCO3 left in the container. Calculate K for the reaction at this temperature: MgCO3 (s) MgO (s) + CO2 (g).

36. When 0.250 mol of CO2 is placed into a container, some of it decomposed according

to the reaction: 2 CO2 (g) 2 CO (g) + O2 (g). At equilibrium, 0.040 mol of CO are found in the container. Calculate K for the reaction.

37. A chemist is studying the reaction, 2 BrCl (g) Br2 (g) + Cl2 (g) K = 0.11.

What will be the equilibrium concentrations of all three species if 0.20 mol of BrCl is placed into a 2.00-L flask and allowed to decompose?

38. At very high temperatures, Cl2 molecules decompose slightly: Cl2 (g) 2 Cl (g).

The equilibrium constant for the reaction is small (K = 1.2 x 10-6 at 1000°C). What will be the concentration of Cl atoms in a flask where the initial concentration of Cl2 molecules is 0.20 M?

39. A chemist was studying the reaction, 2 NO2 (g) NO (g) + NO3 (g) K = 0.085.

If 5.0 mol of NO2 are placed into a 10.0-L vessel, what mass of NO will be in the container when the system reaches equilibrium?

40. 3.00 mol of H2 gas is placed in a 10.0-L flask along with 3.00 mol of F2 gas and

allowed to react until equilibrium was reached: H2 (g) + F2 (g) 2 HF (g) If K = 115 at this temperature, calculate the concentration of HF gas in the container at equilibrium.

41. How is the reaction quotient, Q, calculated? How is it different from K? What is it

used for? 42. The equilibrium constant for the Haber process, N2 (g) + 3 H2 (g) 2 NH3 (g) is

0.060 at a given temperature. A reaction vessel contains 2.00 M nitrogen, 2.00 M hydrogen and 2.00 M ammonia.

a) Use calculations to show that the system is NOT at equilibrium. b) Which direction will the reaction shift to reach equilibrium?


43. Carbon monoxide gas reacts with steam to produce carbon dioxide and hydrogen (all species are gases). At 700 K, the equilibrium constant for this reaction is 5.10. Calculate the equilibrium concentrations of all species if 1.00 mol of each gas is placed into a 2.00-L vessel and allowed to reach equilibrium.


44. While studying the reaction, 2 HI (g) H2 (g) + I2 (g), a chemist placed 1.0 mol of each gas into a 5.00-L flask. If the equilibrium constant for the reaction at this temperature is 22.0, calculate the concentrations of all three gases when the system reaches equilibrium.

45. Consider the equilibrium where ammonia decomposes into its elements:

2 NH3 (g) N2 (g) + 3 H2 (g)

Suppose the original concentrations of the three species are all 1.0 M. Which of the following are possible values for the concentrations at equilibrium?

[NH3] [N2] [H2]

a) 0.8 M 1.1 M 1.4 M b) 1.2 M 0.9 M 0.7 M c) 1.2 M 1.2 M 1.2 M d) 0.6 M 1.2 M 1.6 M e) 0.2 M 1.4 M 2.2 M

46. Given: 2 SO2 (g) + O2 (g) 2 SO3 (g) ∆H = -192 kJ/mol. For this reaction,

describe how the following should be adjusted to give a high equilibrium concentration of SO3 (g).

a) Temperature b) Pressure c) Concentration of SO2 d) Concentration of O2

47. Hydrogen peroxide decomposes according to: H2O2 (aq) H2O (l) + ½O2 (g).

The reaction is endothermic. After equilibrium has been reached, predict what would happen to the amount of hydrogen peroxide in the system if…

a) the system was placed in an ice bath b) a catalyst is added c) more water is added


48. Consider this reaction which is at equilibrium in a closed vessel: Heat + 2 SO3 (g) 2 SO2 (g) + O2 (g)

Explain what would happen to the amount of O2 in the vessel if the following stresses were applied: a) more SO3 is added c) the system is heated e) SO2 is removed b) a catalyst is added d) the volume is doubled f) pressure increases


49. Consider the reaction which is at equilibrium in a closed flask:

Fe3O4 (s) + 4 H2 (g) 3 Fe (s) + 4 H2O (g) ∆H = -52 kJ/mol

Explain what would happen to the mass of iron oxide in the container if the following stresses were applied:

a) hydrogen gas is removed from the flask b) water vapor is added to the flask c) the temperature is raised d) iron metal is added to the flask e) the total pressure in the flask is increased f) hydrogen gas is added to the flask

50. Consider the endothermic reaction at 456 °C:

2 NOBr (g) 2 NO (g) + Br2 (g)

The reaction is at equilibrium in a closed flask. Explain what would happen to the amount of Br2 in the flask if the following stresses were applied:

a) the volume of the flask is doubled b) the flask is cooled down c) NO is removed from the flask d) NOBr is added to the flask e) an inert gas (Argon) is added to the flask, increasing the total pressure f) a catalyst is added to the flask

51. State LeChatelier’s Principle. 52. Explain how changing the volume of a flask affects the equilibrium position of the

system in the flask. Do volume changes always affect the equilibrium position of a system? Explain your answer.

53. For the reaction in question 50, decide whether the following statements are true or

false.

a) ∆H for the reaction is positive b) At equilibrium, increasing [NO] will cause an increase in [NOBr] c) At equilibrium, removing Br2 causes the system to shift left d) Adding more Br2 will change the value of K e) Increasing temperature will change the value of K f) Adding a catalyst will increase the yield of product


AP Chemistry 42S Supplementary Problems Acids & Bases… page 17

Acids & Bases 1. List several properties of acids and bases. 2. Give three examples of commonly encountered acids and three of bases. 3. What does it mean to say that an acid is “strong”? Is this the same as saying the

acid is “concentrated”? Explain. 4. What is a monoprotic acid? a diprotic acid? Give an example of each. 5. Write balanced equations showing the dissociation of the following acids in water. If

the acid is polyprotic, write the equations showing the step-wise dissociation of the acid. a) nitric acid d) carbonic acid b) hydrochloric acid e) phosphoric acid c) acetic acid f) sulfuric acid

6. Write the Ka expressions for the acid dissociations in the previous question. 7. Identify the following acids as strong or weak. Explain.

a) nitric acid d) hydrofluoric acid b) carbonic acid e) hydrochloric acid c) acetic acid f) phosphoric acid

8. What is an “amphoteric” substance? Decide which of the following species are

amphoteric. If the substance is amphoteric, write balanced equations to show the substance acting as an acid and as a base. a) H2O b) PO4

3- c) HCO3- d) H2PO4

- e) S2-

9. What equation is used to calculate “pH”? 10. A solution whose pH is 4.00 has _____ times the concentration of hydronium ion as

a solution whose pH is 7.00. 11. A solution whose pH is 3.50 has 1% the concentration of hydronium ions as a

solution whose pH is _____. 12. Calculate the hydronium ion concentration, [H3O+], in a solution whose pH is…

a) 1.00 d) 3.52 b) 2.00 e) 4.91 c) 8.00 f) 11.22

13. Pure water has a pH of 7.00. What is the [H3O+] in pure water? Write a balanced

chemical equation that shows how the hydronium ions are formed in water.


14. The Kw for water at 50°C is 5.47 x 10-14. What is the pH of pure water at 50°C?


15. Make the following inter-conversions. In each case, tell whether the solution is acidic or basic.

pH [H3O+] [OH-] a) 1.00 ? ? b) 10.50 ? ? c) ? 1.3 x 10-5 ? d) ? ? 2.3 x 10-4

16. Describe an acidic solution at 25°C in terms of a) [H3O+] and [OH-] and b) pH. 17. Write the formula for the conjugate base for each of the following Bronsted-Lowry

acids: a) H2O b) HNO2 c) NH4

+ d) H3PO4 18. In the following equations, identify the Bronsted-Lowry acids (A) and bases (B):

a) HSO4- + H2O SO4

2- + H3O+ b) CH3COO- + HCl Cl- + CH3COOH c) NH3 + H2O NH4

+ + OH- d) HCO3

- + HSO4- SO4

2- + H2CO3 19. Write equations for the following acting as Bronsted-Lowry bases in aqueous

solutions: a) fluoride ion, F- b) methylamine, CH3NH2 c) sulfate, SO4

2-

20. Recall that the products of an acid-base neutralization reaction are a salt and water.

Write balanced equations for the following neutralization reactions. In each case, the neutralization reaction is complete. a) magnesium hydroxide, Mg(OH)2, and hydrobromic acid, HBr b) sulfuric acid, H2SO4, and sodium hydroxide, NaOH c) phosphoric acid, H3PO4, and potassium hydroxide, KOH d) ammonium hydroxide, NH4OH, and hydrogen carbonate, HCO3

- 21. What will be the concentration of sulfuric acid solution, if 25.0 mL of 0.260-M solution

is diluted to 35.0 mL? 22. How much water must be added to 25.0 mL of 6.0-M H2SO4 to make a 1.0-M

solution? 23. Rank the following acids in order of increasing strength: HClO4, HNO2, CH3COOH,

HF 24. Use Le Chatelier’s principle to explain what happens to [OH-] when NH4Cl is added

to a solution of ammonia, NH3. What happens to [H3O+] at the same time? 25. Calculate the pH of 0.020-mol/L HCl. The Ka for hydrochloric acid is large. 26. Calculate the pH of 0.00050-M HNO3. Nitric acid is a strong acid.


27. Calculate the pH of 0.0035-M HBr. Hydrobromic acid is a strong acid.


28. Calculate the pH of 0.0025-M HF. The Ka for hydrofluoric acid is 6.8x 10-4. 29. Calculate the pH of 0.0052-M CH3COOH. The Ka for acetic acid is 1.8 x 10-5. 30. Calculate the pH of 0.042-M benzoic acid, C6H5COOH. The Ka is 6.5 x 10-5. 31. Calculate the pH of a solution of formic acid, HCOOH, made by dissolving 1.00 g of

the acid in enough water to make 500.0 mL of solution. The Ka of formic acid at 25 °C is 1.8 x 10-4.

32. Calculate the pH of a solution of benzoic acid, C6H5COOH, made by dissolving 2.50

g of the acid in enough water to make 100.0 mL of solution. The Ka for benzoic acid is 6.5 x 10-5 at 25°C.

33. Calculate the pH of a 0.00105-M solution of sulfuric acid, H2SO4. Ka1 is large, and

Ka2 is 0.012. Consider each step in the dissociation separately! 34. Calculate the pH of 0.125-M sulfuric acid solution. Ka1 is large, and Ka2 is 0.012. 35. What was the %-dissociation for the second step of the dissociation of sulfuric acid in

the previous question? 36. What assumption can be made in calculating the pH of phosphoric acid solutions?

Why is the assumption valid? 37. Calculate the pH of 0.012-M phosphoric acid. Ka1 is 7.5 x 10-3, Ka2 is 6.2 x 10-8 and

Ka3 is 4.8 x 10-13. 38. Calculate the pH of 0.0025-M H3PO4. Use the Ka values from the previous question. 39. Calculate the percent ionization (dissociation) in each of the following solutions. The

Ka for hydrofluoric acid is 6.8 x 10-4. a) 0.50-M HF b) 0.050-M HF c) 0.0050-M HF

40. Based on your answers to the previous question, describe what happens to the

fraction of acid molecules that have dissociated as an acid solution becomes more dilute.

41. The percent ionization of a 0.10-M solution of a weak monoprotic acid, HA, is found

to be 3.42%. What is the Ka for the acid? 42. The pH of a 0.063-M solution of hypobromous acid, HBrO, is measured to be 4.95.

Calculate the Ka for the acid. 43. The pH of a 0.050-M solution of trichloroacetic acid, CCl3COOH, is 1.40. Calculate

the Ka for the acid. Is this a stronger acid than acetic acid (Ka = 1.8 x 10-5)?


44. A typical sample of vinegar has a pH of 3.0. Assuming that the vinegar is only a dilute solution of acetic acid (Ka = 1.8 x 10-5), calculate the concentration of CH3COOH in the vinegar.


45. A 0.015-M solution of the weak acid, HOCN, has a pH of 2.67. Calculate the Ka of this acid. What is the percent dissociation of the acid?

46. Calculate the pH of the following solutions of strong bases:

a) 0.015 M NaOH b) 0.0020-M Ba(OH)2 47. What is the pH of a saturated solution of calcium hydroxide, Ca(OH)2? This salt is

only sparingly soluble and has a Ksp = 7.9 x 10-6. 48. Calculate the pH of 0.10-M ammonia solution, NH3 (aq). The Kb of ammonia is

1.8 x 10-5 at 25°C. 49. Calculate the pH of 0.015-M ammonia solution. The Kb for ammonia is 1.8 x 10-5. 50. Methylamine is a weak base (Kb = 5.0 x 10-4). What % of CH3NH2 molecules are

ionized in a 0.060-M solution of methylamine? What is the pH of the solution? 51. Pyridine is a weak base (Kb = 1.5 x 10-9). What % of C5H5N molecules are ionized in

a 0.110-M solution of pyridine? What is the pH of the solution? 52. What equation relates the Ka of an acid with the Kb of its conjugate base? 53. Given the following acids and their Ka’s, write the formula of the conjugate bases and

calculate their base dissociation constants, Kb. a) HCN, Ka = 4.0 x 10-10 b) H3PO4, Ka = 0.0075

54. Calculate the pH of a solution made by dissolving 5.00 g of sodium fluoride, NaF, in

50.0 mL water. Is the solution acidic, basic, or neutral? The Ka of HF is 7.2 x 10-4. 55. Calculate the pH of a solution made by dissolving 1.50 g of NaHCO3 in 200.0 mL of

water. Use the Ka values in your text or on the class wall-chart to help answer this. 56. Predict whether the following 0.10-M solutions will be acidic, basic, or neutral. (Do

not actually calculate the pH of the solutions!) a) NaCl b) NH4Br c) K2SO4 d) NH4F

57. Predict whether the following 0.10-M solutions will be acidic, basic, or neutral. (Do

not actually calculate the pH of the solutions!) a) CaI2 b) NaCH3COO c) NH4CH3COO d) Na2CO3

58. What is an acid-base indicator? Describe the color changes involved in Litmus and

in phenolphthalein indicators.


59. In a titration experiment, a 0.925-g sample of potassium hydrogen phthalate (KHP, molar mass 204.22 g/mol) required 22.35 mL of sodium hydroxide solution for neutralization. Calculate the concentration of the NaOH solution. Recall that KHP is a weak monoprotic acid that can be represented as “HA” in an equation.


60. In another experiment, the NaOH solution from the previous question was used to titrate a sample of acetic acid, CH3COOH. If a 25.00-mL sample of the acetic acid required 16.93 mL of the NaOH solution to reach the endpoint of the titration, what was the concentration of the acetic acid?

61. What is the endpoint of a titration? How is it different from the equivalence point of

the titration? 62. How should an indicator be chosen in an acid-base titration? In other words, what

factor(s) must be considered when selecting an acid-base indicator for a titration? 63. A 1.023-g sample of KHP (molar mass, 204.22 g/mol), required 17.33 mL of KOH

solution to reach the endpoint of a titration. Calculate the concentration of the KOH solution. Recall that KHP is a weak monoprotic acid.

64. A 10.00-mL sample of phosphoric acid, H3PO4, required 42.08 mL of the KOH

solution from the previous question to reach the third equivalence point. Calculate the concentration of the phosphoric acid solution.

65. A 25.00-mL sample of ammonia (NH3) required 22.84 mL of 0.0875-M sulfuric acid

(H2SO4) for complete neutralization. Calculate the concentration of the ammonia solution.

66. What is a “standard” solution? What does it mean to “standardize” a solution? 67. You require 36.78 mL of 0.0105-M HCl solution to reach the equivalence point in a

titration of a 0.525-g sample of a weak base. Calculate the molar mass of the base. Assume that the base reacts with the HCl in a 1:1 ratio (the base can be represented as “B” in a chemical equation).


68. Assume that you dissolve 0.235 g of the weak acid benzoic acid (C6H5COOH) in exactly 100.0 mL of water and then titrate a 10.00-mL sample of the solution with 0.108-M ammonia (NH3). What volume of the ammonia will be required to neutralize the benzoic acid sample?

AP Chemistry 42S Supplementary Problems Solubility Equilibrium… page 22

Solubility Equilibrium 1. For equilibrium to exist, two opposing processes must be occurring at the same rate.

Use this definition to explain how a “saturated solution” is an example of equilibrium. 2. Write balanced equations for the dissociation of the following salts. Include the

solubility product expressions for each salt, also.

a) silver phosphate, Ag3PO4 d) lead iodide, PbI2b) magnesium fluoride, MgF2 e) cerium(III) hydroxide, Ce(OH)3c) calcium carbonate, CaCO3 f) aluminum chromate, Al2(CrO4)3

3. What mass of sodium hydroxide, NaOH, must be dissolved to prepare 2.00 L of

0.0500-M solution? 4. What mass of iron(III) nitrate, Fe(NO3)3 , must be dissolved to prepare 500.0 mL of

0.120-M solution? 5. A student used a pipette to transfer 25.00 mL of 0.50-M sulfuric acid to a clean 500-

mL volumetric flask. She then filled the flask with distilled water and shook to mix. What is the concentration of the new solution?

6. A chemistry student mixed 20.0 mL of 0.010-M NaOH with 25.0 mL of 0.025-M

KNO3. What is the concentration of each ion in the mixed solution? No reaction occurs during the mixing.

7. Indicate whether the following statements are true or false.

a) Sodium salts are generally very soluble in water. b) The solubility of salts is generally higher in cold water than in hot water. c) The solubility of gases is generally higher in cold water than in hot water. d) Salts containing the chloride ion are soluble, except silver chloride and lead

chloride. e) Salts containing the nitrate ion are insoluble. f) “Sparingly soluble” is a more accurate description for salts considered to be

“insoluble” in water. g) A salt with a small Ksp is generally very soluble in water. h) To compare solubilities of two salts in water, one can simply compare the sizes

of their solubility product constants. 8. A student mixed 1.0-M solutions of the following salts. In which case(s) will a

precipitate likely form? Identify any precipitates that form.

a) NaI is mixed with KNO3 e) KF is mixed with AgNO3 b) Pb(NO3)2 is mixed with KNO3 f) ZnCl2 is mixed with Pb(NO3)2 c) FeCl3 is mixed with KOH g) CaCl2 is mixed with NaOH


d) CuSO4 is mixed with CaCl2 h) KOH is mixed with NaCl


9. At 25°C, the solubility product constant for silver bromide, AgBr, is 5.0 x 10-13. Calculate the molar solubility of silver bromide at this temperature.

10. Calculate the mass of AgBr that can be dissolved in 2.00 L of water at 25°C. 11. As temperature increases, does Ksp for a salt increase or does it decrease? Justify

your answer. 12. At 25°C, the Ksp of calcium fluoride (CaF2) is 3.9 x 10-11. What is the molar solubility

of calcium fluoride in water at 25°C? 13. The Ksp of magnesium hydroxide, Mg(OH)2, is 7.1 x 10-12 at 25°C. Calculate the

solubility of this salt, expressed in “grams per liter”. 14. Lead sulfate, PbSO4, has a solubility product constant, Ksp, of 6.3 x 10-7 at 25°C.

The Ksp of lead fluoride, PbF2, is smaller, 3.6 x 10-8 at 25°C. Calculate the molar solubility of each salt. Which salt is more soluble in water?

15. The Ksp of aluminum hydroxide, Al(OH)3, is very small (3.0 x 10-34). How much

water is required to dissolve 1.0 g of this salt? 16. Calculate the volume of water needed to dissolve 0.50g of silver chromate. The Ksp

of Ag2CrO4 is 1.2 x 10-12 at 25°C. 17. What mass of lead iodide, PbI2, is needed to make 1.00 L of saturated solution at

25°C? The Ksp of lead iodide is 7.9 x 10-9 at 25°C. 18. What mass of iron(II) hydroxide, Fe(OH)2, is needed to prepare 500.0 mL of

saturated solution? The Ksp of Fe(OH)2 is 7.9 x 10-16 at 25°C. 19. A student added a teaspoon of barium sulfate, BaSO4, to a glass of water and stirred

for several minutes. Use calculations to predict whether the salt will appear to dissolve in the water. The Ksp of barium sulfate is 1.1 x 10-10.

20. Lead iodide, PbI2, is a bright yellow precipitate that makes a colorful demonstration.

The Ksp of lead iodide is 7.9 x 10-9. A chemist added 1.00 g of sodium iodide, NaI, to 2.00 L of 0.100-M lead nitrate, Pb(NO3)2. Will the chemist observe a precipitate?

21. What mass of potassium hydroxide, KOH, can be dissolved in 500.0 mL of 0.20-M

magnesium nitrate, Mg(NO3)2, without causing a precipitate to form? The Ksp of Mg(OH)2 is 7.1 x 10-12.

22. Calculate the mass of Na2CO3 that must be dissolved in 2.00 L of 0.050-M AgNO3 to

make a saturated solution of silver carbonate, Ag2CO3. The Ksp of silver carbonate is 8.1 x 10-12.

23. The solubility of strontium carbonate, SrCO3, is 0.0059 g per 250 mL of water. What

is the Ksp for strontium carbonate?


24. The molar solubility of barium phosphate, Ba3(PO4)2, is 0.0041 mol/L. Calculate the Ksp for this salt.


25. The solubility of calcium hydroxide, Ca(OH)2, is measured to be 0.93 g/L at some temperature. What is the Ksp of the salt at this temperature?

26. You place 2.75 g of BaF2 in 1.00 L of pure water at 25°C. After equilibrium has been

established, the fluoride ion concentration is 0.0150 M. What is the Ksp of the salt? 27. Describe the common ion effect on the solubility of a salt. 28. Compare the solubility of AgCl (Ksp = 1.0 x 10-10) in water to that in 0.10-M NaCl. 29. The Ksp of silver iodide is 8.3 x 10-17. What is the iodide ion concentration of a 1.00-

L saturated solution of AgI, to which 0.020 mol of AgNO3 is added? 30. Calculate the molar solubility of silver thiocyanate, AgSCN, in pure water and in a

solution containing 0.010-M NaSCN. The Ksp of silver thiocyanate is 1.0 x 10-12. 31. What is the solubility, in “grams per milliliter”, of barium fluoride in a solution

containing 1.5 g/L potassium fluoride, KF. The Ksp of BaF2 is 1.7 x 10-6. 32. Calculate the molar solubility of calcium hydroxide, Ca(OH)2, in 0.100-M NaOH. The

Ksp of Ca(OH)2 is 6.5 x 10-6 at 25°C. 33. Rank the following compounds in order of increasing solubility in pure water:

Na2CO3, BaF2, PbCl2, and Ag2CrO4. Refer to the Ksp table in ch. 17 of your text. 34. Rank the following compounds in order of increasing solubility in pure water: CaSO4,

NiCO3, AgOH, and ZnS. Refer to the Ksp table in ch. 17 of your text. 35. Which has a greater solubility in pure water – Ag2SO4 or CaSO4? Refer to the Ksp

table in ch. 17 of your text. 36. Will a precipitate form if 100.0 mL of 0.0030-M NaCl is added to 200.0 mL of 0.0.012-

M AgNO3? The Ksp of silver chloride, AgCl, is 1.0 x 10-10. 37. Will a precipitate form if 10.0 mL of 0.050-M NaOH is added to 500 mL of 0.050-M

Pb(NO3)2? The Ksp of Pb(OH)2 is 1.2 x 10-15. 38. Will a precipitate form if 2.0 g of solid Pb(NO)3 is added to 250 mL of 0.020-M NaI?

The Ksp of PbI2 is 1.4 x 10-8. Assume that no volume change occurs. 39. Will a precipitate form if 50.0 mL of 0.10-M AgNO3 is added to 1.0-L of 0.10-M KCl?

The Ksp of AgCl is 1.0 x 10-10. 40. Will a precipitate form if 25.0 mL of 0.0040-M NaF is added to 75.0 mL of 0.0160-M

Mg(NO3)2? The Ksp of MgF2 is 6.4 x 10-9. 41. Define a “supersaturated solution” in terms of the reaction quotient.


42. A solution is prepared by dissolving 1.40 g of Ag2SO4 in 100.0 mL of hot water. Will a precipitate form if the solution is cooled to 25°C? At 25°C, the Ksp of Ag2SO4 is 1.2 x 10-5.

AP Chemistry 42S Problem Booklet Electrochemistry… page 25

Oxidation, Reduction & Electrochemistry 1. In each of the following reactions, identify the substance being oxidized and the

substance being reduced. Identify the oxidizing agent and the reducing agent.

a) 2 Al + 3 Cl2 → 2 AlCl3 b) FeS + 3 NO3

- + 4 H3O+ → 3 NO + SO42- + Fe3+ + 6 H2O

c) Zn + HCl → ZnCl2 + H2 2. In each of the following reactions, identify the substance being oxidized and the

substance being reduced. Identify the oxidizing agent and the reducing agent.

a) 4 NH3 + 3 O2 → 2 N2 + 6 H2O b) PbO2 + Pb + 2 H2SO4 → 2 PbSO4 + 2 H2O c) Cl2 + 2 I- → 2 Cl- + I2

3. Assign oxidation numbers (states) to each atom in the following species:

a) ClO4- b) NaOCl c) KNO3 d) CH4

4. Assign oxidation states to each atom in the following species:

a) MnO4- b) CrO4

2- c) Cr2O72- d) O3

5. Assign oxidation numbers to the nitrogen atoms below:

a) NO b) N2 c) N3- d) NO2-

e) NO2 f) NH3 g) N2O5 h) N2O3

6. Identify each of the following changes as “oxidation” or “reduction”:

a) MnO4- → MnO2 b) P4O10 → P4O6 c) Fe3+ → Fe2+

7. Identify each of the following changes as “oxidation” or “reduction”.

a) Cr2O72- → Cr3+ b) SO4

2- → SO32- c) Zn → Zn2+

8. Balance the following, using oxidation numbers.

a) HBrO3 + SO2 + H2O → Br2 + H2SO4 b) K2Cr2O7 + H2O + S → KOH + Cr2O3 + SO2 c) PH3 + I2 + H2O → H3PO2 + HI d) Cl2 + KOH → KClO3 + KCl + H2O e) KIO4 + KI + HCl → KCl + I2 + H2O f) HNO2 + HI → NO + I2 + H2O g) As2O3 + Cl2 + H2O → H3AsO4 + HCl


h) MnO2 + H2SO4 + H2C2O4 → MnSO4 + CO2 + H2O


9. Identify which reactions below represent redox reactions. If the reaction is redox, identify what is being reduced & what is being oxidized. Also identify the reducing & oxidizing agents.

a) Li + H2O → LiOH + H2 b) Al + HCl → AlCl3 + H2 c) H2O + SO3 → H2SO4 d) 2 NaOH + H2CO3 → 2 H2O + Na2CO3 e) 2 NaBr + Cl2 → 2 NaCl + Br2 f) BaCl2 + 2 KIO3 → Ba(IO3)2 + 2 KCl

10. Balance the following equations using the half-reactions method, in acidic solution:

a) Cr2O72- + I- → Cr3+ + I2

b) Fe2+ + MnO4- → Fe3+ + Mn2+

c) H2O2 + Br- → H2O + Br2 d) C2H5OH + Cr2O7

2- → CH3COOH + Cr3+ e) Cu + NO3

- → Cu2+ + NO2 11. Balance the following equations using the half-reactions method, in basic solution:

a) ClO- + CrO2- → Cl- + CrO4

2- b) CrO4

2- + S2- → Cr(OH)3 + S c) MnO4

- + SO32- → MnO2 + SO4

2- d) MnO2 + ClO3

- → MnO4- + Cl-

e) Fe(OH)3 + ClO- → FeO42- + Cl-

12. A KMnO4 solution was standardized by titration against As2O3. A 0.1156-g sample of

As2O3 requires 27.08 mL of the KMnO4 for its titration. What is the concentration of the potassium permanganate solution? Assume the reaction occurs in acidic solution.

MnO4- + As2O3 → Mn2+ + H3AsO4

13. The titration of 50.0 mL of saturated solution of sodium oxalate, Na2C2O4, requires

25.8 mL of 0.0214-M K2Cr2O7. Calculate the concentration of the sodium oxalate solution. Assume the reaction occurs in acidic solution.

Cr2O7

2- + C2O42- → Cr3+ + CO2

14. A 1.026-g sample of an iron-containing ore was dissolved in an acid solution and the

iron converted to Fe2+. This solution required 24.35 mL of 0.0195-M KMnO4 to reach the equivalence point of a titration. Calculate the percentage of iron in the sample (by mass).

MnO4- + Fe2+ → Mn2+ + Fe3+

15. A 25.1-mL sample of NaBr requires 24.9 mL of 0.129-M KMnO4 solution to reach the

equivalence point for its oxidation to BrO- in a basic solution. The KMnO4 is reduced to MnO2. Find the concentration of the NaBr solution.

MnO4

- + Br- → MnO2 + BrO-



16. Sketch the electrochemical cells below in your notebook.

i) Label the anode and cathode ii) Label the half-cells as “oxidation” or “reduction” iii) Indicate the direction of electron flow iv) Write the half-reaction occurring in each beaker v) Record E° for each half-cell vi) Calculate E°cell

a) b) 17. What does “standard” refer to in the phrase, “standard electrochemical cell”? 18. What are two common salts used in salt bridges? Describe an alternative device

that could be used in place of a salt bridge in an electrochemical cell. 19. How is the standard reduction potential for a species related to its standard oxidation

potential?


Na2SO4

Zn2+(aq) , 1M

Ag+(aq), 1M

Z INC

S ILVER

Na2SO4

Fe2+(aq) , 1M

Cu2+(aq), 1M

IRON

COPPER


20. Which of the species below is most likely to be reduced? a) Ag, I2 or Zn2+ b) Cl2, Li, or F- c) Cu2+, Cl-, or MnO4

-

21. Which of the species below is most likely to get oxidized?

a) K+, F- or K b) Fe, Zn, or Mg c) H2, Ag, or Na+

22. Draw and completely label each of the electrochemical cells below. Calculate E° for

each cell. a) Al | Al3+ (1 M) || Ag+ (1M) | Ag b) Zn | Zn2+ (1M) || Cl2 (1 atm) | Pt c) Cu | Cu2+ (1M) || Au3+ (1 M) | Au

23. Draw and completely label each of the electrochemical cells below. Calculate E° for

each cell. a) Mg | Mg2+ (1M) || Sn2+ (1M) | Sn b) Al | Al3+ (1M) || Sn2+ (1M) | Sn c) Ni | Ni2+ (1M) || Ag+

(1M) | Ag 24. In each case below, decide if a redox reaction will spontaneously occur. Show

calculations of E° to support your answers. a) A copper wire is placed into a solution of zinc nitrate (contains Zn2+) b) A lead strip is dipped into a solution of copper(II) sulfate (contains Cu2+) c) A silver coin is washed with a solution containing Cl2(aq) d) A nickel coin is dipped into a solution containing Mg2+ e) Sodium iodide (contains I-) is mixed with aqueous bromine, Br2(aq) f) An iron nail is dipped into a solution of nickel(II) sulfate (contains Ni2+)

25. A student dissolved some potassium permanganate in a beaker of water to make a

relatively concentrated solution. She then dipped a piece of magnesium metal into the solution. To her surprise, there was no reaction! a) Calculate E° for the reaction and state whether the reaction should be

spontaneous or non-spontaneous. b) Suggest a reason for the student’s observation.

26. Which of the following metals would be practical ones to use in an experiment to

prepare hydrogen gas from dilute hydrochloric acid solution: iron, sodium, zinc, copper, magnesium. For the metals that you deem to be impractical, explain why.

27. If the Ag/Ag+ half-cell was chosen as the standard reference cell (instead of

hydrogen), what would be the standard reduction potentials of the …

a) Cu2+/Cu half-cell b) H+/H2 half-cell c) Mg2+/Mg half-cell 28. Which is the best reducing agent? the best oxidizing agent? I2, I-, Au, Au3+, Mg, Mg2+ 29. Can a 1.0-M solution of Fe2(SO4)3 be stored in a nickel container? Explain.


30. Which is the best oxidizing agent? the best reducing agent? Sn2+, Sn, K+, K, Cl2, Cl-


31. The following observations were made, and equations for the reactions are given. Using these observations, list the ions with their reduction half-reactions, in order of decreasing ability to gain electrons. (do NOT refer to a table of reduction potentials!)

a) Mn + Zn2+ → Zn + Mn2+ b) Fe + Co2+ → Fe2+ + Co c) Fe + Zn2+ → no reaction

32. In an experiment, strips of gold, silver and tin were placed in solutions containing

solutions of Au3+, Ag+ and Sn2+ ions. The following results were observed:

a) Sn + Au3+ → metallic gold was deposited b) Au + Ag+ → no reaction c) Sn + Ag+ → metallic silver was deposited Arrange the three metal ions as reduction half-reactions, in decreasing ability to gain electrons.

33. Given: H2O, Ni, Cl2, K+, Cl-, Cu+, H2O2

a) Which is most easily oxidized? b) Which is the strongest oxidizing agent? c) Which could reduce Fe3+ to Fe2+? d) Which could oxidize Br- to Br2?

34. Write net ionic equations for these reactions (all aqueous). If no reaction occurs,

write “no reaction”.

a) chlorine + tin(II) nitrate b) iron(II) nitrate + potassium iodide c) silver + hydrochloric acid d) tin + sulfuric acid e) iron(II) nitrate + aluminum nitrate

35. Draw a diagram showing how a copper spoon could be plated with silver metal.

Identify the anode and cathode. Calculate the mass of silver that could be plated by a 1.0-A current flowing for 27 h.

36. How long must an electrolytic cell operate to produce 3.55 g of Cl2(g) with a current

of 6.50 A (from a solution of Cl-). 37. A current of 1.50 A is passed through 250 mL of 0.10-M Cu2+(aq). How long will it

take for all off the copper in the solution to precipitate as copper metal? 38. If you use a current of 2.50 A to deposit nickel from a solution of Ni(NO3)2 for 2.00 h,

how many grams of nickel will be deposited?


39. If you wish to convert 1.00 g of Au3+(aq) to Au(s) in 30 min, what current should you use?


40. A current of 1.00 A is applied to 1.0 L of 1.0-M HCl solution for 24 h, producing hydrogen gas. a) Calculate the number of moles of hydrogen ions that would have been consumed

in the 24 h period. b) What would be the pH of the solution at the end of the 24 h period? c) If the hydrogen gas was collected at 25°C and 100.0 kPa, what volume of gas

would there be? (R = 8.314 kPa·L·mol-1·K-1) 41. A metal forms the chloride compound, MCl4. Electrolysis of the molten metal

chloride by a current of 1.81 A for 25.6 min produced 1.09 g of the metal. Find the molar mass of the metal. Identify the metal from the periodic table.

42. The electrolysis of molten potassium chloride is carried out with a current of 9.78 A.

How long will it take to deposit 17.5 g of potassium metal? 43. What mass of tin can be deposited during the electrolysis of molten SnI2 by a 5.00-A

current for 2 h 30 min? 44. Predict the products at the anode and cathode from the electrolysis of…

a) molten sodium chloride b) molten zinc bromide c) aqueous potassium bromide d) aqueous copper(II) chloride

45. Predict the products at the anode and cathode from the electrolysis of…

a) water b) molten aluminum iodide c) molten silver chloride d) aqueous hydrochloric acid

46. Draw and label the standard electrochemical cell: Mg | Mg2+ (1M) || Ag+ (1M) | Ag

a) Calculate the standard cell potential for the cell. b) Write a balanced equation for the overall reaction occurring in this cell. c) Would the voltage increase, decrease, or stay the same if a bigger magnesium

electrode was used? (hint: LeChatelier’s Principle!) d) Would the voltage increase, decrease, or stay the same if water was added to

the magnesium half-cell? e) Would the voltage increase, decrease, or stay the same if AgNO3 was added to

the silver half-cell? 47. Describe what is meant by “cathodic protection”. Give an example. 48. What processes occur at the anode and cathode in an electrochemical cell. in an

electrolytic cell? 49. What is a battery?


50. What material is the anode in a dry cell (e.g. Duracell)? the cathode?

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