1. Write the Net Ionic Reactions for Each of the Following Reactions

1. Write the net ionic reactions for each of the following reactions.

a. Solid sodium oxide is added to water.

b. A solution of silver nitrate is added to concentrated ammonia

c. A solution of lead (II) nitrate is added to a solution of sodium sulfide.

d. Sulfur dioxide is bubbled into water.

e. Solid aluminum nitrate is added to water.

f. Solid sodium acetate is added to water.

g. A solution of iron (III) chloride is added to a solution of sodium carbonate.

h. Dilute sulfuric acid is added to solid calcium fluoride.

i. Dilute hydrochloric acid is added to a dilute solution of mercury (I) nitrate.

j. Excess concentrated potassium hydroxide solution is added to a precipitate of zinc

hydroxide.

k. The gases boron trifluoride and ammonia are mixed

l. Equimolar amounts of sodium phosphate and hydrochloric acid are mixed

m. Equimolar amounts of phosphoric acid and sodium hydroxide are mixed

n. Dilute hydrochloric acid is added to a solution of potassium sulfite.

2. Methylamine, CH3NH2, is a weak base. (Kb = 4.38 x 10-4). The salts of methylamine are typically salts such as methylammonium nitrate (CH3NH3NO3)

a. What is the pH of 120.0-mL of 0.150 M methylamine?

b. When 3.0-g of methylammonium nitrate, CH3NH3NO3 is added to the solution in (a), what is the final pH? (assume no volume change)

c. To the solution in part (b), 20.0-mL of 0.125-M HCl is added. What is the final

pH?

d. You want to make a buffer of pH 10.5. How many moles of HCl or NaOH must

be added to the solution in part (b) to make this solution.

e. To the solution in part (b), 100.0-mL of water is added. What is the resulting pH?

3. The Ka of lactic acid, HC3H5O3 is 1.38 x 10-4.

a. What is the hydrogen ion concentration of a 0.0250 M solution of lactic acid?

b. When 25.0-mL of 0.0250-M lactic acid is added to 50.0-mL of 0.0125-M NaOH, what is the resulting pH?

c. When 45.0-mL of 0.125-M lactic acid is added to 25.0-mL of 0.200-Msodium

lactate, NaC3H5O3, what is the final pH?

d. A buffer of pH 5.00 is made with lactic acid and sodium lactate. What is the ratio of lactate to lactic acid in the solution?

4. The Ksp of CdCO3 is 5.2 x 10-12 and the Ksp of FeCO3 is 2.1x10-11.

a. What are the molar solubilities of each chemical listed above?

b. A beaker containing solutions of 0.100-M Fe(NO3)2 and 0.0175-M Cd(NO3)2 is titrated with sodium carbonate (Na2CO3). Which solid will precipitate first?

c. What percentage of the first ion to precipitate will remain when the second ion

begins to precipitate?

5. At 125°aC, Kp = 0.25 for the reaction:

2NaHCO3(s) =Na2CO3(s) + CO2(g) + H2O (g)

A 1.00-L flask containing 10.0-g of NaHCO3 is evacuated and heated to 125°C

a. Calculate the partial pressures of CO2 and H2O after equilibrium is established

b. Calculate the masses of NaHCO3 and Na2CO3 present at equilibrium

c. Calculate the minimum container volume necessary for all of the NaHCO3 to

decompose.

6. Ammonium hydrogen sulfide is a crystalline solid that decomposes as follows:

NH4HS(s) = NH3(g) + H2S(g)

a. Some solid NH4HS is placed in an evacuated vessel at 25°VC. After equilibrium is

attained, the total pressure inside the vessel is found to be 0.659 atmospheres.

Some solid NH4HS remains in the vessel at equilibrium. For this decomposition,

write the expression for KP and calculate its numerical value at 25°ëC.

b. Some extra NH3 gas is injected into the vessel containing the sample described in part (a). When equilibrium is reestablished at 25°C, the partial pressure of NH3 in the vessel is twice the partial pressure of H2S. Calculate the numerical value of the partial pressure of NH3 and the partial pressure of H2S in the vessel after the NH3 has been added and the equilibrium has been reestablished.

c. In a different experiment, NH3 gas and H2S gas are introduced into an empty

1.00-liter vessel at 25°äC. The initial partial pressure of each gas is 0.500 atmospheres. Calculate the number of moles of solid NH4HS that is present when equilibrium is established.

7. Three volatile compounds X, Y, and Z each contain element Q. The percent by weight of element Q in each compound was determined. Some of the data obtained are given below:

Compound % by mass of element Q in the compound

Molecular Weight

X 64.8% ?

Y 73.0% 104 g/mol

Z 59.3 % 64.0 g/mol

a. The vapor density of compound X at 27°C and 750. mm Hg was determined to be 3.53 g/L. Calculate the molecular weight of compound X.

b. Determine the mass of element Q contained in 1.00 mole of each of the three

compounds.

c. Calculate the most probable value of the atomic weight of element Q.

d. Compound Z contains carbon, hydrogen, and element Q. When 1.00 gram of

compound Z is oxidized and all of the carbon and hydrogen are converted to

oxides, 1.37 grams of CO2 and 0.281 gram of water are produced. Determine the

most probable molecular formula of compound Z.

8. A mixture of H2(g), O2(g), and 2.00 mL of H2O(l) is present in a 0.500 liter rigid container at 25°_C. The number of moles of H2 and the number of moles of O2 are equal. The total pressure is 1146 mm Hg. (The equilibrium vapor pressure of pure water at 25°C is 24 mmHg.)

The mixture is sparked, and H2 and O2 react until one reactant is completely consumed.

a. Identify the reactant remaining and calculate the number of moles of the reactant remaining.

b. Calculate the total pressure in the container at the conclusion of the reaction if the final temperature is 90°ñC. (The equilibrium vapor pressure of water at 90°C is 526mm Hg.)

c. Calculate the number of moles of water present as vapor in the container at 90°C.

9. Consider the following general equation for a chemical reaction.

A(g) + B(g) C(g) + D(g) ? Delta Hrxn= -10 kJ

a. Describe the two factors that determine whether a collision between molecules of A and B results in a reaction.

b. How would a decrease in temperature affect the rate of the reaction shown above? Explain your answer.

c. Write the rate law expression that would result if the reaction proceeded by the

mechanism shown below.

A + B =[AB] (fast)

[AB] + B =C + D (slow)

d. Explain why a catalyst increases the rate of a reaction but does not change the

value of K.

10. The following results were obtained when the reaction represented below was studied at 25°C.

2 ClO2(g) + F2(g) = 2 ClO2F(g)

Experiment

Initial

[ClO2],

(mol.L–1)

Initial

[F2],

(mol.L–1)

Initial Rate of

Increase of

[ClO2F],

(mol.L–1.sec–1)

1 0.010 0.10 2.4X10–3

2 0.010 0.40 9.6X10–3

3 0.020 0.20 9.6X10–3

a. Write the rate law expression for the reaction above.

b. Calculate the numerical value of the rate constant and specify the units.

c. In experiment 2, what is the initial rate of decrease of [F2]?

d. Which of the following reaction mechanisms is consistent with the rate law

developed in (a). Justify your choice.

I.

ClO2 + F2 = ClO2F2 (fast)

ClO2F2 = ClO2F + F (slow)

ClO2 + F = ClO2F

II.

F2 = 2 F

2 (ClO2 + F = ClO2F)

11. The reaction represented below is a reversible reaction.

BCl3(g) + NH3(g) = Cl3BNH3(s)

a. Predict the sign of the entropy change, Delta S, as the reaction proceeds to the right. Explain your prediction.

b. If the reaction spontaneously proceeds to the right, predict the sign of the enthalpy change, DeltaH. Explain your prediction.

c. The direction in which the reaction spontaneously proceeds changes as the

temperature is increased above a specific temperature. Explain.

d. What is the value of the equilibrium constant at the temperature referred to in (c) that is, the specific temperature at which the direction of the spontaneous reaction changes? Explain.

12.

A. Solid calcium carbonate is heated.

B. Solid calcium sulfite is heated in a vacuum.

C. Solid ammonium nitrate is heated to temperatures above 300C.

D. Solid sodium hydrogen carbonate is strongly heated.

E. A solution of hydrogen peroxide is exposed to sunlight.

F. Tetraphosphorus decoxide powder is sprinkled over distilled water.

G. Solid potassium oxide is added to water.

H. Solid calcium oxide is heated in the presence of sulfur trioxide gas.

I. Calcium metal is heated strongly in nitrogen gas.

J. The gases of boron trifluoride and ammonia are mixed.

K. Carbon dioxide gas is bubbled into magnesium oxide.

L. Excess concentrated sulfuric acid is added to solid calcium phosphate.

M. Hydrogen sulfide gas is bubbled into a solution of mercury (II) chloride.

N. Solutions of manganese (II) sulfate and ammonium sulfide are mixed.

O. An excess of sodium hydroxide solution is added to a solution of magnesium nitrate.

P. Solid lithium hydride is added to water.

Q. Equal volumes of 0.1 M sulfuric acid and 0.1 M potassium hydroxide are mixed.

R. Equal volumes of equimolar solutions of disodium hydrogen phosphate and hydrochloric acid are mixed.

S. Solutions of potassium phosphate and zinc nitrate are mixed.

T. Hydrogen sulfide gas is bubbled into a solution of nickel (II) nitrate.

U. Excess hydrochloric acid solution is added to a solution of potassium sulfite.

V. Dilute sulfuric acid is added to solid calcium fluoride.

W. Equal volumes of equimolar solutions of phosphoric acid and potassium hydroxide are mixed.

X. Solid zinc carbonate is added to 1.0 M sulfuric acid.

Y. Excess sodium cyanide solution is added to a solution of silver nitrate.

Z. Solid aluminum oxide is added to a solution of sodium hydroxide.

AA. Concentrated ammonia solution is added to a solution of zinc iodide.

BB. Excess concentrated potassium hydroxide solution is added to a precipitate of zinc hydroxide.

CC. A solution of ammonium thiocyanate is added to a solution of iron (III) chloride.

DD. A stream of chlorine gas is passed through a solution of cold, dilute sodium hydroxide.

EE. A solution of iron (TI) nitrate is exposed to air for an extended period of time.

FF. A concentrated solution of hydrochloric acid is added to solid potassium permanganate.

GG. A solution of potassium dichromate is added to an acidified solution of iron (II) chloride.

HH. A solution of potassium iodide is electrolyzed.

II. A concentrated solution of hydrochloric acid is added to powdered manganese dioxide and gently heated.

JJ. Solutions of potassium permanganate and sodium oxalate are mixed.

KK. A strip of copper is immersed in dilute nitric acid.

LL. Potassium permanganate solution is added to an acidic solution of hydrogen peroxide.

MM. A strip of silver is immersed in dilute nitric acid.

NN. A solution of iron (II) nitrate is added to a basic solution of hydrogen peroxide.

OO. Ethanol is burned in oxygen gas.

PP. Solid copper (II) sulfide is strongly heated in oxygen gas.

QQ. Propanol is burned in oxygen gas.

RR. Carbon disulfide vapor is burned in excess oxygen.

SS. Ethene gas is burned in air

TT. Butanol is burned in air.

13. (a) What is the pH of a 2.0-molar solution of acetic acid. Ka acetic acid = 1.8x10-5

(b) A buffer solution is prepared by adding 0.10-liter of 2.0-molar acetic acid solution to 0.1 liter of a 1.0-molar sodium hydroxide solution. Compute the hydrogen ion concentration of the buffer solution.

(c) Suppose that 0.10-liter of 0.50-molar hydrochloric acid is added to 0.040-liter of the buffer prepared in (b). Compute the hydrogen ion concentration of the resulting solution.

14.

A sample of 40.0-milliliters of a 0.100-molar HC2H3O2 solution is titrated with a 0.150- Molar NaOH solution. Ka for acetic acid = 1.8x10-5

(a) What volume of NaOH is used in the titration in order to reach the equivalence point?

(b) What is the molar concentration of C2H3O2- at the equivalence point?

(c) What is the pH of the solution at the equivalence point?

15. Predict whether solutions of each of the following salts are acidic, basic, or neutral. Explain your prediction in each case

(a) Al(NO3)3

(b) K2CO3

(c) NaBr

16.

A buffer solution contains 0.40-mole of formic acid, HCOOH, and 0.60-mole of sodium formate, HCOONa, in 1.00-litre of solution. The ionization constant, Ka, of formic acid is 1.8x10-4.

(a) Calculate the pH of this solution.

(b) If 100.-millilitres of this buffer solution is diluted to a volume of 1.00-litre with pure water, the pH does not change. Discuss why the pH remains constant on dilution.

(c) A 5.00-millilitre sample of 1.00-molar HCl is added to 100.-millilitres of the original buffer solution. Calculate the [H3O+] of the resulting solution.

(d) A 800.-milliliter sample of 2.00-molar formic acid is mixed with 200.-milliliters of 4.80 molar NaOH. Calculate the [H3O+] of the resulting solution.

(e) If you wanted to change the original solution so that it would buffer at a pH of 4.3, how many moles of HCl or NaOH must be added to the original solution? Assume no volume change.

17.

Sodium benzoate, C6H5COONa, is the salt of a weak acid, benzoic acid, C6H5COOH. A 0.10 molar solution of sodium benzoate has a pH of 8.60 at room temperature.

(a) Calculate the [OH-] in the sodium benzoate solution described above.

(b) Calculate the value for the equilibrium constant for the reaction:

C6H5COO- + H2O = C6H5COOH + OH-

(c) Calculate the value of Ka, the acid dissociation constant for benzoic acid.

(d) A saturated solution of benzoic acid is prepared by adding excess solid benzoic acid to pure water at room temperature. Since this saturated solution has a pH of 2.88, calculate the molar solubility of benzoic acid at room temperature.

18.

A 30.00-mL sample of a weak monoprotic acid was titrated with a standardized solution of NaOH. A pH meter was used to measure the pH after each increment of NaOH was added, and the curve above was constructed.

(a) Explain how this curve could be used to determine the molarity of the acid.

(b) Explain how this curve could be used to determine the dissociation constant Ka of the weak monoprotic acid.

(c) If you were to repeat the titration using a indicator in the acid to signal the endpoint, which of the following indicators should you select? Give the reason for your choice.

Methyl red Ka = 1x10-5

Cresol red Ka = 1x10-8

Alizarin yellow Ka = 1x10-11

(d) Sketch the titration curve that would result if the weak monoprotic acid were replaced by a strong monoprotic acid, such as HCl of the same molarity. Identify differences between this titration curve and the curve shown above.

19. Ammonium chloride is a crystalline solid that decomposes as follows:

NH4Cl(s) =NH3(g) + HCl (g)

a. Some solid NH4Cl is placed in an evacuated vessel at 25oC. After equilibrium is attained, the total pressure inside the vessel is found to be 0.762 atm. Some solid NH4Cl remains in the vessel at equilibrium. For this decomposition, write the expression for Kp and calculate its value at 25oC.

b. Some extra NH3 gas is injected into the vessel containing the sample described in part

(a). When equilibrium is reestablished at 25oC, the partial pressure of NH3 in the