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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)
(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] + B =
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) =
2F(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 =
2F2 (fast)
ClO2F2 =
2F + F (slow)
ClO2 + F = ClO2F
II.
F2 =
2 (ClO2 + F =
2F)
11. The reaction represented below is a reversible reaction.
BCl3(g) + NH3(g) =
3BNH3(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.
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) =
3(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
vessel is three times the partial pressure of HCl. Calculate the numerical value of
the partial pressure of NH3 and the partial pressure of HCl in the vessel after the
NH3 has been added and equilibrium has been re-established.
c. In a different experiment, NH3 gas and HCl gas are introduced into an empty
1.50 L vessel at 25oC. The initial partial pressure of each gas is 0.700 atm.
Calculate the number of moles of solid NH4Cl that is present when equilibrium is
established.
20. Given the following reaction at equilibrium: N2O4(g) = 2NO2(g)
0.0400 mol N2O4(g) are placed in a previously evacuated, 1.00-L flask and
heated to 100oC. When equilibrium is established at 100oC, the equilibrium
concentration of N2O4(g) is found to be 0.0134 M.
a. Calculate the equilibrium concentration of NO2(g) .
b. Calculate the equilibrium constant, Kc, for the reaction at 100 oC.
c. In another experiment, equilibrium was approached from the other direction by
injecting a quantity of NO2(g) into a previously evacuated flask. After equilibrium
is established in this system at 100oC, the equilibrium concentration of NO2 is
found to be 0.0243 M. Find the equilibrium concentration of N2O4 in the system.
d. What was the initial concentration of NO2 that was injected in the experiment
described in part (c)?
21. The equilibrium constant, Kc, for the following reaction: N2(g) + O2(g)
=2NO(g) is 4.00 x 10-2 at a very high temperature. The reaction is at equilibrium
at this temperature with [N2] = [O2] = 0.100 Mand [NO] = 0.0200 M in a 2.00 liter
flask. If 0.120 mol of NO is suddenly added to the reaction mixture what will be
the concentrations of all species when equilibrium is re-established?
22. CO2(g) =
2(g)
A 1.00 mole sample of CO2 is placed in a 1.00 L container and allowed to come
to equilibrium at 2500K.When equilibrium is reached at 2500K, 17.6% of the
original CO2 has decomposed to CO and O2.
a. Calculate the values of equilibrium constants Kc and Kp for the dissociation
reaction at 2500K.
b. What are the partial pressures of each gas at equilibrium? What is the total
pressure?
23. Given the reaction: NH4HS (s) =
3(g) + H2S(g) delta H = +197.8 kJ.
Suppose the substances in the reaction are at equilibrium at 500 K. State
whether the partial pressure of NH3(g) will have increased, decreased, or
remained the same when equilibrium is re-established after each of the following
disturbances of the original system. Justify each answer with a brief explanation.
a. A small quantity of H2S is added.
b. The temperature of the system is increased.
c. The volume of the system is increased.
d. A quantity of N2 is added.
24. At 460oC, the reaction: SO2(g) + NO2(g) =
3(g)
has Kc = 85.0. What will be the equilibrium concentrations of the four gases if a
mixture of SO2 and NO2 is prepared in which they both have initial concentration
of 0.0750 M?
25. If 500.0 ml of 4.2 x 10-3 M Ce(NO3)3 is mixed with 800.0 ml of 5.6 x 10-3 M
NaIO3 , will the precipitate Ce(IO3)3 (Ksp = 1.9 x 10-10) form?
26. What mass of NaOH must be added to one liter of 0.010 M Mg(NO3)2 in order
to produce the first trace of Mg(OH)2 , Ksp = 1.5 x 10-11?
27. A solution is 0.10 M in Fe2+ and 0.10 M in Co2+ .
a. When H2S is added slowly, what precipitate first forms?
(Ksp of FeS = 1 x 10-17 and Ksp of CoS = 1 x 10-20)
b. What is the concentration of the first cation when the second cation starts to
precipitate?
28. 50.0 ml of 0.01 M Cd(NO3)2 is added to 50.0 ml of 0.1 M H2S.Will a
precipitate form?
(Ksp of CdS = 1 x 10-6 ).
29. In an experiment 1.056 g of a metal carbonate, containing and unknown
metal M, is heated to give the metal oxide and 0.376 g of CO2. What is the
identity of the metal
a. M = Ni
b. M = Cu
c. M = Zn
d. M = Ba
30. An unknown metal reacts with oxygen to give the neutral oxide, MO 2. Identify
the metal based on this information.
Mass of metal = 0.356 g
Mass of sample after converting metal completely to oxide = 0.452 g
31. Titanium (IV) oxide, TiO2, is heated in hydrogen gas to give water and a new
titanium oxide, TixOy. If 1.598 g of TiO2 produces 1.438 g of TixOy, what is the
formula of the new oxide?
32. Suppose you mix 25.0 mL of 0.234 M FeCl3 solution with 42.5 mL of 0.453 M
NaOH.
a. Which reactant is in excess?
b. What is the maximum mass, in grams, of Fe(OH)3 that precipitates?
c. What are the molar concentrations of all aqueous species at the end of the
reaction?
d. What is the % yield if 0.528 g are isolated experimentally?
33. You wish to determine the mass percent of copper in a copper-containing
alloy. After dissolving a sample of an alloy in acid, an excess of KI is added, and
the Cu2+ and I- ions undergo the reaction:
2Cu2+(aq) + 5I-(aq) =
(s) + I3 (aq)
The liberated I3- is titrated with sodium thiosulfate according to the equation:
2I3-(aq) + 2S2O32-(aq) =
4O6 (aq) + 3I (aq)
If 26.32 mL of 0.101 M Na2S2O3 is required for titration to the equivalence point,
what is the mass percent of Cu in 0.251 g of the alloy?
34. A 0.5895 g sample of impure magnesium hydroxide is dissolved in 100.0 mL
of 0.2050 M HCl solution. The excess acid then needed 19.85 mL of 0.1020M
NaOH for neutralization. Calculate the percent by mass of magnesium hydroxide
in the sample, assuming that it is the only substance reacting with the HCl
solution.
35. A compound of Ca, C, N, and S was subjected to quantitative analysis and
formula mass determination, and the following data were obtained. A 0.250 g
sample was mixed with Na2CO3 to convert all of the Ca to 0.160 g of CaCO3. A
0.115 g sample of the compound was carried through a series of reactions until
all of its S was changed to 0.344 g of BaSO4. A 0.712 g sample was processed
to liberate all of its N as NH3, and 0.155 g of NH3 was obtained. The formula
mass was found to be 156. Determine the empirical and molecular
formulas of the compound.
36. A mixture was prepared in a rigid 500.0 mL reaction vessel from 300.0 mL of
O2 (measured at 25C and 740 torr) and 400 mL of H2 (measured at 45C and
1250 torr). The mixture was ignited and the H2 and O2 reacted to form water.
a. What is the limiting reactant?
b. How many grams of liquid water will form?
c. What is the total pressure of the reaction vessel before the reaction occurs if
the internal tempera
d. What is the total pressure of the reaction vessel after the reaction is over, the
internal temperature is 65C?
37. Acetylene gas, C2H2 (g), can be prepared by the reaction of calcium carbide
with water:
CaC2 (s) + 2H2
2 (s) + C2H2 (g)
Calculate the volume of C2H2 that is collected over water at 21C by reaction of
3.26 g of CaC2 if the total pressure of the gas is 748 torr? (The vapor pressure of
water at 21C is 18.65 torr.)
38. Gaseous iodine pentafluoride can be prepared by the reaction of solid iodine
and gaseous fluorine: I2 (s) + 5F2 =
5 (g)
A 5.00-L flask containing 10.0 g I2 is charged with 10.0 g F2, and the reaction
proceeds until
on of the reactants is completely consumed. After the reaction is complete, the
temperature
in the flask is 125C.
a. What is the partial pressure of IF5 in the flask?
b. What is the mole fraction of IF5 in the flask?
39. Compound Z decomposes according to the following elementary process:
Z(g) =
) DeltaH= - 30 kCal
Addition of a catalyst causes the normally slow forward reaction to proceed
rapidly
a. Draw a diagram of potential energy versus reaction coordinate for the
uncatalyzed reaction. On this diagram label:
i. The axes.
ii. The energies of the reactants and the products.
iii. The energy of the activated complex.
iv. All significant energy differences.
b. On the same diagram indicate the change or changes that result from the
addition of the catalyst. Explain the role of the catalyst in changing the rate of the
reaction.
c. If the temperature is increased, will the ratio kf/kr increase, remain the same, or
decrease? Justify your answer with a one- or two-sentence explanation. (kf and kr
are the specific rate constants for the forward and the reverse reactions,
respectively.)
40. The decomposition of N2O5 to NO2 and O2 is a first-order reaction with the
rate constant, k = 1.2 x 10-3 s-1 at 374 K.
a. Determine the half-life of this reaction.
b. At 374 K, how much N2O5 would remain in a 1.0 L flask 5.0 min after the
introduction of 0.05 mol of N2O5?