1.
What is acid and
base by Arrhenius, Bronsted lowry and
Lewis?
2.
Write the chemical
equation for the autoionization of water and the equilibrium law for Kw?
3.
How are acidic, basic,
and neutral solutions in water defined
a. in
terms of [H+] and [OH-] and
b. in
terms of pH ?
4.
At the temperature of
the human body, 37oC, the value of Kw is 2.4 x 10-14.
Calculate the [H+], [OH-], pH and pOH of pure water at
this temperature. What is the relation between pH, pOH, and Kw at this
temperature? Is water neutral at this temperature?
5.
Deuterium oxide, D2O,
ionizes like water. At 20°C its Kw, or
ion product constant analogous to that of water, is 8.9 x 10-16.
Calculate [D+] and [OD-] in deuterium oxide at 20°C.
Calculate also the pD and the pDO.
6.
Calculate the H+
concentration in each of the following solutions in which the hydroxide ion
concentrations are :
a. 0.0024
M
b. 1.4
x 10-5 M
c. 5.6
x 10-9 M
d. 4.2
x 10-13 M
7.
Calculate the OH-
concentration in each of following
solutions in which the hydrogen ion concentrations are
a. 3.5
x 10 -8 M
b. 0.0065
M
c. 2.5
x 10 -13 M
d.
7.5 x 10 -5
M
8.
A certain brand of beer
had a hydrogen ion concentration equal to 1.9 x 10-5 mol L-1.What
is the pH of the beer?
9.
A soft drink was put on
the market with [
]
= 1,4 x 
mol 
. What it’s pH ?
]
= 1,4 x mol . What it’s pH ?
10. Calculate
the pH of each of the solutions in Exercises 5 and 6.
11. Calculate
the molar concentrations of H+ and OH- in solution that
have the following pH values.
a.
3.14
b.
2.78
c.
9.25
d.
13.24
e.
5.70
12. Calculate
the molar concentration of H+ and OH- in solution that
have the following pOH values .
a. 8.26
b. 10.25
c. 4.65
d. 6.18
e. 9.70
13.
What is the pH of
0.010 M HCl ?
14. What
is the pH of 0.0050 M solution of HNO3 ?
15. A
sodium hydroxide solution is prepared by dissolving 6.0 g NaOH in 1.00 L of
solution. What is the pOH and the pH of the solution?
16. A
solution was made by dissolving 0.837 g Ba(OH)2 in 100 mL final
volume. What is the pOH and the pH of the solution?
17. A
solution of Ca(OH)2 has a measured pH of 11.60. What is the molar
concentration of Ca(OH)2 in the solution?
18. A
solution of HCl has a pH of 2.50. How many grams of HCl are there in 250 mL of
this solution.
19. Write
the chemical equation for the ionization of each of the following weak acids in
water (For any polyprotic acids , write only the equation for the first step in
the ionization).
a. HNO2
b. H3PO4
c. HAsO42-
d. (CH3)3NH+
20. For
each of the acids in exercise 18, write the appropriate Ka
expression
21. Write
the chemical equation for the ionization of each of following weak bases in
water.
a. (CH3)3N
b. AsO43-
c. NO2-
d. (CH3)2N2H2
22. For
each of the bases in Exercise 20, write the appopriate Kb
expression.
23. Benzoic acid, C6H5CO2H,
is an organic acid whose sodium salt, C6H5CO2Na,
has long been used as a safe foods additive to protect beverages and many foods
againts harmful yeasts and bacteria. The
acid is monoprotic. Write the equation for it’s Ka !
24. Write
the equation for the equilibrium that the benzoate ion, C6H5CO2-
(review exercise 22), would
produce in water as functions as a Bronsted base. Then write the expression for
the Kb of the conjugate base of benzoic acid.
25.
The pKa of HCN is
9.21 and that of HF is 3.17. Which is the strong Bronsted base CN−
or F−?
26.
The
Ka for HF is 6.8 x 10x. what is the Kb for F-?
27. The
barbiturate ion C4HO has Kb = 1,0 x 10 -10 . What is Ka
for Barbituric acid ?
28. Hydrogen
peroxide, H2O2 is a week acid with Ka = 1.8 x 10-12. What the value of Kb
for the HO2 ion?
29. Methylamine,
CH3NH2 resambles ammonia in odor and basicity. Its Kb
is 4.4 x 10-4. Calculate the Ka of its conjugate acid!
30. Lactic
acid, HC3H5O3, is responsible for the sour
taste of sour milk. At 25oC its Ka = 1.4 x 10-4. What is
the Kb of its conjugate base, tha lactate ion, C3H5O3-
?
31. Iodic
acid, HIO3 has a pKa of 0.77
a. What is the formula an the Kb of
its conjugate base?
b. Its is conjugate base a stronger
or a weaker base than the acetate ion?
32. Periodic
acid,HIO4,is an important oxidizing agent and a moderately strong
acid. In a 0.10 M solution , [H+] = 3.8 x 10-2 mol L-1. Calculate the Ka and pKa for
periodic acid!
33. Choloacetic
acid, HC2H2ClO2, is a stronger monoprotic acid
than acetic acid. In a 0,10 M solution, this acid is 11 % ionized. Calculate
the Ka and pKa for Choloacetic acid.
34. Ethylamine, CH3CH2NH2,
has a strong, pungent odor similar to that ammonia.
Like ammonia, it is a Bronsted base. A 0.10 M solution has a pH of 11.86.
Calculate the Kb and pKb for ethylamine.
35. Hidroxylamine,
HONH2, like ammonia, is a Bronsted base. A 0.15 M solution has a pH
of 10.12. What are Kb and pKb for Hidroxylamine?
36. Refer
to data in the preceding question to calculate the percentage ionization of the
base in 0.15 M HONH2.
37.
What is the pH of
0.125 M pyruvic acid ? It’s Ka is 3.2 x 10-3
38. What
is pH of 0.15 M HN3 ? for HN3, Ka = 1.8 x 10-5
39. What
is the pH of a 1.0 M solution of hydrogen peroxide, H2O2?
For this solute, Ka =
1.8 x 10-2
40. Phenol,
also known as carbolic acid, is sometimes used as a disinfectant. What are the
concentrations of all of the substance in a 0.050 M solution of phenol, HC6H50? What percentage of the
phenol is ionized? For this acid, Ka= 1.3 x 10-10
41. Codeine,
a cough suppressant extracted from crude opium, is a weak base with a pKb of
5.79. What will be the pH of a 0.020 M solution of codeine? (Use Cod as a symbol for codeine)
42. Deuteroammonia,
ND3, is a weak base with a pKb of 4.96 at 25oC. What is
the pH of a 0.20 M solution of this compound?
43. A
solution of acetic acid has a pH of 2.54. What is the concentration of acetic
acid in this solution ?
44. Aspirin
is acetylsalicyclic acid, a monoprotic acid whose Ka value is 3,27 x
10-4. does a solution of the sodium salt of aspirin in water test
acidic, basic, or neutral ? Explain
45. The
Kb value of the oxalate ion, C2O42-,
is 1.9x10-10. Is a solution of K2C2O4
acidic, basic, or neutral? Explain.
46. Consider
the following compounds and suppose that 0.5M solutions are prepared of each :
NaI, KF, (NH4)2SO4, KCN, KC2H3O2,
CsNO3, and KBr. Write the formulas of those that have solutions that
are
a. Acidic,
b. Basic,
and
c. Neutral.
47. Will an aqueous solution
of ALCl3 turn litmus red or blue ? explain?
48. Explain
why the beryllium ion is a more acidic cation than the calcium ion.
49.
Ammonium nitrate is
commonly used in fertilizer mixtures as a source of nitrogen for plant growth.
What effect, if any, will this compound have on the acidity of the moisture in
the ground? Explain.
50.
Calculate the pH of
0.20 M NaCN.
51. Calculate
the pH of 0,04 M KNO2 ?
52. Calculate
the pH of 0.15 M CH3NH3Cl. For CH3NH2,
Kb = 4.4 x 10-4
53. A
weak base B forms the salt BHCl, composed of the ions BH+ and Cl-.
A 0.15 M solution of the salt has a pH of 4.28. What is the value of Kb
for the base B?
54. Calculate
the number of grams of NH4Br that have to be dissolved in 1.00 L of
water at 25oC to have a solution with a pH of 5.16 !
55. The
conjugate acid of a molecular base has a hypohetical formula. BH+,
and has pKa of 5.00. A solution of salt of this cation, BHY, tests slightly
basic. Will the conjugate acid of Y-, HY, have a pKa greater than
5.00 or less than 5.00? explain
56. Many
drugs that are natural Bronsted bases are put into aqueous solution as their
much more soluble salt with strong acids. The powerful painkiller morphine, for
example, is very slightly soluble in water, but morphine nitrate is quite
soluble. We may represent morphine by the symbol Mor and its conjugate acid as
H-Mor+. The pKb of morphine is 6.13. What is the calculated pH of a
0.20 M solution of H-Mor+?
57. Quinine,
an important drug in treating malaria, is a weak Bronsted base that we may
represent as Qu. To make it more soluble in water, it is put into a solution as
its conjugate acid, which we may represent as H-
. What is the calculate pHof a 0,15
M solution of H- ? Its pKa
is 8,52 at 25 0C.
. What is the calculate pHof a 0,15
M solution of H- ? Its pKa
is 8,52 at 25 0C.
58. Generally,
under what conditions are we unable to use the initial concentration of an acid
or base as though it were the equilibrium concentration in the mass action
expression?
59. What
is the percentage ionization in a 0.15 M solution of HF ? What is the pH of the
solution ?
60. What
is the percentage ionization in 0.0010 M acetic acid ? What is the pH of the
solution?
61.
What is the pH of a
1.0 x 10-7 M solution of HCl ?
62. The
hydrogen sulfate ion HSO4-, is a moderately strong
Bronsted acid with a Ka of 1.0x10-2.
a.
Write the chemical
equation for the ionization of the acid and give the appropriate Ka expression.
b.
What is the value of [
H+] in 0.010 M HSO4- (furnished by the salt,
NaHSO4) ? Do NOT make simplifying assumptions; solve the quadratic
equation.
c.
What is the calculate
of [H+] in 0.010 M HSO4-, obtained by using
the usual simplifying assumption?
d.
How much error is
produced by incorrectly using the simplifying assumption?
63. Para-Aminobenzoic
acid (PABA) is a powerful sunscreening agent whose salt were once used widely
in suntanning...... The parent acid, which we may symbolize as H-Paba, is a
weak acid with a pKa of 4.92 (.....oC). What is the [H+] and pH of
0.030 M solution of this acid?
64. Barbituric
acid, HC4H3N2O3 (which we will
abbreviate H-Bar), was discovered by the Nobel Prize-winning organic chemist
Adolph von Baeyer and named after his friend, Barbara. It is the parent
compound of widely sleeping drugs, the barbituretes. Its pKa is 4.01. what is
the [H+] and pH of a 0.050 M solution of H-Bar?
65. Write
ionic equation that illustrate how each pair of compounds can serve as a buffer
pair.
a. H2CO3
and NaHCO3 (the “carbonate” buffer in blood)
b. NaH2PO4
and Na2HPO4 (the “phosphate” buffer in side body cells)
c. NH4Cl
and NH3
66. Which
buffer would be better able to hold a steady pH on the addition of strong acid,
buffer 1 or buffer 2? Explain.
Buffer 1 is a solution containing 0.10 M NH4Cl
and 1 M NH3.
Buffer 2 is a solution containing 1 M NH4Cl
and 0.10 M NH3.
67. What
is the pH of a solution that contains 0.15 M HC2H3O2
and 0.25 M C2H3O2-?
Use Ka = 1.8 x 10-5 for HC2H3O2
Use Ka = 1.8 x 10-5 for HC2H3O2
68. Rework
the preceding problem using the Kb for the acetate ion. ( be sure to write the
poper chemical equation and equilibrium law )
69. By
how much will the pH change if 0.050 mol of HCl is added to 1.00 L off the
buffer in Exercise 66.
70. By
how much will the pH change if 50.0 mL of 0.10 M NaOH is added to 500mL of the
buffer in Exercise 66.
71. A
buffer is prepared containg 0.25 M NH3
and 0.14 M NH4+
a.
calculate the pH of the
buffer using the Kb for NH3
b.
calculate the pH of the
buffer using the Ka for NH4+
72. By
how much will the pH change if 0.020 mL of HCl is added to 1.00 L of the buffer
in Exercise 70?
73.
By how much will
the pH change if 75 ml of 0.10 M KOH is added to 200 ml of the buffer in
exercize 70?
74.
How many grams of
sodium acetat, NaC2H3O2, would have to be
added to 1.0 L of 0.15 M acetic acid (pKa
4.74) to make the solution a buffer for pH 5.00?
75. How
many grams of sodium formate, NaCHO2, would have to be added to 1.0
L of 0.12 M formic acid (Pka
3.74) to make the solution a buffer for pH 3.80 ?
76. What
mole ratio of NH4Cl to NH3 would buffer a solution at pH
9.25?
77. How
many grams of ammonium choride would have to be dissolved in 500 mL of 0.20 M NH3 to prepare a solution
buffered at pH 10.00?
78. How
many grams of ammonium chloride have to be dissolved into 125 mL of 0.10 M NH3
to make it a buffer with a pH of 9.15 ?
79. Suppose
25.00 mL of 0.100 M HCl is added to an
acetate buffer prepared by dissolving 0.100 mol of acetic acidand 0.110 of
sodium acetate in 500 mL of solution. What are the initial and final pH value?
what would be the pH if the same amount of HCl solution were added to 125 mL of
pure water?
80. How
many milliliters of 0.15 M HCl would have to be added to 100 mL of the buffer
described in exercise 78 to make the pH decrease by 0.05 pH unit? How many
milliliters of the same HCl solution would, if added to 100 mL of pure water,
make the pH decrease by 0.05 pH unit?
81. What
can make the titrated solution at the equivalence point in an acid-base
titration have a pH not equal to 7,00 ? Ho w does this possibility affect the
choice of an indicator ?
82. Explain
why ethyl red is a better indicator than phenolphtalein in the titration of
dilute ammonia by dilute hydrochloric acid?
83. What
is a good indicator for titrating potassium hydroxide with hydrobromic acid?
Explain.
84.
In the titration of an
acid with base,what condition concerning the quantities of reactans ought to be
true at the equivalence point?
85.
When 50 mL of 0.10
M formic acid is titrated with 0.10 M sodium hydroxide, what is the pH at the
equivalence point? (Be sure to take into account the change in volume during
the titration). What is a good indicator for this titration?
86. When
25 mL of 0.10 M aqueous
ammonia is titrated with 0.10 M hydrobromic acid, what is the pH at the
equivalence point? What is a good indicator?
87. For
the titratin of 25.00 mL of 0.1000 M HCl with 0.1000 M NaOH, calculate the pH
of the resulting solution after each of the following quantities of base has
been added to the original solution (you must take into account the change in
total volume). Construct a graph showing the titration curve for this
experiment.
a.
0 mL
b.
10.00 mL
c.
24.90 mL
d.
24.99 mL
e.
25.00 mL
f.
25.01 mL
g.
25.10 mL
h.
26.00 mL
i.
50.00 mL
88. For the titration of
25.00 mL of 0.1000 M acetic acid with 0.1000 M NaOH, calculate the pH:
a. Before
the addition of any NaOH solution,
b. After
10.00 mL of the base has been added,
c. After
half of the HC2H302 has been neutralized, and
d. At
the equivalence point.
89. For
the titration of 25.00 mL of 0.1000 M ammonia with 0.1000 M HCl, calculate the
pH
a. before
the addition of any HCl solution,
b. after
10.00 mL of the acid has been added,
c. after
half of the NH3 has been neutralized, and
d. at
the equivalence point
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