Acids, Bases and Salts Unit Plan
Period/Topic
Worksheets
1. Arrhenius, Bronsted Acids, Ka and Strength.
1
2. Arrhenius, Bronsted Bases, Kb and Strength
2
3. Acid & Base Reactions. Amphiprotic. Acid Chart.
3
4. Leveling effect, Anhydrides and Relationships.
4
5. Hydrolysis of Salts. Quiz.
5
6. Acid, Base & Salt Reactions. Hydrolysis.
6
7. Yamada’s Indicator Lab. Hydrolysis.
7
8. Ionization of Water, [H+] & [OH-], pH scale.
8
9. pH Calculations
9
10. pH Calculations for Weak Acids.
10
11. Ka from pH for Weak Acids.
11
Quiz
1
2
3
4
5
12. Indicators Lab.
13. Kbs from Kas for Weak Bases.
12
14. pH for Weak Bases pH [H+] [OH-] Relationships.
13
15. Amphiprotic Ions- Kas and Kbs.
14
16. Titration Lab. Primary Standards.
6
7
Acids Midterm Practice Test
17. Titration Lab
18. Buffers & Indicators
15
8
19. Titration Curves.
16
9
20. Review # 1
Web Site Review
21. Review # 2
Practice Test 2
22. Test
Practice Test 1
Worksheet # 1
Properties of Acids and Bases
1.
List five properties of acids that are in your textbook.
2.
List five properties of bases that are in your textbook.
3.
Make some notes on the commercial acids: HCl and H2SO4 .
HCl
H2SO4
4.
Make some notes on the commercial base NaOH.
5.
Describe the difference between a concentrated and dilute acid (hint: concentration refers
to the molarity). Describe their relative conductivities.
6.
Describe the difference between a strong and weak acid. Use two examples and write
equations to support your answer. Describe their relative conductivities.
7.
Describe a situation where a strong acid would have the same conductivity as a weak acid
(hint: think about concentration)
Complete each acid reaction. Label each reactant and product as an acid or base. The first on is
done for you.
1.
HCN
Acid
+
H2O
Base
⇄
2.
H3C6O7
+
H2O
⇄
3.
H3PO4
+
H2O
⇄
4.
HF
+
H2O
⇄
5.
H2CO3
+
H2O
⇄
6.
NH4+ +
H2O
⇄
7.
CH3COOH +
H2O
⇄
8.
HCl
+
H2O
⇄
9.
HNO3 +
H2O
⇄
H3O+
Acid
+
CNBase
Write the equilibrium expression (Ka) for the first seven above reactions. The first one is done
for you.
10.
Ka =
11.
Ka =
[H3O+][CN-]
[HCN]
14. Ka =
15. Ka =
12.
Ka =
13.
Ka =
17.
Which acids are strong?
18.
What does the term strong acid mean?
19.
Why is it impossible to write an equilibrium expression for a strong acid?
20.
Which acids are weak?
23.
What does the term weak acid mean?
24.
Explain the difference between a strong and weak acid in terms of electrical conductivity.
Acid
16. Ka =
Conjugate Base
Base
Conjugate Acid
14.
16.
18.
HNO2
HSO3H2O2
15.
17.
19.
HCOOIO3NH3
20.
22.
HSH2O
21.
23.
CH3COOH2O
Define:
22.
Bronsted acid-
23.
Bronsted base-
24.
Arrhenius acid-
25.
Arrhenius base-
26.
List the six strong acids.
27.
Rank the acids in order of decreasing strength. HCl
HSO4
28.
What would you rather drink vinegar or hydrochloric acid? Explain.
H 2S
H3PO4
H2CO3
HF
Making a Universal Indicator Lab Activity
Mix the following indicators in a 50 mL beaker. Stir with an eyedropper.
Yamada’s Universal Indicator
5 drops thymol blue
6 drops methyl orange
5 drops phenolphthalein
10 drops bromothymol blue
20 drops of water
Part 1. In a spot plate add two drops of each buffer solution to a cell. Add one drop of Yamada’s
indicator to each. Record each colour on another lab sheet by colouring the cell the same colour.
Make sure you are accurate because you will use this information for future labs and projects.
<---------- Acid Strength Increases -----pH = 1
pH = 3
pH = 5
Neutral
pH = 7
----Base Strength Increases ------->
pH = 9
pH =11
pH = 13
Part 2. Test a drop of HCl, CH3COOH, NaOH, NH3, NaHCO3, H2CO3 and NaCl solution for
conductivity. Test with your Universal Indicator. Record the pH of each. Test with your
Universal Indicator. Explain your results with what you know about acids and bases. Classify
each as a strong or weak acid or base or neutral, acidic, or basic salt. Write an equation for each
to show how they ionize in water using the Bronsted (Chemistry 12) definition of an acid.
Wash and dry your chem plate
Wash and return your eyedropper.
Wash and return your beaker.
Wash your hands.
Compound
Conductivity
pH
Classification
HCl
CH3COOH
NaOH
NH3
NaHCO3
H2CO3
NaCl
Worksheet # 2
Conjugate Acid-Base Pairs
Complete each reaction. Label each reactant and product as an acid or base.
1.
HCN
+
Acid
H2 O
⇄
Base
HCl
+
H2O
⇄
3.
HF
+
H2O
⇄
4.
F-
+
H2O
⇄
5.
HSO4(acid)
+
H2O
⇄
6.
NH4+
+
H2 O
⇄
7.
HPO42(base)
+
H2 O
⇄
Acid
Conjugate Base
8.
10.
HCO3HPO42-
12.
14.
H2O
HS-
16.
Circle the strong bases.
17.
CO32-
NaOH
Sr(OH)2
+
Acid
2.
Fe(OH)3
Zn(OH)2
H3 O+
Base
9.
11.
CH3COOIO3-
13.
15.
NH2C2H5SO73-
CsOH
Ba(OH)2
Rank the following acids from strongest to weakest.
CNBase
Conjugate Acid
CH3COOH
KOH
Ca(OH)2
H2S
18.
H2PO4-
HI
HCl
HF
Rank the following bases from the strongest to weakest.
H2O
19.
CH3COOH
F-
NH3
SO32-
HSO3-
NaOH
i) Write the reaction of H3BO3 with water (remove one H+ only because it is a weak
acid).
ii) Write the Ka expression for the above.
iii) What is the ionization constant for the acid (use your table). Ka =
20.
List six strong acids.
21.
List six strong bases.
22.
List six weak acids in order of decreasing strength (use your acid/base table).
23.
List six weak bases in order of decreasing strength (use your acid/base table).
Worksheet # 3
Using Acid Strength Tables
Acid-base reactions can be considered to be a competition for protons. A stronger acid
can cause a weaker acid to act like a base. Label the acids and bases. Complete the
reaction. State if the reactants or products are favoured.
1.
HSO4-
+
HPO42-
⇄
2.
HCN
+
H2O
⇄
3.
HCO3-
+
H2S
⇄
4.
HPO42-
+
NH4+
⇄
5.
NH3
+
H2O
⇄
6.
H2PO41-
+
NH3
⇄
7.
HCO3-
+
HF
⇄
8.
Complete each equation and indicate if reactants or products are favoured. Label each
acid or base.
HSO4-
+
HCO3-
⇄
H2PO4-
+
HC03-
⇄
HS03-
+
HPO42-
⇄
NH3
+
HC2O4-
⇄
9.
Explain why HF(aq) is a better conductor than HCN(aq).
10.
Which is a stronger acid in water, HCl or HI? Explain!
11.
State the important ion produced by an acid and a base.
12.
Which is the stronger base? Which produces the least OH-? F- or CO3-2
13.
Define a Bronsted/Lowry acid and base.
14.
Define an Arrhenius acid and base.
15.
Complete each reaction and write the equilibrium expression.
HF + H2O
-
F + H2O
⇄
Ka=
⇄
Kb=
→
16.
H2SO4 + NaOH
17.
Define conjugate pairs.
18.
Give conjugate acids for: HS-, NH3, HPO4-2, OH-, H2O, NH3,
19.
Give conjugate bases for: NH4+,
Worksheet # 4
HF,
H2PO4-,
H3O+, OH-,
CO3-2
HCO3-,
H2O
Acid and Basic Anhydrides
1.
What is the strongest acid that can exist in water? Write an equation to show how a
stronger acid would be reduced in strength by the leveling effect of water.
2.
What is the strongest base that can exist in water? Write an equation to show how a
stronger base would be reduced in strength by the leveling effect of water.
3.
List three strong acids and three strong bases.
4.
Rank the acids in decreasing strength:
HClO4
HClO2
5.
Ka is very large
Ka=8.0x10-5
HClO3
HClO
Ka=1.2x10-2
Ka=4.4x10-8
For an oxy acid what is the relationship between the number of O’s and
acid strength? (Compare H2S04 and H2S03)
6.
Which acid is stronger?
HI03 or HIO2
7.
8.
Which produces more H30+?
Which produces more OH-?
H2CO3 or HS04F- or HC03-
9.
Which conducts better NH3 or NaOH (both .1M)? Why?
10.
Which conducts better HF or HCN (both .1M)? Why?
11.
Compare and contrast a strong and weak acid in terms of degree of ionization, size of ka,
conductivity, and concentration of H+.
Classify each formula as an acid anhydride, basic anhydride, strong acid, weak acid, strong, or
weak base. For each formula write an equation to show how it reacts with water. For anhydrides
write two equations.
Formula
12.
Na2O
13.
CaO
14.
SO3
15.
CO2
16.
SO2
17.
HCl
Classification
Reaction
18.
NH3
19.
NaOH
20.
HF
21.
H3PO4
Worksheet # 5
Hydrolysis of Salts and Reactions of Acids and Bases
Describe each as an acid, base, neutral salt, acidic salt, or basic salt. For each salt write a parent
acid-base formation equation, dissociation equation, and hydrolysis equation (only for acidic and
basic salts). For acids and bases write an equation to show how each reacts with water.
1.
NH3
2.
KCl
3.
HNO3
4.
NaHCO3
5.
RbOH
6.
AlCl3
7.
H2C2O4
8.
NaC6H5O
9.
Co(NO3)3
10.
Na2CO3
Worksheet # 6
Hydrolysis of Salts and Reactions of Acids and Bases
Describe each as an acid, base, neutral salt, acidic salt, or basic salt. For each salt write a
dissociation equation and hydrolysis equation (only for acidic and basic salts). For acids and
bases write an equation to show how each reacts with water.
1.
NH3
2.
NaCl
3.
HCl
4.
NaCN
5.
NaOH
6.
FeCl3
7.
HF
8.
LiHCO3
9.
Fe(NO3)3
10.
MgCO3
11.
H2S
12.
HF
13.
CaI2
14.
Mg(OH)2
15.
Ba(OH)2
16.
Describe why Tums (CaCO3) neutralizes stomach acid.
17.
Describe why Mg(OH)2 is used in Milk of Magnesia as an antacid instead of NaOH.
Worksheet # 7
Yamada’s Indicator Activity
Acid, Base and Salt Lab
Purpose:
1)
2)
3)
To use Yamada’s Indicator to determine the pH of various acids, bases and salts.
To classify compounds as strong acids, weak acids, strong bases, weak bases, neutral
salts, acid anhydrides, and basic anhydrides.
To write reactions for each compound to show how each ionizes, hydrolyzes or reacts
with water.
Procedure:
1)
2)
To a cell in a spot plate add one drop of solution or a very tiny amount of solid. Write the
formula of the compound in the data table.
Add two drops of Yamada’s Indicator. Record the pH of the compound.
3)
4)
1.
Classify the compound as a strong acid, weak acid, strong base,
weak base, neutral
salt, acid anhydride, or basic anhydride. Use the formula of the compound as well as the
pH.
Write an equation to show the reaction of anhydrides with water, the hydrolysis of salts,
or the ionization of acids or bases.
Formula of compound
pH
Classification
Reaction or reactions
2.
Formula of compound
pH
Classification
Reaction or reactions
3.
Formula of compound
pH
Classification
Reaction or reactions
4.
Formula of compound
pH
Classification
Reaction or reactions
5.
Formula of compound
pH
Classification
Reaction or reactions
6.
Formula of compound
pH
Classification
Reaction or reactions
7.
Formula of compound
pH
Classification
Reaction or reactions
8.
Formula of compound
pH
Classification
Reaction or reactions
9.
Formula of compound
pH
Classification
Reaction or reactions
10.
Formula of compound
pH
Classification
Reaction or reactions
11.
Formula of compound
pH
Classification
Reaction or reactions
12.
Formula of compound
pH
Classification
Reaction or reactions
13.
Formula of compound
pH
Classification
Reaction or reactions
14.
Formula of compound
pH
Classification
Reaction or reactions
15.
Formula of compound
pH
Classification
Reaction or reactions
16.
Formula of compound
pH
Classification
Reaction or reactions
17.
Formula of compound
pH
Classification
Reaction or reactions
Worksheet # 8
pH and pOH Calculations
Complete the chart:
1.
2.
3.
4.
5.
6.
7.
8.
9.
[H+]
7.00 x 10-3 M
[OH-]
pH
pOH
Acid/base/neutral
8.75 x 10-2 M
7.33
4.00
Neutral (2 sig
figs)
10.7
2.553
5.0 x 10-10 M
4.7 x 10-10 M
10.
Calculate the [H+], [OH-], pH and pOH for a 0.20 M Ba(OH)2 solution.
11.
Calculate the [H+], [OH-], pH and pOH for a 0.030 M HCl solution.
12.
Calculate the [H+], [OH-], pH and pOH for a 0.20 M NaOH solution.
13.
300.0 mL of 0.20 M HCl is added to 500.0 mL of water, calculate the pH of the
solution.
14.
200.0 mL of 0.020 M HCl is diluted to a final volume of 500.0 mL with water, calculate
the pH.
15.
150.0 mL of 0.40 M Ba(OH)2 is placed in a 500.0 mL volumetric flask and filled to the
mark with water, calculate the pH of the solution.
16.
250.0 mL of 0.20M Sr(OH)2 is diluted by adding 350.0 mL of water, calculate the pH of
the solution.
17.
Calculate the pH of a 0.40 solution of Ba(OH)2 when 25.0 mL is added to 25.0 mL of
water.
Worksheet # 9
1.
Calculate the [H+], [OH-], pH, and pOH for 0.20 M HCN.
[H+] =
2.
pH Calculations for Weak Acids
[OH-] =
pH =
Calculate the [H+], [OH-], pH, and pOH for 2.20 M HF.
pOH =
3.
[H+] =
[OH-] =
pH =
pOH =
Calculate the [H+], [OH-], pH, and pOH for 0.805 M CH3COOH.
[H+] =
4.
[OH-] =
pH =
pOH =
Calculate the [H+], [OH-], pH, and pOH for 1.65 M H3BO3.
[H+] =
[OH-] =
pH =
pOH =
5.
Calculate the pH of a saturated solution of Mg(OH)2.
6.
Calculate the pH of a 0.200 M weak diprotic acid with a Ka = 1.8 x 10-6.
7.
350.0 mL of 0.20M Sr(OH)2 is diluted by adding 450.0 mL of water, calculate the pH of
the solution.
Worksheet # 10
pH Calculations for Weak Acids
1.
The pH of 0.20 M HCN is 5.00. Calculate the Ka for HCN. Compare your calculated
value with that in the table.
2.
The pH of 2.20 M HF is 1.56. Calculate the Ka for HF. Compare your calculated value
with that in the table.
3.
The pH of 0.805 M CH3COOH is 2.42. Calculate the Ka for CH3COOH. Compare your
calculated value with that in the table.
4.
The pH of 1.65 M H3BO3 is 4.46. Calculate the Ka for H3BO3. Compare your calculated
value with that in the table.
5.
The pH of a 0.10 M diprotic acid is 3.683, calculate the Ka and identify the acid.
6.
The pH of 0.20 M NH3 is 11.227; calculate the Kb of the Base.
7.
The pH of 0.40 M NaCN is 11.456; calculate the Kb for the basic salt. Start by writing an
equation and an ICE chart.
8.
The pH of a 0.10 M triprotic acid is 5.068, calculate the Ka and identify the acid.
9.
How many grams of CH3COOH are dissolved in 2.00 L of a solution with pH = 2.45?
Use questions 1 to 4 from last assignment to mark questions 1 to 4.
Worksheet # 11
Kb For Weak Bases
Determine the Kb for each weak base. Write the ionization reaction for each. Remember that Kw =
Ka • Kb (the acid and base must be conjugates). Find the base on the right side of the acid table
and use the Ka values that correspond. Be careful with amphiprotic anions! The first one is done
for you.
1.
NaNO2 (the basic ion is NO2-)
Kb(NO2-)
=
Kw
=
Ka(HNO2)
1.0 x 10-14
4.6 x 10-4
2.
KCH3COO (the basic ion is CH3COO-)
3.
NaHCO3
4.
NH3
=
2.2 x 10-11
5.
6.
NaCN
Li2HPO4
7.
KH2PO4
8.
K2CO3
9.
Calculate the [H+], [OH-], pH, and pOH for 0.20 M H2CO3.
[H+] =
10.
[OH-] =
pH =
pOH =
The pH of 0.20 M H2CO3 is 3.53. Calculate the Ka for H2CO3. Compare your calculated
value with that in the table.
11.
Calculate the [H+], [OH-], pH, and pOH for 0.10 M CH3COOH.
[H+] =
[OH-] =
pH =
pOH =
12.
The pH of 0.10 M CH3COOH is 2.87. Calculate the Ka for CH3COOH. Compare your
calculated value with that in the table.
13.
200.0 mL of 0.120 M H2SO4 reacts with 400.0 mL of 0.140 M NaOH. Calculate the pH
of the resulting solution.
Worksheet # 12
Acid and Base pH Calculations
For each weak bases calculate the [OH-], [H+], pOH and pH. Remember that you need to
calculate Kb first.
1.
0.20 M CN-
2. 0.010 M NaHS (the basic ion is HS-)
3.
0.067 M KCH3COO
4.
0.40 M KHCO3
5.
0.60 M NH3
6.
If the pH of a 0.10 M weak acid H2X is 3.683, calculate the Ka for the acid and identify
the acid using your acid chart.
7.
Calculate the [H+], [OH-], pH, and pOH for 0.80 M H3BO3.
[H+] =
8.
pH =
pOH =
Calculate the [H+], [OH-], pH, and pOH for 0.25 M H2CO3.
[H+] =
9.
[OH-] =
[OH-] =
pH =
pOH =
The pH of 1.65 M H3BO3 is 4.46. Calculate the Ka for H3BO3. Compare your calculated
value with that in the table.
10.
The pH of 0.65 M NaX is 12.46. Calculate the Kb for NaX.
11.
Consider the following reaction: 2HCl + Ba(OH)2 → BaCl2 + 2H2O
When 3.16g samples of Ba(OH)2 were titrated to the equivalence point with an HCl
solution, the following data was recorded.
Trial
#1
#2
#3
Volume of HCl added
37.80 mL
35.49 mL
35.51 mL
Calculate the original [HCl]
12.
Calculate the volume of 0.200M H2SO4 required to neutralize 25.0 ml of 0.100M NaOH.
13.
25.0 mL of .200 M HCl is mixed with 50.0 mL 0.100 M NaOH, calculate the pH of he
resulting solution.
14.
15.
10.0 mL of 0.200 M H2SO4 is mixed with 25.0 mL 0.200 M NaOH, calculate the pH of
the resulting solution.
125.0 mL of .200 M HCl is mixed with 350.0 mL 0.100 M NaOH, calculate the pH of the
resulting solution.
16.
Define standard solution and describe two ways to standardize a solution.
17.
What is the [H3O+] in a solution formed by adding 60.0 mL of water to 40.0 mL of 0.040
M KOH solution?
Worksheet # 13
Amphiprotic Ions and Review
1.
List the properties of acids/bases.
2.
Define the following:
Arhenius strong acid
Arhenius weak base
Bronsted strong acid
Bronsted weak base
Conjugate pair
Amphiprotic
3.
Standard solution.
Show by calculation if the following amphiprotic ions are acids or bases:
HCO3H2PO4HPO42-
4.
What is the strongest base in water? What is the strongest acid in water? Write equations
to explain your answer.
5.
Match each equation:
Acid/base complete
HCl + NaOH → NaCl + HOH
Acid/base net ionic
Solubility product
F- + HOH ⇄ HF + OHH+ + OH- → HOH
Hydrolysis
AgCl(s) ⇄ Ag+ + Cl-
Acid/Base formula
Ionization of water
H20 ⇄ H+ + OHH+ + Cl- + Na+ + OH- → Na++ Cl- + H2O
6.
HCl and HF. Describe each acid as:
a)
b)
c)
d)
e)
strong or weak
high or low ionization
large or small Ka
good or poor conductor
strong or weak electrolyte
7.
Out of 0.2 M HCl and 1.0 M HF, which is the most concentrated?
Which is the strongest acid?
8.
Label the scale as strong/weak acid and strong/weak base.
pH
|________________________|_________________________|__
0
7
14
9.
Which ions are amphiprotic?
HPO42HCl
FHS-
H2S
H2 O
10.
Write the net ionic equation between any strong acid and strong base.
11.
Write the ionization equation for water.
12.
Write the Kw expression.
13.
H2SO3 + HS- ⇄
H2S + HSO3-
a)
Are the reactants or products favoured?
b)
Is the Keq large, small or about 1?
Determine the pH Write equations for each first!
14.
.20M HCl
pH=?
15.
0.20M Ba(OH)2
pH=?
16.
0.20M H2CO3
pH=?
17.
0.40M KHCO3
pH=?
18.
The pH increases by 2 units. How does [H+] change?
19.
The pH decreases by 1 unit.
How does [H+] change?
20.
a) For distilled water : pH=
pOH=
[H+]=
[OH-]=
b) For 1M HCl:
pH=
pOH=
[H+]=
[OH-]=
c) For 1M NaOH:
pH=
pOH=
[H+]=
[OH-]=
21.
The pH of 0.20 M NaX is 12.50; calculate the Kb.
22.
The pH of 0.2 M HX is 4.5; calculate the Ka.
23.
100.0 mL of 0.200 M NaOH is mixed with 100.0 mL of 0.180 M HCl. Calculate the pH
of the resulting solution.
24.
How many grams of NaHCO3 are required to make 100.0 mL of 0.200M solution?
25.
What volume of 0.200 M NaOH is required to neutralize 25.0 mL of 0.150 M H2SO4?
26.
In a titration 25.0 mL of .200M H2SO4 is required to neutralize 10.0 mL NaOH. Calculate
the concentration of the base.
27.
Calculate the concentration of a solution of NaCl made by dissolving 50.0 g in 250.0 mL
of water.
28.
SO3(g)
+ H2O(g)
⇄
H2SO4(l)
Equilibrium concentrations are found to be:
[SO3] = 0.400 M
[ H2O] = 0.480 M
Calculate the value of the equilibrium constant.
29.
[H2SO4] = 0.600 M
4.00 moles of SO2 and 5.00 moles O2 are placed in a 2.00 L container at 200º C and
allowed to reach equilibrium. If the equilibrium concentration of O2 is
2.00 M, calculate the Keq.
2SO2(g) +
O2(g)
⇄
2SO3(g)
Worksheet # 14
Buffers and indicators
Buffers
1.
Definition
2.
Acid
Conjugate Base
Salt
HCN
KHCO3
NH3
HF
NaCH3COO
HC2O4-
3.
Write an equation for the first three buffer systems above.
4.
Which buffer could have a pH of 4.0?
a) HCl & NaCl
b) HF & NaF
Which buffer could have a pH of 10.0?
c)
NH3
& NH4Cl
5.
Predict how the buffer of pH = 9.00 will change. Your possible answers are 9.00, 8.98,
9.01, 2.00, and 13.00
Final pH
a)
2 drops of 0.10 M HCl are added
b)
1 drop of 0.10 M NaOH is added
c)
10 mL of 0.10 M HCl are added
6.
Write an equation for the carbonic acid, sodium hydrogencarbonate buffer system. A few
drops of HCl are added. Describe the shift and each concentration change.
Equation:
Shift
[H+] =
[HCO3-] =
[H2CO3] =
Indicators
1.
Definition
2.
Equilibrium equation
3.
Colors for methyl orange
4.
Compare the relative sizes of [HInd] and [Ind-] at the following pH’s for methyl orange.
Color
HInd
Ind-
Relationship
pH = 2.0
pH = 3.7
pH = 5.0
5.
HCl is added to methyl orange, describe if each increases or decreases.
[H+]
[HInd]
[Ind-]
Color Change
6.
NaOH is added to methyl orange, describe if each increases or decreases.
[H+]
[HInd]
[Ind-]
Color Change
7.
State two equations that are true at the transition point of an indicator.
8.
What indicator has a Ka = 4 x 10-8
9.
What is the Ka for methyl orange?
10.
A solution is pink in phenolphthalein and colorless in thymolphthalein. What is the pH of
the solution?
11.
A solution is blue in bromothymol blue, red in phenol red, and yellow in thymol blue.
What is the pH of the solution?
Worksheet # 15
Titration Curves
Choose an indicator and describe the approximate pH of the equivalence point for each titration.
Complete each reaction.
pH
+
NaOH
→
RbOH
→
1.
HCl
2.
HF
+
3.
HI
+ Ba(OH)2
4.
HCN
5.
HClO4
6.
CH3COOH
+
+
→
KOH
→
NH3
→
+
→
LiOH
Indicator
7.
Calculate the Ka of bromothymol blue.
8.
9.
An indicator has a Ka = 1 x 10-10, determine the indicator.
Calculate the Ka of methyl orange.
10.
An indicator has a Ka = 6.3 x 10-13, determine the indicator.
11.
Explain the difference between an equivalence point and a transition point.
Draw a titration curve for each of the following.
12.
Adding 100 ml 1.0 M NaOH to
50 mL 1.0 M HCl
13.
Adding 100 ml 1.0 M NaOH to
50 mL 1.0 M HCN
pH
Volume of base added
Volume of base added
14.
Adding 100 ml 0.10 M HCl
to 50 mL 0.10M NH3
pH
15.
Adding 100 ml .10 M HCl
to 50 mL 0.10 M NaOH
pH
Volume of HCl added
Volume of HCl added
Acids Unit Midterm Practice Test
1.
Consider the following:
I
H2CO3 + F-  HCO3- + HF
II
HCO3- + HC2O4-  H2CO3 + C2O42III
HCO3- + H2C6H6O7-  H2CO3 + HC6H5O72The HCO3- is a base in
A.
B.
C.
D.
2.
I only
I and II only
II and III only
I, II, and III
Consider the following equilibrium for an indicator:
HInd + H2O  Ind- + H3O+
When a few drops of indicator methyl red are added to 1.0 M HCl, the colour of the resulting solution is
A.
B.
C.
red and the products are favoured
red and the reactants are favoured
yellow and the products are favoured
3.
D.
yellow and the reactants are favoured
The volume of 0.200 M Sr(OH)2 needed to neutralize 50.0 mL of 0.200 M HI is
A.
B.
C.
D.
10.0 mL
25.0 mL
50.0 mL
100.0 mL
4.
The pOH of 0.050 M HCl is
A.
0.050
B.
1.30
C.
12.70
D.
13.70
5.
The volume of 0.450 M HCl needed to neutralize 40.0 mL of 0.450 M Sr(OH) 2 is
A.
18.0 mL
B.
20.0 mL
C.
40.0 mL
D.
80.0 mL
Consider the following
I
H3PO4
II
H2PO4III
HPO42IV
Which of the above solutions are amphiprotic?
A.
I and II only
B.
II and III only
C.
I, II, and III only
D.
II, III, and IV only
6.
PO43-
7.
Which of the following solutions will have the largest [H3O+]?
A.
1.0 M HNO2
B.
1.0 M H3BO3
C.
1.0 M H2C2O4
D.
1.0 M HCOOH
8.
Consider the following: H2O + 57 kJ  H3O+ + OHWhen the temperature of the system is increased, the equilibrium shifts
A.
left and the Kw increases
B.
left and the Kw decreases
C.
right and the Kw increases
D.
right and the Kw decreases
9.
Normal rainwater has a pH of approximately 6 as a result of dissolved
A.
oxygen
B.
carbon dioxide
C.
sulphur dioxide
D.
nitrogen dioxide
10.
A 1.0 M solution of sodium dihydrogen phosphate is
A.
acidic and the pH < 7.00
B.
acidic and the pH > 7.00
C.
basic and the pH < 7.00
D.
basic and the pH > 7.00
11.
Consider the following equilibrium for an indicator:
HInd + H2O  Ind- + H3O+
When a few drops of indicator chlorophenol red are added to a colourless solution of pH 4.0, the resulting
solution is
A.
B.
C.
D.
red as [HInd] < [Ind-]
red as [HInd] > [Ind-]
yellow as [HInd] < [Ind-]
yellow as [HInd] > [Ind-]
12.
A Bronsted-Lowry base is defined as a chemical species that
A.
accepts protons
B.
neutralizes acids
C.
donated electrons
D.
produces hydroxides ions in solution
13.
Which of the following solutions will have the greatest electrical conductivity?
A.
1.0 M HCN
B.
1.0 M H2SO4
C.
1.0 M H3PO4
D.
1.0 M CH3COOH
14.
Consider the following equilibrium: HC6H5O72- + HIO3  H2C6H5O7- + IO3The order of Bronsted-Lowry acids and bases is
A.
acid, base, acid, base
B.
acid, base, base, acid
C.
base, acid, acid, base
D.
base, acid, base acid
15.
Consider the following: H2O(l)  H+ + OHWhen a small amount of 1.0 M KOH is added to the above system, the equilibrium
A.
shifts left and [H+] decreases
B.
shifts left and [H+] increases
C.
shifts right and [H+] decreases
D.
shifts right and [H+] increases
16.
Which of the following has the highest pH?
A.
1.0 M NaIO3
B.
1.0 M Na2CO3
C.
1.0 M Na3PO4
D.
1.0 M Na2SO4
17.
In a 100.0 mL sample of 0.0800 M NaOH the [H3O+] is
A.
1.25 x 10-13 M
B.
1.25 x 10-12 M
C.
8.00 x 10-3 M
D.
8.00 x 10-2 M
18.
Consider the following:
I
ammonium nitrate
II
calcium nitrate
III
When dissolved in water, which of these salts would form a neutral solution?
A.
B.
C.
D.
19.
II only
III only
I and III only
I, II, and III
Consider the following: SO42- + HNO2  HSO4- + NO2-
iron III nitrate
Equilibrium would favour the
A.
the products since HSO4- is a weaker acid than HNO2
B.
the reactants since HSO4- is a weaker acid than HNO2
C.
the products since HSO4- is a stronger acid than HNO2
D.
the reactants since HSO4- is a stronger acid than HNO2
20.
The net ionic equation for the hydrolysis of Na2CO3 is
A.
H2O + Na+  NaOH + H+
B.
H2O + 2Na+  Na2O + 2H+
C.
H2O + CO32-  H2CO3 + O2D.
H2O + CO32-  HCO3- + OH-
21.
Consider the following equilibrium: 2H2O(l)  H3O+ + OHA few drops of 1.0 M HCl are added to the above system. When equilibrium is
re-established, the
A.
[H3O+] has increased and the [OH-] has decreased
B.
[H3O+] has increased and the [OH-] has increased
C.
[H3O+] has decreased and the [OH-] has increased
D.
[H3O+] has decreased and the [OH-] has decreased
22.
A basic solution
A.
tastes sour
B.
feels slippery
C.
does not conduct electricity
D.
reacts with metals to release oxygen gas
23.
The balanced formula equation for the neutralization of H2SO4 by KOH is
A.
H2SO4 + KOH → KSO4 + H2O
B.
H2SO4 + KOH → K2SO4 + H2O
C.
H2SO4 + 2KOH → K2SO4 + H2O
D.
H2SO4 + 2KOH → K2SO4 + 2H2O
24.
An Arrhenius base is defined as a substance which
A.
donates protons
B.
donates electrons
C.
produces H+ in solution
D.
produces OH- in solution
25.
HS- + H3PO4  H2S + H2PO4- The order of Bronsted-Lowry acids and bases is
A.
acid, base, acid, base.
B.
acid, base, base, acid
C.
base, acid, acid, base
D.
base, acid, base, acid
26.
The equation representing the reaction of ethanoic acid with water is
A.
CH3COO- + H2O  CH3COOH + OHB.
CH3COO- + H2O  CH3COO2- + H3O+
C.
CH3COOH + H2O  CH3COO- + H3O+
D.
CH3COOH + H2O  CH3COOH2+ + OH-
27.
Consider the following equilibrium: 2H2O + 57kJ  H3O+ + OHWhen the temperature is decreased, the water
A.
stays neutral and the [H3O+] increases
B.
stays neutral and the [H3O+] decreases
C.
D.
28.
29.
becomes basic and [H3O+] decreases
becomes acidic and [H3O+] increases
The equation for the reaction of Cl2O with water is
A.
Cl2O + H2O  2HClO
B.
Cl2O + H2O  2ClO + H2
C.
Cl2O + H2O  Cl2 + H2O2
D.
Cl2O + H2O  Cl2 + O2 + H2
The conjugate acid of C6H50- is
A.
C6H4OB.
C6H5OH
C.
C6H4O2D.
C6H5OH+
30.
Which of the following solutions will have the greatest electrical conductivity?
A.
1.0 M HCl
B.
1.0 M HNO2
C.
1.0 M H3BO3
D.
1.0 M HCOOH
31.
A solution of 1.0 M HF has
A.
a lower pH than a solution of 1.0 M HCl
B.
a higher pOH than a solution of 1.0 M HCl
C.
a higher [OH-] than a solution of 1.0 M HCl
D.
a higher [H3O+] than a solution of 1.0 M HCl
32.
Which of the following is the weakest acid
A.
HIO3
B.
HCN
C.
HNO3
D.
C6H5COOH
33.
Considering the following data
H3AsO4 Ka = 5.0 x 10-5
H2AsO4- Ka = 8.0 x 10-8
HAsO42- Ka = 6.0 x 10-10
The Kb value for H2AsO4- is
A.
2.0 x 10-10
B.
8.0 x 10-8
C.
1.2 x 10-7
D.
1.7 x 10-5
34.
In a solution at 25oC, the [H3O+] is 3.5 x 10-6 M. The [OH-] is
A.
3.5 x 10-20 M
B.
2.9 x 10-9 M
C.
1.0 x 10-7 M
D.
3.5 x 10-6 M
35.
In a solution with a pOH of 4.22, the [OH-] is
A.
1.7 x 10-10 M
B.
6.0 x 10-5 M
C.
6.3 x 10-1 M
D.
1.7 x 104 M
36.
An aqueous solution of NH4CN is
A.
B.
C.
D.
basic because Ka < Kb
basic because Ka > Kb
acidic because Ka < Kb
acidic because Ka > Kb
37.
The net ionic equation for the predominant hydrolysis reaction of KHSO 4 is
A.
HSO4- + H2O  SO42- + H3O+
B.
HSO4- + H2O  H2SO4 + OHC.
KHSO4 + H2O  K+ + SO42- + H3O+
D.
KHSO4 + H2O  K+ + H2SO4 + OH-
38.
The [OH-] in an aqueous solution always equals
A.
Kw x [H3O+]
B.
Kw - [H3O+]
C.
Kw/[H3O+]
D.
[H3O+]/Kw
39.
The [H3O+] in a solution with pOH of 0.253 is
A.
5.58 x 10-15 M
B.
1.79 x 10-14 M
C.
5.58 x 10-1 M
D.
5.97 x 10-1 M
40.
The equilibrium expression for the hydrolysis reaction of 1.0 M K 2HPO4 is
A.
[H2PO4-][OH-]
B.
[H3PO4][OH-]
[HPO42-]
[H2PO4-]
C.
41.
42.
[K+]2 [HPO42-]
[K2HPO4]
2.30
3.30
10.70
11.7
Consider the following equilibrium for an indicator: HInd + H2O  Ind- + H3O+
At the transition point,
A.
[HInd] > [Ind-]
B.
[HInd] = [Ind-]
C.
[HInd] < [Ind-]
D.
[HInd] = [H3O+]
Acids Unit Midterm Practice Test
1.
D.
The solution with the highest pH is
A.
1.0 M NaCl
B.
1.0 M NaCN
C.
1.0 M NaIO3
D.
1.0 M Na2SO4
The pH of 100.0 mL of 0.0050 M NaOH is
A.
B.
C.
D.
43.
[K+] [KHPO4-]
[K2HPO4]
Subjective
a)
Write the net ionic equation for the reaction between NaHSO3 and
NaHC2O4.
b)
Explain why the reactants are favoured in the above reaction.
2.
What is the [H3O+] in a solution formed by adding 60.0 mL of water to 40.0 mL
of 0.400 M KOH?
3.
A solution of 0.100 M HOCN has a pH of 2.24. Calculate the Ka value for the acid.
4.
Calculate the pH in 100.0 mL 0.400 M H3BO3.
5.
Calculate the pH of the solution formed by mixing 20.0 mL of 0.500 M HCl with 30.0 mL of 0.300 M
NaOH.
6.
7.
a)
Write the balanced equation representing the reaction of HF with H 2O.
b)
Identify the Bronsted-Lowry bases in the above equation.
Consider the following data:
Barbituric acid
Sodium propanoate
Propanoic acid
HC4H3N2O3
NaC3H5O2
HC3H5O2
Ka = 9.8 x 10-5
Kb = 7.5 x 10-10
Ka = ?
Which is the stronger acid, propanoic acid or babituric acid? Explain using calculations.
8.
A solution of 0.0100 M lactic acid, HC3H5O3, has a pH of 2.95. Calculate the Ka value.
9.
a)
Write equations showing the amphiprotic nature of water as it reacts with
HCO3-.
b)
10.
Calculate the Kb for HCO3-.
Calculate the [H3O+] in 0.550 M C6H5COOH.
Acids
Quiz #1
Properties of Acids and BasesStrength
1.
Drano®, a commercial product used to clean drains, contains small bits of
metal and
A.
ammonia
B.
acetic acid
C.
hydrochloric acid
D.
sodium hydroxide
2.
A net ionic equation for the reaction between CH3COOH and KOH is
A.
CH3COOH(aq) + K+(aq) ⇄ CH3COOK(aq)
B.
CH3COOH(aq) + OH-(aq) ⇄ H2O(l) + CH3COO-(aq)
C.
CH3COOH(aq) + KOH(aq) ⇄ H2O(l) + CH3COOK(aq)
D.
CH3COOH(aq) + K+(aq) + OH-(aq)⇄ H2O(l) + KCH3COO(aq)
3.
Which equation represents a neutralization reaction?
A.
Pb2+(aq) + 2Cl-(aq) → PbCl2(s)
B.
HCl(aq) + NH3(aq) → NH4Cl(aq)
C.
BaI2(aq) + MgSO4(aq) → BaSO4(s) + MgI2(aq)
D.
MnO4-(aq) + 5Fe2+(aq) +8H+(aq) → Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
4.
An Arrhenius acid is a substance that
A.
accepts a proton
B.
donates a proton
C.
produces H+ in solution
D.
produces OH- in solution
5.
Consider the following data table:
Breaker
1
2
3
4
Volume
15 mL
20 mL
25 mL
50 mL
aluminum
Contents
0.1 M Sr(OH)2
0.2 M NH4OH
0.1 M KOH
0.2 M NaOH
The beaker that requires the smallest volume of 1.0 M HCl for complete neutralization is:
A.
Beaker 1
B.
Beaker 2
C.
Beaker 3
D.
Beaker 4
6.
The net ionic equation for the titration of HClO4(aq) with LiOH(aq) is
A.
H+(aq) + OH-(aq) → H2O(l)
B.
HClO4(aq) + OH-(aq) → ClO4-(aq) + H2O(l)
C.
HClO4(aq) + LiOH(aq) → LiClO4(aq) + H2O(l)
D.
H+(aq) + ClO4-(aq) + Li+(aq) + OH-(aq) → LiClO4(aq) + H2O(l)
7.
8.
The equilibrium constant expression for a sulphurous acid is
A
Ka = [H+][HSO3-]
B.
Ka = [H+][HSO3-]
[H2SO3]
C.
Ka = [2H+][SO32-]
[H2SO3]
D.
Ka = [H+][SO32-]
[H2SO3]
To distinguish between a strong acid and a strong base, an experimenter could use
A.
odor
B.
magnesium
C.
a conductivity test
D.
the common ion test
9.
How many acids from the list below are known to be weak acids?
HCl,
HF,
H2SO3, H2SO4, HNO3,
HNO2
A.
2
B.
3
C.
4
D.
5
10.
There are two beakers on a laboratory desk. One beaker contains 1.0 M HCl and
the other contains tap water. To distinguish the acid solution from the water, one
would use
A.
a piece of copper.
B.
a piece of magnesium
C.
phenolphthalein indicator
D.
a piece or red litmus paper
11.
Caustic soda, NaOH, is found in
A.
fertilizers
B.
beverages
C.
toothpaste
D.
oven cleaners
12.
Which of the following is the strongest acid?
A.
Acetic acid
B.
Oxalic acid
C.
Benzoic acid
D.
Carbonic acid
13.
The acid used in the lead-acid storage battery is
A.
HCl
B.
HNO3
C.
H2SO4
D.
CH3COOH
14.
A definition for a Brønsted-Lowry acid is:
A.
the donation of H+
B.
the donation of OHC.
the acceptance of H+
D.
the acceptance of OH-
15.
What species will form when H+ ions are in the presence of H2O molecules?
A.
HO+
B.
HO2+
C.
D.
16.
17.
H3O+
H2O2
What is the conjugate acid of the base HAsO42-?
A.
AsO43B.
H2AsO42C.
H2AsO4D.
H3AsO4
What species is the hydronium ion?
A.
HO+
B.
HO2+
C.
H2O2
D.
H3O+
18.
Which of the following is a characteristic of acidic solutions?
A.
They accept H+ ions.
B.
They have a pH more than 7.
C.
They react with Mg to produce H2 gas.
D.
They turn limus blue.
19.
Which of the following is a common base found in drain cleaners?
A.
bleach
B.
vinegar
C.
milk of magnesia
D.
sodium hydroxide
20.
Which of the following is a characteristic of basic solutions?
A.
They donate H+ ions.
B.
They have a pH less than 7.
C.
They feel slippery.
D.
They turn limus red.
21.
Which of the following is the weakest acid?
A.
Acetic acid
B.
Oxalic acid
C.
Benzoic acid
D.
Carbonic acid
22.
What is the main difference between a strong acid and a weak acid?
A.
their degree of ionization
B.
their reactivity with platinum
C.
their concentration in solution
D.
their effect on an indicator
23.
Which of the following best describes an acidic solution?
Litmus Colour
Reaction with Zn
A.
red
reaction
B.
red
no reaction
C.
blue
no reaction
D.
blue
reaction
24.
Identify the common acid found in the stomach.
A.
benzoic acid
B.
sulphuric acid
C.
perchloric acid
D.
hydrochloric acid
25.
Which of the following best describes an basic solution?
Litmus Colour
pH
A.
red
less than 7
B.
red
more than 7
C.
blue
less than 7
D.
blue
more than 7
26.
Which of the following is a general characteristic of Arrhenius acids?
A.
They produce H+ in solution.
B.
They accept an H+ from water.
C.
They turn bromothymol blue a blue colour.
D.
They react with H3O+ ions to produce H2.
27.
Identify a conjugate pair from the equilibrium provided:
PO43A.
B.
C.
D.
+ HCO3⇌ HPO422CO3
and
HPO423PO4
and
HCO3PO43- and
HPO42HCO3- and
HPO42-
+
CO32-
28.
Which of the following best describes a weak acid?
A.
pH = 10
B.
It may be very soluble, but only partly ionized.
C.
It must be very soluble and completely ionized.
D.
It must be of low solubility and completely ionized.
29.
Which of the following is a common property of acid solutions?
A.
They have a pH > 7.
B.
They turn red litmus blue.
C.
They have a slippery feeling.
D.
They turn pink phenolphthalein colourless.
30.
What is a general characteristic of bases?
A.
They all donate H+.
B.
They all accept OH- .
C.
They will neutralize acids
D.
They will react with acids to produce H2 gas.
Acids
Quiz #2
1.
A test that could be safely used to distinguish a strong base from a weak base is
A.
taste
B.
touch
C.
litmus paper
D.
electrical conductivity
2.
Identify the two substances that act as Bronsted-Lowry bases in the equation
A.
B.
C.
D.
3.
Conjugates, Reactions, Amphiprotic, Arrhenius, Bronsted Bases, KB, & Strength
HS- + SO42- ⇄ S2- + HSO4HS and S2SO4 2-and S2HS- and HSO4SO42- and HSO4-
The conjugate acid of H2C6H5O-7 is
A.
B.
C.
D.
C6H5O73HC6H5O72H2C6H5O7
H3C6H5O7
4.
Which one of the following substances is/are amphiprotic?
(1) H3PO4
(2) H2PO4(3) HPO42A.
2 only
B.
3 only
C.
1 and 2
D.
2 and 3
5.
The net ionic equation for the neutralization of HBr by Ca(OH) 2 is
A.
H+(aq) + OH-(aq)  H2O(l)
B.
Ca2+(aq) + 2Br-(aq)  CaBr2(s)
C.
2HBr(aq) + Ca(OH)2(aq)  2H2O(l) + CaBr2(s)
D.
2H+(aq) + Ca(OH)2(aq)  2H2O(l) + Ca2+(aq)
6.
If reactants are favored in the following equilibrium, the stronger base must be
HCN + HS - ⇄ H2S + CN A.
B.
C.
D.
H2S
HSCNHCN
7.
The hydronium ion, H3O+ is a water molecule that has
A.
lost a proton
B.
gained a proton
C.
gained a neutron
D.
gained an electron
8.
The complete ionic equation for the neutralization of acetic acid by sodium
hydroxide is
9.
A.
H+ + OH- ⇄ H2O
B.
CH3COOH + OH- ⇄ CH3COOH- + H2O
C.
CH3COOH + NaOH ⇄ NaCH3COOH + H2O
D.
CH3COOH + Na+ + OH- ⇄ Na+ + CH3COO- + H2O
In the following Bronsted – Lowry acid-base equation:
NH4+ (aq) + H2O(l) ⇄ NH3(aq) + H3O+(aq)
The stronger base is
A.
NH4+
B.
H2O
C.
NH3
D.
H3O+
10.
Consider the following equilibrium system:
OCl-(aq) + HC7H5O2(aq) ⇄ HOCl(aq) + C7H5O2-(aq)
Keq= 2.1 x 103
At Equilibrium
A.
products are favored and HOCl is the stronger acid
B.
C.
D.
11.
reactants are favored and HOCl is the stronger acid
products are favored and HC7H5O2 is the stronger acid
reactants are favored and HC7H5O2 is the stronger acid
In the equilibrium system
H2BO3- (aq) + HCO3-(aq) ⇄ H2CO2(aq) + HBO32-(aq)
The two species acting as Bronsted-Lowry acids are
A.
HCO3- and H2CO3
B.
H2BO3- and H2CO3
C.
HCO3- and HBO32D.
H2BO3- and HBO32-
12.
Consider the following equilibrium HS- + H2C2O4 ⇄ HC2O4- + H2S
The stronger acid is
A.
HSB.
H2C2O4
C.
HC2O4D.
H2S
13.
Which species is not amphiprotic?
A.
H2O
B.
H3BO3
C.
HPO42D.
HC6H5O72-
14.
Given the equilibrium: H2BO3- +
Which is the strongest acid?
A.
H2PO4B.
H3BO3
C.
HPO42D.
H2BO3-
15.
Which equation represents the reaction of a Brønsted-Lowry base with water?
A.
2Na + 2H2O
 2NaOH + H2
⇌ H3BO3 +
HPO42-
⇌ P2H5+ + OH-
B.
P2H4 + H2O
C.
H2PO4- + H2O ⇌ H3O+
D.
H2C2O4
+
H2PO4-
H2O
⇌
+
H3O+
HPO42+
HC2O4-
16.
A definition for a Brønsted-Lowry acid is:
A.
the donation of H+
B.
the donation of OHC.
the acceptance of H+
D.
the acceptance of OH-
17.
Which of the following would be a typical pH of soda pop?
A.
1.0
B.
3.2
C.
6.8
D.
7.2
18.
When three 25.0 mL samples of the strong acid H2SO4 were titrated with 0.25 M NaOH
19.
the following results were obtained:
Volume of NaOH(aq)
47.2 mL 39.9 mL 40.1 mL
What is the concentration of the H2SO4 sample?
A.
0.20 M
B.
0.21 M
C.
0.40 M
D.
0.42 M
Which of the following is a common base found in drain cleaners?
A.
baking soda
B.
vinegar
C.
milk of magnesia
D.
sodium hydroxide
20.
Identify the common acid found in the stomach.
A.
benzoic acid
B.
sulphuric acid
C.
perchloric acid
D.
hydrochloric acid
21.
Which of the following solutions would typically show the greatest electrical
conductivity?
A.
1.0 M weak acid
B.
0.8 M weak base
C.
0.5 M strong acid
D.
0.1 M strong base
22.
Which of the following are amphiprotic in aqueous solutions?
I.
H3BO3
II.
H2BO3- III.
HBO32- IV.
A.
I only
B.
IV only
C.
I and II only
D.
II and III only
23.
Which of the following best describes an acidic solution?
Litmus Colour
Reaction with Mg
A.
red
reaction
B.
red
no reaction
C.
blue
no reaction
D.
blue
reaction
24.
What is the Ka expression for H3PO4 ?
A.
Ka = [H+]3
[PO43-]
B.
Ka = [H+]3
[H3PO4]
25.
C.
Ka = [H+][H2PO4-]
[H3PO4]
D.
Ka = [H+][HPO42-]
[H3PO4]
BO33-
What volume of 0.250 M KOH is needed to titrate 0.00230 mol of the acid H2C2O4 ?
A.
1.15 mL
B.
4.60 mL
C.
9.20 mL
D.
26.
27.
18.4 mL
What is the net ionic equation for the titration of H3PO4(aq) with Sr(OH)2(aq) ?
A.
H+(aq) + OH- (aq)  H2O(l)
B.
6H+(aq) + 6OH-(aq)  6H2O(l)
C.
2H3PO4(aq) + 3Sr2+(aq) + 6OH-(aq)  Sr3(PO4)2(s) + 6H2O(l)
D.
6H+(aq) + 2PO43-(aq) + 3Sr2+(aq) + 6OH- (aq)  3Sr2+ + 2PO43-(aq) + 6H2O(l)
Which of the following is a piece of equipment used in acid-base titrations?
A.
burette
B.
test tube
C.
litmus paper
D.
graduated cylinder
28.
Water has the greatest tendency to act as an acid with which of the following?
A.
ClB.
NO2C.
H2PO4D.
CH3COO-
29.
Which net ionic equation best describes the reaction between NaOH and H 2S?
A.
H+(aq) + OH-(aq)  H2O(l)
B.
H2S(aq) + 2OH-  2H2O(l) + S2-(aq)
C.
H2S(aq) + 2NaOH  2H2O(l) + Na2S(aq)
D.
2H+(aq) + S2-(aq) + 2NaOH(aq)  2H2O(l) + 2Na+(aq) + S2-(aq)
30.
Which of the following is a general characteristic of Arrhenius acids?
A.
They produce H+ in solution.
B.
They accept an H+ from water.
C.
They donate a proton.
D.
They accept a proton.
31.
Which of the following is a general characteristic of Bronstead acids?
A.
They produce H+ in solution.
B.
They accept an H+ from water.
C.
They donate a proton.
D.
They accept a proton.
32.
Which of the following is a general characteristic of Arrhenius bases?
A.
They produce OH- in solution.
B.
They produce H3O+ in water.
C.
They donate a proton.
D.
They accept a proton.
33.
Which of the following is a general characteristic of Bronstead base?
A.
They produce H+ in solution.
B.
They accept an H+ from water.
C.
They donate a proton.
D.
They accept a proton
Acids
1.
Quiz #3 Leveling effect, Anhydrides, Hydrolysis, Relationships
Which of the following oxides will form the most acidic solution?
A.
SO2
B.
MgO
C.
Na2O
D.
2.
3.
Al2O3
Which one of the following salts will produce an acidic solution?
A.
KBr
B.
LiCN
C.
NH4Cl
D.
NaCH3COO
The balanced equation for the reaction between sodium oxide and water is
A.
Na2O + H2O → 2NaOH
B.
Na2O + H2O → 2NaH + O2
C.
Na2O + H2O → 2Na + H2O2
D.
Na2O + H2O → 2Na + H2 +O2
4.
Normal rainwater is slightly acidic due to the presence of dissolved
A.
methane
B.
carbon dioxide
C.
sulphur dioxide
D.
nitrogen dioxide
5.
Which of the following oxides would hydrolyze to produce hydroxide ions?
A.
NO
B.
SO2
C.
Cl2O
D.
Na2O
6.
The approximate pH of “normal” rainwater is
A.
0
B.
6
C.
7
D.
8
7.
Which of the following oxides would hydrolyze to produce hydronium ions?
A.
CaO
B.
SO2
C.
MgO
D.
Na2O
8.
Which of the following gasses results in the formation of acid rain?
A.
H2
B.
O3
C.
SO2
D.
NH3
9.
Consider the following acid base solution
HSO3- + HF ⇄ H2SO3 + FThe order of Bronsted-Lowry acids and bases in this equation is
10.
A.
acid + base ⇄ acid + base
B.
acid + base ⇄ base + acid
C.
base + acid ⇄ base + acid
D.
base + acid ⇄ acid + base
The conjugate acid of OH- is
A.
H+
B.
O2-
C.
D.
H2O
H3O+
11.
Which of the following 0.10 M solutions will have the greatest electrical
conductivity?
A.
HF
B.
NH3
C.
NaOH
D.
C6H5COOH
12.
The amphiprotic ion HSeO3- can undergo hydrolysis according to the following equations
Kb
HSeO3- + H2O ⇄ H2SeO3 + OHHSeO3- + H2O ⇄ SeO32-+ H3O+
Ka
An aqueous solution of HSeO3- is found to be acidic. HSeO3- behaves mainly as a
A.
proton donor, and Kb is less than Ka
B.
proton donor, and Kb is greater than Ka
C.
proton acceptor, and Kb is less than Ka
D.
proton acceptor, and Kb is greater than Ka
13.
The Kb expression for HPO42- is
A. [PO43-][H3O+]
B. [HPO42-][OH-]
2[HPO4 ]
[H 2PO4-]
C. [H2PO4 ][OH ]
D. [HPO42-][ H3O+]
2[HPO4 ]
[PO43-]
14.
Given the equilibrium:
H2BO3- + H2PO4- ⇌ H3BO3 + HPO42Which is the strongest acid?
A.
HPO42B.
H3BO3
C.
H2PO4D.
H2BO3-
15.
Which species is not amphiprotic?
A.
H2O
B.
H3BO3
C.
H2PO4D.
H2C6H5O7-
16.
What is produced when CH3NH2 acts as a base in water?
A.
CH3NHB.
CH3NH3+
C.
CH3NH2+
D.
CH2NH2-
17.
What species will form when H+ ions are in the presence of H2O molecules?
A.
HO+
B.
H2O+
C.
H3O+
D.
H2O2+
18.
What is the conjugate acid of the base HAsO42-?
A.
AsO43B.
H2AsO4C.
H4AsO4+
D.
H3AsO4
19.
Which solution will have the greatest electrical conductivity?
A.
0.50 M HCl
B.
0.10 M RbOH
C.
0.50 M K3PO4
D.
2.0 M C6H12O6
20.
The following equilibrium favours the formation of products:
NH2OH + CH3NH2
⇌ CH3NH- + NH3OH+
Which species is the strongest acid?
A.
NH3OH+
B.
NH2OH
C.
CH3NH2
D.
CH3NH-
21.
Which of the following solutions would have the greatest [OH -]?
A.
0.1 M HCO3B.
0.1 M HPO42C.
0.1 M H2PO4D.
0.1 M CO32-
22.
Which of the following amphiprotic ions will act predominantly as a base in
solution?
A.
HSO3B.
HSO4C.
H2PO4D.
HPO42-
23.
Which is the strongest base?
A.
ClB.
NO2C.
HPO42D.
CH3COO-
24.
Water has the greatest tendency to act as an acid with which of the following?
A.
ClB.
NO2C.
HPO42D.
CH3COO-
25.
Water has the greatest tendency to act as a base with which of the following?
A.
HF
B.
H2CO3
C.
H3PO4
D.
CH3COOH
26.
Which of the following will have the smallest Ka value?
A.
HF
B.
H2CO3
C.
H3PO4
D.
27.
28.
CH3COOH
Which of the following will have the smallest Kb value?
A.
IO3B.
NH3
C.
CND.
HPO42What volume of 0.500 M NaOH is required to neutralize 25.0 mL 0.250 M HBr ?
A.
5.00 mL
B.
12.5 mL
C.
20.0 mL
D.
25.0 mL
29.
Which of the following Ka values represents the acid that is strongest?
A.
Ka = 2.8 x 10-18
B.
Ka = 8.2 x 10-18
C.
Ka = 4.4 x 10-22
D.
Ka = 6.4 x 10-22
30.
Which of the following Ka values represents the acid with the strongest conjugate
base?
A.
Ka = 2.8 x 10-18
B.
Ka = 8.2 x 10-18
C.
Ka = 4.4 x 10-22
D.
Ka = 6.4 x 10-22
Acids
Quiz #4
Anhydrides, Hydrolysis
1.
Which of the following pairs of gases are primarily responsible for producing
“acid rain”?
A.
O2 and O3
B.
N2 and O2
C.
CO and CO2
D.
NO2 and SO2
2.
Sodium potassium tartrate (NaKC4H4O6) is used to raise the pH of fruit during
processing. In this process, sodium potassium tartrate is being used as a/an
A.
salt
B.
acid
C.
base
D.
buffer
3.
The net ionic equation for the hydrolysis of the salt, Na2S is
A.
Na2S ⇄ 2Na+ + S2-
B.
S2- + H2O ⇄ OH- + HS+
C.
Na2S + 2H2O ⇄ 2NaOH + H2S
D.
2Na+ + S2- + 2H2O ⇄ 2Na+ + 2OH- + H2S
4.
Which of the following solutions would be acidic?
A.
sodium acetate
B.
iron III chloride
C.
sodium carbonate
D.
potassium chloride
5.
Consider the following salts:
I.
NaF
II.
NaClO4 III.
NaHSO4
Which of these salts, when dissolved in water, would form a basic solution?
A.
I only
B.
I and II only
C.
II and III only
D.
I, II and III
6.
Which of the following, when dissolved in water, forms a basic solution?
A.
KCl
B.
NaClO4
C.
Na2CO3
D.
NH4NO3
7.
Which of the following oxides forms a basic solution?
A.
K2O
B.
CO2
C.
SO3
D.
NO2
8.
Which of the following is amphiprotic in water?
A.
SO2
B.
SO32C.
HSO3D.
H2SO3
9.
Consider the following equilibrium expression
K= [H2S][OH-]
[HS-]
This expression represents the
A.
Kb for H2S
B.
Ka for H2S
C.
Kb for HSD.
Ka for HS-
10.
The reaction of a strong acid with a strong base produces
A.
A salt and a water
B.
A base and an acid
C.
A metallic oxide and water
D.
A non-metallic oxide and water
11.
Consider the following equilibrium:
CH3COOH(aq) + NH3(aq) ⇄ CH3COO-(aq) ⇄ NH4+(aq)
The sequence of Bronsted-Lowry acids and bases is
A.
acid, base, base, acid
B.
acid, base, acid, base
C.
base, acid, base, acid
D.
base, acid, acid, base
12.
The pH range of ‘acid rain’ is often
A.
3 to 6
B.
6 to 8
C.
7 to 9
D.
10 to 12
13.
Water will act as a Bronsted-Lowry acid with
A.
NH3
B.
C.
D.
H2S
HCN
HNO3
14.
Which of the following is a conjugate acid-base pair?
A.
H3PO4 and PO43B.
H2PO4- and PO43C.
H3PO4 and HPO42D.
H2PO4- and HPO42-
15.
What is the net ionic equation for the hydrolysis of NH 4Cl ?
A.
NH4Cl  NH4+ + Cl⇌ HCl + OH-
B.
Cl- +
C.
NH4+ + H2O
⇌ H3O+ + NH3
D.
NH4+ + H2O
⇌ NH5+ + OH-
H2O
16.
What is the approximate pH of the salt NH4Cl?
A.
0.0
B.
5.0
C.
7.0
D.
9.0
17.
Four samples of rain are collected from different geographic regions and the pH is
measured for each sample.
Sample
pH
1
2.8
2
4.0
3
6.2
4
6.8
Which of the above samples would be classified as acid rain?
A.
1 only
B.
1 and 2
C.
1, 2 and 3
D.
1, 2, 3 and 4
18.
Which of the following equations correctly represents the reaction of a metallic
oxide with water?
A.
K2O + H2O  2KOH
B.
SO3 + H2O  H2SO4
C.
Na2O + H2O  Na2O2 + H2
D.
NaOH + HCl  NaCl + HOH
19.
Which of the following equations describes the dissociation of the salt,
ammonium chloride, in water?
A.
NH4Cl(s)  NH4Cl(aq)
B.
NH4+(aq) + Cl-(aq)  NH4Cl(s)
C.
NH4Cl(s)  NH4+(aq) + Cl-(aq)
D.
NH4+(aq) + H2O(l)  H3O+(aq) + NH3(aq)
20.
Which of the following salt solutions is acidic?
A.
NaBr
B.
C.
D.
AlCl3
Li2CO3
NaHCO3
21.
What pH would most likely result when CO2 dissolves naturally in rainwater?
A.
3.0
B.
6.6
C.
7.0
D.
7.5
22.
Which of the following is the net ionic equation that describes the hydrolysis that
occurs in a K2CO3 solution?
A.
K2CO3  2K+ + CO32B.
2K+ + CO32-  K2CO3
C.
CO32- + H2O  HCO3- + OHD.
K2CO3 + 2H2O  H2CO3 + 2KOH
23.
What is the equilibrium expression for the predominant equilibrium in
NaCN(aq)?
A.
Kb
=
[Na+][CN-]
[NaCN]
B.
Kb
=
[Na+][CN-]
[HCN]
C.
Kb
=
[HCN][OH-]
[CN-]
D.
Kb
=
[CN-][OH-]
[HCN]
24.
Which of the following salts will be basic?
A.
KCl
B.
NH4Cl
C.
NaBr
D.
Na2HPO4
25.
Which of the following salts will be acidic?
A.
KCl
B.
NH4Cl
C.
NaBr
D.
Na2HPO4
26.
Which is an acid anhydride?
A.
Na2O
B.
NaCl
C.
CO2
D.
NH4Cl
27.
Which formula has the highest pH when in water?
A.
Na2O
B.
NaCl
C.
CO2
D.
NH4Cl
28.
Which formula has the lowest pH when in water?
A.
B.
C.
D.
Na2O
NaCl
FeCl3
NaHCO3
29.
Which formula has the lowest pH when in water?
A.
NaCN
B.
Na2HPO4
C.
NH4Cl
D.
NaOH
30.
Which of the following properties is true for a solution of NaNO3 ?
A.
It is neutral.
B.
It is very basic.
C.
It is slightly basic.
D.
It is slightly acidic.
Acids
Quiz #5
pH calculations for Strong and Weak Acids
1.
The 1.0 M acidic solution with the highest pH is
A.
H2S
B
HNO2
C.
HNO3
D.
H3BO3
2.
At 25 oC, the equation representing the ionization of water is
A
H2O + H2O ⇄ 2H2 + O2
B.
H2O + H2O ⇄ H2O2
C.
H2O + H2O ⇄
D.
H2O + H2O ⇄ H3O +OH-
4H+
+ 2O
+ H2
2-
+
3.
The pH of a 0.3 M solution of NH3 is approximately
A.
14.0
B.
11.0
C.
6.0
D.
3.0
4.
The pH of an aqueous solution is 4.32. The [OH-] is
A.
6.4 x 10-1 M
B.
4.8 x 10-5 M
C.
2.1 x 10-10 M
D.
1.6 x 10-14 M
5.
The pH of an aqueous solution is 10.32. The [OH-] is
A.
5.0 x 10-12 M
B.
2.0 x 10-11 M
C.
4.8 x 10-11 M
D.
2.1 x 10-4 M
6.
The pH of a 0.025 M HClO4 solution is
A.
0.94
B.
1.60
C.
12.40
D.
7.
13.06
Consider the following equilibrium
H2O(l) + H2O(l) ⇄ H3O+(aq) + OH-(aq)
The equilibrium constant for this system is referred to as
A.
Kw
B.
Ka
C.
Kb
D.
Ksp
8.
The [H3O-] in a solution of pH 0.60 is
A.
4.0 x 10-14 M
B.
2.2 x 10-1 M
C.
2.5 x 10-1 M
D.
6.0 x 10-1 M
9.
A solution is prepared by adding 100 mL of 10 M of HCl to a 1 litre volumetric
flask and filling it to the mark with water. The pH of this solution is
A.
-1
B.
0
C.
1
D.
7
10.
The approximate pH of a 0.06 M solution of CH3COOH is
A.
1
B.
3
C.
11
D.
13
11.
The [OH-] is greater than the [H3O+] in
A.
HCl(aq)
B.
NH3(aq)
C.
H2O(aq)
D.
CH3COOH(aq)
12.
The pH of 0.15 M HCl is
A.
0.15
B.
0.71
C.
0.82
D.
13.18
13.
Which of the following equations correctly relates pH and [H 3O+] ?
A.
pH= log [H3O+]
B.
pH= 14 - [H3O+]
C.
pH= -log [H3O+]
D.
pH= pKw – [H3O+]
14.
The pH of 0.20 M HNO3 is
A.
0.20
B.
0.63
C.
0.70
D.
1.58
15.
The [OH-] in 0.050 M HNO3 at 25oC is
A.
5.0 x 10-16 M
B.
1.0 x 10-14 M
16.
C.
2.0 x 10-13 M
D.
5.0 x 10-2 M
What is the approximate pH of a 1.0 M solution of the salt NH 4Cl?
A.
0.0
B.
5.0
C.
7.0
D.
10.0
17.
What is the approximate pH of a 1.0 M solution of the salt NH 3?
A.
2.0
B.
5.0
C.
7.0
D.
10.0
18.
What is the approximate pH of a 1.0 M solution of the salt NaHCO3?
A.
2.0
B.
5.0
C.
7.0
D.
9.0
19.
What is the approximate pH of a 1.0 M solution of the salt AlCl 3?
A.
0.0
B.
3.0
C.
7.0
D.
9.0
20.
What is the approximate pH of a 1.0 M solution of the salt NaHSO 4?
A.
0.0
B.
3.0
C.
7.0
D.
9.0
21.
Which of the following is a definition of pH?
A.
pH = +Log[H3O+]
B.
pH = -Log[OH-]
C.
pH = -Log[H3O+]
D.
pH = pOH + pKw
22.
What is the mass of NaOH required to prepare 100.0 mL of NaOH (aq) that has a
pH = 13.62 ?
A.
0.38 g
B.
0.42 g
C.
1.67 g
D.
0.14 g
What is the [KOH] in a KOH solution that has a pH =12.00?
A.
0.010 M
B.
0.56 M
C.
2.0 M
D.
2.0 x 10-12 M
23.
24.
What is the [H3O+] in 0.70 M HCN ?
A.
0.70 M
B.
1.9 x 10-5 M
C.
1.0 x 10-6 M
D.
2.4 x 10-10 M
25.
What is the value of pKw for water at 25 0C ?
A.
10 x 10-14
B.
10 x 10-7
C.
7.00
D.
14.00
26.
What pH would most likely result when CO2 dissolves naturally in rainwater?
A.
3.5
B.
6.5
C.
7.2
D.
7.8
27.
What is the pOH of 0.05 M Ba(OH)2 ?
A.
1.0
B.
1.3
C.
12.7
D.
13.0
28.
What is the pH of 0.5 M Sr(OH)2 ?
A.
0.0
B.
0.3
C.
12.7
D.
14.0
29.
Which of the following statements is true for an acidic solution at 25 0C ?
A.
pH > 7.0
B.
pOH < 7.0
C.
[H3O+] < [OH-]
D.
[H3O+] > [OH-]
30.
What is the [OH-] in 0.025 M HCl ?
A.
2.5 x 10-16 M
B.
4.0 x 10-13 M
C.
2.5 x 10-2 M
D.
1.6 M
Acids
Quiz #6 Ka’s from pH
Kb’s from Ka’s
1.
The Kb for the dihydrogen phosphate ion is
A.
1.4 x 10-12
B.
6.3 x 10-8
C.
1.6 x 10-7
D.
7.1 x 10-3
2.
What volume of 0.100 M NaOH is required to neutralize a 10.0 mL sample of 0.200 M
A.
5.0 mL
B.
10.0 mL
C.
20.0 mL
D.
40.0 mL
H2SO4?
3.
Consider the following equilibria:
I
HCO3- + H2O ⇄ H2CO3 + OH-
II
NH4+ + H2O ⇄ H3O+ + NH3
III
HSO3- + H3O ⇄ H2O + H2SO3
Water acts as a Bronsted-Lowry base in
A.
III only
B.
I and II only
C.
II and III only
D.
I, II, and III
4.
Which of the following is represented by a Kb expression?
A.
Al(OH)3(s) ⇄ Al3+(aq) + 3OH-(aq)
B.
HF(aq) + H2O(l) ⇄ H3O+(aq) + F-(aq)
C.
D.
Cr2O72-(aq) + 2OH-(aq) ⇄ 2CrO42-(aq) + H2O(l)
CH3NH2(aq) + H2O(l) ⇄ CH3NH3+ (aq) + OH-(aq)
5.
A student combines 0.25 mol of NaOH and 0.20 mol of HCl in water to make 2.0 liters
of solution. The pH of the solution is
A.
1.3
B.
1.6
C.
12.4
D.
12.7
6.
If OH- is added to a solution, the [H3O+] will
A.
Remain constant
B.
Adjust such that [H3O+]= [OH-]
Kw
C.
Increase such that [H3O+][OH-] = Kw
D.
Decrease such that [H3O+][OH-] = Kw
7.
In a titration, 10.0 mL of H2SO4(aq) is required to neutralize 0.0400 mol of NaOH.
From this data, the [H2SO4] is
A.
0.0200 M
B.
2.00 M
C.
4.00 M
D.
8.00 M
8.
Consider the following equilibrium for an acid-base indicator:
Hlnd ⇄ H+ + IndKa = 1.0 x 10-10
Which of the following statements is correct at pH 7.0?
A.
B.
C.
D.
9.
[Ind-] < [HInd]
[Ind-] = [HInd]
[Ind-] > [HInd]
[Ind-] = [H+] = [HInd]
Which of the following indicators would be yellow at pH 7 and blue at pH 10?
A.
thymol blue
B.
methyl violet
10.
11.
12.
C.
bromthymol blue
D.
bromcresol green
Consider the following equilibrium for phenolphthalein: HInd + H + + IndWhen phenolphthalein is added to 1.0 M NaOH, the color of the resulting solution is
A.
pink as [HInd] is less than [Ind-]
B.
pink as [HInd] is greater than [Ind-]
C.
colorless as [HInd] is less than [Ind-]
D.
colorless as [HInd] is greater than [Ind-]
Water acts as a base when it reacts with
A.
CNB.
NH3
C.
NO2D.
NH4+
What is the pH of a solution prepared by adding 0.50 mol KOH to 1.0 L of 0.30 M
HNO3?
A.
0.20
B.
0.70
C.
13.30
D.
13.80
13.
The 1.0 M acid solution with the largest [H3O+] is
A.
HNO2
B.
H2SO3
C.
H2CO3
D.
H3BO3
14.
Consider the following indicator equilibrium:
HInd(aq) + H2O(l) ⇌ H3O+ + Ind-(aq)
Colourless
blue
What is the effect of adding HCl to a blue sample of this indicator?
Equilibrium Shift Colour Change
A.
left
less blue
B.
left
more blue
C.
right
less blue
D.
right
more blue
15.
Consider the following indicator equilibrium:
HInd(aq) + H2O(l) ⇌ H3O+ + Ind-(aq)
Colourless
blue
What is the effect of adding NaOH to a light blue sample of this indicator?
Equilibrium Shift Colour Change
A.
left
less blue
B.
left
more blue
C.
right
less blue
D.
right
more blue
16.
Consider the following indicator equilibrium:
HInd(aq) + H2O(l) ⇌ H3O+ + Ind-(aq)
yellow
blue
What is the effect of adding FeCl3 to a green sample of this indicator?
Equilibrium Shift
Colour Change
A.
left
yellow
B.
left
blue
C.
right
less blue
D.
right
more blue
17.
An indicator has a Ka = 4 x 10-6. Which of the following is true for this indicator?
pH at Transition Point
Indicator
A.
4.0
methyl orange
B.
4.0
bromcresol green
C.
5.4
methyl red
D.
5.4
bromcresol green
18.
An indicator has a Ka = 6.3 x 10-13. Which of the following is true for this indicator?
pH at Transition Point
Indicator
A.
4.0
methyl orange
B.
4.0
bromcresol green
C.
12.2
alizarin yellow
D.
12.2
indigo carmine
19.
A salt forms in the reaction between HF(aq) and NaOH(aq). What is the net ionic equation
for the hydrolysis of this salt?
A.
NaF(aq)  Na+(aq) + F-(aq)
⇌ H3O+(aq) + F-(aq)
B.
HF(aq) + H2O(l)
C.
D.
F-(aq) + H2O(aq) ⇌ HF(aq) + OH-(aq)
HF(aq) + NaOH(aq)  NaF(aq) + H2O(l)
20.
What is the Ka for the indicator that is yellow in its basic form and blue in its acid form?
A.
6 x 10-13
B.
2 x 10-9
C.
2 x 10-7
D.
3 x 10-5
21.
What term is used to describe the point at which a chemical indicator changes colour?
A.
titration point
B.
transition point
C.
equivalence point
D.
stoichiometric point
22.
What is the net ionic equation in the reaction between HF(aq) and NaOH(aq)?
A.
NaF(aq)  Na+(aq) + F-(aq)
⇌ F-(aq) + H2O(l)
B.
HF(aq) + OH-(aq)
C.
D.
F-(aq) + H2O(aq) ⇌ HF(aq) + OH-(aq)
HF(aq) + NaOH(aq)  NaF(aq) + H2O(l)
23.
Consider the ionization of water: 2H2O ⇌ H3O+ + OHWhat happens to the pH when KOH is added to water?
A.
pH increases since [H3O+] increases.
B.
pH increases since [H3O+] decreases.
C.
pH decreases since [H3O+] increases.
D.
pH decreases since [H3O+] decreases.
24.
Which of the following salt solutions is acidic?
A.
NaBr
B.
FeCl3
C.
LiCN
D.
NaHCO3
25.
Which of the following salt solutions is neutral?
A.
NaBr
B.
FeCl3
C.
LiCN
D.
NaHCO3
26.
What happens to the [ion] in water when a small amount of HBr(aq) is added?
A.
[H3O+] = [OH-] = 1.0 x 10-7 M
B.
[H3O+] and [OH-] both increase
C.
[H3O+] increases and [OH-] decreases
D.
[H3O+] increases and [OH-] is unchanged
27.
What is the Kb value for H2PO4- ?
A.
1.3 x 10-12
B.
6.2 x 10-8
C.
1.6 x 10-7
D.
2.5 x 103
28.
What is one of the Ka values for thymol blue?
A.
2 x 10-9
B.
2 x 10-7
C.
1 x 10-7
D.
5 x 102
29.
Consider the indicator equilibrium:
HInd(aq) + H2O(l) ⇌ H3O+ + Ind-(aq)
yellow
blue
Which of the following is true about the transition point of this indicator?
A.
pH = 7.0
B.
[HInd] = [Ind-]
C.
[HInd] > [Ind-]
D.
moles of H3O+ moles of Ind-
30.
What is the [OH-] in 0.25 M HCl ?
A.
2.5 x 10-15 M
B.
4.0 x 10-14 M
C.
2.5 x 10-1 M
D.
2.5 x 10-13 M
Acids
Quiz #7 pH Relationships, Amphiprotic Calculations
1.
In water, the hydrogen phosphate ion, HPO42-, will act as
A.
An acid because Ka < Kb
B.
An acid because Ka > Kb
C.
A base because Ka < Kb
D.
A base because Ka > Kb
2.
A student records the pH of 1.0 M solution of two acids
Acid
pH
X
4.0
Y
2.0
Which of the following statements can be concluded from the above data?
A.
Acid X is stronger than acid Y
B.
Acid X and acid Y are both weak
3.
4.
C.
Acid X is diprotic while acid Y is monoprotic
D.
Acid X is 100 times more concentrated than acid Y
When added to water, the hydrogen carbonate ion, HCO 3-, produces a solution,
which is
A.
basic because Kb is greater than Ka
B.
basic because Ka is greater than Kb
C.
acidic because Ka is greater than Kb
D.
acidic because Kb is greater than Ka
The concentration, Ka and pH values of four acids are given in the following table
ACID
Concentration
Ka
pH
HA
HB
3.0 M
0.7 M
2.0 x 10-5
4.0 x 10-5
2.1
2.3
HC
HD
4.0 M
1.5 M
1.0 x 10-5
1.3 x 10-5
2.2
2.4
Based on this data, the strongest acid is
A.
HA
B.
HB
C.
HC
D.
HD
5.
Which of the following 0.10 M solutions is the most acidic?
A.
AlCl3
B.
FeCl3
C.
CrCl3
D.
NH4Cl
6.
Which of the following acid-base indicators has a transition point between pH 7
and pH 9?
A.
Ethyl red, Ka = 8 x 10-2
B.
Congo red, Ka = 9.0 x 10-3
C.
Cresol red, Ka = 1.0 x 10-8
D.
Alizarin blue, Ka = 7.0 x 10-11
7.
Which of the following would be a typical pH of lemon juice?
A.
1.0
B.
3.0
C.
6.8
D.
7.2
8.
What is the [KOH] in a KOH solution that has a pH = 13.00?
A.
0.10 M
B.
0.56 M
C.
1.0 M
D.
1.0 x 10-13 M
9.
What is the [H3O+] in 0.70 M HCN ?
A.
0.70 M
B.
1.9 x 10-5 M
C.
1.0 x 10-7 M
D.
3.4 x 10-10 M
10.
The S2- ion is a base with an equilibrium constant of 0.77, what is the Ka value
for HS-?
A.
1.3 x 10-14
B.
9.1 x 10-8
C.
1.1 x 10-7
D.
7.7 x 1013
11.
Which of the following is true as a result of the predominant hydrolysis of
NaHCO3?
Solution
Reason
A.
basic
Ka > Kb
B.
basic
Ka < Kb
C.
acidic
Ka > Kb
D.
acidic
Ka < Kb
12.
Consider the ionization of water: H2O + H2O ⇌ H3O+ + OHWhat happens to the pH when 0.1 M NaOH is added to water?
A.
pH increases since [H3O+] increases.
B.
pH increases since [H3O+] decreases.
C.
pH decreases since [H3O+] increases.
D.
pH decreases since [H3O+] decreases.
13.
Which of the following amphiprotic ions will act predominantly as a base in
solution?
A.
HSO3B.
HSO4C.
H2PO4D.
HPO42-
14.
Which of the following statements is true for a alkaline solution at 25 0C ?
A.
pH < 7.0
B.
pOH > 7.0
C.
[H3O+] > [OH-]
D.
[H3O+] < [OH-]
15.
What is the [OH-] in 0.025 M HCl ?
A.
2.5 x 10-16 M
B.
4.0 x 10-13 M
C.
2.5 x 10-2 M
D.
2.5 x 10-12 M
16.
Which of the following statements is true for an acidic solution at 25 0C ?
A.
pH > 7.0
B.
pOH < 7.0
C.
[H3O+] < [OH-]
D.
[H3O+] > [OH-]
17.
What is the [H+] in 0.025 M NaOH?
A.
2.5 x 10-16 M
B.
4.0 x 10-13 M
C.
2.5 x 10-2 M
D.
2.5 x 10-12 M
18.
What [H3O+] results when 25 mL of 1.0 M HCl is mixed with 10 mL of
2.0 M KOH ?
A.
0.70 M
B.
0.00 M
C.
0.10 M
D.
0.14 M
19.
What is the pH of 0.70 M HCN ?
A.
0.70 M
B.
1.9 x 10-5 M
C.
5.32
D.
4.73
20.
Consider the following indicator equilibrium:
HInd + H2O ⇌ H3O+ +
Indyellow
blue
What is the result of adding FeCl3 to this indicator?
Equilibrium Shift Colour
A.
left
blue
B.
left
yellow
C.
right
blue
D.
right
yellow
21.
An indicator changes colour when 1.0 M HCl is added. If the indicator has a
Ka = 1.0 x 10-10, determine the indicator and the pH at its transition point.
Indicator
pH
A.
phenolphthalein 4.0
B.
phenolphthalein 10.0
C.
thymolphthalein 4.0
D.
thymolphthalein 10.0
22.
What is the Kb value for HC6H5O72- ?
A.
1.0 x 10-14
B.
5.9 x 10-10
C.
2.4 x 10-8
D.
4.1 x 107
23.
What is produced when MgO is added to water?
A.
the metal Mg
B.
the acid HMgO
C.
the base Mg(OH)2
D.
the amphiprotic species HMgO2
24.
Which of the following is a major source of NO2(g), which contributes to
the problem of acid rain?
A.
a smelter
B.
an air conditioner
C.
a nuclear power plant
D.
the automobile engine
25.
The indicator bromothymol blue can be described by the following equilibrium
equation:
HInd + H2O ⇌ H3O+ +
Indyellow
blue
A base is added to the above indicator. After equilibrium has been re-established,
how do the H3O+ and the colour of the solution compare with the original
equilibrium?
H3O+
Colour of Solution
A.
decreases
turns blue
B.
decreases
turns yellow
C.
increases
turns blue
D.
increases
turns yellow
26.
The indicator bromothymol blue can be described by the following equilibrium
equation:
HInd + H2O ⇌ H3O+ +
Indyellow
blue
An acid is added to the above indicator. After equilibrium has been re-established,
how do the H3O+ and the colour of the solution compare with the original
equilibrium?
H3O+
Colour of Solution
A.
decreases
turns blue
B.
decreases
turns yellow
C.
increases
turns blue
D.
increases
turns yellow
27.
What is the Ka value for the indicator neutral red?
A.
1 x 10-14
B.
4 x 10-8
C.
7.4
D.
14.0
28.
What is a common source of SO2(g) ?
A.
automobiles
B.
a car battery
C.
a lead smelter
D.
oxidation of iron
29.
What is the [H3O+] in 0.70 M HCN ?
A.
0.70 M
B.
1.9 x 10-5 M
C.
1.0 x 10-6 M
D.
2.4 x 10-10 M
30.
What is the value of pKw for water at 25 0C ?
A.
1.0 x 10-14
B.
10 x 10-7
C.
7.00
D.
14.00
Acids
1.
Quiz #8 Buffers and Indicators
Consider the following acid solutions:
I. H2CO3II. HClO4
III. HF
Which of the above acids would form a buffer solution when its conjugate base is
A.
B.
C.
D.
2.
I only
II only
I and III only
I, II, and III only
Consider the following base indicator:
HInd ⇄ H+ + IndWhen the indicator is added to different solutions, the following data are obtained:
Solution
1.0 M HCl
1.0 M HA1
1.0 M HA2
Colour
Yellow
Blue
Yellow
The acids HAl, HA2, and HInd listed in the order of decreasing acid strength is
A.
HA2, HInd, HAl
B.
HInd, HAl, HA2
C.
HA2, HAl, HInd
D.
HAl, HInd, HA2
3.
Which of the following compounds, when added to a solution of ammonium nitrate, will
result in the formation of a buffer solution?
A.
Ammonia
B.
Nitric acid
C.
Sodium nitrate
D.
Ammonium chloride
4.
Which of the following represents a buffer equilibrium?
5.
A.
HI + H2O ⇄ H3O+ + I-
B.
HCl + H2O ⇄ H3O + Cl-
C.
HCN + H2O ⇄ H3O+ + CN-
D.
HClO4 + H2O ⇄ H3O + ClO4-
Consider the following equilibrium:
HF(aq) + H2O(l) ⇄ H3O+(aq) + F-(aq)
The above system will behave as a buffer when the [F-] is approximately equal to
A.
Ka
B.
[HF]
C.
[H2O]
D.
[H3O+]
6.
A basic buffer solution can be prepared by mixing equal numbers of moles of
A.
NH4Cl and HCl
B.
NaCl and NaOH
C.
Na2CO3 and NaHCO3
D.
NaCH3COO and CH3COOH
7.
Which of the following would be used to prepare an acidic buffer solution?
A.
HF and H3O+
B.
NaHS and H2S
C.
NH3 and NH4Cl
D.
HCl and NaCl
added?
8.
HCN(aq) + H2O
⇌
H3O+(aq)
+
CN-(aq)
1.0 M
0.0001 M
0.0001 M
Why is the solution above not considered to be a true buffer solution?
A.
excessive [HCN]
B.
excessive [H3O+]
C.
insufficient [CN-]
D.
insufficient [HCN]
9.
What could be added to 1.0 L of this solution in order for it to behave as a true
A.
1.0 mol HCl
B.
1.0 mol HCN
C.
1.0 mol H3O+
D.
1.0 mol NaCN
10.
Which of the following could typically be used to prepare a buffer solution?
A.
NaHS and H2S
B.
H2S and Na2S
C.
HNO3 and NaNO3
D.
HNO3 and NaNO2
11.
What happens to the pH of a buffer solution if a small amount of acid is added?
A.
The pH remains constant.
B.
The pH increases slightly.
C.
The pH decreases slightly.
D.
The pH decreases significantly.
Consider the following buffer equilibrium:
12.
HF
+ H2O ⇌ H3O+ + FWhat concentration would limit the buffering action if acid were added?
A.
[F-]
B.
[HF]
C.
[H2O]
D.
[H3O+]
13.
Consider the following buffer equilibrium:
HF
+ H2O ⇌ H3O+ + FWhat concentration would limit the buffering action if base were added?
A.
[F-]
B.
[HF]
C.
[H2O]
D.
[H3O+]
14.
Which term does the following statement best describe? A mixture of a weak
acid and its conjugate base, each with distinguishing colours.
A.
buffer
B.
titration
C.
indicator
D.
conjugate
15.
The HC2O4-(aq) ion will act as
A.
a base since Ka > Kb
B.
a base since Ka < Kb
C.
an acid since Ka > Kb
D.
an acid since Ka < Kb
buffer?
16.
What do a chemical indicator and a buffer solution typically both contain?
A.
a strong acid and its conjugate acid
B.
a strong acid and its conjugate base
C.
a weak acid and its conjugate acid
D.
a weak acid and its conjugate base
17.
What is the approximate pH and Ka at the transition point for phenol red?
A.
B.
C.
D.
pH
6.0
7.3
7.3
8.0
Ka
6.3
3.1
5.0
1.0
x
x
x
x
10-7
10-14
10-8
10-8
18.
Which of the following acids could not be present in a buffer solution?
A.
HF
B.
HNO2
C.
H2SO3
D.
HCl
19.
The indicator phenol red will be red in which of the following solutions?
A.
1.0 M HF
B.
1.0 M HBr
C.
1.0 M NH4Cl
D.
1.0 M NaHCO3
20.
Which of the following chemical indicators has a Ka = 1.6 x 10 -4 ?
A.
methyl orange
B.
phenolphthalein
C.
thymolphthalein
D.
bromcresol green
21.
Which of the following chemical indicators has a Ka = 7.9 x 10 -10 ?
A.
methyl orange
B.
phenolphthalein
C.
thymolphthalein
D.
bromcresol green
22.
What is the Ka for phenolphthalein?
A.
1 x 10-14
A.
8 x 10-10
A.
1 x 109
A.
1.0 x 10-7
23.
What volume of 0.250 M KOH is required to titrate 0.00230 mol of the weak
acid H2CO4 ?
A.
1.15 mL
B.
4.60 mL
C.
9.20 mL
D.
18.4 mL
24.
What is the net ionic equation for the titration of H3PO4(aq) with Sr(OH)2(aq) ?
A.
H+(aq) + OH- (aq)  H2O(l)
B.
6H+(aq) + 6OH-(aq)  6H2O(l)
C.
2H3PO4(aq) + 3Sr2+(aq) + 6OH-(aq)  Sr3(PO4)2(s) + 6H2O(l)
D.
6H+(aq) + 2PO43-(aq) + 3Sr2+(aq) + 6OH- (aq)  3Sr2+ + 2PO43-(aq) + 6H2O(l)
25.
Which of the following will have the smallest Ka value?
A.
HF
B.
H2CO3
C.
H3PO4
D.
CH3COOH
26.
Which of the following will have the smallest Kb value?
A.
IO3B.
NH3
C.
CND.
HPO42-
27.
What volume of 0.500 M NaOH is required to neutralize 25.0 mL 0.250 M HBr ?
A.
5.00 mL
B.
12.5 mL
C.
20.0 mL
D.
25.0 mL
28.
Which of the following equations correctly represents the reaction of a metallic
oxide with water?
A.
K2O + H2O  2KOH
B.
SO3 + H2O  H2SO4
C.
Na2O + H2O  Na2O2 + H2
D.
NaOH + HCl  NaCl + HOH
29.
Which of the following equations describes the dissociation of the salt,
ammonium chloride, in water?
A.
NH4Cl(s)  NH4Cl(aq)
B.
NH4+(aq) + Cl-(aq)  NH4Cl(s)
C.
NH4Cl(s)  NH4+(aq) + Cl-(aq)
D.
NH4+(aq) + H2O(l)  H3O+(aq) + NH3(aq)
30.
Which of the following salt solutions is acidic?
A.
NaBr
B.
AlCl3
C.
Li2CO3
D.
NaHCO3
Quiz #9 Titrations and Titration Curves
1.
Which of the following indicators would be used when titrating a weak acid with a strong base?
A.
B.
C.
D.
2.
Methyl red
Methyl violet
Indigo carmine
Phenolphthalein
Which of the following acid-base pairs would result in an equivalence point with pH greater than 7.0?
A.
B.
C.
D.
HCl and LiOH
HNO3 and NH3
HClO4 and NaOH
CH3COOH and KOH
3.
Which of the following standardized solutions should a chemist select when titrating a 25.00 mL sample of
0.1 M NH3, using methyl red as an indicator?
A.
B.
C.
D.
4.
0.114 M HCl
6.00 M HNO3
0.105 M NaOH
0.100 M CH3COOH
Consider the following 0.100 M solutions I. H2SO4 II. HCl
III. HF
The equivalence point is reached when 10.00 mL of 0.100 NaOH has been added to 10.00 mL of
solution(s)
A.
B.
C.
D.
II only
I and II only
II and III only
I, II and III
5.
1412108642-
Which pair of 0.10 M solutions would result in the above titration curve?
A.
B.
C.
D.
6.
A suitable indicator for the above titration is
A.
B.
C.
D.
7.
HF and KOH
HCl and NH3
H2S and NaOH
HNO3 and KOH
Methyl violet
Alizarin yellow
Thymolphthalein
Bromcresol green
The pH scale is
A.
B.
C.
D.
direct
inverse
logarithmic
exponential
8.
12108642-
Which of the following indicators should be used in the titration represented by the above titration curve?
A.
B.
C.
D.
9.
Which of the following indicators should be used when 1.0 M HNO 2 is titrated with NaOH(aq)?
A.
B.
C.
D.
10.
0.150 M NaOH
0.001 M NaOH
3.00 M NaOH
6.00 M NaOH
What volume of 0.250 M H2SO4 is required to neutralize 25.00 mL of 2.50 M KOH?
A.
B.
C.
D.
12.
Methyl red
Thymol blue
Methyl orange
Indigo carmine
Which of the following solutions should be used when titrating a 25.00 mL sample of CH 3COOH that is
approximately 0.1 M?
A.
B.
C.
D.
11.
Orange IV
Methyl red
Phenolphthalein
Alizarin yellow
125 mL
150 mL
250 mL
500mL
Which of the following pairs of substances form a buffer system for human blood?
A.
HCl and ClB.
NH3 and NH2C.
H2CO3 and HCO3D.
H2C6H5O7 and HC6H5O72-
Acids
Quiz #10
Review
Answers
1.
How many moles of Mg(OH)2 are required to neutralize 30.00 mL of 0.150 M
HCl?
A.
2.25 x 10-3 mol
B.
4.50 x 10-3 mol
C.
5.00 x 10-3 mol
D.
9.00 x 10-3 mol
2.
The approximate Ka for the indicator phenolphthalein is
A.
6 x 10-19
B.
8 x 10-10
C.
6 x 10-8
D.
2 x 10-2
3.
A new indicator, “B.C. Blue (HInd),” is red in bases and blue in acids. Describe the shift in equilibrium
and the resulting color change if 1.0 M HIO3 is added to a neutral, purple solution of this indicator:
HInd + H2O ⇄ H3O+ + IndA.
Equilibrium shifts left, and colour becomes red
B
Equilibrium shifts left, and colour becomes blue
C.
Equilibrium shifts right, and colour becomes red
D.
Equilibrium shifts right, and colour becomes blue
4.
Which one of the following combinations would act as an acid buffer?
A.
HCl and NaOH
B.
KOH and KBr
C.
NH3 and NH4Cl
D.
CH3COOH and NaCH3COO
5.
What is the pH at the transition point of an indicator if its Ka is 7.9 x 10-3?
A.
0.98
B.
2.10
C.
7.00
D.
11.90
6.
Which of the following curves best represents the titration of sodium hydroxide
with hydrochloric acid?
B
A.
.
pH
C.
pH
D.
pH
pH
7.
A student prepares a buffer by placing ammonium chloride in a solution of
ammonia. Equilibrium is established according to the equation:
NH3 + H2O ⇄ NH4+ + OHWhen a small amount of base is added to the buffer, the base reacts with
A.
NH3 and the pH decreases
B.
NH4+ and the pH decreases
C.
NH3 to keep the pH relatively constant
D.
NH4+ to keep the pH relatively constant
8.
At the equivalence point in a titration involving 1.0 M solutions, which of the
following combinations would have the lowest conductivity?
A.
Nitric acid and barium hydroxide
B.
Acetic acid and sodium hydroxide
C.
Sulphuric acid and barium hydroxide
D.
Hydrochloric acid and sodium hydroxide
9.
An indicator HInd produces a yellow colour in 0.1 M HCl solution and a red
colour in 0.1 M HCN solution. Therefore, the following equilibrium:
HCN + Ind- ⇄ HInd + CNA.
Products are favored and the stronger acid is HInd
B.
Products are favored and the stronger acid is HCN
C.
Reactants are favored and the stronger acid is HInd
D.
Reactants are favored and the stronger acid is HCN
10.
11.
The indicator methyl red is red in a solution of NaH2PO4. Which of the following
equations is consistent with this observation?
A.
H2PO4- + H2O ⇄ HPO42- + H3O+
B.
H2PO4- + H2O ⇄ H3PO4 + OH-
C.
HPO42- + H2O ⇄ PO43- + H3O+
D.
HPO42- + H2O ⇄ H2PO4- + OH-
Consider the following acid-base indicator equilibrium:
HInd(aq) + H2O(l) ⇄ H3O+(aq) ⇄ Ind-(aq)
Which of the following statements describes the conditions that exist in an
indicator equilibrium system at its transition point?
A.
[HInd] = [Ind-]
B.
[Ind-] = [H3O+]
C.
[HInd] = [H3O+]
D.
[H3O+] = [H2O]
12.
Which of the following titrations would have an equivalence point less that pH 7?
A.
NH3 and HCl
B.
NaOH and HNO3
C.
Ba(OH)2 and H2SO4
D.
KOH and CH3COOH
13.
Oxalic acid dihydrate is a pure, stable, crystalline substance. Which of the following describes one of its
uses in acid-base titrations?
A.
buffer
B.
primary standard
C.
indicator
D.
acid anhydride
14.
15.
What is the complete ionic equation that describes the reaction of HCl(aq) with
A.
H+(aq) + OH-(aq)  H2O(l)
B.
2HCl(aq) + Pb(OH)2(s)  PbCl2(s) + 2H2O(l)
C.
2H+(aq) + 2Cl-(aq) + Pb(OH)2(s)  PbCl2(s) + 2H2O (l)
D.
2H+(aq) + 2Cl-(aq) + Pb2+(aq) + 2OH- Pb2+(aq) + 2Cl-(aq) + 2H2O (l)
Pb(OH)2(s)?
What is the formula equation that describes the reaction of HCl(aq) with Pb(OH)2(s)?
A.
H+(aq) + OH-(aq)  H2O(l)
B.
2HCl(aq) + Pb(OH)2(s)  PbCl2(s) + 2H2O(l)
C.
2H+(aq) + 2Cl-(aq) + Pb(OH)2(s)  PbCl2(s) + 2H2O (l)
D.
2H+(aq) + 2Cl-(aq) + Pb2+(aq) + 2OH- Pb2+(aq) + 2Cl-(aq) + 2H2O (l)
16.
What term is used to describe the point at which a chemical indicator changes
colour?
A.
titration point
B.
transition point
C.
equivalence point
D.
neutralization point
17.
What term is used to describe the point in a titration where the acid has
completely reacted with the base.
A.
titration point
B.
transition point
C.
equivalence point
D.
neutralization point
18.
Which of the following is a piece of equipment typically used in acid-base
titrations?
A.
burette
B.
cuvette
C.
litmus paper
D.
graduated cylinder
19.
Identify an environmental problem associated with acid rain.
A.
increasing the pH of lakes
B.
the green house effect
C.
chemical decomposition of rainwater
D.
metals leeching from river rocks accumulating in lakes
20.
A buffer solution is prepared using sufficient amounts of H 2S and NaHS. What
limits this buffer’s effectiveness when NaOH is added?
A.
[H2S]
B.
[HS-]
C.
[OH-]
D.
[H3O+]
21.
Which of the following is not a good use for an acid-base titration curve?
A.
to determine the concentration of the acid
B.
to select a suitable indicator for the titration
C.
to determine whether the acid is strong or weak
D.
to select a suitable primary standard for the titration
22.
A substance which completely produces hydroxide ions in solution is a definition
of which of the following?
A.
a strong Arrhenius acid
B.
a strong Arrhenius base
C.
a weak Arrhenius base
D.
a strong Brønsted-Lowry base
23.
When a strong acid is titrated with a strong base, what will the pH value be at the
equivalence point?
A.
0.0
B.
5.0
C.
7.0
D.
9.0
24.
At a certain point in a strong acid-strong base titration, the moles of H+ are equal
to the moles of OH-. This is a definition of which of the following?
A.
the end point
B.
the titration point
C.
the transition point
D.
the equivalence point
25.
A 25.0 mL sample of H2SO4(aq) is titrated with 15.5 mL of 0.50 M NaOH(aq).
What is the concentration of the H2SO4(aq) ?
A.
0.078 M
B.
0.16 M
C.
0.31 M
D.
0.62 M
26.
When a weak acid is titrated with a strong base, what will the pH value be at the
equivalence point?
A.
0.0
B.
6.7
C.
7.0
D.
8.8
27.
A substance which completely produces accepts a proton from an acid is a
definition of which of the following?
A.
a strong Arrhenius acid
B.
a strong Arrhenius base
C.
a weak Arrhenius base
D.
a strong Brønsted-Lowry base
28.
Which of the following is a piece of equipment typically used in acid-base
titrations?
A.
pipette
B.
test tube
C.
litmus paper
D.
graduated cylinder
29.
Water has the greatest tendency to act as an acid with which of the following?
A.
NO3B.
NO2C.
H2PO4D.
CH3COO-
30.
Which net ionic equation best describes the reaction between NaOH and H2S?
A.
H+(aq) + OH-(aq)  H2O(l)
B.
H2S(aq) + 2OH-  2H2O(l) + S2-(aq)
C.
H2S(aq) + 2NaOH  2H2O(l) + Na2S(aq)
D.
2H+(aq) + S2-(aq) + 2NaOH(aq)  2H2O(l) + 2Na+(aq) + S2-(aq)
Web Review of Acids
1.
List five properties of:
a)
acids
b)
bases
2.
What ion is produced when an acid reacts with water? A base?
3.
Define:
Conjugate
Arhenius strong acid
Bronsted weak acid
Bronsted strong base
Ionization of water
Equivalence point
Transition point
Buffer
Hydrolysis.
4.
Identify the acids or bases in the following equation. Are the reactants or products
favoured?
HC03- + HF ⇄ H2CO3 + F-
5.
Classify each compound as a strong or weak acid or base; acidic or basic anhydride;
acidic, basic, or neutral salt; or buffer system. Write an equation to show how each reacts
with water.
NH3
AlCl3
H2CO3
HClO4
KCN
NH4Cl
KOH
SO2
NaF
HCl
NaI
K2O
NaOH
CO2
NH3 and NH4Cl
NaCH3COO and CH3COOH
6.
H+is short for _______.
7.
Determine the conjugates for each of the bases.
CN-
8.
NH3
F-
OH-
Co(H2O)5(OH)2+
Determine the conjugates for each of the acids.
HF
HCN
Al(H2O)63+
NH4+
HPO42-
9.
Describe a strong and weak acid as well as a strong and weak base in terms of each of the
following:
strong acid
weak acid
strong base
weak base
Conductivity
Size of Ka
Size of Kb
Degree of Ionization
pH.
10.
Why is the strongest acid in water H3O+? Explain!
11.
Why is the strongest base in water OH-? Explain!
12.
Which has the higher pH H2S03 or H3BO3? Explain!
13.
Which has the higher pH NaCN or NaF? Explain!
14.
A buffer has a pH of 9.00. 2 drops of a dilute strong acid are added. Estimate how the pH
changes?
15.
a) Complete the chart by indicating the pairs required to make buffer solutions. For
example HCN (weak acid) and NaCN (salt containing the conjugate of the weak acid)
will make a buffer solution. b) Write an equation the describes the equilibrium for each
buffer. c) Circle the formulas that have high concentrations.
Weak Acid or Base
Salt
HF
NaCH3COO
NH3
NaCN
H2CO3
KH2PO4
HCH3COO
16.
Match each equation with its type:
Acid/base formula equation
F- + HOH(l) ⇄ HF + OH-
Acid/base net ionic equation
HCl + NaOH →NaCl + HOH(l)
Solubility product ionization equation
H+ + OH- → HOH(l)
Hydrolysis of a weak acid
AgCl(s) ⇄ Ag+ + Cl-
Hydrolysis of a weak base
H20(l) ⇄ H+ + OH-
Ionization of water
NH4+ + H20 ⇄ NH3 + H3O+
17.
Write the equilibrium expressions for each of the above equations except for the second
and third reactions.
18.
A student tested the electrical conductivity of two acid solutions. One solution was a
strong acid and the other a weak acid. Both solutions had the same conductivity. Explain
how this could be possible.
19.
Describe in terms of hydrolysis how NaCH3COO can be added to potato chips in order to
produce the vinegar flavour.
20.
Describe what happens to the H+ and the OH- when the pH increases by 2 units.
21.
Describe two gases responsible for acid rain. Write equations to show how they react
with water. What gas naturally lowers the pH of normal rain?
22.
Complete the following reaction using a formula equation, complete ionic quation and
net ionic equation. H2C2O4 + NaOH →
23.
Write the equilibrium expression for phosphoric acid in water.
24.
a) An indicator HInd is red in acid and blue in base. Write the equation for the indicator and
explain the colours.
b) What is true at the transition point?
c)What is the color of this indicator in a solution of AlCl3?
d) In the above solution, what is larger [HInd] or [Ind-] ?
e) Calculate the Ka for Phenolphthalein.
f) What indicator has a Ka of approximately 1.0 X 10-10?
25.
Give an example of a monoprotic, diprotic, and triprotic acid. Write an equation for each
to show how they ionize in water.
26.
Give the approximate pH of the equivalence for each titration. Choose an appropriate
indicator.
Acid
Base
HCl
NaOH
pH of Equivalence Point Indicator
H2SO4 NH3
HF
KOH
27.
Which of the following will have the lowest pH?
HCLO HClO2 HClO3 HClO4
28.
Pick the formulae that are amphiprotic.
H2SO4 H20 F-
HCO3-
CO3-2 KOH H2PO4- HPO4-2
CALCULATIONS
1.
Calculate the quantities required to complete the table. In the first row write the general
equations for each quantity. Watch your significant figures.
[H+] =
[H+]
[OH-] =
pH =
[OH-]
pH
pOH =
pOH
Acid/base/neutral
-3
5.0 x 10 M
1.3 x 105
M
3.1
2.508
neutral (2sig
figs)
2.
What volume of 0.20 M H2SO4 is needed to neutralize 50.00 ml 0.30M NaOH?
3.
What mass of NaF is required to prepare 100.0 ml of 0.300 M solution?
4.
35.5 mL of 0.300 M NaOH is required to neutralize 10.0 mL of H2SO4. What is the acid
concentration?
5.
100.0 mL of .200 M HCl is mixed with 120.0 mL of 0.200M NaOH. Calculate the pH of
the resulting solution.
6.
Calculate the Ka for phenolphthalein.
7.
The Ksp of AgOH is 6.8 x 10-12. Calculate the pH.
8.
The OH- concentration in 0.10M NaCN is 2.7 x 10-6 M. Calculate the Kb from this
information only.
9.
Calculate the pH for 0.20 M HCl.
10.
Calculate the pH for 0.10M Ba(OH)2.
11.
Calculate the pH for 0.40 M HCN.
12.
Calculate the pH for 0.40 M Na2CO3.
13.
What is the pH for 0.30 M NaCl?
14.
Calculate the pH of 0.20 M NH3.
15.
Calculate the pH of 0.20 M NH4Cl.
16.
Show by calculation if H2PO4- is an acid or base (compare the Kb and Ka).
17.
Calculate the pH of a saturated solution of Mg(OH)2 if the Ksp is 1.2 X 10-11.
18.
A 0.50 M NH3 solution is found to have a OH- concentration of 1.86 x 10-3 M. Using this
data only calculate the Kb.
19.
A 0.18 M acid HX has a pH of 2.40. What is the Ka?
20.
The following data were recorded when 25.00 mL of H2SO4 were titrated with 0.1030 M
NaOH. The volumes of NaOH used in three runs were: 46.06 mL, 44.52 mL, 44.54 mL.
Calculate the acid concentration.
21.
The following data were recorded when 10.00ml of NaOH were titrated with 0.1030M
H2SO4. The volumes of H2SO4 used in three runs were: 12.55 mL, 12.55 mL, 12.10 mL.
Calculate the base concentration.
Acids Practice Test # 1
1.
An equation representing the reaction of a weak acid with water is
A.
HCl + H2O ⇄ H3O+ + Cl-
B.
NH3 + H2O ⇄ NH4+ + OH-
C.
HCO3- H2O ⇄ H2CO3 + OH-
D.
HCOOH + H2O ⇄ H3O+ + HCOO-
2.
The equilibrium expression for the ion product constant of water is
A.
Kw = [H3O+][OH-]
[H2O]
B.
Kw = [H3O+]2[O2]
C.
Kw = [H3O+][OH-]
D.
Kw = [H3O+]2[O2-]
3.
Consider the following graph for the titration of 0.1 M NH 3 with 1.0 M HCl.
pH
14
7
0
I
II
III
I
V
Volume HCl added
A buffer solution is present at point
A.
I
B.
II
C.
III
D.
IV
4.
Consider the following equilibrium system for an indicator: HInd + H 2O ⇄ H3O+ + IndWhich two species must be of two different colours in order to be used as an indicator?
A.
HInd and H2O
B.
HInd and IndC.
H3O+ and IndD.
Hind and H3O+
5.
Which of the following indicators is yellow at pH 10.0?
A.
methyl red
B.
phenol red
C.
thymol blue
D.
methyl violet
6.
A sample containing 1.20 x 10-2 mole HCl is completely neutralized by 100.0 mL of Sr(OH) 2. What is the
[Sr(OH)2]?
A.
6.00 x 10-3 M
B.
6.00 x 10-2 M
C.
1.20 x 10-1 M
D.
2.4 x 10-1 M
7.
Which of the following titrations will have the highest pH at the equivalence point?
A.
HBr with NH3
B.
HNO2 with KOH
C.
HCl with Na2CO3
D.
HNO3 with NaOH
8.
An Arrhenius acid is defined as a chemical species that
A.
is a proton donor.
B.
is a proton acceptor.
C.
produces hydrogen ions in solution.
D.
produces hydroxide ions in solution.
9.
Consider the following acid-base equilibrium system: HC2O4- + H2BO3- ⇄ H3BO3 + C2O42Identify the Bronsted-Lowry bases in this equilibrium.
A.
H2BO3- and
H3BO3
B.
HC2O4- and
H3BO3
C.
HC2O4- and
C2O42D.
H2BO3 and
C2O42-
10.
The equation representing the predominant reaction between NaCH3COO with water is
A.
CH3COO- + H2O ⇄ CH3COOH + OH-
B.
CH3COO- + H2O ⇄ H2O + CH2COO2-
C.
CH3COOH + H2O ⇄ H3O+ + CH3COO-
D.
CH3COOH + H2O ⇄ CH3COOH2+ + OH-
11.
Which of the following solutions will have the lowest electrical conductivity?
A.
0.10 M HF
B.
0.10 M NaF
C.
0.10 M H2SO3
D.
0.10 M NaHSO3
12.
Which of the following is the strongest Bronsted-Lowry base?
A.
NH3
B.
CO32C.
HSO3D.
H2BO3-
13.
A 1.0 x 10-4 M solution has a pH of 10.00. The solute is a
A.
weak acid
B.
weak base
C.
strong acid]
D.
strong base
14.
The ionization of water at room temperature is represented by
15.
A.
H2O ⇄ 2H+ + O2-
B.
2H2O ⇄ 2H2 + O2
C.
2H2O ⇄ H2 + 2OH-
D.
2H2O ⇄ H3O+ + OH-
Addition of HCl to water causes
A.
both [H3O+] and [OH-] to increase
B.
both [H3O+] and [OH-] to decrease
C.
[H3O+] to increase and [OH-] to decrease
D.
[H3O+] to decrease and [OH-] to increase
16.
Consider the following:
I.
H2SO4
II.
HSO4III.
SO42Which of the above is/are present in a reagent bottle labeled 1.0 M H 2SO4?
A.
I only
B.
I and II only
C.
II and III only
D.
I, II, and III
17.
The pH of 0.10 M KOH solution is
A.
0.10
B.
1.00
C.
13.00
D.
14.10
18.
An indicator changes colour in the pH range 9.0 to 11.0. What is the value of the Ka for the indicator?
A.
1 x 10-13
B.
1 x 10-10
C.
1 x 10-7
D.
1 x 10-1
19.
Which of the following are amphiprotic in aqueous solution?
I.
HBr
II.
H2O
III.
HCO3IV.
H2C6H5O7A
B.
C.
D.
I and II only
II and IV only
II, III, and IV only
I, II, III, and IV
20.
Which of the following always applies at the transition point for the indicator Hind?
A.
[Ind-] = [OH-]
B.
[HInd] = [Ind-]
C.
[Ind-] = [H3O+]
D.
[HInd] = [H3O+]
21.
Calculate the [H3O+] of a solution prepared by adding 10.0 mL of 2.0 M HCl to 10.0 mL of 1.0 M NaOH.
A.
0.20 M
B.
0.50 M
C.
1.0 M
D.
2.0 M
22.
Both acidic and basic solutions
A.
taste sour
B.
feel slippery
C.
conduct electricity
D.
turn blue litmus red
23.
The conjugate acid of HPO42- is
A.
PO43B.
H2PO4C.
H2PO42D.
H2PO43-
24.
What is the value of the Kw at 25 oC?
A.,
1.0 x 10-14
B.
1.0 x 10-7
C.
7
D.
14
25.
Consider the following equilibrium: 2H2O(l) ⇄ H3O+(aq) + OH-(aq)
A small amount of Fe(H2O)63+ is added to water and equilibrium is re-established. Which of the following
represents the changes in ion concentrations?
[H3O+]
[OH-]
A.
increases
increases
B.
increases
decreases
C.
decreases
decreases
D.
decreases
increases
26.
Consider the following equilibrium for an indicator: HInd + H 2O ⇄ H3O+ + IndIn a solution of pH of 6.8, the colour of bromthymol blue is
A.
blue because [HInd] = [Ind-]
B.
green because [HInd] = [Ind-]
C.
green because [HInd] < [Ind-]
D.
yellow because [HInd] > [Ind-]
27.
The indicator with Ka = 4 x 10-8 is
A.
neutral red
B.
methyl red
C.
indigo carmine
D.
phenolphthalein
28.
In a titration a 25.00 mL sample of Sr(OH)2 is completely neutralized by 28.60 mL of 0.100 M HCl. The
concentration of the Sr(OH)2 is
A.
1.43 x 10-3 M
B.
2.86 x 10-3 M
C.
5.72 x 10-2 M
D.
1.14 x 10-1 M
29.
A student mixes 15.0 mL of 0.100 M NaOH with 10.0 mL of 0.200 M HCl. The resulting solution is
A.
basic
B.
acidic
C.
neutral
D.
amphiprotic
30.
Which of the following salts will dissolve in water to produce a neutral solution?
A.
LiF
B.
CrCl3
C.
KNO3
D.
NH4Cl
31.
What is the value of the Kb for HC6H5O72-?
A.
5.9 x 10-10
B.
2.4 x 10-8
C.
4.1 x 10-7
D.
1.7 x 10-5
32.
The pOH of 0.015 M HCl solution is
A.
0.97
B.
1.80
C.
12.18
D.
13.03
33.
Which of the following will produce an acidic solution?
A.
NaCl
B.
NH4NO3
C.
Ca(NO3)2
D.
Ba(NO3)2
34.
Which of the following salts will dissolve in water to produce an acid solution?
A.
LiF
B.
CrCl3
C.
KNO3
D.
NaCl
35.
Which of the following salts will dissolve in water to produce a basic solution?
A.
LiF
B.
CrCl3
C.
KNO3
D.
NH4Cl
A student mixes 400 mL of 0.100 M NaOH with 100 mL of 0.200 M H 2SO4. The resulting solution has a
pH of
A.
14.000
B.
0.000
C.
13.800
D.
7.000
36.
37.
A student mixes 500 mL of 0.400 M NaOH with 500 mL of 0.100 M H 2SO4. The resulting solution has a
pH of
A.
14.000
B.
0.000
C.
13.000
D.
7.000
38.
The strongest acid in water is
A.
HClO4
B.
HI
C.
HF
D.
H3O+
39.
The formula that has the highest pH in water is
A.
HF
B.
H2CO3
C.
H2C2O4
D.
HCN
The formula that has the highest pH in water is
A.
NaF
B.
NaCl
C.
H2C2O4
D.
NaCN
39.
Subjective
1.
A chemist prepares a solution by dissolving the salt NaCN in water.
a) Write the equation for the dissociation reaction that occurs.
b) Write the equation for the hydrolysis reaction that occurs.
c) Calculate the value of the equilibrium constant for the hydrolysis
2.
A 3.50 x 10-3 M sample of unknown acid, HA has a pH of 2.90. Calculate the value of the Ka and identify
this acid.
3.
Calculate the mass of NaOH needed to prepare 2.0 L of a solution with a pH of 12.00.
4.
A 1.00 M OCl- solution has an [OH-] of 5.75 x 10-4 M. Calculate the Kb for OCl-.
5.
Calculate the pH of a solution prepared by adding 15.0 mL of 0.500 M H2SO4 to 35.0 mL of 0.750 M
NaOH.
6.
Determine the pH of a 0.10 M solution of hydrogen cyanide.
7.
Determine the pH of 0.100 M NH3.
8.
Determine the pH of a saturated solution of Mg(OH) 2.
Acids Practice Test # 2
1.
What colour would 1.0 M HCl be in an indicator mixture consisting of phenol red and thymolphthalein?
A
red
B
blue
C
yellow
D
colourless
2.
During a titration, what volume of 0.500 M KOH is necessary to completely neutralize 10.0 mL of 2.00 M
CH3COOH?
A
10.0 mL
B
20.0 mL
C
25.0 mL
D
40.0 mL
3.
Which indicator has a Ka = 1.0 x 10-6?
A
neutral red
B
thymol blue
C
thymolpthalein
D
chlorophenol red
4.
Acid is added to a buffer solution. When equilibrium is reestablished the buffering effect has resulted in
[H3O+]
A
increasing slightly
B
decreasing slightly
C
increasing considerably
D
decreasing considerably
5.
A buffer solution will form when 0.10 M NaF is mixed with an equal volume of
A
0.10 M HF
B
0.10 M HCl
C
0.10 M NaCl
D
0.10 M NaOH
6.
Which of the following statements applies to 1.0 M NH3(aq) but not to 1.0 M NaOH(aq)?
A
partially ionizes
B
neutralizes an acid
C
has a pH greater than 7
D
turns bromocresol green from yellow to blue
7.
In which of the following are the reactants favoured?
A
HNO2 + CN- ⇄ NO2- + HCN
B
H2S + HCO3- ⇄ HS- + H2CO3
C
H3PO4 + NH3 ⇄ H2PO4- + NH4+
D
CH3COOH + PO43-
⇄ CH3COO- + HPO42-
8.
What is the pOH of a solution prepared by adding 0.50 moles of NaOH to prepare 0.50 L of solution?
A
0.00
B
0.30
C
14.00
D
13.70
9.
What is the [H3O+] in a solution with a pH = 5.20?
A
1.4 x 10-14
B
1.6 x 10-9
C
6.3 x 10-6
D
7.1 x 10-1
10.
Consider the following equilibrium: 2H2O(l) + energy ⇄ H3O+(aq) + OH-(aq)
What will cause the pH to increase and the Kw to decrease?
A
adding a strong acid
B
adding a strong base
C
increasing the temperature
D
decreasing the temperature
11.
The complete neutralization of 15.0 mL of KOH requires 0.0250 moles H 2SO4. The [KOH] was
A
1.50 M
B
1.67 M
C
3.33 M
D
6.67 M
12.
What is the [H3O+] at the equivalence point for the titration between HBr and KOH?
A
1.0 x 10-9 M
B
1.0 x 10-7 M
C
1.0 x 10-5 M
D
0.0 M
13.
Which of the following would form a buffer solution when equal moles are mixed together?
A
HCl and NaCl
B
HCN and NaCN
C
KNO3 and KOH
D
Na2SO4 and NaOH
14.
Which of the following acids has the weakest conjugate base?
A
HIO3
B
HNO2
C
H3PO4
D
CH3COOH
15.
When 10.0 ml of 0.10 M HCl is added to 10.0 mL of water, the concentration of H 3O+ in the final solution
is
A
0.010 M
B
0.050 M
C
0.10 M
D
0.20 M
16.
The conjugate base of an acid is produced by
A
adding a proton to the acid
B
adding an electron to the acid
C
removing a proton from the acid
D
removing an electron from the acid
17.
Which of the following represents the predominant reaction between HCO3- and water?
A
2HCO3- ⇄ H2O + 2CO2
B
HCO3- + H2O ⇄ H2CO3
C
HCO3-
D
2HCO3- + H2O ⇄ H3O+ + CO32-
+ H2O ⇄
+ OH-
H3O+ + CO32+
OH- + CO2
18.
Water acts as an acid when it reacts with
I
CNII
NH3
III
HClO4
IV
CH3COOA
I and IV only
B
II and III only
C
I, II, and IV
D
II, III, and IV
19.
In a solution of 0.10 M H2SO4, the ions present in order of decreasing concentration are
A
[H3O+] > [HSO4-] > [SO42-] > [OH-]
B
[H3O+] > [SO42-] > [HSO4-] > [OH-]
C
[OH-] > [SO42-] > [HSO4-] > [H3O+]
D
[SO42-] > [HSO4-] > [OH-] > [H3O+]
20.
Which of the following will dissolve in water to produce an acidic solution?
A
CO2
B
CaO
C
MgO
D
Na2O
21.
Which of the following solutions will have a pH = 1.00?
I
0.10 M HCl
II
0.10 M HNO2
III
0.10 M NaOH
A
I only
B
II only
C
I and II only
D
I, II, and III
22.
Ka for the acid H2AsO4- is 5.6 x 10-8. What is the value of the Kb for HAsO42-?
A
5.6 x 10-22
B
3.2 x 10-14
C
1.8 x 10-7
B
2.4 x 10-4
23.
In a titration, which of the following has a pH = 7.00 at the equivalence point?
A
NH3 and HNO3
B
KOH and HCl
C
NaF and HCl
D
Ca(OH)2 and CH3COOH
24.
Which of the following salts dissolves to produce a basic solution?
A
KCl
B
NH4Br
C
Fe(NO3)3
D
LiCH3COO
25.
Calculate the pH in a 0.200 M solution of Sr(OH)2.
A
1.40
B
1.70
C
13.30
D
13.60
26.
Which of the following solutions has a pH less than 7.00?
A
NaCl
B
LiOH
C
NH4NO3
D
KCH3COO
27.
Which of the following will form a basic aqueous solution?
A
HSO3B
HSO4C
HPO42D
HC2O4-
28.
What is the approximate Ka value for the indicator chlorophenol red?
A
1 x 10-14
B
1 x 10-8
C
1 x 10-6
D
1 x 10-3
29.
What is the approximate pH of the solution formed when 0.040 mol NaOH is added to 2.00 L of 0.020 M
HCl?
A
0.00
B
1.40
C
1.70
D
7.00
30.
In which one of the following equations are the Bronsted acids and bases all correctly identified?
Acid
+
Base
⇄
Base
2-
-
+
Acid
A
H2O2
SO3
⇄
HO2
HSO3-
B
H2O2
SO32-
⇄
HSO3-
HO2-
C
SO32-
H2O2
⇄
HO2-
HSO3-
D
SO32-
H2O2
⇄
HSO3-
HO2-
31.
Which of the following titrations will always have an equivalence point at a
pH > 7.00?
A
weak acid with a weak base
B
strong acid with a weak base
C
weak acid with a strong base
D
strong acid with a strong base
32.
A buffer solution may contain equal moles of
A
weak acid and strong base
B
strong acid and strong base
C
weak acid and its conjugate base
D
strong acid and its conjugate base
33.
A gas which is produced by burning coal and also contributes to the formation of acid rain is
A
H2
B
O3
C
SO2
D
C3H8
34.
Which of the following 1.0 M salt solutions is acidic?
A
BaS
B
NH4Cl
C
Ca(NO3)2
D
NaCH3COO
35.
Which of the following statements applies to 1.0 M NH3(aq) but not to 1.0 M NaOH(aq)?
A
partially ionizes
B
neutralizes and acid
C
has a pH greater than 7
D
turns bromcresol green from yellow to blue
36.
When the indicator thymol blue is added to 0.10 M solution of an unknown acid, the solution is red. The
acid could be
A
HF
B
H2S
C
HCN
D
HNO3
Subjective
1.
Calculate the pH of the solution prepared by mixing 15.0 mL of 0.50 M HCl with 35.0 mL 0.50 M NaOH.
2. Calculate the [OH-] in 0.50 M NH3(aq).
3.
A titration was performed by adding 0.175 M H2C2O4 to a 25.00 mL sample of NaOH.
The following data was collected.
Trial 1
Trial 2
Final volume of H2C2O4 from burette (mL)
23.00
39.05
20.95
Initial volume of H2C2O4 from burette (mL)
4.85
23.00
5.00
Calculate the [NaOH]
Trial 3
4.
A 250.0 mL sample of HCl with a pH of 2.000 is completely neutralized with 0.200 M NaOH. What
volume of NaOH is required to reach the stoichiometric point.
5.
If the HCl were titrated with 0.200 M NH3(aq) instead of 0.200 M NaOH, how would the volume of base
required to reach the equivalence point compare with the volume calculated in the last question? Explain
your answer.
6.
a)
Consider the following salt ammonium acetate, NH4CH3COO.
Write the equation for the dissociation of NH4CH3COO.
b)
Write the equations for the hydrolysis reactions that occur.
c)
Explain why a solution of NH4CH3COO has a pH =7.00. Support your answer with a calculation.
7.
Consider the following equilibrium:
energy + 2H2O ⇄ H3O+
a) Explain how pure water can have a pH = 7.30.
b) Calculate the value of the Kw for the sample of water with a pH = 7.30.
+
OH-