Acids, Bases and Salts Unit Plan
Period/Topic
Worksheets
1. Properties of Acids, Bases & Salts
WS 1
2. Arrhenius, Bronsted Acids, Ka and Strength.
WS 2
3. Arrhenius, Bronsted Bases, Kb and Strength
WS 3
4. Acid & Base Reactions. Amphiprotic. Acid Chart. WS 4
5. Leveling effect, Anhydrides and Relationships.
WS 5
6. Hydrolysis of Salts. Quiz.
WS 6
7. Acid, Base & Salt Reactions. Hydrolysis.
WS 7
8. Yamada’s Indicator Lab. Hydrolysis.
WS 8
9. Ionization of Water, [H+] & [OH-], pH scale.
WS 9
10. pH Calculations for Weak Acids.
WS 10
11. Ka from pH for Weak Acids.
WS 11
Quiz
1
2
3
4
5
12. Indicators Lab.
13. Kbs from Kas for Weak Bases.
WS 12
6
14. pH for Weak Bases pH [H+] [OH-] Relationships.
WS 13
15. Amphiprotic Ions- Kas and Kbs.
WS 14
16. Titration Lab. Primary Standards.
Acids Midterm Practice Test
7
17. Titration Lab
18. Buffers & Indicators
WS 15
8
19. Titration Curves.
WS 16
9/10
20. Review # 1
Web Site Review
Practice Test 1
21. Review # 2
Practice Test 2
22. Test
Worksheet # 1
1.
Properties of Acids and Bases
Add 1 drop of each solution to 1 drop of the acid-base indicator in a spot plate. Record the colour
in the data table below. Describe each solution as an acid or base in the space provided. Write the
acid colour and base colour in the table below.
Indicator
Phenolphthalein
Litmus
BromothymolBlue
Acid/Base
Solution
HCl
NaOH
Vinegar
Ammonia
(NH3)
Lemon Juice
Seven-up
Baking Soda
(NaHCO3)
Indicator
Acid Colour
Base Colour
Phenolphthalein
Litmus
Bromothymol Blue
Wash and dry your spot plate before going on to step 2.
2.
Wear safety goggles for this experiment. Pour approximately 50 mL of 1 M HCl into a fleaker.
Add one level spoonful of Ca and cover with a plastic funnel. After 1 minute and not before light
the top of the funnel using a match. Write the equation for the reaction below.
Wash and dry your fleaker before going on to step 3.
3.
Taste a lemon and describe the taste in one word
4.
Taste some baking soda and describe the taste in one word.
6.
Test two drops of HCl for conductivity in a spot plate. Result:
Write an equation that accounts for the conductivity of HCl.
7.
Test two drops of NaOH for conductivity in a spot plate. Result:
Write an equation that accounts for the conductivity of NaOH (dissociation).
Clean, dry and put away the spot plate
8.
List five properties of acids that are in your textbook.
9.
List five properties of bases that are in your textbook.
10.
Make some notes on the commercial acids: HCl and H2SO4 .
HCl
H2SO4
11.
Make some notes on the commercial base NaOH.
12.
Describe the difference between a concentrated and dilute acid (hint: concentration refers to the
molarity). Describe their relative conductivities.
13.
Describe the difference between a strong and weak acid. Use two examples and write equations to
support your answer. Describe their relative conductivities.
14.
Describe a situation where a strong acid would have the same conductivity as a weak acid (hint:
think about concentration).
Worksheet # 2
Conjugate Acid-Base Pairs
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.
[H3O+][CN-]
[HCN]
10.
Ka =
14. Ka =
11.
Ka =
15. Ka =
12.
Ka =
16. 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
14.
16.
18.
20.
22.
Conjugate Base
HNO2
HSO3H2O2
HSH2O
Base
15.
17.
19.
21.
23.
Conjugate Acid
HCOOIO3NH3
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
28.
What would you rather drink vinegar or hydrochloric acid? Explain.
H 2S
H3PO4
Making a Universal Indicator Lab Activity
Mix the following indicators in a 50 mL beaker. Stir with an eyedropper.
H2CO3
HF HSO4
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.
Results
Compound
HCl
CH3COOH
NaOH
NH3
NaHCO3
H2CO3
NaCl
Conductivity
pH
Classification
Worksheet # 3
Conjugate Acid-Base Pairs
Complete each reaction. Label each reactant and product as an acid or base.
1.
HCN
+
H2 O
Acid
⇄
Base
HCl
+
H2O
⇄
3.
HF
+
H2O
⇄
4.
F-
+
H2O
⇄
5.
HSO4(acid)
+
H2O
⇄
6.
NH4+
+
H2O
⇄
7.
HPO42(base)
+
H2O
⇄
Acid
Conjugate Base
8.
10.
12.
14.
HCO3HPO42H2O
HS-
16.
Circle the strong bases.
9.
11.
13.
15.
NaOH
Sr(OH)2
CH3COOIO3NH2C2H5SO73-
CsOH
Ba(OH)2
CNBase
Conjugate Acid
CH3COOH
KOH
Ca(OH)2
Rank the following acids from strongest to weakest.
H2S
18.
Base
CO32-
Fe(OH)3
Zn(OH)2
+
Acid
2.
17.
H3 O+
CH3COOH
H2PO4-
HI
HCl
Rank the following bases from the strongest to weakest.
H2O
F-
NH3
SO32-
HSO3-
NaOH
HF
19.
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 # 4
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
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
16.
H2SO4 + NaOH
⇄
Ka=
⇄
Kb=
→
acid or
17.
Define conjugate pairs.
18.
Give conjugate acids for: HS-, NH3, HPO4-2, OH-, H2O, NH3,
19.
Give conjugate bases for: NH4+,
Worksheet # 5
HF,
H2PO4-,
H3O+, OH-,
CO3-2
HCO3-,
H2 O
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
Ka is very large
Ka=8.0x10-5
HClO3
HClO
Ka=1.2x10-2
Ka=4.4x10-8
5.
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
18.
NH3
19.
20.
NaOH
HF
21.
H3PO4
Classification
Worksheet # 6
Reaction
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 # 7
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 # 8
Yamada’s Indicator Activity
Acid, Base and Salt Lab
Purpose:
1)
To use Yamada’s Indicator to determine the pH of various acids, bases and salts.
2)
3)
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)
3)
4)
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.
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.
Data
1.
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 # 9
pH and pOH Calculations
Complete the chart:
1.
2.
[H+]
7.00 x 10-3 M
[OH-]
8.75 x 10-2 M
pH
pOH
Acid/base/neutral
3.
4.
5.
6.
7.
8.
9.
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 # 10
1.
pH Calculations for Weak Acids
Calculate the [H+], [OH-], pH, and pOH for 0.20 M HCN.
[H+] =
[OH-] =
pH =
pOH =
2.
Calculate the [H+], [OH-], pH, and pOH for 2.20 M HF.
3.
[H+] =
[OH-] =
pH =
pOH =
+
Calculate the [H ], [OH ], pH, and pOH for 0.805 M CH3COOH.
[H+] =
4.
[OH-] =
pH =
Calculate the [H+], [OH-], pH, and pOH for 1.65 M H3BO3.
pOH =
[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 # 11
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 # 12
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-)
2.
=
Kw
=
1.0 x 10-14
Ka(HNO2)
4.6 x 10-4
KCH3COO (the basic ion is CH3COO-)
=
2.2 x 10-11
3.
NaHCO3
4.
NH3
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+] =
12.
[OH-] =
pH =
pOH =
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 # 13
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.
[OH-] =
pH =
pOH =
Calculate the [H+], [OH-], pH, and pOH for 0.25 M H2CO3.
[H+] =
[OH-] =
pH =
pOH =
9.
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
Volume of HCl added
37.80 mL
#2
#3
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.
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.
15.
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 # 14
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:
HCl + NaOH → NaCl + HOH
F- + HOH → HF + OHH+ + OH- → HOH
AgCl(s) → Ag+ + ClH20 → H+ + OHH+ + Cl- + Na+ + OH- → Na++ Cl- + H2O
Acid/base complete
Acid/base net ionic
Solubility product
Hydrolysis
Acid/Base formula
Ionization of water
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
9.
|________________________|_________________________|__
0
7
14
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.
[H2SO4] = 0.600 M
29.
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) +
Worksheet # 15
⇄
O2(g)
2SO3(g)
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 # 16
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
7.
Calculate the Ka of bromothymol blue.
+
+
→
KOH
→
NH3
→
+
→
LiOH
Indicator
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
14.
Adding 100 ml 0.10 M HCl
50 mL 0.10M NH3
Volume of base added
15.
Adding 100 ml .10 M HCl
to 50 mL 0.10 M NaOH
to
pH
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
3.
A.
red and the products are favoured
B.
red and the reactants are favoured
C.
yellow and the products are favoured
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.
4.
10.0 mL
25.0 mL
50.0 mL
100.0 mL
The pOH of 0.050 M HCl is
A.
B.
C.
D.
0.050
1.30
12.70
13.70
5.
The volume of 0.450 M HCl needed to neutralize 40.0 mL of 0.450 M Sr(OH)2 is
6.
A.
18.0 mL
B.
20.0 mL
C.
40.0 mL
D.
80.0 mL
Consider the following
I
H3PO4
II
H2PO4-
III
HPO42-
IV
Which of the above solutions are amphiprotic?
A.
B.
C.
D.
7.
Which of the following solutions will have the largest [H3O+]?
A.
B.
C.
D.
8.
I and II only
II and III only
I, II, and III only
II, III, and IV only
1.0 M HNO2
1.0 M HBO3
1.0 M H2C2O4
1.0 M HCOOH
Consider the following:
H2O + 57 kJ  H3O+ + OH-
When the temperature of the system is increased, the equilibrium shifts
9.
A.
left and the Kw increases
B.
left and the Kw decreases
C.
right and the Kw increases
D.
right and the Kw decreases
Normal rainwater has a pH of approximately 6 as a result of dissolved
A.
B.
C.
D.
10.
oxygen
carbon dioxide
sulphur dioxide
nitrogen dioxide
A 1.0 M solution of sodium dihydrogen phosphate is
A.
B.
C.
D.
acidic and the pH < 7.00
acidic and the pH > 7.00
basic and the pH < 7.00
basic and the pH > 7.00
PO43-
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.
12.
A Bronsted-Lowry base is defined as a chemical species that
A.
B.
C.
D.
13.
accepts protons
neutralizes acids
donated electrons
produces hydroxides ions in solution
Which of the following solutions will have the greatest electrical conductivity?
A.
B.
C.
D.
14.
red as [HInd] < [Ind-]
red as [HInd] > [Ind-]
yellow as [HInd] < [Ind-]
yellow as [HInd] > [Ind-]
1.0 M HCN
1.0 M H2SO4
1.0 M H3PO4
1.0 M CH3COOH
Consider the following equilibrium: HC6H5O72- + HIO3  H2C6H5O7- + IO3The order of Bronsted-Lowry acids and bases is
A.
B.
C.
D.
15.
acid, base, acid, base
acid, base, base, acid
base, acid, acid, base
base, acid, base acid
Consider the following: H2O(l)  H+ + OHWhen a small amount of 1.0 M KOH is added to the above system, the equilibrium
A.
B.
C.
D.
shifts left and [H+] decreases
shifts left and [H+] increases
shifts right and [H+] decreases
shifts right and [H+] increases
16.
Which of the following has the highest pH?
A.
B.
C.
D.
17.
In a 100.0 mL sample of 0.0800 M NaOH the [H3O+] is
A.
B.
C.
D.
18.
1.0 M NaIO3
1.0 M Na2CO3
1.0 M Na3PO4
1.0 M Na2SO4
1.25
1.25
8.00
8.00
x
x
x
x
10-13 M
10-12 M
10-3 M
10-2 M
Consider the following:
I
ammonium nitrate
II
calcium nitrate
III
iron III nitrate
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- + NO2Equilibrium would favour the
A.
B.
C.
D.
20.
The net ionic equation for the hydrolysis of Na2CO3 is
A.
B.
C.
D.
21.
the products since HSO4- is a weaker acid than HNO2
the reactants since HSO4- is a weaker acid than HNO2
the products since HSO4- is a stronger acid than HNO2
the reactants since HSO4- is a stronger acid than HNO2
H2O
H2O
H2O
H2O
+
+
+
+
Na+  NaOH + H+
2Na+  Na2O + 2H+
CO32-  H2CO3 + O2CO32-  HCO3- + OH-
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.
B.
[H3O+] has increased and the [OH-] has decreased
[H3O+] has increased and the [OH-] has increased
C.
D.
22.
A basic solution
A.
B.
C.
D.
23.
KOH → KSO4 + H2O
KOH → K2SO4 + H2O
2KOH → K2SO4 + H2O
2KOH → K2SO4 + 2H2O
donates protons
donates electrons
produces H+ in solution
produces OH- in solution
acid, base, acid, base.
acid, base, base, acid
base, acid, acid, base
base, acid, base, acid
The equation representing the reaction of ethanoic acid with water is
A.
B.
C.
D.
27.
+
+
+
+
Consider the following equilibrium: HS- + H3PO4  H2S + H2PO4The order of Bronsted-Lowry acids and bases is
A.
B.
C.
D.
26.
H2SO4
H2SO4
H2SO4
H2SO4
An Arrhenius base is defined as a substance which
A.
B.
C.
D.
25.
tastes sour
feels slippery
does not conduct electricity
reacts with metals to release oxygen gas
The balanced formula equation for the neutralization of H2SO4 by KOH is
A.
B.
C.
D.
24.
[H3O+] has decreased and the [OH-] has increased
[H3O+] has decreased and the [OH-] has decreased
CH3COO- + H2O  CH3COOH + OHCH3COO- + H2O  CH3COO2- + H3O+
CH3COOH + H2O  CH3COO- + H3O+
CH3COOH + H2O  CH3COOH2+ + OH-
Consider the following equilibrium: 2H2O + 57kJ  H3O+ + OHWhen the temperature is decreased, the water
A.
B.
C.
D.
28.
stays neutral and the [H3O+] increases
stays neutral and the [H3O+] decreases
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.
C.
D.
29.
The conjugate acid of C6H50- is
A.
B.
C.
D.
30.
a lower pH than a solution of 1.0 M HCl
a higher pOH than a solution of 1.0 M HCl
a higher [OH-] than a solution of 1.0 M HCl
a higher [H3O+] than a solution of 1.0 M HCl
Which of the following is the weakest acid
A.
B.
C.
D.
33.
1.0 M HCl
1.0 M HNO2
1.0 M H3BO3
1.0 M HCOOH
A solution of 1.0 M HF has
A.
B.
C.
D.
32.
C6H4OC6H5OH
C6H4O2C6H5OH+
Which of the following solutions will have the greatest electrical conductivity?
A.
B.
C.
D.
31.
Cl2O + H2O  2ClO + H2
Cl2O + H2O  Cl2 + H2O2
Cl2O + H2O  Cl2 + O2 + H2
HIO3
HCN
HNO3
C6H5COOH
Considering the following data
H3AsO4
Ka = 5.0 x 10-5
H2AsO4
Ka = 8.0 x 10-8
HAsO42Ka = 6.0 x 10-10
The Kb value for H2AsO4- is
A.
B.
C.
D.
34.
2.0
8.0
1.2
1.7
x
x
x
x
10-10
10-8
10-7
10-5
In a solution at 25oC, the [H3O+] is 3.5 x 10-6 M. The [OH-] is
A.
B.
C.
D.
3.5
2.9
1.0
3.5
x
x
x
x
10-20 M
10-9 M
10-7 M
10-6 M
35.
In a solution with a pOH of 4.22, the [OH-] is
A.
B.
C.
D.
1.7
6.0
6.3
1.7
10-10 M
10-5 M
10-1 M
104 M
x
x
x
x
36.
An aqueous solution of NH4CN is
37.
A.
basic because Ka < Kb
B.
basic because Ka > Kb
C.
acidic because Ka < Kb
D.
acidic because Ka > Kb
The net ionic equation for the predominant hydrolysis reaction of KHSO4 is
A.
B.
C.
D.
38.
The [OH-] in an aqueous solution always equals
A.
B.
C.
D.
39.
41.
5.58
1.79
5.58
5.97
x
x
x
x
10-15 M
10-14 M
10-1 M
10-1 M
The equilibrium expression for the hydrolysis reaction of 1.0 M K2HPO4 is
A.
[H2PO4-][OH-]
[HPO42-]
B.
[H3PO4][OH-]
[H2PO4-]
C.
[K+] [KHPO4-]
[K2HPO4]
D.
[K+]2 [HPO42-]
[K2HPO4]
The solution with the highest pH is
A.
B.
C.
D.
42.
Kw x [H3O+]
Kw - [H3O+]
Kw/[H3O+]
[H3O+]/Kw
The [H3O+] in a solution with pOH of 0.253 is
A.
B.
C.
D.
40.
HSO4- + H2O  SO42- + H3O+
HSO4- + H2O  H2SO4 + OHKHSO4 + H2O  K+ + SO42- + H3O+
KHSO4 + H2O  K+ + H2SO4 + OH-
1.0 M NaCl
1.0 M NaCN
1.0 M NaIO3
1.0 M Na2SO4
The pH of 100.0 mL of 0.0050 M NaOH is
A.
B.
C.
D.
43.
2.30
3.30
10.70
11.70
Consider the following equilibrium for an indicator: HInd + H2O  Ind- + H3O+
At the transition point,
A.
B.
C.
D.
[HInd]
[HInd]
[HInd]
[HInd]
>
=
<
=
[Ind-]
[Ind-]
[Ind-]
[H3O+]
Acids Unit Midterm Practice Test
1.
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.
a)
Write the balanced equation representing the reaction of HF with H2O.
b)
Identify the Bronsted-Lowry bases in the above equation.
7.
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 [H3O+] in 0.550 M C6H5COOH.
Quiz #1
1.
Calculate the Kb for HCO3-.
Properties of Acids, Bases, Salts, Arrhenius Bronsted Acids, Ka, Strength
Drano®, a commercial product used to clean drains, contains small bits of aluminum metal and
A.
B.
C.
ammonia
acetic acid
hydrochloric acid
D.
2.
3.
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)
Which equation represents a neutralization reaction?
A.
B.
C.
D.
4.
Pb2+(aq) + 2Cl-(aq) → PbCl2(s)
HCl(aq) + NH3(aq) → NH4Cl(aq)
BaI2(aq) + MgSO4(aq) → BaSO4(s) + MgI2(aq)
MnO4-(aq) + 5Fe2+(aq) +8H+(aq) → Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
An Arrhenius acid is a substance that
A.
B.
C.
D.
5.
sodium hydroxide
accepts a proton
donates a proton
produces H+ in solution
produces OH- in solution
Consider the following data table:
Breaker
Volume
Contents
1
15 mL
0.1 M Sr(OH)2
2
20 mL
0.2 M NH4OH
3
25 mL
0.1 M KOH
4
50 mL
0.2 M NaOH
Identify the beaker that requires the smallest volume of 1.0 M HCl for complete neutralization
6.
A.
Beaker 1
B.
Beaker 2
C.
Beaker 3
D.
Beaker 4
The net ionic equation for the titration of HClO4(aq) with LiOH(aq) is
A.
B.
C.
D.
7.
H+(aq) + OH-(aq) → H2O(l)
HClO4(aq) + OH-(aq) → ClO4-(aq) + H2O(l)
HClO4(aq) + LiOH(aq) → LiClO4(aq) + H2O(l)
H+(aq) + ClO4-(aq) + Li+(aq) + OH-(aq) → LiClO4(aq) + H2O(l)
The equilibrium constant expression for a sulphurous acid is
A
Ka = [H+][HSO3-]
B.
Ka = [H+][HSO3-]
[H2SO3]
C.
Ka = [2H+][SO32-]
[H2SO3]
Ka = [H+][SO32-]
[H2SO3]
D.
8.
To distinguish between a strong acid and a strong base, an experimenter could use
A.
B.
C.
D.
9.
10.
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
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.
B.
C.
D.
11.
fertilizers
beverages
toothpaste
oven cleaners
Which of the following is the strongest acid?
A.
B.
C.
D.
13.
a piece of copper.
a piece of magnesium
phenolphthalein indicator
a piece or red litmus paper
Caustic soda, NaOH, is found in
A.
B.
C.
D.
12.
odor
magnesium
a conductivity test
the common ion test
Acetic acid
Oxalic acid
Benzoic acid
Carbonic acid
The acid used in the lead-acid storage battery is
A.
B.
HCl
HNO3
C.
D.
Quiz #2
1.
Conjugates, Amphiprotic, Arrhenius, Bronsted Bases, Kb, & Strength
A test that could be safely used to distinguish a strong base from a weak base is
A.
B.
C.
D.
2.
H2SO4
CH3COOH
taste
touch
litmus paper
electrical conductivity
Identify the two substances that act as Bronsted-Lowry bases in the equation
HS- + SO42- ⇄ S2- + HSO4A.
B.
C.
D.
3.
The conjugate acid of H2C6H5O-7 is
A.
B.
C.
D.
4.
6.
C6H5O73HC6H5O72H2C6H5O7
H3C6H5O7
Which one of the following substances is/are amphiprotic?
(1) H3PO4
(2) H2PO4(3) HPO42A.
B.
C.
D.
5.
HS- and S2SO4 2-and S2HS- and HSO4SO42- and HSO4-
2 only
3 only
1 and 2
2 and 3
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) + 2Br -(aq) + Ca2+(aq) + 2OH-(aq) ⇄ 2H2O(l) + Ca2+(aq) + 2Br -(aq)
If reactants are favored in the following equilibrium, the stronger base must be
HCN + HS - ⇄ H2S + CN A.
B.
C.
H2S
HSCN-
D.
7.
The hydronium ion, H3O+ is a water molecule that has
A.
B.
C.
D.
8.
9.
HCN
lost a proton
gained a proton
gained a neutron
gained an electron
The complete ionic equation for the neutralization of acetic acid by sodium hydroxide is
A.
H+ + OH- ⇄ H2O
B.
CH3COOH + OH- ⇄ CH3COO- + 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.
B.
C.
D.
10.
NH4+
H2O
NH3
H3O+
Consider the following equilibrium system:
OCl-(aq) + HC7H5O2(aq) ⇄ HOCl(aq) + C7H5O2-(aq)
Keq= 2.1 x 103
At Equilibrium
A.
B.
C.
D.
11.
products are favored and HOCl is the stronger acid
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) ⇄ H2CO3(aq) + HBO32-(aq)
The two species acting as Bronsted-Lowry acids are
A.
B.
HCO3- and H2CO3
H2BO3- and H2CO3
C.
D.
12.
HCO3- and HBO32H2BO3- and HBO32-
Consider the following equilibrium HS- + H2C2O4 ⇄ HC2O4- + H2S
The stronger acid is
A.
B.
C.
D.
Quiz #3
HSH2C2O4
HC2O4H2S
Leveling effect, Anhydrides, Hydrolysis, Relationships
1.
Which of the following oxides will form the most acidic solution?
2.
A.
SO2
B.
MgO
C.
Na2O
D.
Al2O3
Which one of the following salts will produce an acidic solution?
A.
B.
C.
D.
3.
4.
The balanced equation for the reaction between sodium oxide and water is
A.
B.
C.
Na2O + H2O → 2NaOH
Na2O + H2O → 2NaH + O2
Na2O + H2O → 2Na + H2O2
D.
Na2O + H2O → 2Na + H2 +O2
‘Normal’ rainwater is slightly acidic due to the presence of dissolved
A.
B.
C.
D.
5.
methane
carbon dioxide
sulphur dioxide
nitrogen dioxide
Which of the following oxides would hydrolyze to produce hydroxide ions?
A.
B.
C.
D.
6.
KBr
LiCN
NH4Cl
NaCH3COO
NO
SO2
Cl2O
Na2O
The approximate pH of “normal” rainwater is
A.
0
B.
C.
D.
7.
Which of the following oxides would hydrolyze to produce hydronium ions?
A.
B.
C.
D.
8.
CaO
SO2
MgO
Na2O
Which of the following gasses results in the formation of acid rain?
A.
B.
C.
D.
9.
6
7
8
H2
O3
SO2
NH3
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.
B.
C.
D.
11.
Which of the following 0.10 M solutions will have the greatest electrical conductivity?
A.
B.
C.
D.
12.
H+
O2H2O
H3O+
HF
NH3
NaOH
C6H5COOH
The amphiprotic ion HSeO3- can undergo hydrolysis according to the following equations
HSeO3- + H2O ⇄ H2SeO3 + OH-
K1
HSeO3- + H2O ⇄ SeO32-+ H3O+
K2
An aqueous solution of HSeO3- is found to be acidic. This observation indicates that when it is
added to water, HSeO3- behaves mainly as a
A.
B.
C.
D.
13.
proton donor, and K1 is less than K2
proton donor, and K1 is greater than K2
proton acceptor, and K1 is less than K2
proton acceptor, and K1 is greater than K2
The Kb expression for HPO42- is
A. [PO43-][H3O+]
[HPO42-]
B. [HPO42-][OH-]
[H2PO4-]
C. [H2PO4-][OH-]
[HPO42-]
D. [HPO42-][ H3O+]
[PO43-]
Quiz #4
1.
2.
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
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.
B.
C.
D.
3.
4.
salt
acid
base
buffer
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
Which of the following solutions would be acidic?
A.
B.
C.
D.
5.
Anhydrides, Hydrolysis
sodium acetate
iron III chloride
sodium carbonate
potassium chloride
Consider the following salts:
I. NaF
II. NaClO4
III. NaHSO4
Which of these salts, when dissolved in water, would form a basic solution?
6.
A.
I only
B.
I and II only
C.
II and III only
D.
I, II and III
Which of the following, when dissolved in water, forms a basic solution?
A.
B.
C.
D.
7.
Which of the following oxides forms a basic solution?
A.
B.
C.
D.
8.
Kb for H2S
Ka for H2S
Kb for HSKa for HS-
The reaction of a strong acid with a strong base produces
A.
B.
C.
D.
11.
SO2
SO32HSO3H2SO3
Consider the following equilibrium expression
K= [H2S][OH-]
[HS-]
This expression represents the
A.
B.
C.
D.
10.
K2O
CO2
SO3
NO2
Which of the following is amphiprotic in water?
A.
B.
C.
D.
9.
KCl
NaClO4
Na2CO3
NH4NO3
A salt and a water
A base and an acid
A metallic oxide and water
A non-metallic oxide and water
Consider the following equilibrium:
CH3COOH(aq) + NH3(aq) ⇄ CH3COO- (aq) + NH4+(aq)
The sequence of Bronsted-Lowry acids and bases in the above equilibrium equation is
A.
B.
C.
acid, base, base, acid
acid, base, acid, base
base, acid, base, acid
12.
D.
base, acid, acid, base
The pH range of ‘acid rain’ is often
A.
B.
C.
D.
3 to 6
6 to 8
7 to 9
10 to 12
13.
Water will act as a Bronsted-Lowry acid with
14.
A.
NH3
B.
H2S
C.
HCN
D.
HNO3
Which of the following is a conjugate acid-base pair?
A.
B.
C.
D.
Quiz #5
1.
3.
H2S
HNO2
HNO3
H3BO3
At 25oC, the equation representing the ionization of water is
A
H2O + H2O ⇄ 2H2 + O2
B.
H2O + H2O ⇄ H2O2 + H2
C.
H2O + H2O ⇄ 4H+ + 2O2-
D.
H2O + H2O ⇄ H3O+ +OH-
The pH of a 0.3 M solution of NH3 is approximately
A.
B.
C.
D.
4.
pH calculations for Strong and Weak Acids
The 1.0 M acidic solution with the highest pH is
A.
B
C.
D.
2.
H3PO4 and PO43H2PO4- and PO43H3PO4 and HPO42H2PO4- and HPO42-
14.0
11.0
6.0
3.0
The pH of an aqueous solution is 4.32. The [OH-] is
A.
6.4 x 10-1 M
B.
C.
D.
5.
The pH of an aqueous solution is 10.32. The [OH-] is
A.
B.
C.
D.
6.
4.0 x 10-14 M
2.2 x 10-1 M
2.5 x 10-1 M
6.0 x 10-1 M
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.
B.
C.
D.
10.
Kw
Ka
Kb
Ksp
The [H3O+] in a solution of pH = 0.60 is
A.
B.
C.
D.
9.
0.94
1.60
12.40
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.
B.
C.
D.
8.
5.0 x 10-12 M
2.0 x 10-11 M
4.8 x 10-11 M
2.1 x 10-4 M
The pH of a 0.025 M HClO4 solution is
A.
B.
C.
D.
7.
4.8 x 10-5 M
2.1 x 10-10 M
1.6 x 10-14 M
-1
0
1
7
The approximate pH of a 0.06 M solution of CH3COOH is
A.
B.
C.
D.
1
3
11
13
11.
The [OH-] is greater than the [H3O+] in
A.
B.
C.
D.
12.
The pH of 0.15 M HCl is
A.
B.
C.
D.
13.
Quiz #6
5.0 x 10-16 M
1.0 x 10-14 M
2.0 x 10-13 M
5.0 x 10-2 M
Ka’s from pH
Kb’s from Ka’s
The Kb for the dihydrogen phosphate ion is
A.
B.
C.
D.
2.
0.20
0.63
0.70
1.58
The [OH-] in 0.050 M HNO3 at 25oC is
A.
B.
C.
D.
1.
pH= log [H3O+]
pH= 14 - [H3O+]
pH= -log [H3O+]
pH= pKw – [H3O+]
The pH of 0.20 M HNO3 is
A.
B.
C.
D.
15.
0.15
0.71
0.82
13.18
Which of the following equations correctly relates pH and [H3O+]?
A.
B.
C.
D.
14.
HCl(aq)
NH3(aq)
H2O(aq)
CH3COOH(aq)
1.3 x 10-12
6.3 x 10-8
1.6 x 10-7
7.1 x 10-3
What volume of 0.100 M NaOH is required to neutralize a 10.0 mL sample of 0.200 M H2SO4?
A.
B.
C.
D.
3.
5.0 mL
10.0 mL
20.0 mL
40.0 mL
Consider the following equilibriums:
I
HCO3- + H2O ⇄ H2CO3 + OH-
II
NH4+ + H2O ⇄ H3O+ + NH3
III
HSO3- + H3O+ ⇄ H2O + H2SO3
Water acts as a Bronsted-Lowry base in
A.
B.
C.
D.
4.
5.
6.
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.
Cr2O72-(aq) + 2OH-(aq) ⇄ 2CrO42-(aq) + H2O(l)
D.
CH3NH2(aq) + H2O(l) ⇄ CH3NH3+ (aq) + OH-(aq)
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
If OH- is added to a solution, the [H3O+] will
A.
B.
C.
D.
7.
III only
I and II only
II and III only
I, II, and III
Remain constant
Adjust such that [H3O+]= [OH-]
Kw
+
Increase such that [H3O ][OH-] = Kw
Decrease such that [H3O+][OH-] = Kw
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.
B.
C.
D.
8.
0.0200 M
2.00 M
4.00 M
8.00 M
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.
Which of the following indicators would be yellow at pH 7 and blue at pH 10?
A.
B.
C.
D.
10.
0.20
0.70
13.30
13.80
The 1.0 M acid solution with the largest [H3O+] is
A.
B.
C.
D.
Quiz #7
1.
CNNH3
NO2NH4+
What is the pH of a solution prepared by adding 0.50 mol KOH to 1.0 L of
0.30 M HNO3?
A.
B.
C.
D.
13.
pink as [HInd] is less than [Ind-]
pink as [HInd] is greater than [Ind-]
colorless as [HInd] is less than [Ind-]
colorless as [HInd] is greater than [Ind-]
Water acts as a base when it reacts with
A.
B.
C.
D.
12.
thymol blue
methyl violet
bromthymol blue
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.
B.
C.
D.
11.
[Ind-] < [HInd]
[Ind-] = [HInd]
[Ind-] > [HInd]
[Ind-] = [H+] = [HInd]
HNO2
H2SO3
H2CO3
H3BO3
pH for Weak bases, pH Relationships, Amphiprotic Calculations
In water, the hydrogen sulphide ion, HPO42-, will act as
A.
B.
C.
D.
An acid because Ka < Kb
An acid because Ka > Kb
A base because Ka < Kb
A base because Ka > Kb
2.
A student records the pH of 1.0 M solution of two acids. Which of the following statements can
be concluded from the above data?
A.
B.
C.
D.
3.
When added to water, the hydrogen carbonate ion, HCO3-, produces a solution, which is
A.
B.
C.
D.
4.
Acid X is stronger than acid Y
Acid X and acid Y are both weak
Acid X is diprotic while acid Y is monoprotic
Acid X is 100 times more concentrated than acid Y
basic because Kb is greater than Ka
basic because Ka is greater than Kb
acidic because Ka is greater than Kb
acidic because Kb is greater than Ka
The concentration, Ka and pH values of four acids are given in the following table
ACID Concentration
HA
HB
HC
HD
3.0 M
0.7 M
4.0 M
1.5 M
Ka
pH
2.0 x 10-5
4.0 x 10-5
1.0 x 10-5
1.3 x 10-5
2.1
2.3
2.2
2.4
Based on this data, the strongest acid is
A.
B.
C.
D.
5.
HA
HB
HC
HD
Which of the following 0.10 M solutions is the most acidic?
A.
B.
C.
D.
6.
Acid
X
Y
AlCl3
FeCl3
CrCl3
NH4Cl
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
pH
4.0
2.0
B.
C.
D.
Quiz #8
Solution
Colour
Congo red, Ka = 9.0 x 10-3
Cresol red, Ka = 1.0 x 10-8
Alizarin blue, Ka = 7.0 x 10-11
Buffers and Indicators
1.0 M HCl
Yellow
I. H2CO3
1.0 M HAl
Blue
II. HClO4
1.0 M HA2
Yellow
1.
Consider the following acid
solutions:
III. HF
Which of the above acids would form a buffer solution when its conjugate base is added?
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:
The acids HAl, HA2, and HInd listed in the order of decreasing acid strength is
A.
B.
C.
D.
3.
Which of the following compounds, when added to a solution of ammonium nitrate, will result in
the formation of a buffer solution?
A.
B.
C.
D.
4.
HA2, HInd, HAl
HInd, HAl, HA2
HA2, HAl, HInd
HAl, HInd, HA2
Ammonia
Nitric acid
Sodium nitrate
Ammonium chloride
Which of the following represents a buffer equilibrium?
A.
HI + H2O ⇄ H3O+ + I-
B.
HCl + H2O ⇄ H3O + Cl-
C.
HCN + H2O ⇄ H3O+ + CN-
D.
HClO4 + H2O ⇄ H3O + ClO4-
5.
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.
B.
C.
D.
6.
A basic buffer solution can be prepared by mixing equal numbers of moles of
A.
B.
C.
D.
Quiz #9
1.
HCl and LiOH
HNO3 and NH3
HClO4 and NaOH
CH3COOH and KOH
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.
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.
3.
NH4CL and HCl
NaCl and NaOH
Na2CO3 and NaHCO3
NaCH3COOH and CH3COOH
Titrations and Titration Curves
Which of the following indicators would be used when titrating a weak acid with a strong base?
A.
B.
C.
D.
2.
Ka
[HF]
[H2O]
[H3O+]
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.
1412108642Which 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.
12108642Which 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 HNO2 is titrated with NaOH(aq)?
A.
B.
C.
D.
10.
Methyl red
Thymol blue
Methyl orange
Indigo carmine
Which of the following solutions should be used when titrating a 25.00 mL sample of CH3COOH
that is approximately 0.1 M?
A.
B.
C.
D.
11.
Orange IV
Methyl red
Phenolphthalein
Alizarin yellow
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.
125 mL
150 mL
250 mL
500mL
12.
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-
Quiz #10
1.
Review
How many moles of Mg(OH)2 are required to neutralize 30.00 mL of 0.150 M HCl?
A.
B.
C.
D.
2.
2.25 x 10-3 mol
4.50 x 10-3 mol
5.00 x 10-3 mol
9.00 x 10-3 mol
The approximate Ka for the indicator phenolphthalein is
A.
B.
C.
D.
6 x 10-19
8 x 10-10
6 x 10-8
2 x 10-2
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
3.
this indicator:
A.
B
C.
D.
4.
HInd + H2O ⇄ H3O+ + Ind-
Equilibrium shifts left, and colour becomes red
Equilibrium shifts left, and colour becomes blue
Equilibrium shifts right, and colour becomes red
Equilibrium shifts right, and colour becomes blue
Which one of the following combinations would act as an acid buffer?
A.
B.
C.
D.
HCl and NaOH
KOH and KBr
NH3 and NH4Cl
CH3COOH and NaCH3COO
What is the pH at the transition point of an indicator if its Ka is 7.9 x 10-3?
5.
A.
B.
C.
D.
6.
0.98
2.10
7.00
11.90
Which of the following pH curves best represents the titration of sodium hydroxide with
hydrochloric acid?
B.
A.
7
7
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.
B.
C.
D.
8.
At the equivalence point in a titration involving 1.0 M solutions, which of the following
combinations would have the lowest conductivity?
A.
B.
C.
D.
9.
NH3 and the pH decreases
NH4+ and the pH decreases
NH3 to keep the pH relatively constant
NH4+ to keep the pH relatively constant
Nitric acid and barium hydroxide
Acetic acid and sodium hydroxide
Sulphuric acid and barium hydroxide
Hydrochloric acid and sodium hydroxide
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.
B.
C.
D.
10.
Products are favored and the stronger acid is HInd
Products are favored and the stronger acid is HCN
Reactants are favored and the stronger acid is HInd
Reactants are favored and the stronger acid is HCN
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-
11.
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.
B.
C.
D.
12.
[HInd] = [Ind-]
[Ind-] = [H3O+]
[HInd] = [H3O+]
[H3O+] = [H2O]
Which of the following titrations would have an equivalence point less that pH 7?
A.
B.
C.
D.
NH3 and HCl
NaOH and HNO3
Ba(OH)2 and H2SO4
KOH and CH3COOH
Web Review of Acids
1.
2.
List five properties of:
a)
acids
b)
bases
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.
F-
OH-
Co(H2O)5(OH)2+
Determine the conjugates for each of the acids.
HF
9.
NH3
HCN
Al(H2O)63+
NH4+
HPO42-
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
5.0 x 10-3 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.
2.
3.
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-
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-]
Consider the following graph for the titration of 0.1 M NH3 with 1.0 M HCl.
pH
14
7
Volume HCl added
I
0
II
III
I
V
A buffer solution is present at point
A.
B.
C.
D.
4.
Consider the following equilibrium system for an indicator: HInd + H2O ⇄ H3O+ + IndWhich two species must be of two different colours in order to be used as an indicator?
A.
B.
C.
D.
5.
methyl red
phenol red
thymol blue
methyl violet
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.
B.
C.
D.
7.
HInd and H2O
HInd and IndH3O+ and IndHind and H3O+
Which of the following indicators is yellow at pH 10.0?
A.
B.
C.
D.
6.
I
II
III
IV
6.00 x 10-3 M
6.00 x 10-2 M
1.20 x 10-1 M
2.4 x 10-1 M
Which of the following titrations will have the highest pH at the equivalence point?
A.
B.
C.
HBr with NH3
HNO2 with KOH
HCl with Na2CO3
D.
8.
An Arrhenius acid is defined as a chemical species that
A.
B.
C.
D.
9.
HNO3 with NaOH
is a proton donor.
is a proton acceptor.
produces hydrogen ions in solution.
produces hydroxide ions in solution.
Consider the following acid-base equilibrium system:
HC2O4- + H2BO3- ⇄ H3BO3 + C2O42Identify the Bronsted-Lowry bases in this equilibrium.
A.
B.
C.
D.
10.
11.
CH3COO- + H2O ⇄ CH3COOH + OH-
B.
CH3COO- + H2O ⇄ H2O + CH2COO2-
C.
CH3COOH + H2O ⇄ H3O+ + CH3COO-
D.
CH3COOH + H2O ⇄ CH3COOH2+ + OH-
Which of the following solutions will have the lowest electrical conductivity?
0.10 M HF
0.10 M NaF
0.10 M H2SO3
0.10 M NaHSO3
Which of the following is the strongest Bronsted-Lowry base?
NH3
CO32HSO3H2BO3-
A 1.0 x 10-4 M solution has a pH of 10.00. The solute is a
A.
B.
C.
D.
14.
H3BO3
H3BO3
C2O42C2O42-
A.
A.
B.
C.
D.
13.
and
and
and
and
The equation representing the predominant reaction between NaCH3COO with water is
A.
B.
C.
D.
12.
H2BO3HC2O4HC2O4H2BO3-
weak acid
weak base
strong acid]
strong base
The ionization of water at room temperature is represented by
A.
H2O ⇄ 2H+ + O2-
15.
B.
2H2O ⇄ 2H2 + O2
C.
2H2O ⇄ H2 + 2OH-
D.
2H2O ⇄ H3O+ + OH-
Addition of HCl to water causes
A.
B.
C.
D.
16.
both [H3O+] and [OH-] to increase
both [H3O+] and [OH-] to decrease
[H3O+] to increase and [OH-] to decrease
[H3O+] to decrease and [OH-] to increase
Consider the following:
I.
II.
III.
H2SO4
HSO4SO42-
Which of the above is/are present in a reagent bottle labeled 1.0 M H2SO4?
17.
A.
I only
B.
I and II only
C.
II and III only
D.
I, II, and III
The pH of 0.10 M KOH solution is
A.
B.
C.
D.
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.
B.
C.
D.
19.
20.
0.10
1.00
13.00
14.10
1
1
1
1
x
x
x
x
10-13
10-10
10-7
10-1
Which of the following are amphiprotic in aqueous solution?
I.
II.
III.
IV.
HBr
H2O
HCO3H2C6H5O7-
A
B.
C.
D.
I and II only
II and IV only
II, III, and IV only
I, II, III, and IV
Which of the following always applies at the transition point for the indicator Hind?
A.
B.
C.
D.
[Ind-] = [OH-]
[HInd] = [Ind-]
[Ind-] = [H3O+]
[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.
B.
C.
D.
24.
What is the value of the Kw at 25 oC?
A.,
B.
C.
D.
25.
[H3O+]
[OH-]
increases
increases
decreases
decreases
increases
decreases
decreases
increases
Consider the following equilibrium for an indicator: HInd + H2O ⇄ H3O+ + IndIn a solution of pH of 6.8, the colour of bromthymol blue is
A.
B.
C.
D.
27.
1.0 x 10-14
1.0 x 10-7
7
14
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?
A.
B.
C.
D.
26.
PO43H2PO4H2PO42H2PO43-
blue because [HInd] = [Ind-]
green because [HInd] = [Ind-]
green because [HInd] < [Ind-]
yellow because [HInd] > [Ind-]
The indicator with Ka = 4 x 10-8 is
A.
B.
C.
D.
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.
B.
C.
D.
29.
LiF
CrCl3
KNO3
NH4Cl
5.9
2.4
4.1
1.7
x
x
x
x
10-10
10-8
10-7
10-5
0.97
1.80
12.18
13.03
Which of the following will produce an acidic solution?
A.
B.
C.
D.
34.
basic
acidic
neutral
amphiprotic
The pOH of 0.015 M HCl solution is
A.
B.
C.
D.
33.
10-3 M
10-3 M
10-2 M
10-1 M
What is the value of the Kb for HC6H5O72-?
A.
B.
C.
D.
32.
x
x
x
x
Which of the following salts will dissolve in water to produce a neutral solution?
A.
B.
C.
D.
31.
1.43
2.86
5.72
1.14
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.
B.
C.
D.
30.
neutral red
methyl red
indigo carmine
phenolphthalein
NaCl
NH4NO3
Ca(NO3)2
Ba(NO3)2
Which of the following salts will dissolve in water to produce an acid solution?
A.
LiF
B.
C.
D.
35.
36.
37.
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 H2SO4. The resulting
solution has a pH of
A.
14.000
B.
0.000
C.
13.800
D.
7.000
A student mixes 500 mL of 0.400 M NaOH with 500 mL of 0.100 M H2SO4. The resulting
solution has a pH of
A.
B.
C.
D.
38.
CrCl3
KNO3
NaCl
14.000
0.000
13.000
7.000
The strongest acid in water is
A.
B.
C.
D.
HClO4
HI
HF
H3O+
39.
The formula that has the highest pH in water is
39.
A.
HF
B.
H2CO3
C.
H2C2O4
D.
HCN
The formula that has the highest pH in water is
A.
B.
C.
D.
NaF
NaCl
H2C2O4
NaCN
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
B
C
D
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
B
C
D
3.
neutral red
thymol blue
thymolpthalein
chlorophenol red
Acid is added to a buffer solution. When equilibrium is reestablished the buffering effect has
resulted in [H3O+]
A
B
C
D
5.
10.0 mL
20.0 mL
25.0 mL
40.0 mL
Which indicator has a Ka = 1.0 x 10-6?
A
B
C
D
4.
red
blue
yellow
colourless
increasing slightly
decreasing slightly
increasing considerably
decreasing considerably
A buffer solution will form when 0.10 M NaF is mixed with an equal volume of
A
B
C
D
0.10 M HF
0.10 M HCl
0.10 M NaCl
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
B
C
D
7.
8.
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-
What is the pOH of a solution prepared by adding 0.50 moles of NaOH to prepare 0.50 L of
solution?
A
B
C
D
9.
x
x
x
x
10-14
10-9
10-6
10-1
adding a strong acid
adding a strong base
increasing the temperature
decreasing the temperature
The complete neutralization of 15.0 mL of KOH requires 0.0250 moles H2SO4. The [KOH] was
A
B
C
D
12.
1.4
1.6
6.3
7.1
Consider the following equilibrium: 2H2O(l) + energy ⇄ H3O+(aq) + OH-(aq)
What will cause the pH to increase and the Kw to decrease?
A
B
C
D
11.
0.00
0.30
14.00
13.70
What is the [H3O+] in a solution with a pH = 5.20?
A
B
C
D
10.
partially ionizes
neutralizes an acid
has a pH greater than 7
turns bromocresol green from yellow to blue
1.50 M
1.67 M
3.33 M
6.67 M
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
B
C
D
14.
Which of the following acids has the weakest conjugate base?
A
B
C
D
15.
18.
0.010 M
0.050 M
0.10 M
0.20 M
The conjugate base of an acid is produced by
A
B
C
D
17.
HIO3
HNO2
H3PO4
CH3COOH
When 10.0 ml of 0.10 M HCl is added to 10.0 mL of water, the concentration of H3O+ in the final
solution is
A
B
C
D
16.
HCl and NaCl
HCN and NaCN
KNO3 and KOH
Na2SO4 and NaOH
adding a proton to the acid
adding an electron to the acid
removing a proton from the acid
removing an electron from the acid
Which of the following represents the predominant reaction between HCO3- and water?
A
2HCO3- ⇄ H2O + 2CO2
B
HCO3- + H2O ⇄ H2CO3
C
HCO3- + H2O ⇄
D
2HCO3- + H2O ⇄ H3O+ + CO32- + OH- + CO2
+ OH-
H3O+ + CO32-
Water acts as an acid when it reacts with
I
II
III
IV
CNNH3
HClO4
CH3COO-
A
B
C
I and IV only
II and III only
I, II, and IV
D
19.
II, III, and IV
In a solution of 0.10 M H2SO4, the ions present in order of decreasing concentration are
A
B
C
D
[H3O+] > [HSO4-] > [SO42-] > [OH-]
[H3O+] > [SO42-] > [HSO4-] > [OH-]
[OH-] > [SO42-] > [HSO4-] > [H3O+]
[SO42-] > [HSO4-] > [OH-] > [H3O+]
20. Which of the following will dissolve in water to produce an acidic solution?
A
B
C
D
21.
22.
Which of the following solutions will have a pH = 1.00?
I
II
III
0.10 M HCl
0.10 M HNO2
0.10 M NaOH
A
B
C
D
I only
II only
I and II only
I, II, and III
Ka for the acid H2AsO4- is 5.6 x 10-8. What is the value of the Kb for HAsO42-?
A
B
C
B
23.
5.6
3.2
1.8
2.4
x
x
x
x
10-22
10-14
10-7
10-4
In a titration, which of the following has a pH = 7.00 at the equivalence point?
A
B
C
D
24.
CO2
CaO
MgO
Na2O
NH3 and HNO3
KOH and HCl
NaF and HCl
Ca(OH)2 and CH3COOH
Which of the following salts dissolves to produce a basic solution?
A
B
C
D
KCl
NH4Br
Fe(NO3)3
LiCH3COO
25.
Calculate the pH in a 0.200 M solution of Sr(OH)2.
A
B
C
D
26.
Which of the following solutions has a pH less than 7.00?
A
B
C
D
27.
1
1
1
1
x
x
x
x
10-14
10-8
10-6
10-3
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
B
C
D
30.
HSO3HSO4HPO42HC2O4-
What is the approximate Ka value for the indicator chlorophenol red?
A
B
C
D
29.
NaCl
LiOH
NH4NO3
KCH3COO
Which of the following will form a basic aqueous solution?
A
B
C
D
28.
1.40
1.70
13.30
13.60
0.00
1.40
1.70
7.00
In which one of the following equations are the Bronsted acids and bases all correctly identified?
Acid
+
Base
⇄
Base
+
Acid
A
H2O2
SO32-
⇄
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
B
C
D
32.
A buffer solution may contain equal moles of
A
B
C
D
33.
BaS
NH4Cl
Ca(NO3)2
NaCH3COO
Which of the following statements applies to 1.0 M NH3(aq) but not to 1.0 M NaOH(aq)?
A
B
C
D
36.
H2
O3
SO2
C3H8
Which of the following 1.0 M salt solutions is acidic?
A
B
C
D
35.
weak acid and strong base
strong acid and strong base
weak acid and its conjugate base
strong acid and its conjugate base
A gas which is produced by burning coal and also contributes to the formation of acid rain is
A
B
C
D
34.
weak acid with a weak base
strong acid with a weak base
weak acid with a strong base
strong acid with a strong base
partially ionizes
neutralizes and acid
has a pH greater than 7
turns bromcresol green from yellow to blue
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
B
C
D
Subjective
HF
H2S
HCN
HNO3
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.
Final volume of H2C2O4 from burette (mL)
Initial volume of H2C2O4 from burette (mL)
Trial 1
Trial 2
Trial 3
23.00
4.85
39.05
23.00
20.95
5.00
Calculate the [NaOH]
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.
+ OH-
b) Calculate the value of the Kw for the sample of water with a pH = 7.30.
Acids, Bases and Salts Unit Plan
Notes- double click on the lesson number and download Power Point Viewer if you do not have it.
Worksheets
1. Properties of Acids, Bases & Salts
WS 1
2. Arrhenius, Bronsted Acids, Ka and Strength.
WS 2
3. Arrhenius, Bronsted Bases, Kb and Strength
WS 3
4. Acid & Base Reactions. Amphiprotic. Acid Chart.
WS 4
5. Leveling effect, Anhydrides and Relationships.
WS 5
6. Hydrolysis of Salts. Quiz.
WS 6
7. Acid, Base & Salt Reactions. Hydrolysis.
WS 7
Quiz
1
2
3
ClassifyingEverything Activity
8. Yamada’s Indicator Lab. Hydrolysis.
WS 8
9. Ionization of Water, [H+] & [ OH- ], pH scale.
4
Hydrolysis Quiz
1Hydrolysis Quiz 2Hydrolysis Quiz 3
10. pH for Strong Acids and Bases. Hydrolysis Quiz
WS 9
11. pH Calculations for Weak Acids.
WS 10
12. Ka from pH for Weak Acids.
WS 11
5
13. Indicators Lab.
14. Kbs from Kas for Weak Bases.
WS 12
6
15. pH for Weak Bases pH [H+] [ OH-] Relationships. WS 13
16.Amphiprotic Ions- Kas and Kbs.
WS 14
17. Titration Lab. Primary Standards.
7
Acids
Midtermreview Test
18. Titration Lab
19. Buffers & Indicators
WS 15
8
20. Titration Curves.
WS 16
9/10
21. Review #1
Web Site Review Sheet
Quizmebc
22. Review #2
Practice Test # 1
Practice Test 2
23. Test
Text book
Hebden
Read Unit IV
WS #1 Properties of Acids and Bases
1. Add 1 drop of each solution to 1 drop of the acid-base indicator in a spot plate. Record the colour in the
data table below. Describe each solution as an acid or base in the space provided. Write the acid colour
and base colour in the table below.
Indicator-->
Phenolphthalein
Solution:
HCl
Clear
Red
Yellow
Acid
NaOH
Pink
Blue
Blue
Base
Vinegar
Clear
Red
Yellow
Acid
Ammonia
Pink
(NH3)
Lemon Juice Clear
Blue
Blue
Base
Red
Yellow
Acid
Seven-up
Red
Yellow
Acid
Baking Soda Pink
(NaHCO3)
Blue
Blue
Base
Indicator
Acid Colour
Base Colour
Clear
Pink
Clear
Litmus
Phenolphthalein
Bromothymol Blue
Litmus
Red
Blue
Bromothymol Blue
Yellow
Blue
Acid or Base
Wash and dry your spot plate before going on to step 2.
2. Wear safety goggles for this experiment. Pour approximately 50 mL of 1 M HCl into a fleaker. Add
one level spoonful of Mg and cover with a plastic funnel. After 1 minute and not before light the top of
the funnel using a match. Write the equation for the reaction below.
2HCl(aq) + Mg(s) → H2(g) + MgCl2(aq)
Wash and dry your fleaker before going on to step 3.
3. Taste a lemon and describe the taste in one word
Sour
4. Taste some baking soda and describe the taste in one word.
Bitter
6. Test two drops of HCl for conductivity in a spot plate. Result:
Good Conductor
Write an equation that accounts for the conductivity of HCl.
HCl
→
H+
+
Cl-
7. Test two drops of NaOH for conductivity in a spot plate. Result:
Good Conductor
Write an equation that accounts for the conductivity of NaOH (dissociation).
NaOH → Na+ + OHClean, dry and put away the spot plate
8. List five properties of acids that are in your textbook.
Acids conduct electricity, taste sour, neutralize bases, change the color of indicators, and react with some
metals to produce hydrogen.
9. List five properties of bases that are in your textbook.
Bases conduct electricity, taste bitter, neutralize acids, change the color of indicators, and feel
slippery.
10. Make some brief notes on the commercial acids: HCl and H2SO4 (p 112).
HCl
H2SO4
11. Make some brief notes on the commercial base NaOH (p 114).
12. Describe the difference between a concentrated and dilute acid (hint: concentration refers to the
molarity). Describe their relative conductivities.
Concentrated means relatively high molarity and dilute means relatively low molarity.
13. Describe the difference between a strong and weak acid (p 121-124). Use two examples and write
equations to support your answer. Describe their relative conductivities.
A strong acid completely ionizes and a weak acid partially ionizes.
14. Describe a situation where a strong acid would have the same conductivity as a weak acid (hint: think
about concentration).
A weak acid could have a high molarity and the strong acid could have a low molarity.
Complete this worksheet for next period. Read pages 107-126 for homework.
WS #2 Conjugate Acid-Base Pairs
Complete each acid reaction. Label each reactant or product as an acid or base. The first on is done for
you.
1.
HCN
Acid
+
H2O
Base
⇄
H3O+
Acid
+
CNBase
2.
H3C6O7
acid
+
H2O
base
⇄
H2C6O7- + H3O+
base
acid
3.
H3PO4
acid
+
H2O
base
⇄
H2PO4base
4.
HF
acid
+
H2O
base
⇄
F- +
base
5.
H2CO3
acid
+
H2O
base
⇄
HCO3- +
base
6.
NH4+
acid
+
H2O
base
⇄
NH3 +
base
7.
CH3COOH +
acid
H2O
base
⇄
CH3COO- +
base
8.
HCl
acid
+
H2O
base
→
Cl- +
base
9.
HNO3 +
acid
H2O
base
→
H3 O+
acid
+ H3O+
acid
H 3 O+
acid
H3O+
acid
H 3 O+
acid
H3 O+
acid
H 3 O+
acid
+
NO3base
Write the equilibrium expression (Ka) for the first seven above reactions.
[H3O+] [ CN-]
[HCN]
10.
Ka =
11.
Ka = [H3O+] [H2C6O7-]
[H3C6O7]
12.
Ka =
[H3O+] [H2PO4-]
[H3PO4]
14. Ka =
[H3O+] [HCO3-]
[H2CO3]
15. Ka =
[H3O+] [NH3]
[NH4+]
16. Ka =
[H3O+] [CH3COO-]
[CH3COOH]
[H3O+] [F-]
[HF]
17. Which acids are strong? The six on the top of the acid chart are strong.
18. What does the term strong acid mean? They complete ionization into ions. Such as: HCl + H2O
→ Cl- + H3O+
13.
Ka =
19. Why is it impossible to write an equilibrium expression for a strong acid?
[HCl] is equal to zero and in math numbers divided by zero are undefined.
20. Which acids are weak?
All acids listed on the acid chart below the top six.
23. What does the term weak acid mean?
Ka = [H3O+] [Cl-]
[HCl]
Incomplete ionization. Such as: HF + H2O ⇄ F- + H3O+
24. Explain the difference between a strong and weak acid in terms of electrical conductivity.
A strong acid is a good conductor. A weak acid conducts but not so good.
14.
16.
18.
20.
22.
Acid
Conjugate Base
HNO2
HSO3H2O2
HSH2O
NO2SO32HO2S2OH-
15.
17.
19.
21.
23.
Base
Conjugate Acid
HCOOIO3NH3
CH3COOH2O
HCOOH
HIO3
NH4+
CH3COOH
H3 O+
Define:
22. Bronsted acid- a proton donor
23. Bronsted base- a proton acceptor
24. Arrhenius acid- a substance that ionizes in water to produce H+
25. Arrhenius base- a substance that ionizes in water to produce OH26. List the six strong acids. HCLO4 HI
HBr
HCl
HNO3
H2SO4
27. Rank the acids in order of decreasing strength.
HCl
HSO4H3PO4
HF
H2CO3
H2 S
28. What would you rather drink vinegar or hydrochloric acid? Explain.
Vinegar. It is a weak acid and produces much less H30+ ion which is the corrosive part of an acid.
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
8 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 acetate.
Wash and return your eyedropper.
Wash and return your beaker.
Wash your hands.
Results
Compound
Conductivity
pH
Classification
HCl
good
1
strong acid
CH3COOH
ok
3
weak acid
NaOH
good
13
strong base
NH3
ok
11
weak base
NaHCO3
good
11
weak base
H2CO3
ok
3
weak acid
NaCl
good
7
neutral salt
WS #3 Conjugate Acid-Base Pairs
Complete each reaction. Label each reactant or product as an acid or base.
1.
HCN
+
H2O
⇄
H3 O +
+
CN-
2.
HCl
+
H2O
⇄
H3 O +
+
Cl-
3.
HF
+
H2O
⇄
H3 O +
+
F-
4.
F-
+
H2O
⇄
HF
+
OH-
5.
HSO4-
+ H2O
(acid)
⇄
H3 O +
+
SO42-
6.
NH4+
+
H2O
⇄
H3 O +
+
NH3
7.
HPO42-
+
(base)
H2O
⇄
H2PO4-
+
OH-
Acid
Conjugate Base
8.
10.
12.
14.
HCO3HPO4-2
H2O
HS-
16.
Circle the strong bases.
CO32PO43OHS2-
Fe(OH)3
Zn(OH)2
17.
Conjugate Acid
CH3COOIO3NH2C2H5SO73-
CsOH
Ba(OH)2
CH3COOH
HIO3
NH3
HC2H5SO72-
KOH
Ca(OH)2
Rank the following acids from strongest to weakest.
H2PO46
CH3COOH
4
HI
1
HCl
1
HF
3
Rank the following bases from the strongest to weakest.
F4
H2O
6
19.
9.
11.
13.
15.
NaOH
Sr(OH)2
H2S
5
18.
Base
SO323
NH3
2
HSO35
NaOH
1
i) Write the reaction of H3BO3 with water (remove one H+ only because it is a weak acid).
H3BO3 +
H2O
⇄
H2BO3- +
H3O+
ii) Write the Ka expression for the above.
Ka =
[H3O+] [H2BO3-]
[H3BO3]
iii) What is the ionization constant for the acid (use your table). Ka = 7.3 x 10-10
20.
List six strong acids.
HClO4
HI
KOH
LiOH
HBr
HCl
HNO3
H2SO4
21.
List six strong bases. NaOH
Ba(OH)2
RbOH
22.
List six weak acids in order of decreasing strength (use your acid/base table).
CsOH
HIO3
23.
H2C2O4
HSO4-
H2SO3
H3PO4
HNO2
List six weak bases in order of decreasing strength (use your acid/base table).
PO43-
CO32-
CN-
H2BO3-
NH3
HS-
WS #4 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-
⇄
SO42-
H2PO4-
+
Acid
Base
Base
Acid
Products are favoured as HSO4 is a stronger acid than H2PO4
HCN
+
H2 O
⇄
H3 O + +
CNAcid
Base
Acid
Base
Reactants are favoured as H3O+ is a stronger acid than HCN.
2.
3.
HCO3-
+
⇄
H2 S
HS-
H2CO3 +
Base
Acid
Acid
Reactants are favoured as H2CO3 is a stronger acid than H2S
4.
HPO42-
+
⇄
NH4+
H2PO4-
Base
+
NH3
Base
Acid
Acid
Base
Reactants are favoured as H2PO4 is a stronger acid than NH4+
5.
NH3
+
⇄
H2 O
NH4+
+
Base
Acid
Acid
Reactants are favoured as OH- is a stronger base than NH3
6.
H2PO41Acid
+
⇄
NH3
Base
HPO42Base
OHBase
+
NH4+
Acid
Products are favoured as NH3 is a stronger base than HPO42HCO3+
HF
⇄
H2CO3
Base
Acid
Acid
Products are favoured as HF is a stronger acid than H2CO3
7.
+
FBase
8. Complete each equation and indicate if reactants or products are favoured. Label each acid or base.
HSO4- + HCO3H2PO4favoured
⇄ H2CO3
+ HC03
HS03- + HPO42-
+
⇄ HPO42-
⇄ H2PO4-
+
SO42+
SO32-
products are favoured
H2CO3
reactants are
products are favoured
NH3
+ HC2O4-
⇄ NH4+
C2O42-
+
products are favoured
9. Explain why HF(aq) is a better conductor than HCN(aq).
HF is a stronger acid and creates more ions.
10. Which is a stronger acid in water, HCl or HI? Explain!
Both are strong acids and have the same strength as both completely ionize to from H+.
11. State the important ion produced by an acid and a base.
Acid: H+ or H3O+ Base: OH12. Which is the stronger base? Which produces the least OH-?
F- is the weaker base and produces the least OHCO3-2 is the stronger base
13. Define a Bronsted/Lowry acid and base.
An acid is a proton donor and a base is a proton acceptor.
14. Define an Arrhenius acid and base.
An acid ionizes in water to produce H+ and a base ionizes in water to produce OH-.
15. Complete each reaction and write the equilibrium expression.
HF + H2O
⇄
H3O+ + F-
F- + H2O
⇄
HF + OH-
Ka=
[ H3O+][ F-]
Kb=
[HF][
OH-]
16. H2SO4 + 2NaOH
→ Na2SO4
17. Define conjugate pairs.
Acid base pairs that differ by one proton.
18. Give conjugate acids for:
HS-,
CO3-2
H2S NH4+
HCO3.
19 Give conjugate bases for:
NH4+,
H2O
NH3 FOH-
[F-]
[HF]
+
2HOH
NH3, HPO4-2,
OH-,
H2O,
NH3,
H2PO4-
HOH
H 3 O+
NH4+
H3O+,
OH-,
HCO3-,
HOH
O2-
CO3-2
HF,
H2PO4-,
HPO4-2
WS #5 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.
H3 O +
HCl + H2O
→
H3 O+
+
Cl2. 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.
OHNaOH → Na+ + OH3. List three strong acids and three strong bases.
HCl HI
HClO4
NaOH
KOH
4. Rank the acids in decreasing strength:
LiOH
HClO4
HClO2
1
3
Ka is very large
Ka=8.0x10-5
HClO3
HClO
2
4
Ka=1.2x10-2
Ka=4.4x10-8
5. For an oxy acid what is the relationship between the number of O’s and acid strength? (Compare H2S04
and H2S03)
The more O’s the stronger the acid.
6.Which acid is stronger?
HI03 or HIO2
7.Which produces more H30+?
H2CO3 or HS04-
8.Which produces more OH-?
F- or HC03-
9.Which conducts better NH3 or NaOH (both .1M)? Why?
NaOH is a strong acid.
10.Which conducts better HF or HCN (both .1M)? Why?
HF is a stronger acid.
11. Compare and contrast a strong and weak acid in terms of degree of ionization, size of ka, conductivity,
and concentration of H+.
Strong acid: complete ionization, very large Ka, good conductor, high [H+].
Weak acid: partial ionization, small Ka, OK conductor, low [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
Classification
Reaction
12. Na2O
basic anhydride
Na2O
13. CaO
basic anhydride
CaO
14. SO3
acid anhydride
15. CO2
16. SO2
H2 O →
2NaOH
+
H2 O →
Ca(OH)2
SO3
+
H2 O
→
H2SO4
acid anhydride
CO2
+
H2 O
→
H2CO3
acid anhydride
SO2
+
H2 O
→
H2SO3
+
17. HCl
strong acid
HCl
+
H2 O
→
H3 O+
18. NH3
weak base
NH3
+
H2 O

NH4+ +
19. NaOH
strong base
NaOH
→
Na+
+
20. HF
weak acid
HF
+
H2 O

H3 O+
+
F-
21. H3PO4
weak acid
H3PO4 +
H2 O

H3 O+
+
H2PO4-
+
Cl-
OH-
OH-
WS # 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 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
weak base
NH3
2. KCl
3. HNO3
+
NH4+ +
OH-
KCl
+
neutral salt
HCl
+
KOH
KCl
→
K+
→
H2 O
Cl-
+
strong acid
HNO3 +
4. NaHCO3
⇄
H2 O
H2 O
→
H3 O + +
NO3-
basic salt
H2CO3 +
NaOH
→
NaHCO3
+
H2O
5. RbOH
NaHCO
→
Na+
+
HCO3-
HCO3-
+
H2 O
⇄
H2CO3
→
Rb+
+
OH-
3HCl
+
Al(OH)3
AlCl3
→
Al+3
Al(H2O)63+
⇄
7. H2C2O4
+
H2 O
C6H5OH
+
NaOH
NaC6H5O
→
Na+
+
C6H5O-
C6H5O-
+
H2 O
⇄
3HNO3
+
Co(OH)3
Co(NO3)3
→
Co+3
Co(H2O)63+
⇄
acid salt
9. Co(NO3)3
10. Na2CO3
→ AlCl3
+
3H2O
Al(H2O)5(OH)2+
+
H+
⇄
H3 O+
+
HC2O4-
NaC6H5O
+
H2O
C6H5OH
+
OH-
→ Co(NO3)3
+
3H2O
+
2H2O
3Cl-
+
weak acid
H2C2O4
8. NaC6H5O
OH-
strong base
RbOH
6. AlCl3
+
basic salt
→
acid salt
+
3NO3-
Co(H2O)5(OH)2+
+ H+
basic salt
H2CO3
+
2NaOH
→ Na2CO3
Na2CO3
→
2Na+ +
CO3-2
+
H2 O
WS # 7
⇄
CO3-2
HCO3-
+
OH-
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
weak base
NH3
2. NaCl
+
+
CN-
→
+
Cl-
Na+ +
+ H2 O
+
OH-
CN⇄
HCN
+
OH-
strong base
→
Na+
→
Fe+3 +
acid salt
FeCl3
Fe(H2O)63+
7. HF
H3 O +
Cl-
H2 O
→
NaOH
6. FeCl3
+
basic salt
NaCN
5. NaOH
OH-
strong acid
HCl
4. NaCN
NH4+ +
neutral salt
NaCl → Na+
3. HCl
⇄
H2 O
weak acid
⇄
3Cl-
Fe(H2O)5(OH)2+ + H+
HF
8. LiHCO3
+
H2 O
⇄
→
Li+ +
HCO3- + H2O
10. MgCO3
Fe+3 +
Fe(H2O)63+
⇄
Fe(H2O)5(OH)2+
OH-
→
Mg+2 +
3NO3+ H+
basic salt
CO3-2
⇄
HCO3- +
OH-
H2 O
⇄
H3 O +
+
HS-
H2 O
⇄
H3 O +
+
F-
weak acid
+
weak acid
HF
+
neutral salt
CaI2 → Ca+2
14. Mg(OH)2
H2CO3 +
→
H2 S
13. CaI2
⇄
HCO3-
Fe(NO3)3
CO3-2 + H2O
12. HF
F-
acid salt
MgCO3
11. H2S
+
basic salt
LiHCO3
9. Fe(NO3)3
H3 O +
weak base
+
2I-
Mg(OH)2
15. Ba(OH)2
⇄
Mg+2 +
2OH-
strong base
Ba(OH)2 → Ba+2
+
2OH-
16. Describe why Tums (CaCO3) neutralizes stomach acid. It is a weak base and will neutralize acid.
basic salt
CaCO3
→
Ca+2 +
CO3-2
CO3-2 + H2O
⇄
HCO3- +
OH-
17. Describe why Mg(OH)2 is used in Milk of Magnesia as an antacid instead of NaOH.
Mg(OH)2 is weak base and releases OH- slowly, whereas NaOH is a strong base which releases OHin high concentrations which is corrosive.
Mg(OH)2
NaOH
⇄
→
Mg+2 +
Na+ +
2OHOH-
WS #8 Yamada’s Indicator Activity
Acid, Base and Salt Lab
Purpose:
1) To use Yamada’s Indicator to determine the pH of various acids, bases and salts.
2) To classify compounds as strong acids, weak acids, strong bases, weak bases, neutral salts, acid
anhydrides, and basic anhydrides.
3) To write reactions for each compound to show how each ionizes, hydrolyzes or reacts with water.
Procedure:
1) 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.
2) Add two drops of Yamada’s Indicator. Record the pH of the compound.
3) 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.
4) Write an equation to show the reaction of anhydrides with water, the hydrolysis of salts, or the
ionization of acids or bases.
Data
1.
Formula of compound
pH
Classification
Fe(NO3)3
2
acid salt
Reaction or reactions
Fe(NO3)3
→
Fe+3 +
Fe(H2O)63+
⇄
Fe(H2O)5(OH)2+
3NO3+ H+
2.
Formula of compound
pH
Classification
NaCH3COO
10
basic salt
Reaction or reactions
NaCH3COO →
CH3COO-
3.
Formula of compound
pH
Classification
K2HPO4
10
basic salt
Reaction or reactions
K2HPO4
HPO4-2
4.
5.
6.
Formula of compound
pH
Classification
HCL
0
strong acid
Reaction or reactions
HCl
Formula of compound
pH
Classification
Al2(SO4)3
3
acid salt
Reaction or reactions
CH3COO-
+ H2O ⇄ CH3COOH +
→
+ H2 O
2K+ +
⇄
H2PO4-
H+ +
Cl-
Al2(SO4)3
→
2Al+3 +
3SO4-2
Al(H2O)63+
⇄
Al(H2O)5(OH)2+
→
2Na+ +
Formula of compound
pH
Classification
Na2CO3
12
basic salt
Reaction or reactions
Na2CO3
OH-
HPO4-2
→
CO3-2 + H2O
7.
Na+ +
+ H+
CO3-2
⇄
HCO3- +
Formula of compound
pH
Classification
P2O5
2
acid anhydride
Reaction or reactions
P2O5 + H2O
→
H2P2O6
H2P2O6
⇄
H+ +
OH-
HP2O6-
+
OH-
8.
Formula of compound
pH
Classification
Cu(NO3)2
4
acid salt
Reaction or reactions
Cu(NO3)2
→
Cu2+ +
2NO3Cu(H2O)62+ <⇄
Cu(H2O)5
9.
10.
11.
12.
13.
14.
(OH)+
+
H+
Formula of compound
pH
Classification
Fe2(SO4)3
3
acid salt
Reaction or reactions
Fe2(SO4)3
→
2Fe+3 +
Fe(H2O)63+
⇄
Fe(H2O)5(OH)2+
3SO4-2
Formula of compound
pH
Classification
N2O5
0
acid anhydride
Reaction or reactions
N2O5 + H2O → H2N2O6 → 2HNO3
HNO3 → H+ +
NO3-
Formula of compound
pH
Classification
Ca(OH)2
12
weak base
Reaction or reactions
Ca(OH)2
Formula of compound
pH
Classification
KHSO4
2
acid salt
Reaction or reactions
⇄
Ca+2
+
2OH-
KHSO4
→
K+
+
HSO4-
HSO4-
⇄
H+
+
SO42-
→
Na+
+
HCO3-
Formula of compound
pH
Classification
NaHCO3
10
basic salt
Reaction or reactions
NaHCO3
Formula of compound
pH
Classification
HCO3- + H2O ⇄
CaCO3
10
basic salt
Reaction or reactions
CaCO3
→
H2CO3 +
OH-
Ca+2
CO3-2
+
+ H+
CO3-2 + H2O
15.
Formula of compound
pH
Classification
CaO
12
basic anhydride
Reaction or reactions
CaO
+
Ca(OH)2
16.
17.
<⇄ HCO3- +
OH-
H2O → Ca(OH)2
⇄
Ca+2
+
2OH-
Formula of compound
pH
Classification
Al2(SO4)3
3
acidic salt
Reaction or reactions
Al2(SO4)3
→
2Al+3 +
⇄
Al(H2O)5(OH)2+
Formula of compound
pH
Classification
Al(H2O)63+
NaCl
7
neutral salt
Reaction or reactions
NaCl →
Na+
+
3SO4-2
+ H+
Cl-
WS # 9 - pH and pOH Calculations
Complete the chart:
[H+]
[OH-]
-3
7.00 x 10 M 1.43 x 1012M
1.14 x 108.75 x 10-2
13M
M
-8
4.7x 10 M
2.1 x 10-7M
-10
1.0 x 10 M 1.0 x 10-4M
1.0 x 10-7M
1.0 x 10-7M
1.
2.
3.
4.
5.
5 x 10-4M
2.80 x 103M
5.0 x 10-10 M
2.1 x 10-5M
6.
7.
8.
9.
10.
2 x 10-11M
3.57 x 1012M
2.0 x 10-5M
4.7 x 10-10
M
pH
2.155
pOH
11.845
Acid/base/neutral
acid
12.942
1.058
base
7.33
10.00
7.00
6.67
4.00
7.00
3.3
2.553
10.7
11.447
9.30
4.67
4.70
9.33
base
base
Neutral (2sig
figs)
acid
acid
Calculate the [H+], [OH-] , pH and pOH for a 0.20 M Ba(OH)2 solution.
Ba(OH)2
0.20M
⇄
Ba+2 +
0.20M
2OH0.40M
base
acid
[OH-] = 0.40 M
11.
[H+] = 2.5 x 10-14 M
H+
[H+] = 0.030M
pOH = 0.40
+
Cl0.030M
[OH-] = 3.3 x 10-13 M
pH
= 1.52
pOH = 12.48
Calculate the [H+], [OH-], pH and pOH for a 0.20 M NaOH solution.
NaOH
0.20M
→
[OH-] = 0.20 M
13.
= 13.60
Calculate the [H+], [OH-], pH and pOH for a 0.030 M HCl solution.
HCl →
0.030M
14.
pH
Na+ +
0.20M
OH0.20M
[H+] = 5.0 x 10-14 M
pH
= 13.30
pOH = 0.70
300.0 mL of 0.20 M HCl is added to 500.0 mL of water, calculate the pH of the solution.
300.0 x 0.20 M =
HCl
0.075 M
→
H+
0.075 M
Cl-
+
pH = -Log[H+] =
1.12
800.0
14.
200.0 mL of 0.020 M HCl is diluted to a final volume of 500.0 mL with water, calculate the pH.
HCl
200.0 x 0.020 M = 0.0080 M
→
H+
0.0080 M
Cl-
+
pH = -Log[H+] =
2.10
500.0
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.
Ba(OH)2
150.0 x 0.40 M =
500.0
→
0.12 M
pOH = -Log[OH-] = 0.62
Ba2+
+
0.12 M
0.24 M
2OH-
pH = 14.00 - pOH
= 13.38
16.
250.0 mL of 0.20 M Sr(OH)2 is diluted by adding 350.0 mL of water, calculate the pH of the
solution.
Sr(OH)2
250.0 x 0.20 M =
600.0
0.083 M
pOH = -Log[OH-] = 0.78
→
Sr2+
0.083 M
pH = 14.00 - pOH
+
2OH0.1667 M
= 13.22
17.
water.
Calculate the pH of a saturated solution of 0.40M Ba(OH)2 when 25 mL was added 25.0 mL of
Ba(OH)2
(25)0.40 M
(50)

Ba2+ +
2OH0.20 M 0.40 M
[OH-] = 0.40
pOH = 0.40
pH
= 13.60
WS # 10 pH Calculations for Weak Acids
Calculate the [H+], [OH-], pH, and pOH for 0.20 M HCN.
1.
HCN

I
0.20 M
0
C
x
x
x
E
0.20 - x
x
x
=
4.9 x 10-10
H+
+
CN-
0
x2
0.20 - x
x
= 9.9 x 10-6 M
[H+] = 9.9 x 10-6 M
[OH-] = 1.0 x 10-9 M
pH = 5.00
2.
Calculate the [H+], [OH-], pH, and pOH for 2.20 M HF.
3.
[H+] = 2.8 x 10-2 M
[OH-] = 3.6 x 10-13 M
pH = 1.56
+
Calculate the [H ], [OH ], pH, and pOH for 0.805 M CH3COOH.
pOH = 9.00
pOH = 12.44
[H+] = 3.8 x 10-3 M
[OH-] = 2.6 x 10-12 M
pH = 2.42
+
Calculate the [H ], [OH ], pH, and pOH for 1.65 M H3BO3.
4.
[H+] = 3.5 x 10-5 M
5.
[OH-] = 2.9 x 10-10 M
pOH = 11.58
pH = 4.46 pOH = 9.54
Calculate the pH of a saturated solution of Mg(OH)2.
Mg(OH)2
x

Mg2+ +
x
2OH2x
Ksp = [Mg2+][OH-]2
5.6 x 10-12 = 4x3
[OH-] = 2x = 2.22 x 10-4 M
pH = 10.35
Calculate the pH of a 0.200 M weak diprotic acid with a Ka = 1.8 x 10-6.
6.

H2 X
weak acid!!
H+
+
X-
I
0.200 M
0
0
C
x
x
x
E
0.20 - x
x
x
Small Ka approximation x = 0
x2
=
1.8 x 10-6
0.20
x
= 6.0 x 10-4 M
Note- only lose one proton for any
[H+] = 6.0 x 10-4 M
[OH-] = 1.7 x 10-11 M
pH = 3.22
pOH = 10.78
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.
Sr(OH)2
350.0 x 0.20 M =
800.0
→
0.0875 M
pOH = -Log[OH-] = 0.76
Sr2+
2OH-
+
0.0875 M
pH = 14.00 - pOH
0.175 M
= 13.24
WS # 11 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.
[H+] = 10-pH = 10-5.00 = 0.0000100 M
HCN

I
0.20 M
0
C
0.0000100 M
0.0000100 M
0.0000100 M
E
0.19999
0.0000100 M
0.0000100 M
Ka = (0.0000100)2
H+
+
CN-
0
=
5.0 x 10-10
0.19999
Ka = 5.0 x 10-10
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.
Ka = 3.5 x 10-4
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.
Ka = 1.8 x 10-5
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.
Ka = 7.3 x 10-10
5.
The pH of a 0.10 M diprotic acid is 3.683, calculate the Ka and identify the acid.
[H+] = 10-pH = 10-3.683 = 0.0002075 M
H2 X
one proton.

I
0.10 M
0
C
0.0002075 M
0.0002075 M
0.0002075 M
E
0.09979
0.0002075 M
0.0002075 M
Ka = (0.0002075)2
0.09979
Ka = 4.3 x 10-7
6.
H+
+
HX-
Note a diprotic weak acid only loses
0
=
4.3 x 10-7
Carbonic acid
H2CO3
The pH of 0.20 M NH3 is 11.227; calculate the Kb of the Base.
pOH
= 14.00
[OH-] =
10-pOH
- pH = 2.773
=
0.001686 M
Look up on Ka Table.
NH3
+
H2 O
⇄
NH4+
+
OH-
I
0.20 M
0
0
C
0.001686 M
0.001686 M
0.001686 M
E
0.1983 M
0.001686 M
0.001686 M
Kb= (0.001686)2
0.1983
=
1.4 x 10-5
7.
The pH of 0.40 M NaCN is 11.456; calculate the pH for the basic salt. Start by writing an equation
and an ICE chart.
pOH
= 14.00
[OH-] =
- pH = 2.544
10-pOH
CN-
+
=
0.002858 M
H2 O
⇄
+
0
OH-
I
0.40 M
C
0.002858 M
0.002858 M
0.002858 M
E
0.3971 M
0.002858 M
0.002858 M
Kb= (0.002858)2
0.3971
8.
HCN
=
0
2.0 x 10-5
The pH of a 0.10 M triprotic acid is 5.068, calculate the Ka and identify the acid.
[H+] = 10-pH = 10-5.068 = 8.55 x 10-6 M
H3 X

only loses one proton.
H+
+
I
0.10 M
C
8.55 x 10-6 M8.55 x 10-6 M8.55 x 10-6 M
E
0.10 M
Ka = (8.55 x 10-6)2
0
H2X-
0
8.55 x 10-6 M8.55 x 10-6 M
=
7.3 x 10-10
Note a triprotic weak acid
0.10
Ka = 7.3 x 10-10
9.
Boric acid
H3BO3
Look up on Ka Table.
How many grams of CH3COOH are dissolved in 2.00 L of a solution with pH = 2.45?
[H+]
=
10-2.45 =
⇄
CH3COOH
0.003548 M
H+
+
CH3COO-
I
x
0
0
C
0.003548 M
0.003548 M
0.003548 M
E
x
0.003548 M
0.003548 M
- 0.003548 M
Keq
=
1.8 x 10-5
=
[CH3COOH] =
[H+][CH3COO-]
[CH3COOH]
(0.003548)(0.003548)
[CH3COOH]
0.6994 M
2.00 L
x 0.6994 moles
x
60.0 g =
84 g
1L
1 mole
* Use questions 1 to 4 from last assignment to mark questions 1 to 4.
1.
2.
WS # 12
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!
1. NaNO2 (the basic ion is NO2-)
Kb(NO2-) = Kw
Ka(HNO2)
=
1.0 x 10-14
4.6 x 10-4
3.
Kb = 2.2 x 10-11
2.
2. KCH3COO (the basic ion is CH3COO-)
Kb = 5.6 x 10-10
3.
3. NaHCO3
Kb = 2.3 x 10-8
4. NH3
Kb = 1.8 x 10-5
5. NaCN
Kb = 2.0 x 10-5
6. Li2HPO4 Kb = 1.6 x 10-7
7. KH2PO4
Kb = 1.3 x 10-12
8. K2CO3
Kb = 1.8 x 10-4
9. Calculate the [H+], [OH-], pH, and pOH for 0.20 M H2CO3.
[H+] = 2.9 x 10-4 M
[OH-] = 3.4 x 10-11 M
pH = 3.53
pOH =
10.47
10. The pH of 0.20 M H2CO3 is 3.53. Calculate the Ka for H2CO3. Compare your calculated value with
that in the table.
Ka = 4.4 x 10-7
11. Calculate the [H+], [OH-], pH, and pOH for 0.10 M CH3COOH.
[H+] = 1.3 x 10-3 M
[OH-] = 7.5 x 10-12 M
pH = 2.87
pOH = 11.13
12. The pH of 0.10 M CH3COOH is 2.87. Calculate the Ka.
[H+] =
10-2.87 =
CH3COOH
0.001349 M
⇄
H+
+
CH3COO-
I
0.10 M
0
0
C
0.001349 M
0.001349 M
0.001349 M
E
0.09865 M
0.001349 M
0.001349 M
Ka
=
Ka
=
Ka
=
[H+][CH3COO-]
[CH3COOH]
(0.001349)( 0.001349)
(0.09865)
1.8 x 10-5

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.


H2SO4
+
2NaOH

Na2SO4
+
2HOH


0.200 L x 0.120 mol = 0.0240 mol
0.400 L x 0.140 mol = 0.0560 mol

L
L

Note the loss of sig fig
Beware subtraction!


















I
0.0240 mole
0.0560 mole
C
0.0240 mole
0.0480 mole
E
0
0.0080 mole
[OH-]
=
0.0080 mole =
0.6000 L
0.013 M
Note the final volume is used as the
two solutions are mixed together.
200.0 mL + 400.0 mL = 600.0 mL
pOH =
1.88
pH
12.12
=
WS # 13
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 CNKb(CN-) =
Kw
=
I
C
E
H2 O
x2
=

=
2.0408 x 10-5
4.9 x 10-10
Ka(HCN)
CN- +
0.20
x
0.20 - x
1.0 x 10-14
HCN
0
x
x
+
OH0
x
x
2.0408 x 10-5
0.20 - x
x = [OH-] = 2.0 x 10-3 M
[OH-] = 2.0 x 10-3 M
pOH = 2.69
2. 0.010 M NaHS (the basic ion is HS-)
pH = 11.31
[H+] = 4.9 x 10-12 M
Kb = 1.1 x 10-7 [OH-] = 3.3 x 10-5 M
pOH = 4.48
pH = 9.52
[H+] = 3.0 x 10-10 M
pOH = 5.21
pH = 8.79
[H+] = 1.6 x 10-9 M
pOH = 4.02
pH = 9.98
[H+] = 1.0 x 10-10 M
pOH = 2.49
pH = 11.51
[H+] = 3.1 x 10-12 M
3. 0.067 M KCH3COO
Kb = 5.55 x 10-10 [OH-] = 6.1 x 10-6 M
4. 0.40 M KHCO3
Kb = 2.3 x 10-8 [OH-] = 9.6 x 10-5 M
5. 0.60 M NH3
Kb = 1.786 x 10-5 [OH-] = 3.3 x 10-3 M
6. If the pH of a 0.10 M weak acid HX is 3.683, calculate the Ka for the acid and identify the acid using
your acid chart.
⇄
HX
H+
X-
I
0.100 M
0
0
C
- 0.0002075
0.0002075
0.0002075
E
0.09979
0.0002075
0.0002075
Ka
=
(0.0002075)2
(0.09979)
=
4.3 x 10-7
7. Calculate the [H+], [OH-], pH, and pOH for 0.80 M H3BO3.
Carbonic acid
[H+] = 2.4 x 10-5 M
[OH-] = 4.1 x 10-10 M
pH = 4.62
pOH = 9.38
8. Calculate the [H+], [OH-], pH, and pOH for 0.25 M H2CO3.
[H+] = 3.3 x 10-4 M
[OH-] = 3.0 x 10-11 M
pH = 3.48
pOH = 10.52
9. The pH of 1.65 M H3BO3 is 4.46. Calculate the Ka for H3BO3. Compare your calculated value with that
in the table.
Ka = 7.3 x 10-10
[OH-] = 2.88 x 10-10 M
pH = 4.46
[H+] = 3.47 x 10-5
pOH = 9.54
10. The pH of 0.65 M NaX is 12.46. Calculate the Kb for NaX.
[OH-] = 10-1.54 = 0.02884 M
pOH = 14.00 - 12.46 = 1.54
I
C
E
CN- +
0.65 M
0.02884 M
0.6212 M
H2 O

0
HCN
OH-
+
0
0.02884 M
0.02884 M
0.02884 M
0.02884 M
(0.02884)2
Kb =
(0.6212)
Kb =
1.3 x 10-3
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
Volume of HCl added
#1
37.80 mL
Reject
#2
35.49 mL
#3
35.51 mL
Calculate the original [HCl] = 1.04M
35.50 mL
Average
2HCl
0.03550 L
+
Ba(OH)2 → BaCl2
3.16 g
3.16 g Ba(OH)2 x
Molarity
+ 2H2O
1 mole
x
2 moles HCl
=
171.3g
1 mole Ba(OH)2
0.03550 L
[HCl] = 1.04M
12. Calculate the volume of 0.200M H2SO4 required to neutralize 25.0 ml of 0.100M NaOH.
0.00625 L
13. 25.0 ml of .200M HCl is mixed with 50.0 ml .100M NaOH, calculate the pH of the resulting solution.
No excess pH = 7.000
14. 10 ml of 0.10M H2SO4 is mixed with 25 ml 0.20M NaOH, calculate the pH of the resulting solution.
pH = 12.456
15. 125.0 ml of .200M HCl is mixed with 350.0 ml .100M NaOH, calculate the pH of the resulting
solution.
pH = 12.323
16. Define standard solution and describe two ways to standardize a solution.
A standard solution is one of known molarity. If you make the solution from a weighed amount of solid
and dilute it to a final volume in a volumetric flask it is a standard solution. If you titrate a solution to
determine its concentration it is a standard 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?
[H+] = 6.3 x 10-13 M
WS # 14 Review
1. List the properties of acids/bases.
Acids- conduct electricity, taste sour, change the color of indicators, neutralize bases, react with
active metals like Mg to produce H2 gas.
Bases- conduct electricity, taste bitter, change the color of indicators, neutralize acids, feel slippery.
2. Define the following:
Arhenius strong acid- completely ionizes to form H+
Arhenius weak base- partially ionizes to form OHBronsted strong acid- completely donates a proton to a base
Bronsted weak base- partially accepts a proton to an acid
Conjugate pair – an acid base pair that differs by one proton
Amphiprotic- a chemical species that can be an acid or base
Standard solution- a solution of known molarity
3. Show by calculation if the following amphiprotic ions are acids or bases:
a) HCO3Base
Ka = 5.6 x 10-11
Kb = 2.3 x 10-8
b) H2PO4Acid
Ka = 6.2 x 10-8
Kb = 1.3 x 10-12
2-13
c) HPO4
Base
Ka = 2.2 x 10
Kb = 1.6 x 10-7
4. What is the strongest base in water? What is the strongest acid in water? Write equations to explain
your answer.
Base OHNaOH → Na+ + OHAcid H+
HCl →
H+ +
Cl5. Match each equation:
Acid/base complete
Acid/base net ionic
Solubility product
Hydrolysis
Acid/Base formula
Ionization of water
HCl + NaOH →NaCl + HOH
F- + HOH → HF + OHH+ + OH- → HOH
AgCl(s) → Ag+ + ClH20 → H+ + OHH+ + Cl- + Na+ + OH-→Na++ Cl- + H2O
6. HCl and HF. Describe each acid as:
a) strong/weak b) high/low ionization c) large or small Ka
d) good/poor conductor e) strong or weak electrolyte
7. 0.2M HCl and 1.0M 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
SA
WA
WB
SB
9. Which ions are amphiprotic?
HPO4-2
HCl
FHS-
H2S
H2 O
10. Write the net ionic equation between any acid and base.
11. Write the ionization equation for water.
12. Write the Kw expression.
H+ + OH- → HOH
H20 → H+ + OH-
Kw = [H+][OH-] = 1.0 x 10-14
13. H2SO3 + HS- <====> H2S + HSO3a) Are the reactants or products favoured?
b) Are the Keq large, small or about 1?
14. 0.20M HCl
pH = 0.70
15. 0.20M Ba(OH)2
pH = 13.60
16. 0.20M H2CO3
pH = 3.53
17. 0.40M KHCO3
pH = 9.98
18. The pH increases by 2 units. How does [H+] change?
19. The pH decreases by 1 unit.
Decreases by a factor of 100
How does [H+] change? Increases by a factor of 10
20.
a) For distilled water :
pH = 7.00
pOH =7.00 [H+] = 1.0 x 10-7 M [OH-] = 1.0 x
10-7 M
b) For 1M HCl:
pH = 0.0
pOH =14.0 [H+] = 1 M
[OH-] = 1.0 x
-14
10 M
c) For 1M NaOH
pH = 14.0
pOH =0.0 [H+] = 1.0 x 10-14 M [OH-] = 1 M
21. The pH of .20M NaX is 12.50, calculate the Kb.
Kb = 5.9 x 10-3
22. The pH of .2M HX is 4.5, calculate the Ka.
Ka = 5 x 10-9
23. 100mL of .200M NaOH is mixed with 100ml of .180M HCl. Calculate the pH of the resulting
solution.
pH = 12.00
24. How many grams of NaHCO3 are required to make 100mL of .200M solution?
1.68 g
25. What volume of 0.200M NaOH is required to neutralize 25.0 mL of 0.150M H2SO4?
0.0375 L
26. In a titration 25.0mL of .200M H2SO4 is required to neutralize 10.0mL NaOH. Calculate the
concentration of the base.
1.00 M
27. Calculate the concentration of a solution of NaCl made by dissolving 50.0g in 250mL of water.
3.42 M
28.
SO3(g)
+ H2O(g)
⇄
H2SO4(l)
Equilibrium concentrations are found to be:
[SO3] = 0.400M
[ H2O] = 0.480M
[H2SO4] = 0.600M
Calculate the value of the equilibrium constant.
5.21
29.
2SO2(g) +
O2(g)
⇄
2SO3(g)
4.00 moles of SO2 and 5.00 moles O2 are placed in a 2.00 L container at 200ºC and allowed to reach
equilibrium. If the equilibrium concentration of O2 is 2.00M, calculate the Keq.
Keq = 0.500
Ws # 15 Buffers
1. Definition (buffer) A solution that is made by mixing a weak acid or base with a salt containing
the conjugate which maintains a relatively constant pH.
2.
Acid
Conjugate Base
Salt
HCN
CN-
NaCN
H2CO3
HCO3-
KHCO3
NH4+
NH3
NH4Cl
HF
F-
NaF
CH3COOH
CH3COO-
NaCH3COO
H2C2O4
HC2O4-
Na HC2O4
3. Write an equation for the first three buffer systems above.
HCN
⇄
H+
+
CN-
H2CO3
⇄
H+
+
HCO3-
H2 O
⇄
NH3
+
NH4+ +
OH-
4. Which buffer could have a pH of 4.0 ?
Which buffer could have a pH of 10.0 ?
a) HCl & NaCl
c)
b) HF & NaF
NH3
& NH4Cl
5. Predict how the buffer of pH = 9.00 will change. Your answers are 9.00, 8.98, 9.01, 2.00, and 13.00
a) 2 drops of 0.10M HCl are added
Final pH
8.98
b) 1 drop of 0.10M NaOH is added
9.01
c) 10 mL of 1.0 M HCl are added
2.00
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
left
⇄
H2CO3
[H+] =
increases
H+
HCO3-
+
[H2CO3] =
increases
[HCO3-] =
Indicators
1. Definition (Indicator) A weak acid whose conjugate base is a different color.
HInd ⇄ H+
2. Equilibrium equation
3. Colors for methyl orange
HInd red
Ind-
+
Ind-
yellow
4. Compare the relative sizes of [HInd] and [Ind-] at the following pH’s.
Color
Relationship
pH = 2.0
red
[Hind] > [Ind-]
pH = 3.7
orange
[Hind] = [Ind-]
pH = 5.0
yellow
[Hind] < [Ind-]
5. HCl is added to methyl orange, describe if each increases or decreases.
[H+]
increases
[HInd]
increases
[Ind-]
decreases
Color Change
yellow to red
6. NaOH is added to methyl orange, describe if each increases or decreases.
[H+]
decreases
[HInd]
decreases
[Ind-]
increases
Color Change
red to yellow
7. State two equations that are true at the transition point of an indicator.
[Hind] = [Ind-]
Ka = [H+]
decreases
8. What indicator has a Ka = 4 x 10-8 Neutral Red
9. What is the Ka for methyl orange. 2 x 10-4
10. A solution is pink in phenolphthalein and colorless in thymolphthalein. What is the pH of the
solution?
pH = 10
11. A solution is blue in bromothymol blue, red in phenol red, and yellow in thymol blue. What is the pH
of the solution?
pH = 8
Ws # 16 Titration Curves
Choose an indicator and describe the approximate pH of the equivalence point for each titration.
Complete each reaction.
pH
Indicator
1. HCl
+
NaOH ------->
7
bromothymol blue
2. HF
+
RbOH ------->
9
phenolphthalein
3. HI
+
Ba(OH)2 ------->
7
bromothymol blue
KOH ------>
9
phenolphthalein
+ NH3 ------->
5
bromocresol green
9
phenolphthalein
4. HCN
5. HClO4
+
6. CH3COOH
+
LiOH ------->
7. Calculate the Ka of bromothymol blue.
Ka = 2 x 10-7
8. An indicator has a ka = 1 x 10-10, determine the indicator. Thymolphthalein
9. Calculate the Ka of methyl orange.
Ka = 2 x 10-4
10. An indicator has a ka = 6.3 x 10-13, determine the indicator. Indigo Carmine
11. Explain the difference between an equivalence point and a transition point.
The equivalence point refers to endpoint of a titration (moles acid = moles base) and a transition point
refers to when an indicator changes color.
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
14
pH
7
0
0
50
100
50
0
Volume of base added
100
Volume of base added
14. Adding 100 ml 0.10 M HCl to 50 mL 0.10M NH3
M NaOH
15. Adding 100 ml .10 M HCl to 50 mL 0.10
14
pH
pH
7
0
0
50
100
0
Volume of base added
50
Volume of base added
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