AP Chemistry
AP Chemistry 12: Equilibrium—Acids and Bases Name ____________________________
A. Acids and Bases (16.1 to16.11)
1. three theories
a. Arrhenius: acid (H+) reacts with base (OH-) to form water (neutralization reaction):
H+ + OH- → H2O (not an equilibrium)
b. Brønsted-Lowry: proton donor (HA) transfers proton (H+) to proton acceptor (:B-)
1. includes above definition
2. HA + :B- Δ :A- + HB (equilibrium)
3. free H+ shown as hydronium ion: H3O+
c. Lewis: electron pair acceptor (M+) attaches to an electron pair donor (:B-)
1. includes above definitions
2. M+ + :B Δ M:B+ (complex ion formation)
3. H3B + :NH3 Δ H3BNH3 (incomplete octet)
2. acids
a. neutral molecule made from H+ + anion (HA)
b. name based on anion name
1. –ide anion: acid name is Hydro___ic acid
HCl (hydrochloric acid)
2. –ate anion: acid name is ___ic acid
HNO3 (nitric acid)
3. –ite anion: acid name is ___ous acid
HNO2 (nitrous acid)
c. when A- is Cl-, Br-, I-, NO3-, ClO3-, ClO4, SO42-
1. % ionization = [H+]/[HA]o x 100 = 100 %
2. strong acid = strong electrolyte
d. when A- is any other ion—weak acid (electrolyte)
1. % ionization < 5 %
2. organic acids, HC2H3O2 or CH3COOH
3. A- is oxyanion, HXOy (HClO, HBrO2, H2SO3)
a. H is bonded to O
b. weaker O-H bond = stronger acid
1. ↑ electronegativity ∴ (HClO > HBrO)
2. ↑ oxygen ∴ (HClO2 > HClO)
c. nonmetal oxides: CO2 + H2O → H2CO3
e. polyprotic acids (HxA)—first H+ is easiest to remove from neutral molecule (successive H+ are harder to remove from anion)
f. acid dissociation/ionization
1. HA(aq) Δ H+ + A-
2. Ka = [H+][A-]/[HA] = [H+]2/[HA]
3. acid strength: (larger Ka = stronger acid)
4. polyprotic acids (HxA)
a. H2A Δ H+ + HA-: Ka1 = [H+]2/[H2A]
b. HA- Δ H+ + A2-: Ka2 = [H+][A2-]/[HA-]
[HA-] = [H+] ∴ ka2 = [A2-]
3. bases
a. hydroxides (:OH-)
1. soluble: column 1, Ca2+, Sr2+, Ba2+ = strong
2. insoluble = antacids
3. oxides: M2O(s) + 2 H2O → 2 M+ + 2 OH-
b. ammonia or amines (:NH3, :NH2CH3, etc.)
c. anions from weak acids (:A-)
d. base dissociation/ionization
1. ammonia and its derivatives
NH3(aq) + H2O Δ NH4+ + OH-
Kb = [NH4+][OH-]/[NH3] = [OH-]2/[NH3]
(water is not included because it is liquid)
2. anions other than from strong acids
A- + H2O Δ HA(aq) + OH-
Kb = [HA][OH-]/[A-] = [OH-]2/[A-]
4. relationship between Ka and Kb for a conjugate pair
a. HX Δ H+ + X- Ka
X- + H2O Δ HX + OH- Kb
H2O Δ H+ + OH- Kw
b. ∴ Ka x Kb = Kw = 1 x 10-14
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5. when to use Ka and Kb
a. use Ka for any equation that has H+ (H3O+)
1. HA(aq) Δ H+ + A- Ka
2. H+ + A- Δ HA(aq) 1/Ka
b. use Kb for any equation that has OH-
1. A- + H2O Δ HA(aq) + OH- Kb
2. HA(aq) + OH- Δ A- + H2O 1/Kb
6. autoionization of water
a. H2O Δ H+ + OH-
1. Kw = [H+][OH-] = 1 x 10-14
2. pure water at 25oC: [H+] = [OH-] = 1 x 10-7 M
b. solutions: acidic: [H+] > [OH-], basic: [H+] < [OH-]
c. pH scale
1. pH = -log [H+], pOH = -log[OH-]
2. pH scale: acidic < 7, neutral = 7, basic > 7
3. pH + pOH = pKW = 14
4. significant figures (right of decimal): 3.45 (2 sf)
d. pKa + pKb = pKw = 14
7. pH of ionic compounds (salts) in water—hydrolysis
a. cation (except from strong bases) are acidic
M+ + H+•OH- Δ MOH + H+ ∴ acidic (pH < 7)
b. anion (except from strong acids) are basic
X- + H+•OH- Δ HX + OH- ∴ basic (pH > 7)
c. polyprotic acid anions (HXOy-) are acidic
HXOy- Δ H+ + XOy2- ∴ acidic (pH < 7)
d. overall pH of salt is the result of both ions
1. eliminate ion(s) from strong acids or bases
2. affect from remaining ion(s) is additive
8. acid/base reactions
a. Arrhenius neutralization reaction (produces H2O)
1. HA + MOH: H+ + OH- → H2O
2. HA + M2O: 2 H+ + Na2O(s) → 2 Na+ + 2 H2O
b. Brønsted-Lowry proton transfer reaction
1. HA + :B- Δ HB + :A-
a. conjugate pairs
acid (HA) Δ base (:A-)
base (:B-) Δ acid (HB)
b. stronger acid + stronger base →
weaker acid + weaker base
2. examples (conjugate pairs)
|acid |+ |base |Δ |acid |+ |base |K |
|HA |+ |H2O |Δ |H3O+ |+ |A- |1/Ka |
|H3O+ |+ |B |Δ |HB+ |+ |H2O |Ka |
|HA |+ |OH- |Δ |H2O |+ |A- |1/Kb |
|H2O |+ |B |Δ |HB+ |+ |OH- |Kb |
3. amphiprotic
a. can act as a proton donor and proton acceptor, depending on other reactant
b. HCO3-, HSO3-, H2O, NH3
B. Acid-Base Equilibrium Problems
1. pure acid/base equilibrium problems
a. determine Ka, given pH or [H+]E and [HA]o
• set up "ICE Box" (shaded boxes are given)
|[ ] | HA Δ H+ + A- |
|I |[HA]o |0 |0 |
|C |–[H+]E |+[H+]E |+[H+]E |
|E |[HA]o – [H+]E |[H+]E |[H+]E |
• solve for Ka = [H+]E[A-]E/[HA]E = [H+]E2/([HA]E
determine Kb, given pOH or [OH-]E and [B]o
• set up "ICE Box" (shaded boxes are given)
|[ ] | B + H2O Δ HB+ + OH- |
|I |[B]o | |0 |0 |
|E |
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b. determine all [ ]E, given [HA]o and Ka
• set up "ICE Box" (shaded boxes are given)
| [ ] | HA Δ H+ + A-|
|I |[HA]o |0 |0 |
|C |–x |+x |+x |
|E |≈ [HA]o (x < 5%) |x |x |
• solve for x (Ka = [H+]E[A-]E/[HA]E = x2/[HA]o)
• solve for [ ]E
determine all [ ]E, for a polyprotic acid, given [H2A]o, Ka1 and Ka2
• set up "ICE Box" (shaded boxes are given)
| [ ] | H2A Δ H+ + HA- |
|I |[H2A]o |0 |0 |
|C |–x1 |+x1 |+x1 |
|E |≈ [H2A]o (x1 < 5%) |x1 |x1 |
• solve for x1 (Ka = x12/[H2A]o)
• set up next "ICE Box"
|[ ] | HA- Δ H+ + |
| |A2- |
|I |x1 |x1 |0 |
|C |–x2 |+x2 |+x2 |
|E |≈ x1 |≈ x1 |x2 |
• solve for x2 (Ka2 = (x1)(x2)/(x1) = x2)
• solve for [ ]E
determine [ ]E, given [B]o and Kb
• set up "ICE Box" (shaded boxes are given)
|[ ] | B + H2O Δ HB+ + OH- |
|I |[B]o |
|I |[HA]o |0 |[A-]o |
|C |–x |+x |+x |
|E |≈ [HA]o |x |≈ [A-]o |
• solve for x (Ka = (x)([A-]o)/[HA]o) = [H+]E
determine [OH-]E, given [HA]o, [A-]o and Kb
• set up "ICE Box" (shaded boxes are given)
| [ ] | A- + H2O Δ HA + OH- |
|I |[A-]o |
|I |[HA]o |0 |x |
|C |–[H+]E |+[H+]E |+[H+]E |
|E |≈ [HA]o |[H+]E |≈ x |
• solve for x= [A-]o (Ka = (x)([H+]E/[HA]o)
determine [HA]o, given [A-]o, equilibrium pOH and Kb
• set up "ICE Box" (shaded boxes are given)
|[ ] | A- + H2O Δ HA + OH- |
|I |[A-]o | |x |0 |
|E |
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C. Buffer Systems (17.2)
1. mixture of weak acid or base with its conjugate
2. addition of strong acid or base to buffer
a. addition of acid: A- + H+ → HA(aq)
b. addition of base: HA(aq) + OH- → A- + H2O
c. process is reversible, but not equilibrium
3. determine [H+] or [OH-] after addition of acid or base
|determine original moles of HAo, A-o |
|nHA-o = Vsolution x [HA]o |
|nA-o = Vsolution x [A-]o |
|determine moles of H+ or OH- added |
|nH+ = (VH+)(MH+) |
|nOH- = (VOH-)(MOH-) |
|calculate equilibrium moles of HAE and A-E |
|nHA-E = nHA-o – nOH- or nHA-o + nH+ |
|nA-E = nA-o + nOH- or nA-o – nH+ |
|calculate [H+] or [OH-] |
|moles buffer > moles H+/OH-: [H+] = Ka(nHA/nA-) |
|moles buffer < moles H+/OH- |
|[H+] = (nH+ – nA-o)/Vtot |
|[OH-] = (nOH- – nHAo)/Vtot |
D. Acid-Base Titration (17.3)
1. acid or base is added to a fixed amount of base or acid
a. pH is monitored using a pH meter
b. equivalence when moles of H+ = moles OH-
1. nH+MaVa = nOH-MbVb
2. indicators
a. bromthymol blue (BB) for pH = 7
b. methyl red (MR) for pH < 7
c. phenolphthalein (PP) for pH > 7
c. buffered solution during incomplete neutralization
2. graphs
[pic]
a. upper line: graphs addition of HCl to strong base (equivalence at pH = 7)
b. lower line graphs addition of HCl to weak base
1. equivalence: middle of vertical section (pH = 9)
2. buffer region: mini-plateau (10 mL to 40 mL)
3. [H+] or [OH-] calculation chart
| |1. SA + SB |2. WA + SB |3. WB + SA |
|initial |[H+] = |[H+] = |[OH-] = |
| |[HA] |(Ka[HA])½ |(Kb[B])½ |
|buffer |[H+] = |[H+] = |[OH-] = |
|(2 and 3) |(nHA - nOH-)/Vtot |Ka(nHA/nA-) |Kb(nB/nHB+) |
|equivalence |[H+] = |[OH-] = |[H+] = |
| |1 x 10-7 M |(Kb[A-])½ |(Ka[HB+])½ |
|excess |[OH-] = |[OH-] = |[H+] = |
| |(nOH- - nHA)/Vtot |(nOH- - nHA)/Vtot |(nH+ - nB)/Vtot |
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Experiments
1. Acid-Base Properties (Wear Goggles)—Observe the properties of acids and bases in a variety of indicators and the acid/base properties of salts.
Acid/base indicators: Add 10 drops of 0.1 M HCl to 4 separate wells. Dip red litmus paper into well #1, dip blue litmus paper into well #2, add a drop of universal indicator to well #3 and add a drop of phenolphthalein to well #4; record color colors of each well. Repeat with 0.1 M HC2H3O2, NaOH and NH3.
a. Record the colors in the space provided.
|Acid/Base |red |blue |Universal Indicator |Phenol-phthal|
| |litmus |litmus | |ein |
| | | |Color |pH | |
|HCl | | | | | |
|HC2H3O2 | | | | | |
|NaOH | | | | | |
|NH3 | | | | | |
Determine the pH of some salts: Add 10 drops of the salt listed in table 3 to separate wells. Add 1 drop of universal indicator and record the pH.
b. Complete the table for the four salts tested and compare the predicted with actual pH range.
|Well # |Salt |Spectator |Predicted pH |Actual pH |
| | |ion(s) |Range | |
|1 |Zn(NO3)2 | | | |
|2 |NaCl | | | |
|3 |Na3PO4 | | | |
|4 |NH4C2H3O2 | | | |
c. Complete the table by filling in the cation groups and anion groups that are neutral, acidic or basic.
|Acid-Base Property |Cation |Anion |
|Neutral | | |
|Acidic | | |
|Basic | | |
2. Buffer Lab (Wear Goggles)—Make buffers using two techniques and examine their ability to "buffer" a solution.
Test the effect of adding HCl and NaOH to distilled water: Add 20 mL of distilled water to each of two 50 mL beakers labeled A and B. Measure the pH using the pH meter. Add 20 drops of 0.10 M HCl to beaker A and measure pH. Add 20 drops of 0.10 M NaOH to beaker B and measure pH.
a. Record the pH of pure water and after HCl and NaOH are added.
|Distilled Water |After adding HCl |After adding NaOH |
| | | |
Partially neutralized weak acid technique: Add 50. mL of 0.10 M HC2H3O2 to a 150 mL beaker. Place the pH meter in the solution and gently stir the solution while adding single drops of 6 M NaOH. Stop adding NaOH when the pH = 5.0. Add 20 mL of the pH 5 buffered solution to beaker A and beaker B. Add 20 drops of 0.10 M HCl to beaker A and measure the pH. Add 20 drops of 0.10 M NaOH to beaker B and measure the pH.
b. Record the pH of the buffered solution and after HCl and NaOH are added.
|Buffered solution |After adding HCl |After adding NaOH |
| | | |
c. Calculate the amount of NH4NO3 needed for 50 mL of a pH 9 buffer made with 0.10 M NH3 (Kb = 1.8 x 10-5).
|pOH | |
|[OH-] | |
|[NH4+] | |
|mol NH4+ | |
|mass NH4NO3 | |
Add salt of the conjugate acid to a weak base technique: Add the calculate amount of NH4NO3 to 50 mL of 0.10 M NH3 in a 150 mL beaker. Measure the pH. Add 20 mL of the pH 9 buffer to beaker A and beaker B. Add 20 drops of 0.10 M HCl to beaker A and measure the pH. Add 20 drops of 0.10 M NaOH to beaker B and measure the pH.
d. Record the pH of the buffered solution and after HCl and NaOH are added.
|Buffered solution |After adding HCl |After adding NaOH |
| | | |
e. How effective were the buffered solutions in controlling changes in pH compared to pure water?
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3. pH Profile Lab (Wear Goggles)—Monitor the pH while performing acid-base titrations and compare the pH profile to the theoretical graphs.
Add 25 mL of 0.10 M HCl to a 150 mL beaker. Measure the pH of the solution. Add increments of 10 mL of 0.10 M NaOH until a total of 50 mL of 0.10 M NaOH has been added. Measure the pH after each 10 mL aliquot is added.
a. Record the pH of the solution and after each 10 mL aliquot of NaOH.
|mL of NaOH Added |
|0 |10 |20 |30 |40 |50 |
| | | | | | |
b. Graph the pH vs. mL of NaOH added.
|pH |12 | |
| mL of 0.10 NaOH added to 25 mL of 0.10 M HCl |
Add 25 mL of 0.10 M HC2H3O2 to a 150 mL beaker. Measure the pH of the solution. Add increments of 10 mL of 0.10 M NaOH until a total of 50 mL of 0.10 M NaOH has been added. Measure the pH after each 10 mL aliquot is added.
c. Record the pH of the solution and after each 10 mL aliquot of NaOH.
|mL of NaOH Added |
|0 |10 |20 |30 |40 |50 |
| | | | | | |
d. Graph the pH vs. mL of NaOH added.
|pH |12 | |
| mL of 0.10 NaOH added to 25 mL of 0.10 M HC2H3O2 |
Add 25 mL of 0.10 M NH3 to a 150 mL beaker. Measure the pH of the solution. Add increments of 10 mL of 0.10 M HCl until a total of 50 mL of 0.10 M HCl has been added. Measure the pH after each 10 mL aliquot is added.
e. Record the pH of the solution and after each aliquot.
|mL of HCl Added |
|0 |10 |20 |30 |40 |50 |
| | | | | | |
f. Graph the pH vs. mL of NaOH added.
|pH |12 | |
| mL of 0.10 HCl added to 25 mL of 0.10 M NH3 |
g. Determine the theoretical pH when 10 mL aliquots of 0.10 M NaOH are added to 25 mL of 0.10 M HCl.
|Excess H+ |
|a. VH+ |0.025 |0.025 |0.025 |
|b. VOH- |0 |0.010 |0.020 |
|c. moloH+ = MH(a) | | | |
|d. moloOH- = MOH(b) | | | |
|e. molEH+ = (c – d) | | | |
|f. Vtot = (a + b) | | | |
|g. [H+]E = (e/f) | | | |
|h. pH = -log(g) | | | |
|Equivalence (NaCl(aq)) |
|a. VH+ |0.025 |
|b. VOH- |0.025 |
|c. moloH+ = MH(a) | |
|d. moloOH- = MOH(b) | |
|e. pH of NaCl(aq) |7.00 |
|Excess OH- |
|a. VH+ |0.025 |0.025 |0.025 |
|b. VOH- |0.030 |0.040 |0.050 |
|c. moloH+ = MH(a) | | | |
|d. moloOH- = MOH(b) | | | |
|e. molEOH- = (d – c) | | | |
|f. Vtot = (a + b) | | | |
|g. [OH-]E = (e/f) | | | |
|h. pOH = -log(g) | | | |
|i. pH = 14 – (h) | | | |
h. Graph the pH (y-axis) vs. mL of NaOH added (x-axis). Mark equivalence point and buffer region on the graph.
|pH |12 | |
| mL of 0.10 NaOH added to 25 mL of 0.10 M HCl |
i. Determine the pH when 10 mL aliquots of 0.10 M NaOH are added to 25 mL of 0.10 M HC2H3O2 (Ka = 1.8 x 10-5).
|Excess H+ |
|(Pure HA) (Buffer) |
|a. VHA |0.025 |0.025 |0.025 |
|b. VOH- |0 |0.010 |0.020 |
|c. moloHA = MHA(a) | | | |
|d. moloOH- = MOH(b) | | | |
|e. molEHA = (c – d) | | | |
|f. molEA- = (d) | | | |
|g. [HA]E = (e/a) | | | |
|h. [H+]E = (Ka(g))½ | | | |
|i. [H+]E = Ka(h/a) | | | |
|j. pH = -log(h or i) | | | |
|Equivalence (pure A-) |
|a. VHA |0.025 |
|b. VOH- |0.025 |
|c. moloH+ = MHA(a) | |
|d. moloOH- = MOH(b) | |
|e. molEA- = (d) | |
|f. Vtot = (a + b) | |
|g. [A-]E = (e/f) | |
|h. [OH-]E = (Kb(g))½ | |
|i. pOH = -log(h) | |
|j. pH = 14 – (i) | |
|Excess OH- |
|a. VH+ |0.025 |0.025 |0.025 |
|b. VOH- |0.030 |0.040 |0.050 |
|c. moloH+ = MH(a) | | | |
|d. moloOH- = MOH(b) | | | |
|e. molEOH- = (d – c) | | | |
|f. Vtot = (a + b) | | | |
|g. [OH-]E = (e/f) | | | |
|h. pOH = -log(g) | | | |
|i. pH = 14 – (h) | | | |
j. Graph the pH (y-axis) vs. mL of NaOH (x-axis). Mark equivalence point and buffer region on the graph.
|pH |12 | |
| mL of 0.10 NaOH added to 25 mL of 0.10 M HC2H3O2 |
k. Determine the pH when 10 mL aliquots of 0.10 M HCl are added to 25 mL of 0.10 M NH3 (Kb = 1.8 x 10-5).
|Excess OH- |
|(Pure HA) (Buffer) |
|a. VB |0.025 |0.025 |0.025 |
|b. VH+ |0 |0.010 |0.020 |
|c. moloB = MB(a) | | | |
|d. moloH+ = MH(b) | | | |
|e. molEB = (c – d) | | | |
|f. molEHB+ = (d) | | | |
|g. [B]E = (e/a) | | | |
|h. [OH-]E = (Kb(g))½ | | | |
|i. [OH-]E = Ka(h/a) | | | |
|j. pOH = -log(h or i) | | | |
|k. pH = 14 – (j) | | | |
|Equivalence (pure HB) |
|a. VB |0.025 |
|b. VH+ |0.025 |
|c. moloB = MB(a) | |
|d. moloH+ = MH(b) | |
|e. molEHB+ = (d) | |
|f. Vtot = (a + b) | |
|g. [HB+]E = (e/f) | |
|h. [H+]E = (Ka(g))½ | |
|i. pH = -log(h) | |
|Excess H+ |
|a. VB |0.025 |0.025 |0.025 |
|b. VH+ |0.030 |0.040 |0.050 |
|c. moloB = MB(a) | | | |
|d. moloH+ = MH(b) | | | |
|e. molEH+ = (d – c) | | | |
|f. Vtot = (a + b) | | | |
|g. [H+]E = (e/f) | | | |
|h. pH = -log(g) | | | |
l. Graph the pH (y-axis) vs. mL of HCl added (x-axis). Mark equivalence point and buffer region on the graph.
|pH |12 | |
| mL of 0.10 HCl added to 25 mL of 0.10 M NH3 |
m. Compare the theoretical pH profiles with the actual pH profiles and complete the following chart.
|Experiment |Comparison between Graphs |
| |Equivalence pH |Buffer Region |Overall Shape |
|SA + SB | | | |
|WA + SB | | | |
|SA + WB | | | |
n. Select the indicator that is most effective for each titration given the pH at equivalence.
|Indicator |pH |SA + SB |SA + WB |WA + SB |
|Methyl Red |5 | | | |
|Bromthymol Blue |7 | | | |
|Phenolphthalein |9 | | | |
Practice Problems
A. Acid-Base
1. Define acid and base, and write a general acid-base reaction for each of the three theories.
|Theories |Acid |Base |Reaction |
|Arrhenius | | | |
|Brønsted-Lowry | | | |
|Lewis | | | |
2. Complete the chart with the formula or name of the acid.
|H2SO4 | |HClO2 |
| |hydrofluoric acid | |
|HCl | |HNO2 |
| |nitric acid | |
| |H2SO3 | |
|hydroiodic acid | |perchloric acid |
3. Complete the chart for a 1 M solution of acid.
|Ionization reaction |% ionization |[HX] |[H+] |[X-] |
|HCl(g) → |100 % | | | |
|HF(g) Δ |8 % | | | |
|HCHO2(l) Δ |4 % | | | |
4. Rank the following acids from strongest to weakest:
|Acid |HBrO |HClO |HClO2 |HIO |
|Rank | | | | |
5. Complete the chemical equation and write the Ka expression for the following weak acids.
|Equation |Ka expression |
|HC2H3O2(aq) Δ | |
|H2PO4- Δ | |
|H2CO3(aq) Δ | |
|HCO3-(aq) Δ | |
6. Ka1 for H2SO3 is 1.3 x 10-2 and Ka2 is 6.3 x 10-8. Write each dissociation equation and the overall equation. Calculate K for the reaction H2SO3 Δ 2 H+ + SO32-.
| |
| |
7. Complete the chemical equation and write the Kb expression for the following weak bases.
|Equation |Ka expression |
|F-(aq) + H2O Δ | |
|CH3NH2(aq) + H2O Δ | |
8. Given the following:
HC2H3O2(aq) Δ H+ + C2H3O2- Ka = 1.8 x 10-5
HCN(aq) Δ H+ + CN- Ka = 4.0 x 10-10
a. Calculate Kb for C2H3O2- and CN-
|C2H3O2- | |
| CN- | |
b. Circle which is stronger.
|HC2H3O2 or HCN |C2H3O2- or CN- |
c. Complete the statement about the relative strength of the acid and base of a conjugate pair.
|A strong conjugate acid makes a ________ conjugate base. |
d. What is K for the equilibrium below?
HC2H3O2(ag) + CN-(aq) Δ C2H3O2-(aq) + HCN(aq)
| |
e. Starting with all reactants and products at 1 M, which way will the reaction proceed to reach equilibrium?
| |
f. Complete the statement about the relative strength of the acids and bases in a Brønsted-Lowry system.
|______ acid + ______ base → ______ acid + ______ base |
9. Kb for NH3 is 1.8 x 10-5. Determine
a. Ka for NH4+.
| |
b. K for the reactions.
|Reaction |K |
|NH3(aq) + H2O Δ NH4+ + OH- | |
|NH4+ + OH- Δ NH3(aq) + H2O | |
|NH3(aq) + H+ Δ NH4+ | |
|NH4+ Δ H+ + NH3(aq) | |
10. Ka for HF is 6.9 x 10-4. Determine
a. Kb for F-:
| |
b. K for the reactions.
|Reaction |K |
|HF(aq) Δ H+ + F- | |
|H+ + F- Δ HF(aq) | |
|HF(aq) + OH- Δ F- + H2O | |
|F- + H2O Δ HF(aq) + OH- | |
11. Complete the following:
|H2O(l) Δ ____ + ____ |H+ + H2O(l) → ____ |
|[H+] x [OH-] = 1 x 10-14 ∴ [H+] = [OH-] = 1 x 10-7 M |
|acid solutions [H+] _ [OH-] and base solutions [H+] _ [OH-] |
|pH = ________, pOH = ________ and pH + pOH = ___ |
12. fill in the range of values for each solution.
|Solution |Acid |Neutral |Base |
|pH | | | |
|[H+] | | | |
|[OH-] | | | |
13. Solve for the missing values.
|pH |[H+] |pOH |[OH-] |
|4.20 | | | |
| |3.0 x 10-9 | | |
14. Classify the salts as acidic, basic, neutral or can't tell.
|NaCl |Cu(NO3)2 |KNO2 |NH4F |
| | | | |
|NaClO |BaCl2 |Cu(C2H3O2)2 |LiF |
| | | | |
15. Predict the products, and then balance the following neutralization reactions. Write as net ionic equations.
|nitric acid + | |
|sodium hydroxide | |
|hydrofluoric acid + | |
|strontium hydroxide | |
|sulfuric acid + | |
|ammonia | |
|perchloric acid + calcium | |
|oxide | |
|sulfur trioxide(g) + | |
|potassium hydroxide | |
|hydrochloric acid + | |
|sodium sulfide | |
|hydrobromic acid + sodium | |
|bicarbonate | |
|ammonium nitrate + sodium | |
|hydroxide | |
|nitrogen dioxide(g) + | |
|potassium hydroxide | |
|nitric acid + | |
|methylamine | |
|sufuric acid + calcium | |
|hydroxide | |
|carbon dioxide + | |
|water | |
16. Explain the observations using chemical equations.
a. Statues made of marble (CaCO3) that are displayed in polluted (acidic air) cities lose their definition over time.
| |
b. Milk of Magnesia is a medication that absorbs excess stomach acid contains Mg(OH)2.
| |
17. Complete the equation, label the acids (A) and bases (B), and link the conjugate pairs.
|NH4+ + H2O Δ |NO2- + H2O Δ |
|HNO2 + H2O Δ |HNO2 + NH3 Δ |
|NH3 + H2O Δ |NH4+ + OH- Δ |
|HNO2 + OH- Δ |NH3 + H3O+ Δ |
B. Acid-Base Equilibrium
18. 1.369 g of HClO2 is dissolved in enough water to make 100. mL of solution. The pH is 1.36. Determine
a. Initial concentration of HClO2.
| |
b. Equilibrium concentration of H+.
| |
c. Ka.
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
19. In the first-step ionization of phosphorous acid the acid is 33.3% dissociated in a 0.300 M solution of H3PO3. Calculate Ka for the first-step ionization of H3PO3.
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
20. HC2O4- Δ H+ + C2O42-
7.35% of the HC2O4- is dissociated in a 0.0100 M solution. Calculate Ka for the second-step dissociation of oxalic acid.
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
21. Acetic Acid (HC2H3O2) has Ka = 1.8 x 10-5. Determine
a. The [H+] in 0.100 M solution.
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
b. The percent ionization.
|% = [H+]/[HA]o x 100 = |
c. Was the assumption that the equilibrium concentration of HC2H3O2 is 0.100 M valid? Explain your answer.
| |
22. 0.50 mol of phenol (HOC6H5) in 5.0 L has Ka = 1.6 x 10-10.
a. Determine [H+].
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
b. Determine pH.
| |
23. For H2SO3, Ka1 = 1.3 x 10-2 and Ka2 = 6.3 x 10-8.
a. Calculate [H+] for 0.100 M H2SO3.
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
b. What are the concentrations of HSO3- and SO32-?
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
24. For phosphoric acid (H3PO4) Ka1 = 7.1 x 10-3, Ka2 = 6.2 x 10-8 and Ka3 = 4.5 x 10-13. Determine for 0.100 M H3PO4.
a. [H+] and [H2PO4-]
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
b. [HPO42-]
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
c. [PO43-]
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
25. 0.10 M methylamine, CH3NH2, has Kb = 5.0 x 10-4. Determine [OH-]E.
|[ ] | |
|I | | | | |
|C | | | | |
|E | | | | |
| |
26. 6.6 g of hydroxylamine, HONH2, are in a 1-liter solution.
a. Determine the initial concentration of HONH2.
| |
b. Determine [OH-]E given Kb = 9.1 x 10-9
|[ ] | |
|I | | | | |
|C | | | | |
|E | | | | |
| |
27. 4.00 g of NaF are in 0.500 L solution.
a. Determine the initial concentration of NaF.
| |
b. Determine Kb for F- given that Ka for HF is 6.7 x 10-4.
| |
c. Determine the equilibrium concentration of OH-.
|[ ] | |
|I | | | | |
|C | | | | |
|E | | | | |
| |
| |
d. Determine the pH of the solution.
| |
28. 2.70 g of HCN (Ka = 4.0 x 10-10) and 2.45 g of NaCN are added to water to make 1.00 L of solution
a, What are the initial concentrations of HCN and CN-?
| |
b. What is the pH of the solution?
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
| |
| |
29. Calculate [H+] for a solution that is 0.10 M in both H2CO3 and NaHCO3. The Ka for H2CO3 is 4.2 x 10-7.
|[ ] | |
|I | | | |
|C | | | |
|E | | | |
| |
C. Buffer Systems
30. A liter of benzoic acid (HBen) and sodium benzoate (Ben-) buffer is prepared using 1.00 mol of benzoic acid and 0.50 mol of sodium benzoate. The Ka for HBen is 6.3 x 10-5.
a. 0.100 mol of H+ is added. Calculate
|mol HBen | |
|mol Ben- | |
|[H+]E | |
|pH | |
b. 0.100 mol of OH- is added. Calculate
|mol HBen | |
|mol Ben- | |
|[H+]E | |
|pH | |
31. A buffer is prepared by adding 2.00 mol of HC2H3O2 and 2.00 mol of NaC2H3O2 to enough water to make 5.00 L of solution. Acetic acid has a Ka 1.8 x 10-5. Calculate
a. the pH of the buffer.
| |
| |
| |
b. the pH when 0.50 mol of H+ is added.
|mol HC2H3O2 | |
|mol C2H3O2- | |
|[H+]E | |
| | |
|pH | |
c. the pH when 0.50 mol of OH- is added.
|mol HC2H3O2 | |
|mol C2H3O2- | |
|[H+]E | |
| | |
|pH | |
Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
1. What is the correct formula for sulfurous acid?
(A) H2S (B) HSO3 (C) H2SO3 (D) H2SO4
| | |
2. Which is the net ionic equation for the neutralization reaction between barium hydroxide and ammonium chloride?
(A) H+ + OH- → H2O (C) OH- + NH4+ → NH4OH
(B) H+ + NH3 → NH4+ (D) OH- + NH4+ → NH3 + H2O
| | |
3. Which 1-molar solution is the most basic?
(A) NaNO3 (B) Na2CO3 (C) NaCI (D) NaHSO4
| | |
Questions 4-9 6.0 g HC2H3O2 (MM = 60 g/mol) is dissolved in water to make 50. mL solution.
4. What is the concentration of C2H3O2-?
(A) 5.0 M (B) 2.0 M (C) 0.50 M (D) 0.20 M
| | |
5. The [H+] is 3.0 x 10-3 M, what is the approximate pH?
(A) 2.00 (B) 2.50 (C) 3.00 (D) 4.70
| | |
6. What is Ka for HC2H3O2?
(A) 9.0 x 10-6 (B) 1.8 x 10-5
(C) 4.5 x 10-5 (D) 9.0 x 10-4
| | |
7. What is Kb for C2H3O2-?
(A) 1.1 x 10-9 (B) 5.6 x 10-10
(C) 2.2 x 10-10 (D) 1.1 x 10-11
| | |
8. Enough solid NaC2H3O2 is added to make the solution equal molar in both HC2H3O2 and C2H3O2-. What is the approximate pH of the resulting solution?
(A) 2.00 (B) 2.50 (C) 3.00 (D) 4.70
| | |
9. Additional NaC2H3O2 is added to completely neutralize the HC2H3O2. What is the approximate pH of the resulting solution?
(A) 2 (B) 4 (C) 7 (D) 9
| | |
10. Which is generally true as the number of oxygen atoms increases in a series of acids, such as HXO, HXO2, HXO3?
(A) The acid strength decreases only if X is a nonmetal.
(B) The acid strength decreases only if X is a metal.
(C) The acid strength decreases.
(D) The acid strength increases.
| | |
11. HSO4– + H2O Δ H3O+ + SO42–
In the equilibrium represented above, the species that act as bases include which of the following?
I. HSO4- II. H2O III. SO42-
(A) II only (B) II only (C) III only (D) II and III
| | |
12. HC2H3O2 + CN- Δ HCN + C2H3O2–
The equilibrium constant, K = 3.7 x 104. Which can be concluded from this information?
(A) CN- is a stronger base than C2H3O2-.
(B) HCN is a stronger acid than HC2H3O2.
(C) The conjugate base of CN- is C2H3O2-.
(D) K will increase with an increase in temperature.
| | |
13. Equal volumes of 0.10 M H3PO4 and 0.20 M KOH are mixed. After equilibrium is established, the type of ion in solution in largest concentration, other than the K+ ion, is
(A) H2PO4- (B) HPO42- (C) PO43- (D) OH-
| | |
14. Which of the following species is in the greatest concentration in a 0.100 M solution of H2SO4 in water?
(A) H2SO4 (B) H3O+ (C) HSO4– (D) SO42–
| | |
15. Which of the following reactions does NOT proceed significantly to the right in aqueous solutions?
(A) H3O+ + OH– → 2 H2O
(B) HCN + OH– → H2O + CN–
(C) H2SO4 + H2O → H3O+ + HSO4–
(D) H2O + HSO4– → H2SO4 + OH–
| | |
16. Which of the following is NOT amphiprotic?
(A) HCO3- (B) H2PO4- (C) NH4+ (D) H2O
| | |
17. A molecule or an ion is classified as a Lewis acid if it
(A) accepts a proton from water
(B) accepts a pair of electrons to form a bond
(C) donates a pair of electrons to form a bond
(D) donates a proton to water
| | |
18. When phenolphthalein is used as the indicator in a titration of HCl with NaOH, it undergoes a rapid color change from clear to red at the end point of the titration because
(A) phenolphthalein is a very strong acid that is capable of rapid dissociation
(B) the solution being titrated undergoes a large pH change near the equivalence point of the titration
(C) phenolphthalein undergoes an irreversible reaction in basic solution
(D) OH- acts as a catalyst for the decomposition of phenolphthalein
| | |
19. The pH of 0.1-molar ammonia is approximately
(A) 1 (B) 7 (C) 11 (D) 14
| | |
20. What is NH4+ in the reaction: 2 NH3 Δ NH4+ + NH2-?
(A) a catalyst
(B) both an acid and a base
(C) the conjugate acid of NH3
(D) the reducing agent
| | |
21. At 25°C, an aqueous solutions with a pH of 8 has a [OH-] of
(A) 10-14 M (B) 10-8 M (C) 10-6 M (D) 1 M
| | |
22. In the titration of a weak acid of unknown concentration with a standard solution of a strong base, a pH meter was used to follow the progress of the titration. Which of the following is true for this experiment?
(A) The pH is 7 at the equivalence point.
(B) The pH at the equivalence point depends on the indicator used.
(C) The graph of pH versus volume of base added rises gradually at first and then much more rapidly.
(D) The graph of pH versus volume of base added shows no sharp rise.
| | |
23. How can 100. mL of NaOH solution with a pH of 13 be converted to a NaOH solution with a pH of 12?
(A) Add 10 mL of distilled water to the 100 mL of NaOH.
(B) Add 100 mL of distilled water to the 100 mL of NaOH.
(C) Add 900 mL of distilled water to the 100 mL of NaOH.
(D) Add 100. mL of 0.10 M HCI to the 100 mL of NaOH.
| | |
24. A 0.2 M solution of a weak monoprotic acid, HA, has a pH of 3. The ionization constant of this acid is
(A) 5 x 10-7 (B) 2 x 10-7 (C) 5 x 10-6 (D) 5 x 10-3
| | |
Questions 25-26 Oxalic acid, H2C2O4, is a diprotic acid with
K1 = 5 x 10-2 and K2 = 5 x 10-5.
25. Which equals the equilibrium constant for the reaction: H2C2O4 + 2 H2O Δ 2 H3O+ + C2O42–?
(A) 5 x 10-2 (B) 5 x 10-5 (C) 25 x 10-7 (D) 5 x 10-7
| | |
26. Which species is in highest concentration in 0.1 M H2C2O4?
(A) H2C2O4 (B) H3O+ (C) HC2O4- (D) C2O42-
| | |
27. If the acid dissociation constant, Ka, for an acid HA is
8 x 10-4 at 25oC, what percent of the acid is dissociated in a 0.5 M solution of HA at 25oC?
(A) 0.08% (B) 0.2% (C) 1% (D) 4%
| | |
28. Acid Acid Dissociation Constant, Ka
H3PO4 7 x 10-3
H2PO4– 8 x 10-8
HPO42– 5 x 10-13
On the basis of the information above, a buffer with a
pH = 9 can best be made by using
(A) H3PO4 + H2PO4– (B) H2PO4- + PO42–
(C) H2PO4- + HPO42– (D) HPO42- + PO43–
| | |
Questions 29-31 Assume all solutions are 1 M.
(A) NH3 and NH4CI
(B) H3PO4 and NaH2PO4
(C) HCI and NaCI
(D) NH3 and HC2H3O2 (acetic acid)
29. The solution with a pH = 0 and is not a buffer
| | |
30. The most nearly neutral solution
| | |
31. A buffer at a pH > 8
| | |
32. A buffer at a pH < 6
| | |
Questions 33-34 Ka for HCN is 5.0 x 10-10.
33. What is the H+ concentration in 0.05 M HCN?
(A) 2.5 x 10-11 M (B) 2.5 x 10-10 M
(C) 5.0 x 10-10 M (D) 5.0 x 10-6 M
| | |
34. What is the pH of a 0.02 M solution of HCN?
(A) Between 7 and 10 (B) 7
(C) Between 4 and 7 (D) 4
| | |
Questions 35-36 The graph shows the titration curve when 100. mL of 0.0250 M acetic acid is titrated with 0.100 M NaOH.
[pic]
35. Which of the following indicators is the best for this titration?
Indicator pH Range of Color Change
(A) Methyl orange 3.2 - 4.4
(B) Methyl red 4.8 - 6.0
(C) Bromthymol blue 6.1 - 7.6
(D) Phenolphthalein 8.2 - 10.0
| | |
36. What part of the curve corresponds to the optimum buffer action for the acetic acid/acetate ion pair?
(A) Point V (B) Point X (C) Point Z (D) Along WY
| | |
Practice Free Response
1. Complete the chart with the formula or name of the following acids.
|hydrobromic acid | |nitrous acid |
| |H2SO4 | |
| |carbonic acid | |
|H2S | |HClO4 |
2. Write a net ionic equation for the following aqueous reactions.
a. Sodium hydroxide and nitric acid
| |
b. Sulfur dioxide gas and barium hydroxide
| |
c. Solid potassium oxide and hydrochloric acid
| |
3. Excess nitric acid is added to solid calcium carbonate.
a. Balanced equation:
| |
b. Briefly explain why statues made of marble displayed outdoors in urban areas are deteriorating.
| |
4. Rank from strongest acid (1) to weakest (4).
|HBrO3 |HClO3 |HBrO2 |HIO2 |
| | | | |
5. Consider the weak acid, HCN.
a. 2.70 g of HCN is added to water to make 1.0 L.
(1) Determine the initial concentrations of HCN
| |
(2) The pH at equilibrium is 5.155. Determine [H+]E.
| |
(3) Determine Ka.
| [ ]| | |
|I | | | | |
|C | | | | |
|E | | | | |
(4) What is the percent ionization?
| |
(5) What is Kb for CN-?
| |
b. Complete the equation; label the acids (A) and bases (B), and link the conjugate pairs; and then write the equilibrium expression and determine K.
|Equation |Expression |
|HCN + H2O Δ | |
|HCN + OH- Δ | |
|CN- + H2O Δ | |
|CN- + H3O+ Δ | |
c. Identify the Lewis acid and Base: CN- + H+ → HCN(aq)
|base | |acid | |
d. Determine the pH of a 0.50 M NaCN.
|[ ] | | |
|I | | | | | |
|C | | | | | |
|E | | | | | |
e. A pH-9.00 buffer is made by adding NaCN to 0.500 L of 1.00 M HCN. How many grams of NaCN are added?
| |
| |
f. Calculate the pH when each of the following is added to the pH 9 buffer from part (e).
|0.10 mol H+ | |
|0.20 mol OH- | |
|0.30 mol H+ | |
6. 25 mL of 0.40 M HF (Ka = 7.2 x 10-4) reacts with 15 mL of 0.40 M NaOH according to the reaction.
HF(aq) + OH-(aq) → H2O(l) + F-(aq)
a. Calculate the moles of HF remaining.
| |
b. Calculate the initial concentration of HF.
| |
c. Calculate the initial concentration of F-.
| |
d. Calculate the equilibrium concentration of H+.
|[ ] | | |
|I | | | | |
|C | | | | |
|E | | | | |
e. Calculate the pH.
| |
7. Propanoic acid, HC3H5O2, ionizes in water:
HC3H5O2(aq) Δ C3H5O2-(aq) + H+(aq) Ka = 1.34 x 10-5
a. Write the equilibrium expression for the ionization.
| |
b. Calculate the pH of 0.265 M propanoic acid.
|[ ] | | |
|I | | | | |
|C | | | | |
|E | | | | |
c. 0.496 g of NaC3H5O2 is added to a 50.0 mL of 0.265 M propanoic acid. Calculate
(1) the initial concentration of C3H5O2-.
| |
(2) the equilibrium concentrations of H+.
|[ ] | | |
|I | | | | |
|C | | | | |
|E | | | | |
Methanoic acid forms when methanoate reacts with water:
CHO2-(aq) + H2O(l) Δ HCHO2(aq) + OH-(aq)
d. Given that [OH-]E is 4.18 x 10-6 M in a 0.309 M solution of sodium methanoate, calculate
(1) Kb for CHO2-.
| [ ] | | |
|I | | | | | |
|C | | | | | |
|E | | | | | |
(2) Ka for HCHO2.
| |
e. Which acid is stronger, propanoic acid or methanoic acid? Justify your answer.
| |
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