acid-base equilibria chapter 16. conjugate acid-base pairs eqn 16.7 p 670
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![Page 1: Acid-Base Equilibria Chapter 16. Conjugate Acid-Base Pairs Eqn 16.7 p 670](https://reader035.vdocuments.net/reader035/viewer/2022062407/56649d455503460f94a2274f/html5/thumbnails/1.jpg)
Acid-BaseAcid-BaseEquilibriaEquilibria
Chapter 16Chapter 16
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Conjugate Acid-Base Pairs
Eqn 16.7 p 670
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Conjugate Acid-Base Pairs
Eqn 16.8 p 671
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Fig 16.4 Relative strengths of some conjugate acid-base pairs
The strongerthe acid…
the weaker its conjugate base.
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H2O (l) H+ (aq) + OH- (aq)
The Ion Product of Water
Kc =[H+][OH-]
[H2O][H2O] = 55.6 M = constant
Kc[H2O] = Kw = [H+][OH-]
Ion-product constant (Kw) - the product of the molar concentrations of H+ and OH- ions at a particular temperature.
At 25°C:Kw = [H+][OH-] = 1.0 x 10-14
[H+] = [OH-]
[H+] > [OH-]
[H+] < [OH-]
Solution is:
neutral
acidic
basic
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pH – A Measure of Acidity
pH = −log [H+]
[H+] = [OH-]
[H+] > [OH-]
[H+] < [OH-]
Solution is:
neutral
acidic
basic
[H+] = 1 x 10-7
[H+] > 1 x 10-7
[H+] < 1 x 10-7
pH = 7
pH < 7
pH > 7
At 25°C
pH [H+]
pH = −log [H3O+]
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pOH = -log [OH-]
[H+][OH-] = Kw = 1.0 x 10-14
−log [H+] – log [OH-] = 14.00
pH + pOH = 14.00
Fig 16.5 H+ concentrations and pH of common substances
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The pH of rainwater collected in a certain region of the northeastern United States on a particular day was 4.82. What is the H+ ion concentration of the rainwater?
pH = -log [H+]
[H+] = 10-pH = 10-4.82 = 1.5 x 10-5 M
The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?
pH + pOH = 14.00
pOH = -log [OH-] = -log (2.5 x 10-7) = 6.60
pH = 14.00 – pOH = 14.00 – 6.60 = 7.40
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AcidsandBases
Other “p” Functions
The “p” in pH tells us to take the negative base - 10 logarithm of the quantity (in this case, hydronium ions)
Some similar examples:
pOH = -log [OH−]
pKw = -log Kw
pCl = -log [Cl−]
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AcidsandBases
How Do We Measure pH?
• For less accurate measurements:– Litmus paper
• “Red” paper turns blue above ~pH = 8• “Blue” paper turns red below ~pH = 5
– Indicator:
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AcidsandBases
For more accurate measurements: ◦ pH meter, which measures the voltage in the solution
Fig 16.6 Digital pH meter
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Table 4.2
Strong Electrolyte – 100% dissociation
Strong Acids and Strong Bases are strong electrolytes (p 130)
HCl (aq) + H2O (l) H3O+ (aq) + Cl− (aq)
Strong Acids and Bases
NaOH (aq) Na+ (aq) + OH− (aq)
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HF (aq) + H2O (l) H3O+ (aq) + F- (aq)
Weak Acids are weak electrolytes
HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq)
Weak Acids
Weak electrolytes - only partially ionized in aqueous solution
• Most acidic substances are weak acids
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HA (aq) + H2O (l) H3O+ (aq) + A- (aq)
Weak Acids (HA) and Acid Ionization Constants
HA (aq) H+ (aq) + A- (aq)
Ka =[H+][A-][HA]
Ka ≡ acid ionization constant
Kaweak acidstrength
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16.5
Table 16.2 Some Weak Acids in Water at 25 °C
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Calculating Ka from the pHSample Exercise 16.10 p 682
The pH of a 0.10 M solution of formic acid, HCOOH, at 25Cis 2.38. Calculate Ka for formic acid at this temperature.
HCOOH (aq) H+ (aq) + HCOO- (aq)
pH = -log [H+]
2.38 = -log [H+]
-2.38 = log [H+]
10-2.38 = 10log [H+] = [H+]
4.2 10-3 = [H+] = [HCOO-]
Ka =[H+][HCOO-][HCOOH]
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[HCOOH], M [H3O+], M [HCOO-], M
Initially 0.10 0 0
Change - 4.2 10-3 + 4.2 10-3 + 4.2 10-3
At Equilibrium 0.10 - 4.2 10-3
= 0.0958 = 0.10
4.2 10-3 4.2 10-3
4.2 10-3 = [H+] = [HCOO-]
Calculating Ka from the pH
HCOOH (aq) H+ (aq) + HCOO- (aq)
Ka =[H+][HCOO-][HCOOH]
=[4.2 10-3] [4.2 10-
3][0.10]
= 1.8 10-4
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Percent ionization =
For a monoprotic acid HA:
Percent ionization = [H+]
[HA]0
x 100% [HA]0 = initial concentration
Concentration ionized
Original concentrationx 100%
Fig 16.9The more dilute the acid,the greater the percent
ionization:
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What is the pH of a 0.50 M HF solution (at 25°C)?
HF (aq) H+ (aq) + F- (aq) Ka =[H+][F-][HF]
= 6.8 x 10-4
HF (aq) H+ (aq) + F- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.50 0.00
-x +x
0.50 - x
0.00
+x
x x
Ka =x2
0.50 - x= 6.8 x 10-4
Ka x2
0.50= 6.8 x 10-4
0.50 – x 0.50If [HF] > 100 Ka
x2 = 3.40 x 10-4 x = 0.018 M
[H+] = [F-] = 0.018 M pH = -log [H+] = 1.72
[HF] = 0.50 – x = 0.48 M
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When can I use the approximation?
Then 0.50 – x 0.50
x = 0.0180.018 M0.50 M
x 100% = 3.6%Less than 5%
Approximation ok.
What is the pH of a 0.05 M HF solution (at 25°C)?
Ka x2
0.05= 6.8 x 10-4 x = 0.006 M
0.006 M0.05 M
x 100% = 12%More than 5%
Approximation not ok.
Must solve for x exactly using quadratic equation...
Let’s determine error introduced:
If [HF] > 100 Ka
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Solving weak acid ionization problems:
1. Identify the major species that can affect the pH.
• In most cases, the autoionization of water can be ignored.
• Ignore [OH¯] because it is determined by [H+].
2. Use ICE table to express the equilibrium concentrations in terms of single unknown x.
3. Write Ka in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly.
4. Calculate concentrations of all species and/or pH of the solution.
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What is the pH of a 0.122 M monoprotic acid whose Ka is 5.7 x 10-3?
HA (aq) H+ (aq) + A- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.122 0.00
-x +x
0.122 - x
0.00
+x
x x
Ka =x2
0.122 - x= 5.7 x 10-3
NO!!
Approximation not ok.
Is [HF] > 100 Ka ?
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Ka =x2
0.122 - x= 5.7 x 10-3 x2 + 5.7 X 10-3 x – 6.95 X 10-4 = 0
ax2 + bx + c =0-b ± b2 – 4ac
2ax =
x = 0.0237 x = - 0.0294
HA (aq) H+ (aq) + A- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.122 0.00
-x +x
0.122 - x
0.00
+x
x x
[H+] = x = 0.0237 M pH = -log[H+] = 1.625
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Polyprotic AcidsHave more than one ionizable proton
If difference between the Ka1 and subsequent Ka values > 103, the pH generally depends only on the first dissociation.