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Acid-Base Chemistry
ν There are a couple of ways to define acidsand bases
ν Brønsted-Lowry acids and basesν Acid: H+ ion donor
ν Base: H+ ion acceptor
ν Lewis acids and basesν Acid: electron pair acceptor
ν Base: electron pair donor
Brønsted-Lowry Acids & Bases
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Brønsted-Lowry Acids & Bases
Brønsted-Lowry Acids & Bases
ν In most acid-base systems, water may play arole as either an acid (H+ donor) or a base (H+
acceptor)
Water as an acidH2O(λ) + B(aq) ↔ OH-(aq) + HB+(aq)
Water as a baseH2O(λ) + HA(aq) ↔ H3O+(aq) + A-(aq)
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Conjugate Acids & Bases
ν Acids react with bases and vice versa
ν All acids and bases come with a conjugatepair—a base or acid, respectively, that isformed in conjunction with the originalspecies
ExamplesHCl(aq) + H2O(λ) ↔ H3O+(aq) + Cl-(aq) acid base conjugate conjugate
acid base
Conjugate Acids & Bases
Examples
NaOH(aq) + H2O(λ) ↔ OH-(aq) + H2O(λ) + Na+(aq)base acid conjugate conjugate
base acid
CH3COOH(aq) + H2O(λ) ↔ CH3COO-(aq) + H3O+(aq)
acid base conjugate conjugate base acid
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Strengths of Acids and Bases
ν Strong acids donate H+ ions more easilyν The stronger the acid, the weaker the conjugate
base associated with that acid
ν Strong bases accept H+ ions more easilyν The stronger the base, the weaker the conjugate
acid associated with that base
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Autoionization of Water
ν Water always undergoes some degree ofdissociation to form H3O+ ions and OH- ions
2 H2O(λ) ↔ H3O+(aq) + OH-(aq)
ν The equilibrium constant for this process at25 oC is:
Kw = [H3O+][OH-] = 1.0 x 10-14
ν In pure water
[H3O+] = [OH-] = 1.0 x 10-7 M
Autoionization of Water
ν Kw is temperature dependent—it increaseswith increasing temperature
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Autoionization of Water
Example
Determine [H3O+] in a 0.053 M NaOH solution
Step 1: since NaOH is a strong base,dissociation is complete∴ [OH-] = 0.053 M
Step 2: Use Kw to calculate [H3O+]
M 10 x 1.9 053.0
10 x 1.0
][OH
K ]O[H
10 x 1.0 ]][OHO[H K
13-14-
-w
3
14--3w
===
==
+
+
The pH Scale
ν pH is a measure of the hydronium ion contentof a solution
ν pH is defined as:
pH = -log[H3O+]
log is log base 10, not ln (natural log)
[H3O+] is given in molar units (M)
ν pH of pure water ([H3O+] = 1.0 x 10-7 M):
pH = -log(1.0x10-7) = 7.0
ν pH of last example ([H3O+] = 1.9 x 10-13 M):
pH = -log(1.9x10-13) = 12.7
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The pH Scale
ν Neutral is defined as the pH of pure water:
pH = 7
ν Acidic solutions have pH lower than 7:
pH < 7 ⇒ acidic
ν Basic solutions have pH larger than 7:pH > 7 ⇒ basic
Figure 16-2
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The pH Scale
ν We can also use pOH to describe a solution
ν pOH is defined as:
pOH = -log[OH-]
ν The sum of pH and pOH must equal 14
pH + pOH = 14
assuming room temperature (25 oC)
The pH Scale
Example
Find [H3O+] of a solution that has pH = 9.37
ν Remember that pH = -log[H3O+]
So [H3O+] = 10-pH
Solution:
[H3O+] = 10-pH = 10-9.37 = 4.27 x 10-10 M
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The pH Scale
Example
Find pH of a solution that is a nitric acidconcentration of [HNO3] = 0.0018 M
Solution:
Since HNO3 is a strong acid, it will dissociatecompletely, so the hydronium ionconcentration will also be 0.0018 M
pH = -log[H3O+] = -log[0.0018 M] = 2.74
Salts of Acids and Bases
ν When an acid and a base undergo anexchange reaction, the result is a salt andwater:
HX(aq) + MOH(aq) ↔ MX(aq) + H2O acid base salt
ν If a strong base is neutralized by a strongacid, the resulting solution contains only thesaltHCl(aq) + NaOH(aq) ↔ NaCl(aq) + H2O(λ)
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Acid Rain
ν Carbon dioxide in the air is in equilibrium withH2O in atmospheric water droplets (clouds &fog):
CO2(aq) + H2O ↔ H2CO3(aq)
carbonic acid Ka = 4.2 x 10-7
H2CO3(aq) + H2O ↔ H3O+(aq) + HCO3-(aq)
ν Natural rain water has pH = 5.6
Acid Rain
ν Emitted pollutants can form additional acidsources in clouds:
NO2:
2 NO2(aq) + H2O ↔ HNO3(aq) + HNO2(aq)
nitric acid nitrous acid
strong
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Acid Rain
ν Emitted pollutants can form additional acidsources in clouds:
SO2:
SO2(aq) + H2O ↔ H2SO3(aq)
sulfurous acid2 SO2(g) + O2(g) → 2 SO3(g)
SO3(aq) + H2O ↔ H2SO4(aq)
sulfuric acid
strong
Acid Rain