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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