1 Acids and Bases - the Three Definitions 1. The Arrhenius Definition of an Acid 2. Acid strength...
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Transcript of 1 Acids and Bases - the Three Definitions 1. The Arrhenius Definition of an Acid 2. Acid strength...
1
Acids and Bases - the Three Definitions
1. The Arrhenius Definition of an Acid
2. Acid strength and pKa
3. Ka, pKa, pKb
4. polyprotic acids, pKa1, pKa2, pKa3
5. Kb and pKb
6. Base strength and pKb
7. The pH scale and the
8. Autoionization of water, Kw
9. pH, pOH, and pKa
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Acids and Bases - Simple Definitions
Arrhenius Definition:
Acids: increases [H+] in aqueous solution
Bases: increases [OH-] in aqueous solution
Bronsted-Lowry Definition: (based on proton transfer reactions)
Acids: proton (H+) donor
Bases: proton (H+) acceptor
Lewis Definition:
Acids: electron pair acceptor
Bases: electron pair donor
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Ka and pKa
Acetic Acid is a weak Arrhenius acid, which liberates H+ in solutions
CH3COOH (aq) = CH3COO- (aq) + H+ (aq)
Ka =
[CH3COO-] [H+][CH3COOH]
= 1.76 x 10-5
The pKa is by definition the negative of log10 Ka:
pKa = - log10 (1.76 x 10-
5) = 4.75
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Kb and pKb
Arrhenius bases liberate OH- in solution.
Kb is the equilibrium constant for this reaction.NH4OH (aq) = NH4
+ (aq) + OH- (aq)
Kb =
[NH4+] [OH-]
[NH4OH]
= 1.76 x 10-5
pKb = - log10 Kb (definition)
pKb = - log10 (1.8 x 10-
5) = 4.74
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Ka and Acid Strength
The stronger the acid, the larger the Ka and the smaller the pKa:
CH3COOH (aq) = CH3COO- (aq) + H+ (aq) Ka = 1.76 x 10-5
HCN (aq) = CN- (aq) + H+ (aq) Ka = 6.17 x 10-10
pKa = 4.75
HNO2 (aq) = NO2- (aq) + H+ (aq)
Ka = 4.6 x 10-4 pKa = 3.34
pKa = 9.21
stronger
weaker
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Polyprotic Acids
pKa1, pKa2, pKa3 describe the dissociation of the first, second, and third ionizable protons.H2CO3 (aq) = HCO3
- (aq) + H+ (aq) Ka1 = 4.3 x 10-7HCO3
- (aq) = CO32- (aq) + H+ (aq)
Ka2 = 5.6 x 10-11
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Kb and Base Strength
The stronger the base, the larger the Kb and the smaller the pKb:
NH4OH (aq) = NH4+ (aq) + OH- (aq)
Kb = 1.8 x 10-5 pKb = 4.74
stronger
weaker
PO43- (aq) + H2O (l) = HPO4
2- (aq) + OH- (aq) Kb = 4.5 x 10-2 pKb =
1.34
Conclusion: phosphate anion is a stronger base than NH4OH.
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Acidity/Basicity of a Solution and the pH ScaleThe degree of acidity or basicity of a solution is measured on the pH scale:
pH = -log10 [H+]
[H+] pH
1 M 0
10-2 M 2
10-4 M 4
10-6 M 6
10-8 M 8
10-10 M 10
10-12 M 12
10-14 M 14
low pH (≈0-1) is strongly acid
high pH (≈13-14) is strongly basic
pH 7 is a neutral solution
pOH = -log10 [OH-]
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Self-Ionization of WaterThe concentrations of H+ and OH- are related by the self- ionization of water -
H2O = H+ + OH- Kw = [H+] [OH-] = 10-14 (at 25°C)
What is the [H+] in pure water? If x = the molarity of [H+] , H2O = H+ + OH- Kw = [H+] [OH-] = 10-14
x x x2 = 10-14
Therefore, x = [H+] = [OH-] = 10-7 M
in pure water at 25°CPure water is pH 7.0. pH 7 a neutral solution,
[H+] = [OH-]
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pH and pOH: Measures of Acidic and Basicity
Because of the self-ionization of water, [H+] and [OH-] are not
independent quantities but are related byKw = [H+] [OH-] = 10-14
pOH is a logarithmic measure of the [OH-] concentration
pH + pOH = 14
pOH = -log10 [OH-]
From the expression for Kw:
-Log10 Kw = - log10 [H+] - log10[OH-] = +14
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Acidity, Basicity, pH, and pOH
pH or pOH can be used to measure acidity
[H+] pH pOH
1 M 0 14
10-2 M 2 12
10-4 M 4 10
10-6 M 6 8
10-8 M 8 6
10-10 M 10 4
10-12 M 12 2
10-14 M 14 0
low pH or high pOH is strongly acid
pH 7 is a neutral solution
high pH or low pOH is strongly acid
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Measurement of pH: the pH Meter
pH varies linearly with output voltage and can be
measured over the range pH 0 to
pH 14
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Acids and Bases – Simple Definitions
Arrhenius Definition:
Acids: increases [H+] in aqueous solution
Bases: increases [OH-] in aqueous solution
Bronsted-Lowry Definition: (based on proton transfer reactions)
Acids: proton (H+) donor
Bases: proton (H+) acceptor
Lewis Definition:
Acids: electron pair acceptor
Bases: electron pair donor
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Bronsted-Lowry Definition
Many proton transfer reactions occur in aqueous solution. These are also acid-base neutralizations according to the Bronsted-Lowry definition (but not according to the Arrhenius definition).
For example, weak acids can neutralize weak bases by a proton transfer reaction. In such reactions there are always two acids and two bases.
HNO2(aq) + NH3(aq) = NO2-(aq) +
NH4+(aq)
acid base base acid
The acids are the proton donors, the bases are proton acceptors.
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Bronsted-Lowry Acid/Base Pairs
Each species participating in a proton transfer reaction can exist in a protonated form and a de-protonated form. The protonated form is the Bronsted acid, and the de-protonated form is the Bronsted base. Thus one speaks of “conjugate acid/base pairs”.
In any Bronsted acid/base neutralization, there are
1. Two Bronsted conjugate acids
2. Two Bronsted conjugate bases
3. Two Bronsted conjugate acid/base pairs
HNO2(aq) + NH3(aq) = NO2-(aq) + NH4
+(aq)Bronsted acid Bronsted base conjugate acid/base pair
Bronsted base Bronsted acid conjugate acid/base pair
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Bronsted-Lowry Acid-Base Neutralizations
H2PO4-(aq) + HCO3
-(aq) = HPO42-(aq) +
H2CO3(aq)
HCl(g) + NH3(g) = NH4Cl (s)
2 NH3(g) + CO2(g) NH2CONH2(aq) + H2O(l)
Which of the following reactions are acid/base neutralizations in the Bronsted-Lowry picture?
Pick out the conjugate acid/base pairs.
H3O+(aq) + OH-(aq) = 2 H2O (l)
Note: any Arrhenius acid (or base) is also a Bronsted acid (or base)
Acid BaseBase Acid
Acid AcidBase Base
not Bronsted-Lowry
Bronsted-Lowry
Bronsted-Lowry
Acid Base
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In the Bronsted-Lowry Definition, many species can function either as Acids OR Bases
H2PO4-(aq) + HCO3
-(aq) = HPO42-(aq) +
H2CO3(aq)
H2PO4-(aq) + HCO3
-(aq) = H3PO4(aq) + CO3
2-(aq)
In this reaction, H2PO4- functions as a
Bronsted acid (why?)
In this reaction, H2PO4- functions as a
Bronsted base (why?)
Whether a chemical species is a Bronsted-Lowry acid or base can depend on the reaction it is in.
Is H2PO4- an acid or a base?
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Relationship of Ka and Kb for a Bronsted Acid/Base Pair
A- + H2O = HA + OH-
[HA] [OH-]
[A-] Kb =
HA + H2O = A- + H3O+
[A-] [H3O+]
Ka =
conj acid
conj base
H3O+H2O
conj acid
conj base
OH-H2O
A Ka can be defined for any conjugate acid, and a Kb for its conjugate base.
Note that Ka . Kb
= Kw
[HA]
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