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Acids and Bases - the Three Definitions

Acids and Bases - the Three Definitions. 1. The Arrhenius Definition of an Acid 2. Acid strength and pK a 3. K a , pK a , pK b 4. polyprotic acids, pK a1 , pK a2 , pK a3 5. K b and pK b 6. Base strength and pK b 7. The pH scale and the

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Acids and Bases - the Three Definitions

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  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

  2. 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

  3. [CH3COO-] [H+] Ka = [CH3COOH] Ka and pKa Acetic Acid is a weak Arrhenius acid, which liberates H+ in solutions CH3COOH (aq) = CH3COO- (aq) + H+ (aq) = 1.76 x 10-5 The pKa is by definition the negative of log10 Ka: pKa = - log10 (1.76 x 10-5) = 4.75

  4. Ionization Constants (1)

  5. Ionization Constants (2)

  6. [NH4+] [OH-] Kb = [NH4OH] Kb and pKb Arrhenius bases liberate OH- in solution. Kb is the equilibrium constant for this reaction. NH4OH (aq) = NH4+ (aq) + OH- (aq) = 1.76 x 10-5 pKb = - log10 Kb (definition) pKb = - log10 (1.8 x 10-5) = 4.74

  7. HNO2 (aq) = NO2- (aq) + H+ (aq) Ka = 4.6 x 10-4 pKa = 3.34 Ka and Acid Strength The stronger the acid, the larger the Ka and the smaller the pKa: stronger CH3COOH (aq) = CH3COO- (aq) + H+ (aq) Ka = 1.76 x 10-5 pKa = 4.75 HCN (aq) = CN- (aq) + H+ (aq) Ka = 6.17 x 10-10 pKa = 9.21 weaker

  8. 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 10-7 HCO3- (aq) = CO32- (aq) + H+ (aq) Ka2 = 5.6 x 10-11

  9. NH4OH (aq) = NH4+ (aq) + OH- (aq) Kb = 1.8 x 10-5 PO43- (aq) + H2O (l) = HPO42- (aq) + OH- (aq) Kb = 4.5 x 10-2 pKb = 4.74 pKb = 1.34 Kb and Base Strength The stronger the base, the larger the Kb and the smaller the pKb: stronger Conclusion: phosphate anion is a stronger base than NH4OH. weaker

  10. Acidity/Basicity of a Solution and the pH Scale The degree of acidity or basicity of a solution is measured on the pH scale: pH = -log10 [H+] pOH = -log10 [OH-] [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 pH 7 is a neutral solution high pH (≈13-14) is strongly basic

  11. Self-Ionization of Water The 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 Min pure water at 25°C Pure water is pH 7.0. pH 7 a neutral solution, [H+] = [OH-]

  12. 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 by Kw = [H+] [OH-] = 10-14 pOH is a logarithmic measure of the [OH-] concentration pOH = -log10 [OH-] From the expression for Kw: -Log10 Kw = - log10 [H+] - log10[OH-] = +14 pH + pOH = 14

  13. [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 Acidity, Basicity, pH, and pOH pH or pOH can be used to measure acidity

  14. Measurement of pH: the pH Meter pH varies linearly with output voltage and can be measured over the range pH 0 to pH 14

  15. 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

  16. 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.

  17. 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 Two Bronsted conjugate acids Two Bronsted conjugate bases 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

  18. Note: any Arrhenius acid (or base) is also a Bronsted acid (or base) Bronsted-Lowry Acid-Base Neutralizations Which of the following reactions are acid/base neutralizations in the Bronsted-Lowry picture? Pick out the conjugate acid/base pairs. Bronsted-Lowry H2PO4-(aq) + HCO3-(aq) = HPO42-(aq) + H2CO3(aq) Acid Base Base Acid 2 NH3(g) + CO2(g) NH2CONH2(aq) + H2O(l) not Bronsted-Lowry HCl(g) + NH3(g) = NH4Cl (s) Base Acid H3O+(aq) + OH-(aq) = 2 H2O (l) Acid Acid Base Bronsted-Lowry Base

  19. In the Bronsted-Lowry Definition, many species can function either as Acids OR Bases 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? H2PO4-(aq) + HCO3-(aq) = HPO42-(aq) + H2CO3(aq) In this reaction, H2PO4- functions as a Bronsted acid (why?) H2PO4-(aq) + HCO3-(aq) = H3PO4(aq) + CO32-(aq) In this reaction, H2PO4- functions as a Bronsted base (why?)

  20. conj acid conj base H3O+ H2O [HA] [OH-] [A-] [A-] [H3O+] A- + H2O = HA + OH- HA + H2O = A- + H3O+ Ka = Kb = conj acid conj base OH- H2O Relationship of Ka and Kb for a Bronsted Acid/Base Pair A Ka can be defined for any conjugate acid, and a Kb for its conjugate base. [HA] Note that Ka . Kb = Kw Back to 5

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