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Chapter 14

Chapter 14. Aqueous Equilibria: Acids and Bases. Acid–Base Concepts 01. Arrhenius Acid: A substance which dissociates in water to form hydrogen ions (H + ) in solution. HA( aq ) + H 2 O( l )  H 3 O + ( aq ) + A – ( aq )

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Chapter 14

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  1. Chapter 14 Aqueous Equilibria: Acids and Bases

  2. Acid–Base Concepts 01 • Arrhenius Acid:A substance which dissociates in water to form hydrogen ions (H+) in solution. HA(aq) + H2O(l) H3O+(aq) + A–(aq) • Arrhenius Base:A substance that dissociates in, or reacts with water to form hydroxide ions (OH–). • MOH(aq)  M+(aq) + OH–(aq)

  3. Acid–Base Concepts 02 • Brønsted–Lowry Acid:Substance that can donate H+ • Brønsted–Lowry Base:Substance that can accept H+ • Chemical species whose formulas differ only by one proton are said to be conjugate acid–base pairs.

  4. Acid–Base Concepts 03

  5. Acid–Base Concepts 04

  6. Acid–Base Concepts 05 • A Lewis Acid is an electron-pair acceptor. These are generally cations and neutral molecules with vacant valence orbitals, such as Al3+, Cu2+, H+, BF3. • A Lewis Base is an electron-pair donor. These are generally anions and neutral molecules with available pairs of electrons, such as H2O, NH3, O2–. • The bond formed is called a coordinate bond.

  7. Acid–Base Concepts 06

  8. Acid–Base Concepts 07 • Problems 14.1,14.2 • Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids. (a) H2SO4 (b) HSO4– (c) H3O+ • Problems 14.27 • Identify the Lewis acid and Lewis base in each of the following reactions: (a) AlCl3(s) + Cl–(aq) æ AlCl4–(aq) (b) SO2 (aq) + OH–(aq) æ HSO3–(aq) (c) Ag+(aq) + 2 NH3(aq) æ Ag(NH3)2+(aq)

  9. Dissociation of Water 01 • Water can act as an acid or as a base. H2O(l) æ H+(aq) + OH–(aq) • This is called the autoionization of water. H2O(l) + H2O(l)æ H3O+(aq) + OH–(aq)

  10. Dissociation of Water 02 • This equilibrium gives us the ion product constant for water. Kw = Kc = [H3O+][OH–] = 1.0 x 10–14 If we know either [H3O+] or [OH–] then we can determine the other quantity.

  11. Dissociation of Water 03 • Problem 14.6 The concentration of OH– in a sample of seawater is 5.0 × 10-6 M. Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, or basic. • Problem 14.7 At 50°C the value of Kw is 5.5 × 10-14. What are the concentrations of H3O+ and OH– in a neutral solution at 50°C?

  12. pH – A Measure of Acidity 01 • The pH of a solution is the negative logarithm of the hydrogen ion concentration (in mol/L). pH = –log [H3O+] pOH = –log [OH– ] pH + pOH = 14 Acidic solutions: [H+] > 1.0 x 10–7 M, pH < 7.00Basic solutions: [H+] < 1.0 x 10–7 M, pH > 7.00Neutral solutions: [H+] = 1.0 x 10–7 M, pH = 7.00

  13. pH – A Measure of Acidity 02 • Problem 14.8 Calculate the pH of each of the following Problem 14.9 Calculate the concentrations of H3O+ and OH– in each of the following solutions: (a) Human blood (pH 7.40) (b) A cola beverage (pH 2.8) Problem 14.10 Calculate the pH of • 0.050 M HClO4 (b) 6.0 M HCl (c) 4.0 M KOH (d) 0.010 M Ba(OH)2 Problem 14.11 Calculate the pH of a solution prepared by dissolving 0.25 g of BaO in enough water to make 0.500 L of solution

  14. pH – A Measure of Acidity 03

  15. Strength of Acids and Bases 01 • Strong acids and bases: are strong electrolytes that are assumed to ionize completely in water. • Weak acids and bases: are weak electrolytes that ionize only to a limited extent in water. • Solutions of weak acids and bases contain ionized and non-ionized species.

  16. HClO4 HI HBr HCl H2SO4 HNO3 H3O+ HSO4– Strength of Acids and Bases 02 ACIDCONJ. BASE ACIDCONJ. BASE ClO4– I– Br – Cl – HSO4 – NO3 – H2O SO42– HSO4– HF HNO2 HCOOH NH4+ HCN H2O NH3 SO42– F – NO2 – HCOO – NH3 CN – OH – NH2 – IncreasingAcid Strength Increasing Acid Strength

  17. Strength of Acids and Bases 03 • Stronger acid + stronger base weaker acid + weaker base • Problems 14.4 Predict the direction of the following: HF(aq) + NO3– (aq) æ F–(aq) + HNO3(aq) NH4+(aq) + CO3–2 (aq) æ HCO3– (aq) + NH3(aq)

  18. Acid Ionization Constants 01 • Acid Ionization Constant: the equilibrium constant for the ionization of an acid.HA(aq) + H2O(l) æ H3O+(aq) + A–(aq)

  19. Acid Ionization Constants 02 ACIDKaCONJ. BASE Kb HF HNO2 C9H8O4 (aspirin) HCO2H (formic) C6H8O6 (ascorbic) C6H5CO2H (benzoic) CH3CO2H (acetic) HCN C6H5OH (phenol) 7.1 x 10 –4 4.5 x 10 –4 3.0 x 10 –4 1.7 x 10 –4 8.0 x 10 –5 6.5 x 10 –5 1.8 x 10 –5 4.9 x 10 –10 1.3 x 10 –10 F– NO2 – C9H7O4 – HCO2 – C6H7O6 – C6H5CO2 – CH3CO2 – CN – C6H5O – 1.4 x 10 –11 2.2 x 10 –11 3.3 x 10 –11 5.9 x 10 –11 1.3 x 10 –10 1.5 x 10 –10 5.6 x 10 –10 2.0 x 10 –5 7.7 x 10 –5

  20. Calculating Equilibrium Concentration in Solutions of Weak Acids

  21. Acid Ionization Constants 04 • Initial Change Equilibrium Table: Determine the pH of 0.50M HA solution at 25°C. Ka = 7.1 x 10–4. - + H + A æ HA (aq) (aq) (aq) Initial ( M ) : 0.50 0.00 0.00 Change (M): – x + x + x Equilib 0.50 – x x x (M):

  22. Acid Ionization Constants 05 • pH of a Weak Acid (Cont’d): • Substitute new values into equilibrium expression. • If Ka is significantly (>1000 x) smaller than [HA] the expression (0.50 – x) approximates to (0.50). • The equation can now be solved for x and pH. • If Ka is not significantly smaller than [HA] the quadratic equation must be used to solve for x and pH.

  23. Acid Ionization Constants 06 • The Quadratic Equation: • The expression must first be rearranged to: • The values are substituted into the quadratic and solved for a positive solution to x and pH.

  24. Acid Ionization Constants 08 • Percent Dissociation: A measure of the strength of an acid. • Stronger acids have higher percent dissociation. • Percent dissociation of a weak acid decreases as its concentration increases.

  25. Base Ionization Constants 01 • Base Ionization Constant: The equilibrium constant for the ionization of a base. • The ionization of weak bases is treated in the same way as the ionization of weak acids.B(aq) + H2O(l) æ BH+(aq) + OH–(aq) • Calculations follow the same procedure as used for a weak acid but [OH–] is calculated, not [H+].

  26. Base Ionization Constants 02 BASEKbCONJ. ACID Ka C2H5NH2 (ethylamine) CH3NH2 (methylamine) C8H10N4O2 (caffeine) NH3 (ammonia) C5H5N(pyridine) C6H5NH2 (aniline) NH2CONH2 (urea) 5.6 x 10 –4 4.4 x 10 –4 4.1 x 10 –4 1.8 x 10 –5 1.7 x 10 –9 3.8 x 10 –10 1.5 x 10 –14 C2H5NH3+ CH3NH3+ C8H11N4O2+ NH4+ C5H6N+ C6H5NH3+ NH2CONH3+ 1.8 x 10 –11 2.3 x 10 –11 2.4 x 10 –11 5.6 x 10 –10 5.9 x 10 –6 2.6 x 10 –5 0.67 Note that the positive charge sits on the nitrogen.

  27. Diprotic & Polyprotic Acids 01 • Diprotic and polyprotic acids yield more than one hydrogen ion per molecule. • One proton is lost at a time. Conjugate base of first step is acid of second step. • Ionization constants decrease as protons are removed.

  28. Diprotic & Polyprotic Acids 02 ACIDKaCONJ. BASE Kb H2SO4 HSO4– C2H2O4 C2HO4– H2SO3 HSO3– H2CO3 HCO3– H2S HS– H3PO4 H2PO4– HPO42– Very Large 1.3 x 10 –2 6.5 x 10 –2 6.1 x 10 –5 1.3 x 10 –2 6.3 x 10 –8 4.2 x 10 –7 4.8 x 10 –11 9.5 x 10 –8 1 x 10 –19 7.5 x 10 –3 6.2 x 10 –8 4.8 x 10 –13 HSO4 – SO4 2– C2HO4– C2O42– HSO3 – SO3 2– HCO3– CO3 2– HS– S 2– H2PO4– HPO42– PO43– Very Small 7.7 x 10 –13 1.5 x 10 –13 1.6 x 10 –10 7.7 x 10 –13 1.6 x 10 –7 2.4 x 10 –8 2.1 x 10 –4 1.1 x 10 –7 1 x 10 –5 1.3 x 10 –12 1.6 x 10 –7 2.1 x 10 –2

  29. Molecular Structure and Acid Strength 01 • The strength of an acid depends on its tendency to ionize. • For general acids of the type H–X: • The stronger the bond, the weaker the acid. • The more polar the bond, the stronger the acid. • For the hydrohalic acids, bond strength plays the key role giving: HF < HCl < HBr < HI

  30. Molecular Structure and Acid Strength 02 • The electrostatic potential maps show all the hydrohalic acids are polar. The variation in polarity is less significant than the bond strength which decreases from 567 kJ/mol for HF to 299 kJ/mol for HI.

  31. Molecular Structure and Acid Strength 03 • For binary acids in the same group, H–A bond strength decreases with increasing size of A, so acidity increases. • For binary acids in the same row, H–A polarity increases with increasing electronegativity of A, so acidity increases.

  32. Molecular Structure and Acid Strength 04 • For oxoacids bond polarity is more important. If we consider the main element (Y):Y–O–H • If Y is an electronegative element, or in a high oxidation state, the Y–O bond will be more covalent and the O–H bond more polar and the acid stronger.

  33. Molecular Structure and Acid Strength 05 • For oxoacids with different central atoms that are from the same group of the periodic table and that have the same oxidation number, acid strength increases with increasing electronegativity.

  34. Molecular Structure and Acid Strength 06 • For oxoacids having the same central atom but different numbers of attached groups, acid strength increases with increasing central atom oxidation number. • As shown on the next slide, the number of oxygen atoms increases the positive charge on the chlorine which weakens the O–H bond and increases its polarity.

  35. Oxoacids of Chlorine: Molecular Structure and Acid Strength 07

  36. Molecular Structure and Acid Strength 08 • Predict the relative strengths of the following groups of oxoacids: a) HClO, HBrO, and HIO. b) HNO3 and HNO2. c) H3PO3 and H3PO4.

  37. Acid–Base Properties of Salts 01 • Salts that produce neutral solutions are those formed from strong acids and strong bases. • Salts that produce basic solutions are those formed from weak acids and strong bases. • Salts that produce acidic solutions are those formed from strong acids and weak bases.

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