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Chapter 14

Chapter 14. Aqueous Equilibria: Acids and Bases. Acid–Base Concepts 01. Arrhenius Acid: A substance which dissociates to form hydrogen ions (H + ) in solution. HA( aq )  H + ( 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 to form hydrogen ions (H+) in solution. HA(aq)  H+(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. Strong vs. Weak acids 03

  5. HA(aq) H1+(aq) + A1-(aq) n = 1 H3O1+ n = 2 H5O21+ n = 3 H7O31+ n = 4 H9O41+ Hydrated Protons and Hydronium Ions Due to high reactivity of the hydrogen ion, it is actually hydrated by one or more water molecules. [H(H2O)n]1+ For our purposes, H1+ is equivalent to H3O1+.

  6. Acid–Base Concepts

  7. Lewis Acid–Base Concepts

  8. 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.

  9. Acid–Base Concepts 06 - +

  10. Lewis Acids and Bases Lewis Acid: An electron-pair acceptor. Lewis Base: An electron-pair donor.

  11. Lewis Acids and Bases Lewis Acid: An electron-pair acceptor. Lewis Base: An electron-pair donor.

  12. Acid–Base Concepts 07 • Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids. (a) H2SO4 (b) HSO4– (c) H3O+ • Identify the Lewis acid and Lewis base in each of the following reactions: • (a) SnCl4(s) + 2 Cl–(aq) æ SnCl62–(aq) • (b) Hg2+(aq) + 4 CN–(aq) æ Hg(CN)42–(aq) • (c) Co3+(aq) + 6 NH3(aq) æ Co(NH3)63+(aq)

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

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

  15. Dissociation of Water 03 • The concentration of OH– ions in a certain household ammonia cleaning solution is 0.0025 M. Calculate the concentration of H+ ions. • Calculate the concentration of OH– ions in a HCl solution whose hydrogen ion concentration is 1.3 M.

  16. 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 [H+], [H+] = 10-pH 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

  17. pH – A Measure of Acidity 02 • Nitric acid (HNO3) is used in the production of fertilizer, dyes, drugs, and explosives. Calculate the pH of a HNO3 solution having a hydrogen ion concentration of 0.76 M. • The pH of a certain orange juice is 3.33. Calculate the H+ ion concentration. • The OH– ion concentration of a blood sample is 2.5 x 10–7 M. What is the pH of the blood?

  18. pH – A Measure of Acidity 04 Color of Tea: Polyphenols, Thearubigins Color of Red Cabbage: Anthocyanin 

  19. pH – A Measure of Acidity 04

  20. HClO4 HI HBr HCl H2SO4 HNO3 H3O+ HSO4– Strength of Acids and Bases 03 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

  21. Strength of Acids and Bases 04 • Stronger acid + stronger base weaker acid + weaker base • Predict the direction of the following: • HNO2(aq) + CN–(aq) æ HCN(aq) + NO2–(aq) • HF(aq) + NH3(aq) æ F–(aq) + NH4+(aq)

  22. 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) • Or simply: HA(aq) æ H+(aq) + A–(aq)

  23. A- + H2O(l) HA(aq) + OH−(aq) [HA] [OH−] [A-] [HA] [OH−] [A-] Kb = Kb = Conjugate Base Ionization Const Ka  = Kw Ka Kb = Kw

  24. 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

  25. Strength of Acids and Bases 03 (a) Arrange the three acids in order of increasing value of Ka. (b) Which acid, if any, is a strong acid? (c) Which solution has the highest pH, and which has the lowest? K = (42/2) = 8 12/5= 0.2 Very Large

  26. Acid Ionization ConstantsDetermine the pH of 0.50M HA solution at 25°C. Ka = 7.1 x 10–4 05 • Initial Change Equilibrium Table:. - + 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):

  27. HA (aq) H+(aq) + A-(aq) = 7.1 x 10-4 = 7.1 x 10-4 = 7.1 x 10-4 [H+][A-] x2 x2 Ka Ka = Ka = 0.50 - x [HA] 0.50 HA (aq) H+(aq) + A-(aq) What is the pH of a 0.50M Citric acidsolution (at 250C)? Initial (M) 0.50 0.00 0.00 Change (M) -x +x +x Equilibrium (M) 0.50 - x x x 100•Ka < Co ? 100 x 7.1 x 10-4 = 0.071 < 0.5 0.50 – x 0.50 x2 = 3.55 x 10-4 x = 0.019 M pH = -log [H+] = 1.72 [H+] = [A-] = 0.019 M [HA] = 0.50 – x = 0.48 M

  28. Acid Ionization Constants 06 • pH of a Weak Acid (Cont’d): • Substitute equilibrium concentrations into equilibrium expression. • If 100•Ka < Cothen (C0 – x) approximates to (C0). • The equation can now be solved for x and pH. • If 100•Kais not significantly smaller than Co the quadratic equation must be used to solve for x and pH.

  29. Acid Ionization Constants 07 • 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.

  30. HA(aq) H1+(aq) + A1-(aq) Acid Ionization Constants 09 • 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.

  31. Percent dissociation of a weak acid decreases as its concentration increases • Concentration Dependence:

  32. Weak Bases: 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+].

  33. 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. (caffeine)

  34. Base Ionization Constants 03 • Product of Ka and Kb: multiplying out the expressions for Ka and Kb equals Kw. Ka Kb = Kw

  35. NH3(aq) + H2O(l) NH4+(aq) + OH−(aq) [NH4+] [OH−] [NH3] = 1.8  10−5 Kb = pH of Basic Solutions What is the pH of a 0.15 M solution of NH3?

  36. 1.8  10−5 = (x)2 (0.15 - x) pH of Basic Solutions (1.8  10−5) (0.15) = x2 2.7  10−6 = x2 1.6  10−3 = x2 100 x Kb < C0 ? 1.8  10−3< 0.15 0.15 –x = 0.15

  37. pH of Basic Solutions Therefore, X = [OH−] = 1.6  10−3M pOH = −log (1.6  10−3) pOH = 2.80 pH = 14.00 − 2.80 pH = 11.20

  38. Diprotic & Polyprotic Acids 01 H2SO4 H3PO4 • 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.

  39. 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

  40. 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 567 kJ/mol for HF 299 kJ/mol for HI

  41. 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.

  42. (pm)

  43. 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.

  44. 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, the Y–O bond will pull more electrons, the O–H bond will be more polar and the acid will be stronger.

  45. 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.

  46. Polar Covalent Bonds 02 Pauling Electronegativities Detailed List of Electronegativity; http://environmentalchemistry.com/yogi/periodic/electronegativity.html

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

  48. 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.

  49. Acid-Base Properties of Salts

  50. Strong bases • Strong bases: • The following metals make strong hydroxy base • Alkali metal cations of group 1A • Alkaline earth metal cations of group 2A • except for Be

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