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Equilibrium – Acids and Bases

Equilibrium – Acids and Bases. Review of Acids and Bases. Arrhenius Theory of Acids and Bases An acid is a substance that dissociates in water to produce one or more hydrogen ions (H + ) A base is a substance that dissociates in water to form one or more hydroxide ions. (OH - ) Examples:

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Equilibrium – Acids and Bases

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  1. Equilibrium – Acids and Bases

  2. Review of Acids and Bases • Arrhenius Theory of Acids and Bases • An acid is a substance that dissociates in water to produce one or more hydrogen ions (H+) • A base is a substance that dissociates in water to form one or more hydroxide ions. (OH-) • Examples: • Acid:HCl(aq)H+(aq) + Cl-(aq) • Base: LiOHLi+(aq) + OH-(aq)

  3. Limitations: • Classified based on chemical formula • Some substances do not have OH- in their chemical formulas but still yield OH- when they react with water. E.g. NH3 (ammonia) • Solution?

  4. Bronsted-Lowry Theory of Acids and Bases • An acid is a proton (H+) donor and must have H in its formula. • A base is a proton acceptor and must have a lone pair of electrons to form a bond with H+

  5. Two molecules or ions that are related by the transfer of a proton are called a conjugate acid-base pair. • Conjugate acid of a base is the particle that results when the base receives the proton from the acid. • Conjugate base of the acid is the particle that results when the acid donates a proton.

  6. Practice • Identify the conjugate acid/base pairs in the following: NH3(aq) + H2O(l)  NH4+(ag) + OH-(aq)

  7. Amphiprotic: Can act as either an acid or a base i.e has both a lone pair and an H-atom • Ex: H2O HCO3-(aq)) + H2O(l)  H2CO3(aq) + OH-(aq) HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq)

  8. Strong Acids and Bases • Completely dissociate in water into their ions (quantitative reactions) 100% HCl(aq) + H2O(aq) H3O+(aq) + Cl-(aq) 100% LiOH+ H2O(aq)LiOH(aq) + OH-(aq)

  9. As a result the [H3O+] in a solution of a strong acid is equal to the concentration of the acid. • Strong acids include HClO4 (perchloric), HI, HBr, HCl, H2SO4 (sulfuric), and HNO3 (nitric) • Strong bases include all oxides and hydroxides of alkali metals as well as alkaline earth metal oxides and hydroxides below beryllium. • The stronger the acid, the weaker it’s conjugate base and vice versa

  10. Weak Acids and Bases • Do NOTcompletely dissociate in water into their ions 1% CH3COOH(aq) + H2O(aq)↔ H3O+(aq) + CH3COO-(aq) 1% NH3(aq) + H2O(aq) ↔ NH4+(aq) + OH-(aq) • As a result, the concentration [H3O+] in a solution of a weak acid is always less than the concentration of the dissolved acid.

  11. Percent Ionization • % Ionization for strong acids is 100% • % Ionization for weak acids is < 100%

  12. Polyprotic Acids • Monoprotic acids contain only a single hydrogen ion that can dissociate. • Example: HCl • Polyprotic acids contain more than one hydrogen ions that can dissociate. • Example H2SO4, H3PO4

  13. Autoionization of Water • Water dissociates: H2O(l) + H2O(l) <--> H3O+(aq) + OH-(aq) What is the equilibrium constant (K) of this reaction? Kw = [H3O+][OH-] Kw is the ion product constant of water Kw = 1.0 x 10-14 @ SATP

  14. [H3O+] > [OH-]  acidic • [H3O+] < [OH-]  basic • [H3O+] = [OH-]  neutral

  15. Practice • There is a 0.25 mol/L solution of HBr(aq) • Calculate the hydrogen ion concentration • Calculate the hydroxide ion concentration • Strong acid – ionizes completely • Kw = [H3O+][OH-] = Kw = 1.0 x 10-14

  16. Practice • In a 0.13 mol/L solution of NaOH, what is the [H+] and [OH-]? • NaOH is hydroxide of an alkali metal so it is a STRONG base meaning [OH-]= [base] • Kw = [H3O+][OH-] = Kw = 1.0 x 10-14

  17. The pH Scale • Measures the acidity of a solution. • Measure [H+] in a solution. • Ranges from 0 to 14 • Distilled water is 7 (neutral) • Acids < 7 • Bases > 7 • A logarithmic scale • A pH of 1 is ten times more acidic then a pH of 2

  18. pH equations • pH = -log[H3O+] • [H3O+] = 10-pH • pOH = -log[OH-] • [OH-] = 10-pOH • pH + pOH = 14

  19. Practice • Calculate the pH of a solution of 1.24 x 10-4 M HCl • pH = -log[H3O+] • pH = -log[1.24 x 10-4 mol/L] • pH = 3.91

  20. Practice • If the normal pH of blood is 7.3, then find the pOH, [H3O+] and [OH-] • pH + pOH = 14 • 7.3 + pOH = 14 • pOH = 6.7 • [H3O+] = 10-pH • [H3O+] = 10-7.3 • [H3O+] = 5 x 10-8 • [OH-] = 10-pOH • [OH-] = 10-6.7 • [OH-] = 2 x 10-7

  21. Acid- Base Strength & Dissociation • Recall:Strongacids and bases dissociate quantitatively (>99.9%) in water • Weakacids and bases dissociate partiallyin water • When a weak acid or base is added to water dynamic equilibrium is established

  22. The Acid-Dissociation Constant, Ka For Weak Acids: All concentrations are those at equilibrium Note: the smaller the value of Ka, the weaker the acid

  23. Determine the Ka of propanoic acid (C2H5COOH(aq)) given that a 0.10 mol/L solution has a pH of 2.96. (Hint: use an ICE table) [H3O+] = 10-pH [H3O+] = 10-2.96 [H3O+] =0.00110 C2H5COOH(aq) CH3COO- + H3O+ I 0.10 mol/L o mol/L 0 mol / L C -x +x +x E

  24. The Base-Ionization Constant, Kb For Weak Bases: All concentrations are those at equilibrium Note: the smaller the value of Kb, the weaker the base

  25. Calculate the pH of a 3.6 X 10-3 mol/L solution of quinine (C20H24N2O2(aq)). Kb = 3.3 X 10-6 C20H24N2O2(aq)+ H2O  HC20H24N2O2 + + OH- I C E

  26. Relationship between Ka, Kb, & Kw Example: Consider acid HCN and conjugate base CN- HCN(aq) + H2O(l) H3O+(aq) + CN- (aq) Ka = [H3O+] [CN -] [HCN] Kb = [HCN] [OH -] [CN-] Ka Kb = [H3O+] [CN -][HCN] [OH -] [HCN] [CN-] Ka Kb = [H3O+] [OH -] Ka Kb = Kw

  27. Practice • The Kb for hydrazine, N2H4(g), a rocket fuel, is 1.7 x 10-6. What is the Ka of its conjugate acid, N2H5 (aq)? • KaKb = Kw • Ka(1.7 x 10-6)= 1.0 x 10-14 • Ka= 6.0 x 10-9

  28. Practice • Chloracetic acid, HC2H2O2Cl(aq) is a weak acid. Determine the pH of a 0.0100 mol/L solution of chloracetic acid if the Kb of the conjugate base is Kb= 7.35 x 10-12 . HC2H2O2Cl (aq) C2H2O2Cl- + H3O+ I C E • Ka Kb = Kw • Ka(7.35 x 10-12)= 1.0 x 10-14 • Ka=0.00136

  29. Neutralization Reactions • A salt is an ionic compound that results from a neutralization reaction • Acid + base  salt + water • Salts are strong electrolytes that completely ionize in water • Salts can affect the pH of a solution

  30. Neutral Salt Solutions • Strong acid + strong base • Both will dissociate completely • Therefore… • Salts containing an anion from a strong acid and cation from a strong base will be neutral • Ex: NaOH + HCl NaCl + H2O

  31. Acidic Salt Solutions • Strong acid + weak base • The acid dissociates completely, but the base only dissociates partially • Therefore… • Salts containing an anion from a strong acid and a cation from a weak base will be acidic • Ex: HCl + NH3  NH4Cl  NH4+ + Cl- • NH4+ will act as a weak acid

  32. Basic Salt Solutions • Weak acid + strong base • The base will dissociate completely but the acid will only dissociate partially • Therefore… • Salts containing an anion from a weak acid and a cation from a strong base will be basic • Ex: HC2H3O2 + NaOH NaC2H3O2 + H2O  Na+ + C2H3O2- • C2H3O2- will act as a weak base

  33. Buffers • Resist changes in pH when a moderate amount of acid or base is added • The acid and base components must not react in a neutralization reaction • Solutions of a weak acid and the salt of its conjugate base OR a weak base and the salt of its conjugate acid

  34. Acetic Acid/Sodium Acetate Buffer • Consider a buffered solution made by adding similar molar concentrations of acetic acid (CH3COOH) and its salt, sodium acetate (CH3COONa) • Sodium acetate ionizes completely in water: • When an acid is added to the buffer, the acetate ion reacts with the hydronium ion to neutralize the solution • When a base is added to the buffer, the acetic acid reacts with the hydroxide ions to neutralize the solution

  35. Buffer Examples • It is extremely important for blood to remain near it’s optimal pH of 7.4 • Any change greater than 0.2 is life-threatening • If the blood were not buffered, the acid absorbed by consuming a glass of orange juice would probably kill you

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