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Acid Base Equilibria

Acid Base Equilibria. Types of Acids. H + vs H 3 O + Arrhenius Definition Acids – Increase [H+] when dissolved in water HCl Bases – Increase [OH-] when dissolved in water NaOH Br ø nsted-Lowry Definition Acids – Donates protons Bases – Accepts protons

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Acid Base Equilibria

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  1. Acid Base Equilibria

  2. Types of Acids • H+ vs H3O+ • Arrhenius Definition • Acids – Increase [H+] when dissolved in water • HCl • Bases – Increase [OH-] when dissolved in water • NaOH • Brønsted-Lowry Definition • Acids – Donates protons • Bases – Accepts protons • HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq) • HCl(g) + NH3(g) → NH4Cl(s) • NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)

  3. Conjugate Acid Base Pairs • HX(aq) + H2O(l) ↔ X-(aq) + H3O+(aq) • HNO2(aq) + H2O(l) ↔ NO2-(aq) + H3O+(aq) • What is the conjugate base of each of the following acids? • HClO4, H2S, HCO3-, NH4+ • What is the conjugate acid of each of the following bases? • CN-, SO42 -, H2O, HCO3-

  4. Practice • There are two possible reactions that HSO4- can have with water. Write the reaction in which it acts as an acid and another where it acts as a base. • When lithium oxide is dissolved in water, the solution turns basic from the reaction of the oxide ion(O2-) with water. Write the reaction and identify the conjugate acid base pairs.

  5. Acid Base Strength • Acid strength depends on how easily it gives up a proton • What are the strong acids? • Bases strength depends on how readily it accepts one. • What are the strong bases?

  6. Acid Base Strength • HX + H2O ↔ H3O+ + X- • Equilibrium lies on the side with weaker base • HCl + H2O ↔ H3O+ + Cl- • HC2H3O2 + H2O ↔ H3O+ + C2H3O2-

  7. Practice • Identify whether equilibrium lies predominantly to the left or right. • HSO4- + CO32 - ↔ SO42 - + HCO3- • HPO42 - + H2O ↔ H2PO4- + OH- • NH4+ + OH- ↔ NH3 + H2O Ans: R, L, R

  8. Autoionization of Water • H2O + H2O ↔ H3O+ + OH- • What is the Ke q of this reaction • At 25°C Ke q = 1x101 4 • What are the concentrations of each ion if each the concentration of [H3O+]=[OH-]? • How can we use this relationship to identify if a solution is acidic, basic, or neutral?

  9. pH Scale • pH scale goes from 0 to 14 • pH = - log[H+] • pOH = -log [OH-] • pH + pOH = 14

  10. Strong Acids and Bases • For strong acids and bases no equilibrium is reached. • For monoprotic strong acids: [H+] = [HA] • What is the pH of a 0.040M solution of HClO4 • pH = -log(0.040) = 1.40 • For strong bases: [OH-] = to # of OH-'s x [Base] • What is the pH of a 0.028M solution of NaOH and a 0.0011M solution of Ca(OH)2

  11. Practice • An aqueous solution of HNO3 has a pH of 2.34. What is the concentration of the acid? • What is the concentration of a solution of KOH for which the pH is 11.89; Ca(OH)2 for which the pH is 11.68? • Ans: #1 0.0046M, #2 7.8x10- 3M, 2.4x10- 3M

  12. Weak Acids • Weak acids only partially ionize and therefore reach equilibrium. • Ke q can be used to tell what extent it ionizes, how? • HA(aq) + H2O(l) ↔ A-(aq) + H3O+(aq) HA(aq) ↔ H+(aq) + A-(aq) • What is the equilibrium expression for this reaction? • Ka = acid-dissociation constant • How does Ka relate to acid strength?

  13. Weak Acids

  14. Using pH to find Ka • A Student prepared a 0.10M solution of formic acid (HCHO2) and measured its pH using a pH meter. The pH was found to be 2.38. Calculate the Ka for formic acid at this temperature. What percentage of the acid is ionized? • Ans: 1.8x10- 4, 4.2%

  15. Practice • Niacin, a type of B vitamin has the formula C5H4NCO2H. A 0.020M solution of the vitamin has a pH of 3.26. What percentage of the acid is ionized in this solution? What is the acid dissociation constant Ka for niacin? • Ans: 2.7%, 1.6x10- 5

  16. Using Ka to find pH • We need to know Ka and the initial concentration of the acid. Ka of acetic acid = 1.8x10- 5, [HC2H3O2] = 0.30M • Write the ionization equilibrium for the acid HC2H3O2(aq) ↔ H+(aq) + C2H3O2-(aq) • Write the equilibrium constant expression and value of Ka Ka = [H+][C2H3O2-]/[HC2H3O2] • Find equilibrium concentrations

  17. Using Ka to find pH • 4) Substitute the equilibrium concentrations into the equilibrium expression and solve for x. • ** x may be disregarded as long as it is less than 5% of the initial [ ] • x = 2.3x10- 3 = [H+] • pH = 2.64

  18. Practice • Calculate the pH of a 0.20M solution of HCN. Ka for HCN is 4.9x10- 1 0. Ans: pH = 5.00

  19. Calculating Percent Ionization • Calculate the percentage of HF molecules ionized in (a) 0.10M HF solution; (b) 0.010M HF solution. Find x then divide that by initial concentration. Ans: a) 7.9%, 23%

  20. Polyprotic Acids • Polyprotic acids have more than one ionizable H atoms and thus multiple Ka's. • Easier to remove 1s t proton than the next ones. • Finding pH for polyprotic acids • Compare size of Ka's • H2CO3(aq) ↔ H+(aq) + HCO3-(aq) Ka 1 = 4.3x10- 7 • HCO3-(aq) ↔ H+(aq) + CO32 -(aq) Ka 2 = 5.6x10- 1 1 • For most acids only Ka 1 is important

  21. Polyprotic Acids

  22. pH of Polyprotic Acids • What is the pH of a 0.0037M solution of H2CO3? What is the [CO32 -] in the solution? Ka 1= 4.3x10-7, Ka 2 = 5.6x10- 1 1 Ans: pH = 4.40, [CO32 -] = 5.6x10- 1 1M

  23. Practice • Calculate the pH and concentration of oxalate ion [C2O42 -], in a 0.020M solution of oxalic acid (H2C2O4). Ka 1= 5.9x10- 2, Ka 2= 6.4x10- 5 Ans: pH= 1.80, [C2O42 -]= 6.4x10- 5

  24. Weak Bases • B(aq) + H2O(l) ↔ HB+(aq) + OH-(aq) NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq) • Kb = base-dissociation constant • What is the Kb expression for NH3? • Calculate the [OH-] in a 0.15M solution of NH3. Kb= 1.8x10- 5 Ans: 1.6x10- 3M

  25. Weak Bases

  26. Ka and Kb • Ka x Kb = Kw • The larger Ka the lower the Kb • Calculate the Ka for HF if the Kb of F- is 1.5x10- 1 1

  27. pH of Salt Solutions • An anion that is the conjugate base of a strong acid will not affect the pH of a solution. ex. Br- • An anion that is the conjugate base of a weak acid will cause an increase in pH. ex. CN- • A cation that is the conjugate acid of a weak base will cause a decrease in pH. ex. NH4+ • With the exception of ions of group 1A and heavier members of group 2A, metal ions will cause a decrease in pH. • When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the ion with the largest ionization constant will have the greatest affect on pH.

  28. Practice • Predict whether the salt Na2HPO4 will form an acidic or basic solution on dissolving in water. • Predict whether the K2HC6H5O7 will form an acidic or basic solution in water. (look at citric acid) • Ans: basic, acidic

  29. Acid Strength • Molecules only donate protons if the H—X bond is polar, where the anion X is more electronegative than the H. • As you move from left to right X becomes more electronegative and acid strength increases. • CH4 < NH3 << H2O < HF • Strong H—X bonds are harder to break than weak ones. • The strength of the H—X bond decreases as the size of X increases. • HF versus HCl, H2S versus H2O • The more stable the resulting anion, X-, the stronger the acid.

  30. Binary Acid Strength

  31. Oxyacid Strength • Acids with OH groups and additional oxygen atoms bound to a central atom are called oxyacids. H2SO4 • For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom. • For oxyacids that have the same central atom, acid strength increases as the number of oxygen atoms attached increases. ex. Chloric acids.

  32. Lewis Acids and Bases • Lewis acids – electron pair acceptor • Increases the types of compounds that we can consider acids • Lewis bases – electron pair donor • All bases that are Brønsted-Lowry bases are Lewis bases. • NH3 + BF3 → NH3BF3 • Metal ions reacting with water act as Lewis acids

  33. Homework

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