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. NH 3 (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: NH 3(aq) + H 2 O (l)  NH 4 + (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: H 2 O HCO 3 - (aq) ) + H 2 O (l)  H 2 CO 3(aq) + OH - (aq) HCO 3 - (aq) + H 2 O (l)  CO 3 2- (aq) + H 3 O + (aq)

8. Strong Acids and Bases Completely dissociate in water into their ions (quantitative reactions) 100% HCl (aq) + H 2 O (aq)  H 3 O + (aq) + Cl - (aq) 100% LiOH + H 2 O (aq)  LiOH (aq) + OH - (aq)

9. As a result the [H 3 O + ] in a solution of a strong acid is equal to the concentration of the acid. Strong acids include HClO 4 (perchloric), HI, HBr, HCl, H 2 SO 4 (sulfuric), and HNO 3 (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 NOT completely dissociate in water into their ions 1% CH 3 COOH (aq) + H 2 O (aq) ↔ H 3 O + (aq) + CH 3 COO - (aq) 1% NH 3(aq) + H 2 O (aq ) ↔ NH 4 + (aq) + OH - (aq) As a result, the concentration [H 3 O + ] 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 H 2 SO 4, H 3 PO 4

13. Autoionization of Water Water dissociates: H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH - (aq) What is the equilibrium constant (K) of this reaction? K w = [H 3 O + ][OH - ] K w is the ion product constant of water K w = 1.0 x 10 -14 @ SATP

14. [H 3 O + ] > [OH - ]  acidic [H 3 O + ] < [OH - ]  basic [H 3 O + ] = [OH - ]  neutral

15. Practice There is a 0.25 mol/L solution of HBr (aq) a)Calculate the hydrogen ion concentration b)Calculate the hydroxide ion concentration Strong acid – ionizes completely K w = [H 3 O + ][OH - ] = K w = 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] K w = [H 3 O + ][OH - ] = K w = 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[H 3 O + ] [H 3 O + ] = 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[H 3 O + ] 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, [H 3 O + ] and [OH - ] pH + pOH = 14 7.3 + pOH = 14 pOH = 6.7 [H 3 O + ] = 10 -pH [H 3 O + ] = 10 -7.3 [H 3 O + ] = 5 x 10 -8 [OH - ] = 10 -pOH [OH - ] = 10 -6.7 [OH - ] = 2 x 10 -7

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

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

23. Determine the K a of propanoic acid (C 2 H 5 COOH (aq) ) given that a 0.10 mol/L solution has a pH of 2.96. (Hint: use an ICE table) [H 3 O + ] = 10 -pH [H 3 O + ] = 10 -2.96 [H 3 O + ] =0.00110 C 2 H 5 COOH (aq)  CH 3 COO - + H 3 O + I0.10 mol/Lo mol/L 0 mol / L C-x+x+x E

24. The Base-Ionization Constant, K b 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 (C 20 H 24 N 2 O 2(aq) ). K b = 3.3 X 10 -6 C 20 H 24 N 2 O 2(aq) + H2O  HC 20 H 24 N 2 O 2 + + OH - I C E

26. Relationship between K a, K b, & K w Example: Consider acid HCN and conjugate base CN - HCN (aq) + H 2 O (l)  H 3 O + (aq) + CN - (aq) K a = [H 3 O + ] [CN - ] [HCN] K b = [HCN] [OH - ] [CN - ] K a K b = [H 3 O + ] [CN - ] [HCN] [OH - ] [HCN] [CN - ] K a K b = [H 3 O + ] [OH - ] K a K b = K w

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)? K a K b = K w K a (1.7 x 10 -6 )= 1.0 x 10 -14 K a = 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 K b of the conjugate base is K b = 7.35 x 10 -12. HC2H2O2Cl (aq)  C2H2O2Cl - + H 3 O + I C E K a K b = K w K a (7.35 x 10 -12 )= 1.0 x 10 -14 K a =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 (CH 3 COOH) and its salt, sodium acetate (CH 3 COONa) 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