1. 1 Acid-Base Acid-Base Chemistry acidH + (proton) donor An acid is a H + (proton) donor. acid = H-A General formula of acid = H-A. H-A H + + A - acid proton conjugate base

2. 2 When H-A dissociates to H + and A – in aqueous (water) solution, the “hydronium ion” or H 3 O + forms: H-A+ H 2 O H 3 O + + A – H-A + H 2 O H 3 O + + A – H-AA – H-A + H-O-H H-O-H + A – H HH H + means

3. 3 BaseH + acceptor A Base is a H + (proton) acceptor. baseunshared pair of electrons A base has an unshared pair of electrons to share with H +. base :B :Base. we can represent a base by the general formula :B or :Base. :Base + H + H-Base + from an conjugate acid acid

4. 4 When a Base accepts protons from water, it increases the concentration of HO – (hydroxide ions) in solution: :B B + :B + H-O-H H-B + + – O-H OR :B + H 2 O HB + + – OH

5. 5 Examples H HNO 3 + H 2 O H 3 O + + NO 3 – nitric acid nitrate ion acidconjugate base acid conjugate base H HCl + H 2 O H 3 O + + Cl – hydrochloric acid chloride ion acidconjugate base acid conjugate base

6. 6 :NH 3 :NH 3 + H 2 O + NH 4 + – OH ammonia ammonium ion baseconjugate acid baseconjugate acid

7. 7 Strengths of Acids & Bases strong acid HA completely dissociates A strong acid HA completely dissociates to H + (= H 3 O + ) and A – in H 2 O: H-A H-A + H 2 O H 3 O + + A –. 1 mole to start gives 1 mole H 3 O + + 1 mole A –

8. 8 The common strong acids are: HClhydrochloric acid HBrhydrobromic acid HIhydroiodic acid HNO 3 nitric acid H 2 SO 4 sulfuric acid [HClO 4 perchloric acid] don’t memorize [ ]

9. 9 strong base A strong base is one which completely dissociates to provide HO – in water. These bases are ionic metal hydroxides: Metal cation + hydroxide anion MOH M + + – OH 1 mole 1 mole + 1 mole

10. 10 There are 4 common strong bases: LiOH lithium hydroxide NaOH sodium hydroxide KOH potassium hydroxide Ba(OH) 2 barium hydroxide

11. 11 Some Advice... memorize lists It is best to memorize the lists of strong acids & strong bases -- assume all others are weak.

12. 12 But Also, LEARN HOW to write the equations for the reaction of each chemical as a strong acid strong base. or strong base.

13. 13 What about WEAK acids & bases? weak incompletely ionized A weak acid or weak base is incompletely ionized or dissociated in water solution. The ionization reaction of a weak acid or weak base is a reversible equilibrium!!

14. 14 Example CH 3 CO-H + H 2 O CH 3 CO – + H 3 O + O O acetic acid acetate before equilibrium: 1 mole 0 0 after equilibrium: 0.98 mole 0.02 0.02

15. 15 The position of equilibrium is different for each weak acid and weak base. The equilibrium constant describes how much ionization has occurred.

16. 16 CH 3 CO-H + H 2 O CH 3 CO – + H 3 O + O O For acetic acid: K eq = [CH 3 CO 2 – ][H 3 O + ] [CH 3 CO 2 H][H 2 O] OR K a K eq [H 2 O] = K a = [CH 3 CO 2 – ][H 3 O + ] [CH 3 CO 2 H]

17. 17 We use K a as a measure of the strength of an acid because [H 2 O] is very large and does not effectively change as a result of the ionization reaction. Typical K a values for weak acids are in the range of 1 x 10 -1 to 1 x 10 -10

18. 18 Sample K a values: do NOT memorize!!!! WhatK a = CH 3 CO 2 H acetic acid 1.8x10 -5 H 2 O water 1 x 10 -16 CH 3 CH 2 OH alcohol 1 x 10 -17

19. 19 What do these small K a values mean? * Solutions of weak acids and bases at equilibrium contain both the unionized (H-A) and ionized forms (H + and A – ). * The largest species present is H-A (unionized acid.)

20. 20 And... * Any compound having K a smaller than H 2 O is considered “neutral” = “not acidic in water”. Thus, alcohols are not acidic compounds. They are neutral substances.

21. 21 Monoprotic release one H + per molecule Monoprotic acids are those which can release one H + per molecule of acid. Polyproticrelease more than one H + Polyprotic acids can release more than one H + per molecule of acid.

22. 22 Examples Monoprotic H Monoprotic HCl Hydrochloric Acid DiproticH 2 Diprotic H 2 SO 4 Sulfuric acid TriproticH 3 Triprotic H 3 PO 4 Phosphoric Acid

23. 23 Amphiprotic (Amphoteric) Compounds can act as acids or bases.Examples HCO 3 – bicarbonate ion as acidH + as acid HCO 3 – CO 3 -2 + H + as base H as base HCO 3 – + H + H 2 CO 3

24. 24 Dihydrogen phosphate H 2 PO 4 – amphoteric Dihydrogen phosphate H 2 PO 4 – is another important amphoteric ion. as acid H 2 – H + H 2 PO 4 – HPO 4 -2 + H + as base – H H 2 PO 4 – + H + H 3 PO 4

25. 25 Water is the most important amphoteric compound of all. H H-O-H + H-O-H HO – + H 3 O + “acid” “base” “acid” “base” OR K eq = [HO – ][H 3 O + ] OR [H 2 O][H 2 O] K w = [HO – ][H 3 O + ] K eq [H 2 O] 2 = K w = [HO – ][H 3 O + ]

26. 26 In pure H 2 O at room temperature, K w has a value of 1 x 10 -14. Since K w = [HO – ][H 3 O + ], this means [HO – ]=[H 3 O + ] = 1x10 -7 moles/liter in neutral water.

27. 27 pH is defined as pH = -log ([H 3 O + ]) pH = -log ([H 3 O + ]) of a solution. A logarithm tells what power of 10 = a certain number. y = 10 x log(y) = x That is, if y = 10 x, then log(y) = x. Thus, 100 = 10 2, so log(100) = 2.

28. 28 We saw that for pure H 2 O, [H 3 O + ] = 1 x 10 -7. Since -log (1 x 10 -7 ) = 7, neutral water has pH = 7 we say neutral water has pH = 7.

29. 29 Recall: [HO – ] and [H 3 O + ] are related through the equilibrium 2 H 2 O HO – + H 3 O +. AcidsH + donors increase[H 3 O + ] Acids are H + donors and cause an increase in [H 3 O + ]. So: pH decreases. Also: [HO – ] decreases. Acidic solutions have pH <7. (in the range 0 to 6.99)

30. 30 Again: 2 H 2 O HO – + H 3 O +. Basedecreases [H 3 O + ] A Base increases [HO - ] or decreases [H 3 O + ]. So pH increases: basicsolution has a pH >7 a basic or “alkaline” solution has a pH >7 (in the range 7+ to 14).

31. 31 14 0 7Neutral [HO – ] = [H 3 O + ] pH Acidic [HO – ] < [H 3 O + ] Basic [HO – ] > [H 3 O + ]