1. ACIDS AND BASES …for it cannot be But I am pigeon-liver’d and lack gall To make oppression bitter… Hamlet

2. Learning objectives Name and write formulae for common acids and bases Describe acids and bases according to Arrhenius and Brønsted definitions Identify conjugate acid-base pairs Distinguish between strength and concentration in acid-base terminology Calculate and estimate pH of acids and bases Identify buffer solutions and explain action

3. ACIDS AND BASES The meaning of acid and base has changed over the years Arrhenius acid is one that generates protons when dissolved in water Arrhenius base is one that generates hydroxide ions when dissolved in water

5. Hydronium ion is the active ingredient of an acid Protons do not exist in solution CH 3 CO 2 H + H 2 O = H 3 O + + CH 3 CO 2 - Vinegar in water produces hydronium ions

6. Hydroxide ion is the active ingredient of a base NH 3 + H 2 O = NH 4 + + OH - Ammonia, a base, dissolves in water and produces hydroxide ions

7. Neutralization The mixing of an acid with a base: ACID + BASE = SALT + WATER The reaction of carbonic acid (CO 2 in H 2 O) to give limestone: H 2 CO 3 + Ca(OH) 2 = CaCO 3 + 2H 2 O ACID BASE WATER SALT

8. The essence of neutralization Elimination of the components of acid and base by combination to give H 2 O –Reflection of strong O – H bond H + + OH -  H 2 O ACID BASE

9. Brønsted and Lowry A broader definition of acids and bases In the reaction NH 3 + HCl = NH 4 Cl has all the elements of acid-base neutralization but no H 2 O as would be required in the Arrhenius definition Brønsted acid donates a proton Brønsted base accepts a proton

10. Brønsted acid HCl + H 2 O = H 3 O + + Cl -

11. Brønsted base NH 3 + H 2 O = NH 4 + + OH - water NH 3 + HCl = NH 4 + Cl - No water

13. Substances can be both acids and bases – depends on environment Note that in one instance H 2 O behaves like a base – accepting protons, and in another, behaves like an acid – donating protons HCl + H 2 O = H 3 O + + Cl - In presence of an acid H 2 O is a base NH 3 + H 2 O = NH 4 + + OH - In presence of a base H 2 O is an acid

14. The products are themselves acids and bases

15. Equilibrium: solution contains mixture of all components

16. Conjugate Conjugate acids and bases Conjugate Conjugate acid-base pair differs by H + HA + B ↔ A - + HB + Conjugate acid-base pair Conjugate base Conjugate acidbase acid

17. Strength and concentration Not all acids completely donate the protons to water molecules in solution HA + H 2 O  A - + H 3 O + The degree of ionization is described by strength The total number of moles per unit volume is described by concentration

18. Equilibrium for a weak acid Measuring acid strength with K a –Weak acid has small K a –Strong acid has large K a

20. Changing concentration does not change strength Strength refers to degree of ionization: –Strong is completely ionized (100 %) –Weak is partly ionized (1 % - 1:10 6 ) Concentration refers to number of moles per unit volume An acid (or base) can be strong and concentrated, weak and concentrated, strong and dilute, weak and dilute

21. Relative strength of conjugate pair The conjugate base of a very strong acid is itself very weak The conjugate acid of a very strong base is itself very weak

22. Amphotericity A substance that behaves as an acid and a base is amphoteric. Water is a good example

23. Ionization of water Even in pure water, a fraction of the molecules are ionized and the concentrations of OH - and H 3 O + are equal H 2 O + H 2 O = H 3 O + + OH - [H 3 O + ] = [OH - ] Concentration

24. In all aqueous solutions, product of concentrations is a constant [H 3 O + ][OH - ] = constant Increasing [H 3 O + ] decreases [OH - ] (acidic conditions) Increasing [OH - ] decreases [H 3 O + ] (basic conditions)

25. The pH scale – reduces large range of numbers to small In water K W = [H 3 O + ][OH - ] = 10 -14 pH = - log 10 [H 3 O + ] Range of [H 3 O + ] 10 M – 10 -14 M Range of pH -1 to +14 Low pH = acid; high pH = basic pH = 7 = neutral

26. Relating pH to [H 3 O + ] For pH, take exponent of [H 3 O + ], change sign –10 -1 M (0.1 M) HCl has pH = 1 –Pure water has [H 3 O + ] = 10 -7 M, pH = 7 –Ammonia has [H 3 O + ] =10 -11 M, pH = 11 Note: change of 1 unit in pH is factor of ten

27. Acid strength and pH pH is determined by the ratio of the weak acid:conjugate base and the value of K a

28. Estimating pH Estimating pH is often more useful than doing exact calculations Smaller pH value means larger H + concentration Estimating pH

29. Smaller pH value means larger H + concentration pH = - log 10 [H 3 O + ] [H 3 O + ] = 10 -4 pH = - log 10 (10 -4 ) = 4 [H 3 O + ] = 4 x 10 -4 pH = - log 10 (4 x 10 -4 ) = - log 10 4 - log 10 (10 -4 ) = 4 – log 10 4 =4 – 0.602 = 3.4 Concentration 1 x 10 -4 M pH = 4 Concentration 4 x 10 -4 M pH = 3.4

30. pH and pOH In any aqueous solution, [H + ] x [OH - ] = 10 -14 Obtain [H + ] given [OH - ] pOH = -log 10 [OH - ] pH + pOH = 14

31. Relationships between [H + ], pH and [OH - ]

32. Buffer solutions control pH In biological systems it is often critical for stable function to maintain a specific pH environment. Buffer solutions perform this function – they resist large changes in pH when either acids or alkalis are added

33. Buffer solution contains a combination of a weak acid and its salt

34. Mechanism follows Le Chatelier’s principle –Addition of H + converts CH 3 CO 2 - to CH 3 CO 2 H CH 3 CO 2 - + H +  CH 3 CO 2 H –Addition of OH - converts CH 3 CO 2 H to CH 3 CO 2 - CH 3 CO 2 H + OH -  CH 3 CO 2 - + H 2 O –In general, Add OH - shifts equilibrium to right (HA  A - ) HA + H 2 O  A - + H 3 O + Add H + shifts equilibrium to left (A -  HA)