1. Acid-Base Equilibria Acids and bases are some of the more commonly encountered chemicals Acids and Bases control composition of blood and cell fluids, affect flavors, involved in digestion Bases used in house hold cleaners (NH 3 -based cleansers) Acid rain is an environmental problem Acids and bases are involved in reactions that produce polymers, synthetic fibers, dyes.

2. Arrhenius Acid & Base Acid: produces H + in aqueous solution Base: produces OH - in aqueous solution HCl(aq)  H + (aq) + Cl - (aq) NaOH(aq)  Na + (aq) + OH - (aq) Acid + base neutralization: H + (aq) + OH - (aq)  H 2 O(l) However, an H + cannot exist by itself in water

3. Brønsted-Lowry Acids and Bases Acid: proton donor Base: proton acceptor H + : PROTON since the H + consists of 1 proton and 0 electron HCN(aq) + NH 3 (aq)  NH 4 + (aq) + CN - (aq) acid base The H + is transferred from HCN to NH 3 HCN is said to have an acidic H, a hydrogen that can be donated as a H +

4. HCl has an “acidic” H +, but by itself cannot act as an acid However HCl(aq): HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) H 3 O + : hydronium ion

5. HCN - hydrogen cyanide HCN(aq) + H 2 O(l)  H 3 O + (aq) + CN - (aq) Only a fraction of HCN donate their H + to H 2 O HCN is a weak acid At equilibrium there is both CN - and un-dissociated HCN

6. In the Brønsted-Lowry theory: a strong acid is fully deprotonated in solution HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) a weak acid is only partially deprotonated in solution HCN(aq) + H 2 O(l)  H 3 O + (aq) + CN - (aq) Typically the solvent is water, but not necessarily. An acid that is strong in water, may be weak in another solvent

7. Brønsted-Lowry Base A proton acceptor. In most cases the molecule possesses a lone pair of electrons to which a H + can bond to. Example: Oxide, O 2- O 2- (aq) + H 2 O(l)  2 OH - (aq) Strong base since all O 2- (aq) forms OH - (aq)

8. NH 3 : a Brønsted base. The lone pairs on N in NH 3 can bond with a H +. NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) NH 3 (aq) is a weak base; at equilibrium both undissociated NH 3 (aq) and NH 4 + (aq) exist. A strong base is completely protonated in solution O 2- (aq) + H 2 O(l)  2 OH - (aq) A weak base is partially protonated in solution NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq) Strength depends on solvent

9. Solvent Leveling Since all strong acids are completely de-protonated in water (behave as though they were solutions of H 3 O + ) strong acids are “leveled” in water To compare acidity of acids that are strong acids in water, need to use a solvent in which the “acidity” of the acids differ Strong bases are leveled in water in the same way as strong bases.

10. Arrhenius definition restricted to water as a solvent However Brønsted-Lowry theory includes non-aqueous solvents CH 3 COOH (l) + NH 3 (l)  CH 3 COO - (am) + NH 4 + (am) am - denotes a species dissolved in ammonia Brønsted-Lowry includes acid/base in the absence of solvent Protons can be transferred in the gas phase: HCl(g) + NH 3 (g)  NH 4 Cl(s) Acid-base reaction does not have to involve the solvent HCN(aq) + NH 3 (aq)  NH 4 + (aq) + CN - (aq)

11. Conjugate Acids & Bases HCN(aq) + H 2 O(l)  H 3 O + (aq) + CN - (aq) acid conjugate base CN - (aq) is the conjugate base of HCN Brønsted-Lowry acids form conjugate bases Acid -----------> conjugate base donates H +

12. Brønsted-Lowry bases form conjugate acids NH 3 (aq) is the base; NH 4 + (aq) is the conjugate acid NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH + (aq) base -------------> conjugate acid accepts H +

13. An acid is a proton donor and a base is a proton acceptor. The conjugate base of an acid is the base formed when the acid has donated a proton. The conjugate acid of a base is the acid that forms when the base has accepted a proton.

14. Lewis Acids & Bases A Lewis base donates a lone pair of electrons A Lewis acid accepts a lone pair of electrons Lewis acids/bases are a broader definition than the Brønsted- Lowry definition H + is an electron pair acceptor; a Lewis acid Soluble metal oxides are strong bases

15. NH 3 + H 2 O  NH 4 + + OH - base acid Reactions between electron deficient and electron-rich molecules BF 3 (g) + NH 3 (g)  F 3 B - NH 3 (s) Lewis acid Lewis base

16. B-N bond is called a coordinate covalent bond; formed by the coordination of an electron-pair donor to an electron pair acceptor

17. Amphoterism H 2 O: acts as both an acid and a base - amphoteric H 2 O(l) + H 2 O(l)  H 3 O + (aq) + OH - (aq) OH - conjugate base of H 2 O H 3 O + conjugate acid of H 2 O HCO 3 - is amphoteric HCO 3 - (aq) + H 2 O(l)  H 3 O + (aq) + CO 3 2- (aq) HCO 3 - (aq) + H 2 O(l)  H 2 CO 3 (aq) + OH - (aq)

18. Water is amphiprotic - both an acid and a base When one molecule transfers a proton to another molecule of the same kind - autoprotolysis or autoionization 2 H 2 O (l)  H 3 O + (aq) + OH - (aq) An O-H bond is strong; the fraction of protons transferred is very small.

19. Calculate the equilibrium constant for the autoionization of H 2 O(l) 2 H 2 O (l)  H 3 O + (aq) + OH - (aq) K w = [H 3 O + (aq) ] [OH - (aq) ]  G r o =  G f o (H 3 O + (aq)) +  G f o (OH - (aq)) - 2  G f o (H 2 O(l)) = + 79.89 kJ/mol  G r o = - R T ln K w K w = 1.0 x 10 -14 at 298 K

20. K w = [H 3 O + (aq) ] [OH - (aq) ] [H 3 O + (aq) ] [OH - (aq) ] = 1.0 x 10 -14 K w is an equilibrium constant; the product of the concentrations of H 3 O + and OH - is always equal to K w. In pure water [H 3 O + (aq) ] = [OH - (aq) ] = 1.0 x 10 -7 M at 298 K

21. If the concentration of [OH - (aq)] in increased, then [H 3 O + (aq) ] decreases to maintain K w.

22. What are the molarities of H 3 O + and OH - in 0.0030 M Ba(OH) 2 at 25 o C? Ba(OH) 2 (aq)  Ba 2+ (aq) + 2 OH - (aq) Molarity of [OH - (aq)] = 0.0060 M [H 3 O + (aq)] = K w /[OH - (aq)] = 1.7 x 10 -12 M

23. pH Scale The concentration of H 3 O + can vary over many orders of magnitude A log scale allows a compact description of the H 3 O + concentration. - - - -

24. pH = - log [H 3 O + ] [H 3 O + ] = 10 - pH mol/L For pure water at 25 o C pH = - log (1.0 x 10 -7 ) = 7.00 For a change in pH by 1, H 3 O + concentration changes by 10 Higher pH, lower H 3 O + concentration pH of pure water is 7 pH of an acidic solution is less than 7 pH of a basic solution is greater than 7