1. Acids, Bases and Equilibria

2. Overview Definitions Strong acids pH Water equilibrium Weak acids Buffers Other equilibria LeChatlier’s Principle

3. Defining Acids and Bases Arrhenius model –Acid – Proton donor – e.g. HCl –Base – Hydroxide donor – e.g. NaOH But how about Sodium Carbonate?

4. Defining Acids and Bases - 2 Brønsted-Lowery model –Acid – Proton donor – same as Arrhenius –Base – Proton ACCEPTOR Aha – so Na 2 CO 3 IS basic! Na 2 CO 3 + 2HCl  2NaCl + H 2 CO 3

5. Strong Acids and Bases Ionic solids like NaOH; completely form ions in water: NaOH + H 2 O  Na + + OH - + H 2 O Covalent molecules like HCl completely IONIZE in water: HCl + H 2 O  H 3 O + + Cl - H 3 O + is “hydronium” ion – no bare protons

6. Defining pH Remember pH? –Less than 7 = acid –More than 7 = base But what does it mean? pH is a measure of the concentration of hydronium ion in water pH = - log [H 3 O + ]

7. Translation: - log Suppose we have 0.1M HCl solution Since it is fully ionized, we have 0.1M H 3 O + 0.1 = 10 -1 -log (10 -1 ) = 1! Therefore pH of this acid solution is 1

8. Getting the pH of a base Even in base, pH measures hydronium ion H 3 O + and OH - are related by the equilibrium of water

9. See p. 611

10. So, what’s equilibrium? Second grade analogy – see-saw In an equilibrium situation, reactions or changes go both ways Hold ice and water at 0 o –Water melts and ice freezes at the same time –“Dynamic” equilibrium

11. Equilibrium 2 Form a saturated solution of NaCl –NaCl dissolves; –Same time, NaCl forms new crystals

12. Water is amphoteric H 2 O + H 2 O  H 3 O + + OH - Reaction moves to right at same rate as to the left Water is being both an acid and a base On the other side, “conjugates” are formed –H 3 O + is the conjugate acid of H 2 O –OH - is the conjugate base of H 2 O

13. Water’s “Equilibrium Constant” K = [H 3 O + ][OH - ] K = 10 -14 Square root of 10 -14 = 10 -7 [H 3 O + ] = [OH - ] = 10 -7 Therefore pH of pure water = 7!

14. So now to pH of bases: Find the pH of 0.01M NaOH Fully ionized; therefore 0.01M OH - [OH - ] = 10 -2 K = [H 3 O + ][OH - ] 10 -14 = [H 3 O + ] * 10 -2 10 -12 = [H 3 O + ]; pH = 12 OR pK = pH + POH 14 = pH + 2 12 = pH

15. And Weak Acids (or Bases) A weak acid is one which is NOT fully ionized Acetic Acid == HAc (or CH3COOH) HAc + H 2 O  H 3 O + + Ac - –Acetate ion is the conjugate base of Acetic acid –At equilibrium, HAc is largely NOT ionized Because the reaction goes both ways, Acetate can accept a proton: from H 3 O + OR from H 2 O Ac - + H 2 O  HAc + OH- Yes, a salt made from a weak acid and a strong base is basic!

16. Typical weak acids: Acetic acid CH 3 COOH Carbonic acid H 2 CO 3 Second or third H + of phosphoric: H 2 PO 4 -1, HPO 4 -2

17. So let’s make a “Buffer” A buffer is a solution of a weak acid and the strong base salt of its conjugate base: Acetic acid and sodium acetate 0.1M 0.1M CH 3 COOH + H 2 O  H 3 O + + CH 3 COO -

18. And let’s add some acid First to water: –Add 0.01M HCl to water –pH becomes 2, right? ([H 3 O + ] = 10 -2 ) But add the same acid to the buffer: 0.1M 0.1M CH 3 COOH + H 2 O  H 3 O + + CH 3 COO - 0.11M 0.09M [H 3 O + ] is almost unaffected! pH stays “same”

19. Buffer: definition A buffer is a solution of a weak acid and its conjugate base OR A buffer is a solution of a weak base and its conjugate acid Which resists changes in pH when small amounts of strong acid or base are added Blood is (or contains) a buffer!

20. LeChatlier’s Principle Notice that a buffer takes advantage of a reversible reaction which shifts away from the species we add: H 3 O + or OH - LeChatlier said ANY system in equilibrium will shift in such a way as to minimize the effect of a stress applied

21. Illustration of Principle

22. Wasn’t that fun????? Definitions Strong acids pH Water equilibrium Weak acids Buffers Other equilibria LeChatlier’s Principle