1. 1 Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases

2. Acids & Bases Acids are some of peoples' favorite chemicals. Everyone's favorite soft drink is a dilute acid solution. Your own stomach contains the strong acid : HCl. Citrus fruits contain citric acid. If wine is too aged - exposed to oxygen, it turns sour - it forms acetic acid. Sulfuric Acid is the top commercially produced chemical in the United State. Although much of it is used in the steel and petroleum refining industries, several million tons of sulfuric acid are used to make Jello.

3. PROPERTIES OF ACIDS AND BASES ACIDS BASES

4. Arrhenius definition acid: produces hydronium ion (H 3 O + ) in aqueous solution base: produces hydroxide ion (OH – ) in aqueous solution

5. Brønsted definition acid: base: donates a proton (hydrogen ion, H + ) accepts a proton HF + H 2 O H 3 O + + F – acidconjugate acid baseconjugate base A conjugate acid is formed by adding a proton to something. A conjugate base is formed by removing a proton from something. NH 3 + H 2 O NH 4 + + OH – baseconjugate acid conjugate base

6. HAH 3 O + + A –  If HA is a stronger acid then A – is a weaker base. If HA is a weaker acid then A – is a stronger base. BBH + + OH –  If B is a stronger base then BH + is a weaker acid. If B is a weaker base then BH + is a stronger acid. Relative strengths of conjugate acid-base pairs

7. You should memorize the names and formulas of the 6 STRONG ACIDs, i.e., HCl, HBr, HI, HClO 4, HNO 3, and H 2 SO 4.STRONG ACIDs The organic acid present in vinegar, acetic acid, is a common WEAK ACID. The common STRONG BASES contain the hydroxide ion (OH - ). Ammonia (NH 3 ), a common WEAK BASE, is that smelly stuff your Grandma used in a dilute solution to clean windows

8. Ion Product Constant of Water Water is an important solvent. Universal solvent Biological solvent Small size Density of water is greater than ice Very polar Hydrogen Bonding

9. Self-ionization of Water amphiprotic Water is an amphiprotic substance that can act either as an acid or a base. HC 2 H 3 O 2(aq) + H 2 O (l) H 3 O + + C 2 H 3 O 2 - (aq) acid base acid base H 2 O (l) + NH 3(aq) NH 4 + (aq) + OH - (aq) acid base acid base

10. Self-ionization of Water When water molecules react with one another to form ions. H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH - (aq) (10 -7 M) (10 -7 M) K w = [ H 3 O + ] [ OH - ] = 1.0 x 10 -14 at 25 o C Note: Note: [H 2 O] is constant and is already included in K w. ion product of water ion product of water

11. pH and pOH We need to measure and use acids and bases over a very large concentration range. pH and pOH are systems to keep track of these very large ranges. pH = - log [H 3 O + ] pOH = - log [OH - ] pH + pOH = 14

12. pH Calculations Determine the following. pH = - log [H + ] pH of 6.7x10 -3 M H + = 2.2 pH of 5.2x10 -12 M H + = 11.3 [H + ], if the pH is 4.5 = 3.2 x 10 -5 M H +

13. pOH Examples Determine the following. pOH = - log [OH - ] = 14 - pH pOH of 1.7 x 10 -4 M NaOH pOH = 3.8 pH = 10.2 pOH of 5.2 x 10 -12 M H + pOH = 2.7 pH = 11.3 [OH - ], if the pH is 4.5 pOH = 9.5 [OH - ] = 3.2 x 10 -10 M

14. pH Scale A log based scale used to keep track of the large change important to acids and bases. 14 7 0 10 -14 M 10 -7 M 1 M Very Neutral Very Basic Acidic When you add an acid, the pH gets smaller. When you add a base, the pH gets larger.

15. pH of Some Common Materials Substance pH 1 M HCl0.0 Lemon juice2.3 Coffee5.0 Pure Water7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 1M NaOH 14.0

16. HA + H 2 OH 3 O + + A – K c = [H 3 O + ][A – ] [HA][H 2 O] ~ constant (55 M)  K c ·[H 2 O] = K a = [H 3 O + ][A – ] [HA] pK a = -log K a if pK a = 5 then K a = 10 –5 if pK a = 8 then K a = 10 –8 stronger acid weaker acid acid dissociation constant Definitions K a, pK a

17. Acid Ionization Constant, K a Acid ionization constants let us define weak, moderate and strong acids. K a < 10 -3 ; it is a weak acid. K a = 10 -3 to 1; it is a moderate acid. K a > 1; it is a strong acid.

18. B + H 2 OBH + +OH – K c = [BH + ][OH – ] [B][H 2 O] ~ constant  K c ·[H 2 O] = K b = [BH + ][OH – ] [B] pK b = -log K b if pK b = 4 then K b = 10 –4 if pK b = 9 then K b = 10 –9 stronger base weaker base base dissociation constant Definitions K b, pK b

19. K a and K b Values For weak acids and bases: K a and K b always have values that are smaller than one. Acids with a larger K a are stronger than ones with a smaller K a. Bases with a larger K b are stronger than ones with a smaller K b. K a x K b = K w Most acids and bases are considered weak.

20. pK a and pK b Concepts The negative logarithms of K a and K b are useful in the same way as pH. pK a = - log K a pK b = - log K b pK a + pK b = 14.00 The larger that the value of pK a is, the weaker the acid. The larger that the value of pK b is, the weaker the base.

21. H 2 O + H 2 OH 3 O + + OH – K w = [H 3 O + ][OH – ] = 10 –14 (constant at 25ºC)pK w = pH + pOH = 14  [H 3 O + ] = 10 –14 [OH – ] pH = 14 - pOH K w : autodissociation of water

22. HAH 3 O + + A – K a = [H 3 O + ][A – ] [HA] A–A– HA + OH – K b = [HA][OH – ] [A – ] usually only one or the other given in a table K a · K b = [H 3 O + ][A – ] [HA] · [HA][OH – ] [A – ] = [H 3 O + ][OH – ] = K w K a · K b = K w for a conjugate acid-base pair or K a =or K b = KwKaKwKa KwKbKwKb K a and K b for conjugate acid-base pairs

23. III. pH Calculations A. Strong acids and bases 100% dissociated  for strong acid:[H 3 O + ] eq = [HA] I base:[OH – ] eq = [B] i e.g., 1.0 x 10 –3 M HCl [H 3 O + ] = pH = [OH – ] =pOH = e.g., 2.5 x 10 –2 M NaOH [OH – ] =pOH = [H 3 O + ] = pH = [OH – ]-1.60 + 14 log pH = 12.40

24. III. pH Calculations B. Weak acids and bases HAH 3 O + + A – K a = [H 3 O + ][A – ] [HA] BBH + + OH – K b = [BH + ][OH – ] [B] Solve equilibrium expressions e.g., What is the pH of 0.10 M HC 2 H 3 O 2 ? (K a = 1.8 x 10 –5 ) HC 2 H 3 O 2 H 3 O + + C 2 H 3 O 2 – 1.8 x 10 –5 = x 2 (0.10 - x) assume x << 0.10 1.8 x 10 –5 = x 2 (0.10) x = [H 3 O + ] = 1.3 x 10 –3 M (assumption valid) pH = 2.87  If [HA] i  400·K a then x << [HA] i (If not, then have to solve quadratic.)

25. Buffers Solutions that resist change in pH when small amounts of acid or base are added. Two types: Two types: weak acid and its salt. weak base and its salt. HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) Add OH - Add H + shift to right shift to left Based on Le Châtelier’s Principle.

26. III. pH Calculations C. Polyprotic acids H 2 SO 4, H 2 SO 3, H 2 CO 3, etc. H2AH2AH 3 O + + HA – K a1 = [H 3 O + ][HA – ] [H 2 A] HA – H 3 O + + A 2– K a2 = [H 3 O + ][A 2– ] [HA – ] (Usually, K a1 >> K a2 ) Assume:1) [H 2 A], [H 3 O + ], and [HA – ] can be determined from the 1 st step. (i.e., HA – dissociates only very little.) 2) [A 2– ] can be determined from the 2 nd step. Lose their protons in separate steps:

27. Buffers and Blood Control of blood pH. Control of blood pH. Oxygen is transported primarily by hemoglobin in the red blood cells. CO 2 transported both in plasma and the red blood cells. CO 2 (aq) + 2 H 2 O H 2 CO 3 (aq) H 3 O + (aq) + HCO 3 - (aq) Carbonate Buffer

28. Buffers and Blood The amount of CO 2 helps control blood pH. Too much CO 2 Too much CO 2 - Respiratory arrest, pH goes down, acid level goes up. acidosis acidosis Solution Solution - ventilate and give bicarbonate via IV. Too little CO 2 Too little CO 2 - Hyperventilation, anxiety, pH goes up, acid level goes down. alkalosis alkalosis Solution Solution - re-breathe CO 2 in paper bag to raise level.

29. Quantitative Aspects of Buffers K a for a weak acid: HAH + + A - K a = [H + ] [A - ] [HA] Henderson-Hasselbalch Equation: Henderson-Hasselbalch Equation: pH = pK a + log [anion] [acid]

30. Neutralization The reaction of an acid with a base to produce a salt and water. HCl + NaOHNaCl + H 2 O We do this when we use antacids. - titrations Neutralization can be used to determine the amount of acid or base in a sample. - titrations

31. Titrations Analytical methods based on measurement of volume. If the concentration of an acid is known, the concentration of the base can be found. If we know the concentration of the base, then we can determine the amount of acid. All that is needed is some calibrated glassware and either an indicator or pH meter.

32. TItrations Buret Buret - volumetric glassware used for titrations. It allows you to add a known amount of your titrant to the solution you are testing. If a pH meter is used, the equivalence point can be measured. An indicator will give you the endpoint.

33. Indicator Examples Acid-base indicators are weak acids that undergo a color change at a known pH. bromothymol blue phenolphthaleinmethyl red