1. Acid-Base Equilibria BLB 10 th Chapter 16

2. Examples of acids & bases

3. AcidsBases Sour (like vinegar)Bitter and slippery (like soap) React with bases to neutralize them and form salts React with acids to neutralize them and form salts Change indicator colors in opposite direction from base (e.g. litmus blue to red) Change indicator colors in opposite direction from acid (e.g. litmus red to blue) Aqueous solutions conduct electricity Liberate hydrogen in reactions with active metals React in aqueous solution with salts of heavy metals to form insoluble hydroxides or oxides

4. 16.1 Acids & Bases: A Brief Review  Arrhenius Definitions Acid – a substance that produces hydrogen ions (H + ) in water HA → H + + A - Base – a substance that produces hydroxide ions (OH - ) in water BOH → B + + OH -

5. 16.2 Brønsted-Lowry Acids & Bases  H + (proton) in water: H + + H 2 O → H 3 O + hydronium ion  Hydronium ion can hydrogen bond with more water molecules to form large clusters of hydrated hydronium ions.  H + and H 3 O + are used interchangeably.

6. 16.2 Brønsted-Lowry Acids & Bases  Brønsted-Lowry definitions acid – proton donor Neutral (HNO 3 ), anionic (HCO 3 - ), cationic (NH 4 + ) Must have a removable (acidic) proton base – proton acceptor Neutral (NH 3 ), anionic (CO 3 2- ) Must have a lone pair of electrons

7. Acid-Base Reactions HCl(aq) + H 2 O(l) → H 3 O + (aq) + Cl - (aq) NH 3 (aq) + H 2 O(l) ⇌ NH 4 + (aq) + OH - (aq) HCl(aq) + NH 3 (aq) → NH 4 + (aq) + Cl - (aq)

9. Acid-base reaction in non-aqueous media: HCl(g) + NH 3 (g) → NH 4 Cl(s)

10.  amphiprotic – capable of behaving as a Brønsted acid and Brønsted base  amphoteric – capable of behaving as a Lewis acid and Brønsted base (17.5)  neutralization: acid + base → salt + water  Conjugate acid/base pairs – differ by a single proton HA(aq) + H 2 O(l) → H 3 O + (aq) + A - (aq) acid + base conj. acid + conj. base

13. Relative Acid/Base Strength  Strength is a measure of the ability of an acid (or base) to donate (or accepts) a H +.  Stronger acids donate H + more readily.  Completely dissociate in water  Conjugate bases have negligible tendency to accept protons.  Weaker acids donate H + less readily.  Partially dissociate and establish equilibrium  Conjugate bases have some tendency to accept protons.  The stronger an acid, the weaker its conjugate base and vice versa.

15. p. 672

16.  Acid/base reactions proceed from the stronger acid-base pair to the weaker acid- base pair.  Common strong acids (p. 679): HClO 4, HClO 3, H 2 SO 4, HI, HBr, HCl, HNO 3 Monoprotic acid – capable of donating only one H + Polyprotic acid – capable of donating more than one H +  Common strong bases (p. 680): M(OH) n, where M = group I (n=1) & II (n=2) metals, except Be

17. Acid/Base Reactions

18. 16.3 The Autoionization of Water H 2 O(l) + H 2 O(l) ⇌ H 3 O + (aq) + OH - (aq)  K w = [H 3 O + ][OH - ] = [H + ][OH - ] = 1.0 x 10 -14 (@ 25°C)  K w – ion-product constant (or dissociation constant)  Pure water is neutral. Thus, [H 3 O + ] = [OH - ] = 1.0 x 10 -7 M @ 25°C  For an aqueous solution: [H 3 O + ] > [OH - ]acidic [H 3 O + ] = [OH - ]neutral [H 3 O + ] < [OH - ] basic

19. Working with K w

20. 16.4 The pH Scale  pH represents a solution’s acidity (@ 25°C. 0 ← 7 → 14 acid neutral base  See Table 16.1, p. 678 for summary.  See Figure 16.5, p. 679 for examples.  pH = −log[H 3 O + ] = −log[H + ] [H 3 O + ] = 10 -pH pOH = −log[OH - ] pH + pOH = 14 [OH - ] = 10 -pOH

21. p. 676

22. More common chemicals ChemicalpH Basic Windex10.57 Bleach9.58 Neutral Tap water*7.46 Acidic Alka Seltzer (in tap water)6.43 Distilled water**6.37 Flat Coke2.60 Toilet bowl cleaner1.04 6.0 M HCl−0.29 *CaCO 3 CO 3 - + H 2 O ⇌ HCO 3 - + OH - **CO 2 + H 2 O → H 2 CO 3

23. pH calculations

24. More about pH  pH does not necessarily indicate strength.  Measuring pH pH meters Acid-base indicators

25. p. 679

26. 16.5 Strong Acids and Bases  Strong acids & bases completely ionize. [HA] 0 = [H 3 O + ] → pH [MOH] 0 = [OH - ] → pOH → pH 2[M(OH) 2 ] 0 = [OH - ] → pOH → pH  H 3 O + is the strongest acid that can exist in water. (produced by all acids in water)  OH - is the strongest base that can exist in water. (produced by all bases in water)

27. pH problems End Test #1 material

28. 16.6 Weak Acids & 16.7 Weak Bases  Weak acids & bases do not completely ionize.  Weak acids establish an equilibrium in aqueous solution. HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A - (aq) HA(aq) ⇌ H + (aq) + A - (aq)  They do not readily donate or accept H + ’s.  [HA] 0 ≠ [H 3 O + ] [MOH] 0 ≠ [OH - ]

29. HA(aq) + H 2 O(l) ⇌ H 3 O + (aq) + A - (aq) HA(aq) ⇌ H + (aq) + A - (aq) K a ↑ acid strength ↑ For polyprotic acids: K a1 >> K a2 >> K a3 pK a = −log[K a ] pK a ↑ acid strength↓ Weak Acids & Acid-dissociation Constant

30. From p. 682 + more in Appendix D, p. 1115-1116

32. Weak Bases & Base-dissociation Constant  Weak bases establish an equilibrium in aqueous solution. B(aq) + H 2 O(l) ⇌ BH + (aq) + OH - (aq)

33. From p. 691 + more in Appendix D, p. 1115-1116

34. % Dissociation (or ionization)  % dissociation decreases as concentration increases (p. 686)

35. Weak acid/base Problems 1) K a (or K b ) from equilibrium pH 2) pH from K a (or K b ) 1. Identify as weak acid or base. 2. Write the chemical equilibrium. 3. Write the equilibrium constant expression. 4. Set up concentration table. (Ch. 15.5) 5. Solve for x. 6. Check with 5% rule. If greater than 5%, use quadratic equation. (type 2 only) 7. Complete problem.

36. The pH of a 0.10 M solution of propanoic acid (CH 3 CH 2 CO 2 H) is 2.94. Calculate the K a for propanoic acid.

37. Calculate the pH of a 1.0 M HF solution.

38. Calculate the pH of a 0.0010 M HF solution.

39. Calculate the pH of a 0.20 M solution of triethylamine N(CH 2 CH 3 ) 3.

40. 16.8 Relationship between K a and K b  For a conjugate acid/base pair: K a x K b = K w (derivation p. 693) pK a + pK b = pK w = 14.00

42. 16.9 Acid-Base Properties of Salt Solutions  Salt – ionic compound  Salts dissolve in water to produce ions.  Ions can also affect the pH.  Hydrolysis – reaction between an ion and water to produce H 3 O + or OH - F - (aq) + H 2 O(l) ⇌ HF(aq) + OH - (aq) NH 4 + (aq) + H 2 O(l) ⇌ H 3 O + (aq) + NH 3 (aq)

45. Which ions will undergo hydrolysis, i.e. react with water and affect the pH of the solution?  Anion: Conjugate base of a weak acid ► basic Conjugate base of a monoprotic strong acid ► neutral  Cation: Conjugate acid of a weak base ► acidic Group I & II metal ions ► neutral (exceptions Be 2+ and Mg 2+ ► acidic) Other metal ions ► acidic

46. Cation + Anion ►Acidic, basic, or neutral?

47. 16.10 Acid-Base Behavior and Chemical Structure  Binary Acids (HX) As bond strength increases, acid strength decreases. Group: size of X ↑ acid strength ↓ Period: electronegativity of X ↑ acid strength↑

49.  Oxyacids – acidic H attached to an oxygen atom Same # of OH groups and O atoms: central atom electronegativity ↑ acid strength ↑ HClO > HBrO > HIO Same central atom, Y: # O atoms ↑ acid strength ↑ HClO 4 > HClO 3 > HClO 2 > HClO  Carboxylic acids – contain −COOH or CO 2 H # electronegative atoms ↑ acid strength ↑

50. Oxyacids

51. 16.11 Lewis Acids and Bases  Lewis acid – electron-pair acceptor e - -poor compounds Metal ions  Lewis base – electron-pair donor Amines, NR 3 Ligands (see chapter 24.1) Every Brønsted base is a Lewis base, but not vice versa.

52. Lewis acid & base examples