1. Chapter 15 Acids and Bases.

2. 2 Acid-Base Theories In defining what is considered to be an acid and what is considered to be a base, three theories have been proposed: Arrhenius acid-base theory Brønsted-Lowry acid-base theory Lewis acid-base theory We will see that each subsequent theory builds upon what was stated in the previous theory.

3. 3 The Arrhenius Theory In the Arrhenius theory of acids, an acid dissolved in water increases the concentration of hydronium ions H 3 O + in the solution: Arrhenius acid HA: HA (aq) + H 2 O (l)  H 3 O + (aq) + A - (aq) In this reaction we see all Arrhenius acids contain protons (H + ) that are donated to water

4. 4 Arrhenius Theory – Acid Strength In the Arrhenius theory of acids, a strong acid COMPLETELY reacts with water, so there is no HA left at the end of the reaction: HA (aq) + H 2 O (l) → H 3 O + (aq) + A - (aq)

5. 5 Arrhenius Theory – Acid Strength In the Arrhenius theory of acids, a weak acid reacts with water until an equilibrium is reached where HA is still present in the equilibrium mixture: HA (aq) + H 2 O (l) ⇌ H 3 O + (aq) + A - (aq)

6. 6 The Arrhenius Theory In the Arrhenius theory of bases, a base dissolved in water increases the concentration of hydroxide ions OH - in the solution: Arrhenius base M(OH) x : M(OH) x (aq)  M x+ (aq) + x OH - (aq) In this reaction we see all Arrhenius bases contain hydroxide (OH - )

7. 7 Arrhenius Theory – Base Strength In the Arrhenius theory of acids, a strong base COMPLETELY dissociates in water, so there is no M(OH) x left at the end of the reaction: M(OH) x (aq) → M x+ (aq) + x OH - (aq)

8. 8 Arrhenius Theory – Acid Strength In the Arrhenius theory of bases, a weak base only partially dissociates in water until an equilibrium is reached where M(OH) x is still present in the equilibrium mixture: M(OH) x (aq)  M x+ (aq) + x OH - (aq)

9. 9 Common strong acids and bases.

10. 10 Why do we need to improve on Arrhenius theory? The Arrhenius theory has a drawback! Certain compounds that DO NOT contain hydroxide can still increase the hydroxide concentration when placed in water. Arrhenius theory does not explain this!

11. 11 The Brønsted-Lowry Theory The Brønsted-Lowry Theory: an acid Brønsted-Lowry Theory: an acid is any substance that donates protons (H + ) while a base is any substance that can accept protons. This means that Brønsted-Lowry acid-base reactions are proton transfer reactions.

12. 12 Proton transfer reactions Pairs of compounds are related to each other through Brønsted-Lowry acid-base reactions. These are conjugate acid-base pairs.

13. 13 Proton transfer reactions Generally, an acid HA has a conjugate base A - (an H + has transferred away from the acid). Conversely, a base B has a conjugate acid BH + (an H + has transferred toward the base).

14. 14 Water in BL acid-base reactions When a Brønsted-Lowry acid is placed in water, it donates a proton to the water (which acts as a base) and establishes an acid-base equilibrium.

15. 15 Water in BL acid-base reactions In the reverse reaction of the equilibrium, the acid H 3 O + donates a proton to the base A - to give back water and HA.

16. 16 Water in BL acid-base reactions When a Brønsted-Lowry base is placed in water, it accepts a proton from water (which acts as an acid) and establishes an acid-base equilibrium.

17. 17 Water in BL acid-base reactions In the reverse reaction of the equilibrium, the acid BH + donates a proton to the base OH - to give back water and B.

18. 18 Brønsted-Lowry Bases To accept a proton (to act as a B-L base) requires a molecule to have an unshared pair of electrons which can then be used to create a bond to the H +. All Brønsted-Lowry bases have at least one lone pair of electrons.

19. 19 Brønsted-Lowry Bases In the previous reactions we’ve seen NH 3 has a lone pair of electrons and can act as a B-L base. Also, water has two lone pairs of electrons, and can act as a B-L base.

20. 20 Amphiprotic substances Some substances, like water, have protons that can be donated (BL acid), and lone pairs of electrons that can accept protons (BL base). This is why it can act like an acid OR a base DEPENDING on the other species present. Such substances are said to be amphiprotic.

21. 21 Problem Write a balanced equation for the dissociation of each of the following Brønsted-Lowry acids in water: a) H 2 SO 4 b) HSO 4 - c) H 3 O + d) NH 4 +

22. 22 Problem What is the conjugate acid of each of the following Brønsted- Lowry bases? a) HCO 3 - b) CO 3 2- c) OH - d) H 2 PO 4 -

23. 23 Why do we need to improve on Brønsted-Lowry theory? There are many reactions that behave VERY MUCH LIKE proton transfer reactions that DO NOT involve protons!

24. 24 Lewis Acids and Bases A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. These definitions are more general than the Brønsted-Lowry definitions because protons DO NOT need to be involved in Lewis acid-base reactions.

25. 25 In general LA + :LB  LA-LB BF 3 is a Lewis acid where the B atom can accept an electron pair. The N of the NH 3 has a lone pair that can be donated, making NH 3 a Lewis base.

26. 26 Metal ions as Lewis acids Many metal ions have the ability to act as Lewis acids. The ions are willing to accept electron pairs from LIGANDS (which act as Lewis bases) because this often stabilizes the ion in solution. The result is often called a complex ion.

27. 27 Problem.

28. 28 Comparing theories Since Arrhenius acids must contain protons, then ALL Arrhenius acids ARE ALSO Brønsted-Lowry acids. We’ve already seen that NOT ALL Brønsted-Lowry bases are Arrhenius bases.

29. 29 Comparing theories There are Lewis acids (like metal ions) that ARE NOT Brønsted-Lowry acids. ALL Brønsted-Lowry bases must all have at least one lone pair of electrons, so ALL Brønsted-Lowry bases MUST ALSO BE Lewis bases

30. 30 BL acid and base strength Brønsted-Lowry acid-base equilibria are competitions! The equilibrium is the result of a tug-of war between the two bases in the system as they fight for protons given away by the two acids.

31. 31 Acid Strength and Base Strength The acid that is “better at donating protons” OR the base that is “better at accepting protons” will be found in lesser amounts at equilibrium compared to the other acid (or base).

32. 32 Strong BL acids in water A strong acid (HA) is one that almost completely dissociates in water (which acts as a base). The conjugate base A - will be a very weak base.

33. 33 Strong BL acids in water At equilibrium, there will be very little to no HA present in the system, and the concentration of A - will essentially be the same as the initial concentration of HA.

34. 34 Weak BL acids in water A weak acid (HA) is one that partially dissociates in water (which acts as a base). The acid is not as good at donating protons to the water. The conjugate base (A - ) will be a weak base. Overall

35. 35 Weak BL acids in water At equilibrium, there will be some A - and H 3 O + present in the system. However, the concentration of HA will still be significant at equilibrium.

36. 36.

37. 37. Notice that the strongest acids have the weakest conjugate bases, and the strongest bases have the weakest conjugate acids!

38. 38 Hydrated Protons and Hydronium Ions The ultimate proton-donor is a proton itself! In water there is no such thing as H +. Often more than one water molecule will crowd around the proton to give hydrates with the formula H(H 2 O) n + where n is 1 to 4.

39. 39 Hydrated Protons and Hydronium Ions

40. 40 Dissociation of Water It is possible for one water molecule to act as an acid while another water molecule acts as a base at the same time. This leads to the self-ionization of water equilibrium: H 2 O (l) + H 2 O (l)  H 3 O + (aq) + OH - (aq) The equilibrium constant for this reaction is called the ion-product constant for water, K w. K w = [H 3 O + ][OH - ]

41. 41 At 25 °C, K w = 1.0 x 10 -14 so [H 3 O + ] = [OH - ] = 1.0 x 10 -7 mol/L Relatively few water molecules are dissociated at equilibrium at room temperature! We will always assume that [H 3 O + ] [OH - ] = 1.0 x 10 -14 at 25 °C.

42. 42 Acidic [H 3 O + ] > 1.0 x 10 -7 M or [OH - ] < 1.0 x 10 -7 M Basic [OH - ] > 1.0 x 10 -7 M or [H 3 O + ] < 1.0 x 10 -7 M Neutral [H 3 O + ] = [OH - ] = 1.0 x 10 -7 M At 25  C

43. 43 We also find, since [H 3 O + ] [OH - ] = 1.0 x 10 -14 = K w then [H 3 O + ] = 1.0 x 10 -14 / [OH - ] and [OH - ] = 1.0 x 10 -14 / [H 3 O + ] At 25  C

44. 44. At 25  C

45. 45 Problem The concentration of OH - in a sample of seawater is 5.0 x 10 -6 mol/L. Calculate the concentration of H 3 O + ions, and classify the solution as acidic, neutral, or basic.

46. 46 Problem At 50 °C the value of K w is 5.5 x 10 -14. What are the [H 3 O + ] and [OH - ] in a neutral solution at 50 °C?

47. 47 The pH Scale [H 3 O + ] in water can range from very small (strongly basic) to very large (strongly acidic) it is sometimes easier to use a negative logarithmic (power of 10) scale to express [H 3 O + ] with a term we call the pH of a solution. pH = - log [H 3 O + ]

48. 48 pH and acidity pH > 7 is basic pH < 7 is acidic pH = 7 is neutral At 25  C

49. 49 pOH and acidity pOH < 7 is basic pOH > 7 is acidic pOH = 7 is neutral At 25  C pOH = - log [OH - ] Or [OH - ] = 10 -pOH

50. 50 pH scale

51. 51. K w = 1.0 x 10 -14 = [H 3 O + ] [OH - ] pK w = - log (1.0 x 10 -14 ) = 14.00 (2 sigfigs! The 14 is not significant!) pK w = - log ([H 3 O + ] [OH - ]) = (- log [H 3 O + ]) + (- log [OH - ]) = pH + pOH so 14.00 = pH + pOH (at 25  C)!

52. 52 Problem Calculate the pH of each of the following solutions: a) A sample of seawater that has an OH - concentration of 1.58 x 10 -6 mol/L b) A sample of acid rain that has an H 3 O + concentration of 6.0 x 10 -5 mol/L

53. 53 Problem Calculate [H 3 O + ] and [OH - ] in each of the following solutions: a) Human blood (pH 7.40) b) A cola beverage (pH 2.8)

54. 54 Measuring pH We often measure the pH of a solution with a chemical acid-base indicator. Indicators are B-L acids (symbolized HIn) where the acid form has a different colour than the conjugate base form (In - ) HIn (aq) + H 2 O (l)  H 3 O + (aq) + In - (aq) colour A colour B

55. 55. Indicators tend to change colour only in small pH ranges of about 2 units.

56. 56 Measuring pH To make a universal indicator that covers the pH range from about 1 to 12, a mixture of several different indicators with different pH ranges is used.

57. 57. pH 1 2 3 4 5 6 7 8 9 10 11 Methyl violet Phenolphthalein Bromothymol blue Bromocresol green Universal indicator Methyl orange

58. 58 Universal indicator