Chapter 17 Principles of Chemical Reactivity: The Chemistry of Acids and Bases

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

2. Acids & Bases: A Review • In Chapter 3, you were introduced to two definitions of acids and bases: the Arrhenius and the Brønsted–Lowry definition. • Arrhenius acid: Any substance that when dissolved in water increases the concentration of hydrogen ions, H+. • Arrhenius base: Any substance that increases the concentration of hydroxide ions, OH, when dissolved in water. • A Brønsted–Lowry acid is a proton (H+) donor. • A Brønsted–Lowry base is a proton acceptor.

3. Strong & Weak Acids/Bases • Generally divide acids and bases into STRONG or WEAK ones. STRONG ACID: HNO3(aq) + H2O(liq)  H3O+(aq) + NO3-(aq) HNO3 is about 100% dissociated in water.

4. Strong & Weak Acids/Bases HNO3, HCl, H2SO4 and HClO4 are classified as strong acids.

5. Strong & Weak Acids/Bases CaO • Strong Base: 100% dissociated in water. NaOH(aq)  Na+(aq) + OH-(aq) Other common strong bases include KOH and Ca(OH)2. CaO (lime) + H2O  Ca(OH)2 (slaked lime)

6. Strong & Weak Acids/Bases • Weak base: less than 100% ionized in water An example of a weak base is ammonia NH3(aq) + H2O(liq)  NH4+(aq) + OH-(aq)

7. Strong & Weak Acids/Bases Weak acids are much less than 100% ionized in water. Example: acetic acid = CH3CO2H

8. The Brønsted–Lowry Concept of Acids & Bases Extended • Proton donors may be molecular compounds, cations or anions.

9. The Brønsted–Lowry Concept of Acids & Bases Extended • Proton acceptors may be molecular compounds, cations or anions.

10. The Brønsted–Lowry Concept of Acids & Bases Extended Using the Brønsted definition, NH3 is a BASE in water and water is itself an ACID Proton acceptor Proton donor

11. The Brønsted–Lowry Concept of Acids & Bases Extended • Acids such as HF, HCl, HNO3, and CH3CO2H (acetic acid) are all capable of donating one proton and so are called monoprotic acids. • Other acids, called polyprotic acids are capable of donating two or more protons.

12. Conjugate Acid–Base Pairs • A conjugate acid–base pair consists of two species that differ from each other by the presence of one hydrogen ion. • Every reaction between a Brønsted acid and a Brønsted base involves two conjugate acid–base pairs

13. Conjugate Acid–Base Pairs

14. Water & the pH Scale Water Autoionization and the Water Ionization Constant, Kw: The water autoionization equilibrium lies far to the left side. In fact, in pure water at 25 °C, only about two out of a billion (109) water molecules are ionized at any instant. Even in pure water, there is a small concentration of ions present at all times. [H3O+] = [OH] = 1.00  107

15. Water & the pH Scale H2O can function as both an ACID and a BASE. In pure water there can be AUTOIONIZATION. Equilibrium constant for autoionization = Kw Kw = [H3O+] [OH-] = 1.00 x 10-14at 25 °C

16. Water & the pH Scale • In a neutral solution, [H3O+] = [OH] • Both are equal to 1.00  10 7 M • In an acidic solution, [H3O+] > [OH] • [H3O+] > 1.00  10 7 M and [OH] < 1.00  10 7 M • In a basic solution, [H3O+] < [OH] • [H3O+] < 1.00  10 7 M and [OH] > 1.00  10 7 M

17. The pH Scale

18. The pH Scale • The pH of a solution is defined as the negative of the base (10) logarithm (log) of the hydronium ion concentration. pH =  log[H3O+] • In a similar way, we can define the pOH of a solution as the negative of the base - 10 logarithm of the hydroxide ion concentration. pOH =  log[OH] pH + pOH = pKw = 14

19. The pH Scale • The concentration of acid, [H3O+] is found by taking the antilog of the solutions pH. • In a similar way, [OH] can be found from:

20. The pH Scale Once [H3O+] is known, [OH] can be found from: And vice versa.

21. Equilibrium Constants for Acids & Bases • In Chapter 3, it was stated that acids and bases can be divided roughly into those that are strong electrolytes (such as HCl, HNO3, and NaOH) and those that are weak electrolytes (such as CH3CO2H and NH3) • In this chapter we will discuss the quantitative aspects of dissociation of weak acids and bases. • The relative strengths of weak acids and bases can be ranked based on the magnitude of individual equilibrium constants.

22. Equilibrium Constants for Acids & Bases • Strong acids and bases almost completely ionize in water (~100%): Kstrong >> 1 (product favored) • Weak acids and bases almost completely ionize in water (<<100%): Kweak << 1 (Reactant favored)

23. Equilibrium Constants for Acids & Bases • The relative strength of an acid or base can also be expressed quantitatively with an equilibrium constant, often called an ionization constant. For the general acid HA, we can write: Conjugate acid Conjugate base

24. Equilibrium Constants for Acids & Bases • The relative strength of an acid or base can also be expressed quantitatively with an equilibrium constant, often called an ionization constant. For the general base B, we can write: Conjugate base Conjugate Acid

25. Ionization Constants for Acids/Bases Increase strength Increase strength Conjugate Bases Acids

26. Equilibrium Constants for Acids & Bases • The strongest acids are at the upper left. They have the largest Ka values. • Ka values become smaller on descending the chart as the acid strength declines. • The strongest bases are at the lower right. They have the largest Kb values. • Kb values become larger on descending the chart as base strength increases.

27. Equilibrium Constants for Acids & Bases • The weaker the acid, the stronger its conjugate base: The smaller the value of Ka, the larger the value of Kb. • Aqueous acids that are stronger than H3O+ are completely ionized. • Their conjugate bases (such as NO3) do not produce meaningful concentrations of OH ions, their Kb values are “very small.” • Similar arguments follow for strong bases and their conjugate acids.

28. Equilibrium Constants for Acids & Bases

29. Equilibrium Constants for Acids & Bases

30. Equilibrium Constants for Acids & Bases Ka Values for Polyprotic Acids In general, each successive dissociation produces a weaker acid.

31. Equilibrium Constants for Acids & Bases Logarithmic Scale of Relative Acid Strength, pKa • Many chemists use a logarithmic scale to report and compare relative acid strengths. pKa =  log(Ka) The lower the pKa, the stronger the acid.

32. Equilibrium Constants for Acids & Bases Relating the Ionization Constants for an Acid and Its Conjugate Base

33. Equilibrium Constants for Acids & Bases Relating the Ionization Constants for an Acid and Its Conjugate Base

34. Equilibrium Constants for Acids & Bases Relating the Ionization Constants for an Acid and Its Conjugate Base

35. Equilibrium Constants for Acids & Bases Relating the Ionization Constants for an Acid and Its Conjugate Base

36. Equilibrium Constants for Acids & Bases Relating the Ionization Constants for an Acid and Its Conjugate Base When adding equilibria, multiply the K values.

37. Acid–Base Properties of Salts

38. Acid–Base Properties of Salts Anions that are conjugate bases of strong acids (for examples, Cl or NO3. These species are such weak bases that they have no effect on solution pH.

39. Acid–Base Properties of Salts Anions such as CO3 that are the conjugate bases of weak acids will raise the pH of a solution. Hydroxide ions are produced via “Hydrolysis”.

40. Acid–Base Properties of Salts Anions such as CO3 that are the conjugate bases of weak acids will raise the pH of a solution. Hydroxide ions are produced via “Hydrolysis”. A partially deprotonated anion (such as HCO3) is amphiprotic. Its behavior will depend on the other species in the reaction.

41. Acid–Base Properties of Salts Alkali metal and alkaline earth cations have no measurable effect on solution pH. Since these cations are conjugate acids of strong bases, hydrolysis does not occur.

42. Acid–Base Properties of Salts Basic cations are conjugate bases of acidic cations such as [Al(H2O)6]3+. Acidic cations fall into two categories: (a) metal cations with 2+ and 3+ charges and (b) ammonium ions (and their organic derivatives). All metal cations are hydrated in water, forming ions such as [M(H2O)6]n+.

43. Acid–Base Properties of Salts: Practice

44. Acid–Base Properties of Salts: Practice

45. Predicting the Direction of Acid–Base Reactions • According to the Brønsted–Lowry theory, all acid–base reactions can be written as equilibria involving the acid and base and their conjugates. • All proton transfer reactions proceed from the stronger acid and base to the weaker acid and base.

46. Predicting the Direction of Acid–Base Reactions • When a weak acid is in solution, the products are a stronger conjugate acid and base. Therefore equilibrium lies to the left. • All proton transfer reactions proceed from the stronger acid and base to the weaker acid and base.

47. Will the following acid/base reaction occur spontaneously? Predicting the Direction of Acid–Base Reactions

48. Will the following acid/base reaction occur spontaneously? Predicting the Direction of Acid–Base Reactions Kb = 5.6  1010 Kb = 1.3  1012 Ka = 1.8  105 Ka = 7.5  105

49. Will the following acid/base reaction occur spontaneously? Equilibrium lies to the right since all proton transfer reactions proceed from the stronger acid and base to the weaker acid and base. Predicting the Direction of Acid–Base Reactions Kb = 5.6  1010 Kb = 1.3  1012 Ka = 1.8  105 Ka = 7.5  105 Stronger Acid + Stronger Base Weaker Base + Weaker Acid

50. Strong acid (HCl) + Strong base (NaOH) Net ionic equation Mixing equal molar quantities of a strong acid and strong base produces a neutral solution. Types Acids–Base Reactions