Equilibria of Acids, Bases, and Salts

1. Equilibria of Acids, Bases, and Salts

2. Acids and Bases: Theory Arrhenius theory of acids Arrhenius definition of an acid: any compound that contains hydrogen and produces H+ (H3O+ when reacts with water) ions when dissolved in water. A strong acid is a water-soluble compound that completely dissociates to give H3O+ ions. A weak acid is a water-soluble compound that dissociates only partially, producing few H3O+ ions.

3. Acids and Bases: Theory Arrhenius theory of acids A strong acid is a water-soluble compound that completely dissociates to give H3O+ ions. A weak acid is a water-soluble compound that dissociates only partially, producing few H3O+ ions.

4. Rafa Muñoa Lizardi Institutua Zarautz Acids and Bases: Theory Arrhenius theory of bases Arrhenius definition of a base: any compound that contains a metal and hydroxide (OH-) group producesOH- (hydroxide) ions when dissolved in water. All hydroxides are strong bases because their dissociation reaction go essentially to completion.

5. Acids and Bases: Theory Bronsted-Lowry theory of acids and bases They proposed that acids and bases can be defined in terms of their ability to transfer protons. An acid is a substance (molecule or ion) that can transfer protons to another substance. A base is a substance than can accept a proton.

6. Acids and Bases: Theory Bronsted-Lowry theory of acids and bases An acid and a base always work together to transfer a proton. In other words, a substance can function as an acid only if another substance simultaneously behaves as a base.

7. Acids and Bases: Theory Bronsted-Lowry theory of acids and bases Some substances can act as an acid in one reaction and as a base in another. For example, H2O is a Bronsted-Lowry base in its reaction with HCl and a Bronsted-Lowry acid in its reaction with NH3. Those substances are called amphoteric.

8. Acids and Bases: Theory Bronsted-Lowry theory of acids and bases To be a Bronsted-Lowry acid, a molecule or ion must have a hydrogen atom that it can lose as H+ ion. To be a Bronsted-Lowry base, a molecule or ion must have a nonbonding pair of electrons that it can use to bind the H+ ion.

9. Acids and Bases: Theory Conjugate acid-base pairs In any acid-base equilibrium both the forward reaction (to the right) and reverse reaction (to the left) involve proton transfers. In the forward reaction HX donates a proton to H2O. Therefore, HX is the Bronsted-Lowry acid, and H2O is the Bronsted-Lowry base.

10. Acids and Bases: Theory Conjugate acid-base pairs In the reverse reaction, the H3O+ ion donates a proton to the X- ion; H3O+ is the acid and X- is the base An acid and a base such as HX and X- that differ only in one proton are called conjugated acid-base pair. Every acid has a conjugate base, formed by the removal of a proton from the acid.

11. Acids and Bases: Theory Exercises 1. Write the conjugate base of each of the following acids: HClO4, H2S, HCO3- 2. Write the conjugate acid of each of the following bases: CN-, SO4-2, H2O, HCO3- 3. The hydrogen sulfite ion, HSO3-, is amphoteric. Write its equations with water, acting as both acid and base

12. Acids and Bases: Theory Relative strengths of acids and bases The relationship between the strengths of acids and their conjugate bases is inverse: 1. the strong acids are those that completely transfer their protons to water, leaving no undissociated molecules in solution. Their conjugate bases have a negligible tendency to combine with a proton in aqueous solution.

13. Acids and Bases: Theory Relative strengths of acids and bases 2. the weak acids are those that only partially dissociate in aqueous solution and therefore exist in the solution as a mixture of acid molecules and ions. Their conjugate bases are weak bases, showing a slight ability to remove protons from water.

14. Acids and Bases: Theory Lewis acids and bases Lewis proposed a definition of acid and base that emphasizes the shared electron pair. Lewis acid is defined as an electron-pair acceptor, and a Lewis base is an electron-pair donor.

15. Ionization Constant of a Weak Acid • The term Kais called theacid ionization constant. • The acid ionization constant, Ka , is constant for a specified temperature but has a new value for each new temperature. • Acetic acid, CH3COOH, is a weak acid and most of the CH3COOH molecules remain unionized. • An acetic solution contains CH3COOH molecules, H3O+ ions, and acetate ions, CH3COO.

16. The equilibrium equation for the ionization of acetic acid is • The equation for Ka is • Because water is the solvent, one can assume that the molar concentration of H2O molecules remains constant. The concentration of water is not included in the equilibrium expression.

17. Ionization data and constants for some dilute acetic acid solutions at 25°C are given below. • The numerical value of Kais almost identical for each solution molarity shown.

18. At constant temperature, an increase in the concentration of CH3COO− ions through the addition of sodium acetate, NaCH3COO, disturbs the equilibrium. • This disturbance causes a decrease in [H3O+] and an increase in [CH3COOH]. • The equilibrium is reestablished with the same value of Ka. But there is a higher concentration of nonionized acetic acid molecules and a lower concentration of H3O+ ions. • Changes in the hydronium ion concentration affect pH.

19. Buffers Buffered solutions resist changes in pH. • Buffered solutions contains both a weak acid and a salt of the weak acid • example: CH3COOH and NaCH3COO– • Buffered solution can react with either an acid or a base. When small amounts of acids or bases are added, the pH of the solution remains nearly constant.

21. Buffered Vs. Nonbuffered Solutions

22. If a small amount of acid is added to the acetic acid–sodium acetate solution, acetate ions react with most of the added hydronium ions to form nonionized acetic acid molecules. • The hydronium ion concentration and the pH of the solution remain practically unchanged.

23. If a small amount of a base is added, the OH− ions of the base react with and remove hydronium ions to form nonionized water molecules. Acetic acid molecules then ionize and mostly replace the hydronium ions neutralized by the added OH− ions. • The hydronium ion concentration and the pH of the solution remain practically unchanged.

24. A solution of a weak base containing a salt of the base also behaves as a buffered solution. • Buffer action has many important applications in chemistry and physiology. • Human blood is naturally buffered to maintain a pH of between 7.3 and 7.5.

25. Ionization Constant of Water • The self-ionization of water is an equilibrium reaction. • Equilibrium is established with a very low concentration of H3O+ and OH− ions. • Kw=[H3O+][OH–] = 1.0  10-14

26. Hydrolysis of Salt • Salts are formed during the neutralization reaction between a Brønsted acid and a Brønsted base. • When a salt dissolves in water, it produces • positive ions (cations) of the base from which it was formed • negative ions (anions) of the acid from which it was formed • If the ions formed are from weak acids or bases, they react chemically with the water molecules, and the pH of the solution will have a value other than 7.

27. A reaction between water molecules and ions of a dissolved salt ishydrolysis. • If the anions react with water, the process is anion hydrolysis and results in a more basic solution. • If the cations react with water molecules, the process is cation hydrolysis and results in a more acidic solution.

28. Anion Hydrolysis • In the Brønsted sense, the anion of the salt is the conjugate base of the acid from which it was formed. • It is also a proton acceptor. • If the acid is weak, its conjugate base (the anion) will be strong enough to remove protons from some of the water molecules. • An equilibrium is established in which the net effect of the anion hydrolysis is an increase in the hydroxide ion concentration, [OH−], of the solution.

29. The equilibrium equation for a typical weak acid in water, HA, is • The general equilibrium equation is • The hydrolysis reaction between water and the anion, A−, that is produced by the dissociation of the weak acid, HA, is • The extent of OH− ion formation and the position of the equilibrium depends on the relative strength of the anion, A−. • The lower the Kavalue of HA, the weaker the acid, HA, the stronger its conjugate base, A−

30. Aqueous solutions of sodium carbonate are strongly basic. • The carbonate ions react as a Brønsted base. • The OH− ion concentration increases until equilibrium is established.

31. Cation Hydrolysis • In the Brønsted sense, the cation of the salt is the conjugate acid of the base from which it was formed. • It is also a proton donor. • If the base is weak, the cation is an acid strong enough to donate a proton to a water molecule to form H3O+ ions. • An equilibrium is established in which the net effect of the cation hydrolysis is an increase in the [H3O+] of the solution.

32. The equilibrium equation for a typical weak base, B, is • Kb, the base dissociation constant is • The hydrolysis reaction between water and the cation, BH+, produced by the dissociation of the weak base, B, is • The extent of H3O+ ion formation and the position of the equilibrium depend on the relative strength of the cation, BH+. • The lower the Kbvalue of B, the weaker the base, the stronger its conjugate acid will be.

33. Hydrolysis in Acid-Base Reactions • Hydrolysis can help explain why the end point of a neutralization reaction can occur at a pH other than 7. • Salts can be placed in four general categories, depending on their hydrolysis properties: • strong acid–strong base • strong acid–weak base • weak acid–strong base • weak acid–weak base

34. Salts of strong acids and strong bases produce neutral solutions. • Neither the cation of a strong base nor the anion of a strong acid hydrolyzes appreciably in aqueous solutions. • The aqueous solutions of salts formed from reactions between weak acids and strong bases are basic. • Anions of the dissolved salt are hydrolyzed by the water molecules, and the hydroxide-ion concentration increases. This raises the pH of the solution.

35. Neutralization Curve for a Weak Acid and a Strong Base

36. Salts of strong acids and weak bases produce acidic aqueous solutions. • Cations of the dissolved salt are hydrolyzed in the water solvent, and hydronium ion concentration increases. The pH of the solution is lowered. • Salts of weak acids and weak bases can produce either acidic, neutral, or basic aqueous solutions, depending on the salt dissolved. • Both ions of the dissolved salt are hydrolyzed extensively.

37. Neutralization Curve for a Strong Acid and a Weak Base

38. Equilibrium and pH Calculations

39. References • Rafa Muñoa • W Sautter • McGraw Hill Ryerson 12 • Nelson, Chemistry 12