1 Chapter 13 Acids and Bases

2 Outline 1. Brønsted-Lowry acid-base model 2. The ion product of water3. pH and pOH 4. Weak acids and their equilibrium constants 5. Weak bases and their equilibrium constants 6. Acid-base properties of salt solutions 7. Extending the concept of acids and bases: the Lewis model

3 Review from Chapter 4 The Arrhenius definition of acid and baseAcids produce H+ in water Bases produce OH- in water H+ from acids combines with OH- from bases to produce water in a reaction called a neutralization

4 Brønsted-Lowry Acid-Base ModelJohannes Brønsted ( ) Thomas Lowry ( ) Brønsted-Lowry model focuses on the reaction that takes place between acid and base, rather than on the independent nature of the acid or base, as the Arrhenius model does Acids donate H+ to bases Bases accept H+ from acids

5 The Nature of H+ The H+ ion is the medium of exchange in a Brønsted-Lowry reaction H+ can also be called a proton Acid-base reactions involve proton exchange

6 Conjugate Pairs The species that forms when a proton is removed from an acid is called the conjugate base of the acid If the acid is HB, the conjugate base is B- The only difference between the members of a conjugate acid-base pair is the position of the proton A species that can either accept or donate a proton is called amphiprotic Consider water: OH-  H2O  H3O+ Remove H Add H+

7 Examples of Conjugate Acid-Base PairsConjugate Base HF F- HSO4- SO42- NH4+ NH3

8 The Hydronium Ion Another way to write the H+ ion is as H3O+H3O+ is the hydronium ion H+ exists in water as hydronium ion, since H+ itself would not be stable in water Depending on the reason for writing the reaction, either H+ or H3O+ can be used, and interchangeably The only difference is the inclusion or exclusion of the H2O molecule

9 Example 13.1

10 Example 13.1, (Cont’d)

11 The Ion Product of WaterWater can react with itself in a reaction called autoionization Water can react with itself in an acid-base reaction: H2O + H2O ⇌ H3O+ (aq) + OH- (aq) An alternate way to write the reaction is: H2O ⇌ H+ (aq) + OH- (aq)

12 Equilibrium and the Auto-Ionization of WaterH2O ⇌ H+ (aq) + OH- (aq) Recall that concentrations can be used to write equilibrium constant expressions K for this reaction is [H+][OH-] This K is called the ion product constant of water, Kw Kw = [H+][OH-] At 25 ºC, Kw = 1.0 X 10-14

13 Concentrations of H+ and OH- in pure waterFor water, [H+][OH-] = 1.0 X 10-14 In pure water, the two concentrations are equal: [H+] = 1.0 X 10-7 M [OH-] = 1.0 X 10-7 M Since one concentration must rise as the other falls, If [H+] > 1.0 X 10-7 M, then [OH-] < 1.0 x 10-7 M and the solution is acidic If [OH-] > 1.0 X 10-7 M, then [H+] < 1.0 x 10-7 M and the solution is basic

14 Figure 13.1

15 pH Defining equation for pHThe higher the pH, the less acidic the solution The lower the pH, the more acidic the solution

16 Figure 13.2

17 pOH Defining equation for pHThe higher the pOH, the more basic the solution The lower the pOH, the more acidic the solution

18 Relationship between pH and pOH

19 Table 13.1: pH of Some Common Materials

20 Example 13.2

21 Example 13.2, (Cont’d)

22 Example 13.2, (Cont’d)

23 pH and Blood From the previous example, it is seen that the [H+] in blood is very small, about 4.0 X 10-8 M Small changes in [H+] can have dramatic physiological effects Many biological reactions depend on [H+] An increase in [H+] from 4.0 X 10-8 to 4.0 X 10-7 can increase the reaction rate by a power of 10 Small increases in [H+] can lead to acidosis; small decreases in [H+] can lead to alkalosis Effective control of many physiological reactions depends on pH control

24 pH of Strong Acids Recall from Chapter 4 that some acids are strongHCl, HBr, HI, HClO4, HNO3, H2SO4 These completely ionize in water [H+] is equal to the [H+] of the acid A 0.10 M solution of HCl has [H+] = 0.10 M, so the pH of the solution is 1.00

25 Figure 13.3

26 pH of Strong Bases Recall as well that some bases are strong:LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 These bases ionize completely to OH- pOH is dependent on the concentration of the strong base For an 0.10 M solution of NaOH, [Na+] = [OH-] = 0.10 M pOH = 1.00 pH = 13.00

27 Example 13.3

28 Example 13.3, (Cont’d)

29 Example 13.3, (Cont’d)

30 Measuring pH pH can be measured with a pH meterTranslates [H+] into an electrical signal Signal is shown on an analog or digital meter calibrated in pH units

31 Figure 13.4: pH of Carbonated Soda

32 pH Indicators Universal indicatorMixture of substances that change color depending on the concentration of H+ Less accurate than pH meter Depending on the indicator used, can display pH over a narrow or wide range of [H+] Some plants can act as pH indicators Color of some flowers in plants is dependent on the pH of the soil in which the plant is grown

33 Figure 13.5: Universal Indicator

34 Figure 13.6

35 Weak Acids and their Equilibrium ConstantsWeak acids ionize only partially Prototype reaction HB (aq) + H2O ⇌ H3O+ (aq) + B- (aq) Two types of species that behave as weak acids 1. Molecules with an ionizable hydrogen atom HNO2 (aq) + H2O ⇌ H3O+ (aq) + NO2- (aq) 2. Cations NH4+ (aq) + H2O ⇌ H3O+ (aq) + NH3 (aq)

36 Metal Cations as Acids Many metal cations act as weak acids in water solution as well Zn(H2O)42+ (aq) + H2O ⇌ H3O+ (aq) + Zn(H2O)3(OH)+ (aq) The bond that forms between the oxygen and the metal ion weakens the O-H bond H+ is more easily ionized as a result of the weakened bond

37 Figure 13.7

38 Equilibrium Constants for Weak AcidsHB (aq) + H2O ⇌ H3O+ (aq) + B- (aq) Ka is the acid equilibrium constant Simplifying the above to HB (aq)  H+ (aq) + B- (aq) Ka values are related to the weak acid strength The smaller Ka is, the weaker the acid is

39 Table 13.2

40 pKa pKa = -logKa The smaller pKa is, the stronger the acidpKa follows the trend for pH

41 Example 13.4

42 Example 13.4, (Cont’d)

43 Example 13.5

44 Example 13.5, (Cont’d)

45 Percent Ionization The percent ionization of a weak acid is defined asFor the calculation in example 13.5, the percent ionization is about 12 % Note that the percent ionization depends on the molarity of the weak acid

46 Figure 13.8 - % Ionization and Concentration

47 Example 13.6

48 Calculating [H+] in a Water Solution of a Weak AcidWe can use the process for calculating equilibrium pressure for gaseous reactions that we looked at in Chapter 12 to calculate the equilibrium concentration of [H+] for a weak acid The relationship between [HB], [H+] and [B-] is given in the equilibrium expression itself

49 Example 13.7

50 Example 13.7, (Cont’d)

51 Example 13.7, (Cont’d)

52 Algebra Review – Quadratic EquationsRecall that for a quadratic equation in the form The roots are

53 Approximations Used in CalculationsThe value of Ka is usually known no more accurately than about ±5% When solving for the unknowns used to work the equilibrium problem, for the expression Where a is the initial concentration of weak acid, you can neglect x in the denominator if doing so does not introduce an error of more than 5%, i.e.,

54 Approximations and Percent IonizationWhen Multiplying by 100% will give the percent ionization: If the percent ionization is 5% or less, you may make the approximation. If the percent ionization is greater than 5%, the quadratic formula or the successive approximation method is required

55 Example 13.8

56 Example 13.8, (Cont’d)

57 Example 13.8, (Cont’d)

58 Polyprotic Weak Acids Acids containing more than one ionizable hydrogen are called polyprotic The anion formed in one step produces another H+ in a successive ionization step The equilibrium constant becomes smaller with each successive step

59 Triprotic Acid Phosphoric acid H3PO4 (aq) ⇌ H+ (aq) + H2PO4- (aq) Ka1HPO42- (aq) ⇌ H+ (aq) + PO43- (aq) Ka3 Ka1 > Ka2 > Ka3 With each successive step, the acid becomes progressively weaker

60 Table 13.3

61 Example 13.9

62 Example 13.9, (Cont’d)

63 Weak Bases and their Equilibrium ExpressionsTypes of weak bases Molecules Ammonia, NH3, and amines NH3 (aq) + H2O ⇌ NH4+ (aq) + OH- (aq) Anions Anions derived from weak acids are weak bases I- (aq) + H2O ⇌ HI (aq) + OH- (aq)

64 Example 13.10

65 Weak Base Equilibrium ConstantNH3 (aq) + H2O ⇌ NH4+ (aq) + OH- (aq) The base equilibrium constant, Kb is For a generic weak base where B (aq) + H2O  HB+ (aq) + OH- (aq)

66 Kb by the Numbers As Kb becomes larger, base strength increasesAs with acids and Ka, we can define a pKb: pKb = -logKb As pKb becomes smaller, base strength increases

67 Calculation of [OH-] in a Weak Base SolutionThe process of calculating the [OH-] in a weak base solution is the same as the process for calculating [H+] in a weak acid solution

68 Example 13.11

69 Example 13.11, (Cont’d)

70 Example 13.11, (Cont’d)

71 Example 13.11, (Cont’d)

72 Relation between Ka and KbConsider the relation between a conjugate acid-base pair HB (aq) ⇌ H+ (aq) + B- (aq) KI= Ka of HB B- (aq) + H2O ⇌ HB (aq) + OH- (aq) KII = Kb of B- These add to H2O ⇌ H+ (aq) + OH- (aq) KIII = Kw Since KIKII=KIII, KaKb = Kw = 1.0 X 10-14 for a conjugate acid base pair only In log form, pKa + pKb = pKw = 14.00

73 Notes on Acid-Base StrengthKa and Kb are inversely related The larger Ka is, the smaller Kb is Features Brønsted-Lowry acids Strong acids Weak acids Acids weaker than water (conjugates of strong bases) Brønsted-Lowry bases Strong bases Weak bases Bases weaker than water (conjugates of strong acids)

74 Table 13.4

75 Hydride ion Reaction of water with CaH2H- is the conjugate base of H2, a very weak acid As a result, H- is an extremely strong base

76 Acid-Base Properties of Solutions of SaltsA salt is an ionic solid containing a cation other than H+ and an anion other than OH- or O2- We can predict whether a salt will be acidic, basic or neutral by 1. Deciding what effect the cation has on water Is it acidic or is it neutral? 2. Deciding what effect the anion has on water Is it basic or is it neutral? 3. Combining the two effects to decide the behavior of the salt in water

77 Cations Weak acid or spectator ion? Most cations are acidicThese will change the pH by more than 0.5 pH units in a 0.1 M solution Exceptions – these are spectators Alkali metal cations Heavier alkaline earth cations (Ca2+, Sr2+, Ba2+)

78 Anions Weak base or spectator ion? Many anions are weak basesThese will change the pH by more than 0.5 pH units at 0.1 M Exceptions – these are spectators Anions of very strong acids: Cl-, Br-, I-, NO3-, ClO4-

79 Sodium Chloride Solution - Neutral

80 Table 13.5

81 Salts: Acidic, Basic or NeutralBy comparing the Ka of an acidic cation with the Kb of a basic anion, the salt of both can be classified as acidic, basic or neutral If Ka > Kb, the salt is acidic NH4F, Ka = 5.6 X 10-10; Kb = 1.4 X 10-11 If Kb > Ka, the salt is basic NH4ClO, Ka = 5.6 X 10-10; Kb = 3.6 X 10-7

82 Amphiprotic Anions HCO3- Ka = 4.7 X 10-11 Kb = 2.3 X 10-8Because Kb > Ka, a solution of NaHCO3 will be basic

83 Figure 13.9: Flowchart for Acid-Base Properties of Salts

84 Example 13.12

85 Example 13.12, (Cont’d)

86 Extending the Concept of Acids and BasesThe Lewis Model A Lewis base donates a pair of electrons A Lewis acid accepts a pair of electrons

87 Lewis Bases The concept of a Lewis base does not structurally differ from that of a Brønsted-Lowry base For a species to accept a proton, it must contain an atom that possesses a lone pair Lewis bases are also Brønsted-Lowry bases

88 Lewis Acids The concept of a Lewis acid extends the acid-base modelThe Lewis model greatly expands the number of species considered to be acids For example, metal cations are not Brønsted-Lowry acids but are Lewis acids

89 Table 13.6

90 Key Concepts 1. Classify a substance as a Brønsted-Lowry acid or base and write the net ionic equation to support the classification 2. Given [H+], [OH-], pH or pOH, calculate the three other quantities 3. Given the pH and original concentration of a weak acid, calculate Ka 4. Given the Ka and original concentration of a weak acid, calculate [H+] 5. Given the Kb and original concentration of a weak base, calculate [OH-]

91 Key Concepts, (Cont’d) 6. Given Ka for a weak acid, calculate Kb for its conjugate base (or vice-versa). Predict whether a salt will be acidic, basic or neutral. Understand the similarities and differences between Lewis and Brønsted-Lowry acids and bases