1

Acids and Bases

2

Bronsted Acids & Bases

Acid: substance capable of donating a proton (H+).

Base: substance capable of accepting a proton.(definition not dependent on OH-)

3

Conjugate Pairs

Acid & conjugate base pairBase and conjugate acid pair

CH3COOH + H2Oacetic acid water

CH3COO- + H3O+

acetate ion hydronium ion

Acid: H+ donorBase: H+ acceptor

How do you recognize a conjugate pair?

4

Conjugate Pairs

NH3 + H2O NH4+ + OH-

NH3 NH4+

base conjugate acid

H2O OH-

acid conjugate base

5

Bronsted Base

NH3 H2O

Bronsted bases must have an atom with a lone pair of electrons to accept a proton(“coordinate covalent bond”)

H+

6

Try It !!!

Identify the conjugate pairs in this reaction.

CN- + H2O HCN + OH-

7

Autoionization of Water

or

H2O + H2O H3O+ + OH-

H2O H+ + OH-

What are the conjugate pairs?

shorthand

8

Ion Product of Water

or since [H2O] is ~constant:

Kc[H2O] = Kw = [H+][OH-]

K’c =[H+][OH-]

[H2O]

Remember: H+ is shorthand for H3O+

Kw = [H3O+][OH-]

9

Ion Product of Water

In pure water at 25oC:

[H3O+] = [OH-] = 1.0 x 10-7 M

Thus

Kw = [H3O+][OH-] = 1.0 x 10-14

Kw is temperature dependent, but we will use this value for most calculations.

10

Acid vs. Base

Acid: [H+] > [OH-]

Base: [H+] < [OH-]

Neutral: [H+] = [OH-]

11

[H+] and [OH-] Calculations

1.In an ammonia solution, what is [H+] if [OH-] is 0.0025 M?

2.What is [OH-] in an HCl solution with [H+] of 1.3 M?

12

pH: DefinitionpH = -log[H+] (no units)

pH increases as [H+] decreases

Acidic: [H+] > 10-7 M & pH < 7Basic: [H+] < 10-7 M & pH > 7Neutral: [H+] = 10-7 M & pH = 7*

*Neutral [H+] = [OH-]Neutral is pH = 7 only at 25oC.

13

pH of Common FluidsStomach fluid 1.5Vinegar 3.0Orange juice 3.5Water + air 5.5Blood 7.4Milk of magnesia 11.5

pH is a log function. A pH increase of 1.0 means [H+] decrease by factor of 10.

pH is a log Function

14

Compare the pH of a solution with a pH of 3 to a solution with a pH of 7.

15

pOH pX = –log(X)

pOH = -log[OH-]

Kw = [H+][OH-] = 1.0 x 10-14

Taking negative logs:

-log[H+] - log[OH-] = 14.00

pH + pOH = 14.00

16

pH pHun

•If [H+] = 3.2 x 10-4 M, what is the pH?

•If the pH of “acid rain” is 4.82, what is [H+]?

•If a basic solution has [OH-] of 2.9 x 10-4 M, what is pH?

Temp. Dependence of Kw

17

pKw

temp. (oC)

H2O H+ + OH-

Is this reaction endo-or exothermic?

18

Acid Strength

Strong acids: completely ionized in water.

HCl, HBr, HI, HNO3, H2SO4, HClO3

Strong oxyacid if #O - #H > 2This applies to neutral acids, not ions.

H2SO4 + H2O H3O+ + HSO4

-100%

Know!

Strong Acids

19

HCl + H2O H3O+ + Cl-

The anion of a strong acid (Cl-) doesn’t ‘hydrolyze’. Cl- does not react with water to bond with an H+.

20

Weak Acids

Weak acids incompletely ionize in water.

HF, NH4+, organic acids (CH3COOH)

HF + H2O H3O+ + F-

Equilibrium lies far to the left. [HF] >> [F-]

since F- holds onto the H+

21

Strong Bases Completely ionized in water.

NaOH, KOH, Ba(OH)2

KOH K+(aq) + OH-(aq)H2O

Bronsted base

Group 1 & 2 hydroxides. All ionize completely, even though some are only slightly soluble (e.g. Mg(OH)2.)

Strong Bases

22

KOH K+(aq) + OH-(aq)H2O

The cation of a strong base (K+) doesn’t ‘hydrolyze’. K+ does not react with water to bond with an OH-

.Group 1 & 2 ions do not hydrolyze.

23

Weak Bases

Weak electrolytes.

NH3 + H2O NH4+ + OH-

Unlike an acid, a weak base does not ionize. Rather is causes the ionization of water.

Common Weak Bases

24

Ammonia, NH3

Organic amines, e.g. CH3NH2

Pyridines, e.g.

What do these have in common?

25

General Properties(Table 15.2)

HCl + H2O H3O+ + Cl-

1. The stronger the acid

the weaker the conjugate base.

HA + H2O H3O+ + A-

2. H3O+ is strongest acid in water.

Stronger acids yield H3O+.

26

General Properties(Table 15.2)

3.All weak acids have the equilibrium far to the left.

4.OH- is strongest base in water. Stronger bases produce OH-.

O-2 + H2O 2OH-

HF + H2O H3O+ + F-

27

PredictionUse Table 15.2 to predict whether primarily reactants or products will remain when HNO2 is mixed with KCN.

HNO2 + CN- NO2- + HCN

stronger weakeracid acid

28

Prediction

Use Table 15.2 to predict whether Kc is >1 or <1 for the following reaction:

F- + HNO2 HF + NO2-

weaker strongeracid acid

29

Strong AcidCalculate pH of 1.0 x 10-3 M HCl.

HCl H+ + Cl-

Initial: 1.x10-3 0 0Change: –1.x10-3 1.x10-3 1.x10-3

Equilib: 0 1.x10-3 1.x10-3

pH = -log(1.0x10-3) = 3.00

ICE

For strong acids, [H3O+] is same

as starting acid concentration.

One way arrow

30

Strong Base: Try It

What is the pH of a 0.020 M barium hydroxide solution?

31

Weak Acids

HA + H2O H3O+ + A-

or simply:

HA H+ + A-

Acid ionization constant, Ka

Ka = [H+][A-]

[HA]

Two way arrow

32

Acid Ionization Constant

The strength of a weak acid is measured by Ka.

See Table 15.3.

You will need to look up Ka’s!

33

Problems Involving KaWhat is the pH of 0.050 M HF, given Ka = 7.1 x 10-4 ?

or simply:HF H+ + F-

(ignore autoionization of water)

HF + H2O H3O+ + F-

Problems Involving Ka

34

ICE

Initial: 0.050 0 0Change: -x +x +xEqulib.: 0.050-x x x

HF H+ + F-

Ka = x2

(0.050 – x)= 7.1 x 10-4

35

Ka Problem continued

x2

(0.050 – x)= 7.1 x 10-4

There are 2 ways to solve!!!1. quadratic (yuck!)2. successive approximation

36

1. Quadratic

x2

(0.050 – x)= 7.1 x 10-4

x2 + (7.1x10-4)x -3.6x10-5 = 0

x = -b ± b2 – 4ac

2a

a b c

1

37

1. Quadraticx2 + (7.1x10-4)x -3.6x10-5 = 0

You get 2 answers; only one is reasonable.

x = 5.6 x 10-3 M = [H+]

x = -b ± b2 – 4ac

2a

pH = -log(5.6 x 10-3) = 2.25

38

2. Successive Approximation

Ka = x2

(0.050 – x)= 7.1 x 10-4

Assume x is very small compared to 0.050 so that the degree of ionization is small.

This works for Ka ~ 10-4 or less.

HF H+ + F-

39

2. Successive Approximation

x2

0.050 = 7.1 x 10-4

x = 0.0060 M

(ignore ‘x’ here)

Ka = x2

(0.050 – x)= 7.1 x 10-4

40

2. Successive Approximation

Plug this value of ‘x’ into the denominator of the Ka equation, and recalculate another ‘x’.

Repeat until ‘x’ doesn’t change within 2 sig figs.

41

2. Successive Approximation

Put x = 0.0060 M here & calculate another x.

= 7.1 x 10-4Ka = x2

(.050 - x)

x = 0.0056 M

(repeat again and get same value)

42

3. Successive Approximation

x = [H+] = 0.0056 M

pH = -log(0.0056)

= 2.25

Usually 2 cycles will do it!

same as quadratic!

43

Summary (so far)

1. Express Ka in terms of ‘x’, the change in concentration (ICE).

2.Use successive approximation first, then quadratic if needed.

3.Having ‘x’, calculated all else.4. pH depends on both acid strength

(Ka) and acid concentration.

44

Try It !!!

1.Calculate pH of 0.036 M nitrous acid (Ka = 4.5 x 10-4)

2. The pH of 0.10 M methanoicacid is 2.39. What is Ka ?

45

Percent Ionization (PI) Another measure of the

strength of an acid.

HA H+ + A-

=[H+][HA]o

x 100

PI = Ionized acid conc. at eq.Initial acid conc.

X 100

46

For 0.50 M HF

PI = [H+][HF]o

x 100

For 0.050 M HF

PI = 0.00560.050

x 100 = 11%

0.0190.50

x 100 = 3.8%PI =

More dilute greater PI

(See slide 39)

47

Percent Ionization (PI)

[HA]o

PI

0

100 strong acid

weak acid

LeChatelier’s PrincipleHA + H2O H3O

+ + A-

48

Weak Bases (Kb)Treated in the same way.

NH3 + H2O NH4+ + OH-

K = [NH4

+][OH-]

[NH3][H2O]

Name the conjugate pairs.

K[H2O] = Kb = [NH4

+][OH-]

[NH3]

49

Base Ionization Constant, Kb

Problems are solved the same as with acids except calculate [OH-] first. Try it.

What is the pH of 0.40 M NH3solution? Kb = 1.8 x 10-5

See Table 15.4

50

Acid Conjugate Basee.g. acetic acid (HAc) and

sodium acetate (NaAc)

For acetic acid:

HAc + H2O H3O+ + Ac-

Ka =[H3O

+][Ac-][HAc]

51

Acid Conjugate Base

Then: Ac- + H2O HAc + OH-

For sodium acetate:

Nothing happens to Na+ WHY?

First: NaAc Na+ + Ac-H2O

Kb = [HAc][OH-]

[Ac-]

Group 1 & 2 ions do not hydrolyze.

52

Acid Conjugate Base

For a conjugate pair (HA and A-), show that : Ka x Kb = Kwand that the two reactions add to the autoionization of water.

(recall the addition rule for chemical equilibria)

(Also note: pKa + pKb = pKw = 14)

Conjugate Pairs

53

1. Conjugate ion of a weak acid is a weak base.

HF H+ + F-

2. Conjugate ion of a weak base is a weak acid.NH3 + H2O NH4

+ + OH-

3. The stronger the acid/base, the weaker its conjugate.

54

Acid Conjugate Base

Ka Kb = Kw

Try it.

If Ka for acetic acid is 1.8 x 10-5, what is Kb for sodium acetate?

55

Polyprotic Acids

e.g. H2CO3 stepwise ionization

Step 1:

H2CO3 H+ + HCO3-

Ka1 =[H+][HCO3

-][H2CO3]

56

Polyprotic Acids

Step 2: conjugate base of step 1 is the acid of step 2.

HCO3- H+ + CO3

-2

Ka2 =[H+][CO3

-2][HCO3

-]

57

Polyprotic AcidsKa1 is always much greater than Ka2. Why?

See Table 15.5

Ka1 = 4.2E-7Ka2 = 4.8E-11

Overall K

58

What is the product of Ka1 x Ka2?

H2CO3 2H+ + CO3-2

H2CO3 H+ + HCO3-

HCO3- H+ + CO3

-2

Ka1 x Ka2 = 2.0E-17

59

Fun with Oxalic Acid

C2H2O4 is diprotic with

Ka1 = 6.5 x 10-2

Ka2 = 6.1 x 10-5

Calculate conc. of all species in 0.10 M solution of C2H2O4

H-O-C-C-O-H

O O

60

Oxalic AcidStep 1: 1st ionization.

C2H2O4 H+ + C2HO4-

I 0.10 0 0C -x +x +xE 0.10-x x x

Ka1 = x2

0.10 - x= 6.5 x 10-2

61

Oxalic Acid: Step 1

Solve for ‘x’. Approximation method is so far off that you must use quadratic.

x = 0.054 M

X2 + 0.065 x - 0.0065 = 0

62

Oxalic Acid: Step 1 Results

[H+] = 0.054 M

[C2HO4-] = 0.054 M

[C2H2O4] = (0.10 – 0.054) M

= 0.046 M

Now use these as starting values for 2nd ionization.

63

Oxalic Acid: Step 2

I 0.054 0.054 0C -y +y +yE 0.054-y 0.054+y y

C2HO4- H+ + C2O4

-2

2nd ionization

Ka2 = (0.054+y) y

(0.054-y)= 6.1 x 10-5

64

Oxalic Acid: Step 2

y = [C2O4-2] = 6.1 x 10-5

Same value as Ka2 !

65

Oxalic Acid: Final Results

[C2H2O4] = 0.046 M

[C2HO4-] = (0.054 – 6.1 x 10-5)M

= 0.054 M

[H+] = (0.054 + 6.1 x 10-5)M

= 0.054 M

[C2O4-2] = 6.1 x 10-5 M

1st only

1st only

1st only

Ka2

66

Oxalic Acid

Note these general results:

2.Second stage conjugate base concentration is equal to Ka2.

1.If Ka1 >> Ka2, [H+] & [first

conjugate base] are result of 1st stage ionization only.

67

Molecular Structure vs. Acid Strength

How is molecular structure of the acid related to acid strength?

HX H+ + X-

Acid strength is measured by degree of ionization (Ka).

68

Acid StrengthDetermine by two factors:

1. Bond strength: the higher the bond dissociation energy, the weaker the acid.

H—A

2. Bond polarity: the more polar the bond, the stronger the acid.

Why?

d+ d-

69

Hydrohalic Acids

HF<<HCl<HBr<HIacid strength

Bond E

nerg

y

Bond P

ola

rity

HF

HCl

HBr

HI

weakstrong

strongstrong

Win

s!

d+ d-H—A

70

Strength of OxyacidsDraw Lewis dot structures of:

Carbonic acidNitric acidSulfuric acid

71

Oxyacids

General structure Z—O—H

If Z is highly electronegative or in high oxidation state, it will attract bonding electrons and make the O—H bond more polar, resulting in a stronger acid.

d- d+

72

Oxyacids with Z Atom from same Group

HClO3 vs. HBrO3

Cl more electronegative, therefore HClO3 is stronger.

What is oxidation number of Z ?

H—O—Z—O

O

73

Oxyacids with Same Z

HClO4 >HClO3 >HClO2 >HClO

acid strength

As oxidation state of Cl increases, acid strength increases. Remember the rule for strong acids.

Ox.#+7 +5 +3 +1

74

Oxyacids: Summary

General structure Z—O—H

•High electronegativity•High oxidation number

For element Z:

strongeracid

75

Predict !!!Place in order of decreasing strength and state which are strong/weak:

1.HIO, HBrO, HClO

2.HNO3, HNO2

3.H3PO3, H3PO4

76

Salt “Hydrolysis” Reaction of a cation and/or

anion of a salt with water, changing pH.

Which Ions React with H2O to Change pH?

77

Anions from strong acids do not hydrolyze. E.g. Cl- comes from the strong acid HCl.

Cations from strong bases do not hydrolyze. E.g. Na+ comes from the strong base NaOH.

All other ions hydrolyze.

Know strong acids and bases.

78

“Neutral” SaltsGroup 1 and 2 cations (except Be+2) combined with anions of a strong acid do not hydrolyze and are neutral.

NaNO3 Na+(aq) + NO3-(aq)

pH ~ 7NaOH HNO3

(strong base, (strong acidno hydrolysis) no hydrolysis)

79

“Basic” SaltsSalt of weak acid and strong base is basic.

CH3COONa Na+ + CH3COO-H2O

CH3COO- + H2O CH3COOH + OH-Na+ doesn’t hydrolyze, but:

basic

80

“Basic” Salts

CH3COO- + H2O CH3COOH + OH-

Kb = [CH3COOH ][OH-]

[CH3COO-]

CH3COONa acts like a weak base when added to water.

81

Try It!

What is the pH of a 0.15 M solution of CH3COONa?

(Ka of acetic acid is = 1.8E-5)

82

“Acidic” Salts

Salt of strong acid and weak base is acidic.

NH4Cl NH4+(aq) + Cl-(aq)

H2O

Cl- doesn’t hydrolyze, but:

NH4+ + H2O NH3 + H3O

+

acidic

83

“Acidic” Salts

NH4+ + H2O NH3 + H3O

+

or simply: NH4+ NH3 + H+

Ka =[NH3][H

+][NH4

+]

Can look up Ka of NH4+

(or Kb of NH3) to solve problems.

84

“Acidic” SaltsThe most acidic cations are small and highly charged.

(high charge density)

e.g. Al+3, Cr+3, Fe+3, Bi+3, Be+2

Larger, +1 cations do not hydrolyze (K+, Na+).

85

Try It !!Al+3 ion becomes hydrated in water, forming Al(H2O)6

+3, which then hydrolyzes as follows:

What is the pH of 0.020 M AlCl3?

Al(H2O)6+3 Al(OH)(H2O)5

+2 + H+

Ka = 1.3 x 10-5

86

Salts in Which Both Cation & Anion Hydrolyze

The math is involved, so let’s just make some qualitative predictions.

87

Salts in Which Both Cation & Anion Hydrolyze

Kb for anion < Ka for cation, solution will be acidic.

Kb for anion ~ Ka for cation,solution will be ~ neutral.

Kb for anion > Ka for cation, solution will be basic.

88

PredictAre solutions of the following salts acidic, neutral, or basic?

NH4I CaCl2Fe(NO3)3 KCN

LiClO4 NH4CN

Ka~E-10 Kb~E-5

89

HCO3- Acidic or Basic?

HCO3- + H2O H3O

+ + CO3-2

HCO3- + H2O H2CO3 + OH-

Ka = 4.8 x 10-11

Kb = 2.4 x 10-8

Since Kb > Ka, solution is basic.

HSO4-: An Exception!

90

HSO4- is the only common

acidic anion.

Cl-, NO3-, ClO4

- neutral

F-, CO32-, HCO3

-, H2PO4- basic

HSO4- acidic

Remember this guy!

91

Oxides See Table 15.8

basic amphoteric acidic

1 2 13 14 15 16 17

92

Basic (Metallic) Oxides

Na2O + H2O 2NaOH

BaO + H2O Ba(OH)2

Metal oxides are “anhydrous” bases (Bases with the water removed).

93

Acidic (Nonmetal) Oxides

Nonmetal oxides are anhydrous acids.

CO2 + H2O H2CO3

(water exposed to air)

SO3 + H2O H2SO4

(acid rain)

94

Try It !!!

Write the reaction for:

CO2(g) + NaOH(aq) ?

BaO(s) + HNO3(aq) ?

95

Amphoteric Oxides

Al2O3 + 6HCl 2AlCl3 + 3H2O

Al2O3 + 2NaOH + 3H2O

2NaAl(OH)4

Al2O3 is both basic & acidic!

96

Transition Metal Oxides

Some transition metal oxides are acidic if oxidation number is high.

Mn2O7 + H2O 2HMnO4

CrO3 + H2O H2CrO4

(No change in oxidation no.)

+7 +7

+6 +6

97

pH and Solubility

Solubility of “insoluble” basic salts are greatly affected by pH due to common ion effect.

e.g. Mg(OH)2 and BaF2

Since OH- and F- both react with H+, the solubility of Mg(OH)2 and BaF2 are affected by pH.

98

pH and SolubilityMg(OH)2 has Ksp = 1.2 x 10-11

Its pH in water is given by:Mg(OH)2 Mg2+ + 2OH-

Ksp = 1.2 x 10-11 = (Mg+2)(OH-)2

2s = [OH-] = 2.8 x 10-4 M

pH = 10.45

s = 1.4 x 10-4 M

= (s)(2s)2

99

pH and Solubility

Mg(OH)2(s) Mg+2(aq) + 2OH-(aq)

If H+ is added to lower pH below 10.45, the added H+ reacts with OH-, thus lowering the [OH-]. Thus more Mg(OH)2 will dissolve to keep Ksp constant (LeChatelier)

100

Similarly for a basic salt like BaF2:

BaF2(s) Ba+2(s) + 2F-(s)

pH and Solubility

Thus more BaF2 will dissolve at lower pH (LeChat.)

The added H+ to lower pH reacts with F-, forming the weak acid HF.

H+(aq) + F-(aq) HF(aq)

101

101

Try it !!!

Which are more soluble in acidic media than in water?

CuS

AgCl

PbSO4

pH and Solubility