1. Chemistry 102(01) spring 2009 Instructor: Dr. Upali Siriwardane
Office: CTH 311 Phone
Office Hours: M,W, 8:00-9:00 & 11:00-12:00 a.m.; Tu,Th,F 9: :00 a.m.  
Test Dates: March 25, April 26, and May 18; Comprehensive Final Exam: May 20, :30-10:45 am, CTH 328.
March 30, 2009 (Test 1): Chapter 13
April 27, (Test 2): Chapters 14 & 15
May 18, (Test 3): Chapters 16, 17 & 18
Comprehensive Final Exam: May 20,2009 :Chapters 13, 14, 15, 16, 17 and 18

2. Chapter 16. Acids and Bases
The Brønsted-Lowry Concept of Acids and Bases
Types of acids/bases:Organic Acids and Amines
The Autoionization of Water
The pH Scale
Ionization Constants of Acids and Bases
Problem Solving Using Ka and Kb
Molecular Structure and Acid Strength
Acid-Base Reactions of Salts
Practical Acid-Base Chemistry
Lewis Acid and Bases

3. Types of Reactions a) Precipitation Reactions.
Reactions of ionic compounds or salts
b) Acid/base Reactions.
Reactions of acids and bases
c) Redox Reactions.
reactions of oxidizing & reducing agents

4. What are Acids &Bases? Definition? a) Arrhenius b) Bronsted-Lowry
c) Lewis

5. Arrhenius Definitions
Arrhenius, Svante August ( ), Swedish chemist, 1903 Nobel Prize in chemistry
Acid Anything that produces hydrogen ions in a water solution.
HCl (aq) H+ ( aq) + Cl- ( aq)
Base Anything that producs hydroxide ions in a water solution.
NaOH (aq) Na+ ( aq) + OH- ( aq)
Arrhenius definitions are limited proton acids and hydroxide bases to aqueous solutions.

6. Brønsted-Lowry definitions
Expands the Arrhenius definitions to include many bases other than hydroxides and gas phase reactions
Acid Proton donor
Base Proton acceptor
This definition explains how substances like ammonia can act as bases.
Eg. HCl(g) + NH3(g) > NH4Cl(s)
HCl (acid), NH3 (base).
NH3(g) + H2O(l) NH OH-

7. Lewis Definition
G.N. Lewis was successful in including acid and bases without proton or hydroxyl ions.
Lewis Acid: A substance that accepts an electron pair.
Lewis base: A substance that donates an electron pair.
E.g. BF3(g) + :NH3(g) F3B:NH3(s)
the base donates a pair of electrons to the acid forming a coordinate covalent bond common to coordination compounds. Lewis acids/bases will be discussed later in detail

8. Dissociation Strong Acids: HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq)
Dissociation Equilibrium Weak Acid/base:
H2O(l) + H2O(l) H3+O(aq) + OH-(aq)
This dissociation is called autoionization of water.
HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2-(aq)
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
Equilibrium constants: Ka, Kb and Kw

9. Brønsted-Lowry Definitions
Conjugate acid-base pairs.
Acids and bases that are related by loss or gain of H+ as H3O+ and H2O.
Examples. Acid Base
H3O+ H2O
HC2H3O2 C2H3O2-
NH4+ NH3
H2SO4 HSO4-
HSO4- SO42-

10. Bronsted acid/conjugate base and base/conjugate acid pairs in acid/base equilibria
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
HCl(aq): acid
H2O(l): base
H3+O(aq): conjugate acid
Cl-(aq): conjugate base
H2O/ H3+O: base/conjugate acid pair
HCl/Cl-: acid/conjugate base pair

11. Select acid, base, acid/conjugate base pair, base/conjugate acid pair
H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4-(aq)
acid
base
conjugate acid
conjugate base
base/conjugate acid pair
acid/conjugate base pair

12. Types of Acids and Bases
Binary acids: HCl, HBr, HI, H2S
More than two elements: HCN
Oxyacid: HNO3, H2SO4, H3PO4
Polyprotic acids: H2SO4, H3PO4
Organic acids: R-COOH, R= CH3-, CH3CH2-
Acidic oxides: SO3, NO2, CO2,
Basic oxides: Na2O, CaO
Amine: NH3. R-NH2, R= CH3-, CH3CH2- : primary
R2-NH : secondary, R3-N: tertiary
Lewis acids & bases: BF3 and NH3

13. Strong Acid vs. Weak Acids
completely ionized
Hydrioidic HI Ka ~ pKa = -11
Hydrobromic HBr Ka ~ 109 pKa = -9
Perchloric HClO4 Ka ~ 107 pKa = -7
Hyrdrochloric HCl Ka ~ 107 pKa = -7
Chloric HClO3 Ka ~ 103 pKa = -3
Sulfuric H2SO4 Ka ~ 102 pKa = -2
Nitric HNO3 Ka ~ 20 pKa = -1.3
Weak acid
partially ionized
Hydrofluoric acid HF Ka = 6.6x pKa = 3.18
Formic acid HCOOH Ka = 1.77x pKa = 3.75
Acetic acid CH3COOH Ka = 1.76x pKa = 4.75
Nitrous acid HNO2 Ka = 4.6x pKa = 3.34
Acetyl Salicylic acid C9H8O4 Ka = 3x pKa = 3.52
Hydrocyanic acid HCN Ka = 6.17x pKa = 9.21

14. Strong Base vs. Weak Base
completely ionized
Lithium hydroxide LiOH
Sodium hydroxide NaOH
Potasium hydroxide KOH Kb~
Rubidium hydroxide RbOH
Cesium hydroxide CsOH
Boarder-line Bases
Magnesium hydroxide Mg(OH)2
Calcium hydroxide Ca(OH)2
Strotium hydroxide Sr(OH) Kb~ 0.01 to0.1
Barium hydroxide Ba(OH)2
Weak Base
partially ionized
Ammonia NH3 Kb=1.79x10-5 pKb = 4.74
Ethyl amine CH3CH2NH2 Kb=5.6x10-4 pKb = 3.25

15. Acid and Base Strength
Strong acids Ionize completely in water HCl, HBr, HI, HClO3,
HNO3, HClO4, H2SO4.
Weak acids Partially ionize in water.
Most acids are weak.
Strong bases Ionize completely in water Strong bases are metal hydroxides - NaOH, KOH
Weak bases Partially ionize in water.

16. Common Acids and Bases Acids Formula Molarity* nitric HNO3 16
hydrochloric HCl
sulfuric H2SO4 18
acetic HC2H3O2 18
Bases
ammonia NH3(aq) 15
sodium hydroxide NaOH solid
*undiluted.

17. Autoionization of Water
Autoionization When water molecules react with one another to form ions.
Acids and bases alter the dissociation equilibrium of water based on Le Chaterlier’s principle
Kw = [ H3O+ ] [ OH- ]
= 1.0 x at 25oC
Note: [H2O] is constant and is included in Kw.
H2O(l) + H2O(l) H3O+(aq) + OH-(aq) (10-7M) (10-7M)
ion product
of water

18. pH and other “p” scales Substance pH 1 M HCl 0.0
Gastric juices
Lemon juice
Classic Coke
Coffee
Pure Water
Blood
Milk of Magnesia
Household ammonia
1M NaOH
We need to measure and use acids and bases over a very large concentration range.
pH and pOH are systems to keep track of these very large ranges.
pH = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14

19. pH scale
A logarithmic scale used to keep track of the large changes in [H+].
10-14 M M M
Very Neutral Very
acidic Basic
When you add an acid to, the pH gets smaller.
When you add a base to, the pH gets larger.

20. pH of some common materials
Substance pH
1 M HCl
Gastric juices
Lemon juice
Classic Coke
Coffee
Pure Water
Blood
Milk of Magnesia
Household ammonia
1M NaOH

21. pH of Aqueous Solutions

22. What is pH? Kw = [H3+O][OH-] = 1 x 10-14 [H3+O][OH-] = 10-7 x 10-7
Extreme cases:
Basic medium
[H3+O][OH-] = x 100
Acidic medium
[H3+O][OH-] = 100 x 10-14
pH value is -log[H+]
spans only 0-14 in water.

23. pH, pKw and pOH The relation of pH, Kw and pOH Kw = [H+][OH-]
log Kw = log [H+] + log [OH-]
-log Kw= -log [H+] -log [OH-] ;
previous equation multiplied by -1
pKw = pH + pOH; pKw = 14
since Kw =1 x 10-14
14 = pH + pOH
pH = 14 - pOH
pOH = 14 - pH

24. pH and pOH calculations of acid and base solutions
a) Strong acids/bases
dissociation is complete for strong acid such as HNO3 or base NaOH
[H+] is calculated from molarity (M) of the solution
b) weak acids/bases
needs Ka , Kb or percent(%)dissociation

25. pH of Strong Acid/bases
Substance pH
1 M HCl
Gastric juices
Lemon juice
Classic Coke
Coffee
Pure Water
Blood
Milk of Magnesia
Household ammonia
1M NaOH
HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq)
Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning.
[HNO3] = [H+] = 0.2 mole/L
pH = -log [H+]
= -log(0.2)
pH = 0.699

26. pH of 0.5 M H2SO4 Solution HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq)
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq)
HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq)
[H3+O][HSO4-]
H2SO4 ; Ka1 =
[H2SO4]
[H3+O][SO42-]
H2SO4 ; Ka2 = ; Ka2 ignored
[HSO4-]

27. pH of 0.5 M H2SO4 Solution pH = -log(0.5) pH = 0.30
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq)
the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning.
[H2SO4] = [H+] = 0.5 mole/L
pH = -log [H+]
pH = -log(0.5)
pH = 0.30

28. 1.5 x 10-2 M NaOH.
1.5 x 10-2 M NaOH.
NaOH is also a strong base dissociates completely in water.
[NaOH] = [HO- ] = 1.5 x 10-2 mole/L
pOH = -log[HO-]= -log(1.5 x 10-2)
pOH = 1.82
As defined and derived previously:
pKw= pH + pOH; pKw= 14
pH = pKw + pOH
pH = 14 - pOH
pH = ; pH = 12.18

29. Mixtures of Strong and Weak Acids
the presence of the strong acid retards the dissociation of the weak acid

30. Measuring pH Arnold Beckman inventor of the pH meter
father of electronic instrumentation

31. Equilibrium, Constant, Ka & Kb
Ka: Acid dissociation constant for a equilibrium reaction.
Kb: Base dissociation constant for a equilibrium reaction.
Acid: HA + H2O H3+O + A-
Base: BOH + H2O B+ + OH-
[H3+O][ A-] [B+ ][OH-]
Ka = ; Kb =
[HA] [BOH]

32. Acid Dissociation Constant
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
[H3+O][Cl-]
Ka=
[HCl]
[H+][Cl-]

33. Base Dissociation Constant
NH3 + H2O NH OH-
[NH4+][OH-]
K =
[NH3]

34. Hydrated Metal Ions as Acids

35. Ionization Constants for Acids

36. Comparing Kw and Ka & Kb
Any compound with a Ka value greater than Kw of water will be a an acid in water.
Any compound with a Kb value greater than Kw of water will be a base in water.

37. WEAKER/STRONGER Acids and Bases & Ka and Kb values
A larger value of Ka or Kb indicates an equilibrium favoring product side.
Acidity and basicity increase with increasing Ka or Kb.
pKa = - log Ka and pKb = - log Kb
Acidity and basicity decrease with increasing pKa or pKb.

38. Which is weaker? a. HNO2 ; Ka= 4.0 x 10-4. b. HOCl2 ; Ka= 1.2 x 10-2.
c. HOCl     ;  Ka= 3.5 x 10-8.
d. HCN      ;  Ka= 4.9 x

39. What is Ka1 and Ka2? H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq)
HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq)

40. Ka Examples H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq)
HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq)
[H3+O][HSO4-]
H2SO4 ; Ka1 =
[H2SO4]
[H3+O][SO42-]
H2SO4 ; Ka2 =
[HSO4-]

41. Ka Examples HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2-(aq) [H+][C2H3O2-]
H C2H3O2; Ka=
[H C2H3O2]
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
[NH4+][OH-]
NH3; Kb=
[ NH3]

42. How do you calculate pH of weak acids/bases
From % dissociation
From Ka or Kb
What is % dissociation
Amount dissociated
% Dissoc. = x 100
Initial amount

43. How do you calculate % dissociation from Ka or Kb
1.00 M solution of HCN; Ka = 4.9 x 10-10
What is the % dissociation for the acid?

44. 1.00 M solution of HCN; Ka = 4.9 x 10-10
First write the dissociation equilibrium equation:
HCN(aq) + H 2O(l) <===> H 3+O(aq) + CN-(aq)
[HCN] [H+ ] [CN- ]
Ini. Con M M M
Cha. Con x x x
Eq. Con x x x
[H 3+O ][CN-] x2
Ka = =
[HCN] x

45. 1.00 M solution of HCN; Ka = 4.9 x 10-10
x ~ since x is small x2
Ka = ; Ka = 4.9 x = x2
1.0
x = x = x
Amount disso x
x 100 = x 100
Ini. amount
% Diss. =2.21 x x 100 = %

46. % Dissociation gives x (amount dissociated) need for pH calculation
% Dissoc. = x 100
Initial amount/con. x
% Dissoc. = x 100
concentration

47. Calculate the pH of a weak acid from % dissociation
1 M HF, 2.7% dissociated
Notice the conversion of % dissociation to a fraction (x): 2.7/100=0.027) x=0.027

48. Calculate the pH of a weak acid from % dissociation
HF(aq) + H 2O(l) <===> H 3+O(aq) + F-(aq)
[H+][F-]
Ka =
[HF]
[HF] [H+ ] [F- ]
Ini. Con M 0.0 M 0.00 M
Chg. Con -x x x
Eq.Con
pH = -log [H+]
pH = -log(0.027)
pH = 1.57

49. HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq)
Weak acid Equilibria
Example
Determine the pH of a 0.10 M benzoic acid solution at 25 oC if Ka = x 10-5
HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq)
The first step is to write the equilibrium expression
Ka =
[H3O+][Bz-]
[HBz]

50. Weak acid Equilibria HBz H3O+ Bz- Initial conc., M 0.10 0.00 0.00
Change, DM -x x x
Eq. Conc., M x x x
[H3O+] = [Bz-] = x
We’ll assume that [Bz-] is negligible compared to [HBz]. The contribution of H3O+ from water is also negligible.

51. Weak Acid Equilibria Solve the equilibrium equation in terms of x
Ka = x =
x = (6.28 x )(0.10)
H3O+ = M
pH = 2.60 x2
0.10

52. pH from Ka or Kb 1.00 M solution of HCN; Ka = 4.9 x 10-10
First write the dissociation equilibrium equation:
HCN(aq) + H 2O(l) H 3+O(aq) + CN-(aq)
[HCN] [H+ ] [CN- ]
Ini. Con M M M
Chg. Con x x x
Eq. Con x x x

53. Weak Acid Equilibria [H 3+O ][CN-] x2 1.0 - x ~ 1.00 since x is small
[HCN] x
x ~ since x is small x2
Ka = ; Ka = 4.9 x = x2
1.0
x = x = x
pH = -log [H+]
pH = -log(2.21 x 10-5)
pH = 4.65

54. The Conjugate Partners of Strong Acids and Bases
The conjugate acid/base of a strong base/acid has no net effect on the pH of a solution
The conjugate base of a weak acid hydrolyze in water and basic or
pH of a solution > E.g. Na+C2H3O2- sodium acetate
The conjugate acid of a weak base hydrolyze in water and acidic or
pH of a solution < 7.00 E.g NH4Cl

55. Hydrolysis
Reaction of a basic anion or acidic cation with water is an ordinary Brønsted-Lowry acid-base reaction.
CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq)
NH4+(aq) + H2O(l) NH3 (aq) + H3O+(aq)
This type of reaction is given a special name.
Hydrolysis
The reaction of an anion with water to produce the conjugate acid and OH-.
The reaction of a cation with water to produce the conjugate base and H3O+.

56. Acid-Base Properties of Typical Ions

57. What salt solutions would be acidic, basic and neutral?
1) strong acid + strong base = neutral
2) weak acid + strong base = basic
3) strong acid + weak base = acidic
weak acid + weak base = neutral,
basic or an acidic solution depending on the relative strengths of the acid and the base.

58. What pH? Neutral, basic or acidic?
a)NaCl
neutral
b) NaC2H3O2
basic
c) NaHSO4
acidic
d) NH4Cl

59. How do you calculate pH of a salt solution?
Find out the pH, acidic or basic?
If acidic it should be a salt of weak base
If basic it should be a salt of weak acid
if acidic calculate Ka from Ka= Kw/Kb
if basic calculate Kb from Kb= Kw/Ka
Do a calculation similar to pH of a weak acid or base

60. What is the pH of 0.5 M NH4Cl salt solution? (NH 3; Kb = 1.8 x 10-5)
Find out the pH, acidic
if acidic calculate Ka from Ka= Kw/Kb
Ka= Kw/Kb = 1 x /1.8 x 10-5)
Ka= X 10-10
Do a calculation similar to pH of a weak acid

61. Continued NH4+ + H2O H 3+O + NH3 [NH4+] [H3+O ] [NH3 ]
Ini. Con M M 0.00 M
Change x x x
Eq. Con x x x
[H 3+O ] [NH3 ]
Ka(NH4+) = =
[NH 4+] x2
; appro.: x
( x)

62. Continued x2 Ka(NH4+) = ----------- = 5.56 x 10 -10 0. 5
x2 = x x 0.5 = x
x= x = x 10-5
[H+ ] = x = 1.66 x 10-5 M
pH = -log [H+ ] = - log x 10-5
pH = 4.77
pH of 0.5 M NH4Cl solution is 4.77 (acidic)

63. Types of Acids and Bases
Binary acids
Oxyacid
Organic acids
Acidic oxides
Basic oxides
Amine
Polyprotic acids

64. Influence of Molecular Structure on Acid Strength
Binary Hydrides
hydrogen & one other element
Bond Strengths
weaker the bond, the stronger the acid
Stability of Anion
higher the electronegativity, stronger the acid

65. Binary Acids
Compounds containing acidic protons bonded to a more electronegative atom.
e.g. HF, HCl, HBr, HI, H2S
The acidity of the haloacid
(HX; X = Cl, Br, I, F)
Series increase in the following order:
HF < HCl < HBr < HI

66. Oxyacids Compounds containing acidic - OH groups in the molecule.
Acidity of H2SO4 is greater than H2SO3 because of the extra O (oxygens)
The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend.
HClO4 > HClO3 > HClO2 > HClO
perchloric chloric chlorus hyphochlorus

67. Influence of Molecular Structure on Acid Strength
Oxyacids
hydrogen, oxygen, & one other element
H-O-E
higher the electronegativity on E, stronger the acid as this weakens the bond between the O and H

68. Oxo Acid
<
<
<
<

69. Acidic Oxides
These are usually oxides of non-metallic elements such as P, S and N.
E.g. NO2, SO2, SO3, CO2
They produce oxyacids when dissolved in water
SO3 + H2O ---> H2SO4
CO2 + H2O ---> H2CO3
NO2 + H2O ---> HNO3

70. Basic Oxides
Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water.
e.g. CaO + H2O ---> Ca(OH)2
Na2O + H2O ---> 2 NaOH

71. Protic Acids
Monoprotic Acids: The form protic refers to acidity due to protons. Monoprotic acids have only one acidic proton. e.g. HCl.
Polyprotic Acids: They have more than one acidic proton.
e.g. H2SO diprotic acid
H3PO triprotic acid.

72. Polyprotic Acids
acids where more than one hydrogen per molecule is released

73. Polyprotic Acids

74. Organic or Carboxylic Acids

75. Organic or Carboxylic Acids
FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid).
Acid Ka pKa
HCOOH (formic acid) X
CH3COOH (acetic acid) 1.74 X
CH3CH2COOH (propanoic acid)1.38 x

76. Amines
Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted-Lowery or Lewis acid/base definitions.

77. Amines

78. Acid-Base Chemistry of Some Antacids

79. Acid-Base in the Kitchen
vinegar - acetic acid
lemon juice (citrus juice) - citric acid
baking soda - NaHCO3
milk - lactic acid
baking powder - H2PO4- & HCO3-

80. Household Cleaners

81. Dishwashing Detergent

82. Lewis Definition
G.N. Lewis was successful in including acid and bases without proton or hydroxyl ions.
Lewis Acid: A substance that accepts an electron pair.
Lewis base: A substance that donates an electron pair.
E.g. BF3(g) + :NH3(g) F3B:NH3(s)
the base donates a pair of electrons to the acid forming a coordinate covalent bond common to coordination compounds. Lewis acids/bases will be discussed later in detail

83. Lewis Acids and Bases Reactions
H+ + NH3  NH4+
acid base
Cu NH3  [Cu(NH3)4+2]
acid base

84. What acid base concepts (Arrhenius/Bronsted/Lewis) would best describe the following reactions:
a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)
b)HCl(g) + NH3(g) ---> NH4Cl(s)
c)BF3(g) + NH3(g) ---> F3B:NH3(s)
d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)