1. Chemistry Chapter 20-21
Acids and Bases

2. Electrolytes: Solutions that conduct electrical current easily.
Aqueous solutions of acids and bases are electrolytes.
Ionic solids that dissolve in water are electrolytes.

3. Electrolytes: The ability of a solution to conduct an electric current.
The ammeter measures the flow of electrons (current)
through the circuit.
If the ammeter measures a current, and the bulb
glows, then the solution conducts. (Electrolyte)
If the ammeter fails to measure a current, and the
bulb does not glow, the solution is non-conducting. (Non-electrolyte)

4. NaCl(s)  Na+(aq) + Cl-(aq)
Dissociation: Water surrounds ions and dissolves them.
Dissociation equation:
NaCl(s) 
Na+(aq) + Cl-(aq)

5. Dissociation of HCl HCl  H+ + Cl- HA  H+ + A-
equation.
HCl 
H+ + Cl-
General
Dissociation
equation.
HA 
H+ + A-
A- represents the anion

6. Basic definition:
Acids: Produce hydrogen ions, H+ , when dissolved in water.
Bases: Produce hydroxide ions, OH-, when dissolved in water.

7. Acid-Base Neutralization+
Base 
Water +
Salt
HCl +
NaOH 
HOH
H2O +
NaCl

8. Some Properties of Acids
Acids taste sour
Acids react with certain metals to produce
hydrogen gas.
2 HCl(aq) + Mg(s)  H2(g) + MgCl2(aq)
Acids effect indicators
Blue litmus turns red
Phenolphthalein turns colorless
Methyl orange turns red

9. Acids Effect on Indicators
Blue litmus paper turns red in contact with an acid.

10. Bases Effect on Indicators
Red litmus paper turns blue in contact with a base.
Phenolphthalein turns purple in a base.

11. pH Indicators and their ranges

12. Hydro_______ic acid Naming Acids Acids without oxygen: HF =
Hydrofluoric acid
HCl =
Hydrochloric acid
HBr =
Hydrobromic acid
HI =
Hydroiodic acid
H2S =
Hydrosulfuric acid
HCN =
Hydrocyanic acid
CN- = cyanide anion

13. Naming Acids H+ + (anions with oxygen)
Acids with oxygen: “ate-ic ite-ous”
Anion name Acid name
____ate ____ic acid
___ite ___ous acid
HNO3 =
Nitr__?__ acid
Anion NO3-1 = nitrate
Nitric acid
HNO2 = Nitr__?__ acid
Anion NO2-1 = nitrite
Nitrous acid

14. Naming Acids H+ + (anions with oxygen)
Formula H+ with Anion Acid name
H2SO4
2H+ + SO42- = sulfate
Sulfuric acid
H2SO3
2H+ + SO32- = sulfite
Sulfurous acid
HC2H3O2
H+ + C2H3O2-1 = Acetate
Acetic acid
H2CO3
2H+ + CO32- = Carbonate
Carbonic acid
H3PO4
3H+ + PO43- = Phosphate
Phosphoric acid

15. Naming Acids H+ + (series of Cl and O)
Formula H+ with Anion Acid name
HCl
H+ + Cl = Chloride
Hydrochloric acid
HClO
H+ + ClO-1 = hypochlorite
Hypochlorous acid
HClO2
H+ + ClO2-1 = Chlorite
Chlorous acid
HClO3
H+ + ClO3-1 = Chlorate
Chloric acid
HClO4
H+ + ClO4-1 = perchlorate
Perchloric acid
Hypo = less oxygen
Per = more oxygen

16. Naming Bases Same method as naming ionic compounds
Formula Cation with OH- Base name
NaOH
Na+ + OH-
Sodium hydroxide
(Lye)
KOH
K+ + OH-
Potassium hydroxide
Ca(OH)2
Ca OH-
Calcium hydroxide
(Lime)
Mg(OH)2
Mg OH-
Magnesium hydroxide
(Milk of magnesia)
Al(OH)3
Al OH-
Aluminum hydroxide
NH3
(No ions)
Ammonia
(common name)
Molecules that contain nitrogen are weak bases.

17. (Self-Ionization of Water)
Two water molecules collide to form Hydronium and Hydroxide ions.
H2O + H2O  H3O+ + OH-
1 in 10,000,000 water molecules self-ionize.
This is 10-7 water molecules.

18. Kw – Ionization Constant for Water
In pure water at 25 C:
[H+] = [H3O+] = 1 x 10-7 mol/L 1
10,000,000
(Dissociate)
[OH-] = 1 x 10-7 mol/L
At 25 C, 1 in 10,000,000 water molecules
will dissociate and form H+ and OH-.
Kw = Ion product constant (for water at 25 C)
Kw = [H+] x [OH-]
Kw = (1 x 10-7)(1 x 10-7) = 1 x 10-14

19. Kw = [H+] x [OH-]
Ion-product constant = Hydrogen-ion concentration x hydroxide-ion concentration

20. Ion Concentration in Water

21. Calculating pH “per Hydronium”
pH = -log10[H3O+]
In other words: pH = -log[H+]
[H+] = H+ concentration in (mol/L) or (M)
Example:
[H+] = 1 x 10-7 M
pH = -log(1 x 10-7)
log(1) = 0
log (10-7) = -7
pH = -(0-7) = 7.0

22. Calculating pH Examples: [H+] = 1 x 10-1 pH = -log(1 x 10-1)

23. Acids Have a pH less than 7

24. Bases have a pH greater than 7

25. Kw = [H+] x [OH-]
Ion-product constant = Hydrogen-ion concentration x hydroxide-ion concentration
Basic, neutral and acidic are all ways of describing pH concentrations in a solution.

26. Calculating pH Examples: [H+] = 1 x 10-1 pH = 1.0 pH = 4.0
Acid
pH = 4.0
Acid
[H+] = 1 x 10-4
pH = 7.0
Neutral
[H+] = 1 x 10-7
pH = 9.0
Base
[H+] = 1 x 10-9
pH = 13.0
Base
[H+] = 1 x 10-13
pH = 0
Acid
[H+] = 1 x 10 0

27. Calculating pH Examples: pH + pOH = 14 [OH-] = 1 x 10-2 pOH = 2.0
Base
14.0
[OH-] = 1 x 10-5
pOH = 5.0
pH = 9.0
Base
14.0
[OH-] = 1 x 10-7
pOH = 7.0
pH = 7.0
Neutral
14.0

28. Calculating pH pH = -log [H+] pOH = -log [OH-] pH + pOH = 14
[OH-] = 1 x 10-8
pOH = 8.0
pH = 6.0
Acid
14.0
[OH-] = 1 x 10-10
pOH = 10.0
pH = 4.0
Acid
14.0
[OH-] = 1 x 10-14
pOH = 14.0
pH = 0
Acid
14.0

29. Calculating pH pH = -log [H+] pOH = -log [OH-] pH + pOH = 14
[OH-] = 1 x 10 0
pOH = 0
pH = 14.0
Base
14.0

30. Examples (Using Calculator):
pH = -log [H+]
pOH = -log [OH-]
pH + pOH = 14
Examples (Using Calculator):
[H+] = 2.5 x 10-5
pH = -log(2.5 x 10-5)
pH = 4.60
Acid
[H+] = 3.0 x 10-9
pH = -log(3.0 x 10-9)
pH = 8.52
Base

31. pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 Examples:
[OH-] = 6.0 x 10-3
pOH = -log(6.0 x 10-3)
pOH = 2.22
pH = 11.78
Base
14.0
[OH-] = 9.0 x 10-12
pOH = -log(9.0 x 10-12)
pOH = 11.05
pH = 2.95
Acid
14.0

32. Equation Summary pH = -log [H+] pOH = -log [OH-] pH + pOH = 14
Kw = [H+] x [OH-]
[OH-] = 10-pOH
1 x = [H+] x [OH-]
[H+] = 10-pH
pH + pOH = 14
[OH-] = 10-pOH
pOH = 14 - pH
[OH-] = 10-9
[H+] = 10-5
Acid
pOH = 9
1 x 10-9 M 5
1 x 10-5 M
pOH = ___
pH = ___
[H+] = __________
[OH-] = __________
pH + pOH = 14
[H+] = 10-pH
[OH-] = 10-pOH
pH = 14 - pOH
[OH-] = 10-6
[H+] = 10-8
pH = 6 8
pOH = ___
1 x 10-6 M
pH = ___
[H+] = __________
1 x 10-8 M
[OH-] = __________
Base

33. Equation Summary pH = -log [H+] pOH = -log [OH-] pH + pOH = 14
Kw = [H+] x [OH-]
[OH-] = 10-pOH
1 x = [H+] x [OH-]
[H+] = 10-pH
pH + pOH = 14
[OH-] = 10-pOH
pOH = 14 - pH
[OH-] =
[H+] =
Acid
pOH = 14 – 2.8
[H+] =
11.2
6.3 x M
2.8
1.6 x 10-3 M
pOH = ___
pH = ___
[H+] = __________
[OH-] = __________

34. Acid-Base Theories Theory Acid Base Arrhenius Produces H+
ions in solution
Produces OH-
ions in solution
HCl
HC2H3O2
NaOH
Ca(OH)2

35. Acid-Base Theories Base Acid NH3 + HOH  HOH NH4+ + OH- Acid Base
Theory
Acid
Base
Bronsted-Lowry
Proton (H+ ion)
donor.
Proton (H+ ion) acceptor.
Includes more molecules that do not have OH-
NH3 (ammonia)
Conjugate Acid-Base Pairs
Base
(proton acceptor)
Acid
(Conjugate acid)
proton donor
NH3 +
HOH 
HOH
NH4+ +
OH-
Acid
(proton donor)
Base
(conjugate base)
proton acceptor

36. Ionization of HCl and formation of hydronium ion, H3O+
H2O + HCl 
H3O+ + Cl-
Proton
acceptor
Proton
donor
Hydronium
ion
Chloride
ion
Acid
Conjugate
Base
Bronsted/Lowry (Acid-Base Pair)

37. Ionization of HCl and formation of hydronium ion, H3O+
H2O + HCl 
H3O+ + Cl-
Proton
acceptor
Proton
donor
Hydronium
ion
Chloride
ion
Conjugate
Acid
Base
Another Acid-Base Pair

38. Acids are Proton Donors
Bronsted/Lowry Theory = proton donor idea
Monoprotic acids
Diprotic acids
Triprotic acids
H3PO4
HCl
H2SO4
HNO3
H2CO3
HC2H3O2
Note: a proton is a hydrogen ion.
“protic” = proton
(All theories agree these are acids)

39. Acid-Base Theories Theory Acid Base Lewis (e- pair) acceptor
(e- pair) donor
BCl3
NH3 B N H Cl Cl H H Cl

40. Lewis Acid-Base Theory
Lewis Base
e- pair acceptor
e- pair donor Cl Al Cl N H Cl H H
AlCl3
(aluminum chloride)
NH3
(ammonia)

41. Lewis Acid-Base Theory
Lewis Base
e- pair acceptor
e- pair donor Al Cl N H
AlCl3
(aluminum chloride)
NH3
(ammonia)

42. Acid-Base Theories Summary
Theory
Acid
Base
Arrhenius
Produces H+
ions in solution
Produces OH-
ions in solution
HCl
HC2H3O2
NaOH
Ca(OH)2
HOH
HOH
Bronsted-Lowry
Amphoteric: A substance that can act as an acid or a base.
Water is an example:
H2O = HOH
H+ ion (proton)
donor.
H+ ion (proton) acceptor.
NH3 (ammonia)
Lewis
(e- pair) acceptor
(e- pair) donor
NH3
BCl3

43. Strong Acid Dissociation
H+ + H2O  H3O+
Hydronium is a result of proton transfer (H+ donated to H2O).

44. Strong Acids
Strong acids are assumed to be 100% ionized in solution (good proton donors).
HCl
H Cl-
H2SO4
H HSO4-
HNO3
H NO3-
Arrow in one direction indicates
that molecule dissociates completely.

45. Weak Acids
Weak acids are usually less than 5% ionized in solution (poor proton donors). HF
H F-
H3PO4
H H2PO4-
H2PO4-
H HPO42-
Double arrow = (Reaction goes both ways)
The larger arrow = (Favored side)

46. Weak Acid Dissociation
Weak acids are mostly molecules.
(Only 5% or less dissociate)

47. Acid Dissociation Constant (Weak Acids)
Ka = product concentrations vs. reactant concentrations for weak acids.
[HF]
HF(aq) +
H2O(l)
[H3O+]
H3O+(aq) F-(aq)
[F-]
Ka =
H2O not included because it is in a different phase.

48. Base Dissociation Constant (Weak Bases)
Kb = product concentrations vs. reactant concentrations for weak acids.
[NH3]
NH3(aq) +
HOH(l)
[NH4+]
NH4+(aq) OH-(aq)
[OH-]
Kb =
H2O not included because it is in a different phase.

49. HX(aq) H+(aq) + X-(aq) [H+] [X-] Ka = = = [HX] Initial Change
0.35 M
Change
0.041 M
0.041 M M
Equilibrium
0.35 – = M
0.041 M
0.041 M
[H+]
[X-]
(0.041)
(0.041)
Ka = = = M
[HX]
0.309

50. Methyl Red changes from red to orange in this range.
Phenolphthalein changes to pink in this range.

53. HIn(aq) H+(aq) + In-(aq)

54. Phenolphthalein  Bromthymol blue  Methyl Red
(Weakest acid) (Strongest Acid)

55. H2In(aq)  H+(aq) + HIn-(aq) Red to yellow.
The indicator must be a diprotic acid for 2 color changes.
H2In(aq)  H+(aq) HIn-(aq)
Red to yellow.
HIn-(aq)  H+(aq) In2-(aq)
Yellow to blue.

56. The other terms are all associated with bases.
Kw = [H+] x [OH-]
Ion-product constant = Hydrogen-ion concentration x hydroxide-ion concentration
Basic, neutral and acidic are all ways of describing pH concentrations in a solution.
The other terms are all associated with bases.
The other terms are related to pH or can be calculated from pH.

57. The other terms are all acid-base theories.
Conjugate acids and bases are part of Bronsted-Lowry theory.
The other terms describe ionizable hydrogen ions..
Dissociation constants are used with weak acids and bases.

58. Sulfuric Acid Used in the production of paper
Highest volume production of any chemical in the U.S.
Used in the production of paper
Used in production of fertilizers
Used in petroleum refining

59. Nitric Acid Used in the production of fertilizers
Used in the production of explosives
Nitric acid is a volatile acid – its reactive components evaporate easily
Stains proteins (including skin!)

60. Hydrochloric Acid Used in the pickling of steel
Used to purify magnesium from sea water
Part of gastric juice, it aids in the digestion of protein
Sold commercially as “Muriatic acid”

61. Phosphoric Acid A flavoring agent in sodas
Used in the manufacture of detergents
Used in the manufacture of fertilizers
Not a common laboratory reagent

62. Acetic Acid Used in the manufacture of plastics
Used in making pharmaceuticals
Acetic acid is the acid present in vinegar

63. Examples of Organic Acids
Citric acid in citrus fruit
Malic acid in sour apples
Deoxyribonucleic acid, DNA
Amino acids, the building blocks of protein
Lactic acid in sour milk and sore muscles
Butyric acid in rancid butter

64. Weak Acids: (Organic) Organic acids (Acids that contain carbon)
are always weak acids.
HC2H3O2
H C2H3O2-
CH3COOH
H CH3COO-
Same acid written another way
Carboxylic Acid group: R-COOH

65. Organic Acids
Organic acids all contain the “carboxyl” group called “carboxylic acid”, sometimes several of them.
The carboxyl group is a poor proton donor, so ALL organic acids are weak acids.

66. Properties of Bases Bases effect indicators Red litmus turns blue
Bases taste bitter
Bases effect indicators
Red litmus turns blue
Phenolphthalein turns purple
Bases have a pH greater than 7
Bases are proton (hydrogen ion, H+) acceptors
Solutions of bases feel slippery
Bases neutralize acids

67. Products of Neutralization
HCl + NaOH 
NaCl + H2O
H2SO4 + Ca(OH)2 
CaSO H2O
HNO3 + KOH 
KNO3 + H2O
The products of neutralization are always a ______ and _______.
salt
water

68. Measuring pH with wide-range paper

69. Narrow-Range pH Paper

70. pH Indicators and their ranges

71. Calculating [H+] and [OH-]
[H+] = 10-pH and [OH-] = 10-pOH
Examples:
pH = 2.0
[H+] = 1 x 10-2
Acidic
pH = 12.0
Basic
[H+] = 10-12
pH = 4.3
Acidic
[H+] = = 5 x 10-5
pOH = 4.0
[OH-] = 1 x 10-4
pH=10
Basic
pOH = 8.0
[OH-] = 1 x 10-8
pH=6
Acidic
pOH = 10.0
[OH-] = 1 x 10-10
pH=4
Acidic

72. SUMMARY Calculating pH, pOH
pH = -log[H3O+] or pH = -log[H+]
pOH = -log[OH-]
Relationship between pH and pOH
pH + pOH = 14
Finding [H3O+], [OH-] from pH, pOH
[H3O+] = 10-pH or [H+] = 10-pH
[OH-] = 10-pOH

73. Acids React with Active Metals
Acids react with active metals to form salts and hydrogen gas.
Mg + 2HCl  MgCl2 + H2(g)

74. Acids React with Carbonates
2HC2H3O2 + Na2CO3
2 NaC2H3O2 + H2O + CO2

75. Effects of Acid Rain on Marble (calcium carbonate)
George Washington:
BEFORE
George Washington:
AFTER

76. Acids Neutralize Bases
HCl + NaOH  NaCl + H2O
Neutralization reactions ALWAYS produce a salt and water.

77. Products of Neutralization (Lab 23)
Trial 1 (Strong acid and Strong Base)
HCl(aq) + NaOH(aq) 
NaCl(aq) + H2O(l)
Trial 2 (Weak acid and Strong Base)
HC2H3O2(aq) + NaOH(aq) 
NaC2H3O2 (aq) + H2O(l)
Trial 3 (Strong acid and Weak Base)
HCl(aq) + NH3 (aq) 
NH4+1(aq) + Cl-1(aq)
Trial 4 (Weak acid and Weak Base)
HC2H3O2(aq) + NH3(aq) 
NH4+1(aq) + C2H3O2-1(aq)

78. Titration Curves

79. Bases Neutralize Acids
Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl.
2 HCl + Mg(OH)2
MgCl2 + 2 H2O

80. Titration Calculations