1. The Chemistry of Acids and Bases

2. Acids pH below 7 turns litmus paper red taste sour
reacts with metals to produce H2(g)
generally starts with a hydrogen ion
[H+] > [OH-]
HCl

3. generally contains a hydroxide ion
Bases
pH greater than 7
turns litmus paper blue
taste bitter
feel slippery
generally contains a hydroxide ion
[H+] < [OH-]
NaOH

4. contains H+ and OH- ions
Both Acids and Bases
amphoteric
an electrolyte
contains H+ and OH- ions

5. Acidic, Basic, and Neutral Solutions
Type of Solution
pH Ranges
[H+] versus [OH-]
Example
Acidic
Below 7
[H+] > [OH-]
Orange Juice
Battery Acid
Your Stomach
Neutral
Equals EXACTLY 7
[H+] = [OH-]
Distilled
Water
Basic
Above 7
[H+] < [OH-]
Bleach
Sea Water
Blood

6. Indicators
Indicators are compounds that have one color in acidic solutions and another in basic.

7. Litmus Paper
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

8. Paper or plastic strips that contain combinations of indicators estimate the pH of a solution by simply dipping a piece of pH paper into it and comparing the resulting color with standards printed on the container
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

9. Range and Color Changes of Some Common Acid-Base Indicators
pH Scale 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Indicators
Methyl orange red – yellow
Methyl red red yellow
Bromthymol blue yellow blue
Neutral red red yellow
From F. Brescia et al., Chemistry: A Modern Introduction, W. B. Saunders Co., 1978.
Adapted from R. Bates, Determination of pH, Theory and Practice, John Wiley & Sons, Inc., New York, 1964.
Choosing the correct indicator for an acid-base titration
1. For titrations of strong acids and strong bases (and vice versa), any indicator with a pKin between 4 and 10 will do
2. For the titration of a weak acid, the pH at the equivalence point is greater than 7, and an indicator such as phenolphthalein or thymol blue, with pKin > 7, should be used
3. For the titration of a weak base, where the pH at the equivalence point is less than 7, an indicator such as methyl red or bromcresol blue, with pKin < 7, should be used
Phenolphthalein colorless red colorless beyond 13.0
Bromthymol blue indicator would be used in titrating a strong acid with a strong base.
Phenolpthalein indicator would be used in titrating a weak acid with a strong base.
Methyl orange indicator would be used in titrating a strong acid with a weak base.

10. pH
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

11. Naming Acids Naming Binary Acids 1. Name begins with prefix hydro-
2. Root of second element follows hydro and ends with the suffix -ic

12. BINARY ACIDS
HBr (aq)
Hydrobromic Acid

13. BINARY ACIDS
HCl (aq)
Hydrochloric Acid

14. BINARY ACIDS
HI (aq)
Hydroiodic Acid

15. Examples of BINARY ACIDS
Hydrofluoric acid HF
Hydrochloric acid HCl
Hydrobromic acid HBr
Hydrosulfuric acid H2S

16. -ous polyatomic ending -ic -ate -ite Nitrate Naming Ternary Acids
H + polyatomic ion
Begin with ion without
the
Add suffix if there was an
Add suffix if there was an
HNO3 
polyatomic
ending
-ic
-ate
-ous
-ite
Nitrate
“In the cafeteria, you ATE something ICky”
Nitric acid

17. Naming Ternary Acids TERNARY ACIDS SO42- H2SO4 H2SO4(aq) Sulfuric acid
POLYATOMIC
IONS
PURE
FORMS
TERNARY ACIDS
SO42-
H2SO4
H2SO4(aq)
Sulfuric acid
SO32-
H2SO3
H2SO3(aq)
Sulfurous Acid

18. TERNARY ACIDS Naming Ternary Acids PO43- H3PO4 H3PO4(aq)
POLYATOMIC
IONS
PURE
FORMS
TERNARY ACIDS
PO43-
H3PO4
H3PO4(aq)
Phosphoric acid
PO33-
H3PO3
H3PO3(aq)
Phosphorus Acid

19. Acid Nomenclature Flowchart

20. Naming Bases polyatomic Sodium hydroxide Calcium hydroxide
Use the same rules as for ions (name the cation, then name the anion)
NaOH 
Ca(OH)2 
KOH 
Sodium hydroxide
Calcium hydroxide
Potassium hydroxide

21. Naming Bases
NaOH
Sodium Hydroxide

22. Naming Bases
MgOH
Magnesium Hydroxide

23. Naming Bases
KOH
Potassium Hydroxide

24. Common Bases Lye, Drano Potash Cleaners Baking soda Tums /Rolaids
Milk of Magnesia
Baking soda
Tums /Rolaids
Limestone, shells
Soaps
Detergents
NaOH
KOH
NH3 or NH4OH
Mg(OH)2
NaHCO3
CaCO3
NaC16O2H31
NaC12O4H25S
Sodium Hydroxide
Potassium Hydroxide
Ammonia
Magnesium Hydroxide
Sodium Bicarbonate
Calcium Carbonate
Sodium Palmitate
Sodium Lauryl Sulfate

25. Common Acids Battery Acid Stomach Acid Coca Cola Carbonated Water
H2SO4
HCl
H3PO4
H2CO3
HC2H3O2
H3C6O7H8
HC6O6H7
H2C4O6H4
H2C9O4H8
Battery Acid
Stomach Acid
Coca Cola
Carbonated Water
Vinegar
Citrus fruits
Vitamin C
Grapes
Aspirin
Sulfuric Acid
Hydrochloric Acid
Phosphoric Acid
Carbonic Acid
Acetic Acid
Citric Acid
Ascorbic Acid
Tartaric Acid
Acetyl Salicylic Acid

26. Definitions for Acids & Bases
Arrhenius
Brønsted-Lowry
Definition for Acids
Definition for Bases
Key Examples
a proton producer in an aqueous solution
a proton donor
a hydroxide producer in an aqueous solution
a proton acceptor
Acid – HCl
Base - NaOH
Acid – HCl
Base – NH3
H+ = proton

27. What to Focus On?
Arrhenius was the most restrictive definition. This definition required:
the solutions to be aqueous and
a base to contain a hydroxide (OH-) ion.
Bronsted-Lowry’s definition is the most commonly used. It is helpful to remember:
acids tend to “lose“ an H+ ion, while
bases tend to “gain“ an H+ ion.
Under this definition, ammonia (NH3) is considered a base even though it is NOT an Arrhenius base.

28. Examples HCl(aq) + KOH(s) H2O(l) + KCl(aq)
3 Ca(OH)2(aq) + 2 H3PO4(aq)  Ca3(PO4)2(s) + H2O(l)
F-(aq) + H2O(l)  HF(aq) + OH-(aq)
HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq)
NH4+(aq) + CO32-(aq)  NH3(aq) + HCO3-(aq)
acid
base
base
acid
base
acid
acid
base
acid
base

29. Remember Electrolytes?
Ionic
Covalent
C6H12O6
Na+
NaCl
Cl-

30. Acids and bases are both strong or weak electrolytes (conduct electricity)
• Electrolytes = dissociate (break apart into ions) when dissolved
• Strong = completely Weak = partially Non = not at all
Weak
Strong H+
HC2H3O2
C2H3O21- H+
H-Cl
Cl-
Only a few Ions
Lots of Ions

31. Strong Electrolytes
PICTURE
WORD DESCRIPTION
Completely breaks apart into its ions
Are good conductors of electricity
Will produce a bright light bulb
Examples of Acids and Bases that are Strong Electrolytes
Strong Acids Strong Bases
H2SO4 NaOH
HCl Ba(OH)2
Notice that all of the ions are separated or dissociated.
= positive ions
= negative ions
The title is hyperlinked to a NCSSM animation of HCl dissolving in water.

32. Weak Electrolytes
WORD DESCRIPTION
Partially breaks apart into its ions
Are poor conductors of electricity
Will produce a dim light bulb
Examples of Acids and Bases that are Weak Electrolytes
Weak Acid
HC2H3O2 (Vinegar)
Weak Base
NH3 (Ammonia)
Notice that only some of the ions are separated or dissociated.
PICTURE
= positive ions
= negative ions
The title is hyperlinked to a NCSSM animation showing HA (a weak acid) dissolving in water.

33. What makes a strong acid or a strong base?
Strong electrolytes make strong acids and bases
Strong Acids
HCl - hydrochloric acid
HNO3 - nitric acid
H2SO4 - sulfuric acid
HBr - hydrobromic acid
HI - hydroiodic acid
HClO4 - perchloric acid
Strong Bases
The hydroxides of the
Group I and Group II
LiOH - lithium hydroxide
NaOH - sodium hydroxide
KOH - potassium hydroxide
*Ca(OH)2 - calcium hydroxide
*Sr(OH)2 - strontium hydroxide
*Ba(OH)2 - barium hydroxide
The title is hyperlinked to a NCSSM animation of HCl dissolving in water.

34. pH Concept

35. pH Scale Pouvoir hydrogéne (hydrogen power) Range is 0-14 Acids 0-7
Bases 7-14
Neutral = 7.0

36. Relationships between pH, [H+], and [OH-]
Click on the picture to get to the animation.
As pH increases…
The [H+] (increases or decreases).
The [OH-] (increases or decreases).
The solution becomes more (acidic or basic).

37. Relationships between pH, [H+], and [OH-]
What happens as pH decreases?
As pH decreases…
The [H+] (increases or decreases).
The [OH-] (increases or decreases).
The solution becomes more (acidic or basic).

38. The pH Scale The value of pH is unitless.
Solutions with a pH less than 7 are acidic and solutions greater than 7 are basic.
If a solution is equal to 7 it is neutral.
Here is a typical pH scale.

39. pH of Common Substances

40. pH Calculations Given Solving for Formula to Use [H+] pH
pH = - log[H+]
[OH-]
pOH
pOH = - log[OH-]
[H+] is the concentration of H+ ions, in mol/L.

41. Logarithms Use your calculator!
If you have a log button, you’re all set.
Each calculator can have its
own method for entering logs.
If you don’t know what to do
your calculator manual
should give examples.
E -2
9 - 43

42. Logarithms If your calculator has a ln button - Don’t use it.
Its for taking natural logs.
This is different than base 10.
E -2
9 - 44

43. Calculating pH
If [H+] is written in scientific notation and has a coefficient of 1, then the pH of the solution equals the absolute value of the exponent
Ex x 10-4 M
pH = 4.0

44. Calculating pH Problem 1: If [H+] = 3.40 x 10-5 M, what is the pH?
Given Unknown Equation
[H+] = 3.40 x 10-5 M pH pH = - log[H+]
Solve:
pH = -log (3.40 x 10-5)
pH = 4.47

45. Calculating pH Problem 2: If [H+] = 1 X 10-10, what is the pH?
Given Unknown Equation
[H+] = 1 X pH pH = - log[H+]
Solve:
pH = - log 1 X 10-10
pH = - (- 10)
pH = 10

46. Calculating pH Problem 3: If [H+] = 1.8 X 10-5, what is the pH?
Given Unknown Equation
[H+] = 1.8 X pH pH = - log[H+]
Solve:
pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74

47. Calculating pOH Solve: Problem 1: pOH = -log (2.3 x 10-12) pOH = 11.6
If [OH-] = 2.30 x M, what is the pOH?
Given Unknown Equation
[OH-] = 2.30 x M pOH pOH = - log[OH-]
Solve:
pOH = -log (2.3 x 10-12)
pOH = 11.6

48. Calculating pOH
If [OH-] is written in scientific notation and has a coefficient of 1, then the pOH of the solution equals the absolute value of the exponent
Problem 2:
If [OH-] = 1.0 x 10-9 M, what is the pH?
pOH = 9.0

49. What’s in a glass of water?
distilled

50. Distilled H2O at the Molecular Level
What’s in a glass of distilled water?
Water Molecules (H2O)
Hydronium Ions (H3O+)
Hydroxide Ions (OH-)
What’s happens in the glass of water?
H2O + H2O ⇆ H3O+ + OH-
This is called the self-ionization of water.

51. pH + pOH = 14 Water Water ionizes- falls apart into ions.
H2O ® H+ + OH-.
Only a small amount.
[H+ ] = [OH-] = 1 x 10-7M
A neutral solution.
In water Kw = [H+ ] x [OH-] = 1 x 10-14
Kw is called the ion product constant.
pH + pOH = 14
Amphoteric
a molecule or ion that can react as an acid as well as a base
Ex: H2O, NH3

52. Calculating pOH from pH
Problem 1:
If the pH is 3.25, what is the pOH?
Given
pH = 3.25
Unknown
pOH ?
Equation
pH + pOH = 14
Substitute and solve :
pOH = 14
(- 3.25) +pOH = 14 (- 3.25)
pOH = 10.8

53. Calculating pH from pOH
Problem 2:
What is the pH of a solution if [OH-] = 4.0 x M?
Given
[OH-] = 4.0 x M
Unknown
pH?
Equation
pH + pOH = 14
Step 1: Find pOH
pOH = -log [OH]
pOH= -log[4.0 x ] = 10.4
Step 2: Calculate pH
pH + pOH= 14; pH = 14 – 10.4
pH = 3.6

54. Looking at the Math
Given
Solving for
Formula to Use pH
[H+]
pOH
[OH-]

55. Calculating [H+] from pH
If the pH of Coke is 3.12, [H+] = ??? [
H+] = = 7.6 x 10-4 M
*** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button
Known
pH = 3.12
Unknown
[H+] ?
Analysis
[H+] = 10 -pH
Substitute and solve :

56. Calculating [H+] from pH
The pH of an unknown solution is What is its [H+]?
Known
pH = 6.00
Unknown
[H+] ?
Analysis
[H+] = 10 -pH
Substitute and solve :
[H+] = 1x M

57. Calculating [H+] from pH
A solution has a pH of What is the Molarity of hydrogen ions in the solution?
Known
pH = 8.5
Unknown
[H+] ?
Analysis
[H+] = 10 -pH
Substitute and solve :
[H+] =
3.16 X M

58. Types of Chemical Reactions
Solution Reactions are classified as follows:
Precipitation reactions
When two solutions are mixed, an insoluble solid forms
B. Acid-Base reactions
A soluble hydroxide and a soluble acid react to form water and a salt
C. Oxidation-Reduction reactions (redox rxns)
Reactions in which one or more electrons are transferred

59. A. Precipitation Reactions
When two solutions are mixed, an insoluble substance (solid) sometimes forms and separates from the solution. Such a reaction is called a precipitation reaction, and the solid that forms is called a precipitate.
Example:
K2CrO4(aq) Ba(NO3)2(aq)  KNO3 (aq) BaCrO4 (s)
(Soluble) (Soluble) (Soluble) (Precipitate)
* Double replacement reactions

60. A. Precipitation Reactions
Mixing KI (aq) and Pb(NO3)2 (aq) leading to precipitation of PbI2

61. B. Acid-Base Reactions or Neutralization Reactions
acid + base  water + salt
HBr(aq) + NaOH(aq) 
H2SO4(aq) + KOH(aq) 
H3PO4(aq) + Ba(OH)2(aq) 
H2O + NaBr
H2O + K2SO4
H2O + Ba3(PO4)2
* Double replacement reactions

62. C. Oxidation-Reduction Reactions
In oxidation-reduction (redox) reaction, one species loses electrons while another species gains electrons.

63. C. Oxidation-Reduction Reactions
Oxidation Numbers
The concept of oxidation numbers is a simple way of keeping track of electrons in a reaction.
The oxidation number (or oxidation state) of an atom in a substance is the actual charge of the atom.
It is hypothetical charge assigned to the atom in the substance by simple rules.

64. C. Oxidation-Reduction Reactions

65. C. Oxidation-Reduction Reactions
Loss of 2 e-1 oxidation
Gain of 2 e-1 reduction
* NOT Double replacement reactions

66. Types of Reactions Tips Always ends with a salt and water
Check to see if one of the product is insoluble in water and look for (s) solid product
Look for transitional metals that are able to change their oxidation number

67. Types of Reactions
The reaction of HNO3 (aq) + KOH(aq) → KNO3 (aq) + H2 O(l) is best classified as a(n) a) acid-base neutralization reaction b) oxidation-reduction reaction c) precipitation reaction *

68. Na3 PO4 (aq) + 3 AgNO3 (aq) → Ag 3PO4 (s) + 3 NaNO3 (aq)
Types of Reactions
The reaction of
Na3 PO4 (aq) + 3 AgNO3 (aq) → Ag 3PO4 (s) + 3 NaNO3 (aq)
is best classified as a(n)
a) acid-base neutralization reaction
b) oxidation-reduction reaction
c) precipitation reaction *

69. Types of Reactions The reaction of is best classified as a(n)
C2 H12 O6 (s) + 6 O2 (g) → 6 CO2 (g) + 6 H2O(l)
is best classified as a(n)
a) acid-base neutralization reaction
b) precipitation reaction
c) oxidation-reduction reaction *