1. Acid-Base Equilibria

2. ATTACK METHOD Identify what you have.
Identify where you are on the titration curve.
When mixing acids and bases, determine which species wins the mole war.
Make sure your answer makes sense.
Prepare by reviewing notes every night.

3. Strong Acid/Strong Base Titration
Excess Base
moles of acid =moles of base
Excess Acid

4. WHAT IS THE DIFFERENCE BETWEEN END POINT AND EQUIVALENCE POINT?

6. Strong Acid/Strong Base Titration
Determine the pH if 50 ml of 0.25 M HCl is mixed with 30 ml of 0.25 M NaOH

7. Strong Acid/Strong Base Titration
Determine the pH when 30.0 ml of M HCl is added to 60 ml of M NaOH.

8. Strong Acid/Strong Base Titration
Determine the pH when 25.0 ml of M HCl is added to 85.0 ml of M NaOH.

9. Strong Acid/Strong Base Titration
Determine the pH if enough M NaOH is added to ml of .500 M HCl to reach the equivalence point.
Determine the volume of NaOH required to reach the equivalence point.

10. Learning Check Strong Acid/Strong Base Titration
What is the pH at the equivalence point for this titration?
Label parts A, B, and C
If 30.0 ml of M NaOH is needed to reach the equivalence point, how many moles of acid were used? B A

11. Weak Acid/Strong Base Titration
Excess base region
Moles of base is greater than acid
Determine the moles of strong base left over . Divide by total volume gives [OH- ] concentration.
All strong base is converted to CB
at the equivalence point
Use Kb = X2
new molarity– x
Hydrolysis region
Moles of acid = moles of base at this point
Buffer Region
pH =pKa
H+ = Ka
Use [H+ ] = Ka [ moles acid ] [ moles base]
You are in in the buffer region when the moles of weak are greater than the moles of strong.

12. Weak Acid/Strong Base Titration
All strong base is converted to CB
Use Kb = X2
new molarity – x

14. HENDERSON-HASSELBALCH EQUATION (Buffers)
Derivative is
H+ = Ka [HB]
[B- ]
If you have a buffer , use the above
equation. Makes the math easier.

15. HENDERSON-HASSELBALCH EQUATION (Buffers)
poH = pkb + log [Acid/Conj. Base]
H+ = Ka [B-]
[HB]

16. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M CH3COOH with M NaOH.
1. Before the addition of any base (initial pH).
2 After the addition of 8.00 mL of M NaOH (buffer region).
3. After the addition of mL of M NaOH
(half-neutralization point).
4. After the addition of mL of M NaOH
(equivalence point).
5. After the addition of mL of M NaOH
(beyond the equivalence point).

17. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M CH3COOH with M NaOH.
1. Before the addition of any base .

18. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M CH3COOH with M NaOH
2. after the addition of 8.00 mL of M NaOH

19. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M CH3COOH
3. after the addition of mL of M NaOH

20. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M CH3COOH
4. after the addition of mL of M NaOH

21. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M CH3COOH
5. After the addition of mL of M NaOH

22. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Determine the pH of a solution when 25.0 ml of 2.00 M HC2H3O2 is titrated halfway to the equivalence point with NaOH

23. Neutralization Reaction Weak Acid - Strong Base 5 Positions on the Titration Curve
Determine the pH of a solution when 25.0 ml of 2.00 M HC2H3O2 is titrated with 1.50 M NaOH to the equivalence point.

24. LEARNING CHECK WEAK ACID STRONG BASE TITRATION
What is the pH at the equivalence point of a weak acid/strong base titration?
When does pKa = pH?
How do you know you have a buffer?
What occurs at the equivalence point?
How do you know you are at the equivalence point?
How do you know you are beyond the equivalence point?
When does H+ = Ka?
Write reaction for a mixture of HF and NaOH when enough NaOH is added to reach the equivalence point.
What equation would you use to solve for pH if you are halfway to the stiochiometric point?

26. Buffer Region Excess Acid Region WEAK BASE STRONG ACID TITRATION
You are in in the buffer region when the moles of weak is greater than strong
Buffer Region
Use [H+ ] = Ka [ moles acid ] [ moles base]
pH = pKa
H+ = Ka
All strong acid is converted to CA ate the equivalence point Use Ka = X2
New molarity– x
Excess Acid
Region
Moles of acid = moles of base
Moles of strong are greater than
weak Determine the moles of acid left over then divide by total volume. Gives H+ concentration

27. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M ammonia with M HCl
1. Before the addition of any acid (initial pH).
2 After the addition of 8.00 mL of M HCl (buffer region).
3. After the addition of mL of M HCl
(half-neutralization point).
4. After the addition of mL of M HCl
(equivalence point).
5. After the addition of mL of M HCl
(beyond the equivalence point).

28. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M ammonia
1. Before the addition of any acid.

29. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M ammonia
2. After the addition of 8.00 mL of M HCl

30. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M ammonia
3. After the addition of mL of M HCl

31. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
4. After the addition of mL of M HCl
Calculate the pH at the following points in the titration of
20.00 mL of M ammonia

32. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Calculate the pH at the following points in the titration of
20.00 mL of M ammonia
5. After the addition of mL of M HCl

33. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Determine the pH of a solution when 25.0 ml of 2.00 M NH3 is titrated halfway to the equivalence point with HCl

34. Neutralization Reaction Weak Base - Strong Acid 5 Positions on the Titration Curve
Determine the pH of a solution when 25.0 ml of 2.00 M NH3 is titrated to the equivalence point with M HCl.

35. LEARNING CHECK WEAK ACID STRONG BASE TITRATION
What is the pH at the equivalence point of a weak acid/strong base titration?
When does pKa = pH?
How do you know you have a buffer?
What occurs at the equivalence point?
How do you know you are at the equivalence point?
How do you know you are beyond the equivalence point?
When does H+ = Ka?
Write reaction for a mixture of NH3 and HCl when enough HCl is added to reach the equivalence point.
What equation would you use to solve for pH if you are halfway to the stiochiometric point?

36. Titration of a Polyprotic Weak Acid

37. Titration of a Polyprotic Strong Acid
2nd equivalence point
1st equivalence point

38. pH Review pH = - log [H+] H+ is really a proton Range is from 0 - 14
If [H+] is high, the solution is acidic; pH < 7
If [H+] is low, the solution is basic or alkaline ; pH > 7

39. The Body and pH Homeostasis of pH is tightly controlled
Extracellular fluid = 7.4
Blood = 7.35 – 7.45
< 6.8 or > 8.0 death occurs
Acidosis (acidemia) below 7.35
Alkalosis (alkalemia) above 7.45

40. Small changes in pH can produce major disturbances
Most enzymes function only with narrow pH ranges
Acid-base balance can also affect electrolytes (Na+, K+, Cl-)
Can also affect hormones

41. Acidosis
Principal effect of acidosis is depression of the CNS through ↓ in synaptic transmission.
Generalized weakness
Deranged CNS function the greatest threat
Severe acidosis causes
Disorientation
coma
death

42. Alkalosis
Alkalosis causes over excitability of the central and peripheral nervous systems.
Numbness
Lightheadedness
It can cause :
Nervousness
muscle spasms or tetany
Convulsions
Loss of consciousness
Death

44. Buffers

45. Preparing a Buffer
A buffer must contain or produce a conjugate acid/base pair.
Example:
HA and A-
HF and F-

46. Preparing a Buffer A buffer can be prepared by
Using a weak acid and ½ the moles of a strong base or less
Using a weak base and ½ the moles of a strong acid or less
A weak acid and the salt of a weak base (ex HF & NaF)
A weak base and the salt of a weak acid (ex NH3 & NH4Cl

47. Buffer Solutions Remember:
The function of a buffer is to resist changes in the pH of a solution.

48. Buffer Solutions How do they resist changes in pH?
ACID USES UP ADDED OH-
BASE USES UP ADDED H+
Equal moles of HF and NaF are used.
HF + H2O F H+
LeChatelier's Principle

49. Why does this work? Preparing a Buffer
Using a weak acid and ½ the moles of a strong base or less
Why does this work?
30.0 ml of M acetic acid and 15.0 ml of
0.200 M sodium hydroxide
R C2H3O2- + OH H C2H3O2 + H2O I C E

50. Buffer Solutions H+ = Ka [HB] [B- ]
If you have a buffer , use the above
equation. Makes the math easier.
You have a buffer if have a conjugate acid base pair or the moles of weak (acid/base) is greater than the strong (acid/base)

51. Buffer Solutions Which of the following mixtures would result
in a buffered solution when 1.0 L of each
of the two solutions are mixed?
0.2 M HNO3 and M NaNO3
0.2 M HNO3 and M HF
0.2 M HNO3 and M NaF
0.2 M HNO3 and M NaOH

52. Buffer Solutions How does adding water to a buffer system
affect the pH?
H+ = Ka [HB]
[B- ]
CONCENTRATION of the acid and conjugate base are not important.
It is the RATIO OF THE NUMBER OF MOLES of each. The volumes cancel You can ignore the volume and use
H+ = Ka HB moles
B- moles

53. Buffer Solutions
Problem: What is the pH of a solution that has [HOAc] = M and [NaOAc-] = M? Ka = 1.8 x 10-5

54. Adding an Acid to a Buffer
What is the pH when 1.00 mL of 1.00 M HCl is added to 1.00 L of buffer that has [HOAc] = M and [OAc-] = M
HINT: I would use moles.

55. Learning Check
30.0 ml of M NH3 is mixed with 60.0 ml of Determine the pH.

56. Learning Check
50.0 ml of 1 M HCl is mixed with 100. ml of 1 M NH moles of NaOH is added to the mixture determine the pH.

57. Adding an Base to a Buffer
What is the pH when 1.00 mL of 1.00 M NaOH is added to 1.00 L of buffer that has [HOAc] = M and [NaOAc-] = M .

58. IS IT A BUFFER OR NOT?
50 ml of 2.00 M NaOH is titrated with 60.0 ml of 0.1 M HC2H3O2

59. Buffer Solutions
How many moles of HCl must be added to 1.0 L of 1.0 M sodium acetate to produce a solution buffered at
pH= pKa
pH =4.20

60. Buffer Solutions
Consider a solution that contains both ammonia and ammonium chloride. Calculate the ratio of [NH3 ]/[NH4 Cl ] at a pH of 5.23.

61. Preparing a Buffer Solution
Buffer prepared from
HCO3- weak acid
CO32- conjugate base
HCO H2O H3O+ + CO32-

62. Preparing a Buffer You want to buffer a solution at pH = 4.30.
This means [H3O+] = x 10-5 M
It is best to choose an acid such that [H3O+] is about equal to Ka (or pH ­ pKa).
—then you get the exact [H3O+] by adjusting the ratio of acid to conjugate base.

63. Preparing a Buffer You want to buffer a solution at pH = 4.30 or
[H3O+] = 5.0 x 10-5 M
POSSIBLE ACIDS Ka
HSO4- / SO x 10-2
HOAc / OAc x 10-5
HCN / CN x 10-10
Best choice is acetic acid / acetate.
Remember pKa = -log Ka
pH = pKa when moles acid = moles of base

64. Preparing a Buffer You want to buffer a solution at pH = 4.30 or
[H3O+] = 5.0 x 10-5 M

65. Preparing a Buffer You want to buffer a solution at pH = 4.30 or
[H3O+] = 5.0 x 10-5 M

66. Preparing a Buffer You want to buffer a solution at pH = 4.30 or
[H3O+] = 5.0 x 10-5 M
Solve for [HOAc]/[OAc-] ratio
= 2.78/ 1
Therefore, if you use mol of NaOAc and mol of HOAc, you will have pH = 4.30.

67. Buffer capacity
The buffer capacity is the measure of the
amount of acid or base that the solution can
absorb without significant change in pH
Depend on how many moles of the conjugate
acid-base pair are present
- for solutions having the same concentration,
the greater the volume  greater buffer
capacity

68. Buffer capacity
To increase the buffering capacity of a buffer system, increase the moles of acid and moles of
base
CONCENTRATION of the acid and conjugate base are
not important. It is the RATIO OF THE NUMBER OF MOLES
of each.

69. Buffer Capacity And Buffer Range
There is a limit to the capacity of a buffer solution to neutralize added acid or base, and this limit is reached before all of one of the buffer components has been consumed.
In general, the more concentrated the buffer components in a solution, the more added acid or base the solution can neutralize.
As a rule, a buffer is most effective if the concentrations of the buffer acid and its conjugate base are equal (pH=pKa).

70. Buffer Capacity And Buffer Range
In this course we assume the buffer system is a small system.

71. Acid-Base Indicators
An acid-base indicator is a weak acid having one color and the conjugate base of the acid having a different color. One of the “colors” may be colorless.
HIn + H2O º H3O+ + In-
color color 2
Acid-base indicators are often used for applications in which a precise pH reading isn’t necessary.
A common indicator used in introductory chemistry laboratories is litmus.

72. 17-3 Acid-Base Indicators
Color of some substances depends on the pH.
HIn + H2O In- + H3O+
In the acid form the color appears to be the acid color.
In the base form the color appears to be the base color.
Intermediate color is seen in between these two states.
The complete color change occurs over about 2 pH units.
Remember….LeChatelier’s Principle

73. Indicator Colors and Ranges
Slide 27 of 45 General Chemistry: Chapter 17 Prentice-Hall © 2007

74. Titration Curve For Strong Acid - Strong Base

75. Acid-Base Titrations/Indicators
Indicators should change color close to the equivalence point.

76. Acid-Base Titrations/Indicators
Phenolphthalein is used in a strong acid/strong base titration the solution turns pink upon the addition of a certain amount of base. Upon the addition of more base the solution turns colorless.
Explain

77. 2012 AP QUESTION 1
(a) Explain how the data in the table above provide evidence that HA is a weak acid rather than a strong acid. (b) Write the balanced net-ionic equation for the reaction that occurs when the solution of NaOH is added to the solution of HA .

78. 2012 AP QUESTION 1
(c) Calculate the number of moles of HA that were titrated.

79. 2012 AP QUESTION 1
(d) Calculate the molar mass of HA .

80. 2012 AP QUESTION 1
(e) Assume that the initial concentration of the HA solution (before any NaOH solution was added) is M. Determine the pH of the initial HA solution.

81. 2012 AP QUESTION 1
(f) Calculate the value of [H3O+] in the solution after 30.0 mL of NaOH solution is added and the total volume of the solution is 80.0 mL.

82. 2006 AP QUESTION 1 FORM B

83. 2006 AP QUESTION 1 FORM B

84. 2006 AP QUESTION 1 FORM B

85. 2006 AP QUESTION 1 FORM B

87. NET IONIC EQUATIONS ACIDS AND BASES
Frustrated Teacher

88. Pre- AP & NET IONIC EQUATIONS
This is a great resource it also comes with a teachers edition with answers. The Ultimate Chemical Equations Handbook by George R Hague, Jr. and Jane D. Smith
I have typed the equations into exam view

89. NET IONIC EQUATIONS Which one(s)? Anhydrides Combustion (Hydrocarbons)
Decomposition (simple)
Double replacement
Redox (Single replacement)
Warm Ups and quizzes

90. ANYDRIDES (ACIDIC & BASIC)
Solid sodium oxide is added to water
Solid lithium oxide is added to water

91. ANYDRIDES (ACIDIC) Carbon dioxide is bubbled through water
Sulfur dioxide is bubbled through water

92. ANYDRIDES (SALT)
Solid sodium oxide is treated with gaseous carbon dioxide
Solid calcium oxide is treated with gaseous sulfur dioxide

93. ANYDRIDES (MISC.)
Dilute nitric acid is added to crystals of pure calcium oxide.
Carbon dioxide gas is bubbled through a solution of concentrated sodium hydroxide.
Carbon dioxide gas is bubbled through a dilute solution of sodium hydrioxde.

94. ANYDRIDES (MISC.)
Gaseous sulfur dioxide is bubble through a dilute solution of potassium hydroxide.

95. GROUP I & CBS Group II IN WATER (form basic solutions)
Solid sodium is added to water.
Solid lithium is added to water.
Solid calcium is added to water

96. Hydroxides
Solid ammonium chloride is added to a solution of sodium hydroxide.
Solid barium carbonate is added to concentrated nitric acid
Solutions of cobalt (II) nitrate and sodium hydroxide are mixed.

97. Trivial Redox (Single Replacement)
Solid zinc is placed in hydrochloric acid
Solid calcium is added to hydrochloric acid
Concentrated hydrochloric acid is added to solid manganese(II) sulfide

98. NON-TRIVIAL REDOX ACIDIC/BASIC
Some essential principles to keep in mind:
• elements in their highest positive oxidation state (same as group #, whether A or B) can ONLY be reduced
• elements in their lowest oxidation state (0 for metals, negative for non-metals, corresponding to distance from noble gases) can ONLY be oxidized
• intermediate oxidation states can go either way!!!
Quizzes are cumlativ

99. NON-TRIVIAL REDOX ACIDIC/BASIC
Some essential principles to keep in mind:
• if the mixture is acidic, H+ should be included as a reactant; water is
one product
• if the mixture is basic, OH- should be included as a reactant; water
is one product
• occasionally the acid anion or base cation may precipitate with a
product ion
Quizzes are cumlativ

100. Non-Trivial Redox oxidizers [remember, oxidizers will become reduced]
MnO4- (in acid) → Mn2+
MnO4- (in neutral or basic) → MnO2
MnO2 (in acid) → Mn2+
Cr2O72- (in acid) → Cr3+
HNO3 (conc.) → NO2
HNO3 (dilute) → NO
H2SO4 (hot, concentrated) → SO2 [if not hot and conc., this acts like HCl or other normal acids]
metal cations → lower charge cations or (rarely) free metals
free halogens → halide ions
reducers [remember, reducers will become oxidized]
halide ions → free halogens
free metals → metal cations
SO32- (or SO2) → SO42-
NO2- → NO3-
free halogens (dil. basic) → hypohalite ions [like XO-]
free halogens (conc. basic) → halate ions [like XO3-]
metal cations → higher charge cations
Quizzes are cumlativ

101. NON-TRIVIAL REDOX
A solution of potassium iodide is added to an acidified solution of potassium dichromate.
A solution of potassium dichromate is aadddede to a solution of sodium hydroxide.
Potassium dichromate solution is added to a solution of sodium sulfite.
Quizzes are cumlativ

102. NON-TRIVIAL REDOX
A solution of potassium iodide is added to an acidified solution of potassium dichromate
Liquid mercury is added to concentrated nitric acid
Potassium permanganate solution is added to concentrated hydrochloric acid
Quizzes are cumlativ

103. Complex Ions – Hydroxides/Ammonia
Concentrated ammonia solution is added to a solution of copper (II) hydroxide.
Excess concentrated ammonia solution is added to a suspension of silver chloride
Excess concentrated potassium hydroxide solution is added to a precipitate of zinc hydroxide
Silver chloride is dissolved in excess ammonia solution.
Quizzes are cumlativ

104. Acid/Base Neutralization
Excess potassium hydroxide solution is added to a solution of sodium hydroxide
Ammonia gas is bubbled into a solution of potassium hydrogen carbonate.
Solutions of ammonia and hydrofluoric acid are mixed
Hydrogen sulfide gas is bubbled through excess potassium hydroxide solution.
Quizzes are cumlativ