1. Chapter 15 Acid-Base Equilibria
CHEMISTRY
Chapter 15 Acid-Base Equilibria

2. Acids and Bases Arrhenius’ Definition:
Acids - are substances that produce hydrogen ions (protons or H+) in solution.
Bases - are substances that produce hydroxide ions in solution.
Strong Acids and Strong Bases – totally ionize in solution
Weak Acids and Weak Bases – partially ionize in solution

3. Acid Dissociation in Water
General Rxn. when Acid dissolves in H2O
HCl + H2O H3O Cl acid base conj. Acid conj. base

4. Properties of H2O @ 25 oC H2O (l) D H+ (aq) + OH- (aq)
Neutral, can act as an acid and a base
Kw = [H+][OH-] = 1.0 x oC
Kw = water dissociation constant

5. Acidity vs. Basicity If [H+] >[OH-] , solution is acidic
If [H+] <[OH-] , solution is basic
The term pX = -log [concentration of X]
So: pH = -log [concentration of H+]
pOH = -log [concentration of OH-]
pH = power of hydrogen; the power of H to which 10 is raised

6. pH pKw = [-log H+] + [-log OH-]= - [log 1.0 x 10-14] = 14
Kw = [H+][OH-] = 1.0 x 25 oC
pKw = [-log H+] + [-log OH-]= - [log 1.0 x 10-14] = 14
pH = -log [H+]
pOH = -log [OH-]

7. Properties of H2O @ 25 oC pKw = - [log 1.0 x 10-14] = 14 pH + pOH = 14

8. Things to Remember pKw = - [log 1.0 x 10-14] = 14 @ 25 oC
pH + pOH = 14
pH <7 ; acidic
pH > 7; basic
pH is between

9. Broensted-Lowry’s Definition
Acid – is a proton (H+) donor.
Base – is a proton (H+) acceptor.
* Broensted-Lowry Definition is more general
It even applies to bases that have no –OH such as NH3.

10. Terminologies H+ = proton OH- = hydroxide ion H3O+ = hydronium ion
Conjugate base –acid minus proton
Conjugate acid – base plus proton

11. More Terminologies Conjugate acid-base pair
Consists of 2 substances related to each other by the donation and acceptance of a single proton (H+).
Acid Dissociation Constant (Ka)

12. Acid Equilibrium Equilibrium Expression for the reaction:
HCl + H2O H3O Cl acid base conj. Acid conj. Base
Ka = [H3O+ ][Cl- ] = [H+][Cl-] [HCl][H2O] [HCl]

13. Equations pH = - log [H+] pOH = - log [OH-] [H+] = 10 – pH
[OH-] = 10 - pOH
Kw = 10 - pKw
pKw = pH + pOH
pKw = - log [Kw]
Pw = [H+][OH-]

14. Sample Problem At 40 oC, a solution has Kw = 2.916 x 10-14; pH = 7.51
Calculate the following:
A. pOH of the solution
B. hydrogen ion concentration [H+]
C. hydroxide ion concentration [OH-]
D. pKw
E. Is the solution acidic basic or neutral?

15. Equilibrium K = [H3O+ ][Cl- ] = [H+][Cl-] [HCl][H2O] [HCl]
H2O removed from top and bottom since H3O+ is simply H+ dissolved in water.
Remember: Keq = [products] [reactants]

16. Terminologies H+ = proton OH- = hydroxide ion H3O+ = hydronium ion
Conjugate base –acid minus proton
Conjugate acid – base plus proton

17. Problems on Acid Dissociation
Write the simple dissociation reaction for each of the following acids. Omit water.
A.) HNO3
B.) CH3COOH (acetic acid)
C.) NH4+
D.) [Al(H2O)3]3+

18. Acid Strength
Strength of acid is given by the equilibrium position of the dissociation reaction:
HA (aq) H2O (l) H3O A-
Strong acid – totally ionized and equilibrium lies far to the right
Weak acid – only partially ionized and equilibrium lies far to the left

19. Strong Acid vs. Weak Acid
Strong Acid – yields a weak conjugate base (one that has weak affinity for proton; weaker than H2O)
Weak Acid – yields a strong conjugate base (one that has strong affinity for proton; stronger than H2O)

20. Comparison Property Strong Acid Weak Acid Ka value Large Ka Small Ka
Equil. Position
Far to the right
Far to the Left
Equil. Concn
[H+] = [HA]0
[H+] << [HA]0
Conj. Base Strength vs H2O
A- much weaker base than H2O
A- much stronger base than H2O

21. Please Note!
Tuesday’s experiment is Experiment 29: Choice I.

22. Sample Problems
Given [OH-] = 1.0 x M, calculate pH. Is the solution basic, acidic or neutral?
Given [H+] = 4.30 x 10-6 M, calculate pH. Is the solution basic, acidic or neutral?

23. Strong Acids and Bases
If the molarity of the acid or base is less than 10-6 M then the autoionization of water needs to be taken into account. In other words, water is the primary source of H+ and OH-, so the pH would be neutral.

24. Acids and Bases: A Brief Review
Acids: taste sour and cause dyes to change color.
Bases: taste bitter and feel soapy.
Arrhenius: acids increase [H+]; bases increase [OH-] in solution.
Arrhenius: acid + base  salt + water.
Problem: the definition confines us to aqueous solution.

25. Brønsted-Lowry Acids and Bases
Conjugate Acid-Base Pairs
Whatever is left of the acid after the proton is donated is called its conjugate base.
Similarly, whatever remains of the base after it accepts a proton is called a conjugate acid.
Consider
After HA (acid) loses its proton it is converted into A- (base). Therefore HA and A- are conjugate acid-base pairs.
After H2O (base) gains a proton it is converted into H3O+ (acid). Therefore, H2O and H3O+ are conjugate acid-base pairs.
Conjugate acid-base pairs differ by only one proton.

26. Brønsted-Lowry Acids and Bases
Relative Strengths of Acids and Bases
The stronger the acid, the weaker the conjugate base.
H+ is the strongest acid that can exist in equilibrium in aqueous solution.
OH- is the strongest base that can exist in equilibrium in aqueous solution.

28. Brønsted-Lowry Acids and Bases
Relative Strengths of Acids and Bases
The conjugate base of a strong acid (e.g. Cl-) has negligible acid-base properties.
Similarly, the conjugate acid of a strong base has negligible acid-base properties.

29. The Autoionization of Water
The Ion Product of Water
In pure water the following equilibrium is established
at 25 C
The above is called the autoionization of water.

30. The pH Scale In most solutions [H+(aq)] is quite small. We define
In neutral water at 25 C, pH = pOH = 7.00.
In acidic solutions, [H+] > 1.0  10-7, so pH < 7.00.
In basic solutions, [H+] < 1.0  10-7, so pH > 7.00.
The higher the pH, the lower the pOH, the more basic the solution.

31. The pH Scale Most pH and pOH values fall between 0 and 14.
There are no theoretical limits on the values of pH or pOH. (e.g. pH of 2.0 M HCl is )

33. The pH Scale Other “p” Scales In general for a number X,
For example, pKw = -log Kw.

34. The pH Scale Measuring pH
Most accurate method to measure pH is to use a pH meter.
However, certain dyes change color as pH changes. These are indicators.
Indicators are less precise than pH meters.
Many indicators do not have a sharp color change as a function of pH.
Most indicators tend to be red in more acidic solutions.

35. The pH Scale

36. Strong Acids and Bases Strong Acids
The strongest common acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4.
are strong electrolytes.
All strong acids ionize completely in solution:

37. Strong Acids and Bases Strong Acids
The strongest common acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4.
Strong acids are strong electrolytes.
All strong acids ionize completely in solution:
HNO3(aq) + H2O(l)  H3O+(aq) + NO3-(aq)
Since H+ and H3O+ are used interchangeably, we write
HNO3(aq)  H+(aq) + NO3-(aq)

38. Strong Acids and Bases Strong Acids
In solutions the strong acid is usually the only source of H+. (If the molarity of the acid is less than 10-6 M then the autoionization of water needs to be taken into account.)
Therefore, the pH of the solution is the initial molarity of the acid.
Strong Bases
Most ionic hydroxides are strong bases (e.g. NaOH, KOH, and Ca(OH)2).

39. Strong Acids and Bases
If the molarity of the acid or base is less than 10-6 M then the autoionization of water needs to be taken into account. In other words, water is the primary source of H+ and OH-, so the pH would be neutral.

40. Strong Acids and Bases Strong Bases
Strong bases are strong electrolytes and dissociate completely in solution.
The pOH (and hence pH) of a strong base is given by the initial molarity of the base. Be careful of stoichiometry.
In order for a hydroxide to be a base, it must be soluble.
Bases do not have to contain the OH- ion:
O2-(aq) + H2O(l)  2OH-(aq)
H-(aq) + H2O(l)  H2(g) + OH-(aq)
N3-(aq) + H2O(l)  NH3(aq) + 3OH-(aq)

41. pH of Strong Acids and Bases
The pH (and hence pOH) of a strong acid is given by the initial molarity of the acid.
The pOH (and hence pH) of a strong base is given by the initial molarity of the base.
Be careful of stoichiometric ratios!

42. Please Note!
Tuesday’s experiment is Experiment 29: Choice I.

43. Bronsted-Lowry Acids and Bases
Bronsted-Lowry acids – compounds that donate a proton (H+)
Bronsted-Lowry Bases – compounds that accept a proton (H+)
Note that Bronsted-Lowry bases need not have the –OH group on the formula

44. Weak Acids Weak acids are only partially ionized in solution.
There is a mixture of ions and unionized acid in solution.
Therefore, weak acids are in equilibrium:

45. Weak Acids Ka is the acid dissociation constant.
Note [H2O] is omitted from the Ka expression. (H2O is a pure liquid.)
The larger the Ka the stronger the acid (i.e. the more ions are present at equilibrium relative to unionized molecules).
If Ka >> 1, then the acid is completely ionized and the acid is a strong acid.

46. NOTE
For Weak Acids and Weak Bases:
USE ICE to determine H+, OH-, pH and pOH.!

47. Sample A Problem
A solution of 0.10 M formic acid (HCOOH) has a pH of 2.38 at 25 oC.
A. Calculate Ka for formic acid at this temperature.
B. What percent of this solution is ionized?

48. Important Reminder Please Take Note:
Kw = 1.0 x is only true at 25 oC
Therefore, pH + pOH = 14 is also true ONLY at 25 oC
If the temperature is not 25 oC, then Kw will be equal to something else and pKw will not be equal to 14.

49. Sample Problem
At the freezing point of water which is 0 oC, Kw = 1.2 x
Calculate [H+] and [OH-] for a neutral solution at this temperature.

50. Sample Problem The Ka of acetic acid is 1.8 x 10-5.
A. Calculate the pH of a 0.30 M solution of CH3COOH.
B. Calculate OH- and pOH.
C. Calculate Kb.
Calculate % ionization.

51. A Simple Trick
Use of approximation: eliminates the difficulty of quadratic equations.
Approximation is Valid if:
X_______ x < 5 %
[Initial Concn.]

52. Relationship between Ka and Kb
Ka x Kb = 1.0 x only at 25 oC.

53. pH of polyprotic acids Treat polyprotic acids as separate steps!
#1. H2A (aq) D H+ (aq) + HA- (aq) Ka1
# 2. HA- (aq) D H+ (aq) + A-2 (aq) Ka2
Initial [H+] in Step 2 is Equil. [H+] from Step 1.
Total [H+] = SUM from Steps 1 & 2

54. HOMEWORK
What is the pH of a 1.00 M solution of tartaric acid, H2C4H4O6 (aq.) at 25.0 oC?
Answer: pH = 1.49

55. Sample Problem
The Ka of acetic acid is 1.8 x Calculate the Kb of of CH3COOH.

56. Sample Problem
The Ka of ammonia is 1.8 x Calculate the pH of a 0.15 M solution of NH3.

57. Sample Problem
Calculate the concentration of an aqueous solution of NaOH that has a pH of

58. HOMEWORK
What is the pH of a 1.00 M solution of tartaric acid, H2C4H4O6 (aq.) at 25.0 oC?
Answer: pH = 1.49

59. Weak Acids Calculating Ka from pH
Weak acids are simply equilibrium calculations.
The pH gives the equilibrium concentration of H+.
Using Ka, the concentration of H+ (and hence the pH) can be calculated.
Write the balanced chemical equation clearly showing the equilibrium.
Write the equilibrium expression. Find the value for Ka.
Write down the initial and equilibrium concentrations for everything except pure water. We usually assume that the change in concentration of H+ is x.

60. Weak Acids Calculating Ka from pH
Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary.
Using Ka to Calculate pH
Percent ionization is another method to assess acid strength.
For the reaction

61. Weak Acids Using Ka to Calculate pH
Percent ionization relates the equilibrium H+ concentration, [H+]eqm, to the initial HA concentration, [HA]0.
The higher percent ionization, the stronger the acid.
Percent ionization of a weak acid decreases as the molarity of the solution increases.
For acetic acid, 0.05 M solution is 2.0 % ionized whereas a 0.15 M solution is 1.0 % ionized.

63. Weak Acids Polyprotic Acids
Polyprotic acids have more than one ionizable proton.
The protons are removed in steps not all at once:
It is always easier to remove the first proton in a polyprotic acid than the second.
Therefore, Ka1 > Ka2 > Ka3 etc.

64. Weak Acids
Polyprotic Acids

65. Weak Bases Weak bases remove protons from substances.
There is an equilibrium between the base and the resulting ions:
Example:
The base dissociation constant, Kb, is defined as

66. Weak Bases Types of Weak Bases
Bases generally have lone pairs or negative charges in order to attack protons.
Most neutral weak bases contain nitrogen.
Amines are related to ammonia and have one or more N-H bonds replaced with N-C bonds (e.g., CH3NH2 is methylamine).
Anions of weak acids are also weak bases. Example: OCl- is the conjugate base of HOCl (weak acid):

67. Relationship Between Ka and Kb
We need to quantify the relationship between strength of acid and conjugate base.
When two reactions are added to give a third, the equilibrium constant for the third reaction is the product of the equilibrium constants for the first two:
Reaction 1 + reaction 2 = reaction 3
has

68. Relationship Between Ka and Kb
For a conjugate acid-base pair
Therefore, the larger the Ka, the smaller the Kb. That is, the stronger the acid, the weaker the conjugate base.
Taking negative logarithms:

69. Acid-Base Properties of Salt Solutions
Nearly all salts are strong electrolytes.
Therefore, salts exist entirely of ions in solution.
Acid-base properties of salts are a consequence of the reaction of their ions in solution.
The reaction in which ions produce H+ or OH- in water is called hydrolysis.
Anions from weak acids are basic.
Anions from strong acids are neutral.

70. Acid-Base Properties of Salt Solutions
An Anion’s Ability to React with Water
Anions, X-, can be considered conjugate bases from acids, HX.
If X- comes from a strong acid, then it is neutral.
If X- comes from a weak acid, then
The pH of the solution can be calculated using equilibrium!

71. Acid-Base Properties of Salt Solutions
An Cation’s Ability to React with Water
Polyatomic cations with ionizable protons can be considered conjugate acids of weak bases.
Some metal ions react in solution to lower pH.
Combined Effect of Cation and Anion in Solution
An anion from a strong acid has no acid-base properties.
An anion that is the conjugate base of a weak acid will cause an increase in pH.

72. Sample Problem
A solution of NH3 in water has a pH of What is the initial molarity of the solution?

73. Other Weak Bases Amines ex. Methylamine (CH3NH2) carbonate ion (CO32-)
hypochlorite ion (ClO-1)

74. Weak Bases Also use ICE! Calculation is the same as for weak acids!
Main difference is that you get [OH-] and pOH first.

75. Effects of Salts on pH
Conjugate bases of strong acids have no effect on pH.
Conjugate acids of strong bases have no effect on pH.
Conjugate bases of weak acids increase pH (more basic).
Ex. F- (aq) H2O(l) D HF (aq) OH- (aq)
Conjugate acids of weak bases decrease pH (more acidic).
NH4+(aq) H2O (l) D NH3 (aq) H3O+ (aq)

76. Relationship Between Ka and Kb

77. Acid-Base Properties of Salt Solutions
Combined Effect of Cation and Anion in Solution
A cation that is the conjugate acid of a weak base will cause a decrease in the pH of the solution.
Metal ions will cause a decrease in pH except for the alkali metals (Grp. I) and alkaline earth metals.(Grp.II)
When a solution contains both cations and anions from weak acids and bases, use Ka and Kb to determine the final pH of the solution.

78. Sample Problem
Determine whether the resulting solution in water will be acidic, basic or neutral.
A. K+ClO3-
B. Na+CH3COO-
C. Na2HPO Ka for HPO4- = 4.2 x 10-13
D. NH4+Cl-

79. Sample Problem
Predict whether the potassium salt of citric acid (K2+HC6H5O7-) will form an acidic, basic or neutral solution in water.

80. Weak Acids Polyprotic Acids
Polyprotic acids have more than one ionizable proton.
The protons are removed in steps not all at once:
It is always easier to remove the first proton in a polyprotic acid than the second.
Therefore, Ka1 > Ka2 > Ka3 etc.

81. Weak Acids
Polyprotic Acids

82. Sample Problem
The solubility of CO2 in pure water at 25 oC and 0.1 atm is M. The common practice is to assume that all of the dissolved CO2 is in the form of carbonic acid (H2CO3), which is produced by the reaction between the CO2 and H2O.
What is the pH of a M solution of H2CO3?
Ka1 = 4.3 x 10-7
Ka2 = 5.6 x

83. Answer
pH = 4.4
x1 = 4.0 x 10-5 M
[CO3-] = x M

84. Sample Problem
Calculate the pH and concentration of oxalate ion (C2O42-), in a M solution of oxalic acid (H2C2O4)

85. Answer
pH = 1.8
[oxalate] = 6.4 x 10-5 M

86. Acid-Base Behavior and Chemical Structure
Factors that Affect Acid Strength
Consider H-X. For this substance to be an acid we need:
H-X bond to be polar with H+ and X- (if X is a metal then the bond polarity is H-, X+ and the substance is a base),
the H-X bond must be weak enough to be broken,
the conjugate base, X-, must be stable.

87. Acid-Base Behavior and Chemical Structure
Binary Acids
Acid strength increases across a period and down a group.
Conversely, base strength decreases across a period and down a group.
HF is a weak acid because the bond energy is high.
The electronegativity difference between C and H is so small that the C-H bond is non-polar and CH4 is neither an acid nor a base.

88. Acid-Base Behavior and Chemical Structure
Binary Acids

89. Acid-Base Behavior and Chemical Structure
Oxyacids
Oxyacids contain O-H bonds.
All oxyacids have the general structure Y-O-H.
The strength of the acid depends on Y and the atoms attached to Y.
If Y is a metal (low electronegativity), then the substances are bases.
If Y has intermediate electronegativity (e.g. I, EN = 2.5), the electrons are between Y and O and the substance is a weak oxyacid.

90. Acid-Base Behavior and Chemical Structure
Oxyacids
If Y has a large electronegativity (e.g. Cl, EN = 3.0), the electrons are located closer to Y than O and the O-H bond is polarized to lose H+.
The number of O atoms attached to Y increase the O-H bond polarity and the strength of the acid increases (e.g. HOCl is a weaker acid than HClO2 which is weaker than HClO3 which is weaker than HClO4 which is a strong acid).

91. Acid-Base Behavior and Chemical Structure
Oxyacids

92. Acid-Base Behavior and Chemical Structure
Carboxylic Acids
Carboxylic acids all contain the COOH group.
All carboxylic acids are weak acids.
When the carboxylic acid loses a proton, it generate the carboxylate anion, COO-.

93. Lewis Acids and Bases Brønsted-Lowry acid is a proton donor.
Focusing on electrons: a Brønsted-Lowry acid can be considered as an electron pair acceptor.
Lewis acid: electron pair acceptor.
Lewis base: electron pair donor.
Note: Lewis acids and bases do not need to contain protons.
Therefore, the Lewis definition is the most general definition of acids and bases.

94. Lewis Acids and Bases
Lewis acids generally have an incomplete octet (e.g. BF3).
Transition metal ions are generally Lewis acids.
Lewis acids must have a vacant orbital (into which the electron pairs can be donated).
Compounds with p-bonds can act as Lewis acids:
H2O(l) + CO2(g)  H2CO3(aq)

95. Lewis Acids and Bases Hydrolysis of Metal Ions
Metal ions are positively charged and attract water molecules (via the lone pairs on O).
The higher the charge, the smaller the metal ion and the stronger the M-OH2 interaction.
Hydrated metal ions act as acids:
The pH increases as the size of the ion increases (e.g. Ca2+ vs. Zn2+) and as the charge increases (Na+ vs. Ca2+ and Zn2+ vs. Al3+).

96. Lewis Acids and Bases
Hydrolysis of Metal Ions

97. End of Chapter 16 Acid-Base Equilibria

98. Problem 1
Give the conjugate base of the following Bronsted-Lowry acids:
H2SO3
H2AsO4-
NH4+

99. Problem 2 By what factor does [H+] change for a pH change of:
A units
B units

100. Problem 3 Calculate [OH-] and pH for:
A.) 1.5 x 10-3 M Sr(OH)2. Sr(OH)2 is a strong base.
B.) a solution formed by adding 10 mL of M HBr to 20.0 mL of M HCl

101. Problem
Calculate the pH of a solution made by adding grams of NaH in enough water to make 2.5 L of solution

102. Problem
Write the ionization and equilibrium expressions for HBrO2.

103. Problem
A particular sample of vinegar has a pH of Assuming acetic acid is the only acid in the vinegar, find the initial concentration of acetic acid in the vinegar.

104. Problem
The acid dissociation constant for benzoic acid (HC7H5O2) is 6.3 x Calculate the equilibrium concentrations of H3O+, C7H5O2- and HC7H5O2 if the initial concentration of HC7H5O2 is M.

105. Problem
Calculate the pH of M pyridine (C5H5N). Kb for pyridine is 1.7 x 10-9.

106. Problem
A M solution of a weak acid, HA is 9.4% ionized. Using this information, calculate [H+], [A-], [HA] and Ka for HA.

107. Problem
An unknown salt is either NaF, NaCl, or NaOCl. When 0.05 mole of the salt is dissolved in water to form L of solution, the pH of the solution is What is the identity of the salt?

108. Problem
Write the chemical equation and the Kb expression for the ionization of the following bases in aqueous solution:
A. Dimethylamine (CH3)2NH
B. Formate ion (HCOO-)
C. Carbonate ion (CO32-)

109. Problem
Calculate the molar concentration of OH- ions in a 0.075M solution of ethylamine.
Kb of C2H5NH2 = 6.4 x 10-4.
Calculate the pH of this solution.

110. Problem
Ka for acetic acid (CH3COOH) is 1.8 x 10-5 while Ka for hypochlorous (HClO) ion is 3.0 x 10-8.
A. Which is the stronger acid?
B. Which is the stronger conjugate base? Acetate ion (CH3COO-) or chlorous (ClO-) ion?
C. Calculate kb values for CH3COO- and ClO-.

111. Solubility vs. Ksp
Solubility – refers to the quantity that dissolves to form a saturated solution. Unit is gm/liter or moles/liter for molar solubility.
- solubility if affected by temperature
Solubility product constant – is the equilibrium constant for the equilibrium that exists between the ionic solute and its saturated aqueous solution

112. Ksp
Solubility product constant – the equilibrium constant indicating how soluble the product is in water.
Example: CaF2 (s) D Ca2+ (aq) F- (aq)
Ksp = [Ca2+][F-]2

113. Problem 1
Give the ionization equation and Ksp expression for the reaction:
Ag2CrO4 (s) D ? ?

114. Problem 2
The Ksp for CaF2 is 3.9 x at 25 oC. Assuming that CaF2 dissociates completely upon dissolving and that there are no other important equilibria affecting its solubility:
a. calculate the solubility of CaF2 in moles per liter.
b. calculate the solubility of CaF2 in grams per liter.

115. Problem 3
The Ksp for LaF3 is 2.0 x in What is the solubility of LaF3 in water in moles per liter?
What is the solubility of LaF3 in water in grams per liter?

116. Answer
A x 10-6 M

117. Factors Affecting Solubility
Common-Ion Effect
Concentration

118. Problem 4
Calculate the molar solubility of CaF2 at 25 oC in a solution that is:
A M in Ca(NO3)2
B M in NaF

119. Precipitation of Ions Remember Q, the reaction quotient?
If: Q > Ksp, prepitations occurs until Q = Ksp Q = Ksp, equilibrium exists (saturated solution) Q < Ksp, solid dissolves until Q = Ksp.

120. Problem 1
A solution contains 1.0 x M Ag+ and 2.0 x 10-2 M Pb2+. When Cl- is added, both AgCl and Ksp precipitate from the solution.
What concentration of Cl- is necessary to begin the precipitation of each salt?
Which salt precipitates first?

121. Insoluble Chlorides
Of the common metals ions, only Ag+, Hg2 2+, Pb 2+ form insoluble chlorides.

122. Qualitative Analysis
Order of separation of ions: Cl-  S2-  (OH)-  (PO4) 3-  NH4+
1st step: add 6M HCl
2nd step: add H2S and 0.20 M HCl
3rd step: add (NH4)2S
4th step: add (NH4)2HPO4 and NH3