1. Chapter 4 Type of Chemical Reactions and Solution Stoichiometric
Water, Nature of aqueous solutions, types of electrolytes, dilution.
Types of chemical reactions: precipitation, acid-base and oxidation reactions.
Stoichiometry of reactions and balancing the chemical equations.

2. Water is the dissolving medium, or solvent.
Aqueous Solutions
Water is the dissolving medium, or solvent.

3. Figure 4. 1: (Left) The water molecule is polar
Figure 4.1: (Left) The water molecule is polar. (Right) A space-filling model of the water molecule.

4. Figure 4.2: Polar water molecules interact with the positive and negative ions of a salt assisting in the dissolving process.

5. Some Properties of Water
Water is “bent” or V-shaped.
The O-H bonds are covalent.
Water is a polar molecule.
Hydration occurs when salts dissolve in water.

6. Figure 4.3: (a) The ethanol molecule contains a polar O—H bond similar to those in the water molecule. (b) The polar water molecule interacts strongly with the polar O—H bond in ethanol. This is a case of "like dissolving like."

7. A Solute dissolves in water (or other “solvent”)
changes phase (if different from the solvent)
is present in lesser amount (if the same phase as the solvent)

8. A Solvent retains its phase (if different from the solute)
is present in greater amount (if the same phase as the solute)

9. General Rule for dissolution
Like dissolve like
Polar dissolve polar (water dissolve ethanol)
Non-polar dissolve nonpolar (benzene dissolve fat)

10. Figure 4.5: When solid NaCl dissolves, the Na+ and Cl- ions are randomly dispersed in the water.

11. pure water, sugar solution
Electrolytes
Strong - conduct current efficiently
NaCl, HNO3
Weak - conduct only a small current
vinegar, tap water
Non - no current flows
pure water, sugar solution

12. Figure 4.4: Electrical conductivity of aqueous solutions.

13. hydrochloric and sulfuric acid
Acids
Strong acids - dissociate completely to produce H+ in solution
hydrochloric and sulfuric acid
HCl , H2SO4
Weak acids - dissociate to a slight extent to give H+ in solution
acetic and formic acid
CH3COOH, CH2O

14. Bases Strong bases - react completely with water to give OH ions.
sodium hydroxide
Weak bases - react only slightly with water to give OH ions.
ammonia

15. Figure 4.6: HCl(aq) is completely ionized.

16. Figure 4.7: An aqueous solution of sodium hydroxide.

17. Figure 4.8: Acetic acid (HC2H3O2) exists in water mostly as undissociated molecules. Only a small percentage of the molecules are ionized.

18. Molarity
Molarity (M) = moles of solute per volume of solution in liters:

19. Common Terms of Solution Concentration
Stock - routinely used solutions prepared in concentrated form.
Concentrated - relatively large ratio of solute to solvent. (5.0 M NaCl)
Dilute - relatively small ratio of solute to solvent. (0.01 M NaCl): (MV)initial=(MV)Final

20. Figure 4.10: Steps involved in the preparation of a standard aqueous solution.

21. Figure 4.12: Dilution Procedure (a) A measuring pipet is used to transfer 28.7mL of 17.4 M acetic acid solution to a volumetric flask. (b) Water is added to the flask to the calibration mark. (c) The resulting solution is 1.00 M acetic acid.

22. Practice Example
How many moles are in 18.2 g of CO2?

23. Practice Example Consider the reaction N2 + 3H2 = 2NH3
How many moles of H2 are needed to completely react 56 g of N2?

24. Practice Example
How many grams are in mole of caffeine C8H10N4O2

25. Practice Example
A solution containing Ni2+ is prepared by dissolving g of pure nickel in nitric acid and diluting to 1.00 L. A mL aliquot is then diluted to mL. What is the molarity of the final solution? (Atomic weight: Ni = 58.70).

26. Practice Example
Calculate the number of molecules of vitamin A, C20H30O in 1.5 mg of this compound.

27. Practice Example
What is the mass percent of hydrogen in acetic acid HC2H3O2

28. Types of Solution Reactions
Precipitation reactions
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Acid-base reactions
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
Oxidation-reduction reactions
Fe2O3(s) + Al(s)  Fe(l) + Al2O3(s)

29. Simple Rules for Solubility
1. Most nitrate (NO3) salts are soluble.
2. Most alkali (group 1A) salts and NH4+ are soluble.
3. Most Cl, Br, and I salts are soluble (NOT Ag+, Pb2+, Hg22+)
4. Most sulfate salts are soluble (NOT BaSO4, PbSO4, HgSO4, CaSO4)
5. Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble)
6. Most S2, CO32, CrO42, PO43 salts are only slightly soluble.

30. Figure 4.13: When yellow aqueous potassium chromate is added to a colorless barium nitrate solution, yellow barium chromate precipitates.

31. Describing Reactions in Solution Precipitation
1. Molecular equation (reactants and products as compounds)
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
2. Complete ionic equation (all strong electrolytes shown as ions)
Ag+(aq) + NO3- (aq) + Na+ (aq) + Cl(aq) AgCl(s) + Na+ (aq) + NO3- (aq)

32. Describing Reactions in Solution (continued)
3. Net ionic equation (show only components that actually react)
Ag+(aq) + Cl(aq)  AgCl(s)
Na+ and NO3 are spectator ions.

33. Performing Calculations for Acid-Base Reactions
1. List initial species and predict reaction.
2. Write balanced net ionic reaction.
3. Calculate moles of reactants.
4. Determine limiting reactant.
5. Calculate moles of required reactant/product.
6. Convert to grams or volume, as required.
Remember: n H+ = n OH-
(MV) H+ = (MV) OH-

34. Neutralization Reaction
acid + base salt + water
HCl (aq) + NaOH (aq) NaCl (aq) + H2O
H+ + Cl- + Na+ + OH Na+ + Cl- + H2O
H+ + OH H2O
4.3

35. Key Titration Terms
Titrant - solution of known concentration used in titration
Analyte - substance being analyzed
Equivalence point - enough titrant added to react exactly with the analyte
Endpoint - the indicator changes color so you can tell the equivalence point has been reached.
movie

36. Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg (s) + O2 (g) MgO (s)
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
Reduction half-reaction (gain e-)
O2 + 4e O2-
2Mg + O2 + 4e Mg2+ + 2O2- + 4e-
2Mg + O MgO

38. Redox Reactions
Many practical or everyday examples of redox reactions:
Corrosion of iron (rust formation)
Forest fire
Charcoal grill
Natural gas burning
Batteries
Production of Al metal from Al2O3 (alumina)
Metabolic processes
combustion

39. Rules for Assigning Oxidation States
1. Oxidation state of an atom in an element = 0
2. Oxidation state of monatomic element = charge
3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1)
4. H = +1 in covalent compounds
5. Fluorine = 1 in compounds
6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion

41. Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn Zn2+ + 2e-
Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e Cu
Cu2+ is reduced
Cu2+ is the oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)
Cu Cu2+ + 2e-
Ag+ + 1e Ag
Ag+ is reduced
Ag+ is the oxidizing agent

42. IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2
Oxidation numbers of all the elements in the following ?
F = -1
7x(-1) + ? = 0
I = +7
K2Cr2O7
NaIO3
Na = +1
O = -2
O = -2
K = +1
3x(-2) ? = 0
7x(-2) + 2x(+1) + 2x(?) = 0
I = +5
Cr = +6

43. Balancing by Half-Reaction Method
1. Write separate reduction, oxidation reactions.
2. For each half-reaction:
 Balance elements (except H, O)
 Balance O using H2O
 Balance H using H+
 Balance charge using electrons

44. Balancing by Half-Reaction Method (continued)
3. If necessary, multiply by integer to equalize electron count.
4. Add half-reactions.
5. Check that elements and charges are balanced.

45. Half-Reaction Method - Balancing in Base
1. Balance as in acid.
2. Add OH that equals H+ ions (both sides!)
3. Form water by combining H+, OH.
4. Check elements and charges for balance.

46. Balancing Redox Equations
Example: Balance the following redox reaction:
Cr2O Fe Cr3+ + Fe3+ (acidic soln)
1) Break into half reactions:
Cr2O Cr3+
Fe Fe3+

47. Balancing Redox Equations
2) Balance each half reaction:
Cr2O Cr3+
Cr2O Cr3+
Cr2O Cr H2O
Cr2O H Cr H2O
6 e- + Cr2O H Cr H2O

48. Balancing Redox Equations
2) Balance each half reaction (cont)
Fe Fe3+
Fe Fe e-

49. Balancing Redox Reactions
3) Multiply by integer so e- lost = e- gained
6 e- + Cr2O H Cr H2O
x 6
Fe Fe e-

50. Balancing Redox Reactions
3) Multiply by integer so e- lost = e- gained
6 e- + Cr2O H Cr H2O
6 Fe Fe e-
4) Add both half reactions
Cr2O Fe H Cr Fe H2O

51. Balancing Redox Reactions
5) Check the equation
Cr2O Fe H Cr Fe H2O
2 Cr 2 Cr
7 O 7 O
6 Fe 6 Fe
14 H 14 H

52. Balancing Redox Reactions
Procedure for Basic Solutions:
Divide the equation into 2 incomplete half reactions
one for oxidation
one for reduction

53. Balancing Redox Reactions
Balance each half-reaction:
balance elements except H and O
balance O atoms by adding H2O
balance H atoms by adding H+
add 1 OH- to both sides for every H+ added
combine H+ and OH- on same side to make H2O
cancel the same # of H2O from each side
balance charge by adding e- to side with greater overall + charge
different

54. Balancing Redox Equations
Multiply each half reaction by an integer so that
# e- lost = # e- gained
Add the half reactions together.
Simply where possible by canceling species appearing on both sides of equation
Check the equation
# of atoms
total charge on each side

55. Balancing Redox Reactions
Example: Balance the following redox reaction.
NH3 + ClO Cl2 + N2H4 (basic soln)
1) Break into half reactions:
NH N2H4
ClO Cl2

56. Balancing Redox Reactions
2) Balance each half reaction:
NH N2H4
2 NH N2H4
2 NH N2H H+
2 H2O
+ 2 OH- + 2 OH-
2 NH OH N2H H2O
2 NH OH N2H H2O + 2 e-

57. Balancing Redox Reactions
2) Balance each half reaction:
ClO Cl2
2 ClO Cl2
2 ClO Cl H2O
2 ClO H Cl H2O
+ 4 OH- + 4 OH-
2 ClO H2O Cl H2O + 4 OH-
2 ClO H2O Cl2 + 4 OH-
2 e- + 2 ClO H2O Cl2 + 4 OH-

58. Balancing Redox Reactions
3) Multiply by integer so # e- lost = # e- gained
2 NH OH N2H H2O + 2 e-
2 e- + 2 ClO H2O Cl2 + 4 OH-
4) Add both half reactions
2 NH OH- + 2ClO- + 2 H2O N2H H2O + Cl2 + 4 OH-

59. Balancing Redox Reactions
5) Cancel out common species
2 NH OH ClO- + 2 H2O N2H H2O + Cl2 + 4 OH- 2
6) Check final equation:
2 NH ClO N2H4 + Cl2 + 2 OH-
2 N 2 N
6 H 6 H
2 Cl 2 Cl
2 O 2 O
-2 -2

62. Practice Example In the following the oxidizing agent is:
5H2O2 + 2MnO4- + 6H+  2Mn2+ + 8H2O + 5O2
a. MnO4-
b. H2O2
c. H+
d. Mn2+
e. O2

63. Practice Example Determine the coefficient of Sn in acidic solution
Sn + HNO3  SnO2 + NO2 + H2O 1

64. Practice Example
The sum of the coefficients when they are whole numbers in basic solution:
Bi(OH)3 + SnO22-  Bi + SnO32- 13

65. http://www. chemistrycoach. com/balancing_redox_in_acid

68. QUESTION

69. ANSWER

70. QUESTION

71. ANSWER

72. QUESTION

73. ANSWER

74. QUESTION

75. ANSWER

76. QUESTION

77. ANSWER STUDENT CD: Visualization: Conductivities of Aqueous Solutions
HMCLASS PRESENT: Video: Conductivities of Aqueous Solutions
HMCLASS PREP: Figure 4.4

78. QUESTION

79. ANSWER

80. QUESTION

81. ANSWER

82. QUESTION

83. ANSWER HMCLASS PREP: Figure 4.2
STUDENT CD: Visualization: The Dissolution of a Solid in a Liquid
HMCLASS PRESENT: Animation: The Dissolution of a Solid in a Liquid

84. QUESTION

85. ANSWER

86. QUESTION

87. QUESTION (continued)

88. ANSWER
STUDENT CD: Understanding Concepts: Precipitation Reactions

89. QUESTION

90. ANSWER HMCLASS PREP: Table 4.1
HMCLASS PRESENT: Animation: The Precipitation of CaCO3

91. QUESTION

92. ANSWER

93. QUESTION

94. ANSWER
STUDENT CD: Understanding Concepts: Acid-Base Reactions

95. QUESTION

96. ANSWER

97. QUESTION

98. ANSWER
STUDENT CD: Understanding Concepts: Oxidation-Reduction Reactions

99. QUESTION

100. ANSWER
HMCLASS PREP: Figure 4.19

101. QUESTION

102. ANSWER

103. QUESTION

104. QUESTION (continued)

105. ANSWER