1. Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

2. Water, the Common Solvent
4.1
Water, the Common Solvent

3. Water
Aqueous solutions- solutions in which water is the dissolving medium, or solvent.
One of the most valuable properties of water is its ability to dissolve many different substances. (Universal Solvent)
An individual H2O molecule is “bent” or V- shaped, with an HOOOH angle of approximately 105 degrees.
Covalent bonds, but electrons not shared equally. Water is a polar molecule

4. cause a salt to “fall apart” in the water, or to dissolve.
Hydration
Hydration- “positive ends” of the water molecules are attracted to the negatively charged anions and that the “negative ends” are attracted to the positively charged cations.
cause a salt to “fall apart” in the water, or to dissolve.
ionic substances (salts) dissolve in water, they break up into the individual cations and anions.

5. Solubility
Solubility
the most important thing to remember at this point is that when an ionic solid does dissolve in water, the ions become hydrated and are dispersed (move around independently).

6. Nonpolar molecules that dissolve
Water also dissolves many nonionic substances like Ethanol (C2H5OH), for example. (in beer, wine and mixed drinks)
The molecule contains a polar O-H bond like those in water, which makes it very compatible with water.

7. Solubility Rule of Thumb
Many substances do not dissolve in water.
Pure water will not dissolve animal fat
fat molecules are nonpolar and do not interact with polar water molecules.
In general, polar and ionic substances are expected to be more soluble in water than nonpolar substances.
“Like dissolves like”

8. CRITICAL THINKING What if no ionic solids were soluble in water
CRITICAL THINKING What if no ionic solids were soluble in water? How would this affect the way reactions occur in aqueous solutions?

9. Strong and Weak Electrolytes
4.2
Strong and Weak Electrolytes

10. Strong and Weak Electrolytes
a substance, the solute, is dissolved in liquid water, the solvent.
electrical conductivity, its ability to conduct an electric current.
strong electrolytes- solutions conduct current very efficiently
weak electrolytes- solutions conduct only a small current
Nonelectrolytes- no current to flow, and the bulb remains unlit

11. Strong and Weak Electrolytes
sodium chloride readily produces ions in aqueous solution and thus are strong electrolytes.
acetic acid produces relatively few ions when dissolved in water and are weak electrolytes.
Sugar forms virtually no ions when dissolved in water and are nonelectrolytes.

12. (1) soluble salts, (2) strong acids, and (3) strong bases.
Strong Electrolytes
Strong electrolytes are substances that are completely ionized when they are dissolved in water
(1) soluble salts, (2) strong acids, and (3) strong bases.

13. Acids as Electrolytes
Arrhenius
Acidity was first associated with the sour taste of citrus fruits. Acidus- latin for sour
an acid is a substance that produces H+ ions (protons) when it is dissolved in water.

14. Acids as Strong Electrolytes
HCl, HNO3, and H2SO4 are placed in water, virtually every molecule ionizes.
These substances are strong electrolytes and are thus called strong acids.

15. Strong Acids as Strong Electrolytes

16. Strong Bases as Strong Electrolytes
strong bases-soluble ionic compounds containing the hydroxide ion (OH-)
dissolved in water, the cations and OH- ions separate and move independently
bases have a bitter taste and a slippery feel

17. The most common weak electrolytes are weak acids and weak bases.
Weak electrolytes are substances that exhibit a small degree of ionization in water
The most common weak electrolytes are weak acids and weak bases.

18. Types of Hydrogens in acids
main acidic component of vinegar is acetic acid (HC2H3O2).
The formula is written to indicate that two chemically distinct types of hydrogen atoms.

19. Weak Acids as Weak Electrolytes
0.1 mole of HC2H3O2 per liter, for every 100 molecules of HC2H3O2 originally dissolved in water, approximately 99 molecules of HC2H3O2 remain intact
the double arrow indicates the reaction can occur in either direction.

20. Weak Acid Definition
weak electrolyte, it is called a weak acid. Any acid, such as acetic acid, that dissociates (ionizes) only to a slight extent in aqueous solutions is called a weak acid.

21. The most common weak base is ammonia (NH3)
Weak Base Definition
The most common weak base is ammonia (NH3)
The solution is basic because OH- ions are produced.
Ammonia is called a weak base because the resulting solution is a weak electrolyte; that is, very few ions are formed.

22. Nonelectrolyte Definition
Nonelectrolytes are substances that dissolve in water but do not produce any ions
An example of a nonelectrolyte is ethanol and table sugar (sucrose)

23. Homework
CH4 HW1: 1, 10, 15, 18, 24, 49

24. The Composition of Solutions
4.3
The Composition of Solutions

25. Chemical reactions often take place when two solutions are mixed.
Stoich in Solutions
Chemical reactions often take place when two solutions are mixed.
To perform stoichiometric calculations
(1) the nature of the reaction, which depends on the exact forms the chemicals take when dissolved,
(2) the amounts of chemicals present in the solutions, usually expressed as concentrations

26. Molarity
molarity (M), which is defined as moles of solute per volume of solution in liters

27. Practice Problem
Calculate the molarity of a solution prepared by dissolving g of solid NaOH in enough water to make 1.50 L of solution.

28. Practice Problem
Calculate the molarity of a solution prepared by dissolving g of gaseous HCl in enough water to make 26.8 mL of solution.

29. Give the concentration of each type of ion in the following solutions:
Practice Problem
Give the concentration of each type of ion in the following solutions:
a M Co(NO3)2
b. 1 M Fe(ClO4)3

30. Calculate the number of moles of Cl- ions in 1. 75 L of 1
Calculate the number of moles of Cl- ions in 1.75 L of 1.0 × 10-3 M ZnCl2

31. To analyze the alcohol content of a certain wine, a chemist needs 1
To analyze the alcohol content of a certain wine, a chemist needs 1.00 L of an aqueous M K2Cr2O7 (potassium dichromate) solution. How much solid K2Cr2O7 must be weighed out to make this solution?

32. Dilution- solutions are often purchased or prepared in concentrated form (called stock solutions). Water is then added to achieve the molarity desired for a particular solution.
Moles of solute after dilution = moles of solute before dilution

33. suppose we need to prepare 500. mL of 1
suppose we need to prepare 500. mL of M acetic acid (HC2H3O2) from a M stock solution of acetic acid.

34. Do First: Pg , 28, 36, 46

35. What volume of 16 M sulfuric acid must be used to prepare 1.5 L of a 0.10-M H2SO4 solution?

36. M1V1 = M2V2
M1 × V1 = mol solute before dilution = mol solute after dilution = M2 × V2

37. Types of Chemical Reactions
4.4
Types of Chemical Reactions

38. Types of Solution Reactions
» Precipitation reactions
» Acid–base reactions
» Oxidation–reduction reactions

39. Precipitation Reactions
4.5
Precipitation Reactions

40. A precipitation reaction also can be called a double displacement reaction.
Precipitation reaction- When two solutions are mixed, an insoluble substance sometimes forms; that is, a solid forms and separates from the solution

44. What happened to the K+ and NO3- ions?
these ions are left dissolved in the solution; KNO3 does not form a solid in water.
How could we isolate the KNO3?

48. Using the solubility rules in Table 4
Using the solubility rules in Table 4.1, predict what will happen when the following pairs of solutions are mixed.
a. KNO3(aq) and BaCl2(aq)
b. Na2SO4(aq) and Pb(NO3)2(aq)
c. KOH(aq) and Fe(NO3)3(aq)

49. Describing Reactions in Solution
4.6
Describing Reactions in Solution

50. formula equation
complete ionic equation
all substances that are strong electrolytes are represented as ions

51. spectator ions- The ions that do not participate directly in the reaction
net ionic equation- only those solution components directly involved in the reaction

53. For each of the following reactions, write the formula equation, the complete ionic equation, and the net ionic equation.
a. Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride precipitate plus aqueous potassium nitrate.
b. Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate to form a precipitate of iron(III) hydroxide and aqueous potassium nitrate.

54. Stoichiometry of Precipitation Reactions
4.7
Stoichiometry of Precipitation Reactions

55. Calculate the mass of solid NaCl that must be added to 1. 50 L of a 0
Calculate the mass of solid NaCl that must be added to 1.50 L of a M AgNO3solution to precipitate all the Ag+ ions in the form of AgCl.

56. Chapter 11 Properties of Solutions

57. 11.1
Solution Composition

58. Molarity or Mass Percent or Mole Fraction or Molality
Solution Composition
Molarity or Mass Percent or Mole Fraction or Molality

60. A solution is prepared by mixing 1. 00 g ethanol (C2H5OH) with 100
A solution is prepared by mixing 1.00 g ethanol (C2H5OH) with g water to give a final volume of 101 mL. Calculate the molarity, mass percent, mole fraction, and molality of ethanol in this solution.

61. A solution is prepared by mixing 1. 00 g ethanol (C2H5OH) with 100
A solution is prepared by mixing 1.00 g ethanol (C2H5OH) with g water to give a final volume of 101 mL. Calculate the molarity, mass percent, mole fraction, and molality of ethanol in this solution.

62. A solution is prepared by mixing 1. 00 g ethanol (C2H5OH) with 100
A solution is prepared by mixing 1.00 g ethanol (C2H5OH) with g water to give a final volume of 101 mL. Calculate the molarity, mass percent, mole fraction, and molality of ethanol in this solution.

63. A solution is prepared by mixing 1. 00 g ethanol (C2H5OH) with 100
A solution is prepared by mixing 1.00 g ethanol (C2H5OH) with g water to give a final volume of 101 mL. Calculate the molarity, mass percent, mole fraction, and molality of ethanol in this solution.

64. Practice Problem
A solution of phosphoric acid was made by dissolving 10.0 g H3PO4 in mL water. The resulting volume was 104 mL. Calculate the density, mole fraction, molarity, and molality of the solution. Assume water has a density of 1.00 g/cm3.

65. Normality

66. The Energies of Solution Formation
11.2
The Energies of Solution Formation

67. Formation of Liquid Solution
1. Separating the solute into its individual components (expanding the solute)
2. Overcoming intermolecular forces in the solvent to make room for the solute
(expanding the solvent)
3. Allowing the solute and solvent to interact to form the solution

68. Enthalpy (heat) of solution (ΔHsoln)
Steps 1 and 2 require energy, since forces must be overcome to expand the solute and solvent. Step 3 usually releases energy.

69. Enthalpy (heat) of solution (ΔHsoln)
Enthalpy (heat) of solution (ΔHsoln)- enthalpy change associated with the formation of the solution

70. pg 439
Example 11.3

71. Factors Affecting Solubility
11.3
Factors Affecting Solubility

72. Structure Effects

73. Henry’s law: C = kP Pressure Effects
Henry’s law states that the amount of a gas dissolved in
a solution is directly proportional to the pressure of the gas above the solution

74. Example 11.4

76. Temperature Effects (for Aqueous Solutions)

77. The Vapor Pressures of Solutions
11.4
The Vapor Pressures of Solutions

79. Vapor Pressure

80. Example 11.5

82. HW3 CH4: 64, 21, 54 CH11: 30, 33, 38, 40, 41, 44, 47, 48, 92,