Chemistry of Solutions

1
Chemistry of Solutions

• Chapter 7

2
Types of Solutions

• Although there are many examples of solutions in
  different phases gases in gases gases,
  liquids, or solids in liquids and liquids or
  solids in solids the most frequent situation in
  chemistry is working with something dissolved in
  a liquid.
• A solution is a homogeneous mixture i.e., no
  separation of solute and solvent, concentration
  the same everywhere.

3
Water

• Water is the most common solvent in a chemistry
  laboratory. Dissolves many materials because of
  its ability to form hydrogen bonds or because of
  its polarity.
• However, water has trouble dissolving many
  non-polar substances, particularly organic
  compounds.

4
Like Dissolves Like

• Polar solvents like to dissolve polar or ionic
  solutes salt in water, acetic acid in water,
  methanol in water, acetic acid in methanol
• Nonpolar solvents like to dissolve nonpolar
  solutes toluene in hexane, hexane in carbon
  tetrachloride
• Note that surfactants work by having a nonpolar
  end that is attracted to nonpolar grease and an
  opposite polar end attracted to water to carry
  the grease away. Also a model for cell walls
  (lipid chemistry).

5
Electrolytes and Nonelectrolytes

• An electrolyte is a solute that separates into
  ions in water.
• Differentiated by labels strong and weak.
• Strong dissociate 100 into ions. (NaCl)
• Weak stays mostly as intact molecules. Only a
  small portion dissociates into ions (acetic
  acid, phenol, ammonia)
• A nonelectrolyte does not dissociate at all
  (sugar, ethanol). Stays as intact molecules.
• Conductivity of a solution is a good measure of
  strength of an electrolyte.

6
Equivalents

• Used to describe electrolyte concentrations
  examples in book are taken from medical
  applications.
• Def an equivalent is the number of moles of an
  ion providing one mole of positive or negative
  charge.
• equivalents of ion moles of ion Absolute
  value of charge of the ion
• e.g., 0.4 moles Ca2 0.8 equivalents Ca2

7
Example on Equivalents

• A solution contains 40 mEq/L Cl- and 15 mEq/L of
  HPO42- . If Na is the only cation in the
  solution, what is the sodium ion concentration in
  milliequivalents per liter?
• What are the molar concentrations of each
  component of the solution?

8
(No Transcript)
9
Solubility

• Not every solution system is completely miscible.
  It is possible to saturate a solution. A
  saturated solution has the maximum amount of
  solute dissolved in a solvent at a given
  temperature. We see this all the time with the
  solubility of, for example, sugar in water.
• Solubility usually increases with temperature.
  Hence, more sugar dissolves in hot tea than in
  iced tea. This is because most solution
  processes are endothermic they absorb heat to
  make them go.

10
Solubility Example

• The solubility of KCl in water
• At 20 deg C, 34 g KCl will dissolve in 100 g
  water
• At 50 deg C, 43 g KCl will dissolve in 100 g
  water
• A solution containing 80. g of KCl in 200. g of
  water at 50 deg C is cooled to 20 deg C. How
  many grams of KCl remain in solution at 20 deg C?
  How many grams of KCl crystallized from solution
  after cooling?

11
(No Transcript)
12
Concentrations

• Defined in the form

13
Percent concentrations

• Mass Percent most common, except in medical
  applications
• Volume Percent (volumes not strictly additive)
• Mass / Volume Percent (using grams of solute, ml
  of solution) seems to be commonly used in
  medical applications

14
Example

• A patient needs 100. g of glucose in the next 12
  hours. How many liters of a 5 (m/v) glucose
  solution must be given?

15
(No Transcript)
16
Molarity

• Most common in the chemistry laboratory
• Gives the number of moles of solute present in a
  given volume. Easy to relate back to chemical
  equations which operate based on moles.

17
Example 1

• Calculate the molarity of 5.85 g of sodium
  chloride in 400. ml of solution.

18
Example 2

• Calculate the number of grams of solute needed to
  make 175 ml of 3.00 M sodium nitrate?

19
Example 3

• How many milliliters of 0.800 M calcium nitrate
  contain 0.0500 moles of this solute?

20
Example 4

• What is the final concentration in molarity of a
  solution in which water is added to 25 ml of a
  25 (m/v) solution of sulfuric acid until the
  final volume is 100.0 ml?

21
Example 5

• How many liters of 0.50 M phosphoric acid can be
  made from 0.500 liter of a 6.0 M phosphoric acid
  stock solution?

22
Example 6

• Lead(II) nitrate reacts with potassium chloride
  to produce lead(II) chloride and potassium
  nitrate. The lead(II) chloride precipitates as a
  solid and is removed from the reaction as it is
  formed.
• Write a balanced equation for this reaction.
• How many grams of lead(II) chloride will be
  formed from 50.0 ml of 1.50 M potassium chloride
  and excess lead(II) nitrate?
• How many milliliters of 2.00 M lead(II) nitrate
  are needed to react completely with 50.0 ml of
  1.50 M potassium chloride?

23
(No Transcript)
24
(No Transcript)