mc_ch12.ppt

Preview
• Objectives
• Solutions
• Suspensions
• Colloids
• Solutes: Electrolytes Versus Nonelectrolytes
Chapter 12

Objectives
• Distinguish between electrolytes and
nonelectrolytes.
• List three different solute-solvent combinations.
• Compare the properties of suspensions, colloids, and
solutions.
• Distinguish between electrolytes and
nonelectrolytes.
Chapter 12 Section 1 Types of Mixtures

Solutions
• You know from experience that sugar dissolves in
water. Sugar is described as “soluble in water.” By
soluble we mean capable of being dissolved.
• When sugar dissolves, all its molecules become
uniformly distributed among the water molecules.
The solid sugar is no longer visible.
• Such a mixture is called a solution. A solution is a
homogeneous mixture of two or more substances in
a single phase.

Click below to watch the Visual Concept.
Visual Concept
Chapter 12
Solutions
Section 1 Types of Mixtures

Solutions, continued
• The dissolving medium in a solution is called the
solvent, and the substance dissolved in a solution is
called the solute.
• Solutions may exist as gases, liquids, or solids.
There are many possible solute-solvent combinations
between gases, liquids, and solids.
• example: Alloys are solid solutions in which the atoms of
two or more metals are uniformly mixed.
• Brass is made from zinc and copper.
• Sterling silver is made from silver and copper.

Visual Concept
Chapter 12
Solutes, Solvents, and Solutions

Visual Concepts
Chapter 12
Types of Solutions

Particle Models for Gold and Gold Alloy

Suspensions
• If the particles in a solvent are so large that they
settle out unless the mixture is constantly stirred or
agitated, the mixture is called a suspension.
• For example, a jar of muddy water consists of soil particles
suspended in water. The soil particles will eventually all
collect on the bottom of the jar, because the soil particles
are denser than the solvent, water.
• Particles over 1000 nm in diameter—1000 times as
large as atoms, molecules or ions—form
suspensions.

Visual Concept
Chapter 12
Suspensions

Colloids
• Particles that are intermediate in size between those
in solutions and suspensions form mixtures known
as colloidal dispersions, or simply colloids.
• The particles in a colloid are small enough to be suspended
throughout the solvent by the constant movement of the
surrounding molecules.
• Colloidal particles make up the dispersed phase, and
water is the dispersing medium.
• example: Mayonnaise is a colloid.
• It is an emulsion of oil droplets in water.

Colloids, continued
Tyndall Effect
• Many colloids look similar to solutions because their
particles cannot be seen.
• The Tyndall effect occurs when light is scattered by
colloidal particles dispersed in a transparent medium.
• example: a headlight beam is visible from the
side on a foggy night.
• The Tyndall effect can be used to distinguish
between a solution and a colloid.

Visual Concepts
Colloids
Chapter 12

Visual Concepts
Emulsions
Chapter 12

Properties of Solutions, Colloids, and
Suspensions

Solutes: Electrolytes Versus Nonelectrolytes
• A substance that dissolves in water to give a solution
that conducts electric current is called an
electrolyte.
• Any soluble ionic compound, such as sodium
chloride, NaCl, is an electrolyte.
• The positive and negative ions separate from
each other in solution and are free to move, making
it possible for an electric current to pass through
the solution.

Solutes: Electrolytes Versus Nonelectrolytes,
continued
• A substance that dissolves in water to give a
solution that does not conduct electric current is
called a nonelectrolyte.
• Sugar is an example of a nonelectrolyte.
• Neutral solute molecules do not contain mobile
charged particles, so a solution of a nonelectrolyte
cannot conduct electric current.

Electrical Conductivity of Solutions

Preview
• Objectives
• Factors Affecting the Rate of Dissolution
• Solubility
• Solute-Solvent Interactions
• Enthalpies of Solution
Chapter 12 Section 2 The Solution Process

Objectives
• List and explain three factors that affect the rate at
which a solid solute dissolves in a liquid solvent.
• Explain solution equilibrium, and distinguish among
saturated, unsaturated, and supersaturated
solutions.
• Explain the meaning of “like dissolves like” in terms
of polar and nonpolar substances.

Objectives, continued
• List the three interactions that contribute to the
enthalpy of a solution, and explain how they combine
to cause dissolution to be exothermic or
endothermic.
• Compare the effects of temperature and pressure on
solubility.

Chapter 12
Dissolving Process
Section 2 The Solution Process

• Because the dissolution process occurs at the
surface of the solute, it can be speeded up if the
surface area of the solute is increased.
• Stirring or shaking helps to disperse solute particles
and increase contact between the solvent and solute
surface. This speeds up the dissolving process.
Factors Affecting the Rate of Dissolution
• At higher temperatures, collisions between solvent
molecules and solvent are more frequent and of
higher energy. This helps to disperse solute
molecules among the solvent molecules, and
speed up the dissolving process.

Visual Concept
Factors Affecting the Rate of Dissolution

Solubility
• If you add spoonful after spoonful of sugar to tea,
eventually no more sugar will dissolve.
• This illustrates the fact that for every combination
of solvent with a solid solute at a given temperature,
there is a limit to the amount of solid that can
be dissolved.
• The point at which this limit is reached for any solute-
solvent combination depends on the nature of the
solute, the nature of the solvent, and the
temperature.

Particle Model
for Soluble
and Insoluble
Substances

Particle Model for Soluble and Insoluble
Substances

Solubility, continued
• When a solute is first added to a solvent, solute
molecules leave the solid surface and move about at
random in the solvent.
• As more solute is added, more collisions occur
between dissolved solute particles. Some of the solute
molecules return to the crystal.
• When maximum solubility is reached, molecules are
returning to the solid form at the same rate at which
they are going into solution.

• Solution equilibrium is the physical state in which the
opposing processes of dissolution and crystallization of
a solute occur at the same rates.

Visual Concept
Solution Equilibrium

Saturated Versus Unsaturated Solutions
• A solution that contains the maximum amount of
dissolved solute is described as a saturated solution.
• If more solute is added to a saturated solution, it falls to the
bottom of the container and does not dissolve.
• This is because an equilibrium has been established between
ions leaving and entering the solid phase.
• A solution that contains less solute than a
saturated solution under the same conditions is
an unsaturated solution.

Mass of Solute Added Versus Mass of Solute
Dissolved

Supersaturated Solutions
• When a saturated solution is cooled, the excess solute
usually comes out of solution, leaving the solution
saturated at the lower temperature.
• But sometimes the excess solute does not separate,
and a supersaturated solution is produced, which is a
solution that contains more dissolved solute than a
saturated solution contains under the same conditions.
• A supersaturated solution will form crystals of
solute if disturbed or more solute is added.

Solubility Values
• The solubility of a substance is the amount of that
substance required to form a saturated solution with a
specific amount of solvent at a specified temperature.
• example: The solubility of sugar is 204 g per 100 g of water
at 20°C.
• Solubilities vary widely, and must be determined
experimentally.
• They can be found in chemical handbooks and are usually
given as grams of solute per 100 g of solvent at a given
temperature.

Solubility of Common Compounds

Visual Concept
Chapter 12
Solubility of a Solid in a Liquid

Solute-Solvent Interactions
• Solubility varies greatly with the type of
compounds involved.
• “Like dissolves like” is a rough but useful rule
for predicting whether one substance will dissolve
in another.
• What makes substances similar depends on:
• type of bonding
• polarity or nonpolarity of molecules
• intermolecular forces between the solute and
solvent

Visual Concept
Chapter 12
Like Dissolves Like

Solute-Solvent Interactions, continued
Dissolving Ionic Compounds in Aqueous Solution
• The polarity of water molecules plays an important
role in the formation of solutions of ionic compounds
in water.
• The slightly charged parts of water molecules attract
the ions in the ionic compounds and surround them,
separating them from the crystal surface and drawing
them into the solution.
• This solution process with water as the solvent
is referred to as hydration. The ions are said to
be hydrated.

Dissolving Ionic Compounds in Aqueous Solution
The hydration of the ionic solute lithium chloride is shown below.

Nonpolar Solvents
• Ionic compounds are generally not soluble in
nonpolar solvents such as carbon tetrachloride, CCl4,
and toluene, C6H5CH3.
• The nonpolar solvent molecules do not attract the
ions of the crystal strongly enough to overcome the
forces holding the crystal together.
• Ionic and nonpolar substances differ widely in
bonding type, polarity, and intermolecular forces,
so their particles cannot intermingle very much.

Liquid Solutes and Solvents
• Oil and water do not mix because oil is nonpolar
whereas water is polar. The hydrogen bonding
between water molecules squeezes out whatever oil
molecules may come between them.
• Two polar substances, or two nonpolar substances,
on the other hand, form solutions together easily
because their intermolecular forces match.
• Liquids that are not soluble in each other are
immiscible. Liquids that dissolve freely in one
another in any proportion are miscible.

Visual Concept
Chapter 12
Comparing Miscible and Immiscible Liquids

Effects of Pressure on Solubility
• Changes in pressure have very little effect on the
solubilities of liquids or solids in liquid solvents.
However, increases in pressure increase gas
solubilities in liquids.
• An equilibrium is established between a gas above a
liquid solvent and the gas dissolved in a liquid.
• As long as this equilibrium is undisturbed, the
solubility of the gas in the liquid is unchanged at a
given pressure:




gas + solvent solution

Effects of Pressure on Solubility, continued
• Increasing the pressure of the solute gas above the
solution causes gas particles to collide with the liquid
surface more often. This causes more gas particles
to dissolve in the liquid.


 



 

• Decreasing the pressure of the solute gas above the
solution allows more dissolved gas particles to
escape from solution.

Visual Concept
Chapter 12
Pressure, Temperature, and Solubility of Gases

Henry’s Law
• Henry’s law states that the solubility of a gas in a
liquid is directly proportional to the partial pressure of
that gas on the surface of the liquid.
• In carbonated beverages, the solubility of carbon dioxide is
increased by increasing the pressure. The sealed
containers contain CO2 at high pressure, which keeps the
CO2 dissolved in the beverage, above the liquid.
• When the beverage container is opened, the pressure
above the solution is reduced, and CO2 begins to escape
from the solution.
• The rapid escape of a gas from a liquid in which it
is dissolved is known as effervescence.

Visual Concept
Chapter 12
Henry’s Law

Visual Concept
Chapter 12
Effervescence

Effects of Temperature on Solubility
• Increasing the temperature usually decreases gas
solubility.
• As temperature increases, the average kinetic energy of
molecules increases.
• A greater number of solute molecules are therefore able to
escape from the attraction of solvent molecules and return
to the gas phase.
• At higher temperatures, therefore, equilibrium is reached
with fewer gas molecules in solution, and gases are
generally less soluble.

Effects of Temperature on Solubility
• Increasing the temperature usually increases
solubility of solids in liquids, as mentioned previously.
• The effect of temperature on solubility for a given
solute is difficult to predict.
• The solubilities of some solutes vary greatly over
different temperatures, and those for other solutes
hardly change at all.
• A few solid solutes are actually less soluble at
higher temperatures.

Solubility vs.
Temperature

Enthalpies of Solution
• The formation of a solution is accompanied by an
energy change.
• If you dissolve some potassium iodide, KI, in
water, you will find that the outside of the
container feels cold to the touch.
• But if you dissolve some sodium hydroxide,
NaOH, in the same way, the outside of the
container feels hot.
• The formation of a solid-liquid solution can
either absorb energy (KI in water) or release
energy as heat (NaOH in water)

Enthalpies of Solution, continued
• Before dissolving begins, solute particles are held
together by intermolecular forces. Solvent particles
are also held together by intermolecular forces.
• Energy changes occur during solution formation
because energy is required to separate solute
molecules and solvent molecules from their
neighbors.
• A solute particle that is surrounded by solvent
molecules is said to be solvated.

The diagram above shows the enthalpy changes
that occur during the formation of a solution.

• The net amount of energy absorbed as heat by the
solution when a specific amount of solute dissolves
in a solvent is the enthalpy of solution.
• The enthalpy of solution is negative (energy is
released) when the sum of attractions from Steps 1
and 2 is less than Step 3, from the diagram on the
previous slide.
• The enthalpy of solution is positive (energy is
absorbed) when the sum of attractions from Steps 1
and 2 is greater than Step 3.

Preview
• Objectives
• Concentration
• Molarity
• Molality
Chapter 12
Section 3 Concentration of
Solutions

Objectives
• Given the mass of solute and volume of solvent,
calculate the concentration of solution.
• Given the concentration of a solution, determine the
amount of solute in a given amount of solution.
• Given the concentration of a solution, determine the
amount of solution that contains a given amount of
solute.
Chapter 12
Solutions

Concentration
• The concentration of a solution is a measure of
the amount of solute in a given amount of solvent
or solution.
• Concentration is a ratio: any amount of a given
solution has the same concentration.
• The opposite of concentrated is dilute.
• These terms are unrelated to the degree to which a
solution is saturated: a saturated solution of a solute
that is not very soluble might be very dilute.
Chapter 12
Solutions

Concentration Units
Solutions
Chapter 12

Visual Concept
Chapter 12
Concentration
Solutions

Molarity
• Molarity is the number of moles of solute in one liter
of solution.
• For example, a “one molar” solution of sodium
hydroxide contains one mole of NaOH in every liter
of solution.
• The symbol for molarity is M. The concentration of a
one molar NaOH solution is written 1 M NaOH.
Chapter 12
Solutions

Molarity, continued
• To calculate molarity, you must know the amount of
solute in moles and the volume of solution in liters.
• When weighing out the solute, this means you will
need to know the molar mass of the solute in order to
convert mass to moles.
• example: One mole of NaOH has a mass of
40.0 g. If this quantity of NaOH is dissolved in
enough water to make 1.00 L of solution, it is a
1.00 M solution.
Chapter 12
Solutions

Molarity, continued
• The molarity of any solution can be calculated by
dividing the number of moles of solute by the number
of liters of solution:
amount of solute (mol)
molarity (M)
volume of solution (L)

Chapter 12
Solutions
• Note that a 1 M solution is not made by adding 1
mol of solute to 1 L of solvent. In such a case, the
final total volume of the solution might not be 1 L.
• Solvent must be added carefully while dissolving to
ensure a final volume of 1 L.

Visual Concept
Chapter 12
Preparation of a Solution Based on Molarity
Solutions

Molarity, continued
Sample Problem A
You have 3.50 L of solution that contains 90.0 g of
sodium chloride, NaCl. What is the molarity of that
solution?
Chapter 12
Solutions

Molarity, continued
Sample Problem A Solution
Given: solute mass = 90.0 g NaCl
solution volume = 3.50 L
Unknown: molarity of NaCl solution
Solution:
1 mol NaCl
90.0 g NaCl 1.54 mol NaCl
58.44 g NaCl
 
1.54 mol NaCl
3.50 L of
0.44
solu
0 M
ti
l
on
NaC

Chapter 12
Solutions

Molarity, continued
Sample Problem B
You have 0.8 L of a 0.5 M HCl solution. How many
moles of HCl does this solution contain?
Chapter 12
Solutions

Molarity, continued
Sample Problem B Solution
Given: volume of solution = 0.8 L
concentration of solution = 0.5 M HCl
Unknown: moles of HCl in a given volume
Solution:
0.5 mol HCl
0.8 L of solution
1.0 L of solut
0.4 mol
ion
HCl
 
Chapter 12
Solutions

Molarity, continued
Sample Problem C
To produce 40.0 g of silver chromate, you will need at
least 23.4 g of potassium chromate in solution as a
reactant. All you have on hand is 5 L of a 6.0 M
K2CrO4 solution. What volume of the solution is
needed to give you the 23.4 g K2CrO4 needed for the
reaction?
Chapter 12
Solutions

Sample Problem C Solution
Given: volume of solution = 5 L
concentration of solution = 6.0 M K2CrO4
mass of solute = 23.4 K2CrO4
mass of product = 40.0 g Ag2CrO4
Unknown: volume of K2CrO4 solution in L
Molarity, continued
Chapter 12
Solutions

2 4
2 4
2 4
0.120 mol K CrO
6.0 M K CrO
L K CrO solution
x

Sample Problem C Solution, continued
Solution:
Molarity, continued
2 4 2 4
1 mol K CrO 194.2 g K CrO

2 4
2 4 2 4
2 4
1 mol K CrO
23.4 g K CrO 0.120 mol K CrO
194.2 g K CrO
 
2 4
0.020 L K CrO solution
x 
Chapter 12
Solutions

Molality
• Molality is the concentration of a solution expressed
in moles of solute per kilogram of solvent.
• A solution that contains 1 mol of solute dissolved in 1
kg of solvent is a “one molal” solution.
• The symbol for molality is m, and the concentration
of this solution is written as 1 m NaOH.
Chapter 12
Solutions

• The molality of any solution can be calculated by
dividing the number of moles of solute by the number
of kilograms of solvent:
amount of solute (mol)
molality ( )
mass of solvent (kg)
m 
Chapter 12
Solutions
Molality, continued
• Unlike molarity, which is a ratio of which the
denominator is liters of solution, molality is per
kilograms of solvent.
• Molality is used when studying properties of solutions
related to vapor pressure and temperature changes,
because molality does not change with temperature.

Visual Concept
Chapter 12
Comparing Molarity and Molality
Solutions

Making a Molal Solution
Chapter 12
Solutions

Molality, continued
Sample Problem D
A solution was prepared by dissolving 17.1 g of sucrose
(table sugar, C12H22O11) in 125 g of water. Find the
molal concentration of this solution.
Chapter 12
Solutions

Molality, continued
Sample Problem D Solution
Given: solute mass = 17.1 C12H22O11
solvent mass = 125 g H2O
Unknown: molal concentration
Solution: First, convert grams of solute to moles and
grams of solvent to kilograms.
12 22 11
12 22 11 12 22 11
12 22 11
1 mol C H O
17.1 g C H O 0.0500 mol C H O
342.34 g C H O
 
2
2
125 g H O
0.125 kg H O
1000 g/kg

Chapter 12
Solutions

Molality, continued
Sample Problem D Solution, continued
Then, divide moles of solute by kilograms of solvent.
12 22 1
12 22 11
1
2
0.0500 mol C H O
0.125
0.400 C
kg H O
H O
m

Chapter 12
Solutions

Molality, continued
Sample Problem E
A solution of iodine, I2, in carbon tetrachloride, CCl4, is
used when iodine is needed for certain chemical tests.
How much iodine must be added to prepare a 0.480 m
solution of iodine in CCl4 if 100.0 g of CCl4 is used?
Chapter 12
Solutions

Molality, continued
Sample Problem E Solution
Given: molality of solution = 0.480 m I2
mass of solvent = 100.0 g CCl4
Unknown: mass of solute
Solution: First, convert grams of solvent to kilograms.
4
4
100.0 g CCl
0.100 kg CCl
1000 g/kg

Chapter 12
Solutions

2 2
2
2
253.8 g I
0.480 mol I
mol
12.2 I
I
g
 
Molality, continued
Sample Problem E Solution, continued
Solution, continued: Then, use the equation for
molality to solve for moles of solute.
 
x
m x
2
2
2
mol I
0.480 0.0480 mol I
0.1 kg H O
Chapter 12
Solutions
Finally, convert moles of solute to grams of solute.

mc_ch12.ppt

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