Chapter 11: Solutions and Solution Properties
11.1 Solution Composition
Concentrations:
So far our focus in terms of solution concentration has been on molarity, which is the most commonly used
by chemists.
Concentration
There are however, several other useful ways for measuring concentration. Here are some of them:
1. % by mass
2. % by volume
3. Molarity (M)
4. molality (m)
5. normality (N)
6. mole fraction (X)
7. ppm
8. ppb
Be sure you know the definitions for each and when they are appropriately used.
The terms “dilute” and “concentrated” are relative to the solute in question.
Types of Solutions
A solution can be formed in many ways other than solids in liquids. Put in an example of each of the
solutions in the table below:
Solute phase
Solvent phase
Example
Gas
Gas
Gas
Liquid
Gas
Solid
Liquid
Gas
None
Liquid
Liquid
Liquid
Solid
Solid
Gas
None
Solid
Liquid
Solid
Solid
11.2 The Energies of Solution Formation
Attractive Forces and Solubility
The surrounding of a solute particle by solvent molecules is called _________________, or in the case of
water as the solvent, _____________________.
There are three sets of attractive forces involved in the dissolving process:
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1. Attractions between solvent molecules. The solvent must expand to make room for the solute
particles.
2. Attractions between solute particles. The solute expands.
3. Attractions between solvent molecules and solute particles form the solution.
The 3 energy changes may be represented as H1, H2, and H3. Heat of solution is then calculated:
Hsln = H1 + H2 + H3.
If the attractions between solute particles (or between solvent molecules) are greater than the attractions of
the solvent molecules to the solute particles, the solute will not dissolve.
If however, the solvent-solute attractions are strong enough, the solute will be pulled apart by the solvent
molecules particle by particle and the solute will dissolve.

Because truly dissolved solute particles are on the order of 1 nm or less in diameter, they are
too small to reflect or refract light. Therefore, except for metal alloys, true solutions are
transparent. (Transparent and colorless do NOT mean the same thing.)
Eventually, when all the solvent molecules are involved surrounding solute particles, so that overall no more
solute can dissolve, the solution is __________________.
Actually, although it appears that dissolving has stopped, the system has really reached a dynamic
equilibrium in which the rate of crystallization has caught up to the rate of dissolving.
Solute + Solvent
Solution
 is dissolving,  is crystallizing (precipitating)
We can easily see also that when 2 solutions are poured together in which oppositely charged ions (a cation
from one solution and an anion from the other) have very strong attraction for each other (greater than their
attractions to the solvent molecules) a ___________________ will form.
It is possible for a solution to have more dissolved solute than theoretically possible.
circumstance is referred to as a _____________________ solution.
This special
Heat of Solution
When dissolving occurs an energy change will accompany the process, as witnessed by a change in solution
temperature. Sometimes the change is slight, and at others it is quite dramatic.
Hsln is measured in kJ/mol of solute dissolved.
If there is a -Hsln value, it means dissolving is ______________ and solution temperature will
____________.
If the value is +Hsln , then dissolving is ______________ and solution temperature will _______________.
Please think about this carefully, as it often seems backwards in terms of the temperature. It may help to
think about the fact that it is the solvent that is giving up or absorbing the energy involved in the dissolving
process.

An example of a solute with a –Hsln is _______________________.
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

An example of a solute with a +Hsln is _______________________.
A salt in which there is very little change in temperature with dissolving is ___________.
When might knowledge of a substance’s heat of solution be useful?
The solubility of a solute at various temps is called a solubility curve. A solute with a curve of increasing
slope indicates a ____Hsln.
A decreasing slope would indicate a ____Hsln.
A standard, simple equation for calculating a H value based on mass of substance and temperature change
is H = mcT where “m” is the mass of the substance that is changing temperature (T) and “c” is the
specific heat capacity of that substance. You may remember this from our calorimetric problems from an
earlier unit.
Please remember that there are many solvents besides water. But the same principles apply whether
for calculating concentrations, heats of solution or colligative properties.
11.3 Factors Affecting Solubility
Factors Affecting Rate of Dissolving
Please be very careful not to confuse these ideas.
The RATE at which a solid solute dissolves can be increased by
1. _________________ the solute.
2. _________________ the mixture.
Solubility Factors
Besides the actual nature of the solvent and solute, the only things that affect SOLUBILITY (how MUCH
solute will dissolve are
1. Temperature and
2. Pressure (mostly for gases in liquids.)
Like Dissolves Like
The more alike solvent and solute molecules (particles) are in terms of their polarity, the easier dissolving
will be.
Water is our most important polar solvent. Water molecules are polar and capable of hydrogen bonding.
It dissolves polar solutes and many ionic solutes well. Most any solute with –OH or –NH2 groups can
hydrogen bond to water and will likely dissolve very well.
Except for small alcohols, many organic solvents are non-polar and will not dissolve things that water
dissolves easily.
Examples of organic solvents: Polar: Methanol, ethanol, ethylene glycol, glycerol.
Non-polar: ethers, acetone, benzene, methylchloride, naphthalene,
gasoline
Another nice example can be found in vitamins. Some vitamins are water soluble (B and C) and others are
fat soluble (A, D, E and K). The body can store fat soluble vitamins, and taking too much can actually be
harmful. The body does not store water soluble vitamins and they must be consumed daily.
A look at the molecules makes it easy to see why their solubilities differ.
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Vitamin A
Vitamin C
It’s pretty easy to guess the polarity of substances (that are not ionic) by seeing whether they will dissolve in
water or not. For example, the ink in a ballpoint pen is non-polar while table sugar is a polar molecule.
Another way to make a pretty intelligent guess about a substance’s polarity is to see what physical state it is
in under normal conditions, or how easily it is converted to a gas or liquid.
The reason all of the main gases of air (except water vapor) are gases is that they are _______________.
If a substance has a relatively low freezing and boiling point, it is _________________.
Pressure Effects
The solubility of gases in liquids is affected by pressure.
Henry’s Law: Cg = kPg
C is the gas solubility, P is the pressure of the gas and k is Henry’s law proportionality constant.
Simply put, gases dissolved best at high pressure. Gas solubility is directly proportional to the pressure of
the gas above the liquid.
That’s why soft drinks and other carbonated beverages are bottled at 4 to 5 times normal atmospheric
pressure. Carbon dioxide is non-polar and doesn’t dissolve well in water. High pressure forces the gas to
become more soluble.
Temperature Effects
Gas solubility also depends heavily on temperature. Gases become ____________soluble as temperature
_________________.
“Thermal pollution” in environmental waters is a serious problem, because as waters get warmer, CO2 that is
needed by _________________ and O2, needed by ________________ both become less soluble.
Even in the very coldest water, the maximum dissolved oxygen concentration is about __________.
Why do all marine mammals breathe air (rather than having gills)?
Most substances have a + heat of solution, becoming more soluble as solution temperature increases.
11.4 Vapor Pressure
Colligative Properties
These are the physical properties of solutions. They include
1. Vapor pressure (VP)
2. Boiling Point (BP)
3. Freezing Point (FP)
4. Osmotic Pressure (OP)
Colligative properties change with the concentration of dissolved particles in liquid solvents.
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Vapor pressure is the gaseous pressure caused by evaporating solvent molecules. For a pure solvent, VP
depends on 2 things: the type of solvent and temperature. As temp. increases, VP _________________.
When a nonvolatile solute is dissolved in a solvent, fewer molecules of solvent are at the surface of the
liquid where they have the chance to escape to the vapor phase. The result is a lower VP for the solution.
Raoult’s Law calculates the new vapor pressure of a solution based on solute concentration.
PA = XAPA where PA is the VP of the solution, XA is the mole fraction of the solvent in the solution and PA
is the VP of the pure solvent at a particular temperature. Your text has example problems of Raoult’s Law.
Raoult’s Law predicts ideal solution behavior. Most real solutions do not follow predictions exactly, but
some have higher pressures that expected while others have lower depending on the intermolecular forces
existing between solvent and solute particles.
11.5 Boiling Point Elevation and Freezing Point Depression
Boiling Point Elevation
Boiling point is defined as the temperature at which the vapor pressure of a liquid equals the pressure of the
gas above the liquid. (Remember how boiling points are affected by pressure: triple point diagrams.)
If the presence of a solute reduces the solvent’s VP, the liquid must be hotter to boil. It has farther to go to
get the VP up to the pressure of the gas above the liquid resulting in boiling point elevation.
Calculating changes in BP is easy: Tb = Kbmi
 Tb is the change or increase in the BP (normally 100C for water).
 Kb is the molal boiling point constant for the solvent. For water, the value is 0.51C/m, but it is
different for other solvents. The AP loves to give this kind of problem using a solvent other than
water.
 m is the molality of the solution (moles solute/1000g of solvent).
 i is a value called the van’t Hoff factor that will be discussed shortly.
Freezing Point Depression
Freezing occurs when the solvent molecules slow down enough that intermolecular forces of attractions
begin to lock them into solid form.
The freezing of solutions is a difficult thing to define. Solutions, especially concentrated ones don’t seem to
freeze cleanly or solidly. Freezing point changes are calculated the same way as BP changes.
Tf = Kfmi Kf is the molal freezing point constant. (The value for water is 1.86C/m and the normal FP
for water is 0C.
The presence of solute particles “interferes” with solvent molecules getting next to and locking onto each
other in the freezing process. As the solvent particles lock together, crystals of pure solvent begin to freeze
out of the solution. As the liquid solvent is removed to the solid form, the remaining solution becomes more
and more concentrated, its freezing point continuing to decline. Solutions therefore often turn to “slush” and
may not ever freeze completely solid. Ever suck on a Popsicle?
11.7 Osmotic Pressure: Don’t worry too much about osmotic pressure. That is more of biological interest
and will probably not be on the AP exam. Osmotic pressure is measured in atmospheres and is calculated
 = MRT where is osmotic pressure in atmospheres, M is molarity, R is 0.0821 L atm/mol K and T is
temperature in K.
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But be sure you can solve problems that find molar mass from freezing point depression or boiling point
elevation.
11.7 Colligative Properties of Electrolyte Solutions
The van’t Hoff Factor (i)
When strong electrolytes dissolve, they produce 2 or more ions from each dissolved molecule or
formula unit.
The van’t Hoff factor tells us how many particles there will be and therefore how may times the
expected effect the solute will have on the BP or FP.




NaCl  __________________
i = _____
HBr  ___________________
i = _____
Ca(NO3)2  ______________________
i = _____
Al2(SO4)3  ______________________
i = _____
Technically, the van’t Hoff factor is calculated as
i = T actual/T calculated for the solute as a non-electrolyte
What we determined above was the ideal van’t Hoff factor. The actual changes in temperature are
never ideal because when oppositely charged ions come together, they temporarily “stick” to each
other acting like one particle instead of two. So the factor is always somewhat less than predicted
ideally. For weak electrolytes (like acetic acid), it will be much less than predicted, because these
electrolytes do not ionize completely. The factor for acetic acid may only be 1.1.
11.8 Colloids
Colloids
When you have a strange substance and you can’t seem to come up with a good answer when you ask
yourself if it’s a solid, liquid or gas, it is probably a colloid.
Colloids include things that seem to defy simple description, like Jello, clouds and Silly Putty.
A colloid is a mixture, but it is not a solution. Where the dispersed particles of a true are less than
one nanometer in diameter and too small to settle out by gravity or reflect light, the dispersed
particles of a colloid are ______________ nanometers in diameter. Such particles may include
macromolecules (proteins, starch or DNA), cells (red blood cells) or small liquid drops, gas bubbles
or solid particles (like smoke).
Even though not truly dissolved, colloidal particles resist settling because of
1. Brownian movement- the constant bombardment of the particles by surrounding molecules,
and
2. like electrostatic charges, that cause the colloidal particles to repel each other.
These particles are too small to be separated by gravity or normal filtration, but are large enough to
reflect/refract light.
This quality results in the Tyndall Effect. Whereas a beam of light is undetectable as it passes
through a true solution, a beam is visible in a colloidal suspension. A car’s headlights are invisible in
dry, clear air, but are visible in air that is foggy or dusty. The small droplets of water or particles of
solid dust make the air a colloid.
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Dispersing Phase
(Solvent like
phase)
Gas
Gas
Gas
Liquid
Liquid
Liquid
Solid
Solid
Solid
Dispersed
Phase
Colloid
Type
Example
Gas
Liquid
Solid
Gas
Liquid
Solid
Gas
Liquid
Solid
None
Aerosol
Dry aerosol
Foam
Emulsion
Sol
Solid foam
Solid emulsion
Solid sol
All are true solutions.
Fog
Smoke
Whipped cream
Milk, mayonnaise
Paint
Marshmallow/Styrofoam
Butter
Ruby glass
Liquid in liquid emulsions are common and important colloids. Two liquids that will dissolve in
each other in any proportion (e.g. water and ethanol) are called miscible. Liquids that won’t mix and
dissolve (e. g. water and oil) are immiscible. Substances which are added to immiscible liquids that
allow them to mix and stay mixed are called emulsifying agents. Mayonnaise is made of vegetable
oil and water, but requires the proteins in egg albumin (white) to stay mixed. Detergents are
emulsifying agents that allow the grease and oil on clothing and dishes to mix with water be carried
away.
Then, of course, the picture becomes even more interesting with substances like milk or blood, which
are both true solutions and colloids at the same time.
Removal of Colloidal Particles
If the electrostatic charges on the colloidal particles can be masked or removed, the colloid breaks
down. Lightening discharges colloidal water droplets. The droplets coalesce into raindrops. Salt
ruins Jello because the ions mask the static charges of the colloidal protein molecules in the Jello.
Salts are also used to remove precipitated “flocculent” from wastewater. Heat also destroys colloids
nicely. Again, the Jello collapses when heated.
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