Chapter 13
Solutions
Chemistry 100
Definition
When one substances is dispersed
uniformly throughout another at the
molecular level, we have a solution.
 Each substance in a solution is a
component.
 The solvent is generally the component
present in the greatest amount.
 The other components are called solutes

Types of Solutions







All gases mix. Generally, we do not think of a
mixture of two gases as a solution.
Gases dissolve in liquids - oxygen in water
Liquids dissolve in liquids - alcohol in water
Solids dissolve in liquids - sugar in tea
Gases dissolve in solids - hydrogen in palladium
Liquids dissolve in solids (rare) - Hg in silver
Solids dissolve in solids - alloys (brass, solder,
etc)
Energy considerations
Begin with solid S and liquid L
 End with S dissolved in L
 The particles in the solid must separate to
dissolve.

 Overcome
intermolecular forces or
electrostatic forces in ionic compounds!

Let H1 be the enthalpy for this process
Solvent-Solvent
Interactions
The solvent molecules must separate to
allow room for the solute molecules.
Intermolecular forces must be overcome in
L
 Let H2 be the enthalpy for this process.

 H1
and H2 will be positive - energy is
added to the system
Solute-Solvent Interactions
The solute and solvent particles interact.
 Example: Ions dissolved in water give rise
to ion-dipole interactions as the water
molecules surround the ions.

 Hydration

Let H3 be the enthalpy for this process
 H3
will be negative - energy is released
when two objects that attract get closer
together
Enthalpy of Solution
For the overall process of S + L 
Solution
Hsoln = H1 + H2 + H3
 The first two will be endothermic
processes (positive) values since they
need to take in energy.
 The last process is exothermic (negative
value) since heat is released when two
attracting molecules are brought together.

Like Dissolves Like

The enthalpy of solution will be
if H3 has a small (negative) value.
 negative when H3 has a large (negative) value.
 positive

H3 will have a
 large
value with an ionic solid and a polar liquid
 small value with a non-polar solute is and a polar
solvent
 If Hsoln very endothermic the solution will not form
Spontaneity
A process is said to be spontaneous when
it occurs without external help
 Endothermic process can be spontaneous
provides H is not too large (more later)
 Example:

NaOH plus water: H = -44 kJ/mole
NH4NO3 plus water: H = +26 kJ/mol
Entropy of Solution






State A consists of solid and liquid as separate phases
State B consists of the solid dissolved in the liquid.
State B has a higher entropy!
Which has least energy? That depends on the the H
values discussed earlier - the net value is the enthalpy of
solution
If enthalpy of solution is negative (exothermic), state B
(solution) is preferred
If it is only slightly endothermic, state B can still be
preferred
Concentration terms
Dilute - not a lot of solute. Concentrated a large amount of solute. Chemists try not
to talk about a strong solution - as in: a
strong cup of tea
 Concentration can be expressed
quantitatively is many ways:
 molarity, molality, percentage, mole
fraction, etc

Molarity

The molarity is the number of moles of
solute in 1 litre of solution.
M
= moles of solute / volume (litres) solution
A 1 M solution is fairly concentrated (often
100 s of grams of solute in a litre)
 Millimolar:

solution is 110-3moles/L
 or 1 millimole/L or 1mM
 A 0.001M
Molality

The molality is the number of moles of
solute dissolved in 1 kg of solvent.
m
= moles of solute / volume (litres) solution
Molalities are related to mass fractions (%
mass, see later).
 Molality is independent of the temperature.

 A 1.0
m (for molal) solution is 1.0 moles of
solute dissolved in 1.00 kg of solvent.
Normality


Chemists has stopped using normality as a
concentration unit but the health profession has
not.
Simple rule to convert molarity to normality
 Multiply
molarity by ion charge to get normality
 A 0.12 M Na+ solution is a 0.12 N Na+ solution
 A 0.25 M Ca2+ solution is a 0.50 N Ca2+ soln.

This will not be on the exam but nurses need to
know it.
Mass percentage




A solution that is 36% HCl by mass, contains 36
g HCl in 100 g of solution.
Milk is not a true solution, but 2% milk refers to 2
g of milk fat in 100 g of milk.
You may see figures on packing such as
3% w/w where w is short for weight (mass)
% by mass 
mass of component
total mass of solution
 100
Parts per million
Percentage refers to parts in 100. Thus
3% means 3 parts solute in 100 parts of
solution.
 Parts per million - how many parts of
solute in a million parts of solution

ppm 
mass of component
total mass of solution
 1,000,000
ppm
Let our solution have that has a Ag+ ion
concentration of 3 ppm
3 g Ag+ in 1,000,000 g of solution
3 mg Ag+ in 1,000 g of solution (1/1000th of
each)
But 1,000 g of water solution has a volume of
approximately 1 litre.
So our solution has 3mg Ag+ / litre
ppm is approximately the same as mg/L
Ppb - parts per billion
This unit is used for highly toxic materials
We saw that 1 ppm = 1 mg / L
In the same way, 1 ppb = 1 g /L
where 1 g is 1 microgram or 110-6g
Percentage by volume
% by volume 
v olume of component
total v olume of solution
 100
When a solute is a liquid, solution
concentrations s are sometimes
expressed as percentages by volume
 Thus a 20%v/v solution of hydrogen
peroxide contains 20 ml of H2O2 in every
100 ml of solution.

Mole fraction
If a solution consists of mA moles of A, mB
moles of B, mC moles of C, etc.
Mole fraction of A
XA = mA / (mA + mB + mC)
Clearly,
XA + XB + XC = 1
Terminology
Solute
 Solvent
 Solubility
 Saturated solution
 Unsaturated solution

Factors That Affect Solubility
Solute-solvent interaction
 Pressure effect when solute is a gas
Henry’s Law: Cg = kPg
k is a constant for the gas at a given temp.
 Temperature effect
most solids are more soluble in warm
water
gases are less soluble in warm water

Colligative Properties

Properties that depend on solute
concentration (collection) not the kind of
solute.
 Depression
of the vapour pressure
 Depression of the freezing point
 Elevation of the boiling point
 Osmosis
Explanation of lowering of VP
There are serious problems with this so-called explanation
Depression of VP
Raoult’s Law:
How VP of solution relates to VP of pure solvent
PA = XAP0A
Note that when XA = 1 (pure solvent), PA = P0A
Solutions that obey Raoult’s law
are called ideal solutions
Colligative Properties
Freezing & boiling point changes
For dilute solutions of non-volatile solutes
Tf = mKf
Kf varies with solvent
Tb = mKb
Kb varies with solvent
For electrolytes, m is the molality of particles:
1m NaCl is 2m in particles (n=2)
1m CaCl2 is 3m in particles (n=3)
Use T = nmK for solutions of electrolytes
Osmosis
Osmosis

The movement of water through a semipermeable membrane from dilute side to
concentrated side
 the
movement is such that the two sides might
end up with the same concentration

Osmotic pressure: the pressure required
to prevent this movement
Osmosis
Ideal gas law:
PV = nRT
V = nRT
 = (n/V)RT
 = MRT
Remember to use osmotic pressure in ATMS
so use can use R = 0.0821 Lt.-atm/mole-K
Osmosis Law:
Terminology





Isotonic: having the same osmotic pressure
Hypertonic: having a higher osmotic pressure
Hypotonic: having a lower osmotic pressure
Hemolysis: the process that ruptures a cell
placed in a solution that is hypotonic to the
cell’s fluid
Crenation: the opposite effect
Colloids
In a solution the particles are small: single
molecules or ions
 In colloids the particles are large:
collection of molecules or a giant molecule
 Particle size 10 to 2000Å
 Colloids scatter light because of large
particles

Types of Colloids
Gas as the dispersion agent
Dispersed phase
Gas
Liquid
Solid
none
aerosol
aerosol
(all are solutions)
fog
smoke
Types of Colloids
Liquid as the dispersion agent
Dispersed phase
Gas
Liquid
Solid
foam
emulsion
sol
whipped cream
milk
paint
Types of Colloids
Solid as the dispersion agent
Dispersed phase
Gas
solid foam
marshmellow
Liquid
solid emulsion butter
Solid
solid sol
ruby glass
Stability of colloids
What prevents the dispersed particles
from coalescing?
 Hydrophobic colloids: the dispersed phase
reacts with the water with forces similar to
those present in solutions
 Hydrophobic: The particles have ions on
the surface; the charges keep the particles
from coming together

Detergents
Long chain molecules
 One end is hydrophobic
 The other is hydrophilic
 The hydrophobic ends attach to dirt
particles (grease)
 The hydrophilic ends keep the particles in
solution
