Solution chemistry

S H A R D A P U B L I C S C H O O L , A L M O R A 1
SOLUTION CHEMISTRY
CLASS XII, UNIT:2
SUBJECT EXPERT
DR. TANUJA NAUTIYAL
SHARDA PUBLIC SCHOOL
ALMORA

Solutions are homogeneous mixtures of two or more pure substances.
In a solution, the solute is dispersed uniformly throughout the solvent.
Table gives examples of each type.
Solution

How Does a Solution Form?
1. Solvent molecules attracted to surface ions.
2. Each ion is surrounded by solvent molecules.
3. Enthalpy (DH) changes with each interaction broken or
formed.
If the solvent is water, the ions are hydrated. The
intermolecular force here is ion-dipole. The ions are
solvated (surrounded by solvent)
.

How Does a Solution Form?

To determine the enthalpy change, we divide
the process into 3 steps.
1. Separation of solute particles.
2. Separation of solvent particles to make ‘holes’.
3. Formation of new interactions between solute
and solvent.
ENERGY CHANGES IN SOLUTION

ENERGY CHANGES IN SOLUTION
DH1Separation of
solute molecules
DH2 Separation of
solvent molecules
DH3 Formation of solute-
solvent interactions

ENTHALPY CHANGES IN SOLUTION

The intermolecular attractions that hold molecules together
in liquids and solids also play a central role in the formation
of solutions. When one substance (the solute) dissolves in
another (the solvent), particles of the solute disperse
throughout the solvent.
The solute particles occupy positions that are normally
taken by solvent molecules. The ease with which a solute
particle replaces a solvent molecule depends on the relative
strengths of three types of interactions:
solvent-solvent interaction, solute-solute interaction,
solvent-solute interaction

ENTHALPY CHANGES DURING DISSOLUTION
DHSoln = DH1+ DH2+ DH3
This equation is an application of Hess’s law.
The enthalpy of solution, DHsoln, can be either positive or
negative.
 DHsoln (MgSO4) = -91.2 kJ/mol --> Exothermic
 DHsoln (NH4NO3) = 26.4 kJ/mol --> Endothermic

 If the solute-solvent attraction is stronger than
the solvent-solvent attraction and solute-solute
attraction, the solution process is favourable, or
exothermic (ΔHsoln < 0).
 If the solute-solvent interaction is weaker than
the solvent-solvent and solute-solute
interactions, then the solution process is
endothermic (ΔHsoln > 0).

 A saturated solution contains the maximum amount of
a solute that will dissolve in a given solvent at a specific
temperature.
 An unsaturated solution contains less solute than it
has the capacity to dissolve.
 A third type, a supersaturated solution, contains more
solute than is present in a saturated solution.
Supersaturated solutions are not very stable.
Types of Solution

Factors Affecting Solubility
 Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.
Example: Vitamin A is soluble in nonpolar compounds (like fats).
Vitamin C Vitamin C is soluble in water.
Nature of Solute and solvent: “like dissolves like”

Factors Affecting Solubility
Glucose (which has hydrogen bonding) is very soluble in water.
 Cyclohexane (which only has dispersion forces) is not water-
soluble.
Intermolecular forces = H-bonds; dipole-dipole; dispersion Ions in
water also have ion-dipole forces soluble in water.
The stronger the intermolecular attractions between solute
and solvent, the more likely the solute will dissolve.

Gases in Solution:
the solubility of gases in water increases with increasing mass. The
reason is Larger molecules have stronger dispersion forces.
Solubility of some gases in water at 200C, with 1 atm Gas Pressure
N2
Solubility (M)
0.69 X 10-3
CO
Solubility (M)
1.04 X 10-3
O2
Solubility (M)
1.38 X 10-3
Ar
Solubility (M)
1.50 X 10-3
Kr
Solubility (M)
2.79 X 10-3

Effect of pressure on solubility:
The solubility of liquids and solids does not change appreciably
with pressure. But, the solubility of a gas in a liquid is directly
proportional to its pressure.
Henry’s Law:
Sg is the solubility of the gas;
k is the Henry’s law constant for that gas in that solvent;
Pg is the partial pressure of the gas above the liquid.
Sg = kPg

Effect of temperature:
Generally, the solubility of solid solutes in liquid
solvents increases with increasing temperature.

Effect of temperature:
The opposite is true of gases. Higher temperature drives
gases out of solution.
Carbonated soft drinks are more “bubbly” if stored
in the refrigerator.
Warm lakes have less O2 dissolved in them than
cool lakes.

Ways of Expressing Concentrations of
Solutions
 Mass percentage:
= (mass of A in solution / total mass of solution) X 100
 Parts per Million (ppm):
= (mass of A in solution/total mass of solution) X106
 Parts per Billion (ppb)
= (mass of A in solution/total mass of solution) X109
 Mole Fraction (X): The mole fraction of a component of a
solution, say, component A, is written XA
XA = (moles of A / total moles in solution)

Ways of Expressing Concentrations of
Solutions
Molarity (M):the number of moles of solute in 1 L of solution;
that is,
(moles of solute / L of solution)
Molality (m): Molality is the number of moles of solute dissolved
in 1 kg (1000 g) of solvent—that is, (moles of solute / kg of
solvent)
Normality (N): the gram equivalents in 1 L of solution;
(gram equivalents of solute / L of solution)

COLLIGATIVE PROPERTIES
Colligative properties depend only on the number of solute particles
present, not on the identity of the solute particles. Among colligative
properties are
Vapor pressure lowering
Boiling point elevation
Melting point depression
Osmotic pressure

VAPOUR PRESSURE
As solute molecules are added to a solution, the
solvent become less volatile (=decreased vapor
pressure).
Solute-solvent interactions contribute to this effect.
Therefore, the vapor pressure of a solution is lower
than that of the pure solvent.
If a solute is non volatile (that is, it does not have a
measurable vapor pressure), the vapor pressure of its
solution is always less than that of the pure solvent.

RAOULT’S LAW
Thus, the relationship between solution vapor pressure and solvent
vapor pressure depends on the concentration of the solute in the
solution. This relationship is expressed by Raoult’s law,
Raoult’s Law: It states that the vapor pressure of a solvent over a
solution, PA is given by the vapor pressure of the pure solvent,
P°A, times the mole fraction of the solvent in the solution, = XA
PA = XAPA

RAOULT’S LAW
There are two cases we can consider.
1. Volatile solute – both solvent and solute are found in the vapor above the
solution. A solution of ethanol in is an example.
2. Non-volatile solute – only the solvent has a vapor pressure. The solute does
not contribute to the pressure so there is a “vapor pressure lowering”.

TOTAL PRESSURE FOR IDEAL SOLUTION
P1 = X1P01 Ptotal = P1 + P2
P1 = X1P01 P2 = X2P02

FOR AN IDEAL SOLUTION
The total vapor pressure over an ideal solution is given by
Ptotal = P1 + P2
In a solution containing only one solute, X1 = 1 – X2, where X2 is the mole
fraction of the solute.
= X1P1 + X2P2
= (1 - X2)P1 + X2P2
= P1+ X2 (P2- P1)
A plot of the total pressure has the form of a straight line.

BOILING POINT ELEVATION
Solute-solvent interactions also cause solutions to have higher boiling points
and lower freezing points than the pure solvent.
The boiling point elevation (DTb ) is defined as the boiling point of the solution
(Tb) minus the boiling point of the pure solvent (T°b):
DTb = Tb - T°b
The value of DTb is proportional to the vapor-pressure lowering, and so is also
proportional to the concentration (molality) of the solution. That is, D Tb α m
The change in boiling point is proportional to the molality of the solution:
where Kb is the molal boiling point elevation constant, a property of the
solvent.
DTb = Kb  m

FREEZING POINT DEPRESSION
A non scientist may remain forever unaware of the boiling-point
elevation phenomenon, but a careful observer living in a cold climate
is familiar with freezing-point depression. Ice on frozen roads and
sidewalks melts when sprinkled with salts such as NaCl or CaCl2.
This method of thawing succeeds because it depresses the freezing
point of water.

The freezing point depression (DTf) is defined as the freezing point of the pure
solvent (T °f) minus the freezing point of the solution (T f):
DTf = Tf - T °f
Because T°f . Tf, DT°f is a positive quantity. Again, DTf is proportional to the
concentration of the solution:DTf α m
DTf = Kf  m
where m is the concentration of the solute in molality units, and Kf is the molal
freezing-point depression constant. We use molality since the solvent properties
do matter and therefore, we would like to calculate the result in relation to a fixed
amount of solvent (1 kg). The constant Kf depends on the solvent and is a
function of the enthalpy of freezing and the molar mass.

DTb = Kb  m
DTf = Kf  m
Here Kf is the molal freezing point depression constant of the solvent.
 The similarity between freezing point depression and boiling point ele
vation is not accidental.
 They both depend on the shift in the equilibrium caused when a solut
ion is formed and changes the equilibrium of the liquid with respect t
o the solid or vapor.
 In both equations, DT does not depend on what the solute is, but
only on how many particles are dissolved.

MOLAL BOILING-POINT ELEVATION AND FREEZING-POINT DEPRESSION
CONSTANTS OF SEVERAL COMMON LIQUIDS
Solvent Freezing Kf Boiling Kb
(°C)* (°C/m) (°C)* (°C/m)
Water 0 1 .86 100 0.52
Benzene 5.5 5.12 80.1 2.53
Ethanol −117.3 1.99 78.4 1.22
Acetic acid 16.6 3.90 117.9 2.93
Cyclohexane 6.6 20.0 80.7 2.79

COLLIGATIVE PROPERTIES OF ELECTROLYTES
The study of colligative properties of electrolytes requires a slightly
different approach than the one used for the colligative properties of
nonelectrolytes. The reason is that electrolytes dissociate into ions in
solution, and so one unit of an electrolyte compound separates into
two or more particles when it dissolves. Because these properties
depend on the number of particles dissolved, solutions of electrolytes
(which dissociate in solution) show greater changes than those of
nonelectrolytes. e.g. NaCl dissociates to form 2 ion particles; its
limiting Van’t Hoff factor is 2.

VAN’T - HOFF FACTOR:
One mole of NaCl in water does not really give rise to two moles of ions.
Some Na+ and Cl− re associate as hydrated ion pairs, so the true
concentration of particles is somewhat less than two times the
concentration of NaCl. Some Na+ and Cl− re associate as hydrated ion
pairs, so the true concentration of particles is somewhat less than two
times the concentration of NaCl.
Re association is more likely at higher concentration.
Therefore, the number of particles present is concentration dependent.

THE VAN’T HOFF FACTOR
We modify the previous equations by multiplying by the van’t Hoff factor, i
i = actual number of particles in solution after dissociation /number of formula
units initially dissolved in solution.
i = 1 for nonelectrolytes
DTf = Kf .m.i

OSMOSIS
 Semipermeable membranes allow some particles to pass through while
blocking others. In biological systems, most semipermeable
membranes (such as cell walls) allow water to pass through, but block
solutes.
 In osmosis, there is net movement of solvent from the area of higher
solvent concentration (lower solute concentration) to the are of lower
solvent concentration (higher solute concentration.
 Water tries to equalize the concentration on both sides until pressure is
too high.

OSMOSIS

MOLAR MASS FROM COLLIGATIVE PROPERTIES
We can use the effects of a colligative property such as osmotic pressure to
determine the molar mass of a compound.
The Vant Hoff’s equation can be modified to form used for the determination of
molar mass by osmometry.
𝝅 = 𝑪𝑹𝑻
𝝅 =
𝒘
𝑴 𝒎
. 𝐑𝐓
Here we related to the concentration C in mol/lt to the concentration w in gms/lt
and the molar mass Mm in grams/mole.
The experimental configuration uses the measurement of height as an estimate
of the osmotic pressure. The equation  = ρ gh is used (h =  / ρ g).

OSMOTIC PRESSURE
The pressure required to stop osmosis, known as osmotic pressure, , is
where M is the molarity, n expresses the number of moles of solute, an d n/V, of
the solution, R is the gas constant (0.0821 L.atm/K.mol),. This equation is called
the Van't Hoff equation for osmotic pressure.
 If the osmotic pressure is the same on both sides of a membrane (i.e., the
concentrations are the same), the solutions are isotonic.
 If the solute concentration outside the cell is greater than that inside the cell,
the solution is hypertonic.
 = MRT=
𝒏
𝑽
𝑹𝑻

OSMOSIS IN CELLS
 If the solute concentration outside the cell is less than that inside the
cell, the solution is hypotonic.
 Water will flow into the cell, and hemolysis results.

COLLOIDS
Suspensions of particles larger than individual ions or molecules, but too
small to be settled out by gravity.

Among the most important colloids are those in which the dispersing
medium is water. Such colloids are divided into two categories called
hydrophilic, or water-loving, and hydrophobic, or water-fearing.
Hydrophilic colloids are usually solutions containing extremely large
molecules such as proteins. In the aqueous phase, a protein like
haemoglobin folds in such a way that the hydrophilic parts of the
molecule, the parts that can interact favourably with water molecules
by ion-dipole forces or hydrogen-bond formation, are on the outside
surface

TYNDALL EFFECT
Colloidal suspensions can scatter rays of light. This phenomenon is known
as the Tyndall Effect
When a beam of light passes through a colloid, it is scattered by the
dispersed phase No such scattering is observed with ordinary solutions
because the solute molecules are too small to interact with visible light.
Another demonstration of the Tyndall effect is the scattering of sunlight by
dust or smoke in the air.

COLLOIDS IN BIOLOGICAL SYSTEMS
Some molecules have a polar, hydrophilic (water-loving) end and a
nonpolar, hydrophobic (water-hating) end.

USING COLLIGATIVE PROPERTIES TO DETERMINE MOLAR MASS
The colligative properties of nonelectrolyte solutions provide a means of
determining the molar mass of a solute. Theoretically, any of the four
colligative properties is suitable for this purpose.
DTb = Kb . m
DTf = Kf . m
 = MRT=
𝒏
𝑽
𝑹𝑻
PA = XAPA

SUMMARY
 Solutions are homogeneous mixtures of two or more substances, which
may be solids, liquids, or gases.
 The ease of dissolving a solute in a solvent is governed by
intermolecular forces. Energy and the disorder that results when
molecules of the solute and solvent mix to form a solution are the forces
driving the solution process.
 The concentration of a solution can be expressed as percent by mass,
mole fraction, molarity, and molality.
 Increasing temperature usually increases the solubility of solid and liquid
substances and usually decreases the solubility of gases in water.

SUMMARY
 According to Henry’s law, the solubility of a gas in a liquid is directly
proportional to the partial pressure of the gas over the solution.
 Raoult’s law states that the partial pressure of a substance
A over a solution is equal to the mole fraction (XA) of A times the vapor
pressure (P°A) of pure A. An ideal solution obeys Raoult’s law over the
entire range of concentration. In practice, very few solutions exhibit ideal
behaviour.
 Vapor-pressure lowering, boiling-point elevation, freezing point
depression, and osmotic pressure are colligative properties of solutions;
that is, they depend only on the number of solute particles that are
present and not on their nature.

SUMMARY
 In electrolyte solutions, the interaction between ions leads to the
formation of ion pairs.
 The Van’t Hoff factor provides a measure of the extent of
dissociation of electrolytes in solution.
 A colloid is a dispersion of particles (about 1 X103 pm to 1 X 106
pm) of one substance in another substance.
 A colloid is distinguished from a regular solution by the Tyndall
effect, which is the scattering of visible light by colloidal particles.
 Colloids in water are classified as hydrophilic colloids and
hydrophobic colloids.

SHARDA PUBLIC SCHOOL
Presentation by
Dr. Tanuja Nautiyal

Solution chemistry

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