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SOLUTIONS
Introduction:
In this chapter, we will discuss about liquid solutions and their
formation. This will be followed by studying the properties of solutions, like
vapour pressure and colligative properties. We will begin with types of
solutions and expressions for concentration of solutions in different units.
Thereafter, we will state and explain Henry’s law and Raoult’s law,
distinguish between ideal and non-ideal solution and deviation of real
solutions from Raoult’s law. We will also discuss abnormal colligative
properties alongwith association and dissociation of solute.
Types of Solutions
All the three states of matter (solid, liquid and gas) may behave either
as solvent or solute. When a solution is composed of only two chemical
substances, it is termed as binary solution. Depending upon the state of solute
or solvent, binary solutions can be classified as
Some Important Definitions
·
Mixture - When two or more chemically non-reacting substances are mixed, they
form mixture.
·
Heterogeneous
Mixture - It consists of distinct phases, and the observed properties are just
the sum of the properties of individual phases.
·
Homogeneous
Mixture - It consists of a single phase which has properties that may differ from
one of the individual components.
·
Solution - The homogeneous mixture of two or more components such that at least
one component is a liquid is called solution.
·
Solvent - It is the constituent of solution which has same physical state as that
of solution and generally present in greater amount than all the other
components.
·
Solute - The component of a solution other than solvent is called solute, may or
may not have same physical state as that of solution. Generally it is in
smaller amount.
Example - In a sugar syrup (liquid solution) containing 60% sugar (solid) and 40%
water (liquid), water is termed as solvent, due to same physical state as that
of solution.
Expressing the Strength of Solution
For a given solution the amount of solute dissolved per unit volume of
solution is called concentration of solute. Strength of solution is the amount
of solute in grams dissolved in one litre of solution. It is generally
expressed in g/litre.
Other methods of expressing the strength of solution are:
1. Mass percentage –
2. Volume percentage –
3.
Molality (m) - It is no. of moles of solute dissolved in 1 kg of the solvent.
4.
Molarity (M) - It is no. of moles of solute dissolved in 1 litre of solution.
5.
Normality (N) - It is no. of gram-equivalents of solute dissolved in 1 litre of solution
6.
Formality - Ionic solutes do not exist in the form of molecules. These molecular
mass is expressed as Gram-formula mass. Molarity for ionic compounds is
actually called as formality.
7.
Mole fraction –
For a binary
solution,
mole fraction of solute + mole fraction of solvent = 1.
8.
Parts per million
(ppm) –
It is defined in two ways
ppm = mass fraction
× 106
ppm = mole fraction × 106Colligative Properties
The properties of dilute solution which depends only on number of
particles of solute (molecules or ions) present in the solution and not on
their nature, are called colligative properties. The important colligative properties are;
1.
Relative lowering of vapour pressure
2.
Elevation of boiling point
3.
Depression in freezing point
4.
Osmotic pressure
i.
Relative
Lowering of Vapour Pressure
When a non-volatile solute is added to a solvent, its vapour pressure
gets lowered. If this pressure is divided by pressure of pure solvent, this is called
relativelowering
of vapour pressure.
According to
Raoult’s law,
where, WA = weight of solute
WB = weight of solvent
MA = molecular weight of solute
MB = molecular weight of solvent
ii. Relative Elevation of boiling point
A liquid boils at the temperature at which its vapour pressure is equal
to the atmospheric pressure. The boiling point of a solution of non-volatile
solute is always higher than that of the boiling point of pure solvent in which
the solution is prepared. Similar to lowering of vapour pressure, the elevation
of boiling point also depends on the number of solute particles rather than
their nature.
Let T° be the boiling point of pure solvent and T be the boiling point of
solution. The increase in boiling point ΔTb = T – T° is known as elevation in boiling point.
For dilute solutions, the ΔTb is directly proportional to the
molal concentration of the solute in a solution. Thus
ΔTb∝ m
ΔTb = Kbm
Kb is molal elevation
constant (Ebullioscopic constant). The unit of Kb
is K kg mol–1.
Substituting the value of molality in above equation, we get
Where, w1 = mass of solvent, w2 = mass of solute
and M2 = molar mass of solute
iii. Depression in freezing point
Freezing point is the temperature at which vapour pressure of liquid
phase becomes same as that of solid phase. The decrease in freezing point of a
solvent on the addition of a non-volatile solute is known as depression in freezing point.
Let T° be the freezing point of pure solvent and T be the freezing point
of solution. The decrease in freezing point ΔTf = T° -T is known as
depression in freezing point.
For dilute solutions, the ΔTf is directly proportional to the
molal concentration of the solute in a solution. Thus
ΔTf∝ m
ΔTf = Kf⋅ m
Here Kf is molal depression constant or cryoscopic constant
Substituting the value of molality in above equation, we get
Where, w1 = mass of solvent, w2 = mass of solute
and M2 = molar mass of solute
iv. Osmotic pressure
Osmosis is the spontaneous flow of the solvent molecules from a less
concentrated solution (dilute) to a more concentrated solution through a semi-permeable membrane. The driving force of osmosis is called osmotic
pressure. Osmotic pressure may be defined as
“the minimum excess pressure that has to be applied on the solution to prevent
the osmosis".
Osmotic pressure of a solution ∝ molar concentration
of solute in that solution
π ∝ c
π = cRT
where, R = Gas constant = 0.0821 lit atm K-1 mole-1
T = Temperature
c = Molar concentration
WB = wt. of solute
MB = Molar mass of solute
van’t Hoff Factor
To calculate the extent of association or dissociation, van’t Hoff in
1886 introduced a factor ‘i’ called van’t Hoff factor. van’t Hoff factor ‘i’ is defined as ratio of the experimental value of colligative
property to the calculated value of colligative property.
Solubility
Solubility of a substance is its maximum amount that can be dissolved in a
specified amount of solvent at a specified temperature. It depends upon the
nature of solute and solvent as well as temperature and pressure. Let us
consider the effect of these factors in solution of a solid or a gas in a liquid.
1. Solubility of Solid in Liquid
A solute dissolves in a solvent if the
intermolecular interactions are similar in them, i.e., like dissolves like.
Polar solute dissolves in polar solvent and non-polar solute in non-polar
solvent. For e.g., sodium chloride and sugar dissolves readily in water and
napthalene and anthracene dissolves readily in benzene.
Solute + Solvent ⟶ Solution
i.
Dissolution: When a solid solute is added to the solvent, some solute dissolves and
its concentration increases in solution. This process is called dissolution.
ii.
Crystallization: Some solute particles collide with solute particles in solution and get
separated out. This process is called crystallization.
iii.
Saturated
solution: Such a solution in which no more solute can be dissolved at the same
temperature and pressure is called a saturated
solution.
iv.
Unsaturated
solution: An unsaturated solution is one in which more solute can be dissolved at
the same temperature.
v.
Effect of
temperature: In general, if in a nearly saturated solution, the
dissolution process is endothermic, the solubility should increase with rise in
temperature, if it is exothermic, the solubility should decrease with rise in
temperature.
vi.
Effect of
pressure: Solids and liquids are highly incompressible, so pressure does not have
any significant effect on solubility of solids and liquids.
vii.
Supersaturated
solution: When more solute can be dissolved at higher temperature in a saturated
solution, then the solution becomes supersaturated.
2.
Solubility
of Gas in Liquid
All
gases are soluble in water as well as in other liquids to a greater or lesser
extent. The solubility of a gas in liquid depends upon the following factors
Nature of the gas, Nature of solvent, Temperature and Pressure.
Generally, the gases which can be easily
liquified are more soluble in common solvents. For e.g., CO2 is more soluble
than hydrogen or oxygen in water. The gases which are capable of forming ions
in aqueous solutions are much more soluble in water than other solvents. For
e.g., HCl and NH3 are highly soluble in water but not in organic solvents (like
benzene) in which they do not ionize.
i.
Effect of
temperature: Solubility of most of the gases in liquids decreases
with rise in temperature. In dissolution of a gas in liquid, heat is evolved
and thus this is an exothermic process. The dissolution process involves
dynamic equilibrium and thus follows Le Chatelier’s principle. As dissolution is exothermic the solubility of gas should
decrease with rise in temperature.
ii.
Effect of
pressure: Henry’s law: At constant temperature, the
solubility of a gas in a liquid is directly proportional to the pressure of the
gas.
p = KH x,
KH = Henry’s law
constant.
Applications of Henry’s law
1. In manufacture of soft drinks and soda water, CO2 is passed at
high pressure to increase its solubility.
2. To minimise the painful effects accompanying the decompression of deep
sea divers. O2 diluted with less soluble. He gas is used as
breathing gas.
3. At high altitudes, the partial pressure of O2 is less then
that at the ground level. This leads to low concentrations of O2 in
the blood of climbers which causes ‘anoxia’.
Vapour Pressure of Solution
It is the
pressure exerted by vapour on the surface layer of liquid at equilibrium
between vapour and liquid.
Factors affecting
Vapour Pressure
i.
Nature of liquid
- Liquid with higher intermolecular attraction forces form less amount of
vapour and hence lower vapour pressure and vice-versa.
ii.
Temperature - Vapour pressure increases with temperature of liquid. This is because, as
temperature increases, kinetic energy of the molecules increases, hence, more
molecules leave the surface of the liquid and come into vapour phase.
Raoult’s Law
According to Raoult’s law, for a solution of volatile liquids, the
relative lowering of vapour pressure of solution is directly proportional to
its mole fraction of dissolved solvent in solute.
Ideal
Solution
The solution which obeys Raoult’s law at all compositions of solute and solvent and at all temperature is
called an ideal solution.
Ex- Benzene and Toluene, n-hexane and
n-heptane.
Characteristics
of an ideal Solution
1.
Raoult’s law is obeyed by it.
2.
ΔHmixing = 0 i.e., no heat
should be absorbed or evolved during mixing.
3.
ΔVmixing = 0, i.e., no change
in volume (expansion or contraction) on mixing.
Non-ideal Solution
Those solutions which deviate from ideal behaviour are called non-ideal solutions or real solutions. Acetone and CS2, Acetone and C2H5OH
Characteristics
of an non-ideal Solution
1.
Raoult’s law is not obeyed by it.
2.
ΔHmixing ≠ 0 i.e., solution
may absorb or release heat.
3.
ΔVmixing ≠ 0 i.e., solution
may expand or contract on mixing of solute and solvent.
Azeotropic Mixture
At the constant boiling temperature, liquid mixture vapouries without
change in composition and the condensate contains same composition, i.e.,
mixture distills like a pure liquid, which has same composition. At this point,
solution or mixture is called an azeotropic
mixture.