Chemistry class 12 ncert solutions chapter 2 notes -
Solution
Solutions are homogeneous mixtures of two or more than two components. The substances forming the solution are called components of the solution. The component present in smaller amount is called solute and the other present in larger amounts in which solute is dissolved, is called solvent. The solutions containing two components are binary solutions, e.g. salt solution.
1.1. Types of Solution
Depending upon the nature of solute and solvent, solutions are classified as: Gaseous solutions (in which gas acts as solvent, e.g, mixture of O2 and N2), liquid solutions (in which liquid acts as solvent, e.g. O, dissolved in water), solid solutions (in which solid acts as solvent or present in large amounts, e.g. alloys).
Different Methods to Express the Concentration of Solution
The concentration of the solutions can be expressed as follows:
Molarity (M) -
It is defined as the number of moles of solute dissolved in one liter (or one cubic decimetre) of the solution.
Molarity (M) = Number of moles of solute x 1000 / Volume of solution (mL)
Moles of solute = W₂/M2,
Where,
W₂ = mass of solute(g)
M₂ = molar mass of solute)
Volume: Mass/Density
•Molarity is a function of temperature and changes with change in temperature because volume depends upon temperature.
Molality (m) -
It is defined as the number of moles of solute per kilogram of the solvent.
Molality (m)= Number of moles of solute x 1000/ Mass of solvent (g)
Normality (N) -
It is the number of gram equivalents of the solute dissolved in one liter of the solution. Number of gram equivalents of solutex 1000
Normality (N)= Number of gram equivalent of solute x 1000 / Volume of solution (mL)
Mole Fraction(x) -
It is the number of moles of one component to the total number of moles of all the components present in the solution. For a binary solution having solvent 1 and solute 2.
NOTE - Mass %, ppm, molality and mole fraction do not change with change in temperature while molarity decreases with rise in temperature.
Parts per million (ppm) -
When a solute is present in trace quantities, the concentration is expressed in parts per million.
Parts per million =Number of parts of the component x 10⁶ / Total number of parts of all the components of the solution
Mass percent (w/w) -
The mass percentage of a component in a given solution is the mass of the component per 100 g of the solution.
Mass percent = mass of component × 100 / total mass of solution
Volume per cent (V/V) -
The volume percentage is the volume of the component per 100 parts by volume of the solution.
Volume percent= Volume of the component 100 / total volume of solution
Mass by volume percentage [ w/V] -
Mass by volume percentage is the mass of solute dissolved in 100 ml of the solution.
Mass by volume % = mass of solute × 100
Volume of solution
Relation between molarity and molality
Molarity(m) = M × 1000
(1000×d) - (M×M2)
1.2 Solubility
Solubility of a substance is its maximum amount of solute that can be dissolved in a specified amount of solvent (at a specified temperature). It depends on the nature of solute, solvent, temperature and pressure. Depending on solubility, the solution can be saturated or unsaturated.
Saturated Solution
Saturated solution is the solution in which no more solute can be dissolved at the same temperature and pressure.
Unsaturated Solution
An unsaturated solution is the one in which more solute can be dissolved at the same temperature.
NOTE-
On dissolving the solid solute in a solvent, its concentration increases, this is dissolution. While when some solute particles in solution collide with the other solid solute particles and get separated out of the solution, this process is called crystallisation.
Dynamic Equilibrium
Dynamic equilibrium is the condition when number of solute particles going into the solution is equal to the solute particles separating out, i.e. dissolution and crystallisation occur at the same rate.
Solute + Solvent →←solution
From Le-Chatelier's principle, if in a nearly saturated solution, the dissolution process is endothermic (Agol H >0), the solubility should increase with rise in temperature and if it is exothermic (AsolH <0), the solubility should decrease.
1.3 Henry's Law
• It states that, at a constant temperature, the solubility of a gas in liquid is directly proportional to the partial pressure of the gas present above the surface of liquid or solution.
S∞p or p= K(h).S
Unit of solubility is same as concentration.
• The partial pressure of the gas in vapour phase (p) is directly proportional to the mole fraction (x) of the gas in the solution.
pxx or p = KH X (where, KH is Henry's law constant)
If we draw a graph between partial pressure of the gas versus mole fraction of the gas in solution, we get straight line whose slope is given by KH
Higher the value of KÏ€ at a given pressure, the lower is the solubility of the gas in the liquid. Solubility of gases increases with increase of pressure.
Applications of Henry's Law
(i) To increase the solubility of CO₂ in soft drinks and soda water, the bottle is sealed under high pressure.
(ii) To avoid bends and the toxic effects of high concentrations of N₂ in the blood, the cylinders used by scuba divers are filled with air diluted with He.
2.1 Raoult's Law for Volatile Solute
This law states that for a solution of volatile liquids, the partial vapour pressure of each component in the solution is directly proportional to its mole
fraction. For component 1, p, o x, or p₁ = pix₁ Similarly, for component 2, P₂ = P2x20
Ptotal = P1 + P₂ = P1x1 + P2x2;
Ptotal = (1-X2) Pi + x2P2
Ptotal = Pi + (p2-pi)x, as (x, + x2 =1)
where,
p₁ and p2 = vapor pressures of pure component 1 and 2 respectively.
If y, and y₂ are the mole fractions of the component 1 and 2 respectively in vapour phase, then
P₁ = y1 x Ptotal.
Similarly, P₂ = y2 × Ptotal
Raoult's Law a Special Case of Henry's Law -
In the solution of a gas in a liquid, if one of the component is so volatile that it exists as a gas, then we can say that Raoult's law becomes a special case of Henry's law in which KH becomes equal to p°.
2.2 Ideal and Non-ideal Solutions
If a non-volatile solute is added to a solvent, then the vapour pressure of the solution decreases. This is because in solution, the surface has both solute and solvent molecules there by the fraction of the surface covered by the solvent molecules gets reduced, thus reducing the vapour pressure. On this basis, solution can be classified as ideal or non-ideal.
Ideal Solutions
Ideal solutions obey Raoult's law over entire range of concentration. For these solutions, Amix H =0 and Amix V=0. In binary solutions, if A-B interactions are nearly equal to A-A or B-B interactions then it is an ideal solution. Solutions of n-hexane and n-heptane, bromoethane and chloroethane, benzene and toluene, etc., are nearly ideal in behaviour.
Non-ideal Solutions
Non-ideal solutions do not obey Raoult's law over entire range of concentration. For such solutions,
Amix H #0 and Amix V#0 The vapour pressure is either higher (positive deviation from Raoult's law) or lower (negative deviation from Raoult's law) than that predicted by Raoult's law.
These are as follows:
Positive Deviation and Negative Deviation
In case of positive deviation from Raoult's law (e.g. mixture of ethanol and acetone, carbon disulphide and acetone), A-B (i.e. solute-solvent) interactions are weaker than those of A-A (solute-solute) or B-B (solvent-solvent) interactions, while in case of negative deviation from Raoult's law (e.g. mixture of phenol and aniline, chloroform and acetone), A-B interactions are stronger than those of A-A or B--B interactions.
The solutions showing positive deviation and negative deviation from Raoult's law are shown in fig.
(a) and (b) respectively.
(i) For positive deviation,
ΔΗ = Positive, ∆V Positive. = mix mix
(ii)For negative deviation,
ΔΗ = Negative, ∆Vmix = Negative. mix
Azeotropes
The binary mixtures (solutions) that have the same composition in liquid and vapour phase and boil at constant temperature like a pure liquid are called azeotropes or
azeotropic mixtures. The solutions which show large negative deviation from Raoult's law, form maximum boiling azeotropes. e.g. Nitric acid-water mixture and the solutions which show large positive deviation from Raoult's law, form minimum boiling azeotropes. e.g. ethanol-water mixture.
3.1 Colligative Properties
The properties of solutions which depend only on the number of solute particles, irrespective of their nature relative to the total number of particles present in solution are known as colligative properties.
Colligative properties & number of particles in the solution ∞ 1
Molar mass of solute
Colligative properties are as follows:
Relative Lowering of Vapour Pressure
When a non-volatile solute is dissolved in a solvent, vapour pressure of the solution becomes lower than that of the pure solvent which is known as lowering of vapour pressure. The relative lowering of vapour pressure of a solution containing the non-volatile solute is equal to the mole fraction of the solute at a given temperature.
Elevation in Boiling Point
The boiling point of a liquid is that temperature at which its vapour pressure becomes equal to the atmospheric pressure. The boiling point of a solution is always higher than the boiling point of the pure solvent in which the solution is prepared. This increase in boiling point is termed as elevation in boiling point.
If T°b is boiling point of pure solvent and Tb, is the boiling point of solution, then the elevation in boiling point is represented as,
∆Tb = Tb -T°b
Elevation in boiling point,
∆Tb ∞ m (where m is molality)
or ∆Tb = Kbm
or. ∆Tb = Kb × W2 × 1000
M2 × W1(g)
where, K, = boiling point elevation constant or molal elevation constant or ebullioscopic constant having unit K kg mol-1.
Depression in Freezing Point
The freezing point of a substance is that temperature at which the vapour pressure of the substance in its liquid phase is equal to its vapour pressure in the solid phase. When a non-volatile solute is added to a solvent, the freezing point of the solution is always lower than that of pure solvent as the vapour pressure of the solvent decreases in the presence of non-volatile solute. This difference in freezing point is known as depression of freezing point, i.e.
∆Tf = T°f - Tf.
where, T°f is the freezing point of pure solvent and Tf is the freezing point of solution.
Depression of freezing point,
∆Tf ∞ m or ∆Tf = Kfm
∆Tf = Kf × W2 × 1000
M2 × WÃ (g)
where, K freezing point depression constant or molal depression constant or cryoscopic constant, having unit K kg mol¹
M2 = Kf × W2 × 1000
∆Tf × W1(g)
Also, Kf = R × M × (T°f)²
1000 × ∆fysH
where, R and M, are gas constant and molar mass of the solvent, respectively. T°f and T°b are the freezing and boiling point of the pure solvent, respectively (in K). ∆fusH and ∆vapH are enthalpies for the fusion and vaporization of the solvent respectively.
Osmosis and Osmotic Pressure
The process of flow of solvent molecules from solution of lower concentration to solution of higher concentration through semipermeable membrane is known as osmosis. The hydrostatic pressure which develops on account of osmosis is called osmotic pressure or the excess pressure that must be applied on the solution to prevent osmosis is called osmotic pressure.
Osmotic pressure (Ï€) is directly proportional to molarity (C) of the solution at a given temperature T.
Ï€ = C
Ï€ = CRT or Ï€ = n2 × RT
V
Osmotic pressure is used to determine molar masses of proteins, polymers and other macromolecules.
(i) Two solutions having same osmotic pressure at a given temperature are called isotonic solutions.
(ii) A solution having lower osmotic pressure than the other solution is called hypotonic while, the one with higher osmotic pressure is called hypertonic.
(iii) People taking salty food, experience water retention in tissue cells and intercellular spaces due to osmosis. The resulting puffiness is called edema.
(iv) Reverse osmosis If a pressure larger than the osmotic pressure is applied to the solution side, then the pure solvent flows out of the solution through the semipermeable membrane. This phenomenon is called reverse osmosis. It is used for the desalination of sea water.
Abnormal Molar Mass
For the substances undergoing association, dissociation, etc. in the solution, molecular mass determined from colligative properties is different (either lower or higher) from expected value. This is known as abnormal molar mass. This change can be known by using the van't Hoff factor.
3.2 van't Hoff Factor
It is the ratio of the experimental value of colligative property to the calculated value of the colligative property.
It is used to find out the extent of dissociation or association.
van't Hoff factor,
i = Normal molar mass
Abnormal molar mass
= Total number of moles of particles after association/ dissociation
Number of moles of particles before association/dissociation
= Observed value of colligative property
Calculated value of colligative property