Almost all processes in body occur in some kind of liquid solutions.

Objectives

After studying this Unit, you will be able to

• describe the formation of different types of solutions;
• express concentration of solution in different units;
• state and explain Henry’s law and Raoult’s law;
• distinguish between ideal and non-ideal solutions;
• explain deviations of real solutions from Raoult’s law;
• describe colligative properties of solutions and correlate these with molar masses of the solutes;
• explain abnormal colligative properties exhibited by some solutes in solutions.

In normal life we rarely come across pure substances.

Most of these are mixtures containing two or more pure substances. Their utility or importance in life depends on their composition. For example, the properties of brass (mixture of copper and zinc) are quite different from those of German silver (mixture of copper, zinc and nickel) or bronze (mixture of copper and tin);

1 part per million (ppm) of fluoride ions in water prevents tooth decay, while 1.5 ppm causes the tooth to become mottled and high concentrations of fluoride ions can be poisonous (for example, sodium fluoride is used in rat poison); intravenous injections are always dissolved in water containing salts at particular ionic concentrations that match with blood plasma concentrations and so on.

In this Unit, we will consider mostly liquid solutions and their formation. This will be followed by studying the properties of the solutions, like vapour

pressure and colligative properties. We will begin with types of solutions and then various alternatives in which concentrations of a solute can be expressed in liquid solution.

Solutions are homogeneous mixtures of two or more than two components.
By homogenous mixture we mean that its composition and properties are uniform throughout the mixture. Generally, the component that is present in the largest quantity is known as solvent.

Solvent determines the physical state in which solution exists. One or more components present in the solution other than solvent are called solutes. In this Unit we shall consider only binary solutions (i.e., consisting of two components).

Here each component may be solid, liquid or in gaseous state and are summarised in Table 1.1.

Type of Solution	Solute	Solvent	Common Examples
Gaseous Solutions	Gas Liquid Solid	Gas Gas Gas	Mixture of oxygen and nitrogen gases. Chloroform mixed with nitrogen gas. Camphor in nitrogen gas.
Liquid Solutions	Gas Liquid Solid	Liquid Liquid Liquid	Oxygen dissolved in water. Ethanol dissolved in water. Glucose dissolved in water.
Solid Solutions	Gas Liquid Solid	Solid Solid Solid	Solution of hydrogen in palladium. Amalgam of mercury with sodium. Copper dissolved in gold.

Composition of a solution can be described by expressing its concentration.
The latter can be expressed either qualitatively or quantitatively.
For example, qualitatively we can say that the solution is dilute (i.e., relatively very small quantity of solute) or it is concentrated (i.e., relatively very large quantity of solute). But in real life these kinds of description can add to lot of confusion and thus the need for a quantitative description of the solution.

There are several ways by which we can describe the concentration of the solution quantitatively.

(i) Mass percentage (w/w): The mass percentage of a component of a solution is defined as:

Mass % of a component

Mass % of a component = Mass of the component in the solution Total mass of the solution × 100

For example, if a solution is described by 10% glucose in water by mass, it means that 10 g of glucose is dissolved in 90 g of water resulting in a 100 g solution.
Concentration described by mass percentage is commonly used in industrial chemical applications.

For example, commercial bleaching solution contains 3.62 mass percentage of sodium hypochlorite in water.

ii) Volume percentage (V/V): The volume percentage is defined as:

Volume % of a component = Volume of the component Total volume of solution × 100

For example, 10% ethanol solution in water means that 10 mL of ethanol is dissolved in water such that the total volume of the solution is 100 mL. Solutions containing liquids are commonly expressed in this unit. For example, a 35% (v/v) solution of ethylene glycol, an antifreeze, is used in cars for cooling the engine. At this concentration the antifreeze lowers the freezing point of water to 255.4K (-17.6°C).

Mole fraction unit is very useful in relating some physical properties of solutions, say vapour pressure with the concentration of the solution and quite useful in describing the calculations involving gas mixtures.

(vi) Molarity: Molarity (M) is defined as number of moles of solute dissolved in one litre (or one cubic decimetre) of solution,

For example, 0.25 mol L^-1 (or 0.25 M) solution of NaOH means that 0.25 mol of NaOH has been dissolved in one litre (or one cubic decimetre).

temperature. The solubility of gases increase with increase of pressure. for solution of gases in a solvent, consider a system as shown in Fig. 1.1 (a). The lower part is solution and the upper part is gaseous system at pressure p and temperature T. Assume this system to be in a state of dynamic equilibrium, i.e., under these conditions rate of gaseous particles entering and leaving the solution phase is the same.

Now increase the pressure over the solution phase by compressing the gas to a smaller volume [Fig. 1.1 (b)]. This will increase the number of gaseous particles per unit volume over the solution and also the rate at which the gaseous particles are striking the surface of solution to enter it. The solubility of the gas will increase until a new equilibrium is reached resulting in an increase in the pressure of a gas above the solution and thus its solubility increases.

…pressure of the gas. If we use the mole fraction of a gas in the solution as a measure of its solubility, then it can be said that the mole fraction of gas in the solution is proportional to the partial pressure of the gas over the solution.

The most commonly used form of Henry’s law states that “the partial pressure of the gas in vapour phase (p) is proportional to the mole fraction of the gas (x) in the solution” and is expressed as:

Here K_H is the Henry’s law constant. If we draw a graph between partial pressure of the gas versus mole fraction of the gas in solution, then we should get a plot of the type as shown in Fig. 1.2.

Different gases have different K_H values at the same temperature (Table 1.2). This suggests that K_H is a function of the nature of the gas.

It is obvious from equation (1.11) that higher the value of K_H at a given pressure, the lower is the solubility of the gas in the liquid. It can be seen from Table 1.2 that K_H values for both N₂ and O₂ increase with increase of temperature indicating that the solubility of gases…

Gas	Temperature/K	KH /kbar	Gas	Temperature/K	KH /kbar
He	293	144.97	Argon	298	40.3
H2	293	69.16	CO2	298	1.67
N2	293	76.48	Formaldehyde	298	1.83 × 10-5
N2	303	88.84	Methane	298	0.413
O2	293	34.86	Vinyl chloride	298	0.611
O2	303	46.82

increases with decrease of temperature.
It is due to this reason that aquatic species are more comfortable in cold waters rather than in warm waters.