Solutions

After studying this Unit, you will be able to

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

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:

$\text{Mass \% of a component}=\frac{\text{Mass of the component in the solution}}{\text{Total mass of the solution}}\times 100(1.1)$

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:

$\text{Volume \% of a component}=\frac{\text{Volume of the component}}{\text{Total volume of solution}}\times 100$

$(1.2)$

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 255.4 K (–17.6°C). (iii) Mass by volume percentage (w/V): Another unit which is commonly used in medicine and pharmacy is mass by volume percentage. It is the mass of solute dissolved in 100 mL of the solution. (iv) Parts per million: When a solute is present in trace quantities, it is convenient to express concentration in parts per million (ppm) and is defined as:

$\text{Parts per million}=\frac{\text{Number of parts of the component}}{\text{Total number of parts of all components of the solution}}\times 10^6(1.3)$

$\text{Mole fraction of a component}=\frac{\text{Number of moles of the component}}{\text{Total number of moles of all the components}}(1.4)$

For example, in a binary mixture, if the number of moles of A and B are $n_A$ and $n_B$ respectively, the mole fraction of A will be

$x_A=\frac{n_A}{n_A+n_B}(1.5)$

For a solution containing $i$ number of components, we have:

$x_i=\frac{n_i}{n_1+n_2+\dots +n_i}=\frac{n_i}{\sum n_i}(1.6)$

It can be shown that in a given solution sum of all the mole fractions is unity, i.e. $x_1+x_2+\dots +x_i=1(1.7)$

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.

Calculate the mole fraction of ethylene glycol (\ce{C2H6O2}) in a solution containing 20% of \ce{C2H6O2} by mass.

Assume that we have 100 g of solution (one can start with any amount of solution because the results obtained will be the same). Solution will contain 20 g of ethylene glycol and 80 g of water. Molar mass of $\text{ce}C2H6O2$ = 12 \times 2 + 1 \times 6 + 16 \times 2 = 62 g mol${}^{-1}$.

Mole fraction of water can also be calculated as: 1 – 0.068 = 0.932

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

$\text{Molarity}=\frac{\text{Moles of solute}}{\text{Volume of solution in litre}}(1.8)$

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