Is Matter Around Us Pure

All the things around us are made up of matter. Matter can be classified as either a pure substance or a mixture.

A pure substance is made up of only one kind of particle and has a fixed composition throughout, so every sample of it has the same set of properties (such as melting point and boiling point). A pure substance can be an element (e.g., iron, oxygen) or a compound (e.g., water, H₂O, or carbon dioxide, CO₂).

A mixture is formed when two or more pure substances (elements and/or compounds) are combined in any ratio, without forming a new substance. For example, lemonade (sugar, salt, lemon juice and water) and air (nitrogen, oxygen, carbon dioxide and other gases) are mixtures.

Mixtures are of two types:

• Homogeneous mixture: has a uniform (same) composition throughout, with no visible boundaries between its components. Example: a solution of salt in water, or air.
• Heterogeneous mixture: has a non-uniform composition, with visible boundaries between its components. Example: a mixture of sand and water, or oil and water.

Unlike a pure substance, the components of a mixture keep their individual properties and can be present in variable proportions.

A solution is a homogeneous mixture of two or more substances. It is made up of a solvent (the component present in the larger amount, in which the other substance dissolves) and one or more solutes (the component(s) present in the smaller amount, which get dissolved).

For example, in a solution of sugar in water, sugar is the solute and water is the solvent. Solutions can involve any state of matter: sea water (solids dissolved in liquid water), soda water (a gas dissolved in a liquid), and alloys such as brass (solid copper and zinc) are all solutions.

A solution is stable (the solute does not settle down on standing), and its particles are so small (less than 1 nm) that they cannot be seen even with a powerful microscope and do not scatter a beam of light.

Depending on how much solute is dissolved at a given temperature, a solution may be:

• An unsaturated solution, in which more solute can still be dissolved at that temperature.
• A saturated solution, in which no more solute can be dissolved at that temperature, so undissolved solute remains at the bottom.

The amount of solute that must be dissolved in a given amount of solvent to make a saturated solution at a given temperature is called the solubility of that solute; solubility generally changes with temperature (usually more solid solute dissolves in hotter water).

The concentration of a solution tells us how much solute is present in a given amount of solution. It can be expressed as:

1. Mass by mass percentage = (Mass of solute ÷ Mass of solution) × 100
2. Mass by volume percentage = (Mass of solute ÷ Volume of solution) × 100
3. Volume by volume percentage = (Volume of solute ÷ Volume of solution) × 100, used when both the solute and the solvent are liquids.

For example, if 20 g of salt is dissolved in 80 g of water, the mass of the solution is 100 g, so the mass by mass percentage of the solution is (20 ÷ 100) × 100 = 20%. Similarly, if 20 mL of ethanol is made up to 200 mL of solution with water, the volume by volume percentage is (20 ÷ 200) × 100 = 10%.

Not every mixture is a clear solution. Based on the size of the particles of their components, mixtures can also be classified as suspensions and colloids.

A suspension is a heterogeneous mixture in which solute-like particles are large enough (bigger than 1000 nm) to be seen with the naked eye. These particles do not dissolve, remain suspended, and can be seen to settle down when the suspension is left undisturbed; they can also be separated from the liquid by filtration. Example: chalk powder mixed in water, or muddy river water.

A colloid (or colloidal solution) is a mixture in which the particle size is between that of a true solution and a suspension, roughly between 1 nm and 1000 nm. Colloidal particles are too small to be seen with the naked eye, do not settle down on standing, and cannot be separated by ordinary filter paper, but they are large enough to scatter a beam of light passing through them. Examples: milk, ink, shaving foam and smoke.

A colloid has two parts: the dispersed phase (the solute-like component present in a smaller amount) and the dispersion medium (the component in which the dispersed phase is spread out).

Two important properties help identify colloids:

• The Tyndall effect: colloidal particles scatter a beam of light passed through them, making the path of the light visible, as when a beam of sunlight enters a dusty room through a small hole.
• Brownian motion: colloidal particles show a continuous, random zig-zag movement because they are constantly struck by the fast-moving molecules of the dispersion medium; this movement helps keep them from settling down.

Matter can undergo two broad kinds of change: physical change and chemical change.

In a physical change, only the physical properties of a substance (such as its state, shape, size or colour) change, but no new substance is formed. The composition of the substance stays the same, and most physical changes can be reversed. For example, melting of wax, dissolving sugar in water, cutting a sheet of paper, and the melting of ice into water are all physical changes.

In a chemical change, one or more new substances with entirely different properties are formed, and the original substance(s) usually cannot be recovered easily; most chemical changes are irreversible. For example, burning of paper, rusting of an iron almirah (which forms rust, a form of iron oxide), and the curdling of milk are all chemical changes.

Some processes, such as the burning of a candle, involve both kinds of change at once: the melting of the solid wax near the flame is a physical change, while the burning of the wax vapour to form carbon dioxide and water is a chemical change.

The components of a mixture can be separated by physical methods, chosen according to the nature of the mixture and the properties of its components.

• Evaporation: used to recover a dissolved solid from a solution by heating it so that the liquid solvent turns into vapour and escapes, leaving the solid behind. Example: recovering common salt from salt water.
• Distillation: used to separate a liquid solvent from a solution (or two miscible liquids with sufficiently different boiling points) by heating the mixture; the liquid with the lower boiling point vaporises first and is then cooled and condensed back into a liquid in a separate container. Example: obtaining pure water from salty water.
• Separating funnel: used to separate two immiscible liquids (liquids that do not mix and form separate layers), based on the difference in their densities. The denser liquid settles at the bottom and is drained off first through the stopcock, leaving the lighter liquid behind. Example: separating kerosene oil from water, or oil from water.
• Chromatography: used to separate components of a mixture that are present in very small quantities but dissolve in the same solvent, based on their different solubility or movement through a medium such as filter paper. Example: separating the different coloured dyes present in black ink or a flower extract.
• Centrifugation: used to quickly separate insoluble solid particles suspended in a liquid by spinning the mixture at high speed; the denser particles settle towards the bottom while the lighter liquid stays near the top. Example: separating cream from milk.
• Coagulation: used to separate very fine suspended or colloidal particles from a liquid by adding a coagulating substance (such as alum/fitkari), which makes the fine particles clump together into larger, heavier masses that settle down and can then be removed. Example: purifying muddy water by adding alum.
• Sublimation: used to separate a substance that changes directly from the solid to the gaseous state on heating (and back to solid on cooling) from a mixture containing a substance that does not sublime. Example: separating ammonium chloride or camphor from common salt.
• Crystallisation: used to obtain pure, well-formed crystals of a solid solute from its solution; the solution is heated to become saturated and then allowed to cool slowly so that pure crystals separate out. Crystallisation gives a purer product than simple evaporation, because prolonged strong heating during evaporation can decompose the solute or trap other soluble impurities in it. Example: obtaining pure crystals of copper sulphate from an impure sample.