Chemical Reactions & Equations

A chemical reaction is a change in which one or more new substances, with new properties, are formed. The substances you start with change their chemical identity, so you cannot get them back just by cooling or filtering. For example, when a magnesium ribbon burns, the shiny silvery metal turns into a white powder that is nothing like the metal you began with.

How do you know a chemical reaction has actually taken place? Watch for one or more of these tell-tale signs:

• a change of state (solid, liquid or gas)
• a change of colour
• the evolution of a gas (bubbles or fizzing)
• a change in temperature (the mixture becomes hotter or colder)
• the formation of a precipitate (an insoluble solid that settles out of a solution)

The substances that take part in the reaction are the reactants. The new substances that are formed are the products. We always write the reactants on the left and the products on the right.

We first describe the reaction in words as a word equation, and then rewrite it as a chemical (skeletal) equation using symbols and formulae. The arrow (→) points from reactants to products and is read as "produces" or "yields". Take magnesium burning in air to form magnesium oxide:

• Word equation: Magnesium + Oxygen → Magnesium oxide
• Skeletal equation: Mg(s) + O₂(g) → MgO(s)

This skeletal equation uses the correct formulae, but it is not yet balanced — the number of atoms on the two sides is not equal. Making the atoms equal on both sides is the very next step you will learn.

In every chemical reaction, mass is neither created nor destroyed. This is the law of conservation of mass. So the total mass of the reactants must always equal the total mass of the products. Atoms are only rearranged, not made or lost — which means the number of atoms of each element must be the same on both sides of the equation.

An equation that shows the correct formulae of reactants and products, but does not yet have equal atoms on both sides, is called a skeletal (unbalanced) equation. Because it breaks the law of conservation of mass, we must balance it. A balanced equation has the same number of atoms of every element on the left-hand side (LHS) and the right-hand side (RHS).

Golden rule: balance an equation by changing only the coefficients — the big numbers written in front of a formula. Never change a subscript inside a formula, because that changes the substance itself. For example, you may write 2H₂O (two water molecules), but you must never turn it into H₂O₂ (that is hydrogen peroxide — a completely different compound). This method of adjusting coefficients by trial is called the hit-and-trial method.

Worked example — balancing Fe + H₂O → Fe₃O₄ + H₂ (iron reacting with steam):

1. Write the skeletal equation and count the atoms. LHS: Fe = 1, H = 2, O = 1. RHS: Fe = 3, H = 2, O = 4. Clearly not balanced.
2. Balance oxygen first, since it appears in the largest number (in Fe₃O₄). Put a 4 before H₂O: Fe + 4H₂O → Fe₃O₄ + H₂. Now O = 4 on both sides.
3. Balance hydrogen. The LHS now has 8 H atoms, so put a 4 before H₂ on the RHS: Fe + 4H₂O → Fe₃O₄ + 4H₂. Now H = 8 on both sides.
4. Balance iron. The RHS has 3 Fe, so put a 3 before Fe: 3Fe + 4H₂O → Fe₃O₄ + 4H₂.
5. Final check — Fe: 3 = 3, H: 8 = 8, O: 4 = 4. The equation is balanced.

The last step is to make the equation more informative by adding the physical states in brackets and any conditions of the reaction:

• (s) for solid, (l) for liquid, (g) for gas, and (aq) for a substance dissolved in water (aqueous).
• Conditions such as heat (Δ), a catalyst, pressure, temperature or light are written above or below the arrow.
• With states added, our example becomes 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g). Water is used here as steam, so it is written (g).

A combination reaction (also called a synthesis reaction) is one in which two or more reactants combine to form a single product. The quick test: if you see many substances on the left of the arrow but only one substance on the right, it is a combination reaction.

A classic example is quicklime reacting with water. Calcium oxide (quicklime) reacts vigorously with water to form calcium hydroxide (slaked lime), and a large amount of heat is given out:

CaO(s) + H₂O(l) → Ca(OH)₂(aq)

The slaked lime, Ca(OH)₂, is the solution used to whitewash walls. Here are a few more combination reactions you must know:

• Burning of carbon (coal): C(s) + O₂(g) → CO₂(g)
• Burning of hydrogen: 2H₂(g) + O₂(g) → 2H₂O(l)
• Burning of magnesium ribbon: 2Mg(s) + O₂(g) → 2MgO(s)

Notice that all of these reactions give out heat. A reaction in which energy — usually heat, and sometimes light — is released along with the products is called an exothermic reaction. Most combination reactions are exothermic.

Exothermic reactions are all around us. Respiration is an exothermic process: the glucose we get from food is slowly broken down using oxygen to release the energy our body needs. The burning of fuels (like LPG, coal or wood) also releases heat and light. Remember, though — respiration and burning give out energy, but they are not themselves combination reactions.

A decomposition reaction is one in which a single reactant breaks down into two or more products. It is the exact opposite of a combination reaction, where two or more reactants join to form a single product.

To pull a compound apart, you must supply energy. So most decomposition reactions absorb energy from the surroundings — they are endothermic. We name them by the source of that energy: heat, electricity or light.

1. Thermal decomposition — breakdown by heat. When limestone (calcium carbonate) is heated strongly, it breaks into quicklime and carbon dioxide (the heat symbol Δ is written over the arrow):

• CaCO₃(s) → CaO(s) + CO₂(g). The quicklime (CaO) formed is very useful — for example in making cement and slaked lime.
• 2FeSO₄(s) → Fe₂O₃(s) + SO₂(g) + SO₃(g). Green ferrous sulphate crystals lose water and colour on heating, leaving reddish-brown ferric oxide, with the smell of burning sulphur (from SO₂ and SO₃).

2. Electrolytic decomposition — breakdown by passing electricity. When electric current is passed through acidified water, it splits into hydrogen and oxygen:

• 2H₂O(l) → 2H₂(g) + O₂(g). Notice that the volume of hydrogen collected is twice the volume of oxygen.

3. Photolytic (photo) decomposition — breakdown by light. Silver halides break down in sunlight, leaving behind grey silver metal:

• 2AgCl(s) → 2Ag(s) + Cl₂(g). White silver chloride slowly turns grey in sunlight.
• 2AgBr(s) → 2Ag(s) + Br₂(l). This sensitivity of silver halides to light is used in black-and-white photography.

A displacement reaction is one in which a more reactive element takes the place of a less reactive element in a compound (usually a salt solution). The more reactive metal goes into the solution as a new salt, and the less reactive metal comes out as the free element.

The classic example is putting an iron nail into blue copper sulphate solution. Iron is more reactive than copper, so iron displaces copper:

• Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

You can actually see this happening. The blue colour of the copper sulphate solution slowly fades (it turns pale green as FeSO₄ forms), and a reddish-brown deposit of copper collects on the iron nail. Two more examples work the same way, because both zinc and lead are more reactive than copper:

• Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
• Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)

Whether displacement happens depends on the reactivity series of metals. A metal higher in the reactivity series can displace a metal lower than it from that metal's salt solution. So iron, zinc and lead (all above copper) can displace copper. The reverse never happens on its own: copper cannot displace iron, zinc or lead, because copper is the less reactive metal.

In a double displacement reaction, two compounds react by exchanging their ions. Each compound gives up its ions, they swap partners, and you get two brand-new compounds.

A precipitation reaction is a special double displacement reaction in which one of the new compounds is an insoluble solid. This solid does not dissolve, so it settles out of the solution. The standard NCERT example is:

Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)

Here sodium sulphate and barium chloride exchange ions. Barium sulphate (BaSO₄) is a white precipitate — that is why it carries the state symbol (s). Sodium chloride stays dissolved, so it keeps the symbol (aq).

A precipitate is simply an insoluble solid that separates out of a solution during a reaction. In equations we show it with the symbol (s).

Neutralisation is also a double displacement reaction. An acid reacts with a base to give a salt and water:

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

The ions of the acid and the base swap partners to form common salt and water, so acid + base → salt + water is just another ion exchange.

• Double displacement: two compounds exchange ions → two new compounds.
• Precipitation: a double displacement in which an insoluble solid (precipitate) is formed.
• A white solid settling down (like BaSO₄) is a clear sign a precipitate has formed.

In many reactions, oxygen (or hydrogen) is passed from one substance to another. Two changes go on side by side, and each one has a name.

Oxidation is the gain of oxygen or the loss of hydrogen by a substance. Reduction is the loss of oxygen or the gain of hydrogen. (In terms of electrons, oxidation is loss of electrons and reduction is gain of electrons — remember OIL RIG: Oxidation Is Loss, Reduction Is Gain.)

A redox reaction is one in which oxidation and reduction happen together, at the same time. One substance can be oxidised only because another substance is reduced.

Look at CuO(s) + H₂(g) → Cu(s) + H₂O(l):

• Copper(II) oxide loses oxygen, so CuO is reduced to copper (Cu).
• Hydrogen gains oxygen, so H₂ is oxidised to water (H₂O).

The substance that supplies oxygen (or removes hydrogen) is the oxidising agent, and it gets reduced itself. The substance that removes oxygen (or supplies hydrogen) is the reducing agent, and it gets oxidised itself. So in the reaction above, CuO is the oxidising agent and H₂ is the reducing agent.

Another example is ZnO(s) + C(s) → Zn(s) + CO(g): zinc oxide loses oxygen, so it is reduced to zinc (Zn), while carbon gains oxygen and is oxidised to carbon monoxide (CO). Here ZnO is the oxidising agent and carbon is the reducing agent.

Corrosion is the slow eating away of a metal by the action of air, moisture and chemicals on its surface. It is a harmful effect of oxidation because it slowly destroys metal articles, tools, bridges, machines and vehicles.

The most common example is the rusting of iron. Rusting needs both oxygen (from air) and water (moisture) together — if either one is missing, iron will not rust. The reddish-brown rust that forms is hydrated iron(III) oxide, Fe₂O₃·xH₂O, which slowly flakes off and eats deeper into the metal.

Other metals corrode too. Copper slowly loses its shiny surface and gains a green coating of basic copper carbonate in moist air, and silver turns black because it forms a coating of silver sulphide (Ag₂S) with sulphur compounds present in the air.

We can prevent rusting by keeping air and moisture away from the metal surface, or by changing the metal itself:

• Painting the surface.
• Oiling or greasing the parts.
• Galvanising — coating iron with a layer of zinc.
• Electroplating — depositing a thin layer of another metal on the surface.
• Alloying — mixing metals to make a rust-resistant alloy, for example stainless steel.

Rancidity is the oxidation of the fats and oils present in food. When food is left open for a long time, its fats and oils get oxidised, and the food develops an unpleasant smell and taste. We slow down rancidity by keeping oxygen away from the food:

• Adding antioxidants — substances that get oxidised in place of the food.
• Flushing the pack with nitrogen gas so that no oxygen is left inside (as in packets of chips).
• Storing food in airtight containers.
• Refrigerating the food to slow down the oxidation.