Chemical Reactions and Equations — Notes

A chemical reaction turns one or more substances (reactants) into new substances (products) with different properties. A physical change alters only the form or state; no new substance forms and it is usually easy to reverse (melting ghee, dissolving sugar, freezing water).

Signs that a reaction has happened — look for one or more of:

• a change in colour (a cut apple browning; blue copper sulphate fading when iron is added)
• evolution of a gas (bubbles in fermenting idli batter; fizzing of an antacid tablet in water)
• a change in temperature (quicklime and water getting hot)
• formation of a precipitate (soap scum in hard water; yellow lead iodide)
• a change in state (wood burning away into gases and ash)

One sign alone is not proof — ice melting changes state but is physical. Always ask: is a new substance formed?

Burning magnesium: a ribbon rubbed clean with sandpaper (to remove the oxide coating) burns with a dazzling white flame and leaves white magnesium oxide powder. Light, heat and a new white solid all point to a chemical change.

Word equation → skeletal equation (formulae, not yet balanced) → balanced equation → add state symbols and conditions.

Why balance? The law of conservation of mass: mass is neither created nor destroyed in a chemical reaction, so the number of atoms of each element must be the same on both sides.

Hit-and-trial method

1. Write correct formulae; list the atoms of each element on both sides.
2. Start with the compound that has the most atoms; balance one element at a time using coefficients (numbers in front).
3. Leave free elements such as $\ce{O2}$ or $\ce{H2}$ for last.
4. Recount every element. Use the smallest whole numbers.
5. Never change subscripts — that would change the substance.

State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution. Conditions (heat, light, electricity, catalyst, temperature, pressure) go above or below the arrow.

Combination: two or more reactants form a single product, $\ce{A + B -> AB}$.

• Quicklime and water (whitewash): $\ce{CaO(s) + H2O(l) -> Ca(OH)2(aq)}$ + heat. Over the next few days the slaked lime on the wall reacts with carbon dioxide in air to form a thin, shiny layer of calcium carbonate.
• Burning of coal and of hydrogen; ammonia + hydrogen chloride → white fumes of ammonium chloride.

Decomposition: a single reactant breaks into two or more simpler products, $\ce{AB -> A + B}$. It needs energy:

• Thermal (heat): limestone in a lime kiln; ferrous sulphate crystals (pale green → reddish-brown $\ce{Fe2O3}$, smell of burning sulphur); lead nitrate (brown fumes of nitrogen dioxide).
• Electrolytic (electricity): water with a few drops of dilute sulphuric acid → hydrogen at the cathode (−) and oxygen at the anode (+), 2 : 1 by volume. Hydrogen burns with a pop; oxygen relights a glowing splinter.
• Photolytic (light): white silver chloride turns grey in sunlight; silver bromide behaves the same (black-and-white photography). Silver salts are stored in dark bottles.

Displacement: a more reactive element replaces a less reactive one from its compound, $\ce{A + BC -> AC + B}$.

Reactivity series (most → least reactive): K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au.

• Iron nail in copper sulphate: blue → pale green; reddish-brown copper coats the nail.
• Zinc in copper sulphate: blue → colourless; copper deposits.
• Copper in silver nitrate: solution turns blue; shiny silver deposits.
• Copper in iron sulphate: no reaction (copper is less reactive than iron).

Double displacement: two compounds exchange ions, $\ce{AB + CD -> AD + CB}$. If a product is insoluble it forms a precipitate (a precipitation reaction).

• Sodium sulphate + barium chloride → white barium sulphate.
• Lead nitrate + potassium iodide → yellow lead iodide.
• Silver nitrate + sodium chloride → white, curdy silver chloride.
• Copper sulphate + sodium hydroxide → blue copper hydroxide.
• Acid + base (neutralisation) is double displacement without a precipitate.

Exothermic reactions give out heat (surroundings warm up): burning of fuels (LPG, coal), respiration, quicklime + water, decomposition of vegetable matter into compost, neutralisation.

Endothermic reactions take in heat (energy must be supplied, or the surroundings cool): thermal decomposition of calcium carbonate, photosynthesis, electrolysis of water. Most decomposition reactions are endothermic because energy is needed to break the reactant apart.

Respiration is exothermic: glucose from digested food is oxidised in cells, releasing the energy that keeps us warm and active. Photosynthesis is endothermic: plants absorb light energy and store it in glucose.

• Oxidation: gain of oxygen or loss of hydrogen.
• Reduction: loss of oxygen or gain of hydrogen.
• They always occur together: a redox reaction.
• Oxidising agent: gives oxygen or removes hydrogen; it is itself reduced.
• Reducing agent: removes oxygen or gives hydrogen; it is itself oxidised.

Heating brown copper in air gives black copper oxide (copper oxidised). Passing hydrogen over hot copper oxide turns it brown again (copper oxide reduced, hydrogen oxidised).

Method for any redox question: (1) find the substance that loses oxygen/gains hydrogen → reduced, and it is the oxidising agent; (2) find the substance that gains oxygen/loses hydrogen → oxidised, and it is the reducing agent.

Corrosion: a metal is slowly attacked by substances around it such as air, moisture and acids.

• Iron → reddish-brown rust (hydrated iron(III) oxide, $\ce{Fe2O3.xH2O}$). Rusting needs both oxygen and water; salt water speeds it up.
• Copper → green coating (basic copper carbonate); silver → black coating (silver sulphide).
• Prevention: painting, oiling/greasing, galvanising (zinc coating), chrome plating, alloying (stainless steel).

Rancidity: oxidation of fats and oils in food, giving a bad smell and taste.

• Prevention: airtight containers, flushing packets with nitrogen, adding antioxidants, refrigeration, keeping food away from light.