Class 10 Science Notes Chapter 1 — Chemical Reactions and Equations#
Signs of a chemical reaction#
A chemical reaction forms new substances. You can tell one has happened from:
- a change in state
- a change in colour
- evolution of a gas
- a change in temperature
- formation of a precipitate
Example: magnesium ribbon burns with a dazzling white flame and forms white magnesium oxide powder.
$$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$$Chemical equations#
- Reactants are written on the left, products on the right, with an arrow between them.
- Balanced equation — the number of atoms of each element is the same on both sides. This follows the law of conservation of mass: mass is neither created nor destroyed in a reaction.
- Balance by changing the coefficients in front of formulae, never the subscripts inside them.
- States: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution. Conditions such as heat, pressure or a catalyst are written above or below the arrow.
Types of chemical reactions#
| Type | Pattern | Example |
|---|---|---|
| Combination | $A + B \rightarrow AB$ | $\text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2$ |
| Decomposition | $AB \rightarrow A + B$ | $\text{CaCO}_3 \xrightarrow{\text{heat}} \text{CaO} + \text{CO}_2$ |
| Displacement | $A + BC \rightarrow AC + B$ | $\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$ |
| Double displacement | $AB + CD \rightarrow AD + CB$ | $\text{Na}_2\text{SO}_4 + \text{BaCl}_2 \rightarrow \text{BaSO}_4 + 2\text{NaCl}$ |
Combination reactions#
- Quick lime + water: $\text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_2$ + heat. Calcium oxide (quick lime) forms calcium hydroxide (slaked lime). A solution of slaked lime is used for whitewashing: it reacts slowly with CO₂ in air to form a shiny coat of calcium carbonate.
- Burning of coal: $\text{C} + \text{O}_2 \rightarrow \text{CO}_2$
- Formation of water: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$
Decomposition reactions#
They need energy, given as heat, light or electricity.
- Thermal: green ferrous sulphate crystals lose water and decompose, with the smell of burning sulphur: $$2\text{FeSO}_4 \xrightarrow{\text{heat}} \text{Fe}_2\text{O}_3 + \text{SO}_2 + \text{SO}_3$$
- Thermal: lead nitrate gives off brown fumes of nitrogen dioxide: $$2\text{Pb(NO}_3)_2 \xrightarrow{\text{heat}} 2\text{PbO} + 4\text{NO}_2 + \text{O}_2$$
- Thermal: limestone to quick lime (used in making cement): $$\text{CaCO}_3 \xrightarrow{\text{heat}} \text{CaO} + \text{CO}_2$$
- Electrolysis of water: hydrogen and oxygen are formed, with the volume of hydrogen double that of oxygen. $$2\text{H}_2\text{O} \xrightarrow{\text{electricity}} 2\text{H}_2 + \text{O}_2$$
- Light: white silver chloride turns grey in sunlight. These reactions are used in black-and-white photography. $$2\text{AgCl} \xrightarrow{\text{sunlight}} 2\text{Ag} + \text{Cl}_2$$ $$2\text{AgBr} \xrightarrow{\text{sunlight}} 2\text{Ag} + \text{Br}_2$$
Displacement reactions#
A more reactive metal displaces a less reactive metal from its salt solution.
- Iron nail in copper sulphate: the nail turns brownish and the blue solution fades. $$\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$$
- $\text{Zn} + \text{CuSO}_4 \rightarrow \text{ZnSO}_4 + \text{Cu}$
- $\text{Pb} + \text{CuCl}_2 \rightarrow \text{PbCl}_2 + \text{Cu}$
Double displacement and precipitation#
Two compounds exchange ions. A reaction that forms an insoluble solid (precipitate) is a precipitation reaction.
- White precipitate of barium sulphate: $$\text{Na}_2\text{SO}_4 + \text{BaCl}_2 \rightarrow \text{BaSO}_4\downarrow + 2\text{NaCl}$$
- Yellow precipitate of lead iodide: $$\text{Pb(NO}_3)_2 + 2\text{KI} \rightarrow \text{PbI}_2\downarrow + 2\text{KNO}_3$$
Exothermic and endothermic reactions#
- Exothermic — heat is given out. Examples: quick lime with water, burning of natural gas, respiration: $$\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy}$$ Decomposition of vegetable matter into compost is also exothermic.
- Endothermic — heat is taken in. Most decomposition reactions are endothermic.
Oxidation and reduction#
| Oxidation | Reduction | |
|---|---|---|
| Oxygen | Gain | Loss |
| Hydrogen | Loss | Gain |
- Redox reaction — one substance is oxidised and another is reduced in the same reaction.
- Heating copper powder in air gives a black coating of copper oxide. Passing hydrogen over it turns it brown again: $$\text{CuO} + \text{H}_2 \xrightarrow{\text{heat}} \text{Cu} + \text{H}_2\text{O}$$ Here CuO is reduced and H₂ is oxidised.
- $\text{ZnO} + \text{C} \rightarrow \text{Zn} + \text{CO}$: ZnO is reduced, C is oxidised.
- $\text{MnO}_2 + 4\text{HCl} \rightarrow \text{MnCl}_2 + 2\text{H}_2\text{O} + \text{Cl}_2$: MnO₂ is reduced, HCl is oxidised.
- The substance that gets oxidised is the reducing agent. The substance that gets reduced is the oxidising agent.
Oxidation in everyday life#
- Corrosion — a metal is attacked by moisture, air or acids around it. Examples: rusting of iron (reddish-brown), a black coating on silver, a green coating on copper. It damages cars, bridges, railings and ships.
- Rancidity — fats and oils in food get oxidised, so the smell and taste change. Prevented by adding antioxidants, storing in airtight containers, and filling chip packets with nitrogen.
Common mistakes#
- Balancing by changing a subscript (H₂O to H₂O₂). Change only the coefficients.
- Calling every reaction that gives off a gas a decomposition reaction.
- Mixing up the agents: the substance that is oxidised is the reducing agent.
- Forgetting that respiration is exothermic.