CBSE Class 10 Science Advanced Chemistry Chemical Reactions and Equations Notes Set 01

Chemical Reactions and Equations

A. Chemical Changes and the Characteristics of Chemical Reactions

Chemistry is the branch of science that studies what matter is made of, how it behaves and how it changes when it interacts with other matter.

(a) Physical change and chemical change

Examples. Physical changes: melting of ice, evaporation of water, dissolving salt in water, stretching a rubber band. Chemical changes: setting of milk into curd, digestion of food in the body, burning of petrol, rusting of iron.

A new substance forms in a chemical change, and energy such as heat, electricity or radiation is either taken in or given out. The total mass of the substances stays the same before and after the change.

(b) Chemical reaction

If sugar crystals are heated, they first melt, then give off steam-like vapour and leave a brownish-black mass. The crystals do not return on cooling. So heating sugar is a chemical change, and the process that brings about a chemical change is a chemical reaction.

(c) Characteristics of a chemical reaction

1. The starting substances are the reactants and the substances formed are the products.

2. Old chemical bonds break and new bonds form.

3. The products are new substances with different names, formulae and properties from the reactants.

4. A chemical reaction is hard to reverse.

5. Besides heat, light and electricity can also be used to start reactions.

(d) Observations that show a chemical reaction has taken place

1. Evolution of a gas. Some reactions give off a gas.

(a) A metal such as zinc, magnesium or iron with dilute sulphuric acid gives hydrogen gas: \( \ce{Zn(s) + H2SO4 "(dil)" -> ZnSO4(aq) + H2(g)} \)

(b) Iron sulphide with dilute sulphuric acid gives hydrogen sulphide gas: \( \ce{FeS(s) + H2SO4 "(dil)" -> FeSO4(aq) + H2S(g)} \)

(c) Heating a mixture of potassium chlorate and manganese dioxide gives oxygen gas: \( \ce{2KClO3(s) ->[\Delta, MnO2] 2KCl(s) + 3O2(g)} \). Here \( \ce{MnO2} \) is only a catalyst. This is the laboratory method of preparing oxygen.

2. Change of colour. Many reactions change the colour of the substances.

(a) Red lead oxide on strong heating becomes yellow lead monoxide and gives off oxygen: \( \ce{2Pb3O4(s) ->[heat] 6PbO(s) + O2(g)} \)

(b) Green copper carbonate on strong heating leaves a black residue: \( \ce{CuCO3(s) ->[heat] CuO(s) + CO2(g)} \)

(c) White lead(II) nitrate gives yellow lead(II) oxide and brown nitrogen dioxide gas: \( \ce{2Pb(NO3)2(s) ->[heat] 2PbO(s) + 4NO2(g) + O2(g)} \)

(d) White sugar turns into black carbon: \( \ce{C12H22O11(s) -> 12C(s) + 11H2O(g)} \)

3. Formation of a precipitate. When two clear solutions are mixed, an insoluble solid, called a precipitate, may appear.

(a) Silver nitrate with sodium chloride gives white silver chloride: \( \ce{AgNO3(aq) + NaCl(aq) -> NaNO3(aq) + AgCl(s)} \). Both reactants and sodium nitrate are colourless in solution.

(b) Ferrous sulphate with sodium hydroxide gives a dirty green precipitate of ferrous hydroxide: \( \ce{FeSO4(aq) + 2NaOH(aq) -> Na2SO4(aq) + Fe(OH)2(s)} \)

(c) Barium chloride with dilute sulphuric acid gives white barium sulphate: \( \ce{BaCl2(aq) + H2SO4(aq) -> BaSO4(s) + 2HCl(aq)} \)

4. Energy changes. Every chemical reaction either takes in energy or gives it out, so reactions are of two kinds.

(a) Endothermic reaction: a reaction that absorbs energy. Examples: \( \ce{C(s) + 2S(s) ->[heat] CS2(l)} \), and photosynthesis, in which green plants use light energy to make food from carbon dioxide and water.

(b) Exothermic reaction: a reaction that releases energy. Examples: a magnesium wire heated at its tip burns with a dazzling white flame, giving heat and light, \( \ce{2Mg(s) + O2(g) ->[heat] 2MgO(s) + "energy"} \); and quick lime dropped in water makes the water very hot, sometimes making it boil, \( \ce{CaO(s) + H2O(l) -> Ca(OH)2(aq) + "heat"} \).

5. Change of state. Some reactions change the state of the substances. Two volumes of hydrogen with one volume of oxygen give liquid water, \( \ce{2H2(g) + O2(g) -> 2H2O(l)} \), and passing an electric current through water splits it again, \( \ce{2H2O(l) ->[electric current] 2H2(g) + O2(g)} \). Ammonia and hydrogen chloride gases give solid ammonium chloride: \( \ce{NH3(g) + HCl(g) -> NH4Cl(s)} \).

Solved Examples

Example 1. A shiny brown-coloured element 'X' on heating in air becomes black in colour. Name the element 'X' and the black coloured compound formed.
Answer:
X is copper (Cu) and the black compound is copper oxide (CuO): \( \ce{2Cu + O2 ->[heat] 2CuO} \)

Example 2. What do you mean by a precipitation reaction ? Explain by giving examples.
Answer:
A reaction in which an insoluble solid, the precipitate, is formed is a precipitation reaction.
\( \ce{Na2CO3(aq) + CaCl2(aq) -> CaCO3(s) + 2NaCl(aq)} \) (calcium carbonate is the precipitate)
\( \ce{Na2SO4(aq) + BaCl2(aq) -> BaSO4(s) + 2NaCl(aq)} \) (barium sulphate is the precipitate)

B. Chemical Equations and Balancing

A chemical reaction can be described in words, but a sentence would be long. A chemical equation is a short way of writing a chemical reaction with the symbols and formulae of the substances involved. For example, zinc + hydrochloric acid gives zinc chloride + hydrogen, where zinc and hydrochloric acid are the reactants and the other two are the products.

(a) Types of chemical equations

1. Word equation. It writes the names of the reactants and products. The magnesium ribbon burning in oxygen: Magnesium + Oxygen → Magnesium oxide. Another example is Zinc + Sulphuric acid → Zinc sulphate + Hydrogen. In a word equation: reactants go on the left with a + sign between them, products go on the right with a + sign between them, an arrow separates the two sides and its head points toward the products. A word equation is useful but does not give the true picture of the reaction.

2. Symbol (formula) equation. It uses symbols and formulae in place of names, always written from the word equation. For burning of magnesium: \( \ce{Mg + O2 -> MgO} \)

(b) Unbalanced and balanced equations

If the number of atoms of any element is different on the two sides, the equation is unbalanced (also called a skeletal equation). In \( \ce{Mg + O2 -> MgO} \) both sides have one Mg atom, but oxygen is two on the left and one on the right. In a balanced equation every element has the same number of atoms on both sides: \( \ce{2Mg + O2 -> 2MgO} \)

Why balance ? Atoms are neither created nor destroyed in a reaction, which is the law of conservation of mass, so the equation must show the same number of atoms of each element on both sides.

(c) Balancing by hit and trial

We adjust the numbers (coefficients) in front of the formulae until every element is equal on both sides. Steps:

1. Write the word equation. 2. Convert it into a symbol equation with correct formulae. 3. Balance it by adjusting coefficients. 4. Never change the formula (the subscripts) of a substance. 5. Add the extra information (states, conditions) if possible.

Worked examples

Example 1. Zinc reacts with dilute sulphuric acid to give zinc sulphate and hydrogen. Symbol equation: \( \ce{Zn + H2SO4 -> ZnSO4 + H2} \)

All counts are equal, so the equation is already balanced.

Example 2. Iron reacts with steam to form iron(II,III) oxide and hydrogen. Word equation: Iron + Water → Iron(II,III) oxide + Hydrogen. Symbol equation: \( \ce{Fe + H2O -> Fe3O4 + H2} \)

Step I. Count atoms.

Hydrogen is equal, but Fe and O are not.

Step II. Oxygen is 1 on the left and 4 on the right, so write 4 before water: \( \ce{Fe + 4H2O -> Fe3O4 + H2} \)

Step III. Hydrogen is now 8 on the left, so write 4 before H₂: \( \ce{Fe + 4H2O -> Fe3O4 + 4H2} \)

Step IV. Iron is 3 on the right, so write 3 before Fe: \( \ce{3Fe + 4H2O -> Fe3O4 + 4H2} \)

Step V. Final check: Fe 3 and 3, H 8 and 8, O 4 and 4. The equation is balanced.

(d) What a chemical equation tells us

1. Which substances react and which are formed.

2. How many molecules of each substance take part.

3. The masses involved. In \( \ce{CaCO3 -> CaO + CO2} \), 100 g of calcium carbonate gives 56 g of calcium oxide and 44 g of carbon dioxide. The mass on both sides is 100 g because matter is not lost.

4. The volume of a gas formed. In the same reaction one mole of CO₂ occupies 22.4 litres at STP.

5. The valency of an element. In \( \ce{Mg + 2HCl -> MgCl2 + H2} \), one Mg atom displaces two H atoms, so the valency of magnesium is 2.

Unless something else is stated, equations are understood to be written under ordinary conditions of temperature and pressure.

(e) Limitations of equations and how to remove them

A bare equation does not tell us: the physical states; the concentration; the speed of the reaction; the favourable conditions (temperature, pressure, catalyst); whether energy is absorbed or released; whether the reaction is reversible; or the precautions needed. We make equations more informative as follows.

1. State symbols. (s) solid, (l) liquid, (g) gas, (aq) aqueous solution (dissolved in water). A gas formed in the reaction may also be shown by an upward arrow (↑) and a precipitate by a downward arrow (↓) or by 'ppt.'. They matter most in reactions which give out or take in heat: \( \ce{2H2(g) + O2(g) -> 2H2O(l) + "572 kJ"} \) and \( \ce{2H2(g) + O2(g) -> 2H2O(g) + "484 kJ"} \). Both are exothermic, but the heat released is different when the water formed is a liquid and when it is a vapour.

Examples: \( \ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)} \); \( \ce{Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)} \); \( \ce{AgNO3(aq) + NaCl(aq) -> AgCl(s) + NaNO3(aq)} \)

2. Concentration. Write 'dilute' or 'conc.' below the formula, for example \( \ce{Mg + H2SO4 -> MgSO4 + H2} \) with H₂SO₄ marked dilute.

3. Conditions. Write temperature, pressure and catalyst above and below the arrow. The Haber process: \( \ce{N2 + 3H2 <=> 2NH3} \) at 500 °C, 200 atm with Fe/Mo catalyst, giving 22400 calories of heat.

4. Direction. A reversible reaction has a double (half) arrow and an irreversible reaction has a single arrow.

5. Heat change. Heat released is written on the right side with a + sign (exothermic): \( \ce{N2 + 3H2 <=> 2NH3} \) + @@22400 cal@@. Heat absorbed is written with a - sign (endothermic): \( \ce{N2 + O2 <=> 2NO} \) - @@43200 cal@@.

Solved Examples

Example 1. Write the balanced equation for the following chemical reactions. (i) Hydrogen + Chlorine → Hydrogen chloride (ii) Barium chloride + Aluminium sulphate → Barium sulphate + Aluminium chloride (iii) Sodium + Water → Sodium hydroxide + Hydrogen
Answer:
(i) \( \ce{H2(g) + Cl2(g) -> 2HCl(g)} \)
(ii) \( \ce{3BaCl2 + Al2(SO4)3 -> 3BaSO4 + 2AlCl3} \)
(iii) \( \ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)} \)

Example 2. Write a balanced chemical equation with state symbols for the following reactions. (i) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride. (ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water.
Answer:
(i) \( \ce{BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)} \)
(ii) \( \ce{NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l)} \)

C. Types of Chemical Reactions

(a) Slow and fast reactions

Reactions may be classified on the basis of how fast they go.

Fast reactions are over in an instant on mixing the reactants. They are mostly ionic reactions. The reaction between a strong acid and a strong base is complete in about 10^-10 second: \( \ce{NaOH + HCl -> NaCl + H2O} \). Silver nitrate and HCl give white silver chloride at once: \( \ce{AgNO3 + HCl -> AgCl + HNO3} \). The book also describes photosynthesis in plants as a very fast reaction with a half-life of about 10^-12 second. (The half-life of a reaction is the time in which half of the reactants have changed into products.)

Slow reactions take hours, days or even years. Rusting of iron is the best example: \( \ce{4Fe + 3O2 + 6H2O -> 2Fe2O3.3H2O} \) (rust). Other examples are the heating of potassium chlorate, \( \ce{2KClO3 -> 2KCl + 3O2} \), and the formation of ethyl acetate from acetic acid and ethanol, \( \ce{CH3COOH + C2H5OH -> CH3COOC2H5 + H2O} \).

(b) Irreversible and reversible reactions

Irreversible reactions go only in the forward direction. The reactants are used up and the products do not turn back into reactants. They are written with a single arrow. Examples: \( \ce{C + O2 -> CO2} \) (coal burning) and \( \ce{CH4 + 2O2 -> CO2 + 2H2O} \) (methane burning). The products are so stable that they do not react to give back methane.

Reversible reactions go in both directions at the same time, so some of the reactants are always present. They are written with a double half-arrow: \( \ce{A + B <=> C + D} \). The left-to-right change is the forward reaction and the right-to-left change is the backward reaction.

Neutralisation and the pH of the solution. When a strong acid and a strong base of the same strength react in equal amounts, both ionise fully. All the H⁺ ions combine with all the OH^- ions to give water, so the solution is neutral with pH 7. \( \ce{HCl + NaOH -> NaCl + H2O} \), or in ionic form \( \ce{H+ + Cl- + Na+ + OH- -> Na+ + Cl- + H2O} \), and the net change is \( \ce{H+ + OH- -> H2O} \).

A strong acid with a weak base, \( \ce{HCl + NH4OH -> NH4Cl + H2O} \): the weak base ionises only partly, so after reaction extra H⁺ remains and the pH is less than 7.

A weak acid with a strong base, \( \ce{CH3COOH + NaOH -> CH3COONa + H2O} \): the weak acid ionises only partly, so extra OH^- remains and the pH is more than 7.

(c) Addition (combination) reactions

Two or more substances unite to form a single new substance. Heat, light, electricity or pressure may bring it about. Example: \( \ce{H2 + Cl2 -> 2HCl} \). They can occur in three ways.

1. Element + element. This is also a synthesis reaction, a reaction in which a substance is built from its elements, for instance ammonia by the Haber process, \( \ce{N2 + 3H2 -> 2NH3} \). Others: \( \ce{2H2 + O2 -> 2H2O} \), \( \ce{2Mg + O2 -> 2MgO} \), \( \ce{2Na + Cl2 -> 2NaCl} \). Every synthesis reaction is an addition reaction, but not every addition reaction is a synthesis.

2. Compound + compound: \( \ce{NH3 + HCl -> NH4Cl} \), \( \ce{CaO + CO2 -> CaCO3} \), and \( \ce{CH2=CH2 + Br2 -> CH2Br-CH2Br} \) (1,2-dibromoethane).

3. Element + compound: \( \ce{2CO + O2 -> 2CO2} \) and \( \ce{2SO2 + O2 -> 2SO3} \).

In an addition reaction only a single product is formed.

(d) Decomposition reactions

One substance breaks into two or more simpler substances. It is the opposite of an addition reaction. The kinds are:

1. Thermal decomposition (by heat): \( \ce{CaCO3 ->[\Delta] CaO + CO2} \), \( \ce{2Pb(NO3)2 ->[\Delta] 2PbO + 4NO2 + O2} \).

2. Electrolysis (by electric current): \( \ce{2H2O ->[electricity] 2H2 + O2} \), \( \ce{2NaCl "(molten)" ->[electricity] 2Na + Cl2} \), \( \ce{2Al2O3 ->[electricity] 4Al + 3O2} \).

3. Photodecomposition (by light): \( \ce{2AgBr ->[light] 2Ag + Br2} \), \( \ce{2AgCl ->[light] 2Ag + Cl2} \).

4. Analysis reaction: a compound breaks into its elements, for example \( \ce{2HgO ->[\Delta] 2Hg + O2} \) and \( \ce{2HI ->[\Delta] H2 + I2} \). Every analysis reaction is a decomposition reaction, but every decomposition is not an analysis reaction.

(e) Displacement reactions

One element takes the place of another in a compound. In general: a more active metal displaces a less active metal from its compound; some active non-metals displace less active non-metals; and some metals displace hydrogen (a non-metal) from acids.

1. A more active metal displaces a less active metal. An iron nail dipped in copper sulphate solution gets a coating of copper: \( \ce{Fe(s) + CuSO4(aq) -> FeSO4(aq) + Cu(s)} \). Iron has taken the place of copper, so iron is more reactive than copper.

2. A more active non-metal displaces a less active non-metal. Chlorine passed into sodium bromide solution gives sodium chloride and bromine: \( \ce{2NaBr(aq) + Cl2(g) -> 2NaCl(aq) + Br2(aq)} \), so chlorine is more reactive than bromine. Chlorine passed into sodium iodide solution gives iodine: \( \ce{2NaI(aq) + Cl2(g) -> 2NaCl(aq) + I2(aq)} \), so chlorine is more reactive than iodine.

3. An active metal displaces hydrogen from an acid. \( \ce{Zn(s) + H2SO4(aq) -> ZnSO4(aq) + H2(g)} \), so zinc is more reactive than hydrogen.

(f) Double displacement reactions

The two compounds exchange their ions (radicals) and two new compounds are formed. Examples: \( \ce{NaCl(aq) + AgNO3(aq) -> AgCl(s) + NaNO3(aq)} \) and \( \ce{BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)} \). All acid-base neutralisation reactions are double displacement reactions.

(g) Oxidation-reduction (redox) reactions

Rusting of iron, fading of clothes, and burning of cooking gas, wood and coal are all redox reactions. At first, oxidation and reduction were defined only through oxygen and hydrogen, but the idea was later widened to include other elements and finally electrons.

Oxidation (the classical definitions)

Oxidation of a substance means any of the following:

1. Addition of oxygen. \( \ce{C(s) + O2(g) -> CO2(g)} \); \( \ce{2Mg(s) + O2(g) -> 2MgO(s)} \); \( \ce{P4(s) + 5O2(g) -> 2P2O5(s)} \); \( \ce{2H2(g) + O2(g) -> 2H2O(l)} \); \( \ce{2Cu(s) + O2(g) -> 2CuO(s)} \); \( \ce{2CO(g) + O2(g) -> 2CO2(g)} \)

2. Removal of hydrogen. \( \ce{H2S -> H2 + S} \); ethanol with a copper catalyst at 300 °C: \( \ce{CH3CH2OH ->[Cu, 300 C] CH3CHO + H2} \) (ethanol has 6 H atoms and ethanal 4, so two H atoms are lost); \( \ce{H2S(aq) + Br2(aq) -> 2HBr(aq) + S} \); \( \ce{4HCl + MnO2 ->[heat] MnCl2 + 2H2O + Cl2} \); \( \ce{2H2S + O2 -> 2H2O + 2S} \). In each of these, hydrogen is removed from H₂S or HCl, which is oxidised.

3. Addition of an electronegative element. \( \ce{2Fe + 3Cl2 -> 2FeCl3} \) and \( \ce{SnCl2 + Cl2 -> SnCl4} \): the amount of chlorine combined with the substance increases.

4. Removal of an electropositive element. \( \ce{2KI + Cl2 -> 2KCl + I2} \): potassium is lost from KI, so KI is oxidised. Likewise \( \ce{2NaBr + Cl2 -> 2NaCl + Br2} \): NaBr is oxidised to Br₂.

In short, a substance is oxidised if it gains oxygen or an electronegative element, or loses hydrogen or an electropositive element.

Reduction (the classical definitions)

1. Addition of hydrogen. \( \ce{H2 + S -> H2S} \); \( \ce{Cl2 + H2S -> 2HCl + S} \) (here chlorine is reduced to HCl while H₂S is oxidised to sulphur); \( \ce{H2 + Cl2 -> 2HCl} \); \( \ce{2Na + H2 -> 2NaH} \); \( \ce{C2H4 + H2 -> C2H6} \)

2. Removal of oxygen. \( \ce{H2O + C -> CO + H2} \) (steam and coke give water gas; water is reduced to hydrogen); \( \ce{ZnO + H2 -> Zn + H2O} \); \( \ce{ZnO + C -> Zn + CO} \); \( \ce{CuO + H2 -> Cu + H2O} \); \( \ce{2MgO -> 2Mg + O2} \)

3. Addition of an electropositive element. \( \ce{Cl2 + Mg -> MgCl2} \) and \( \ce{I2 + Hg -> HgI2} \): Cl₂ and I₂ are reduced.

4. Removal of an electronegative element. \( \ce{2FeCl3 -> 2FeCl2 + Cl2} \) and \( \ce{FeS -> Fe + S} \).

So a substance is reduced if it gains hydrogen or an electropositive element, or loses oxygen or an electronegative element.

Oxidising and reducing agents

An oxidising agent is a substance that causes oxidation of another substance, by giving it oxygen or an electronegative element, or by removing hydrogen or an electropositive element from it. A reducing agent is a substance that causes reduction of another, by removing oxygen or an electronegative element or by giving hydrogen or an electropositive element to it.

In \( \ce{CuO + H2 -> Cu + H2O} \), hydrogen takes oxygen away from copper oxide. So CuO is reduced and H₂ is the reducing agent. At the same time CuO gives oxygen to hydrogen, so H₂ is oxidised and CuO is the oxidising agent.

Points to remember: (i) the substance that gains oxygen or loses hydrogen is oxidised; (ii) the substance that loses oxygen or gains hydrogen is reduced; (iii) the substance that is oxidised is the reducing agent; (iv) the substance that is reduced is the oxidising agent.

Modern (electronic) concept

There are two modern ideas, the electronic concept and the oxidation number concept. The electronic concept is as follows.

Oxidation is the loss of electrons by an atom, ion or molecule (de-electronation).

1. Atoms: \( \ce{Na -> Na+ + e-} \), \( \ce{Ca -> Ca^2+ + 2e-} \), \( \ce{Zn -> Zn^2+ + 2e-} \)

2. Cations: \( \ce{Fe^2+ -> Fe^3+ + e-} \), \( \ce{Sn^2+ -> Sn^4+ + 2e-} \), \( \ce{Cu+ -> Cu^2+ + e-} \)

3. Anions: \( \ce{2Cl- -> Cl2 + 2e-} \), \( \ce{MnO4^2- -> MnO4- + e-} \), \( \ce{2O^2- -> O2 + 4e-} \)

4. Molecules: \( \ce{H2S -> H2 + S} \) or \( \ce{H2S -> 2H+ + S^2-} \)

In oxidation the positive charge rises or the negative charge falls.

Reduction is the gain of one or more electrons by an atom, ion or molecule (electronation).

1. Atoms: \( \ce{Cl + e- -> Cl-} \), \( \ce{S + 2e- -> S^2-} \)

2. Ions: \( \ce{Mg^2+ + 2e- -> Mg} \), \( \ce{Pb^4+ + 2e- -> Pb^2+} \), \( \ce{Cr2O7^2- -> 2Cr^3+} \)

3. Molecules: \( \ce{MnO2 -> Mn^2+} \)

In reduction the negative charge rises or the positive charge falls.

Oxidation and reduction compared

Oxidation number

The oxidation number (O.N.) of an atom is the charge it has in its ionic form, or the charge it appears to have when all the other atoms are removed from it as ions. Rules:

1. The O.N. of an element in its free state is zero, as in N₂, Cl₂, P₄, S₈, He and Fe. It is zero in any allotropic form and for all the atoms in an alloy.

2. The O.N. of a monoatomic ion equals its charge: Na⁺ is +1, Mg²⁺ is +2, Al³⁺ is +3, Cl^- is -1, S^2- is -2, N^3- is -3.

3. Hydrogen is +1 when combined with non-metals and -1 in hydrides with active metals.

4. Oxygen is -2 in most compounds, but -1 in peroxides. In OF₂ it is +2, since fluorine, the most electronegative element, is -1.

5. In binary compounds of a metal and a non-metal, the metal is positive and the non-metal negative. Alkali metals are always +1 and alkaline earth metals always +2. Halogens are usually -1 and fluorine is always -1, but other halogens can be positive with more electronegative atoms: Cl in Cl₂O is +1, in Cl₂O₇ is +7, and iodine in IF₇ is +7.

6. When two non-metals combine, the one with higher electronegativity gets the negative number. Nitrogen is -3 in NH₃ but +3 in NCl₃.

7. The O.N. of all atoms in a neutral molecule add up to zero.

8. In an ion, the O.N. add up to the charge on the ion; in SO₄^2- the sum is -2.

Method. Write the formula, put the known O.N. above each atom and x for the unknown one, multiply by the number of atoms, equate the total to 0 (molecule) or to the charge (ion), and solve for x.

Example. O.N. of S in H₂S: \( 2(+1) + x = 0 \Rightarrow x = -2 \).

Worked examples on oxidation numbers

Redox reactions in terms of electrons

Oxidation and reduction occur together: while one atom, ion or molecule loses electrons, another gains them. Reactions in which both happen are redox reactions.

1. \( \ce{2Mg + O2 -> 2MgO} \) (that is Mg²⁺ and O^2-): Mg loses electrons (oxidation) and oxygen gains them (reduction).

2. \( \ce{Fe + S ->[\Delta] FeS} \) (Fe²⁺, S^2-): Fe is oxidised and S is reduced.

3. \( \ce{Cu + I2 -> CuI2} \) (as given in the book): copper loses electrons and iodine gains them.

4. \( \ce{2Al + 6HCl -> 2AlCl3 + 3H2} \), which in ionic form is \( \ce{2Al + 6H+ -> 2Al^3+ + 3H2} \): aluminium loses electrons and hydrogen ions gain them.

By the electronic concept, an oxidising agent gains electrons (is reduced) and a reducing agent loses electrons (is oxidised). In \( \ce{CuSO4 + Fe -> FeSO4 + Cu} \), written ionically as \( \ce{Cu^2+ + Fe -> Fe^2+ + Cu} \), iron loses electrons and is oxidised, so it is the reducing agent, whereas Cu²⁺ of copper sulphate gains electrons and is reduced, so copper sulphate is the oxidising agent.

Solved Examples

Example 1. Write the balanced chemical equation for the following and identify the type of reaction in each case. (a) Potassium bromide(aq) + Barium iodide(aq) → Potassium iodide(aq) + Barium bromide(s) (b) Zinc carbonate(s) → Zinc oxide(s) + Carbon dioxide(g) (c) Hydrogen(g) + Chlorine(g) → Hydrogen chloride(g) (d) Magnesium(s) + Hydrochloric acid(aq) → Magnesium chloride(aq) + Hydrogen(g)
Answer:
(a) \( \ce{2KBr(aq) + BaI2(aq) -> 2KI(aq) + BaBr2(s)} \); double displacement reaction
(b) \( \ce{ZnCO3(s) -> ZnO(s) + CO2(g)} \); decomposition reaction
(c) \( \ce{H2(g) + Cl2(g) -> 2HCl(g)} \); combination reaction
(d) \( \ce{Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)} \); displacement reaction

Example 2. Identify the substances that are oxidised and the substances that are reduced in the following reactions. (i) \( \ce{4Na(s) + O2(g) -> 2Na2O(s)} \) (ii) \( \ce{CuO(s) + H2(g) -> Cu(s) + H2O(l)} \)
Answer:
(i) Sodium is oxidised as it gains oxygen, and oxygen is reduced.
(ii) Copper oxide is reduced to copper, while hydrogen is oxidised to water.

D. Effects of Oxidation Reactions in Daily Life

Oxygen is vital for life, and its reactions touch our lives in many ways, mostly useful and some harmful.

(a) Combustion

A reaction in which a substance burns, that is, is oxidised in air or oxygen, is called combustion. Kerosene, coal, charcoal and wood all undergo combustion. Methane of natural gas burns in excess oxygen: \( \ce{CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l) + "energy"} \). Butane, the chief constituent of LPG, burns as: \( \ce{2C4H10(g) + 13O2(g) -> 8CO2(g) + 10H2O(g) + "energy"} \). (The book writes only 5H₂O on the right, which does not balance; the correct number is 10.) All combustion reactions are exothermic. The body works like a furnace in which food is oxidised, and carbohydrates such as glucose, fructose and starch are the main fuel: \( \ce{C6H12O6(s) + 6O2(g) -> 6CO2(g) + 6H2O(l) + "energy"} \). Not all combustion shows a flame; combustion is simply oxidation with release of energy.

(b) Respiration

Respiration is the main biochemical reaction that releases energy in the cells. Oxygen of the air we breathe in enters the lungs, passes into tiny air sacs (alveoli) and diffuses into the blood, where it binds to haemoglobin in red blood cells and is carried to all the cells. There, glucose is oxidised to carbon dioxide and water, and the energy released runs the cell activities, heart and muscles and warms the body. Carbon dioxide and water go back into the blood and are breathed out. All living things respire. Fish take dissolved oxygen through their gills and plants take in air through the pores (stomata) of their leaves.

(c) Harmful effects of oxidation

1. Pollution. Combustion of fuels releases poisonous gases such as carbon monoxide, sulphur dioxide, sulphur trioxide and oxides of nitrogen into the air.

2. Corrosion. It is the slow eating away of a metal by air gases (oxygen, carbon dioxide, hydrogen sulphide) and water vapour, which convert the metal to its oxide, carbonate or sulphide.

(a) Rusting of iron. Iron left in damp air for long gets a brown flaky coating of rust, which is mainly hydrated ferric oxide, \( \ce{Fe2O3.xH2O} \): \( \ce{2Fe(s) + 3/2 O2(g) + xH2O(l) -> Fe2O3.xH2O(s)} \)

(b) Copper loses its shine and gets a green coating of basic copper carbonate, \( \ce{CuCO3.Cu(OH)2} \), by the action of oxygen, carbon dioxide and water vapour: \( \ce{2Cu(s) + CO2(g) + O2(g) + H2O(l) -> CuCO3.Cu(OH)2} \)

(c) Silver tarnishes because hydrogen sulphide in the air forms a black coating of silver sulphide: \( \ce{2Ag(s) + H2S(g) -> Ag2S(s) + H2(g)} \)

Rusting is a serious problem because it weakens bridges, railings and machine parts and costs a great deal to replace. The rust does not stick to the surface and flakes off to expose fresh iron, so rusting never stops by itself.

Prevention of rusting: keep iron away from damp air by (i) painting, (ii) greasing or oiling, (iii) galvanisation (a thin coating of zinc).

3. Rancidity. When foods containing fats and oils are stored for long, air oxidises them and they develop a bad smell and taste. This is rancidity and it makes the food unfit to eat. Ways to prevent or slow it:

1. Add antioxidants, which are reducing agents that stop oxidation. BHA (butylated hydroxy anisole) and BHT (butylated hydroxy toluene) are common ones.

2. Pack the food in nitrogen gas, which is unreactive and keeps out oxygen. Chips bags are filled with nitrogen for this reason.

3. Keep it in a refrigerator; the low temperature slows the oxidation.

4. Store it in air-tight containers, which limit contact with oxygen.

5. Keep it away from light, since light speeds oxidation.