Chemistry

Chemical Reactions and Equations

act 1. burning of a magnesium ribbon in air & collection of

magnesium oxide in a watch-glass

the ribbon burnt with a white flame & changes into magnesium oxide
(white powder) due to the reaction between magnesium & oxygen in
the air

magnesium + oxygen -> magnesium oxide

(reactants) (products)

skeletal equation - 𝑀𝑔 + 𝑂2 −> 𝑀𝑔𝑂 (also oxidation reaction)

act 2. lead nitrate + potassium iodide -> lead iodide + potassium nitrate (also double displacement reaction)

act 3. formation of hydrogen gas by the action of dilute sulphuric acid on

zinc

sulphuric acid + zinc (granules) -> zinc sulphate + hydrogen

a chemical change occurs if there’s

○ change in state
○ change in colour
○ evolution of a gas
○ change in temperature

Chemical Equations
reactants - substances that undergo chemical change in the reaction

product - the new substance formed during the reaction

A word-equation shows change of reactants to products through an arrow placed between them. The reactants are
written on the left-hand side (LHS) with a plus sign (+) between them. Similarly, products are written on the
right-hand side (RHS) with a plus sign (+) between them. The arrowhead points towards the products, and shows the
direction of the reaction

✽ Writing a Chemical Equation

balanced equation - equation with the same number of atoms of each element on both sides
skeletal equation - unbalanced equation with different masses on both sides of the equation

✽ Balanced Chemical Equations

according to the law of conservation of mass, the total mass of the elements present in the products of a chemical
reaction has to be equal to the total mass of elements present in the reactants. The number of atoms of each
element remains the same, before & after a chemical reaction

act 3. 𝑍𝑛 + 𝐻2𝑆𝑂4 −> 𝑍𝑛𝑆𝑂4 + 𝐻2

element number of atoms in reactants (LHS) number of atoms in products (RHS)

Zn 1 1

H 2 2

S 1 1

O 4 4

ex. 𝐹𝑒 + 𝐻2𝑂 −> 𝐹𝑒3𝑂4 + 𝐻2

step 1: draw boxes around each formula & don’t change anything inside the boxes

𝐹𝑒 + 𝐻2𝑂 −> 𝐹𝑒3𝑂4 + 𝐻2

step 2: list the number of atoms of different elements present in the unbalanced equation

element number of atoms in reactants (LHS) number of atoms in products (RHS)

Fe 1 3

H 2 2

O 1 4

step 3: start balancing with the compound that contains the maximum number of atoms (reactant / product). In that
compound, select the element with the maximum number of atoms

here, select 𝐹𝑒3𝑂4 & the element oxygen in it. There are 4 oxygen atoms on the RHS & 1 on the LHS

atoms of oxygen in reactants in products

initial 1 (in 𝐻2𝑂) 4 (in 𝐹𝑒3𝑂4)

to balance 1x4 4

Note: to balance oxygen atoms, the coefficient is ‘4’ as 4𝐻2𝑂 & not 𝐻2𝑂4 or (𝐻2𝑂)4

𝐹𝑒 + 4 𝐻2𝑂 −> 𝐹𝑒3𝑂4 + 𝐻2 (partly balanced)

step 4: Fe & H atoms aren’t balanced

to balance the number of H atoms, make the number of molecules of H as 4 on the RHS
 atoms of hydrogen in reactants in products

initial 8 (in 4 𝐻2𝑂) 2 (in 𝐻2 )

to balance 8 2x4

(partly balanced)

step 5: Fe is the only element left to be balanced

atoms of iron in reactants in products

initial 1 (in 𝐹𝑒 ) 3 (in 𝐹𝑒 𝑂 )

3 4

to balance 1x3 3

to balance Fe, we take 3 atoms of Fe on the LHS

the fully balanced equation is 3 𝐹𝑒 + 4 𝐻2𝑂 −> 𝐹𝑒3𝑂4 + 4 𝐻2

step 6: representing the physical states of compounds to make the chemical equations more informative
○ gaseous (g)
○ liquid (l)
○ aqueous (aq) - a solution in water
○ solid (s)

the balanced equation becomes 3 𝐹𝑒 (𝑠) + 4 𝐻2𝑂 (𝑔) −> 𝐹𝑒3𝑂4 (𝑠) + 4 𝐻2 (𝑔)

sometimes, the reaction conditions (temperature, pressure, catalyst) are indicated above / or below the arrow

340 𝑎𝑡𝑚
ex. 𝐶𝑂 (𝑔) + 2𝐻2 (𝑔) −−−−−−−> 𝐶𝐻3𝑂𝐻 (𝑙)

𝑠𝑢𝑛𝑙𝑖𝑔ℎ𝑡
ex. 6𝐶𝑂 (𝑎𝑞) + 12𝐻2𝑂 (𝑙) −−−−−−−−−> 𝐶6𝐻12𝑂6 (𝑎𝑞) + 6𝑂2 (𝑎𝑞) + 6𝐻2𝑂 (𝑙)
2
𝑐ℎ𝑙𝑜𝑟𝑜𝑝ℎ𝑦𝑙𝑙

Types of Chemical Reactions

chemical reactions involve the breaking & making of bonds between atoms to produce new substances

✽ Combination Reaction
act 4. formation of slaked lime by the reaction of calcium oxide with water

calcium oxide reacts vigorously with water to produce slaked lime (calcium
hydroxide), releasing a large amount of heat

𝐶𝑎𝑂 (𝑠) + 𝐻2𝑂 (𝑙) −> 𝐶𝑎(𝑂𝐻)2 (𝑎𝑞) + 𝐻𝑒𝑎𝑡

(𝑞𝑢𝑖𝑐𝑘 𝑙𝑖𝑚𝑒) (𝑠𝑙𝑎𝑘𝑒𝑑 𝑙𝑖𝑚𝑒)
Calcium oxide & water combine to form a single product, calcium hydroxide

combination reaction - reactions in which a single product is formed from 2 or more reactants
○ burning of coal - 𝐶 (𝑠) + 𝑂2 −> 𝐶𝑂2 (𝑔)

○ formation of water from 𝐻2 (𝑔) & 𝑂𝑠 (𝑔) - 2𝐻 (𝑔) + 02 (𝑔) −> 2𝐻2𝑂 (𝑙)
2

exothermic chemical reactions - reactions in which heat is released along with the formation of products
○ burning of natural gas - 𝐶𝐻 (𝑔) + 2𝑂2 (𝑔) −> 𝐶𝑂2 (𝑔) + 2𝐻2𝑂 (𝑔)
4

○ respiration - 𝐶 𝐻12𝑂6 (𝑎𝑞) + 6𝑂2 (𝑎𝑞) −−−−−−−−−> 6𝐻2𝑂 (𝑙) + 6𝐶𝑂2 (𝑎𝑞) + 𝑒𝑛𝑒𝑟𝑔𝑦
6
during digestion, food is broken down into simpler substances to form glucose, which combined with oxygen
in the cells of our body to provide energy
○ decomposition of vegetable matter into compost

𝐶𝑎𝐶𝑂3

𝐶𝑎(𝑂𝐻)2 (𝑎𝑞) + 𝐶𝑂2 (𝑔) −> 𝐶𝑎𝐶𝑂3 (𝑠) + 𝐻2𝑂 (𝑙)

𝑐𝑎𝑙𝑐𝑖𝑢𝑚 ℎ𝑦𝑑𝑟𝑜𝑥𝑖𝑑𝑒 𝑐𝑎𝑙𝑐𝑖𝑢𝑚 𝑐𝑎𝑟𝑏𝑜𝑛𝑎𝑡𝑒

✽ Decomposition reaction
decomposition reaction - a single reactant breaks down to give simpler products

correct way of heating the boiling tube containing crystals of

ferrous sulphate & of smelling the odour

the green colour of the ferrous sulphate crystals changes & there’s
an odour of burning sulphur

ℎ𝑒𝑎𝑡
2𝐹𝑒𝑆𝑂4 (𝑠) −−−−> 𝐹𝑒2𝑂3 (𝑠) + 𝑆𝑂2 (𝑔) + 𝑆03 (𝑔)
(𝑓𝑒𝑟𝑟𝑜𝑢𝑠 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒) (𝑓𝑒𝑟𝑟𝑖𝑐 𝑜𝑥𝑖𝑑𝑒)

Ferrous sulphate crystals (𝐹𝑒𝑆𝑂4 , 7𝐻2𝑂) lose water when heated &
the colour of the crystals changes, then decomposing into ferric
oxide (𝐹𝑒2𝑂3), sulphur dioxide (𝑆𝑂2 ) & sulphur trioxide (𝑆03 )

endothermic reactions - decomposition reactions in which energy (heat, light, electricity) is absorbed

thermal decomposition - when a decomposition reaction is carried out by heating

○ decomposition of calcium carbonate to calcium oxide (quick lime) & carbon dioxide -
ℎ𝑒𝑎𝑡
𝐶𝑎𝐶𝑂3 (𝑠) −−−−> 𝐶𝑎𝑂 (𝑠) + 𝐶𝑂2 (𝑔)
(𝑙𝑖𝑚𝑒𝑠𝑡𝑜𝑛𝑒) (𝑞𝑢𝑖𝑐𝑘 𝑙𝑖𝑚𝑒)
 it’s used in the manufacture of cement
○ heating of lead nitrate & emission of nitrogen dioxide -
ℎ𝑒𝑎𝑡
2𝑃𝑏(𝑁𝑂3)2 (𝑠) −−−−> 2𝑃𝑏𝑂 (𝑠) + 4𝑁𝑂2 (𝑔) + 𝑂2 (𝑔)

(𝑙𝑒𝑎𝑑 𝑛𝑖𝑡𝑟𝑎𝑡𝑒) (𝑙𝑒𝑎𝑑 𝑜𝑥𝑖𝑑𝑒) (𝑛𝑖𝑡𝑟𝑜𝑔𝑒𝑛 𝑑𝑖𝑜𝑥𝑖𝑑𝑒) (𝑜𝑥𝑦𝑔𝑒𝑛)

there’s an emission of brown fumes of nitrogen dioxide (𝑁𝑂2)

electrolysis of water

1. insert carbon electrodes in the rubber stoppers & connect to a 6 volt

battery
2. add a few drops of sulphuric acid (to ionise the water to conduct
electricity)
3. switch on the current & leave undisturbed

there’s a formation of bubbles at both electrodes, displacing water in

the test tubes. Hydrogen, a cation (positive ion) is formed at the
cathode (negative electrode). Oxygen, an anion (negative ion) is
formed at the anode (positive electrode)

silver chloride turns grey in sunlight to form silver metal

it decomposes into silver & chlorine in light

𝑠𝑢𝑛𝑙𝑖𝑔ℎ𝑡
2𝐴𝑔𝐶𝑙 (𝑠) −−−−−−> 2𝐴𝑔 (𝑠) + 𝐶𝑙2 (𝑔)

silver bromide also behaves the same way

𝑠𝑢𝑛𝑙𝑖𝑔ℎ𝑡
2𝐴𝑔𝐵𝑟 (𝑠) −−−−−−> 2𝐴𝑔 (𝑠) + 𝐵𝑟2 (𝑔)

these reactions are used in black & white photography

act 5. barium hydroxide & ammonium chloride absorb heat to react and form barium chloride, ammonia and water.
The bottom of the test tube becomes cold

𝐵𝑎(𝑂𝐻)2 (𝑎𝑞) + 2𝑁𝐻4𝐶𝑙 (𝑎𝑞) −> 𝐵𝑎𝐶𝑙2 (𝑎𝑞) + 2𝑁𝐻3 (𝑔) + 𝐻2𝑂 (𝑙)
 ✽ Displacement reaction
displacement reaction - when a more reactive element displaces (removes another element) from a solution

reactivity series

iron nails dipped in copper sulphate solution

1. in one test tube, immerse 2 iron nails in a solution of copper sulphate.

Leave another test tube of the solution aside

2. after 20 minutes, compare the intensity of the blue colour in both the test
tubes & compare the colour of the iron nails dipped in the solution with a
regular one

the iron nail becomes brown in colour & the blue colour of the copper
sulphate solution fades

iron nails & copper sulphate solutions compared before & after the experiment

iron displaces copper from the solution

𝐹𝑒 (𝑠) + 𝐶𝑢𝑆𝑂4 (𝑎𝑞) −> 𝐹𝑒𝑆𝑂4 (𝑎𝑞) + 𝐶𝑢 (𝑠)

(𝑐𝑜𝑝𝑝𝑒𝑟 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒) (𝑖𝑟𝑜𝑛 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒)

ex. 𝑍𝑛 (𝑠) + 𝐶𝑢𝑆𝑂4 (𝑎𝑞) −> 𝑍𝑛𝑆𝑂4 (𝑎𝑞) + 𝐶𝑢 (𝑠)

(𝑐𝑜𝑝𝑝𝑒𝑟 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒) (𝑧𝑖𝑛𝑐 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒)
ex. 𝑃𝑏 (𝑠) + 𝐶𝑢𝐶𝑙2 (𝑎𝑞) −> 𝑃𝑏𝐶𝑙2 (𝑎𝑞) + 𝐶𝑢 (𝑠)
(𝑐𝑜𝑝𝑝𝑒𝑟 𝑐ℎ𝑙𝑜𝑟𝑖𝑑𝑒) (𝑙𝑒𝑎𝑑 𝑐ℎ𝑙𝑜𝑟𝑖𝑑𝑒)

✽ Double Displacement reaction

double displacement reactions - reactions where there’s an exchange of ions between the reactants

precipitation reaction - any reaction that produces a precipitate

formation of barium sulphate & sodium chloride

a precipitate (white substance, insoluble in water) 𝐵𝑎𝑆𝑂4 is formed by the

2− 2+
reaction of 𝑆𝑂4 & 𝐵𝑎 . The other product formed is sodium chloride which
remains in the solution

𝑁𝑎2𝑆𝑂4 (𝑎𝑞) + 𝐵𝑎𝐶𝑙2 (𝑎𝑞) −> 𝐵𝑎𝑆𝑂4 (𝑠) + 2𝑁𝑎𝐶𝑙 (𝑎𝑞)

(𝑠𝑜𝑑𝑖𝑢𝑚 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒) (𝑏𝑎𝑟𝑖𝑢𝑚 𝑐ℎ𝑙𝑜𝑟𝑖𝑑𝑒) (𝑏𝑎𝑟𝑖𝑢𝑚 𝑠𝑢𝑙𝑝ℎ𝑎𝑡𝑒) (𝑠𝑜𝑑𝑖𝑢𝑚 𝑐ℎ𝑙𝑜𝑟𝑖𝑑𝑒)

✽ Oxidation & Reduction

oxidation - when a substance gains oxygen / loses hydrogen during a reaction

oxidising agent / oxidant - reactant that helps in oxidation

reduction - when a substance loses oxygen / gains hydrogen during a reaction

reducing agent / reductant - reactant that helps in reduction

redox (oxidation-reduction) reaction - when one reactant is oxidised & the other is reduced during a reaction

oxidation of copper to copper oxide

the surface of the copper powder becomes coated with black copper (II)
oxide because oxygen is added to copper

ℎ𝑒𝑎𝑡
2𝐶𝑢 + 𝑂2 −−−−> 2𝐶𝑢𝑂

if hydrogen is passed over this heated material (CuO), the black coating turns brown as the reverse reaction takes
place & copper is obtained

ℎ𝑒𝑎𝑡
𝐶𝑢𝑂 + 𝐻2 −−−−> 𝐶𝑢 + 𝐻2

during this reaction, the copper (II) oxide is losing oxygen & is being reduced. The hydrogen is gaining oxygen & is
being oxidised

𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛

ℎ𝑒𝑎𝑡
𝐶𝑢𝑂 + 𝐻2 −−−−> 𝐶𝑢 + 𝐻2

𝑟𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛
ex. 𝑍𝑛𝑂 + 𝐶 −> 𝑍𝑛 + 𝐶𝑂
carbon is oxidised to CO & ZnO is reduced to Zn

ex. 𝑀𝑛𝑂 + 4𝐻𝐶𝑙 −> 𝑀𝑛𝐶𝑙2 + 2𝐻2𝑂 + 𝐶𝑙2

2
HCl is oxidised to 𝐶𝑙2 & 𝑀𝑛𝑂2 is reduced to 𝑀𝑛𝐶𝑙2

Effects of Oxidation in Real Life

✽ Corrosion
corrosion - when a metal is attacked by substances around it (moisture, acids) & corrodes
○ rusting of iron - coating of iron articles with rust (red-brown powder)
○ tarnishing of silver - black coating on silver articles
○ oxidation of copper - green coating on copper articles

It causes damage to car bodies, bridges, iron railings, ships & all metal objects, especially iron. Huge amounts of
money is spent to replace damaged iron every year

✽ Rancidity
rancidity - when fats & oils are oxidised, they become rancid & their smell / taste changes

Substances that prevent oxidation (antioxidants) are added to foods containing fats & oil. Keeping food in air tight
containers helps to slow down oxidation

Chips manufacturers flush bags of chips with gas (nitrogen) to prevent the chips from getting oxidised