What is a chemical reaction?

The transformation of chemical substance into another chemical substance is known as Chemical Reaction. For example: Rusting of iron, the setting of milk into curd, digestion of food, respiration, etc.

In a chemical reaction, a new substance is formed which is completely different in properties from the original substance, so in a chemical reaction, a chemical change takes place.

Only a rearrangement of atoms takes place in a chemical reaction.

The substances which take part in a chemical reaction are called reactants.
The new substances produced as a result of a chemical reaction are called products.

Example: The burning of magnesium in the air to form magnesium oxide is an example of a chemical reaction.
2Mg(s) + $O_2$ (g) $△→$ 2MgO(s)

Before burning in air, the magnesium ribbon is cleaned by rubbing with sandpaper.
This is done to remove the protective layer of basic magnesium carbonate from the surface of the magnesium ribbon.

Reactant: Substances which take part in a chemical reaction are called reactants.
Example: Mg and $O_2$ .

Product: New substance formed after a chemical reaction is called a product.
Example: MgO.

Characteristics of Chemical Reactions :

(i) Evolution of gas: The chemical reaction between zinc and dilute sulphuric acid is characterised by the evolution of hydrogen gas.
Zn(s) + $H_2SO_4$ (aq) → $ZnSO_4$ (aq) + $H_2$ (g) ↑

(ii) Change in Colour: The chemical reaction between citric acid and purple coloured potassium permanganate solution is characterised by a change in colour from purple to colourless.

The chemical reaction between sulphur dioxide gas and acidified potassium dichromate solution is characterized by a change in colour from orange to green.

(iii) Change in state of substance: The combustion reaction of candle wax is characterised by a change in state from solid to liquid and gas (because the wax is a solid, water formed by the combustion of wax is a liquid at room temperature whereas, carbon dioxide produced by the combustion of wax is a gas). There are some chemical reactions which can show more than one characteristics.

(iv) Change in temperature: The chemical reaction between quick lime water to form slaked lime is characterized by a change in temperature (which is a rise in temperature).
The chemical reaction between zinc granules and dilute sulphuric acid is also characterised by a change in temperature (which is a rise in temperature).

(v) Formation of precipitate: The chemical reaction between sulphuric acid and barium chloride solution is characterised by the formation of a white precipitate of barium sulphate.
$BaCl_2(aq) + H_2SO_4(aq) → BaSO_4(s) (ppt) + 2HCl(aq)$

3. Types of Chemical Reactions

Type	Definition	General Form	Example
Combination	Two or more reactants form one product.	$A + B \rightarrow AB$	$CaO + H_2O \rightarrow Ca(OH)_2$ (Quicklime to Slaked lime)
Decomposition	A single reactant breaks down into simpler products.	$AB \rightarrow A + B$	$2FeSO_4 \xrightarrow{\Delta} Fe_2O_3 + SO_2 + SO_3$
Displacement	A more reactive element displaces a less reactive one.	$A + BC \rightarrow AC + B$	$Fe + CuSO_4 \rightarrow FeSO_4 + Cu$
Double Displacement	Exchange of ions between two reactants.	$AB + CD \rightarrow AD + CB$	$Na_2SO_4 + BaCl_2 \rightarrow BaSO_4 + 2NaCl$

Reactions involved

Reaction Equation	Type of Reaction	Explanation & Key Properties	Physical Changes Observed
$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$	Combination (Oxidation)	Magnesium ribbon burns in oxygen. $MgO$ is a basic oxide.	Dazzling white flame; white powder ( $MgO$ ) formed.
$CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + \text{Heat}$	Combination (Exothermic)	Quicklime ( $CaO$ ) reacts vigorously with water to form Slaked Lime.	Hissing sound; container becomes very hot.
$2FeSO_4(s) \xrightarrow{\Delta} Fe_2O_3(s) + SO_2(g) + SO_3(g)$	Thermal Decomposition	Green Ferrous Sulphate crystals lose water and decompose.	Color change: Green to Brown; Smell of burning sulfur.
$CaCO_3(s) \xrightarrow{\Delta} CaO(s) + CO_2(g)$	Thermal Decomposition	Limestone is heated to produce Quicklime for cement industry.	Evolution of colorless, odorless gas ( $CO_2$ ).
$2Pb(NO_3)_2(s) \xrightarrow{\Delta} 2PbO(s) + 4NO_2(g) + O_2(g)$	Thermal Decomposition	Lead nitrate powder is heated in a boiling tube.	Brown fumes of $NO_2$ ; yellow residue of $PbO$ left.
$2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_2(g)$	Photolytic Decomposition	Silver chloride is sensitive to light (used in black & white photography).	White powder turns Grey due to silver metal.
$Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$	Displacement	Iron is more reactive than Copper and pushes it out of solution.	Blue solution turns Light Green; reddish-brown coating on iron.
$Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$	Displacement	Granulated zinc reacts with dilute acid.	Evolution of gas; Pop sound when tested with a match.
$Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) \downarrow + 2NaCl(aq)$	Double Displacement	Exchange of ions between Sodium Sulphate and Barium Chloride.	Formation of a White Precipitate ( $BaSO_4$ ).
$2Cu + O_2 \xrightarrow{\Delta} 2CuO$	Oxidation	Copper powder is heated in a china dish.	Reddish-brown surface turns Black (Copper II Oxide).
$CuO + H_2 \xrightarrow{\Delta} Cu + H_2O$	Redox	Hydrogen gas is passed over heated Black Copper Oxide.	Black surface turns Reddish-brown again.

Activities

Activity	Reaction & Equation	Reaction Type	Observation & Physical Changes
1.1	Burning Magnesium Ribbon $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$	Combination / Oxidation	Reactant: Silvery metal. Product: Dazzling white flame; transforms into white powder.
1.2	Lead Nitrate + Potassium Iodide $Pb(NO_3)_2 + 2KI \rightarrow PbI_2 + 2KNO_3$	Double Displacement	Reactant: Two colorless solutions. Product: Formation of a Yellow Precipitate ( $PbI_2$ ).
1.3	Zinc + Dilute Acid $Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2 \uparrow$	Displacement	Reactant: Grey zinc granules. Product: Gas bubbles; conical flask becomes hot (Exothermic).
1.4	Quicklime + Water $CaO + H_2O \rightarrow Ca(OH)_2 + \text{Heat}$	Combination	Reactant: White solid chunks. Product: Vigorous reaction; hissing sound; huge temperature rise.
1.5	Heating Ferrous Sulphate $2FeSO_4 \xrightarrow{\Delta} Fe_2O_3 + SO_2 + SO_3$	Thermal Decomposition	Reactant: Green crystals ( $FeSO_4 \cdot 7H_2O$ ). Product: Turns Brown/Red; smell of burning sulfur.
1.6	Heating Lead Nitrate $2Pb(NO_3)_2 \xrightarrow{\Delta} 2PbO + 4NO_2 + O_2$	Thermal Decomposition	Reactant: White powder. Product: Brown fumes ( $NO_2$ ); Yellow residue ( $PbO$ ) left behind.
1.7	Electrolysis of Water $2H_2O(l) \xrightarrow{\text{Elec.}} 2H_2(g) + O_2(g)$	Electrolytic Decomposition	Product: Volume of $H_2$ collected is double the volume of $O_2$ .
1.8	Silver Chloride in Sunlight $2AgCl(s) \xrightarrow{h\nu} 2Ag(s) + Cl_2(g)$	Photolytic Decomposition	Reactant: White solid. Product: Turns Grey as silver metal forms; Chlorine gas escapes.
1.9	Iron Nails in Copper Sulphate $Fe + CuSO_4 \rightarrow FeSO_4 + Cu$	Displacement	Reactant: Blue solution ( $CuSO_4$ ). Product: Solution turns Pale Green; nail becomes reddish-brown.
1.10	Sodium Sulphate + Barium Chloride $Na_2SO_4 + BaCl_2 \rightarrow BaSO_4 \downarrow + 2NaCl$	Double Displacement	Reactant: Colorless solutions. Product: Instant White Precipitate ( $BaSO_4$ ) formed.
1.11	Heating Copper Powder $2Cu + O_2 \xrightarrow{\Delta} 2CuO$	Oxidation	Reactant: Brown copper powder. Product: Surface becomes coated with Black Copper (II) Oxide.

Writing a Chemical equation

Step	Action	Why?	Example (Formation of Water)
1	Identify Reactants & Products	To know the starting materials and the result.	Reactants: Hydrogen, Oxygen. Product: Water.
2	Write the Skeleton Equation	Use chemical formulas instead of names for brevity.	$H_2 + O_2 \rightarrow H_2O$
3	Balance the Equation	To satisfy the Law of Conservation of Mass.	$2H_2 + O_2 \rightarrow 2H_2O$
4	Indicate Physical States	To provide clarity on the form of the substances.	$2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$
5	Add Reaction Conditions	Mention heat, pressure, or catalysts above/below the arrow ( $\rightarrow$ ).	$2H_2(g) + O_2(g) \xrightarrow{\text{Electric Spark}} 2H_2O(l)$

Corrosion and Rancidity

Corrosion and rancidity are two common phenomena observed in daily life that are caused by the process of oxidation.

1. Corrosion

Corrosion is the process in which metals are gradually eaten up by the action of air, moisture, or a chemical (such as an acid) on their surface. It is essentially the oxidation of metals.

Rusting of Iron: This is the most common form of corrosion. When iron is exposed to moist air for a long time, it acquires a coating of a brown flaky substance called rust.
Black Coating on Silver: Silver articles become black after some time when exposed to air because they react with sulfur in the air to form a coating of silver sulfide.
Green Coating on Copper: Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface, gaining a green coat of basic copper carbonate.
Prevention: Corrosion can be prevented by methods such as painting iron articles, oiling, greasing, galvanizing, or making alloys.

2. Rancidity

Rancidity refers to the condition produced by the aerial oxidation of fats and oils in foods, which is marked by an unpleasant smell and taste.

Mechanism: When fats and oils present in food are oxidized, they become rancid. Their smell and taste change, making the food unfit for consumption.
Prevention Methods:
- Antioxidants: Adding substances which prevent oxidation (antioxidants) to foods containing fats and oils.
- Nitrogen Flushing: Keeping food in an atmosphere of unreactive gas like nitrogen to prevent the chips or snacks from getting oxidized.
- Airtight Containers: Keeping food in airtight containers helps to slow down the process of oxidation.
- Refrigeration: Storing food at low temperatures also helps slow down rancidity.

Valencies

Atoms

Atom	Symbol	Valency	Notes
Hydrogen	H	1	Forms H⁺ or shares 1 electron
Oxygen	O	2	Needs 2 electrons to complete octet
Nitrogen	N	3	Often forms NH₃, can vary (3 or 5)
Carbon	C	4	Forms 4 covalent bonds (CH₄, CO₂)
Sodium	Na	1	Loses 1 electron → Na⁺
Potassium	K	1	Loses 1 electron → K⁺
Calcium	Ca	2	Loses 2 electrons → Ca²⁺
Magnesium	Mg	2	Loses 2 electrons → Mg²⁺
Zinc	Zn	2	Loses 2 electrons → Zn²⁺
Iron	Fe	2 or 3	Can form Fe²⁺ or Fe³⁺
Copper	Cu	1 or 2	Can form Cu⁺ or Cu²⁺
Silver	Ag	1	Usually forms Ag⁺
Lead	Pb	2 or 4	Can form Pb²⁺ or Pb⁴⁺
Sulphur	S	2, 4, or 6	Commonly 2 (H₂S), 6 (SO₄²⁻)
Chlorine	Cl	1	Gains 1 electron → Cl⁻
Iodine	I	1	Gains 1 electron → I⁻
Barium	Ba	2	Loses 2 electrons → Ba²⁺

Radicals

Radical	Formula	Valency	Notes
Hydroxide	OH⁻	1	Present in bases like NaOH, Ca(OH)₂
Nitrate	NO₃⁻	1	Found in salts like KNO₃, AgNO₃
Sulphate	SO₄²⁻	2	Found in salts like CuSO₄, ZnSO₄
Carbonate	CO₃²⁻	2	Present in CaCO₃, Na₂CO₃
Bicarbonate	HCO₃⁻	1	Found in NaHCO₃ (baking soda)
Chlorate	ClO₃⁻	1	Less common, but appears in some salts
Phosphate	PO₄³⁻	3	Found in fertilizers like Ca₃(PO₄)₂
Ammonium	NH₄⁺	1	Found in NH₄Cl, NH₄NO₃
Sulphite	SO₃²⁻	2	Found in Na₂SO₃
Nitrite	NO₂⁻	1	Found in NaNO₂
Permanganate	MnO₄⁻	1	Found in KMnO₄ (oxidizing agent)
Dichromate	Cr₂O₇²⁻	2	Found in K₂Cr₂O₇ (oxidizing agent)

Cross Multiplication

Elements	Valencies	Cross Multiplication	Formula
Na (1), O (2)	Na (1), O (2)	Na₂O	Sodium oxide
Al (3), Cl (1)	Al (3), Cl (1)	AlCl₃	Aluminium chloride
Ca (2), OH (1)	Ca (2), OH (1)	Ca(OH)₂	Calcium hydroxide
Mg (2), N (3)	Mg (2), N (3)	Mg₃N₂	Magnesium nitride

Valence of all the elements in the Periodic Table | Dynamic Interactive Periodic Table

1. Chemical Reactions and Equations