Class 10 Science Chapter 1 Notes: Chemical Reactions and Equations

Contents hide

1What Is a Chemical Reaction?

2Chemical Equations

2.1 Why an Equation Must Be Balanced

2.2 Balancing an Equation: the Hit-and-Trial Method

2.3 Step VII: Physical States and Reaction Conditions

3Types of Chemical Reactions

3.1 Combination Reaction

3.2 Decomposition Reaction

3.3 Displacement Reaction

3.4 Double Displacement Reaction

3.5 Oxidation and Reduction

4Effects of Oxidation in Everyday Life

4.1 Corrosion

4.2 Rancidity

What you will learn in this chapter

How to write a chemical equation and balance it using the hit-and-trial method
How to show physical states (s, l, g, aq) and reaction conditions in an equation
The types of chemical reaction (combination, decomposition, displacement, double displacement, and oxidation-reduction) and how to spot each from an equation
The difference between an exothermic and an endothermic reaction
Oxidation and reduction, in terms of gain or loss of oxygen and hydrogen
How to recognise corrosion and rancidity in everyday life, and how rancidity is slowed down

chemical equationbalanced equationcombination reaction
decomposition reactiondisplacement reactiondouble displacement reaction
precipitationoxidationreductionexothermicendothermic
corrosionrancidity

1What Is a Chemical Reaction?

Milk turns sour in summer. An iron tawa (griddle) rusts. Grapes ferment, food cooks, food is digested inside the body, and we breathe. In every one of these, the starting substance changes into something with a different nature and identity. Whenever this kind of chemical change happens, we say a chemical reaction has taken place.

Activity 1.1

Clean a magnesium ribbon with sandpaper, burn it over a spirit lamp, and collect the ash formed in a watch-glass. The magnesium combines with oxygen from the air to form magnesium oxide (a white ash): this is itself an example of a chemical reaction.

Activities 1.2 and 1.3

Adding potassium iodide solution to lead nitrate solution gives a yellow solid (Activity 1.2). Adding dilute hydrochloric or sulphuric acid to zinc granules gives bubbles of hydrogen gas around the granules, and the flask becomes warm (Activity 1.3).

Figure 1 Zinc granules with dilute sulphuric acid in a conical flask give bubbles of hydrogen gas, and the flask becomes warm (Activity 1.3, book Figure 1.2) — **Figure 1** Zinc granules with dilute sulphuric acid in a conical flask give bubbles of hydrogen gas, and the flask becomes warm (Activity 1.3, book Figure 1.2)

How to tell that a chemical reaction has taken place

From these three activities, any of the following observations tells us that a chemical reaction has taken place:

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

2Chemical Equations

Writing out a reaction in a full sentence gets long, so it is written in short as a word-equation instead:

Word-equation

Magnesium + Oxygen → Magnesium oxide

Substances on the left (LHS) are called the reactants; the new substance formed on the right (RHS) is called the product. The arrow points towards the product.

Learn by heartDefinition 1

A chemical equation is the short way of representing a chemical reaction using chemical formulae instead of words: for example $\ce{Mg + O2 -> MgO}$ .

When the number of atoms on the two sides of an equation is not equal, the equation is called a skeletal chemical equation. The equation above, $\ce{Mg + O2 -> MgO}$ , is a skeletal equation as written: it has 2 oxygen atoms on the left but only 1 on the right.

2.1 Why an Equation Must Be Balanced

Learn by heartDefinition 2

The law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction. So the total number of atoms of each element present in the reactants must equal the number in the products. An equation is called balanced only when this is true.

Exam Tip

“Why should chemical equations be balanced?” is Exercise question 4 of this chapter. Always name the law of conservation of mass in the answer: just saying “atoms must be equal” does not give the reason.

Check the reaction of Activity 1.3, $\ce{Zn + H2SO4 -> ZnSO4 + H2}$ : Zn 1 = 1, H 2 = 2, S 1 = 1, O 4 = 4. Every element has the same number of atoms on both sides, so this equation is already balanced.

2.2 Balancing an Equation: the Hit-and-Trial Method

Example: balance $\ce{Fe + H2O -> Fe3O4 + H2}$ .

Step IDraw an imaginary box around each formula: nothing inside a box may be changed while balancing

↓

Step IICount the atoms of every element on both sides of the unbalanced equation: Fe has 1 atom on the left and 3 on the right, H has 2 and 2, O has 1 and 4

↓

Step IIIStart with the compound that has the most atoms (

\ce{Fe3O4}

): balance its oxygen first, giving

4\,\ce{H2O}

↓

Step IVNow balance hydrogen: put

4\,\ce{H2}

on the right

↓

Step VFinally balance the one element left, iron: put

3\,\ce{Fe}

on the left

↓

Step VI

\ce{3Fe + 4H2O -> Fe3O4 + 4H2}

: every element now has equal atoms on both sides: verified and balanced

Common Mistake

Never change a compound’s formula to balance atoms. To balance oxygen you write a coefficient “4” in front of $\ce{H2O}$ , giving $4\,\ce{H2O}$ . You never change the formula itself to $\ce{H2O4}$ or $\ce{(H2O)4}$ .

Try it yourself: the book asks you to balance the skeletal equation $\ce{Mg + O2 -> MgO}$ by the same steps. Oxygen first: there are 2 O atoms on the left, so write $2\,\ce{MgO}$ on the right. Now there are 2 Mg atoms on the right, so write $2\,\ce{Mg}$ on the left: $\ce{2Mg + O2 -> 2MgO}$ , balanced.

2.3 Step VII: Physical States and Reaction Conditions

To make an equation more informative, the physical state of each substance is written in brackets after it:

Symbol	State
(s)	solid
(l)	liquid
(g)	gas
(aq)	aqueous (dissolved in water)

Fully balanced equation (with states)

\ce{3Fe(s) + 4H2O(g) -> Fe3O4(s) + 4H2(g)}

Here

\ce{H2O(g)}

tells us the water is used as steam. Reaction conditions (temperature, pressure, catalyst) are written above or below the arrow: for example

\ce{CO(g) + 2H2(g) ->[340\ atm] CH3OH(l)}

, and photosynthesis:

6\,\ce{CO2(aq)} + 12\,\ce{H2O(l)} \xrightarrow{\text{sunlight, chlorophyll}} \ce{C6H12O6(aq)} + 6\,\ce{O2(aq)} + 6\,\ce{H2O(l)}

. (A catalyst is a substance that speeds up a reaction without being used up itself, which is why it is written on the arrow rather than as a reactant.)

3Types of Chemical Reactions

During any chemical reaction, atoms of one element never turn into atoms of another element: bonds between atoms simply break and re-form to give new substances. The book studies the types of reaction one by one: combination, decomposition, displacement, double displacement, and oxidation-reduction.

3.1 Combination Reaction

Activity 1.4

Adding water slowly to calcium oxide (quicklime) makes the beaker warm, and slaked lime (calcium hydroxide) is formed, with heat also released.

Learn by heartDefinition 3

A combination reaction is one in which two or more substances (elements or compounds) combine to form a single product.

\ce{CaO(s) + H2O(l) -> Ca(OH)2(aq)}

+ heat

\ce{C(s) + O2(g) -> CO2(g)}

2\,\ce{H2(g) + O2(g) -> 2H2O(l)}

The burning of magnesium in Activity 1.1 ( $\ce{2Mg + O2 -> 2MgO}$ ) is also a combination reaction: a single product is formed, and heat is given out along with it.

Do You Know?

Whitewashing walls uses a solution of slaked lime. It slowly reacts with carbon dioxide in the air to form a thin layer of calcium carbonate $\ce{CaCO3}$ : $\ce{Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)}$ .
It is this layer that makes the wall shine after 2-3 days. Marble has the same chemical formula, $\ce{CaCO3}$ .

Learn by heartDefinition 4

An exothermic reaction is one in which heat is given out along with the product. Examples: burning natural gas ( $\ce{CH4(g) + 2O2(g) -> CO2(g) + 2H2O(g)}$ ), respiration
( $\ce{C6H12O6(aq) + 6O2(aq) -> 6CO2(aq) + 6H2O(l)}$ + energy), and vegetable matter turning into compost.

3.2 Decomposition Reaction

Learn by heartDefinition 5

A decomposition reaction is the exact opposite of a combination reaction: a single reactant breaks down into two or more simpler products.

Thermal decomposition (Activity 1.5)

Decomposition caused by heating is called thermal decomposition. Heating green ferrous sulphate crystals changes their colour and gives off the characteristic odour of burning sulphur:
$2\,\ce{FeSO4(s) ->[\Delta] Fe2O3(s) + SO2(g) + SO3(g)}$ .
Ferrous sulphate crystals ( $\ce{FeSO4.7H2O}$ ) first lose water when heated, and their colour changes. Then the compound decomposes to ferric oxide ( $\ce{Fe2O3}$ ), sulphur dioxide ( $\ce{SO2}$ ) and sulphur trioxide ( $\ce{SO3}$ ). Ferric oxide is a solid, while $\ce{SO2}$ and $\ce{SO3}$ are gases.

Thermal decomposition of limestone

$\ce{CaCO3(s) ->[\Delta] CaO(s) + CO2(g)}$ : an important decomposition reaction used in various industries. Calcium oxide is called lime or quick lime; one of its many uses is in the manufacture of cement.

Activity 1.6

Heating lead nitrate produces brown fumes of nitrogen dioxide:
$2\,\ce{Pb(NO3)2(s) ->[\Delta] 2PbO(s) + 4NO2(g) + O2(g)}$

Electrolysis of water (Activity 1.7)

Passing electric current through acidified water forms hydrogen and oxygen gas: the volume of hydrogen collected is exactly double that of oxygen ( $\ce{H2 : O2 = 2 : 1}$ ).

Photolytic decomposition (Activity 1.8)

Decomposition caused by light is called photolytic decomposition. White silver chloride turns grey on standing in sunlight: $2\,\ce{AgCl(s) ->[\text{sunlight}] 2Ag(s) + Cl2(g)}$ .
Silver bromide behaves the same way: $2\,\ce{AgBr(s) ->[\text{sunlight}] 2Ag(s) + Br2(g)}$ .
This is the principle used in black-and-white photography.

Figure 2 (a) Heating ferrous sulphate crystals in a dry boiling tube, with its mouth pointed away from everyone, and wafting the gas gently towards the nose (Activity 1.5, book Figure 1.4). (b) Heating lead nitrate held with a pair of tongs gives brown fumes of nitrogen dioxide (Activity 1.6, book Figure 1.5) — **Figure 2** (a) Heating ferrous sulphate crystals in a dry boiling tube, with its mouth pointed away from everyone, and wafting the gas gently towards the nose (Activity 1.5, book Figure 1.4). (b) Heating lead nitrate held with a pair of tongs gives brown fumes of nitrogen dioxide (Activity 1.6, book Figure 1.5)

Figure 3 Electrolysis of water (Activity 1.7, book Figure 1.6). Hydrogen collects in the test tube over the cathode (−) and oxygen over the anode (+); the volume of hydrogen is double that of oxygen — **Figure 3** Electrolysis of water (Activity 1.7, book Figure 1.6). Hydrogen collects in the test tube over the cathode (−) and oxygen over the anode (+); the volume of hydrogen is double that of oxygen

Every decomposition reaction above needs energy, in the form of heat, light or electricity, to break down the reactants.

Learn by heartDefinition 6

An endothermic reaction is a reaction in which energy is absorbed.

Carry out the following Activity

Take about 2 g barium hydroxide in a test tube, add 1 g of ammonium chloride and mix with a glass rod. Touch the bottom of the test tube with your palm: it feels cold. The reaction absorbs heat from its surroundings (including your hand), so it is an endothermic reaction.

3.3 Displacement Reaction

Activity 1.9

Dipping iron nails in a copper sulphate solution turns the nails brownish, and the solution’s blue colour fades: the iron has displaced the copper from its compound.

Learn by heartDefinition 7

A displacement reaction is one in which a more reactive element displaces a less reactive element from its compound.

\ce{Fe(s) + CuSO4(aq) -> FeSO4(aq) + Cu(s)}

\ce{Zn(s) + CuSO4(aq) -> ZnSO4(aq) + Cu(s)}

\ce{Pb(s) + CuCl2(aq) -> PbCl2(aq) + Cu(s)}

Zinc and lead are both more reactive than copper, so each one displaces copper from its compounds.

3.4 Double Displacement Reaction

Activity 1.10

Mixing solutions of sodium sulphate and barium chloride produces a white precipitate immediately.

\ce{Na2SO4(aq) + BaCl2(aq) -> BaSO4(s) + 2NaCl(aq)}

The white substance formed is insoluble in water; an insoluble substance formed in this way is known as a precipitate. The white precipitate of $\ce{BaSO4}$ is formed by the reaction of $\ce{SO4^2-}$ and $\ce{Ba^2+}$ ions; the other product, sodium chloride, remains in the solution. Any reaction that produces a precipitate can be called a precipitation reaction.

Learn by heartDefinition 8

A double displacement reaction is a reaction in which there is an exchange of ions between the reactants.

Activity 1.2 (mixing lead(II) nitrate and potassium iodide solutions) is also a double displacement reaction: a yellow precipitate of lead iodide forms, $\ce{Pb(NO3)2(aq) + 2KI(aq) -> PbI2(s) + 2KNO3(aq)}$ .

3.5 Oxidation and Reduction

Activity 1.11

Heating copper powder in air deposits a black coating (copper(II) oxide) on it: the copper has gained oxygen. Passing hydrogen gas over this same black powder turns it back to reddish-brown copper: here the copper oxide has lost oxygen.

Oxidation

Gain of oxygen
or loss of hydrogen
$2\,\ce{Cu + O2 -> 2CuO}$ : the copper is oxidised

Reduction

Loss of oxygen
or gain of hydrogen
$\ce{CuO + H2 -> Cu + H2O}$ : the copper oxide is reduced

Learn by heartDefinition 9

Oxidation = gain of oxygen or loss of hydrogen. Reduction = loss of oxygen or gain of hydrogen.

In $\ce{CuO + H2 -> Cu + H2O}$ , copper(II) oxide loses oxygen and is reduced, while hydrogen gains oxygen and is oxidised. So in this reaction one reactant gets oxidised while the other gets reduced. Such reactions are called oxidation-reduction or redox reactions. Other examples:

\ce{ZnO + C -> Zn + CO}

(carbon is oxidised to CO, ZnO is reduced to Zn)

\ce{MnO2 + 4HCl -> MnCl2 + 2H2O + Cl2}

(HCl is oxidised to

\ce{Cl2}

\ce{MnO2}

is reduced to

\ce{MnCl2}

)

In Activity 1.1, magnesium gains oxygen to form magnesium oxide, so magnesium is oxidised.

4Effects of Oxidation in Everyday Life

4.1 Corrosion

Learn by heartDefinition 10

When a metal is attacked by substances around it, such as moisture and acids, it is said to corrode, and this process is called corrosion. Examples: the reddish-brown coating on iron (rusting), the black coating on silver and the green coating on copper.

Figure 6 Rust (iron oxide) forming on iron is the most familiar everyday example of corrosion. Image: Laitr Keiows / Wikimedia Commons / CC BY-SA 3.0 — **Figure 6** Rust (iron oxide) forming on iron is the most familiar everyday example of corrosion. Image: Laitr Keiows / Wikimedia Commons / CC BY-SA 3.0

Corrosion causes damage to car bodies, bridges, iron railings, ships and all objects made of metals, especially those of iron; every year an enormous amount of money is spent to replace damaged iron. Painting iron articles keeps moisture and air away from the iron surface, which prevents rusting. (Corrosion is studied in more detail in Chapter 3.)

4.2 Rancidity

Learn by heartDefinition 11

Rancidity is the process by which fats and oils in food are oxidised by oxygen in the air, changing their smell and taste (“going stale” or “off”).

Adding antioxidants

Substances that slow down oxidation are added to fat- and oil-containing foods.

Airtight containers

Keeping food away from contact with air (oxygen) slows oxidation down.

Nitrogen flushing

Packets of oily/fatty food, such as chips, are flushed with nitrogen gas to push oxygen out.

Every chemical reaction either joins something together (combination), breaks something apart (decomposition), or swaps something’s place (displacement/double displacement). And in almost every one, energy is either given out or taken in.

This chapter in one line

Check yourself before the exam

Can I state every step of the hit-and-trial method for balancing an equation, without looking?
Can I identify each type of reaction (combination, decomposition, displacement, double displacement, redox) from one example equation?
Can I define oxidation and reduction without getting them confused?
Can I explain why iron articles are painted, and name the three ways rancidity is slowed down?

All definitions in one place

Chemical equationThe short way of representing a chemical reaction using chemical formulae

Law of conservation of massMass can neither be created nor destroyed in a reaction

Combination reactionTwo or more substances combine to form a single product

Exothermic reactionHeat is given out along with the product

Decomposition reactionOne reactant breaks down into two or more simpler products

Endothermic reactionEnergy is absorbed

Displacement reactionA more reactive element displaces a less reactive one from its compound

Double displacement reactionIons are exchanged between the reactants

Precipitation reactionAny reaction that produces a precipitate (an insoluble substance)

Redox reactionOne reactant is oxidised while the other is reduced

OxidationGain of oxygen or loss of hydrogen

ReductionLoss of oxygen or gain of hydrogen

CorrosionA metal is attacked by substances around it, such as moisture and acids

RancidityFat/oil-containing food goes stale from oxidation

Quick Revision: read this the night before the exam

Balanced equation = equal atoms of every element on both sides (law of conservation of mass)
Types: combination (A+B→AB), decomposition (AB→A+B), displacement (A+BC→AC+B), double displacement (exchange of ions), redox (one reactant oxidised, the other reduced)
Oxidation = gain of O/loss of H; reduction = loss of O/gain of H
Exothermic = energy given out; endothermic = energy absorbed
Corrosion (rust, blackening) and rancidity (stale food) are both everyday examples of oxidation

More for this chapter

Important Questions MCQ / Objective NCERT Solutions Short Notes