NCERT Class 8 Chemistry Chapter 6 Chemical Reactions Notes

Chapter 6 Chemical Reactions

We have discussed a large number of chemical reactions. Some are accompanied by a change in colour, some by the evolution of a gas, some by the formation of a precipitate, and so on. At the molecular level, different types of changes occur. In some reactions, smaller parts add up to form bigger entities, whereas in some others, bigger entities break into smaller ones.

Again, in some reactions, one element displaces another from a compound whereas in some others, compounds exchange radicals among themselves. For a systematic study, therefore, it is essential that we classify the reactions into different types.

NCERT Solutions for Chemical Reactions Class 8

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However, we will discuss first the change in energy in chemical reactions, which is common to all types

Change In Energy In A Reaction

All reactions are accompanied by a change in energy. Energy (in the form of heat or light) is either given out or taken in during a chemical reaction.

• A chemical reaction in which heat or light is evolved is called an exothermic reaction.
• A chemical reaction in which heat or light is absorbed is called an endothermic reaction.
• In an exothermic process, heat or light is given out to the surroundings. That is why you feel hot when you accidentally touch anything burning.
• On the other hand, in an endothermic process, heat or light is taken in, i.e., absorbed from the surroundings. That is why a cube of ice is cold to the touch.
• The ice tends to melt when you touch it. The melting of ice is an endothermic process, and the heat required for it is drawn from your hand.
• And so the cold feeling. Remember that the melting of ice is a physical change, not a chemical reaction

Why A Change In Energy In A Reaction

As we know, any two atoms in a molecule are held together by a force of attraction called a chemical bond. And also that a chemical bond is much stronger than an intermolecular force. We also know that it is the atoms that take part in a chemical reaction.

Thus, we can consider a chemical reaction to be a result of the following phenomena.

• The breaking of bonds of the reactant molecules to set the atoms free for a new combination. (Energy is absorbed in the process.)
• The formation of fresh bonds between the new partners so in the product molecules. (Energy is given out in the process.)

The energy required for the breaking of the bonds of the reactant molecules is not the same as that given out in the formation of the bonds between the new partners.

• When the energy given out in the formation of the bonds is more than that used for the breaking of the bonds, the surplus energy is evolved and the process is exothermic.

However, when the energy required for the breaking of the bonds is greater than that given out in the formation of the bonds, the difference in energy has to be compensated for from outside.

In other words, energy will be absorbed from the surroundings for the reaction to happen, and the process will be endothermic.

NCERT Class 8 Chemistry Chapter 6 Chemical Reactions

Examples of Exothermic Reactions

We will now discuss some common exothermic reactions

1. Burning:

The burning of a substance is exothermic. It is the exothermicity of burning that makes a combustible substance a fuel. For example, hydrogen, carbon (coal), CNG(CH₄), and LPG(C₄H₁₀) are fuels.

2. How some active metals react with water:

The reaction of a highly active metal like K, Na or Ca is vigorous and exothermic. Hydrogen is liberated, and the metal hydroxide is formed in the reaction.

2K(s) (potassium) + 2H₂O(l) → 2 KOH(aq)(Potassium hydroxide) + H₂(g)

2Na(s) + 2H₂O(l) → 2NaOH(aq)(sodiumhydroxide) + H₂(g)

Ca(s) + 2H₂O(l) → Ca (OH)₂(calcium hydroxide)(aq) + H₂(g)

3. The dissolution of metal oxides (bases) in water:

Bases like sodium oxide and calcium oxide vigorously react with water to form their hydroxides, releasing a large amount of heat.

The hydroxides dissolve in an excess of water. The reaction of calcium oxide (CaO, commercially known as quicklime) with water, giving calcium hydroxide [Ca(OH)₂, commercially known as slaked lime] is called the slaking of lime.

4. The dissolution of nonmetal oxides (acidic) in water:

The oxides of nonmetals are acidic and many of them dissolve in water to form acids exothermically.

⇒ \(\begin{aligned}
&\mathrm{CO}_2(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq}){ carbonic acid }\\\end{aligned}\)

⇒ \(\begin{aligned}
&\mathrm{SO}_2(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \mathrm{H}_2 \mathrm{SO}_3(\mathrm{aq})sulphurous acid \\\end{aligned}\)

⇒ \(\underset{\text { sulphur trioxide }}{\mathrm{SO}_3(\mathrm{l})}+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \underset{\text { sulphuric acid }}{\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq})}\)

5. The reaction between a basic and an acidic oxide:

A basic oxide reacts with an acidic oxide, forming a salt and evolving heat

⇒ \(\begin{aligned}
&\mathrm{Na}_2 \mathrm{O}(\mathrm{~s})+\mathrm{CO}_2(\mathrm{~g}) \longrightarrow \mathrm{Na}_2 \mathrm{CO}_3(\mathrm{~s})(sodium carbonate)\\
&\text { }
\end{aligned}\)

⇒ \(\mathrm{CaO}(\mathrm{~s})+\mathrm{CO}_2(\mathrm{~g}) \longrightarrow \underset{\substack{\text { calcium carbonate } \\ \text { (limestone) }}}{\mathrm{CaCO}_3(\mathrm{~s})}\)

6. Neutralisation reactions:

In a neutralisation reaction, an acid reacts with a base, forming a salt and water. All such reactions are exothermic

⇒ \(\text { Base }+ \text { acid } \longrightarrow \text { salt + water }\)

⇒ \(\mathrm{NaOH}(\mathrm{aq})+\mathrm{HCl}(\mathrm{aq}) \longrightarrow \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

⇒ \(\mathrm{KOH}(\mathrm{aq})+\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq}) \longrightarrow \mathrm{K}_2 \mathrm{SO}_4(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

Chemical Reactions Class 8 NCERT Notes

7. Reaction of an acid with a metal carbonate:

By performing a similar activity as above, you can conclude that heat is evolved when an acid acts on a carbonate or a hydrogencarbonate, producing CO₂

⇒ \(\mathrm{Na}_2 \mathrm{CO}_3(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq})\) → \(2 \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})+\mathrm{CO}_2(\mathrm{~g})\)

⇒ \(\mathrm{CaCO}_3(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \longrightarrow\)\(\mathrm{CaCl}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})+\mathrm{CO}_2(\mathrm{~g})\)

8. Respiration:

You know that, in respiration, the glucose formed in plants and animals combines with the oxygen of the air to give C02 and H20. The reaction, being exothermic, provides living beings with the energy to sustain their life processes.

⇒ \(\begin{aligned}
&\mathrm{C}_6 \mathrm{H}_{12} \mathrm{O}_6(\mathrm{aq})+6 \mathrm{O}_2(\mathrm{~g}) \longrightarrow 6 \mathrm{CO}_2(\mathrm{~g})+6 \mathrm{H}_2 \mathrm{O}(\mathrm{l})\\
&\text { glucose }
\end{aligned}\)

9. Rusting

In rusting, iron slowly combines with oxygen (air) in the presence of moisture to form the brownish red hydrated iron(III) oxide, called rust

⇒ \(4 \mathrm{Fe}(\mathrm{~s})+3 \mathrm{O}_2(\mathrm{~g})+2 x \mathrm{H}_2 \mathrm{O}(\mathrm{~g}) \longrightarrow\) \(\underset{\text { rust }}{2\left[\mathrm{Fe}_2 \mathrm{O}_3 \cdot x \mathrm{H}_2 \mathrm{O}\right](\mathrm{s})}\)

The reaction is exothermic, but too slow to let you feel its heat at any point in time.

Examples of Endothermic Reactions

Endothermic reactions are much fewer in number than exothermic ones. Some examples are given below

1. Combination of nitrogen with oxygen:

Nitrogen combines with oxygenforming nitric oxide (NO) when an electric spark is passed through a mixture of the two gases (as during lightning). The reaction is endothermic.

⇒ \(\mathrm{N}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow[\text { spark }]{\text { electric }} 2 \mathrm{NO}(\mathrm{~g})\)

2. Thermal decomposition of metal carbonates:

The breaking down of a substance into simpler substances on being heated is called thermal decomposition. The carbonates of many metals decompose on being heated to give the metal oxides and carbon dioxide. These are endothermic reactions. Remember that, among the common carbonates, sodium carbonate (Na2C03) and potassium carbonate (K2C03) do not undergo such a reaction

⇒ \(\mathrm{CaCO}_3(\mathrm{~s}) \xrightarrow{\text { heat }} \mathrm{CaO}(\mathrm{~s})+\mathrm{CO}_2(\mathrm{~g})\)

⇒ \(\mathrm{MgCO}_3(\mathrm{~s}) \xrightarrow{\text { heat }} \mathrm{MgO}(\mathrm{~s})+\mathrm{CO}_2(\mathrm{~g})\)

3. Photosynthesis:

We know that a plant prepares its food, glucose, from atmospheric C02 and soilmoisture in the presence of chlorophyll by absorbing sunlight.

The reaction is endothermic

⇒ \(6 \mathrm{CO}_2(\mathrm{~g})+6 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \xrightarrow[\text { chlorophyll }]{\text { sunlight }}\) \(\mathrm{C}_6 \mathrm{H}_{12} \mathrm{O}_6(\mathrm{aq})+6 \mathrm{O}_2(\mathrm{~g})\)

Endothermic physical processes:

Many physical processes are endothermic. We have learnt that the melting of ice requires heat from outside and so does the boiling of water.

• Water gets converted to steam at 100 °C, but we have to heat it all along, i.e., until it is converted into steam.
• The dissolution of ammonium chloride (NH₄Cl), ammonium nitrate (NH₄NO₃) or glucose in water is also a common example of such a process. Stir any of these substances in a glass of water and hold the glass in your palm. You will feel that it has become cold

Some common observations:

Chemical reactions have the following general characteristics.

1. Exothermic reactions are more common than endothermic reactions:

This is because exothermic reactions do not need energy from the surroundings.

Types of Chemical Reactions Class 8

2. Once they begin, exothermic reactions continue on their own but endothermic reactions do not:

This is because some heat is required to keep a reaction going. And that heat is generated on its own during an exothermic reaction.

For example:

When you light a combustible substance, it burns on its own.

• When it starts rusting, a piece of iron continues to do so, though slowly, till all the iron is eaten up.
• However, this does not happen in an endothermic reaction because such a reaction does not generate any heat.
• So, the reaction goes on only as long as energy is supplied from outside.

For example:

Nitrogen combines with oxygen to form nitric oxide (NO) only as long as an electric spark is passed through a mixture of nitrogen and oxygen. The reaction stops in the absence of a spark

⇒ \(\mathrm{N}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow[\text { spark }]{\text { electric }} 2 \mathrm{NO}(\mathrm{~g})\)

Also, photosynthesis takes place in the presence of sunlight. Does it ever happen in the dark?

3. Most exothermic reactions too have to be initiated:

• A reaction, even exothermic, does not generally start simply when the reactants are mixed.
• Quite often, it has to be initiated by heating the reactants, igniting them, passing an electric spark through them, exposing them to UV light, etc.

For example:

A substance burns when ignited, hydrogen reacts with oxygen when ignited, and so on. You will learn the reason for this in higher classes.

Types Of Chemical Reactions

Among the different types of chemical reactions, combination, decomposition, displacement, double displacement and neutralization reactions are the ones we commonly come across. We will discuss them here.

Combination Reactions

In a combination reaction (synthesis), two or more reactants combine to form a product. Some examples are given below. The burning of some elements Elements like H, C, S, P, Ca and Mg, on being burnt in air, form their oxides

By direct combination:

1. \(2 \mathrm{H}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { burn }} 2 \mathrm{H}_2 \mathrm{O}(\mathrm{~g})\)

2.\(\mathrm{C}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g})  \xrightarrow{\text { burn }} \mathrm{CO}_2(\mathrm{~g}) \\\)

3. \(\mathrm{S}(\mathrm{~s}) +\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { burn }} \mathrm{SO}_2(\mathrm{~g}) \\\)

4. \(2 \mathrm{Ca}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { burn }} 2 \mathrm{CaO}(\mathrm{~s})\\\)

5. \(2 \mathrm{Mg}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { burn }} 2 \mathrm{MgO}(\mathrm{~s})\)

The burning of carbon:

Burn some charcoal (carbon) in a deflagrating spoon. Introduce the spoon into an open gas jar near the mouth of the jar so that the burning charcoal gets a clear supply of air.

• Close the mouth of the jar with a lid when all the carbon has burned.
• Invert the jar and place it in a trough containing clear limewater and open its mouth inside the liquid.
• The limewater rises in the jar and becomes milky. This proves that carbon dioxide is formed

The burning of magnesium:

When ignited in air, say in an open gas jar, it burns with a dazzling white flame. The magnesium oxide formed deposits as a white, smoky solid on the walls of the jar

Chemical Reactions and Equations NCERT Notes

Rusting

A piece of iron, when left in moist air, slowly gets a crust of a brown-red solid called rust. Rust —a hydrated iron(III) oxide product of the direct combination of iron with the oxygen and moisture of the air.

4Fe(s) + 3O₂(g) + Moisture → 2[Fe₂O₃. xH₂O](s)(rust brown- red)

(x can vary).

The combination of elements with chlorine

Chlorine is a very active nonmetal—a greenish-yellow gas. Most metals and some nonmetals combine with chlorine to form their chlorides

⇒\(2 \mathrm{Na}(\mathrm{~s})+\mathrm{Cl}_2(\mathrm{~g}) \longrightarrow \underset{\text { sodium chlorid }}{2 \mathrm{NaCl}(\mathrm{~s})}\)

⇒ \(2 \mathrm{Fe}(\mathrm{~s})+3 \mathrm{Cl}_2(\mathrm{~g}) \xrightarrow{\text { heat }} \underset{\text { iron }(\text { III }) \text { chloride }}{2 \mathrm{FeCl}_3(\mathrm{~s})}\)

⇒ \(\mathrm{H}_2(\mathrm{~g})+\mathrm{Cl}_2(\mathrm{~g}) \xrightarrow{\text { ignite }} \underset{\text { hydrogen chloride }}{2 \mathrm{HCl}(\mathrm{~g})}\)

⇒ \(\underset{\text { phosphorus }}{\mathrm{P}_4(\mathrm{~s})}+10 \mathrm{Cl}_2(\mathrm{~g}) \xrightarrow{\text { burn }} \underset{\begin{array}{c}
\text { phosphorus } \\
\text { pentachloride }
\end{array}}{4 \mathrm{PCl}_5(\mathrm{~s})}\)

The combination of iron with sulphur

You have already learnt that iron combines with sulphur, when heated, to form iron(II) sulphide.

⇒  \(\underset{\substack{\text { iron } \\ \text { (grey) }}}{\mathrm{Fe}}+\underset{\substack{\text { sulphur } \\ \text { (yellow) }}}{\mathrm{S}} \xrightarrow{\text { heat }} \underset{\substack{\text { iron(II) sulphide } \\ \text { (greyish black) }}}{\mathrm{FeS}}\)

The combination of oxides with water

The oxide of a nonmetal generally dissolves in water to form an acid.

⇒  \(\begin{aligned}
&\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{H}_2 \mathrm{CO}_3 (carbonic acid )\\
&\text { }
\end{aligned}\)

⇒  \(\begin{aligned}
&\mathrm{SO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{SO}_3( sulphurous acid)\\
&\text { }
\end{aligned}\)

Balancing Chemical Reactions Class 8

And the oxides of many metals combine with water to form the basic hydroxides.

⇒  \(\underset{\text { sodium oxide }}{\mathrm{Na}_2 \mathrm{O}}+\mathrm{H}_2 \mathrm{O} \longrightarrow \underset{\text { sodium hydroxide }}{2 \mathrm{NaOH}}\)

⇒  \(\underset{\text { calcium oxide }}{\mathrm{CaO}}+\mathrm{H}_2 \mathrm{O} \longrightarrow \underset{\text { calcium hydroxide }}{\mathrm{Ca}(\mathrm{OH})_2}\)

When carbon dioxide is passed through water for some time, the water turns blue litmus wine red, showing that it contains an acid.

• At the same time, limewater, which is prepared by adding water to quicklime (CaO), is a solution of Ca(OH)2 and turns red litmus blue. Thus, Ca(OH)2 is basic.
• Remember that acidic substances turn blue litmus red and basic substances turn red litmus blue

Decomposition Reactions

In a decomposition reaction, one substance breaks down into two or more simpler substances.

Some examples are given below:

Decomposition of water on electrolysis Electrolysis is a process in which a substance is decomposed, or broken down into simpler substances, by passing an electric current through it

Water mixed with a very small amount of an acid breaks down into hydrogen and oxygen on electrolysis

⇒  \(\underset{\text { water }}{\mathrm{H}_2 \mathrm{O}} \xrightarrow{\text { electrolysis }} \underset{\substack{\text { hydrogen } \\
\text { (negative } \\
\text { electrode) }}}{2 \mathrm{H}_2}+\underset{\begin{array}{c}
\text { oxygen } \\
\text { (positive } \\
\text { electrode) }
\end{array}}{\mathrm{O}_2}\)

Take some water, mixed with a few drops of dilute sulphuric acid, in a beaker. Invert two test tubes full of water into it. Remove the insulation from the ends of two thick wires. Introduce them into the test tubes as shown in. Connect the wires to a battery and pass current for some time. Gases start collecting in the test tubes. You will observe that the volume of the gas collected over the negative electrode is twice that of the gas collected over the positive electrode.

Stop passing current when there is enough gas in the test tubes. Cork the test tubes inside the water and take them out. Perform the following experiments.

The gas collected at the negative electrode:

Bring a lighted match near the mouth of the tube and open its mouth. The gas burns with a ‘pop’, and so it is hydrogen.

The gas collected at the positive electrode:

Bring a glowing matchstick near the mouth of the test tube and remove the cork. The matchstick gets lighted. So, the gas is oxygen

The experiment also shows that the volume ratio of hydrogen and oxygen in water is 2:1

Chemical Reactions and Their Types Class 8

The decomposition of baking soda

Sodium hydrogencarbonate (NaHC03) is called baking soda. On strong heating, baking soda decomposes into sodium carbonate, water (vapour) and carbon dioxide.

⇒  \(\underset{\begin{array}{c}
\text { sodium } \\
\text { hydrogencarbonate }
\end{array}}{2 \mathrm{NaHCO}_3(\mathrm{~s})} \xrightarrow{\text { heat }} \underset{\begin{array}{c}
\text { sodium } \\
\text { carbonate }
\end{array}}{\mathrm{Na}_2 \mathrm{CO}_3(\mathrm{~s})}+\underset{\begin{array}{c}
\text { water } \\
\text { vapour }
\end{array}}{\mathrm{H}_2 \mathrm{O}(\mathrm{~g})}\)

The decomposition of potassium chlorate

When strongly heated, potassium chlorate gives potassium chloride and oxygen.

⇒  \(\underset{\begin{array}{c}
\text { potassium } \\
\text { chlorate }
\end{array}}{2 \mathrm{KClO}_3(\mathrm{~s})} \xrightarrow{\text { heat }} \underset{\begin{array}{c}
\text { potassium } \\
\text { chloride }
\end{array}}{2 \mathrm{KCl}(\mathrm{~s})}+\underset{\text { oxygen }}{3 \mathrm{O}_2(\mathrm{~g})}\)

Displacement Reactions

In a displacement reaction, one element displaces another from its compound and takes its place therein.

Displacement reactions are best studied using the activity series. In the activity series, metals, along with hydrogen, are arranged in order of their activity. The higher an element is placed in the series, the more active it is. Thus, you can easily find out the relative activity of the metals along with hydrogen by consulting the series. For example, Na is more active than M,g and Mg is more active than Fe and so they fall in the following order of activity:

Na > Mg > Fe.

Could you say in what order will the activity of copper, silver, magnesium, iron and aluminium decrease? It has been observed that a more active metal displaces a less active one from its compounds.

Some common examples are given below.

The displacement of hydrogen from water by a metal:

An active metal like potassium (K), sodium (Na) or calcium (Ca) reacts with water, displacing hydrogen even in cold conditions.

2K(s) + 2H₂O(1) — 2KOH(aq) + H₂(g)

2Na(s) + 2H₂O(1) — 2NaOH(aq) + H₂(g)

Ca(s) + 2H₂O(1) — Ca(OH)₂(aq) + H₂ (g)

Magnesium, which is less active than the above-mentioned metals, reacts only at high temperatures.

When placed in steam, a burning magnesium ribbon continues to bum, forming magnesium oxide and hydrogen.

⇒ \(\underset{\text { magnesium }}{\mathrm{Mg}(\mathrm{~s})}+\underset{\text { water }(\text { steam })}{\mathrm{H}_2 \mathrm{O}(\mathrm{~g})} \longrightarrow \underset{\text { magnesium oxide }}{\mathrm{MgO}(\mathrm{~s})}+\underset{\text { hydrogen }}{\mathrm{H}_2(\mathrm{~g})}\)

In this reaction, a magnesium atom displaces two hydrogen atoms and takes their place

A less active metal like magnesium, zinc or iron easily displaces hydrogen from dilute hydrochloric or sulphuric acid.

⇒ \(\underset{\text { zinc }}{\mathrm{Zn}(\mathrm{~s})}+2 \mathrm{HCl}(\mathrm{aq}) \longrightarrow \underset{\text { zinc chloride }}{\mathrm{ZnCl}}(\mathrm{aq})+\mathrm{H}_2(\mathrm{~g})\)

⇒ \(\underset{\substack{\text { iron }}}{\mathrm{Fe}(\mathrm{~s})}+\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq}) \longrightarrow \underset{\text { iron(II) sulphate }}{\mathrm{FeSO}_4(\mathrm{aq})}+\mathrm{H}_2(\mathrm{~g})\)

The displacement of copper from copper sulphate by iron or zinc:

When an iron knife or nail is placed in a copper sulphate solution, there is a brown-red deposit of copper over the iron object

After some time, the blue colour of the solution changes to green, owing to the formation of iron(II) sulphate

⇒ \(\underset{\substack{\text { iron } \\ \text { igrey) }}}{\mathrm{Fe}(\mathrm{~s})}+\underset{\substack{\text { copper sulphate } \\ \text { (blue) }}}{\mathrm{CuSO}_4(\mathrm{aq})} \longrightarrow \underset{\text { copper }}{\text { (brown-red) }} \underset{\text { cuper }}{\mathrm{Cu}(\mathrm{~s})}+\underset{\text { (iron(I) sulphate }}{\mathrm{FeSO}_4}(\mathrm{aq})\)

Reactants and Products Class 8 Chemistry

Also, zinc displaces copper from copper sulphate, but the resultant solution is colourless. You can use a piece of granulated zinc in place of iron and it will be coated by copper

⇒ \(\underset{\substack{\text { zinc } \\ \text { (white) }}}{\mathrm{Zn}(\mathrm{~s})}+\underset{\substack{\text { copper sulphate } \\ \text { (blue) }}}{\mathrm{CuSO}_4(\mathrm{aq})} \rightarrow \underset{\substack{\text { copper } \\ \text { (brown-red) }}}{\mathrm{Cu}(\mathrm{~s})}+\underset{\substack{\text { zinc sulphate } \\ \text { (colourless) }}}{\mathrm{ZnSO}_4(\mathrm{aq})}\)

Double Displacement Reactions

In a double displacement reaction, the positive and negative radicals of two reactants are exchanged, leading to the precipitation of a product.

• Double displacement reactions are very fast, and the precipitate is formed as soon as the reactants come in contact.
• Quite often, the colour of the precipitate is different from that of the reactants. The following are some common examples
• The reaction between silver nitrate and sodium chloride
• When an aqueous solution of silver nitrate is mixed with that of sodium chloride, a white precipitate of silver chloride is formed. The sodium nitrate formed remains in solution.

AgNO₃ (aq) silver nitrate (colourless) + NaCl(aq) sodium chloride (colourless) → AgCl(s) + NaNO₃ (aq)

The reaction between barium chloride and sodium sulphate:

When a solution of barium chloride is mixed with that of sodium sulphate, a white precipitate of barium sulphate is formed along with a solution of sodium chloride.

The solution initially appears white but gradually becomes colourless as the precipitate settles down

\(\underset{\substack{\text { barium chloride } \\ \text { solution } \\ \text { (colourless) }}}{\mathrm{BaCl}_2(\mathrm{aq})}+\underset{\substack{\text { sodium sulphate } \\ \text { (colution } \\ \text { (colourless) }}}{\mathrm{Na}_2 \mathrm{SO}_4(\mathrm{aq})}\) → \(\underset{\substack{\text { barium sulphate } \\ \text { precipitate } \\ \text { (white) }}}{\mathrm{BaSO}_4(\mathrm{~s})}+\underset{\substack{\text { sodium chloride } \\ \text { solution } \\ \text { (colourless) }}}{2 \mathrm{NaCl}(\mathrm{aq})}\)

Neutralisation Reactions

In a neutralisation reaction, an acid reacts with a base, forming a salt and water. The acid as well as the base lose their properties to form a salt and water, which are neither acidic nor basic, i.e., they are neutral. So, the reaction is called a neutralisation reaction.

Some examples are given below:

Acid + Base  →  Salt+ Water

HCl(aq) + NaOH(aq)- NaCl(aq) + H₂O(I)

H₂SO₄(aq) + 2KOH(aq)→ K₂SO₄(aq) + 2H₂O(I)

H₂SO₄(aq) + 2NaOH(aq)→ Na₂SO₄(aq) + 2H₂O(I)

2HCl(aq) + Ca(OH)₂(aq)→CaCl₂(aq) + 2H₂O(I)

3HCl(aq) + Al(OH)₃(aq) → AlCl₃(aq) + 3H₂O(I)

Activity:

Take about 5 ml of dilute sodium hydroxide in a conical flask and dilute it with a test tube of water. Swirl the flask to ensure thorough mixing. Add a drop of phenolphthalein solution to it.

The contents of the flask turn pink. (Phenolphthalein is an indicator which turns red in a basic solution and colourless in a neutral or acidic solution.)

Add dilute hydrochloric acid dropwise with the help of a long dropper and swirl the contents after each addition. Add the acid till the pink colour just vanishes. The solution in the flask is neutral —neither acidic nor basic —and the reaction is neutralisation.’

Chemical Reactions in Everyday Life Class 8

Oxides And Their Nature

Oxides are an important class of compounds. We know that the two simple oxides, water (H₂O) and carbon dioxide (CO₂), play important roles in our life processes, agriculture and industry. Let us discuss the preparation and the acid-base nature of oxides in general.

Preparation

They can be prepared by the following processes.

By direct combination:

An element generally forms its oxide when heated or burned in air or oxygen. The reaction is exothermic and hence continues on its own when initiated.

Nonmetals:

⇒ \(2 \mathrm{H}_2(\mathrm{~g})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { ignite }} 2 \mathrm{H}_2 \mathrm{O}(\mathrm{~g})\)

⇒ \(\mathrm{C}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { burn }} \mathrm{CO}_2(\mathrm{~g}) \\\)

⇒ \(\mathrm{S}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { burn }} \mathrm{SO}_2(\mathrm{~g})\)

Metals:

⇒ \(2 \mathrm{Mg}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { ignite }} 2 \mathrm{MgO}(\mathrm{~s}) \\\)

⇒ \(2 \mathrm{Ca}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \xrightarrow{\text { ignite }} 2 \mathrm{CaO}(\mathrm{~s})\)

By the thermal decomposition of some compounds

The oxides of both metals and nonmetals are formed when the hydroxides, carbonates, sulphates and nitrates of metals are strongly heated

1. Hydroxides:

1. \(\mathrm{Zn}(\mathrm{OH})_2(\mathrm{~s}) \xrightarrow{\text { heat }} \mathrm{ZnO}(\mathrm{~s})+\mathrm{H}_2 \mathrm{O}(\mathrm{~g}) \\\)

2. \({\text {(bluish white) }}{\mathrm{Cu}(\mathrm{OH})_2(\mathrm{~s}) \xrightarrow{\text { heat }} \underset{\text { (black) }}{\mathrm{CuO}(\mathrm{~s})}+\mathrm{H}_2 \mathrm{O}(\mathrm{~g})}\)

2. Carbonates:

1.\(\underset{\text { (white) }}{\mathrm{CaCO}_3(\mathrm{~s})} \xrightarrow{\text { heat }} \underset{\text { (white) }}{\mathrm{CaO}(\mathrm{~s})}+\mathrm{CO}_2(\mathrm{~g}) \\\)

2.\(\underset{\text { (white) }}{\mathrm{ZnCO}_3(\mathrm{~s})} \xrightarrow{\text { heat }} \underset{\text { (white) }}{\mathrm{ZnO}(\mathrm{~s})}+\mathrm{CO}_2(\mathrm{~g}) \\\)

3.\(\underset{\text { (white) }}{\mathrm{PbCO}_3(\mathrm{~s})} \xrightarrow{\text { heat }} \underset{\text { (yellow) }}{\mathrm{PbO}(\mathrm{~s})}+\mathrm{CO}_2(\mathrm{~g}) \\\)

4.\( \mathrm{CuCO}_3(\mathrm{~s}) \xrightarrow[\text { (green) }]{\text { heat }} \underset{\text { (black) }}{\mathrm{CuO}(\mathrm{~s})}+\mathrm{CO}_2 \text { (s) }\)

Exceptions: Na₂CO₃ and K₂CO₃

3. Sulphates:

1.\(\mathrm{Fe}_2\left(\mathrm{SO}_4\right)_3(\mathrm{~s}) \xrightarrow{\text{heat}}\mathrm{Fe}_2\mathrm{O}_3(\mathrm{~s})+3 \mathrm{SO}_3(\mathrm{~g}) \\\)

2. \(\underset{\text { (white) }}{\mathrm{CuSO}_4(\mathrm{~s})} \xrightarrow{\text { heat }} \underset{\text { (black) }}{\mathrm{CuO}(\mathrm{~s})}+\mathrm{SO}_3(\mathrm{~g})\)

Exceptions: Na₂SO₄, K₂SO₄, CaSO₄ and many others

4. Nitrates:

1. \(\begin{aligned}\mathrm{Ca}\left(\mathrm{NO}_3\right)_2(\mathrm{~s}) \xrightarrow{\text { heat }} 2 \mathrm{CaO}(\mathrm{~s})+4 \mathrm{NO}_2(\mathrm{~g})&\text {(brown) }+\mathrm{O}_2(\mathrm{~g})\\\end{aligned}\)

2.\(\underset{\text { (colourless) }}{2 \mathrm{~Pb}\left(\mathrm{NO}_3\right)_2(\mathrm{~s})} \xrightarrow[\text {}]{\text { (heat) }} \underset{\text { (yellow) }}{2 \mathrm{PbO}(\mathrm{~s})}+\underset{(brown)}{4 \mathrm{NO}_2(\mathrm{~g})}+\mathrm{O}_2(\mathrm{~g})\)

Exceptions:

NaNO₃ and KNO₃, on being heated, decompose to form sodium nitrite (NaNO₂) and potassium nitrite (KNO₂), respectively, plus oxygen.

It is left to you to write the chemical equations for the reactions.

Types of Reactions: Combination, Decomposition, and Displacement

Classification

Oxides have been classified in many ways. Here we will take up a classification based on their acid-base character

Neutral oxides:

Neutral oxides, which are only four in number, are neither acidic nor basic. They are:

• Water (H₂O)
• Carbon monoxide (CO)
• Nitrous oxide (N₂O), and
• Nitric oxide (NO)

They do not change the character of an acid or a base.

Acidic oxides:

The oxides of nonmetals, in general, are acidic. Reaction with water They dissolve in water to form acids, which neutralise bases to form salts and water.

⇒ \(\mathrm{CO}_2(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \underset{\text { Carbonic acid }}{\mathrm{H}_2 \mathrm{CO}_3(\mathrm{aq})}\)

⇒  \(\mathrm{SO}_2(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \underset{\text { Sulphurous acid }}{\mathrm{H}_2 \mathrm{SO}_3(\mathrm{aq})}\)

⇒  \(\underset{\substack{\text { sulphur } \\ \text { trioxide }}}{\mathrm{SO}_3(\mathrm{l})}+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \underset{\text { sulphuric acid }}{\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq})}\)

⇒  \(\underset{\text { dinitric pentoxide }}{\mathrm{N}_2 \mathrm{O}_5(\mathrm{~s})}+\underset{\text { }}{\text { }} \mathrm{H}_2 \mathrm{O}(\mathrm{l}) \longrightarrow \underset{\text { }}{2 \mathrm{HNO}_3(\mathrm{aq})}\)

A compound that is obtained by removing the elements of water from an acid is called the anhydride of the acid. In other words, the anhydride of an acid is the compound which adds water to form the acid. Thus, these oxides are acid anhydrides

You know that an acid turns blue litmus red. Acidic oxides turn moistened blue litmus paper red—they first add up water (moisture of the litmus paper) to form the acids. Else you can use a solution of blue litmus.

Reaction with bases: They directly react with bases to form salts.

⇒  \(\mathrm{CO}_2(\mathrm{~g})+\mathrm{Na}_2 \mathrm{O}(\mathrm{~s}) \longrightarrow \underset{\text { Sodium carbonate}}{\mathrm{Na}_2 \mathrm{CO}_3(\mathrm{~s})}\)

⇒ \(\mathrm{SO}_2(\mathrm{~g})+2 \mathrm{KOH}(\mathrm{~s}) \longrightarrow \underset{\substack{\text { potassium } \\ \text { sulphite }}}{\mathrm{K}_2 \mathrm{SO}_3(\mathrm{~s})}+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

Basic oxides

The oxides of metals are generally basic and react with acids to form salts and water. They also react with acidic oxides to form salts.

⇒ \(\mathrm{Na}_2 \mathrm{O}(\mathrm{~s})+\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq}) \longrightarrow \mathrm{Na}_2 \mathrm{SO}_4(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \\\)

⇒ \(\mathrm{MgO}(\mathrm{~s})+2 \mathrm{HCl}(\mathrm{aq}) \longrightarrow \mathrm{MgCl}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

⇒ \(\mathrm{CaO}(\mathrm{~s})+\mathrm{CO}_2(\mathrm{~g}) \longrightarrow \mathrm{CaCO}_3(\mathrm{~s})\)

Basic oxides which are soluble at least to some extent turn moistened red litmus paper blue. Thus, a freely soluble metal oxide like Na₂O, a less soluble oxide like CaO and an only slightly soluble oxide like MgO turn moistened red litmus paper blue, but the insoluble CuO does not. All these are bases

Amphoteric oxides

An amphoteric oxide is one which behaves like a base in the presence of an acid and like an acid in the presence of a base. These are metal oxides like Al₂O₃, ZnO and PbO. (Remember that the corresponding hydroxides are also amphoteric.)

An amphoteric oxide reacts with an acid as well as with a base, forming a salt and water.

⇒ \(\text { Base }+ \text { acid } \rightarrow \text { salt }+ \text { water }\)

⇒ \(\mathrm{ZnO}(\mathrm{~s})+\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq}) \rightarrow \underset{\text { zinc sulphate }}{\mathrm{ZnSO}_4(\mathrm{aq})}+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

⇒ \(\mathrm{PbO}(\mathrm{~s})+2 \mathrm{HNO}_3(\mathrm{l}) \rightarrow \underset{\text { lead nitrate }}{\mathrm{Pb}\left(\mathrm{NO}_3\right)_2(\mathrm{aq})}+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

⇒ \(\text { Base }+ \text { acid } \rightarrow \text { salt }+ \text { water }\)

⇒ \(2 \mathrm{NaOH}(\mathrm{aq})+\mathrm{ZnO}(\mathrm{~s}) \rightarrow \underset{\text { sodium zincate }}{\mathrm{Na}_2 \mathrm{ZnO}_2(\mathrm{aq})}+\mathrm{H}_2 \mathrm{O}(\mathrm{l})\)

⇒ \(2 \mathrm{NaOH}(\mathrm{aq})+\mathrm{PbO}(\mathrm{~s}) \rightarrow \underset{\text { sodium plumbite }}{\mathrm{Na}_2 \mathrm{PbO}_2(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l})}\)

Activities:

1 . Take about 3-5 g of zinc carbonate (ZnCO₃) in a dry test tube. Heat it strongly on a flame till the white solid turns yellow. Cool the solid and it will turn white again. This colour change can be observed again and again on heating and cooling. This is the special property of zinc oxide (ZnO). On being strongly heated, ZnCO₃ decomposes into Zno (and CO₂).

Now, divide the residue of ZnO into two parts and treat them as follows.

2. Add some dilute sulphuric acid (or dilute hydrochloric acid) to one part. The solid quickly dissolves (to form the soluble salt ZnSO₄)

Add some dilute sodium hydroxide solution to the other part and shake well. The solid will dissolve (forming the soluble salt Na₂ZnO₂).

2. Heat, as above, about 3-5 g of PbCO₃. This time, the residue is the yellow PbO yellow in hot as well as cold conditions. Divide the residue into two parts and treat them as follows

Add dilute nitric acid to one part. The yellow solid dissolves to form a colourless solution of Pb(NO₃)₂. Do not use dilute HCI or H₂SO₄, because the salts PbCI₂ or PbSO₄ would be formed, which, being insoluble, would form a crust over the solid PbO, preventing it from reacting

Add some dilute NaOH solution to the other part and shake well. The solid will dissolve to form a colourless solution of sodium plumbite, Na₂PbO₂.