Indicators Definitions, Equations and Examples

Indicators

These are substances that tell us whether a substance is an acid or a base by change in colour or odour when added into the substance.

Types of Indicators:

Indicators are classified in the following ways:
Natural Indicators
Natural indicators are the substances obtained from natural sources such as a dye, flower, leaf, etc., in plants.
Examples: Litmus (a purple coloured dye extracted from the lichen plant), red cabbage leaf extract, etc.

The table below shows the effect of some common natural indicators on acids and bases:
Indicators Definitions, Equations and Examples 1
Important
The yellow stain of curry on a white cloth, which turns reddish-brown when scrubbed with soap is due to the fact that soap solution changes the colour of turmeric to red- brown as it is basic in nature. And when the cloth is rinsed with water, the basic soap gets removed and the stain turns to yellow again.

Indicators Definitions, Equations and Examples

Example 1.
You have been provided with three test tubes. One of them contains distilled water and the other two contain an acidic solution and a basic solution, respectively. If you are given only red litmus paper, how will you identify the contents of each test tube?
Answer:
Label the three test tubes as A, B and C. Take a small amount of each of the three given liquids in these test tubes A, B and C. First, dip red litmus paper one by one in each test tube. The liquid which will turn red litmus paper blue is basic in nature.

Remove this test tube as it has been identified as containing a basic solution.
Next, put a drop of solution from the remaining two test tubes on the blue litmus paper (which was earlier red in colour but turned blue on adding basic solution to it).

The solution which turns blue litmus paper red is acidic in nature. The solution that does not change the colour of either red litmus or blue litmus is neutral in nature, which is distilled water.

Synthetic Indicators:

Synthetic indicators are the compounds synthesized in a chemistry lab or industrially rather than compounds found in nature.
Examples: Phenolphthalein, Methyl orange.

The table below shows the effect of some common synthetic indicators on acids and bases:
Indicators Definitions, Equations and Examples 2

Indicators Definitions, Equations and Examples

Olfactory Indicators:

These are the substances whose odour or smell changes in acidic or basic medium.
Examples: Onion extract, vanilla extract. The table beLow shows the effect of some common olfactory indicators on acids and bases:
Indicators Definitions, Equations and Examples 3
Indicators Definitions, Equations and Examples 4
Important
Olfactory indicators ore used to ensure the participation of visually impaired students ¡n the Laboratory for distinguishing between acidic and basic substances

Class 10 Science Notes

Reaction of Acids With Bases Definitions, Equations and Examples

The reaction of Acids With Bases:

Acids react with bases to form salt and water. This is known as neutralization reaction as the effect of a base is nullified by acid and vice-versa.
Example: NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

Neutralization Reaction:

The reaction between an add and a base to give salt and water is known as a neutralization reaction. In general, a neutralization reaction can be written as
Acid + Base → Salt + Water

Example 1.
10 mL of a solution of NaOH is found to be completely neutralized by 8 mL of a given solution of HCl. If we take 20 ml of the same solution of NaOH, the amount of HCl solution (the same solution as before) required to neutralize it will be
(a) 4 mL
(b) 8 mL
(c) 12 mL
(d) 16 mL
Answer:
(d) 16 mL

Explanation: As 10 mL of NaOH solution is completely neutralized by 8 mL of HCl solution, 16 mL of HCl solution will be required to neutralize 20 mL of NaOH solution.

Reaction of Acids With Bases Definitions, Equations and Examples

Similarities among Acids and Bases:

All acids have similar chemical properties. All acids generate hydrogen gas on reacting with metals, so hydrogen is common to all acids.
Acids produce hydrogen ions, H+(aq), in solution, which are responsible for their acidic properties.

Solutions of acids and bases conduct electricity as both acids and bases dissociate into hydrogen ions (H+) and hydroxide ions (OH–) respectively, which are responsible for the conduction of electric current through the solution.

Example 2.
Case-Based:
Take solutions of glucose, alcohol, hydrochloric acid, sulphuric acid, etc. Fix two nails on a cork, and place the cork in a 100 mL beaker. Connect the nails to the two terminals of a 6-volt battery through a bulb and a switch, as shown in Fig. below.
Reaction of Acids With Bases Definitions, Equations and Examples 1
Now pour some dilute HCl into the beaker and switch on the current. Repeat with other solutions one by one.

(A) Four students recorded the following observations. Identify the correct observation:

Electrolyte Observation
(a) Dilute HCl The bulb does not glow
(b) Glucose Bulb glows
(c) Dilute NaOH Bulb glows

Answer:
(c) Electrolyte: Dilute HCl; Observation: Bulb glows
Explanation: The bulb glows when the electrolyte is an acid or a base as in both cases ions are produced in an aqueous solution due to which current flows through the solution.

(B) The activity is repeated with alcohol as the electrolyte. Which of the following statement(s) is (are) correct?
(I) Bulb will not glow because the electrolyte is not acidic.
(II) Bulb will glow because alcohol contains Hydrogen and provides ions for the conduction of electricity.
(III) Bulb will not glow because alcohol does not provide ions for the conduction of electricity.
(IV) Bulb will glow because it depends upon the type of electrolytic solution.
(a) (I) and (III)
(b) (II) and (IV)
(c) (III) only
(d) (IV) only
Answer:
(c) (III) only

(C) Why do HCl, HNO3, etc., show acidic characters in aqueous solutions while solutions of compounds like alcohol and glucose do not show acidic character?
Answer:
Solutions like HCl, HNO3, etc. show acidic characters in aqueous solutions as they contain the cation H+ and produce hydrogen ions, H+(aq) in solution. Whereas solutions of compounds such as alcohol and glucose do not form any such ions and hence they do not show acidic character.

(D) Why does an aqueous solution of acid conduct electricity?
Answer:
An aqueous solution of an acid conducts electricity as acids dissociate into hydrogen ions (H+) in the solution, which is responsible for the conduction of electricity through the solution.

(E) Assertion: An aqueous solution of glucose conducts electricity.
Reason: Electric current is carried through the solution by ions.
(a) Both (A)and R are true, and (R) is the correct explanation of the assertion (A).
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the assertion (A).
(c) (A) is true, but (R) is false.
(d) (A) is false, but (R) is true.
Answer:
(d) (A) is false, but (R) is true.
Explanation: An aqueous solution of glucose does not conduct electricity as glucose does not dissociate into ions even in solution. Electric current will flow in a solution only by ions, as happens in the case of an aqueous solution of acids and bases.

Reaction of Acids With Bases Definitions, Equations and Examples

Example 3.
Why does distilled water not conduct electricity, whereas rainwater does?
Answer:
Electric current is conducted in solutions by ions. As distilled water does not contain any ions, it does not conduct electricity. However, rainwater conducts electricity due to the presence of small amounts of acidic oxides in rainwater such as SO2, NO2 which make It a better conductor of electricity.

Acid and Bases in Water Solution:

Hydrogen or Hydronium Ions
All acids produce hydrogen ions (H+) only in presence of water. The separation of H+ ions from HCl molecules cannot occur in the absence of water.
HCl + H2O → H3O+ + Cl–

Hydrogen ions cannot exist alone, but they exist after combining with water molecules. Thus hydrogen ions must always be shown as H+(aq) or hydronium ion (H3O+).
H+ + H2O → H3O+

Hydroxide Ions:

Bases produce hydroxide ions (OH–) in water. Example: When sodium hydroxide is dissolved in water, it dissociates to form Na+ ions and OH– ions. NaOH(aq) → Na (aq) + OH

Alkalies:

All bases are not soluble in water. The bases which are soluble in water are called alkalis.

Reaction of Acids With Bases Definitions, Equations and Examples

Example 4.
How is the concentration of hydroxide ions (OH–) affected when excess base is dissolved in a solution of sodium hydroxide?
Answer:
A solution of sodium hydroxide is strongly basic due to the formation of hydroxide ions (OH-).
NaOH(s) → Na (aq) + OH (aq)

When the excess base is dissolved in such a solution, the concentration of hydroxide ions will increase further.

Explanation of Neutralization Reactions:

Neutralization reaction between acids and bases can be explained in terms of hydrogen ions produced by acids and hydroxide ions produced by bases as follows:
Reaction of Acids With Bases Definitions, Equations and Examples 2

Dilution of Acids and Bases:

When an acid or a base is added to water, lot of heat is produced. The process of dissolving an acid or a base in water is a highly exothermic reaction.

The process of mixing an acid or base with water results in a decrease in the concentration of ions (H3O+/OH–) per unit volume and this process is called dilution. The acid or the base is said to be diluted.

The acid must always be added slowly to water with constant stirring. If water is added to a concentrated
acid, the heat generated may cause the mixture to splash out and cause burns.

Example 5.
While diluting an acid, why is it recommended that the acid should be added to water and not water to the acid?
Answer:
Dilution of concentrated acid is a highly exothermic process. If we add water to the concentrated acid, the mixture may splash and cause severe burn injuries due to a large amount of heat produced. It is therefore recommended that acid should be added to the water drop by drop with constant stirring so that the heat evolved may be absorbed by the water.

Reaction of Acids With Bases Definitions, Equations and Examples

Example 6.
How is the concentration of hydronium ions (H3O+) affected when a solution of an acid is diluted?
Answer:
When a solution of an acid is diluted by adding acid to the water, the concentration of hydronium ions decreases due to the increase in the volume of the solution. As the number of hydronium ions remains the same on diluting the acid, their concentration decreases.

Class 10 Science Notes

Bases Definitions, Equations and Examples

Bases:

Bases are chemical substances that have a bitter taste, are soapy to touch, and turn red litmus blue. A base is a substance that dissolves in water to produce hydroxide ions (OH–) in the solution.
NaOH(s) → Na+(aq) + OH–(aq)

Classification of Bases:

Bases are of two types:
Strong Bases:
A base that completely ionizes in water and thus produces a large amount of hydroxide ions is called a strong base.
Examples: Sodium hydroxide and potassium hydroxide.

Weak Bases:
A base that is partially ionized in water and thus produces a small amount of hydroxide ions is called a weak base.
Examples: Ammonium hydroxide and calcium hydroxide.

Bases Definitions, Equations and Examples

Properties of Bases (Physical Properties):

  1. Bases have a bitter taste.
  2. Bases feel soapy to touch.
  3. Bases turn red litmus to blue.
  4. Bases conduct electricity in solution due to the presence of ions in water.

Chemical Properties:

1. Reaction with metals
Bases react with some metals to form hydrogen gas. The gas evolved burns with a pop sound showing that it is hydrogen gas.
Base + Metal → Salt + Hydrogen gas
2NaOH(aq) + Zn(s) → Na2ZnO2(aq) + H2(g)

2. Reaction of a Non-metallic Oxide with Base:
Bases react with non-metallic oxides to produce salt and water.
Example: Carbon dioxide, which is a non-metallic oxide reacts with Calcium hydroxide, which is a base, to produce salt and water.
Since this is similar to the reaction between a base and an acid, we can say that nonmetallic oxides are acidic in nature.

Base + Non-metallic oxide → Salt + Water
Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

Class 10 Science Notes

Acids Definitions, Equations and Examples

Acids:

On the basis of their chemical properties, all the chemical compounds can be classified into:

  • Acids
  • Bases
  • Salts

In this chapter, we will be studying the various properties of acids and bases, synthetic, natural, and olfactory indicators, behavior of acids and bases in solutions, strength of acids and bases, pH and universal indicator, the importance of pH in everyday life, family of salts, chemicals obtained from common salt, their manufacture, properties and uses, and water of crystallization.

Acids are those chemical substances which have a sour taste and change the colour of blue litmus to red.

Classifications of Acids:

(A) Based on the source, acids can be classified as mineral acids and organic acids.
Mineral Acids:
The acids prepared from the minerals of the earth are called mineral acids. They are man-made acids. Examples: Hydrochloric acid, sulphuric acid, nitric acid, etc.

Organic Acids:
The acids present in plant materials and animals are called organic acids or naturally occurring acids.

Few examples of naturally occurring acids along with their source are given below:

Name of Acid Source
Lactic Acid Curd
Methanoic Acid Insect sting
Oxalic Acid Tomato
Tartaric Acid Tamarind
Citric Acid Citrus fruits such as lemon, oranges, etc.
Acetic Acid Vinegar

(B) Based on the degree of dissociation in water, acids can be classified as strong acids and weak acids.
Strong Acids
Acids that give rise to more H+ ions in water are said to be strong acids. In other words, strong acids dissociate completely into hydrogen ions in water. Examples: Hydrochloric acid, sulphuric acid.

Weak Acids:
Acids that are partially ionized in water and give less H+ ions in water are said to be weak acids. In other words, their degree of dissociation into hydrogen ions in water is not much.

Examples: Acetic acid, carbonic acid.

Acids Definitions, Equations and Examples

Properties of Acids:

Physical Properties

  1. Acids have a sour taste.
  2. Acids are soluble in water.
  3. Acids turn blue litmus red.
  4. Acid solutions conduct electricity.

Chemical Properties of Acids:

1. Reaction with metals: Acids react with metals to form hydrogen gas. When a metal reacts with an acid, it displaces hydrogen from the acids. The metal combines with the remaining part of the acid and forms a compound called salt.
Acid + Metal → Salt + Hydrogen gas
Zn + H2SO4 → ZnSO4 + H2

Example 1.
Why should curd and sour substances not be kept in brass and copper vessels?
Answer:
Curd is basically lactic acid. Similarly, sour substances also contain acids. We know that acids react with metals to form salt and water. Therefore, if curd and sour substances are stored in brass and copper vessels, the acid will react with the metal and may produce some salts which are poisonous and hence unfit for human consumption.

2. Reaction with metal carbonates and metal hydrogen carbonates: Acids react with metal carbonates and metal hydrogen carbonates to form carbon dioxide gas.
Metal carbonate/hydrogen carbonate + Acid → Salt + carbon dioxide + water
Na2CO3(S) + 2HCl(aq) → 2NaCl(aq) + H2O(l) + CO2(g)
NaHCO3(s) + HCl(aq) → NaCl(aq) + H2O(l) + CO2(g)

Important
Test for carbon dioxide gas:
When carbon dioxide is passed through lime water, the lime water turns milky due to the formation of a white precipitate of calcium carbonate, CaCO(3)
Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

On passing excess carbon dioxide through lime water, the white precipitate formed first dissolves due to the formation of a soluble salt calcium hydrogen carbonate and the solution becomes clear again.
CaCO3(s) + CO2(g) + H2O(l) → Ca(HCO3)2(aq)

Acids Definitions, Equations and Examples

Example 2.
A solution reacts with crushed eggshells to give a gas that turns lime-water milky. The solution contains:
(a) NaCl
(b) HCl
(c) Li Cl
(d) KCl
Answer:
(b) HCl
Explanation: As the gas evolved turns lime water milky, the gas is carbon dioxide. Moreover, the crushed eggshells contain calcium carbonate, which reacts with dilute HCl to give carbon dioxide, salt, and water.

The equation for the reaction taking place is:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(aq) + CO2(g)

3. Reaction with bases: Acids react with bases to form salt and water.
Base + Acid → Salt + Water
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

4. Reaction with metal oxides: Acids react with metal oxides to form salt and water.
Metal Oxide + Acid → Salt + Water
CuO(s) + 2HCl(aq) → CuCl2(aq) + H2O(l)

Example 3.
Case Based:
About 5 ml of dilute sulphuric acid is taken in a test tube ‘A’ and a few pieces of zinc granules are added to it. The gas evolved is passed through the soap solution. The activity is repeated with some more acids like HCl, HNO3 and CH3COOH.
Acids Definitions, Equations and Examples 1
Soap bubble filled with hydrogen -Soap solution
Two more test tubes labeled as ‘B’ and ‘C are taken and about 0.5 g of sodium carbonate (Na2CCO3) is taken in test tube ‘B’ and about 0.5 g of sodium hydrogen carbonate (NaHCO) in test tube ‘Cl About 2 mL of dilute HCl is added to both the test tubes. The gas produced in each case is passed through lime water (calcium hydroxide solution).
Acids Definitions, Equations and Examples 2
(A) A student recorded his observations on the gas produced in test tubes as:
Acids Definitions, Equations and Examples 3
Identify the correct observation:
Answer:
(d) Test Tube A: hydrogen gas; Test Tube B: Carbon dioxide; Test Tube C: carbon dioxide Explanation: The gas produced in test tube ‘A’ is hydrogen and that in test tubes ‘B’ and ‘C’ is carbon dioxide as the reactions taking place in the test tubes are given below:
Test tube A: Zn + H2SO4 → ZnSCO4 + H2
Test tube ‘B’: Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + H2O(l) + CO2(g)
Test tube ‘C’: NaHCO3(s) + HCl(aq) → NaCl(aq) + H2O(l) + CO2(g)

(B) Which of the following observations are incorrect?
(I) When the gas evolved in test tube ‘A’ is passed through a soap solution, the soap bubbles formed extinguish a burning candle.
(II) When the gas evolved in test tube ‘A’ is passed through a soap solution, the soap bubbles burn with a pop sound when a burning candle is brought near a bubble.
(III) The gas evolved in test tube ‘B’ turns lime water milky.
(IV) The gas evolved in test tube ‘C’ burns with a pop sound.
(a) Both (I) and (IV)
(b) Both (II) and (III)
(c) (I), (III), and (IV)
(d) (II), (III) and (IV)
Answer:
(a) Both (I) and (IV)
Explanation: The gas evolved in test tube A’ is hydrogen gas, which burns with a pop sound. The gas evolved in test tubes ‘B’ and ‘C’ is carbon dioxide which turns lime water milky.

Acids Definitions, Equations and Examples

(C) Are the observations in test tube ‘A’ same when the activity is repeated with more acids like HCl, HNO3 and CH3COOH?
Answer:
Yes, the same observations are seen when the activity is repeated with more acids like HCl, HNO3, and CH3COOH as metals react with acids and form hydrogen gas by displacing hydrogen atoms from the acid and also forms a salt.
For example, Zn + 2HCl → ZnCl2 + H2

(D) Metal compound A reacts with dilute hydrochloric acid to produce effervescence. The gas evolved extinguishes a burning candle. Write a balanced chemical equation for the reaction if one of the compounds formed is calcium chloride.
Answer:
As the gas evolved on reacting the metal compound with dilute hydrochloric acid produces effervescence and also extinguishes a burning candle, the gas evolved is carbon dioxide.

Since one of the compounds formed is calcium chloride, metal A is calcium chloride, since metal carbonates react with acid to produce carbon dioxide gas, metal salt, and water.
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)

(E) Assertion: When an acid reacts with a metal carbonate or metal hydrogen carbonate, the gas evolved turns lime water milky.
Reason: When excess carbon dioxide gas is
passed through lime water, calcium hydrogen carbonate solution is formed.
(a) Both (A) and (R) are true and (R) is the correct explanation of the (A).
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the (A).
(c) (A) is true, but (R) is false.
(d) (A) is false, but (R) is true.
Answer:
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the (A).

Class 10 Science Notes

Effects of Oxidation Reactions in Everyday Life Definitions, Equations and Examples

Corrosion:

The process in which metals are eaten up gradually by the action of air, moisture or a chemical such as acids on their surface is called corrosion.
For example, rusting of iron metal is the most common form of corrosion.
4Fe + 3CO2 + 2x H2O → 2Fe2CO3. xH2O
The rusting of iron is a redox reaction.

The black coating on silver and the green coating on copper are other examples of corrosion.
Effects of corrosion: Damage to bridges, iron railings, bodies of vehicles, ships and all objects made of metals like iron.

Effects of Oxidation Reactions in Everyday Life Definitions, Equations and Examples

Example 1.
Why do we apply paint on iron articles?
Answer:
If iron particles are left exposed, their surface comes in contact with the atmospheric oxygen and in the presence of moisture it forms Iron(III) oxide or rust. They are therefore painted to prevent them from rusting as the paint does not allow the surface to come in contact with oxygen and moisture and thus prevents rusting.

Rancidity:

The slow oxidation of fats and oils in foods marked by unpleasant smell and taste is called rancidity.

Prevention of Rancidity

  1. By adding anti-oxidants to foods containing fats and oils: Usually substances that prevent oxidation (antioxidants) are added to foods containing fats and oil. The two common anti-oxidants are BFIA (Butylated Hydroxy-Anisole) and BHT (Butylated Hydroxy-Toluene).
  2. Keeping food in air tight containers helps to slow down oxidation.
  3. By packaging fat and oil-containing foods in nitrogen gas: Chips manufacturers usually flush bags of chips with gas such as nitrogen to prevent the chips from getting oxidized.
  4. By keeping food in a refrigerator.
  5. By storing foods away from light.

Effects of Oxidation Reactions in Everyday Life Definitions, Equations and Examples

Example 2.
Oil and fat-containing food items are flushed with nitrogen. Why?
Answer:
Fried food items containing oil and fat get oxidized and become rancid in the presence of air or oxygen. As nitrogen is a comparatively unreactive gas as compared to oxygen, it prevents the oxidation of food items and hence prevents food from becoming rancid. That is why oil and fat containing food items are flushed with nitrogen gas.

Class 10 Science Notes

Types of Chemical Reactions Definitions, Equations and Examples

Types Of Chemical Reactions

Chemical reactions involve the breaking of bonds present in reactant atoms and forming new bonds in products. Some of the important types of chemical reactions are described next:

Combination Reaction (A + B → AB)
Those reactions in which two or more reactants combine to form a single product are known as combination reactions.

In a combination reaction:

  • Two or more elements can combine to form a compound.
  • Two or more compounds can combine to form a new compound.
  • An element and a compound can combine to form a new compound.

Examples of Combination Reactions

Example of Combination Reaction Chemical Equation
1. Burning of hydrogen in air to form water 2 H2(g) + O2(g) → 2H2O
2. Burning of carbon (coal) in air to form carbon dioxide C(s) + O2(g) → CO2(g)
3. Reaction of hydrogen                with chlorine to form hydrogen chloride H2(g) + Cl2(g) → 2HCl(g)
4. Burning of sodium metal in chlorine to form sodium chloride. 2Na(s) + Cl2 (g) → 2NaCl(s)
5. Heating of iron with sulphur to form iron sulphide Fe(s) + S(s)  → FeS(s)

Types of Chemical Reactions Definitions, Equations and Examples

Example 1.
A solution of a substance ‘X’ is used for whitewashing.
(A) Name the substance ‘X’ and write its formula.
Answer:
The substance ‘X’ which is used in whitewashing is quick lime or Calcium Oxide. It’s chemical formula is CaO.

(B) Write the reaction of the substance ‘X’ named in above with water.
Answer:
Calcium oxide reacts with water to form slaked lime or calcium hydroxide and a large amount of heat is released in this process.
Types Of Chemical Reactions Definitions, Equations and Examples 1

Exothermic and Endothermic Reactions:

Chemical reactions are accompanied with a change in temperature of the reaction mixture. There can be either a rise in temperature or a fall in temperature.

Exothermic Reactions:

Reactant(s) → Produces + Heat
Those reactions in which heat is evolved are known as exothermic reactions. Some examples of exothermic reactions are:

  • Burning of carbon in air to form carbon dioxide:
    C(S) + O2(g) → CO2(g) + Heat
  • Burning of natural gas:
    CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + Heat
  • Reaction between calcium oxide with water to form slaked lime:
    CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat

Important
The reaction between calcium oxide with water to form slaked lime is used for whitewashing walls. The reaction taking place is:
Ca(OH)2(aq) + CO2(g) → CaCO3(S) + H2O(l)

  • Respiration: It is the release of energy by the oxidation of glucose in the cells of our body.
    C6H12O6 + 6O2 → 6CO2 + 6H2O + Energy
  • Decomposition of vegetable matter into compost.
  • All combustion reactions.

Types of Chemical Reactions Definitions, Equations and Examples

Endothermic Reactions: (Reactant(s) + Heat → Products(s))

Those reactions in which heat is absorbed are known as endothermic reactions. Some examples of endothermic reactions are:

  1. Formation of nitrogen monoxide by heating nitrogen and oxygen to a temperature of about 3000°C.
    N2(g) + O2(g) + Heat → 2NO(g)
  2. All decomposition reactions are endothermic reactions since the decomposition reactions require energy either in the form of heat, light or electricity for breaking down the reactants.
    For example, when calcium carbonate is heated, it decomposes to form calcium oxide and carbon dioxide:
    CaCO3(S) + Heat → CaO(s) + CO2(g)
  3. Photosynthesis is an endothermic process as sunlight is absorbed by green plants during photosynthesis resulting in the formation of glucose, water and oxygen.
    Types Of Chemical Reactions Definitions, Equations and Examples 2

Decomposition Reaction (AB → A + B):

Those reactions in which a compound splits up into two or simpler substances are known as decomposition reactions.

Depending upon the form of energy required for the reaction, there are three types of decomposition reactions:

Thermal Decomposition Reactions:
When a decomposition reaction is carried out by heating, it is called thermal decomposition.
For example, decomposition of calcium carbonate to calcium oxide and carbon dioxide on heating is an important decomposition reaction used in various industries.
Types Of Chemical Reactions Definitions, Equations and Examples 3

Example 2.
Case Based:
Two boiting tubes were taken, about 2 grams of a green coloured metal salt A’ was taken in the first tube and 2 grams of a white coloured metal salt ‘B’ was taken in the second tube. Both the tubes were heated by holding them with a pair of tongs. Smell of burning sulphur was observed in first test tube whereas brown gas was emitted in the second test tube.
(A) The salts ‘A’ and ‘B’ are:
(a) Ferrous nitrate and lead sulphate respectively
(b) Ferric oxide and Lead nitrate respectively
(c) Ferrous sulphate and lead nitrate respectively
(d) Ferric oxide and lead sulphate respectively
Answer:
(c) Ferrous sulphate and lead nitrate respectively

Explanation: The green coloured metal salt ‘A’ in the first test tube is ferrous sulphate as it is green in colour and when heated decomposes to form ferric oxide, sulphur dioxide and sulphur trioxide, which smell of burning sulphur.

Equation for the reaction talcing place is:
Types Of Chemical Reactions Definitions, Equations and Examples 5
The white coloured metal salt ‘B’ in the second test tube is lead nitrate as brown fumes of a gas is emitted when it is heated, which are of nitrogen dioxide.

Equation for the reaction talcing place is:
2Pb(NO3)2(S) → 2PbO(s) + 4NO2(g) + O2(g)

(B) During the experiment of heating of ferrous sulphate, four students recorded their observation as
(I) Green colour changes to brown-black colour.
(II) Brownish-yellow gas is evolved.
(III) Blue colour changes to green colour.
(IV) Smell of burning sulphur is observed. Which of the above observations are incorrect?
(a) Both (I) and (II)
(b) Both (I) and (IV)
(c) Both (II) and (III)
(d) Both (III) and (IV)
Answer:
(b) Both (I) and (IV)
Explanation: Ferrous sulphate crystals are green in colour. When they are heated, they lose their molecules of water of crystallization and form ferric oxide, sulphur dioxide and sulphur trioxide gases. The green colour of ferrous sulphate crystal changes to brown-black and smell of burning sulphur is observed.

(C) What are the products formed when ferrous sulphate is heated?
(a) SO3, SO2, FeO
(b) SO3, Fe2, O3
(c) SO3, SO2, FeO3
(d) SO2, FeO
Answer:
(c) SO3, SO2, FeO3
Explanation: When ferrous sulphate crystals are heated, it undergoes decomposition reaction to form ferric oxide, sulphur dioxide and sulphur trioxide.

Equation for the reaction taking place is:
Types Of Chemical Reactions Definitions, Equations and Examples 5

(D) On heating lead nitrate, two gases are evolved, one is colourless and the other is brown in colour. Which gases are they?
(a) Oxygen and nitrogen dioxide
(b) Hydrogen and lead oxide
(c) Hydrogen and oxygen
(d) Oxygen and nitrous oxide
Answer:
(a) Oxygen and nitrogen dioxide Explanation: When lead nitrate is heated, it undergoes decomposition reaction to form lead oxide, along with the evolution of gases nitrogen dioxide and oxygen.

(E) In which of the following category will you put the reaction of heating of ferrous sulphate and lead nitrate?
(I) Decomposition Reaction
(II) Combination Reaction
(III) Endothermic Reaction
(IV) Exothermic Reaction
(a) Only (I)
(b) Only (II)
(c) Both (I) and (III)
(d) Both (II) and (IV)
Answer:
(c) Both (I) and (III)

Explanation: As both ferrous sulphate and lead nitrate absorb heat and undergo decomposition, the reactions are endothermic reactions and are also known as thermal decomposition reactions.

The decomposition reactions that are carried out by using electricity are known as electrolytic decomposition reactions.

For example, the electrolysis of water to form hydrogen and oxygen gas.
Types Of Chemical Reactions Definitions, Equations and Examples 6

Types of Chemical Reactions Definitions, Equations and Examples

Example 3.
Why is the amount of gas collected in one of the test tubes in Activity 1.7 double of the amount collected in the other? Name this gas.
Answer:
When water undergoes electrolysis, it decomposes to form hydrogen gas and oxygen gas as per equation below:
2H2O → 2H2 + O2
As the electrolysis of water produces 2 volumes of hydrogen gas and 1 volume of oxygen gas, it can therefore be concluded that the ratio of hydrogen and oxygen elements in water is 2 :1 by volume.

Decomposition Reaction by Light Energy:

The decomposition reactions that occur in the presence of sunlight are known as photodecomposition reactions or photolysis.
For example, silver chloride (white) decomposes into silver (grey) and chlorine in the presence of sunlight.
Types Of Chemical Reactions Definitions, Equations and Examples 7
A brief summary of some decomposition reactions is given below:

Examples of Decomposition Reaction Chemical Equation
1. Heating of potassium chlorate in presence of manganese dioxide (catalyst) to give potassium chloride and oxygen 2KClO3(s) → 2KCl(S) + 3O2(g)
2.Decomposition of molten sodium chloride on passing electric current to form sodium metal and chlorine gas 2NaCl(l) → 2Na(s) + Cl2(g)
3. Decomposition of hydrogen peroxide (H2O2) in the presence of light to water (H2O) and oxygen gas (O2) 2H2O2 → 2 H2O + O2

Example 4.
Why are decomposition reactions called the opposite of combination reactions? Write equations for decomposition reactions.
Answer:
In a decomposition reaction, a single substance (reactant) splits into two or more products whereas in a combination reaction, two or more reactants combine to form a single product. Thus, decomposition reaction is the opposite of combination reaction.

Consider the reactions:

  • ZnCO3 → ZnO + CO2
  • 2H2 + O2 → 2H2O

Reaction (1) is a decomposition reaction as a single reactant (ZnCO3) splits into two products ZnO and CO2.
Reaction (2) is just the opposite of reaction (1) as two reactants H2 and O2 combine to form a single product (H2O).

Displacement Reaction (A + BC → AC + B):
Those reactions in which one element takes the place of another element in a compound, are known as displacement reactions. In general, a more reactive element displaces a less reactive element from its compound.
Types Of Chemical Reactions Definitions, Equations and Examples 8

In this reaction, iron has displaced or removed another element, copper, from copper sulphate solution. This reaction is known as the displacement reaction.

The reaction between most metals and acids is also a displacement reaction as the metal displaces hydrogen from the acid.
For example. Mg + H2SO4 → MgSO4 + H2

The hydrogen gas is collected by the downward displacement of water as it is lighter than water. Moreover, it is also insoluble in water.

Examples of Displacement Reaction Chemical Equation
1. When a strip of lead metal is placed in a solution of copper chLoride, the green colour of copper chloride solution fades due to the formation of a colourless solution of lead chloride. Lead is able to displace copper from copper chloride solution because lead is more reactive than copper. CuCl2(aq) + Pb(s) → PbCl2(aq) + Cu(s)
2. When iron (III) oxide is heated with aluminium powder, then aluminium oxide and iron metal are formed as a more reactive metal aluminium displaces a less reactive metal iron from its oxide Fe2O3(s) + 2Al(s) → Al2O3(s) + 2Fe(l)

Types of Chemical Reactions Definitions, Equations and Examples

Example 5.
Why does the colour of copper sulphate solution change when an iron nail is dipped in it?
Answer:
When an iron nail is dipped in copper sulphate solution, iron displaces copper from copper sulphate solution as iron is more reactive than copper. Therefore, the colour of the copper sulphate solution changes from blue to green. The reaction taking place is:

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Example 6.
What happens when dilute hydrochloric acid is added to iron filings? Tick the cor¬rect answer.
(a) Hydrogen gas and iron chloride are produced.
(b) Chlorine gas and iron hydroxide are produced.
(c) No reaction takes place.
(d) iron salt and water are produced.
Answer:
(a) Hydrogen gas and iron chloride are produced
Explanation: As iron is more reactive than hydrogen, it displaces chlorine from dil hydrochloric acid and forms hydrogen gas and a salt, iron chloride.

The equation for the reaction taking place is:
2HCl(aq) + Fe(s) → FeCl2(aq) + H2(g)

Double Displacement Reaction (AB + CD → AD + CB)

Those reactions in which there is an exchange of ions between two compounds to form two new compounds are called double displacement reactions. For example, when silver nitrate solution is added to sodium chloride solution, a white precipitate of silver chloride is formed along with sodium nitrate solution.
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

Precipitation Reaction:

Any reaction in which an insoluble solid (called precipitate) is formed that separates from the solution is called a precipitation reaction.
For example, when ammonium hydroxide solution is added to aluminium chloride solution, a white precipitate of aluminium hydroxide: formed along with ammonium chloride solution.
AlCl3(aq) + 3NH4OH(aq) → Al(OH)3(s) + 3NH4Cl(aq)

Neutralization Reactions:

The reactions in which acids or acidic oxides (oxides of non-metals) react with bases or basic oxides (oxides of metals) to form salt and water are called neutralization reactions. For example,
NaOH + HCl → NaCl + H2O

Example of Double Displacement Reaction

Chemical Equation
1. Lead (II) nitrate reacts with sodium chloride to produce sodium nitrate and a white precipitate of lead (II) chloride Pb(N03)2(aq) + 2NaCl(aq) → 2NaN03(aq) + PbCl2(s)
2. Aluminium sulphate reacts with calcium hydroxide solution to produce a white precipitate of aluminium hydroxide and calcium sulphate solution. Al2(S04)3(aq) + 3Ca(OH)2(aq) → 2Al(OH)3(s) +3CaSO4(aq)

Example 7.
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)
Answer:
2KBr(aq) + Bal2(oq) → 2KI(aq) + BaBr2(s)
(It is a double Displacement reaction as there is an exchange of K+ and Ba2+ ions between the reactants)

(B) Zinc carbonate(s) → Zinc oxide(s) + Carbon dioxide(g)
Answer:
ZnCO3(s) → ZnO(s) + CO2(g)
(It is a decomposition reaction as a single reactant splits to form two products)

(C) Hydrogen(g) + Chlorine(g) → Hydrogen chloride(g)
Answer:
H2(g) + Cl2(g) → 2HCl(g)
(It is a combination reaction as two reactants combine to form a single product)

(D) Magnesium(s) + Hydrochloric acid(aq) → Magnesium chloride(aq) + Hydrogen(g)
Answer:
Mg(s) + 2HCl(aq) → MgCl2(oq) + H2(g)
(It is a displacement reaction as Mg displaces hydrogen from HCl as it is more reactive than H)

Types of Chemical Reactions Definitions, Equations and Examples

Example 8.
What is the difference between displacement and double displacement reactions? Write relevant equations for the above.
Answer:
In a displacement reaction a more reactive substance displaces a less reactive substance from its compound or salt solution whereas in a double displacement reaction exchange of ions takes place between the two reactant compounds. In a displacement reaction, only a single displacement takes place whereas in a double displacement reaction, two displacements take place between the molecules.

For example,
In the reaction,
Mg + 2HCl → MgCl2 + H2
Mg displaces hydrogen from HCl as Mg is more reactive than H. It is therefore an example of Displacement reaction In the reaction,
BaCl2(aq) + K2SO4(oq) → 2BaSO4(s) + 2KCl(aq)
Both K and Ba displace each other from their compounds by exchanging ions. It is therefore an example of a double displacement reaction.

Oxidation and Reduction:

Those reactions in which addition of oxygen to a substance or removal of hydrogen from a substance takes place or the loss of electron from an element, are called oxidation reactions.

Those reactions in which addition of hydrogen to a substance or removal of oxygen from a substance takes place or the gain of electron in an element are called reduction reactions.

For example, in the reaction
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O

Hydrochloric acid is oxidized to chlorine (as hydrogen is removed from HCl) and Manganese dioxide is reduced to manganese dichloride (as oxygen is removed from MnO2).

HCl is the reducing agent and MnO2 is the oxidizing agent.
Types Of Chemical Reactions Definitions, Equations and Examples 9
Oxidising Agent
The substance which provides oxygen for oxidation or the substance which removes hydrogen.

Reducing Agent
The substance which gives hydrogen for reduction or the substance which removes oxygen.

Important

  • The substance which gets oxidised is the reducing agent.
  • The substance which gets reduced is the oxidising agent.

In other words, one reactant gets oxidised white the other gets reduced during a reaction. Such reactions are coiled oxidation-reduction reactions or Redax reactions.
Types Of Chemical Reactions Definitions, Equations and Examples 10

Types of Chemical Reactions Definitions, Equations and Examples

Example 9.
Identify the substances that are oxidized and that are reduced in the following equation.
(A) 4Na(s) + O2(g) → 2Na2O(s)
(B) CuO(S) + H2(g) → Cu(s) + H2O(l)
Answer:
In the equation 4Na(S) + O2(g) → 2Na2O(s)
Na is oxidized to Na2O as there is an addition of oxygen and O2 gets reduced.

(B) In the equation,
CuO(s) + H2(g) → Cu(S) + H2O(l)
CuO gets reduced to Cu by losing oxygen and H2 gets oxidized to H2O by gaining oxygen.

Example 10.
Which of the statements about the reaction below are incorrect?
2PbO(s) + C(S) → 2Pb(s) + CO2(g)
(I) Lead is getting reduced
(II) Carbon Dioxide is getting oxidised
(III) Carbon is getting oxidised
(IV) Lead oxide is getting reduced
(a) (I) and (II)
(d) (I) and (III)
(c) (I), (II) and (III)
(e) all
Answer:
(a) (I) and (II)
Explanation: In the given reaction, PbO loses oxygen to form Pb. Therefore, PbO (lead oxide) is reduced.
Moreover, C is oxidized to CO2 by the addition of oxygen. Therefore, (III) and (IV) are correct and (I) and (II) are incorrect.

Class 10 Science Notes

Chemical Equations Definitions, Equations and Examples

Chemical Equations:

Chemical equations are the symbolic representation of chemical reactions where the reactants and products are written in the form of symbols and formulae on the Left hand side and right hand side of an arrow respectively.
Example of a chemical equation:
CaCl2 + 2AgNO3 → Ca(NO3)2 + 2AgCl
In the above equation, the reactants calcium chloride and silver nitrate react to form the products calcium nitrate and silver chloride.

Word Equations:

One way of representing chemicals reactions is word equation which is a short and simplest way of writing chemical reactions as compared to the description of a chemical reaction in a sentence form which is quite long.
The word equation for the reaction between magnesium and oxygen gas would be:
Chemical Equations Definitions, Equations and Examples 1

Representation of Chemical Reactions:

To make chemical equations more concise and useful we use chemical formulae instead of words. A chemical equation represents a chemical reaction. The word equation shown above can be written as:
Mg + O2 → MgO

Chemical Equations Definitions, Equations and Examples

Balanced and Unbalanced Chemical Equations:

A balanced chemical equation has an equal number of atoms of different elements in the reactants and products in accordance with the law of conservation of mass.

Number of atoms of each element in reactant side = Number of atoms of each element in product side
An unbalanced chemical equation has an unequal number of atoms of one or more elements in the reactants and products because the mass is not the same on both sides of the equation. Such a chemical equation is a skeletal chemical equation for a reaction. The equation
Mg + O2 → MgO
is a skeletal chemical equation as the number of atoms of oxygen is not equal in reactant side (2 atoms) and product side (1 atom).

Balanced Chemical Equations:

We need to balance chemical equations to satisfy the law of conservation of mass in chemical reactions. The process of making the number of different types of atoms equal on both the sides of an equation is called balancing of equation.

Hit-and-trial Method of Balancing Chemical Equations
Consider the chemical equation for the reaction between methane and oxygen to yield carbon dioxide and water:
CH4 + O2 → CO2 + H2O
The above equation is an unbalanced equation as the number of atoms of both hydrogen and oxygen are not equal on reactant side and product side.
Step 1: First draw boxes around each formula or put a bracket around each formula. Do not change anything inside the boxes or brackets while balancing the equation.
Chemical Equations Definitions, Equations and Examples 2

Chemical Equations Definitions, Equations and Examples

Step II: List the number of atoms of different elements present in the unbalanced equation
Chemical Equations Definitions, Equations and Examples 3

Step III: Start balancing with the compound that contains the maximum number of atoms. It may be a reactant or a product. In that compound, select the element which has the maximum number of atoms.

Here, we select CH4 and the element hydrogen in it.
Chemical Equations Definitions, Equations and Examples 4
To make the number of atoms of hydrogen equal, we put the coefficient ‘2’ as ‘2H2O’

→ To equalize the number of atoms, we cannot alter the formulae of the compounds or elements involved in the reactions.
Chemical Equations Definitions, Equations and Examples 5
This is a partly balanced equation as the number of oxygen atoms are not balanced.

Step IV: Pick any of those elements whose atoms are still not balanced to proceed further.
As number of oxygen atoms are not equal on reactant and product sides, we will balance O atoms now.
Chemical Equations Definitions, Equations and Examples 6
Step V: Count atoms of each element on both sides of the equation to check the correctness of the balanced equation.
CH4 + 2O2 → CO2 + 2H2O
Chemical Equations Definitions, Equations and Examples 7
The numbers of atoms of elements on both sides of above equation are equal. This equation is now balanced.

Step VI: Writing Symbols of Physical States:
To make a chemical equation more informative, the physical states of the reactants and products are mentioned along with their chemical formulae. For example, the burning of methane in oxygen can be written as:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + Heat

Chemical Equations Definitions, Equations and Examples

Example 1.
Balance the chemical equation Ca + H2O → Ca(OH)2 + H2.
Answer:
Steps to balance the equation are as follows:
Step I: First draw boxes around each formula or put a bracket around each formula.
Ca + H2O → Ca(OH)2 + H2

Step II: List the number of atoms of different elements present in the unbalanced equation
Chemical Equations Definitions, Equations and Examples 11
We note here that both H and O atoms are not equal on both sides.

Step III: We first equalize atoms of H.
Chemical Equations Definitions, Equations and Examples 8
To make the number of atoms of hydrogen equal, we put the coefficient ‘2’ as 2H20
Ca + 2H2O → Ca(OH)2 + H2
By putting the coefficient 2 before H2O, the atoms of both H and O get multiplied by 2.

We now count atoms of each element on both sides of the equation to check the correctness of the equation.
Chemical Equations Definitions, Equations and Examples 9
This is a balanced equation as the number of hydrogen and oxygen atoms are now balanced.

So, the balanced equation is:
Ca + 2H2O → Ca(OH)2 + H2

Example 2.
Balance the following chemical equations.
(A) HNO3 + Ca(OH)2 → Ca(NO3)2 + H2O
Answer:
2HNO3 + 2Ca(OH)2 → Ca(NO3)2 + 2H2O

(B) NaOH + H2SO4 → Na2SO4 + H2O
Answer:
2NaOH + 3H2SO4 → Na2SO4 + 2H2O

(C) NaCl + AgNO3 → AgCl + NaNO3
Answer:
NaCl + AgNO3 → AgCl+NaNO3

(D) BaCl2 + H2SO4 → BaSO4 + HCl
Answer:
BaCl2 + H2SO4 → BaSO4 + 2HCl

Chemical Equations Definitions, Equations and Examples

Example 3.
Write the balanced chemical equations for the following reactions.
(A) Calcium hydroxide + Carbon dioxide → Calcium carbonate + Water
Answer:
Ca(OH)2 + CO2 → CaCO3 + H2O

(B) Zinc + Silver nitrate → Zinc nitrate + Silver
Answer:
Zn + 2AgNO → Zn(NO3)2 + 2Ag

(C) Aluminium + Copper chloride → Aluminium chloride + Copper
Answer:
2Al + 3CuCl2 → 2AlCl3 + 3Cu

(D) Barium chloride + Potassium sulphate → Barium sulphate + Potassium chloride
Answer:
BaCl2 + K2SO4 → BaSO4 + 2KCl

Caution
Students should not alter the formulae of compounds or elements involved in the reactions to equalize the number of atoms while balancing a chemical equation.

For example, for balancing the equation
AlCl3(aq) + NH4OH(oq) → Al(OH)3(s) + NH4Cl(aq)

we observe that number of Chlorine atoms am not equal on both sides. But we cannot balance the equation by writing NH4Cl3 in place of NH4Cl as in
AlCl3(aq) + NH4OH(aq) → Al(OH)3(s) + NH4Cl3(aq).

Instead, we have to balance as:
AlCl3(aq) + 3NH4OH(aq) → Al(OH)3(s) + 3NH4Cl(aq)

We should write the formulae of each compound by using their correct symbols and valencies. If the formula of even a single compound is incorrect, the entire equation will be incorrect.

For example, for writing a balanced equation for the reaction: when copper oxide is heated with magnesium powder, magnesium oxide and copper are formed as magnesium displaces copper from its oxide, if we write formula of copper oxide as CuO2 or Cu20 instead of CuO, the equation will be incorrect.

The correct equation is:
Cu(s) + Mg(s) → MgO(s) + Cu(s)

In order to write and balance the equation correctly, it is important to remember the seven elements that exist in nature as diatomic molecules (H2, N2, O2, F2, Cl2, Br2, and I2)

Whenever word equations are given along with the physical states of reactants and products, the same should also be written in the chemical equations.

Indicate the Physical States of Reactants and Products in an Equation

  1. The Solid-state is indicated by the symbol (s). Formation of precipitate is indicated by writing (s) or downward pointing arrow (l)
  2. Liquid state is indicated by the symbol (l)
  3. Aqueous state (solution made in water) is indicated by the symbol (aq)
  4. Gaseous state is indicated by the symbol (g) or an upward pointing arrow (t).

For example, in the chemical equation
Fe(S) + CuSCO4(aq) → FeSO4(aq)+ Cu

Fe is in solid-state, CuSO4 is in the aqueous state, FeSO4 is in an aqueous state and Cu is in solid state.

Indicate the Heat Changes in an Equation:

  1. An exothermic reaction (reaction in which heat is evolved) is indicated by writing “+ Heat” on the products side of an equation.
  2. An endothermic reaction (reaction in which heat is absorbed) is indicated by writing “+ Heat” on the reactants side of an equation.

Indicate the Conditions under which the Reaction Takes Place
The reaction conditions, such as temperature, pressure, catalyst, etc., for the reaction may be indicated above and/or below the arrow in the equation.

For example, the temperature at which sodium hydrogen carbonate is heated to obtain sodium carbonate, carbon dioxide and water can be represented by the chemical equation,
Chemical Equations Definitions, Equations and Examples 10

Chemical Equations Definitions, Equations and Examples

Example 4.
Write a balanced chemical equation with state symbols for the following reactions.
(A) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride.
Answer:
BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl (aq)

(B) Sodium hydroxide solution (in water) reacts with a hydrochloric acid solution (in water) to produce sodium chloride solution and water.
Answer:
NaOH(aq)) + HCl(aq) → NaCL(oq) + H2O(l)

Example 5.
Translate the following statements into chemical equations and balance them.
(A) Hydrogen gas combines with nitrogen to form ammonia.
Answer:
3H2(g) + N2(g) → 2NH3(g)

(B) Hydrogen sulphide gas burns in air to give water and sulphur dioxide.
Answer:
2H2S(g) + 3O2(g) → 2H2O(l) + 2SO2(g)

(C) Barium chloride reacts with aluminium sulphate to give Aluminium chloride and a precipitate of barium sulphate.
Answer:
3BaCl2(aq) + Al2(SO4)3(aq) → 2AlCl3(oq) + 3BaSO4(S)

(D) Potassium metal reacts with water to give potassium hydroxide and hydrogen gas.
Answer:
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)

Class 10 Science Notes

Chemical Reactions Definitions, Equations and Examples

Chemical Reactions:

We come across a variety of changes that take place around us in our daily life. These changes may be physical changes or chemical changes.

In this chapter we will study chemical reactions and their characteristics, balancing chemical equations, various types of chemical reactions such as combination reactions, decomposition reactions, displacement reactions, double decomposition reactions, exothermic and endothermic reactions, redox reactions and effects of oxidation reactions in our daily life.

Physical changes are the changes which bring about changes in physical characteristics of a substance without changing the chemical composition of the substance.

Examples: Melting of ice to form water, boiling of water to form water vapour, dissolving salt or sugar in water, mixing carbon dioxide gas in soda etc.

Chemical changes are the changes in which new substances are formed having properties entirely different from the original substance(s).

Chemical Reactions Definitions, Equations and Examples

Examples: Cooking of food, digestion, respiration, burning of coal, rusting of iron, souring of milk etc. Chemical reactions are the processes involving a chemical change in which new substances having new properties are formed from original substances.

Chemical reactions involve breaking of old bonds and formation of new bonds.

Reactants and Products:

The substances which take part in a chemical reaction are called reactants while the new substances produced as a result of chemical reaction are called
products.
Examples:
Chemical Reactions Definitions, Equations and Examples 1

Characteristics of Chemical Reactions
A chemical reaction is characterized by any of the following observations:

Change in State:

Consider the reaction between ammonia gas with
hydrogen chloride gas to produce solid ammonium chloride. Change of state takes place in this reaction.
NH3(g) + HCl(g) → NH4Cl(s)

Chemical Reactions Definitions, Equations and Examples

Change in Colour:

In the reaction between silver (silvery white in colour) and hydrogen sulphide to form silver sulphide (black) and hydrogen gas, change of colour takes place.

Evolution of a Gas:

The reaction between zinc and dilsulphuric acid is characterized by the evolution of hydrogen gas.
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)

Change in Temperature:

A chemical reaction can be exothermic or endothermic in nature.
The reaction between zinc and dil. sulphuric acid to form zinc sulphate and hydrogen gas is an exothermic reaction as heat is evolved.
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) + Heat
However, the reaction between ammonium chloride and Barium hydroxide is an endothermic reaction as heat is absorbed.
Ba(OH)2(aq) + 2NH4Cl(s) + Heat → BaCl2(aq) + 2NH3 + 2H2O(l)

Formation of a Precipitate

When barium chloride solution reacts with sulphuric acid, barium sulphate is formed which is an insoluble white substance also known as precipitate, along with hydrochloric acid.
BaCl2(oq) + H2SO4(oq) → BaSO4(s) ↓(White ppt) + 2HCl(aq)

Example 1.
Case Based:
A 2 cm long thin ribbon of a metal X was taken and first cleaned with sandpaper. It was then burnt using a spirit lamp or burner by holding it with a pair of tongs. The ribbon burnt with a dazzling white flame and formed a powder Y which was collected in a watch glass.

(A) Which option correctly identifies both X and Y?
Chemical Reactions Definitions, Equations and Examples 2
Answer:
(c) X is magnesium and Y is magnesium oxide.
Explanation: When magnesium ribbon is burnt, it forms a white powder of magnesium oxide (MgO) due to the reaction between magnesium and oxygen present in the air.
2Mg + O2 → 2MgO

(B) The colour of powder or ash formed when a magnesium ribbon is burnt in the air is:
(a) grey
(b) black
(c) white
(d) yellow.
Answer:
(c) white
Explanation: When a magnesium ribbon is burnt in air, it forms magnesium oxide, which is a white coloured powder.

(C) Why should a magnesium ribbon be cleaned before burning in air?
Answer:
Magnesium ribbon should be cleaned before burning in air to remove the layer of magnesium oxide that may have formed on the ribbon due to the reaction of magnesium with oxygen.

Chemical Reactions Definitions, Equations and Examples

(D) Is burning of magnesium ribbon a physical change or a chemical change? Justify your answer.
Answer:
Burning of ribbon is a chemical change as magnesium burns in air to form a new substance magnesium oxide. Moreover, it also involves change in temperature as a lot of heat and light is produced during this change.

(E) Assertion: Magnesium ribbon burns with a dazzling white flame.
Reason: When magnesium ribbon burns in air only heat is evolved.
(a) Both (A) and (R) are true, and (R) is correct explanation of the assertion.
(b) Both (A) and (R) are true, but (R) is not the correct explanation of the assertion.
(c) (A) is true, but (R) is false.
(d) (A) is false, but (R) is true.
Answer:
(c) (A) is true, but (R) is false.

Explanation: Magnesium ribbon burns in presence of air to form magnesium oxide. A Large amount of heat and light are produced due to which we see a dazzling white flame.

Class 10 Science Notes