CBSE NOTES CLASS 10 SCIENCE CHAPTER 1
Chemical Reactions And Equations
A chemical change or reaction involves formation of one or more new substance.
We say that a chemical reaction has taken place if one of the following has taken place.
- change in state
- change in colour
- evolution of a gas
- change in temperature.
Examples of chemical reactions
- milk turns sour if left at room temperature during summers.
- an iron tawa/pan/nail gets rusted if it is left exposed to humid atmosphere.
- grapes get fermented.
- food is cooked.
- food gets digested in our body.
- we respire.
- Formation of hydrogen gas by the action of dilute sulphuric acid on zinc
- Burning of magnesium ribbon in air, giving dazzling white light.
[Magnesium ribbon must be cleaned before burning in air so that the layer of magnesium oxide (which is formed due to reaction of magnesium with air) can be removed in order to get the desired chemical reaction.]
The description of a chemical reaction can be done in a sentence form which is quite long.
It can also be written as word-equation.
The substances that undergo chemical change in the reaction, for example, magnesium and oxygen, are called reactants.
The new substances formed during the reaction, for example magnesium oxide, are called the products.
A word-equation shows change of reactants to products by placing an arrow between them.
Symbolic Form of Chemical Equation
A chemical equation is the symbolic representation of a chemical reaction in the form of symbols and formulae.
The reactants are written on the left-hand side (LHS) separated by plus sign (+) from each other.
Products are written on the right-hand side (RHS) separated by plus sign (+) from each other.
The arrowhead points towards the products, and shows the direction of the reaction.
Balanced Chemical Equation
The chemical equation needs to be balanced so that it follows the law of conservation of mass.
The total mass of the elements present in the products of a chemical reaction has to be equal to the total mass of the elements present in the reactants.
This means that the number of atoms of each element remains the same, before and after a chemical reaction.
For example the word-equation
Zinc + Sulphuric acid → Zinc sulphate + Hydrogen
May be represented by chemical equation,
Zn + H2SO4 → ZnSO4 + H2
Balancing Chemical Equations
Let us try to balance the following chemical equation –
Fe + H2O → Fe3O4 + H2
Step I: Draw boxes around each formula. Do not change anything inside the boxes while balancing the equation.
Fe + H2O → Fe3O4 + H2
Step II: List the number of atoms of different elements present in the unbalanced equation
Number of atoms in reactants (LHS)
Number of atoms
in products (RHS)
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. Hence we select Fe3O4 and the element oxygen in it. There are four oxygen atoms on the RHS and only one on the LHS.
To balance the oxygen atoms – we need to multiply H2O by 4. We cannot alter the formulae of the compounds or elements involved in the reactions (within the boxes)
Fe + 4 H2O → Fe3O4 + H2
Step IV: Fe and H atoms are still not balanced. Let us balance hydrogen atoms. There are 8 atoms of H on the LHS (in H2O) and 2 on the RHS (in H2). To equalise the number of H atoms, make the number of molecules of hydrogen as four on the RHS.
Fe + 4 H2O → Fe3O4 + 4 H2
Step V: There are now 1 Fe atom on LHS and 3 on the RHS. Multiply Fe on the left by 3 to balance this.
4 Fe + 4 H2O → Fe3O4 + 4 H2
Step VI: Finally, to check the correctness of the balanced equation, the balanced equation is,
3 Fe + 4 H2O → Fe3O4 + 4 H2
This method of balancing chemical equations is called hit-and-trial method as we make trials to balance the equation by using the smallest whole number coefficient.
Step VII: 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. The gaseous, liquid, aqueous and solid states of reactants and products are represented by the notations (g), (l), (aq) and (s), respectively. The word aqueous (aq) is written if the reactant or product is present as a solution in water. Usually physical states are not included in a chemical equation unless it is necessary to specify them.
The reaction conditions, such as temperature, pressure, catalyst, etc., for the reaction are indicated above and/or below the arrow in the equation. For example –
Precautions to be taken while doing expermients
(i) Do not hold the test tube by hand. Use tongs for the purpose.
(ii) Do not point the mouth of the boiling tube at yourself or at your neighbours.
(iii) Do not smell the gas directly. Waft the gas gently towards the nose.
(iv) Do not touch the chemicals by hand. Use spoon or spatula.
(v) Do not heat the volatile substances, ike alcohol, directly over burner. Use water bath instead.
(vi) Do not add water to acid or base. Add acid or to base slowly to the water, with constant stirring for diluting to desired level. This is because the dissolution of acid or base in water is highly exhothermic. The heat evolved may cause explosion or may break the glass tube, which in turn may cause bodily harm.
Types of Chemical Reactions
A reaction in which a single product is formed from two or more reactants is known as a combination reaction.
For example calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide) releasing a large amount of heat.
A solution of slaked lime is used for white washing walls. Calcium hydroxide reacts slowly with the carbon dioxide in air to form a thin layer of calcium carbonate on the walls. Calcium carbonate is formed after two to three days of white washing and gives a shiny finish to the walls. The chemical formula for marble is also CaCO3.
More examples of combination reactions
(i) Burning of coal
C (s) + O2 (g) → CO2 (g)
(ii) Formation of water from H2 (g) and O2 (g)
2H2 (g) + O2 (g) → 2H2O (l)
Exothermic Reactions: Reactions in which heat is released along with the formation of products are called exothermic chemical reactions.
- Combination reactions are mostly exothermic.
- All the above reactions are exothermic.
- Respiration is an exothermic process. In this glucose combines with oxygen in the cells of our body and provides energy.
- The decomposition of vegetable matter into compost is also an example of an exothermic reaction.
The reaction, in which one compound decomposes or breaks into two or more different products, is called decomposition reaction.
- When ferrous sulphate crystals are heated, the green colour changes to yellowish brown.
- Decomposition of calcium carbonate to calcium oxide and carbon dioxide takes place on heating. Calcium oxide is called lime or quick lime.
- Heating of lead nitrate results in emission of nitrogen dioxide (brown fumes),
- Thermal Decomposition -When a decomposition reaction is carried out by heating, it is called thermal decomposition.
- Endothermic Reaction – A reaction in which energy is absorbed is known as endothermic reaction. Decomposition reactions are mostly endothermic.
- Electrolytic Decomposition of Water
Electrolysis – Passing of electric current through acidified water to decompose it.
An electrode is an electrical conductor used to make contact with a nonmetallic part of a circuit (e.g. a semiconductor, an electrolyte, a vacuum or air)
Cathode: The negatively charged electrode by which electrons enter an electrical device or the positively charged electrode of an electrical device, such as a primary cell, that supplies current.
Anode: The positively charged electrode by which the electrons leave an electrical device or the negatively charged electrode of an electrical device, such as a primary cell, that supplies current.
Cation: A positively charged ion, i.e. one that would be attracted to the cathode in electrolysis.
Anion: A negatively charged ion, i.e. one that would be attracted to the anode in electrolysis.
Activity to demonstrate electrolysis of water
- Take a plastic mug. Drill two holes at its base and fit rubber stoppers in these holes.
- Insert carbon electrodes in these rubber stoppers as shown in figure.
- Connect these electrodes to a 6 volt battery.
- Fill the mug with water such that the electrodes are immersed. Add a few drops of dilute sulphuric acid to the water.
Water in its pure form does not conduct electricity. To make the water conducting, we add acid to it.
- Take two test tubes filled with water and invert them over the two carbon electrodes.
- Switch on the current and leave the apparatus undisturbed for some time.
- Observation: We observe the formation of bubbles at both the electrodes. These bubbles displace water in the test tubes.
- The volume of the gas collected in the test tube at cathode is double that in the test tube at anode.
Testing the evolved gases:
Once the test tubes are filled with the respective gases, remove them carefully.
Gas at anode: If we bring a burning candle close to the mouth of the test tube, it starts burning even more brightly. The gas is OXYGEN.
Gas at cathode: If we bring a burning candle close to the mouth of the test tube, it gets extinguished and the gas burns with a distinctive 'squeaky pop' sound. The gas is HYDROGEN.
The volume of hydrogen liberated is double that of oxygen.
Reaction in the solution
H2O (l) H+ (aq) + OH- (aq)
Reaction at Cathode (reduction)
H+ (aq) + 2e− → H2 (g)
Reaction at Anode (oxidation)
4OH− (aq) → O2 (g) + 2H2O (l) + 4e−
- White silver chloride turns grey in sunlight due to the decomposition into silver and chlorine.
2AgCl (s) (white) 2Ag (s) (Grey) + Cl2 (g) (Green gas)
2AgBr (s) (Pale yellow) 2Ag (s) + Br2 (g) (Brown fumes)
The above reactions are used in black and white photography.
Displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its compound. (Remember the reactivity series)
- Zinc and lead are more reactive elements than copper. They displace copper from its compounds.
Double Displacement Reaction
Reactions in which there is an exchange of ions between the reactants are called double displacement reactions.
A substance, which is insoluble in water, is known as a precipitate. Any reaction that produces a precipitate can be called a precipitation reaction.
Oxidation and Reduction
If a substance gains oxygen or losses hydrogen during a reaction, it is said to be oxidized.
2Cu (reddish brown) + O2 2CuO (black)
If a substance loses oxygen or gains hydrogen during a reaction, it is said to be reduced.
CuO + H2 Cu + H2O
One reactant gets oxidised while the other gets reduced simultaneously during a reaction. Such reactions are called oxidation-reduction reactions or redox reactions.
Few Other Examples:
ZnO + C → Zn +CO
MnO2 + 4HCl → MnCl2 + 2H2O + Cl2
CorrosionIt is the process in which metals are slowly eaten up by the action of air, moisture or chemicals.
- For example rusting is a form of corrosion in which iron is eaten up by the action of air and moisture and a reddish brown coating of iron oxide, called rust, is formed as shown in the following chemical reaction.
3Fe + 4H2O → F3O4 + H2
- Silver articles become black after some time when exposed to air. This is because it reacts with sulphur in the air to form a coating of silver sulphide.
- Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is copper carbonate.
- Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of metals, especially those of iron. Corrosion of iron is a serious problem. Every year an enormous amount of money is spent to replace damaged iron.
Methods of protecting the metals from corrosion
- Painting, oiling, greasing, galvanising, Electroplating (e.g., chrome plating), anodising or making alloys are methods of protecting metals from corrosion.
- Galvanisation is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc.
The galvanised article is protected against rusting even if the zinc coating is broken, how?
The outer layer of zinc of any galvanized material reacts with the atmospheric oxygen to form Zinc Oxide (ZnO), which is stronger than Zinc. Thus, even if the outer layer of zinc undergoes corrosion, the material is getting coated with a stronger substance (ZnO), and thus is better able to resist corrosion.
- Alloying is a very good method of improving the properties of a metal.
The properties of any metal can be changed if it is mixed with some other substance.
The substance added may be a metal or a non-metal.
An alloy is a homogeneous mixture of two or more metals, or a metal and a nonmetal.
It is prepared by first melting the primary metal, and then, dissolving the other elements in it in definite proportions. It is then cooled to room temperature.
For example, iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretches easily when hot. But, if it is mixed with a small amount of carbon (about 0.05%), it becomes hard and strong. When iron is mixed with nickel and chromium, we get stainless steel, which is hard and does not rust. Thus, if iron is mixed with some other substance, its properties change.
When the substances containing oils and fats are exposed to air they get oxidized and become rancid due to which their smell, taste and colour change. This process is known as rancidity. For example when when butter is kept open for a long time, its smell and taste gets changed.
- Antioxidants are added to foods containing fats and oil.
- Keeping food in air tight containers helps to slow down oxidation.
- Bags of patoto chips etc. are flushed with gas such as nitrogen to prevent the chips from getting oxidized.