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1. Chemical Reactions and Equations
Consider the following situations of daily life and think what happens when –
- milk is left at room temperature during summers.
- an iron tawa/pan/nail is left exposed to humid atmosphere.
- grapes get fermented.
- food is cooked.
- food gets digested in our body.
- we respire.
In all the above situations, the nature and the identity of the initial substance have somewhat changed. We have already learnt about physical and chemical changes of matter in our previous classes. Whenever a chemical change occurs, we can say that a chemical reaction has taken place.
You may perhaps be wondering as to what is actually meant by a chemical reaction. How do we come to know that a chemical reaction has taken place? Let us perform some activities to find the answer to these questions.
Figure 1.1 Burning of a magnesium ribbon in air and collection of magnesium oxide in a watch-glass
You must have observed that magnesium ribbon burns with a dazzling white flame and changes into a white powder. This powder is magnesium oxide. It is formed due to the reaction between magnesium and oxygen present in the air.
Figure 1.2 Formation of hydrogen gas by the action of dilute sulphuric acid on zinc
From the above three activities, we can say that any of the following observations helps us to determine whether a chemical reaction has taken place –
- change in state
- change in colour
- evolution of a gas
- change in temperature.
As we observe the changes around us, we can see that there is a large variety of chemical reactions taking place around us. We will study about the various types of chemical reactions and their symbolic representation in this Chapter.
1.1 CHEMICAL EQUATIONS
Activity 1.1 can be described as – when a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide. This description of a chemical reaction in a sentence form is quite long. It can be written in a shorter form. The simplest way to do this is to write it in the form of a word-equation.
The word-equation for the above reaction would be –
The substances that undergo chemical change in the reaction (1.1), magnesium and oxygen, are the reactants. The new substance, magnesium oxide, formed during the reaction, is the product.
A word-equation shows change of reactants to products through an arrow placed between them. The reactants are written on the left-hand side (LHS) with a plus sign (+) between them. Similarly, products are written on the right-hand side (RHS) with a plus sign (+) between them. The arrowhead points towards the products, and shows the direction of the reaction.
1.1.1 Writing a Chemical Equation
Is there any other shorter way for representing chemical equations? Chemical equations can be made more concise and useful if we use chemical formulae instead of words. A chemical equation represents a chemical reaction. If you recall formulae of magnesium, oxygen and magnesium oxide, the above word-equation can be written as –
Count and compare the number of atoms of each element on the LHS and RHS of the arrow. Is the number of atoms of each element the same on both the sides? If not, then the equation is unbalanced 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. Equation (1.2) is a skeletal chemical equation for the burning of magnesium in air.
1.1.2 Balanced Chemical Equations
Recall the law of conservation of mass that you studied in Class IX; mass can neither be created nor destroyed in a chemical reaction. That is, 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.
In other words, the number of atoms of each element remains the same, before and after a chemical reaction. Hence, we need to balance a skeletal chemical equation. Is the chemical Eq. (1.2) balanced? Let us learn about balancing a chemical equation step by step.
The word-equation for Activity 1.3 may be represented as –
The above word-equation may be represented by the following chemical equation –
Let us examine the number of atoms of different elements on both sides of the arrow.
|Element||Number of atoms in reactants (LHS)||Number of atoms in products (RHS)|
As the number of atoms of each element is the same on both sides of the arrow, Eq. (1.3) is a balanced chemical equation.
Let us try to balance the following chemical equation –
Step I: To balance a chemical equation, first draw boxes around each formula. Do not change anything inside the boxes while balancing the equation.
Step II: List the number of atoms of different elements present in the unbalanced equation (1.5).
|Element||Number of atoms in reactants (LHS)||Number of atoms in products (RHS)|
Step III: It is often convenient to 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. Using these criteria, 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 –
|Atoms of oxygen||In reactants||In products|
|Initial||1 (in H2O)||4 (in Fe3O4)|
To equalise the number of atoms, it must be remembered that we cannot alter the formulae of the compounds or elements involved in the reactions. For example, to balance oxygen atoms we can put coefficient ‘4’ as 4 H2O and not H2O4 or (H2O)4. Now the partly balanced equation becomes –
(partly balanced equation)
Step IV: Fe and H atoms are still not balanced. Pick any of these elements to proceed further. Let us balance hydrogen atoms in the partly balanced equation.
To equalize the number of H atoms, make the number of molecules of hydrogen as four on the RHS.
|Atoms of Hydrogen||In reactants||In products|
|Initial||8(in 4 H2O)||2(in H2)|
The equation would be –
(partly balanced equation)
Step V: Examine the above equation and pick up the third element which is not balanced. You find that only one element is left to be balanced, that is, iron.
|Atoms of Iron||In reactants||In products|
|Initial||1(in Fe)||3(in Fe3O4)|
To equalize Fe, we take three atoms of Fe on the LHS.
Step VI: Finally, to check the correctness of the balanced equation, we count atoms of each element on both sides of the equation.
The numbers of atoms of elements on both sides of Eq. (1.9) are equal. This equation is now balanced. 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 Carefully examine the above balanced Eq. (1.9). Does this equation tell us anything about the physical state of each reactant and product? No information has been given in this equation about their 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.
The balanced Eq. (1.9) becomes
Note that the symbol (g) is used with H2O to indicate that in this reaction water is used in the form of steam.
Usually physical states are not included in a chemical equation unless it is necessary to specify them.
Sometimes 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 –
Using these steps, can you balance Eq. (1.2) given in the text earlier?
- Why should a magnesium ribbon be cleaned before burning in air?
- Write the balanced equation for the following chemical reactions.
- Write a balanced chemical equation with state symbols for the following reactions.
(i) Hydrogen + Chlorine → Hydrogen chloride
(ii) Barium chloride + Aluminium sulphate → Barium sulphate +Aluminium chloride
(iii) Sodium + Water → Sodium hydroxide + Hydrogen
(i) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride.
(ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water.
1.2 TYPES OF CHEMICAL REACTIONS
We have learnt in Class IX that during a chemical reaction atoms of one element do not change into those of another element. Nor do atoms disappear from the mixture or appear from elsewhere. Actually, chemical reactions involve the breaking and making of bonds between atoms to produce new substances. You will study about types of bonds formed between atoms in Chapters 3 and 4.
1.2.1 Combination Reaction
Figure 1.3 Formation of slaked lime by the reaction of calcium oxide with water
Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide) releasing a large amount of heat.
(Quick lime) (Slaked lime)
In this reaction, calcium oxide and water combine to form a single product, calcium hydroxide. Such a reaction in which a single product is formed from two or more reactants is known as a combination reaction.
Let us discuss some more examples of combination reactions.
(i) Burning of coal
(ii) Formation of water from H2(g) and O2(g)
In simple language we can say that when two or more substances (elements or compounds) combine to form a single product, the reactions are called combination reactions.
In Activity 1.4, we also observed that a large amount of heat is evolved. This makes the reaction mixture warm. Reactions in which heat is released along with the formation of products are called exothermic chemical reactions.
Other examples of exothermic reactions are –
(i) Burning of natural gas
(ii) Do you know that respiration is an exothermic process?
We all know that we need energy to stay alive. We get this energy from the food we eat. During digestion, food is broken down into simpler substances. For example, rice, potatoes and bread contain carbohydrates. These carbohydrates are broken down to form glucose. This glucose combines with oxygen in the cells of our body and provides energy. The special name of this reaction is respiration, the process of which you will study in Chapter 6.
(iii) The decomposition of vegetable matter into compost is also an example of an exothermic reaction.
Identify the type of the reaction taking place in Activity 1.1, where heat is given out along with the formation of a single product.
1.2.2 Decomposition Reaction
Figure 1.4 Correct way of heating the boiling tube containing crystals of ferrous sulphate and of smelling the odour
Have you noticed that the green colour of the ferrous sulphate crystals has changed? You can also smell the characteristic odour of burning sulphur.
In this reaction you can observe that a single reactant breaks down to give simpler products. This is a decomposition reaction. Ferrous sulphate crystals (FeSO4, 7H2O) lose water when heated and the colour of the crystals changes. It then decomposes to ferric oxide (Fe2O3), sulphur dioxide (SO2) and sulphur trioxide (SO3). Ferric oxide is a solid, while SO2 and SO3 are gases.
Decomposition of calcium carbonate to calcium oxide and carbon dioxide on heating is an important decomposition reaction used in various industries. Calcium oxide is called lime or quick lime. It has many uses – one is in the manufacture of cement. When a decomposition reaction is carried out by heating, it is called thermal decomposition.
Another example of a thermal decomposition reaction is given in Activity 1.6.
Figure 1.5 Heating of lead nitrate and emission of nitrogen dioxide
You will observe the emission of brown fumes. These fumes are of nitrogen dioxide (NO2). The reaction that takes place is –
Let us perform some more decomposition reactions as given in Activities 1.7 and 1.8.
Figure 1.6 Electrolysis of water
Figure 1.7 Silver chloride turns grey in sunlight to form silver metal
You will see that white silver chloride turns grey in sunlight. This is due to the decomposition of silver chloride into silver and chlorine by light.
Silver bromide also behaves in the same way.
The above reactions are used in black and white photography.
What form of energy is causing these decomposition reactions ?
We have seen that the decomposition reactions require energy either in the form of heat, light or electricity for breaking down the reactants. Reactions in which energy is absorbed are known as endothermic reactions.
- A solution of a substance ‘X’ is used for white washing.
(i) Name the substance ‘X’ and write its formula.
(ii) Write the reaction of the substance ‘X’ named in (i) above with water.
1.2.3 Displacement Reaction
Figure 1.8 (a) Iron nails dipped in copper sulphate solution
Figure 1.8 (b) Iron nails and copper sulphate solutions compared before and after the experiment
Why does the iron nail become brownish in colour and the blue colour of copper sulphate solution fade?
The following chemical reaction takes place in this Activity–
(Copper sulphate) (Iron sulphate)
In this reaction, iron has displaced or removed another element, copper, from copper sulphate solution. This reaction is known as displacement reaction.
Other examples of displacement reactions are
(Copper sulphate) (Zinc sulphate)
(Copper chloride) (Lead chloride)
Zinc and lead are more reactive elements than copper. They displace copper from its compounds.
1.2.4 Double Displacement Reaction
Figure 1.9 Formation of barium sulphate and sodium chloride
You will observe that a white substance, which is insoluble in water, is formed. This insoluble substance formed is known as a precipitate. Any reaction that produces a precipitate can be called a precipitation reaction.
What causes this? The white precipitate of BaSO4 is formed by the reaction of SO42-and Ba2+. The other product formed is sodium chloride which remains in the solution. Such reactions in which there is an exchange of ions between the reactants are called double displacement reactions.
1.2.5 Oxidation and Reduction
Figure 1.10 Oxidation of copper to copper oxide
The surface of copper powder becomes coated with black copper(II) oxide. Why has this black substance formed?
This is because oxygen is added to copper and copper oxide is formed.
If hydrogen gas is passed over this heated material (CuO), the black coating on the surface turns brown as the reverse reaction takes place and copper is obtained.
If a substance gains oxygen during a reaction, it is said to be oxidised. If a substance loses oxygen during a reaction, it is said to be reduced.
During this reaction (1.29), the copper(II) oxide is losing oxygen and is being reduced. The hydrogen is gaining oxygen and is being oxidised. In other words, one reactant gets oxidised while the other gets reduced during a reaction. Such reactions are called oxidation-reduction reactions or redox reactions.
Some other examples of redox reactions are:
In reaction (1.31) carbon is oxidised to CO and ZnO is reduced to Zn.
In reaction (1.32) HCl is oxidised to Cl2 whereas MnO2 is reduced to MnCl2.
From the above examples we can say that if a substance gains oxygen or loses hydrogen during a reaction, it is oxidised. If a substance loses oxygen or gains hydrogen during a reaction, it is reduced.
1.3 HAVE YOU OBSERVED THE EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE?
You must have observed that iron articles are shiny when new, but get coated with a reddish brown powder when left for some time. This process is commonly known as rusting of iron. Some other metals also get tarnished in this manner. Have you noticed the colour of the coating formed on copper and silver? When a metal is attacked by substances around it such as moisture, acids, etc., it is said to corrode and this process is called corrosion. The black coating on silver and the green coating on copper are other examples of corrosion.
Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of metals, specially those of iron. Corrosion of iron is a serious problem. Every year an enormous amount of money is spent to replace damaged iron. You will learn more about corrosion in Chapter 3.
Have you ever tasted or smelt the fat/oil containing food materials left for a long time?
When fats and oils are oxidised, they become rancid and their smell and taste change. Usually substances which prevent oxidation (antioxidants) are added to foods containing fats and oil. Keeping food in air tight containers helps to slow down oxidation. Do you know that chips manufacturers usually flush bags of chips with gas such as nitrogen to prevent the chips from getting oxidised ?
- Why does the colour of copper sulphate solution change when an iron nail is dipped in it?
- Give an example of a double displacement reaction other than the one given in Activity 1.10.
- Identify the substances that are oxidised and the substances that are reduced in the following reactions.
(i) 4Na(s) + O2(g) → 2Na2O(s)
(ii) CuO(s) + H2(g) → Cu(s) + H2O(l)