|Chemical Reactions and Equations|
Chemistry is defined as that branch of science which deals with the composition and properties of matter and the changes that matter undergoes by various interactions.
A chemical compound is formed as a result of a chemical change and in this process different types of energies such as heat, electrical energy, radiation etc. are either absorbed or evolved. The total mass of the substance remains the same throughout the chemical change.
1.1 CHANGE IN MATTER
Change in matter can be studied in two major ways. The two types of changes are as follows:
1.2 DIFFERENCE BETWEEN PHYSICAL CHANGE AND CHEMICAL CHANGE
|Physical Change||Chemical Change|
|(i) Those changes in which no new substances are formed are called physical changes.||(i) Those changes in which the original substances lose their chemical nature and identity and form new chemical substances with different properties are called chemical changes.|
|(ii) It is a temporary change||(ii) It is a permanent change.|
|(iii) It is easily reversible||(iii) It is usually irreversible.|
|(iv) In a physical change the mass of substance does not alter.||(iv) In a chemical change the mass of the substance does alter.|
The process involving a chemical change is called a chemical reaction.
A chemical reaction is a process which transforms one or more substances into new substances.
The process in which a substance or substances undergo change, to produce new substances with new properties, is known as chemical reaction.
Reactants: The substances which take part in a chemical reaction are called reactants.
Products: The new substances formed as a result of chemical reaction are called products.
(Reactant) (Reactant) (Product)
In the above chemical reaction hydrogen and oxygen which are written on the left hand side are reactants and water which is written on the right hand side is a product.
|Identify the reactant and product in the following reaction|
2.1 EXAMPLES OF SOME CHEMICAL REACTIONS
(i) The burning of magnesium in air to form magnesium oxide
Take a magnesium ribbon and clear it by rubbing with a sand paper. Hold it with a pair of tongs. Burn it using a spirit lamp or burner and collect the white ash so formed in a watch-glass as shown in figure.
Burning of a magnesium ribbon in air and collection of magnesium oxide in a watch-glass
Magnesium ribbon burns with dazzling light and a white substance is formed which is magnesium oxide. This happens due to the following chemical reaction
2 Mg (s) + O2 (g) 2 MgO (s)
Magnesium Oxygen (from air) Magnesium oxide
Thus a chemical reaction has taken place in which magnesium has combined with oxygen of the air to form a new chemical substance, magnesium oxide (MgO). Here Mg and O2 are reactants and MgO formed is product.
| A magnesium ribbon is cleaned to remove the protective layer of basic magnesium carbonate from its surface so that it may readily combine with the oxygen of air (on heating).|
(ii) Reaction between lead nitrate and potassium iodide.
Take lead nitrate solution in a test tube and add some potassium iodide solution to this.
A yellow solid namely lead iodide in the form of precipitate appears. Another substance namely potassium nitrate is also formed which we cannot see as it remains in the solution.
This happens due to the following chemical reaction.
In this reaction lead nitrate and potassium iodide are the reactants while lead iodide and potassium nitrate are the products.
(iii) Reaction between zinc and dilute sulphuric acid (or hydrochloric acid)
| Take a few zinc granules in a conical flask and add some dilute hydrochloric acid (HCl) or sulphuric acid (H2SO4) to this. A gas is evolved very briskly. If we touch the flask, it is found to be hot.
On bringing a lighted candle near the upper end of the tube fitted in the flask (figure 2) the gas burns with a popping sound. This confirms that the gas evolved is hydrogen.
|Reaction between zinc and dilute sulphuric acid (or hydrochloric acid) giving out hydrogen gas
During this reaction, zinc sulphate or zinc chloride is also formed which we cannot see as it remains in the solution. The following chemical reaction takes place.
Zn + 2HCl ZnCl2 + H2
With H2SO4 :
Zn + H2SO4 ZnSO4 + H2
In this chemical reaction Zinc and hydrochloric acid (or sulphuric acid) are the reactants while hydrogen gas and zinc chloride (or zinc sulphate) are the products
2.2 CHARACTERISTICS OF CHEMICAL REACTIONS
The easily observable changes which take place as a result of chemical reactions are known as characteristics of chemical reactions.
When a chemical reaction takes place, any one or more of the above changes or characteristics are observed. Let us discuss these characteristics in detail.
(i) Change in state (formation of precipitate): In the above discussed chemical reactions, in the burning of magnesium ribbon, white powder of magnesium oxide is formed. In the chemical reaction between lead nitrate and potassium iodide yellow precipitate of lead iodide is formed. Similarly when barium chloride solution reacts with sodium sulphate solution, a white precipitate of barium sulphate is obtained. Similarly in the chemical reaction of silver nitrate solution with sodium chloride solution, white precipitate of silver chloride is obtained.
(ii) Change in colour : Some chemical reactions are characterized by change in colour
When citric acid of citrus fruits (lemon and orange), reacts with potassium permagnate solution (purple), the colour changes and a colourless solution is produced.
As the purple colour of potassium permagnate solution changes to colourless by addition of citric acid (lemon juice) it is an example of CHANGE IN COLOUR.
When sulphur dioxide gas reacts with acidified pot. dichromate solution, the orange colour of pot. dichromate changes to green.
When ferrous sulphate crystals are heated, they lose their green colour and become white.
|What changes would you expect to observe in the reaction when iron nail is put in copper sulphate solution :
(a) Change in colour of iron nail (b) Change in solution colour
(c) Both A and B (d) None of these
(iii) Evolution of Gas: Some chemical reactions, when take place, a gas is evolved.
In chemical reaction between zinc and dilute hydrochloric acid or dilute sulphuric acid, hydrogen gas is evolved.
i.e. Zn + HCl ZnCl2 + H2
Zn + H2SO4 ZnSO4 + H2
Metals like sodium and calcium react with water and hydrogen gas is evolved
2Na + 2H2O 2 NaOH + H2
Ca + 2H2O Ca(OH)2 + H2
In the chemical reaction of calcium carbonate with dilute hydrochloric acid, carbon dioxide gas is evolved.
CaCO3 + 2HCl CaCl2 + CO2 + H2O
On heating calcium carbonate, carbon dioxide gas is evolved
(iv) Change in Temperature : Some chemical reactions are known by a change in temperature. This change in temperature may be rise in temperature or fall in temperature. For example
When zinc pieces react with dilute hydrochloric acid or dilute sulphuric acid in a flask, it becomes hot (rise in temperature).
Zn + 2HCl ZnCl2 + H2 + heat
Zn + H2SO4 ZnSO4 + H2 + heat
In the reaction of quick lime with water, a lot of heat is produced. Which causes a rise in its temperature.
CaO + H2O Ca(OH)2 + Heat.
In the chemical reaction of barium hydroxide [Ba(OH)2] with ammonium chloride (NH4Cl) to form barium chloride (BaCl2), ammonia (NH3) and water, a lot of heat energy is absorbed due to which the temperature of the mixture falls and it becomes very cold.
Ba(OH)2 + NH4Cl + Heat BaCl2 + NH3 + H2O
A chemical reaction is said to be ……………………………… if it absorbs energy.
A chemical equation is a symbolic representation of an actual chemical change.
The short-hand method of representing a chemical reaction in terms of symbols and formulae of the different reactants and products is called a chemical equation.
A chemical reaction can be represented in two different ways :
3.1 STEPS FOR WRITING A CHEMICAL EQUATION
Writing of a chemical equation involves the following steps :
(i) The symbols and formulae of the reactants are written on the left hand side with plus (+) sign between them.
(ii) The symbols and formulae of the products are written on the right hand side with + sign between them.
(iii) An arrow sign (®) is put between the reactants and the products, pointing from reactants towards products.
3.2 BALANCED AND UN BALANCED CHEMICAL EQUATIONS
(i) Balanced chemical equation
The equation in which the number of atoms of each element in the reactants, and the products sides are equal, is called a balanced chemical equation.
The chemical equations are balanced to satisfy the law of conservation of mass in chemical reactions.
| Law of conservation of mass states that, 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.|
For example, in a chemical reaction between zinc and dilute sulphuric acid, giving zinc sulphate and hydrogen. The chemical equation can be written as
Zn + H2SO4 ZnSO4 + H2.
Here the No. of atoms of each element in the reactant and products sides are equal. i.e.
In reactants In products
No. of Zn atoms 1 1
No. of H atoms 2 2
No. of S atoms 1 1
No. of O atoms 4 4
Hence it is a balanced chemical equation.
(ii) Unbalanced chemical equation (skeletal equation)
The equation in which the number of atoms of different elements on the reactants and the product sides are not equal is called an unbalanced chemical equation. The unbalanced chemical equation is also known as skeletal equation.
The burning of aluminium in oxygen to form aluminium oxide can be written as :
Al + O2 Al2 O3
Here the No. of atoms of each element in the reactants and products side are not equal. i.e.
In reactants In products
No. of Al atoms 1 1
No. of O atoms 2 3
Hence it is an unbalanced chemical equation.
3.3 BALANCING A CHEMICAL EQUATION
Steps involved in the balancing of a chemical equation are as follows :
(i) Write the equation in the word form by keeping the reactants on the left hand side and the products on the right hand side of the arrow sign (®)
| This step is not required if the equation is already given in terms of symbols and formulae.|
(ii) Write the symbols and formulae of all the reactants and the products in the word equation to get the skeletal chemical equation.
(iii) List the no. of atoms of different elements of reactants and products.
(iv) To balance the chemical equations, select the compound which has maximum number of atoms, irrespective of the fact whether it is a reactant or product. Multiply the symbols and formulae by the smallest possible number to balance the element having maximum atoms.
(v) Also balance the others elements one by one. Do not change the formulae to balance the chemical equation.
Write the balanced chemical equation for the following reaction. Copper sulphate reacts with sodium hydroxide to form copper hydroxide and sodium sulphate.
(i) Writing the equation in the word-form.
Copper + Sodium Copper + Sodium
Sulphate Hydroxide Hydroxide Sulphate
(ii) Writing the skeletal chemical equation and enclosing the formulae in boxes.
(iii) Element No. of atoms in LHS No. of atoms in RHS
Cu 1 1
S 1 1
O 5 6
Na 1 2
H 1 2
(iv) Selecting the biggest formula (i.e. Na2 SO4) and balancing the element with highest number of atoms i.e. oxygen.
There are 5 atoms of oxygen in LHS but 6 in RHS. To balance oxygen multiply NaOH by 2.
CuSO4 + 2 NaOH Cu(OH)2 + Na2SO4
By doing so sodium and hydrogen also get balanced.
(v) Checking the correctness of the balanced equation.
Element In Reactants In Products
Cu 1 1
S 1 1
O 6 6
Na 2 2
H 2 2
Hence the final balanced chemical equation may be written as
CuSO4 + 2NaOH Cu(OH)2 + Na2SO4
| The method of enclosing the formulae in boxes is only for the beginners. Once you get into the habit of remembering that the formula of any reactant or product cannot be changed, boxes are not required. Similarly, listing of atoms of different elements is also not required.|
3.4 LIMITATIONS OF A CHEMICAL EQUATION
It gives us no information about the following :
(i) The physical state of the reactants.
(ii) The concentration of the reactants.
(iii) The time taken for the reaction to complete.
(iv) The rate at which the reaction proceeds.
(v) The conditions necessary to start and carry on the reaction e.g., Is any catalyst required? What is the temperature needed to start and continue the reaction?
(vi) Is the reaction exothermic or endothermic, i.e, is heat evolved or absorbed during the reaction?
3.5 ESSENTIALS OF A CHEMICAL EQUATION
A true chemical equation, therefore, must be in accordance with the following essentials :
(i) It should represent an actual chemical change.
(ii) It should be balanced, i.e., number of atoms of different elements on the two sides of the equation must be equal.
(iii) It should be molecular i.e., all the substances concerned should be expressed as molecules.
3.6 TO MAKE EQUATION MORE INFORMATIVE
It is quite helpful if an equation gives an idea about the physical state, heat changes and the conditions under which the reaction takes place.
This can be done in the following three ways.
(i) To indicate the physical states of reactants and products : By indicating the physical states of the reactants and products of a chemical reactions by writing letters (g) for gaseous state, (l) for liquid state, (s) for solid state and (aq) for aqueous solutions i.e. solutions in water, just after the formulae in an equation as shown in the examples below.
Zn(s) + H2SO4(aq) ® ZnSO4(aq) + H2(g).
An arrow pointing upwards () may also be used to indicate a gaseous product in the equation as follows :
Zn (s) + H2SO4(aq) ® Zn SO4(aq) + H2 .
An arrow pointing downwards (¯) is used to indicate an insoluble product or precipitate (ppt) in an equation as follows :
AgNO3 (aq) + NaCl (aq) ® AgCl(s) ¯ + NaNO3(aq)
Silver Nitrate Sodium chloride Silver chloride Sodium Nitrate
(ii) To indicate the heat changes (thermo chemical equations) : Chemical equations for exothermic or endothermic reactions representing the heat evolved or absorbed during the reaction are called thermo chemical equations.
(a) An exothermic reaction is indicated by writing + heat, or + heat energy or + energy, on the products side of an equation as follows :
C(s) + O2(g) ® CO2(g) + heat
CaO (s) + H2O(l) ® Ca(OH)2 + heat
CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) + heat
(b) An endothermic reaction is indicated by writing “+ heat” or “+ heat energy” or “+ energy” on the reactants side of an equation as shown below.
CaCO3(s) + heat ® CaO(s) + CO2(g)
Calcium carbonate Calcium oxide Carbondioxide
N2(g) + O2(g) + heat ® 2NO(g)
Nitrogen Oxygen Nitric oxide
Ba(OH)2 + 2 NH4Cl + heat ® BaCl2 + 2 NH4OH
Barium hydroxide Ammonium chloride Barium chloride Ammonium hydroxide
| Heat absorbed in a chemical reaction can also be represented by the symbol delta (D) as follows:|
(iii) To indicate the conditions under which the reaction takes place: The conditions of temperature, pressure and the presence of catalyst if any are represented by writing these conditions above or below the arrow sign as follows:
Here 500°C temperature, 200 atm pressure and Fe as catalyst are the condition for the reaction to take place.
Here 300°C temperature, 300 atm pressure and zinc acid (ZnO) and chromium oxide CrO3 as catalyst are the conditions needed for the reaction to take place.
If heat is required for the reaction to take place, then heat sign delta (D) is written over the arrow of the equation as follows :
Here D shows the heat needed and MnO2 is the catalyst needed for the reaction to take place.
|Which of the following chemical equation gives us complete information about reactant and products?
4.1 COMBINATION REACTIONS
Those reactions in which two or more substances (reactants) combine together to form a single substance (product) are called the combination reactions.
Example 1 : Formation of water from H2(g) and O2(g)
2H2(g) + O2(g) ® 2H2O(l)
Hydrogen Oxygen Water
In this reaction, two substances hydrogen and oxygen (reactants) combine together to form a single substance i.e. water (product). So it is a combination reaction.
Example 2 :
2CO (g) + O2(g) ® 2CO2 (g)
Carbon (Coal) Oxygen Carbondioxide
In this reaction again two substances namely, carbon monoxide and oxygen combine together to form a single substance, carbon dioxide (product). Thus it is also a combination reaction.
Example 3 : Reaction of ammonia and hydrogen chloride
NH3(g) + HCl(g) ® NH4Cl(s)
Ammonia Hydrogen chloride Ammonium chloride
This reaction is a combination reaction, as ammonia and Hydrogen chloride combine together to form ammonium chloride as a single product.
Example 4 : Reaction of water on quick lime :
Experiment : If we take a small amount of quick lime (calcium oxide) in a beaker and add some water to it slowly, then they combine vigorously to form slaked lime i.e. (Calcium hydroxide).
| The reaction may be represented by :
In this reaction, two substances (reactants), namely, calcium oxide and water have combined to form a single substance i.e., calcium hydroxide. Thus, it is a combination reaction.
Since a lot of heat is produced along with slaked lime so it is also called an exothermic reaction.
|The above reaction is an example of
(a) Combination between two or more elements
(b) Combination between an element and a compound
(c) Combination between two or more compounds
4.2 DECOMPOSITION REACTIONS
Those reactions in which a single substance (reactant) splits up into two or more simpler substances (products) are known as decomposition reactions.
These reactions are carried out by supplying energy in form of heat, electricity or light which breaks that substance into simpler substances. Thus decomposition reactions are classified as:
(i) Thermolysis or thermal decomposition reactions (decomposition by heat).
(ii) Electrolysis or electrolytic decomposition reactions (decomposition by electricity)
(iii) Photolysis or photodecomposition reactions (decomposition by light).
| (i) Thermal decomposition reactions : Decomposition reactions which are carried out by heating are called thermal decomposition reactions.
Example 1 : Thermal decomposition of ferrous sulphate
Experiment : Take a small amount of ferrous sulphate crystals (FeSO4 . 7H2O) in a test tube. These crystals are green in colour. Heat the test tube over a burner. The green colour of crystals first changes to white due to formation of Anhydrous Ferrous sulphate (FeSO4), as on heating, ferrous sulphate crystals lose the water molecules. On further heating anhydrous ferrous sulphate decomposes to brown solid (which is ferric oxide) along with characteristic smell of burning sulphur (fig. 4).
|Decomposion reaction of ferrous sulphate
The reactions involved are :
In this Reaction, A single substance splits up into three simpler substances. Thus it is a decomposition reaction : Since this decomposition is brought about by heat, so, it is an example of thermal decomposition reaction.
Example 2 : Thermal decomposition of Potassium Chlorate : When potassium chlorate is heated in the presence of manganese dioxide as a catalyst, it decomposes to give potassium chloride and oxygen.
In this reaction, a single substance splits up into two simpler substances on heating. Thus it is a thermal decomposition reaction.
| Example 3 : Thermal decomposition of Lead nitrate
Experiment: Take about 2 gm of powdered lead nitrate (colourless) in a dry test tube. Heat the tube over a burner. We observe that brown fumes of nitrogen dioxide are found to evolve and a yellow residue of lead oxide is left behind in the tube (fig. 5). If we put a burning candle over the mouth of test tube, it catches fire and starts burning again. This shows that the oxygen gas is also evolved during this reaction.
The reaction involved is :
In this reaction a single substance i.e. lead nitrate decomposes to three simpler substances (i.e. lead oxide, nitrogen dioxide and oxygen).
|Decomposition reaction of lead nitrate
Thus it is a decomposition reaction, Since this decomposition is carried about by heat, therefore, it is an example of thermal decomposition reaction.
Example 4: Thermal Decomposition of limestone: When calcium carbonate (limestone) is heated, it decomposes to give calcium oxide and carbon dioxide:
In this reaction, a single substance, limestone breaks up into two simpler substances, calcium oxide and carbon dioxide, so it is a decomposition reaction. Since the decomposition is carried out by heating, so it is an example of thermal decomposition reaction.
Example 5 : Thermal decomposition of zinc carbonate : When zinc carbonate is heated, it decomposes to give zinc oxide and carbon dioxide gas :
In this reaction zinc carbonate on heating decomposes to give two simpler substances, zinc oxide and carbon dioxide, thus it is an example of thermal decomposition reaction.
| (ii) Electrolytic decomposition Reactions (or electrolysis)
The decomposition reactions which are carried out by using electricity, are called electrolytic decomposition reactions (or electrolysis)
Example 1: Electrolytic decomposition of water or electrolysis of water
Experiment: Take a plastic mug and make two holes at the base of mug. Fit two rubber stoppers in these holes, with carbon electrodes. Connect the electrodes to a 6-V battery (fig. 6). Fill the mug with water so that electrodes get immersed in it. Add few drops of dil H2SO4 to make it good conductor of electricity. Fill the two test tubes with water and invert them over the electrodes. Now switch on the battery and keep the apparatus undisturbed for some time.
|Experiment set-up for electrolysis of water
- There is formation of bubbles of two different gases over the electrodes which displaces the water in test tubes.
- The gas collected in the test tube covering the cathode is double than that of gas collected in tube covering the anode.
Keep on passing current till the test tubes are completely filled with respective gases. Now switch off the battery.
On testing the gases by bringing a burning candle near mouth of tubes, it is found that the gas with double volume, burns with poping sound, so it is hydrogen gas. The other gas with lesser volume, makes the burning candle to glow more, so, it is oxygen gas.
Thus above experiment shows that on supplying electricity, water decomposes into hydrogen and oxygen according to the reaction :
Since in this reaction decomposition is carried out by using electricity, so it is an example of electrolytic decomposition reaction or electrolysis.
| Volume of hydrogen collected is double as that of oxygen because two atoms of hydrogen and one atom of oxygen makeup one molecule of water (H2O).|
Example 2: Electrolytic decomposition of molten sodium chloride : On passing electric current through molten sodium chloride, it decomposes to give sodium metal and chlorine gas :
In this reaction on passing electric current a single substance decomposes to give two simpler substances. Thus it an example of electrolytic decomposition reaction.
Example 3: Electrolytic decomposition of molten alumina (aluminium oxide) : On passing electric current through molten alumina, it decomposes to give aluminium metal and oxygen gas :
In this reaction, on passing electric current, a single substance i.e. alumina decomposes to give two simpler substances, aluminium metal and oxygen gas, thus it is an example of electrolytic decomposition reaction.
(iii) Photo-decomposition reactions (or photolysis) : The decomposition reactions which take place by absorption of light are called the photo-decomposition reactions or photolysis.
Example 1: Photo-decomposition of silver chloride or Photolysis of Silver chloride
Experiment: Take a pinch of silver chloride on a watch glass and keep it in sunlight for some time.
Decomposition of silver chloride in the presence of light
It is observed that white silver chloride turns gray due to formation of silver metal.
In this reaction a single substance i.e. silver chloride decomposes (in pressure of sunlight) into two simpler substances, silver metal and chlorine gas. Thus this is an example of photo-decomposition reaction or photolysis.
Example 2: Photolysis of hydrogen iodide: Hydrogen iodide decomposes in the presence of ultraviolet light into hydrogen and iodine :
In this reaction single substance i.e., hydrogen iodide on absorption of ultraviolet light decomposes into two simpler substances, hydrogen and iodine. Thus it is an example of photo-decomposition reaction or photolysis.
Example 3: Photolysis of hydrogen peroxide : In presence of light, hydrogen peroxide decomposes into water and oxygen.
In this reaction, a single substance decomposes in the presence of light, into two simpler substances. Thus it is an example of photo-decomposition reaction.
| Since hydrogen peroxide decomposes in the presence of light, that is why, hydrogen peroxide is kept in coloured bottles so as to cut off light.|
4.2.1 Decomposition reactions are called the opposite of combination reactions
In a decomposition reaction, one substance breaks up into two or more chemical substances, while in a combination reaction two or more substances combine to form one single substance. So these two reactions are called opposite of each other.
4.2.2 Uses of Decomposition Reactions
The decomposition reactions are used in the extraction of metals in the following ways :
(i) Metals are extracted from their molten salts by electrolytic decomposition e.g. sodium from molten sodium chloride and aluminium from alumina (molten aluminium oxide).
(ii) Thermal decomposition reactions form one of the steps in extraction of metals.
For example, Zinc carbonate (the naturally occurring ore of zinc) is first decomposed to give zinc oxide and then reduced to obtain zinc metal i.e.,
4.2.3 Decomposition Reactions in our body
The digestion of food in the body is an example of decomposition reaction. When we eat foods like wheat, rice or potatoes, then the starch present in them decomposes to give simple sugars like glucose in the body and proteins decompose to form amino acids :
4.2.4 Decomposition Reactions are endothermic reactions
All decomposition reactions require energy either in form of heat, light or electricity. Hence all decomposition reactions are endothermic (heat absorbing) reactions.
4.3 DISPLACEMENT REACTIONS
Those reactions in which more active element displaces a less active element from its compound are called displacement reactions.
| The elements involved may be metals or non-metals. In displacement reactions more reactive metal may displace a less reactive metal or a more reactive non-metal may displace a less reactive non-metal from its compound|
4.3.1 Relative activities or reactivities of metals
Metals have been arranged in decreasing order of their activities (or reactivities) in the activity series as follows:
It is clear from the series that the metals lying above the hydrogen are more reactive than the metals lying below the hydrogen. Thus any metal can displace the metals lying below it from its solution.
| Hydrogen is included in the series because, like metals, hydrogen can lose an electron to form positive (H+) ion. As the metals lying above hydrogen are more reactive, they can displace hydrogen from acids or water, i.e. they can react with an acid to give out hydrogen gas. On the other hand, metals lying below hydrogen being less reactive than hydrogen cannot displace hydrogen from acids to give out hydrogen gas.|
4.3.2 Relative activities (or reactivities) of Non-metals
Relative activities of non-metals like halogens is in the order:
F > Cl > Br > I
Thus, fluorine is most reactive and iodine is least reactive. So fluorine (F2) can displace chlorine (Cl2) from NaCl, Bromine (Br2) from NaBr and so on. Similarly, chlorine can displace bromine (Br2) from KBr and iodine (I2) from KI and so on.
Let us discuss some displacement reactions on the basis of reactivity of metals and non-metals.
Example 1: Reaction between iron and copper sulphate solution :
Experiment : Take about 10 ml of copper sulphate solution in a test tube. It is deep blue in colour. Take two iron nails and clean their surface by rubbing with a sand paper.
Now put the one nail in test tube containing CuSO4 solution. Keep another iron nail aside for comparison.
Displacement reaction between iron (nail) and copper sulphate solution.
After 30 minutes we observe the following changes :
(i) Blue colour of CuSO4 has faded and it changes into light green due to formation of iron sulphate (FeSO4)
(ii) The iron nail is covered with a red brown layer of copper metal.
These changes show that the following reaction has taken place :
Thus, more reactive metal iron, (Fe), displaces the less reactive metal copper from copper sulphate solution, so this is an example of displacement reaction.
Example 1 : Displacement reactions in which a more reactive metal displaces a less reactive metal from its compound :
Thus, when zinc pieces are added to copper sulphate solution, then, zinc being more reactive metal than copper, displaces copper from its solution (CuSO4) so that Cu is set free. The blue colour of CuSO4 solution fades due to formation of ZnSO4 (colourless). A reddish brown deposit of copper metal is formed on the surface of zinc. Therefore it is an example of displacement reaction.
Thus, when a copper wire is dipped in silver nitrate solution, copper, being more reactive metal than silver, displaces silver from its solution (AgNO3) so that silver is liberated. This silver is deposited on the copper wire giving it a white shining surface. The solution forms a blue colour due to formation of copper nitrate. Thus it is an example of displacement reaction.
In this reaction sodium, being more active than hydrogen, displaces hydrogen from water so that hydrogen gas is liberated along with the formation of sodium hydroxide. Thus, it is an example of displacement reaction.
Example 2 : The displacement reactions in which more reactive non-metal displaces less reactive non-metal from its compound :
Thus, when Cl2 gas is passed through an aqueous solution of potassium bromide, Chlorine being more reactive than bromine, displaces bromine from KBr so that bromine gas (Br2) is liberated. The solution forms light brown colour due to dissolution of Br2 gas in it. So it is an example of displacement reaction.
Thus, when chlorine gas is passed through potassium iodide solution, chlorine, being more reactive than iodine, displaces iodine from KI and liberates I2 gas.
The solution acquires violet colour due to dissolution of I2 gas in it. So it is also an example of displacement reaction.
(c) Fe(s) + CuSO4 (aq) ® (d)
Will these reactions takes place or not.
4.3.3 Uses of Displacement Reactions
Displacement reactions are used in the extraction of silver and gold. Silver or gold ore is dissolved in sodium cyanide solution. When zinc granules are added to the solution of the compound formed, zinc, being more active than silver and gold, displaces silver and gold from the solution of their compounds and thus silver and gold are extracted.
4.3.4 Displacement reactions are exothermic
All displacement reactions are exothermic (heat producing) reactions, For example :
(i) In the displacement reaction between zinc and dilute hydrochloric acid or dilute sulphuric acid, there is production of heat along with evolution of gas.
(ii) In displacement reaction between iron and copper sulphate solution, there is an increase in temperature due to production of heat.
Thus displacement reactions are heat producing or exothermic reactions.
How would you say that silver is chemically less reactive than copper ?
4.4 DOUBLE DISPLACEMENT REACTIONS
Those reactions in which two different atoms or groups of atoms are exchanged are called the double displacement reactions.
| These reactions generally occur between two ionic compounds in the solution. So they may be defined as :
“Those reactions in which two ionic compounds in the solution react by exchange of their ions to form new compounds are called double displacement reactions”
| Example 1 : Reaction between Barium Chloride solution and sodium sulphate solution.
Experiment: Take sodium sulphate solution in a test tube and add few drops of solution of barium chloride to it and mix the two solution. There is formation of white precipitate of Barium sulphate: The reaction is as follows :
|Formation of white ppt. of barium sulphate on adding barium chloride solution to sodium sulphate
In this reaction, there is double displacement or exchange of ions, i.e. chloride ions of BaCl2 have been replaced by sulphate ions of Na2SO4 whereas sodium ions of Na2SO4 have been replaced by chloride ions of BaCl2. Hence it is a double displacement reaction.
| Above reaction is also called precipitation reaction as white precipitate of BaSO4 is formed.|
| Example 2: Reaction between lead nitrate solution and potassium Iodide solution.
Experiment : Take lead nitrate solution in a test tube and add potassium Iodide solution to it and mix the two solutions. There is formation of yellow precipitate of lead iodide (fig. 11).
The reaction involved is :
|Formation of yellow ppt. of lead iodide on adding potassium iodide solution to lead nitrate
In this reaction, there is double displacement or exchange of ions, i.e. lead ions of Pb(NO3)2 have been replaced by potassium ions of KI and iodide ions of KI have been replaced by nitrate ions of Pb(NO3)2. Thus it is a double displacement reaction.
| Above reaction is also called precipitation reaction as yellow precipitate of PbI2 is formed.|
Some more examples of double displacement reaction are :
All the above reactions are double displacement reactions, as there is double exchange or displacement of ions.
| Since these reactions produce precipitates, thus they are also called precipitation reactions. The reaction that produces a precipitate can be called as precipitation reaction.|
Is an acid-base reaction a double displacement reaction?
4.5 OXIDATION – REDUCTION REACTIONS OR REDOX REACTIONS
Oxidation can be defined in three ways :
(i) Oxidation in terms of oxygen or electronegative element :
Oxidation is a process in which oxygen or electronegative element is added to a substance for example:
(ii) Oxidation in terms of hydrogen or electropositive element
Oxidation is defined as a process in which hydrogen or an electropositive element is removed from the substance i.e.
(a) Removal of hydrogen
(b) Removal of electropositive element
(iii) Oxidation in terms of electronic concept
Oxidation is a process in which loss of electrons take place. or
Oxidation is a process in which electrons are lost by an atom, ion or a group of atoms taking part in the chemical reaction.
As a result, there is increase in positive charge or decrease in negative charge of the atom or group of atoms in an oxidation reaction e.g.,
(d) (Decrease in negative charge) or (loss of electron)
(e) (Decrease in negative charge) or (loss of electron)
Reduction can also be defined in three ways :
(i) Reduction in terms of oxygen or electronegative element: Reduction is a process in which oxygen or electronegative element is removed from a substance i.e.
(ii) Reduction in terms of hydrogen or electropositive element: Reduction is a process in which hydrogen or electropositive element is added to a substance i.e.
(iii) Reduction in terms of electronic concept : Reduction is a process in which gain of electrons takes place
Reduction is a process in which electrons are gained by an atom, ion or a group of atoms. As a result, there is an increase in negative charge or decrease in positive charge e.g.,
Give one example of a redox reaction which is also a combination reaction.
4.6 REDOX REACTIONS
In the term ‘redox’, ‘red’ stands for reduction and ‘ox’ stands for oxidation.
Thus the reactions in which oxidation and reduction take place simultaneously are called Redox reactions, i.e. in redox reactions one substance is oxidized and other is reduced.
4.6.1 Oxidizing agent
It is a substance which itself gets reduced but oxidizes the other substance by
(i) Adding oxygen or electronegative element to other substance.
(ii) Removing hydrogen or electropositive element from other substance.
(iii) Gaining electrons from other substances.
4.6.2 Reducing agent
It is a substance which itself gets oxidized but reduces the other substance by :
(i) Adding hydrogen or electropositive element to substance.
(ii) Removing oxygen or electronegative element from a substance.
(iii) Losing or donating electrons to other substances.
Let us discuss oxidizing agents and reducing agents by taking some examples of oxidation-reduction or redox reactions.
|Antioxidant are :
(a) Reducing agents (b) Oxidising agents (c) Salts (d) None of these
| Example 1: Oxidation of copper to copper oxide and reduction of copper oxide to copper.
Experiment: Take about 1 g of copper powder in a china dish and heat it as shown in figure.
It is observed that surface of copper powder becomes coated with black copper (II) oxide and the reaction taking place is
It is clear from the reaction that there is addition of oxygen to copper, as a result copper has been oxidized to copper (II) oxide. Thus, it is an oxidation reaction.
Now, if hydrogen gas is passed over Cu(II) oxide, the black Copper (II) oxide changes to brown copper:-
It is clear from the reaction that there is removal of oxygen from copper oxide and addition of oxygen to hydrogen, as a result, copper oxide has been reduced to copper and hydrogen has been oxidized to water. Since oxidation and reduction occur simultaneously, in this reaction, so it is an example of redox reaction. Which can be shown more clearly as follows :
In the above reaction copper oxide (CuO) is giving oxygen required for the oxidation of hydrogen, therefore, CuO is oxidizing agent. Hydrogen is responsible for removing oxygen from CuO, therefore, hydrogen is the reducing agent.
(i) Substance oxidized : H2
(ii) Substance reduced : CuO
(iii) Oxidizing agent : CuO
(iv) Reducing agent : H2
| In reaction between copper oxide and hydrogen, hydrogen is more active than copper so, hydrogen replaces copper from copper oxide so that copper is formed, free. Thus it is also an example of displacement reaction.|
Example 2: When hydrogen sulphide reacts with chlorine, then sulphur and hydrogen chloride are formed as:
In this reaction there is removal of hydrogen from H2S, which is being oxidized to S and Cl2 is being reduced to HCl. Since both oxidation and reduction occur together, so it is an example of redox reaction. Which can be shown very clearly as follows:-
In the above reaction, H2S is giving hydrogen required for reduction of Cl2, therefore, H2S is reducing agent. Cl2 is responsible for removing hydrogen from H2S, therefore Cl2 is the oxidizing agent.
(i) Substance oxidized : H2S
(ii) Substance reduced : Cl2
(iii) Oxidizing agent : Cl2
(iv) Reducing agent : H2S
| Oxidizing agent, itself gets reduced, but oxidizes the other substance and reducing agent itself gets oxidized, but reduces the other substance.|
Some more examples of oxidation reduction or Redox reactions:-
|Which of the following is an example of both oxidation and reduction reaction
(c) (d) All of these
4.6.3 The effects of oxidation reactions in everyday life
The damaging effect of oxidation on metals is studied as corrosion and that on food is studied as rancidity. Thus, there are two common effects of oxidation reactions which we observe in everyday life:
(i) Corrosion of metals
(ii) Rancidity of food.
(i) Corrosion : The slow process of eating up of metals due to attack of atmospheric gases such as oxygen, carbon dioxide, hydrogen sulphide, water vapour etc. on the surface of the metals so as to convert the metal into oxide, carbonate, sulphide etc. is known as corrosion.
Example : Corrosion of iron (Rusting) :
The most common example of corrosion is rusting i.e., corrosion of iron. When an iron object is left in the moist air for a long time, its surface is covered with a brown, flaky (non-sticky) substance called rust. Rust is mainly hydrated ferric oxide (Fe2O3.xH2O). It is formed due to attack of oxygen and water vapour present in moist air on the surface of iron :-
This reaction is called corrosion of iron or rusting.
| Rusting involves unwanted oxidation of iron metal which occurs in nature on its own.|
Effect of Rusting in everyday life : Rusting is a serious problem because it weakens the structure of bridges, iron railings, automobile parts etc. Every year, a lot of money is spent to replace rusted iron and steel structures. The reason is that raddish brown crust of rust does not stick to the surface. It falls down exposing fresh surface for rusting. Thus, corrosion of iron or rusting is a continuous process which ultimately eats up the whole iron object.
Methods to prevent rusting or Prevention of Rusting : Rusting can be prevented if iron objects are not allowed to come in contact with the moist air. It can be done by :
(a) Painting the iron objects such as iron gates, steel furniture, bodies of cars etc.
(b) Greasing and oiling the iron objects such as machine parts etc.
(c) Galvanisation is coating the surface of iron objects with a thin layer of zinc, which is more resistant to corrosion.
SOME MORE EXAMPLES OF CORROSION OF METALS
- Corrosion of copper : When a copper object (shiny brown) is left in moist air for a long time, then its surface is covered with green coating of basic carbonate, CuCO3 . Cu(OH)2. This is due to attack of O2, CO2 and water vapour present in moist air on the surface of copper :
This reaction is called corrosion of copper.
| Corrosion of copper involves unwanted oxidation of copper metal which occurs in nature on its own.|
- Corrosion of Silver : When a silver object (shiny white) is kept in air for a long time, its surface is covered with coating of black silver sulphide (Ag2S). This is due to the attack of H2S gas present in air on the surface of silver.
This reaction is called the corrosion of silver.
| Corrosion of silver also involves unwanted oxidation of silver metal which occurs in nature on its own.|
(ii) Rancidity : When fats and oils in food are oxidized, their smell and taste changes and they become rancid. This phenomenon is called rancidity.
Oxidation of oils or fats in a food resulting into a bad smell and bad taste is called rancidity.
Prevention from Rancidity :
(a) adding antioxidants (Reducing agents) like BHA (Butylated Hydroxy Anisole) and BHT (Butylated Hydroxy-Toluene) to foods containing fats and oils.
(b) Packaging fat and oil containing foods in nitrogen gas (inert gas)
(c) Keeping food in a refrigerator
(d) Storing food in air tight containers.
(e) Storing foods away from light.
Which antioxidant is used to prevent rancidity in foods ?
Depending upon the kind of heat change during a reaction, the chemical reactions are classified into two types :
(i) Exothermic Reactions
(ii) Endothermic Reactions
(i) Exothermic Reactions : The term exothermic is taken from Greek word exotherm (exo-out, therm-heat) which means heat goes out. Thus a reaction in which heat is released or produced is called an exothermic reaction.
| In exothermic reaction heat is released to the surroundings, as a result the temperature increases.|
(a) When nitrogen and hydrogen combine then, ammonia (NH3) is formed and heat is liberated. Thus, formation of ammonia is an exothermic reaction.
(b) When methane gas is burnt, heat is liberated, along with CO2 and H2O. Thus, combustion of methane gas is an exothermic reaction.
(c) Burning of Coke in air is an exothermic reaction.
(d) Formation of water from H2(g) and O2(g) is an exothermic reaction.
The steam H2O(g) so produced condenses into liquid water and 45 kJ heat is released.
(e) During respiration process glucose undergoes combustion to form CO2 and H2O. The reaction is exothermic.
WHY AND WHEN A CHEMICAL REACTION IS EXOTHERMIC
We know that in a chemical reaction, heat is supplied to break the bonds in the reactants and heat is released due to bond formation in the products. When the heat released due to bond formation in products is greater than the heat supplied to break the bonds in reactants, then the reaction is exothermic.
For example formation of H2O from (H2) and (O2) is exothermic because the energy released in the formation of 2 covalent bonds in H2O is more than the energy, required to break
H – H bonds in H2 and O = O bonds in O2.
5.1 ENDOTHERMIC REACTIONS
The term endothermic is taken from the Greek word endotherm (endo-in, therm-heat) which means heat is taken in. Thus a reaction in which heat is absorbed is called the endothermic reaction.
| In endothermic reaction heat is taken from the surrounding as a result the temperature decrease.|
(a) Reaction between coke and steam is endothermic i.e., heat is absorbed by reactants to give the products :
(b) Thionyl chloride (SOCl2) is a useful drying agent. Its reaction with water is endothermic.
(c) When nitrogen and oxygen combine to form nitric oxide (NO), heat is absorbed. This reaction is endothermic :
(d) Melting of ice into liquid water is an endothermic reaction (heat absorbing reaction)
WHY AND WHEN A CHEMICAL REACTION IS ENDOTHERMIC
In a chemical reaction when heat required to break the bonds in all reactants is greater than the heat released in the bond formation of the products, then the reaction is said to be endothermic.
For example formation of nitric oxide from N2 and O2 is endothermic because the heat required to break N – N bonds in N2 and (O=O) bonds in O2 is more than the heat released in the formation of 2 covalent bonds in NO.
- The different types of reactions are: combination reactions, decomposition reactions, displacement reactions and redox reactions.
- Reaction in which two or more substances combine to form another compound are combination reactions.
- Reactions in which a compound breaks up into simple substances are decomposition reactions.
- Reactions in which an atom or group of atoms in a molecule is replaced by another atom or group of atoms are displacement reactions.
- Reactions in which two compounds react to form two other compounds by mutual exchange of their ions are double displacement reactions.
- Oxidation is a process which involves addition of oxygen or removal of hydrogen. According to electronic concept, oxidation is a process which involves loss of electrons.
- Reduction is a process which involves addition of hydrogen or removal of oxygen. According to electronic concept, reduction is a process which involves gain of electrons.
- Oxidising agent is a substance which gives oxygen or removes hydrogen and causes oxidation of other substances. It gets itself reduced.
- Reducing agent is a substance which gives hydrogen or removes oxygen and causes reduction of other substance. It gets itself oxidised.
- Irreversible reaction is that which reaches to completion.
- Reversible reaction is that which do not reaches upto completion and occur in both directions.
- At equilibrium state both opposing process takes place with equal rates.
- Rate of reaction depends upon the active masses of reactants.
- Any change in the state of equilibrium leading with equilibrium shift.
- In the refining of silver, the recovery of silver from silver nitrate solution involved displacement by copper metal. Write down the reaction involved.
- Why are bags of fat and oil containing food items (like chips) flushed with nitrogen?
- When is a substance said to be reduced?
- What type of reactions are represented by the following equation?
(a) NH4Cl ® NH3 + HCl
(b) 2H2 + O2 ® 2H2O
(c) BaCl2 + Na2SO4 ® BaSO4 + 2NaCl
(d) Mg + CuSO4 ® MgSO4 + Cu
- Can silver nitrate (AgNO3) solution be stored in an iron container? Explain your answer.
- Balance the following chemical equations:
(i) Fe + H2O ¾® Fe3O4 + H2 (ii) Na + H2O ¾® NaOH + H2
- Give two examples of reduction reaction?
- Why do we apply paint on iron articles?
- What do you mean by precipitation reaction? Give its example.
- Write a balanced chemical equation with state symbols for the following reactions:
(i) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate.
(ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water.
- Write an activity to show the change in the state of matter and change in temperature during a chemical reaction
- Write one activity to show the decomposition of a chemical compound with the evolution of a gas.
- Write an activity to show the electrolysis of water, as an example of decomposition reaction.
- Name different types of chemical reactions. Define them and give their examples.
- What is a balanced chemical equation? Why should chemical equations be balanced?
Objective Type Questions
- In the balanced equation:
The values of a, b, c, d are respectively
(a) 1, 1, 2, 3 (b) 1, 1, 1, 1
(c) 1, 3, 2, 3 (d) 1, 2, 2, 3
- Which of the following reactions is not balanced?
(c) 2Al + 6H2O (d) 4NH3 + 5O2 4NO + 6H2O
- The equation
The values of x and y are
(a) 3 and 5 (b) 8 and 6
(c) 4 and 2 (d) 7 and 1
- Neutralization reaction is an example of –
(a) exothermic reaction (b) endothermic reaction
(c) oxidation (d) none of these
- Which of the following statements is/are true?
(a) The total mass of the substance remains same in a chemical change.
(b) A chemical change is permanent and irreversible.
(c) A physical change is temporary and reversible.
(d) All of these.
- Which of the following statements is not correct?
(a) A chemical equation tells us about the substances involved in a reaction.
(b) A chemical equation informs us about the symbols and formulae of the substances involved in a reaction.
(c) A chemical equation tells us about the, atoms or molecules of the reactants and products involved in a reaction.
(d) All are correct.
- Zn(s) + H2SO4 (aq) ZnSO4 (aq) + H2 (g) is an example of
(a) precipitation reaction (b) endothermic reaction
(c) evolution of gas (d) change in colour
- When dilute hydrochloric acid is added to iron fillings
(a) hydrogen gas and ferric chloride are produced.
(b) chlorine gas and ferric hydroxide are produced.
(c) no reaction takes place.
(d) iron salt and water are produced.
- In the reaction xPb(NO3)2
x, y and z are-
- A redox reaction is one in which
(a) both the reactants are reduced.
(b) both the reactants are oxidised.
(c) an acid is neutralised by the base.
(d) one substance is oxidised, while the other is reduced.
- Chemical reaction 2Na + CI2 2NaCl is an example of
(a) combination reaction (b) decomposition reaction
(c) displacement reaction (d) double displacement reaction
- Which of the following equations is representing combination of two elements?
Objective Type Questions
|1. (c)||2. (b)||3. (c)||4. (a)||5. (d)|
|6. (d)||7. (c)||8. (a)||9. (b)||10. (d)|
|11. (a)||12. (a)|