(d) When copper carbonate is heated, copper dioxide is produced.
(e) One of the products is hydrogen gas when zinc is treated with dilute sulphuric acid,
what is the next product?
(f) Magnesium is burning in air.
6. Fill in the blanks and balance the following chemical equations.
(a) Na + …………. → NaOH + ………….
(b) Na2CO3 + …………. → NaNO3 + …………. + H2O
(c) Mg + …………. → MgCl2 + H2
(d) Ca(OH)2 + …………. → CaCO3 + H2O
(e) Zn + H2SO4 → …………. + ………….
(f) ………. + …………. → MgSO4 + H2
(g) KClO3 → …………. + 3O2
7. Balance the following chemical equations (Fill in the blanks if necessary)
(a) CH4 + O2 → CO2 + H2O
(b) NH4 + Na →
(c) FeS NaNH2 + H2
(d) Fe + O2 → F2O3 + SO2
+ ………….
(e) H2 → FeSO4 + Cu
(f) Zn
+ Cl2 → ………….
+ S
→ ………….
(g) Fe + …………. → Fe2O3
→ Hg
(h) HgO + ………….
(i) Zn + H2SO4 → …………. + H2
(j) Al + …………. → AlCl3 + H2
(k) HNO3 + NaOH → …………. + ………….
(l) H2SO4 + …………. → K2SO4 + ………….
8. There are one or more mistakes in the following chemical equations. Correct and balance
them.
(a) Al + HCl → AlCl3 + H
(b) C + O → CO
(c) Na + Cl → 2NaCl
(d) H2O2 → H2O
(e) H2SO4 + NaOH → Na2SO4 + H3O
Blooming Science Book 9 151
Chapter SOLUBILITY
10
Learning Outcomes Estimated Periods: 5+3
On the completion of this unit, the students will be able to:
• prepare saturated and unsaturated solution.
• define super saturated solution.
• define solubility and solubility curve.
• explain the relationship between temperature and solubility
• explain the process of crystallization.
• solve some simple numerical problems related to solubility.
Introduction
In our daily life, we come across different kinds of mixture. The water we drink is a solution
of different kinds of minerals dissolved in it. We have seen the mixtures of sand and rice coats
in rice, soda water, sugar solution, air and muddy water, etc. In these examples, two or more
substances are mixed with each other in any proportion by weight. Their particles are in intimate
contact but do not react at all. Thus, they are called mixtures.
A mixture is defined as a mass obtained by mixing up two or more substances in any proportion
by weight so that each of which retains its identity and own properties. The substances, which
take part in the formation of mixtures, are called components of mixtures. The properties of each
of the component in a mixture remain unchanged. The components of a mixture may be present
in any of the three states, i.e. solid, liquid or gas. Depending upon the nature of the components
of mixture, there are two types of mixtures. They are:
1. Homogeneous mixture 2. Heterogeneous mixture
1. Homogeneous Mixture is defined as a mass in which the particles of components of mixture
are equally distributed. The components of homogeneous mixture are present as molecular
particles so the component particles cannot be seen in it by the naked eye. In salt-water
mixture, salt is dissolved in water and the salt particles cannot be seen. Other examples of
homogenous mixture are sugar solution, air, alcohol water, brass, soda water, etc.
2. Heterogeneous Mixture is defined as a mass in which particles of components of mixture
are not equally distributed. The component particles of this mixture can be seen by the
eye. In the muddy water, the soil particles can be seen by the eyes. Some of the other
examples of heterogeneous mixture are smoke, blood, oily water, sandy water, milk, etc.
On the basis of the size of the particles, mixtures are following three types:
1. Solution
2. Colloids
3. Suspension
152 Blooming Science Book 9
Solution
In our daily life, we come across different kinds of solutions. The water we drink is a solution
of various minerals dissolved in it. The air we breathe in is a solution of different gases. Our
morning tea is a solution of sugar, tea, milk and water.
The homogenous mixture of two or more different substances is called solution. The size of the
solid particles in solution is 10-5cm or less.
When a pinch of well-crushed copper sulphate is shaken up with water in a test-tube, the solid
particles of copper sulphate finally disappear and a clear blue liquid is obtained. Copper sulphate
is said to have dissolved in water and the resulting mixture of the two is called a solution of copper
sulphate in water. The intensity of the blue colour of copper sulphate is the same throughout the
solution. This shows that the solution is homogenous, i.e. its composition distributes equally.
The amount of solute that can be dissolved in the solvent is different. Here it is discussed about
the solution, the extent to which the solute dissolves in water to make a solution and the effect of
temperature in the solution.
Some common liquid solvents other than water are: alcohol, benzene, carbon disulphide, acetone,
ether, etc. Iodine crystals when dissolved in alcohol, the solution is called tincture of iodine.
Solution is defined as the homogenous mixture of solvent and solute. For example, salt and water
form sugar solution. Solution = Solute + Solvent
Solvent is the substance which allows a solute to dissolve in it to form a solution. Water and
alcohol are solvents.
Solute is the substance that dissolves in some other substances to form solution.
In solid-liquid type of solution, the amount of solute is always less than the amount of solvent.
The solid is always a solute and the liquid is the solvent. Sugar and salt are the solutes. But in
liquid-liquid type of solution, the one which is less in amount is the solute.
Activity
To prepare a solution of Copper Sulphate
1. Take a few crystals of copper sulphate (CuSo4) and put it in a beaker containing
water.
2. Shake it well.
3. The solid copper sulphate finally disappears.
Observation
The uniform blue colour of copper sulphate is observed.
Colloids
It is a homogeneous mixture in which the diameter of particles of components range in between
that of particles of solution and suspension, i.e. a particle’s size ranges from 10-5cm to 10-7 cm.
The particles can pass easily through the hole of filter paper and cannot be seen under a simple
microscope. Some examples of colloids are blood, gum, milk, wax, glue, etc.
Blooming Science Book 9 153
Suspensions
It is a heterogeneous mixture of solid in a solid or solid in a gas in which the diameter of the
particles is 10-5cm or larger. The particles or suspensions are visible under a simple microscope
as well as to the naked eyes. The components of suspensions can be separated by decantation and
filtration methods. Some examples of suspensions are sand water, muddy water, smoke in air, etc.
Importance of Solution in Daily Life
It is very important to study about solutions.
1. Plants take mineral salts in the form of solution from the soil.
2. We can also take some foods and medicine as solutions.
3. Aquatic animals take oxygen dissolved in water.
4. Some chemical reactions are possible only in the state of solution so that it can play
important role in industries for testing of raw materials and products.
Saturated and Unsaturated Solutions
If a well-powdered solute, such as copper sulphate is added little by little to a definite volume
of water at room temperature with constant stirring, it dissolves continuously, till a point comes,
when no more of the copper sulphate will dissolve. The excess of the copper sulphate settle down
to the bottom. The solution at this stage is said to be a saturated solution at that temperature.
Thus, in a saturated solution at a particular temperature contains as much solute as it can dissolve
at that temperature.
At that temperature solvent and solute are in a state of equilibrium. This state of equilibrium
or saturation can be changed either by adding more water (solvent) to it or by increasing the
temperature of solution. For example, the saturated solution at a given temperature becomes
unsaturated when heated, because more of solute will be required to make the solution saturated
at higher temperature. Therefore, different amounts of the same solute are required to prepare
saturated solutions at different temperatures. In other words, a solution saturated at one
temperature may not be saturated at another temperature.
A solution which is unable to dissolve any more of the solute at a particular temperature is called
a saturated solution at that temperature. At every step, before obtaining a saturated solution, a
solution is always unsaturated with respect to the solute. An unsaturated solution contains less
solute than it can dissolve at that temperature. It becomes more unsaturated with the rise of
temperature.
A solution which can dissolve more solute at a given temperature is called an unsaturated solution
at that temperature.
154 Blooming Science Book 9
Differences between Saturated and Unsaturated Solutions
Unsaturated Solution Saturated Solution
1. An unsaturated solution contains as 1. A saturated solution contains as
less solute as it can dissolve at that much solute as it can dissolve at that
temperature. That is, it can dissolve more. temperature.
2. With the increase in temperature, it 2. With the increase in temperature it
becomes more unsaturated. becomes unsaturated.
Supersaturated Solutions
A saturated solution prepared at a higher temperature contains more of a given solute than does
a saturated solution prepared at a lower temperature. So, when a saturated solution prepare at a
higher temperature is cooled to room temperature will tend to throw out the excess of the solid
from the solution in the form of crystals. It takes, however, sometime for the excess solute to
come out. During this interval solution holds in it more solute than is required so saturate it. Such
a solution is called a supersaturated solution. Thus, a supersaturated solution is one which has
more of the solute than a saturated solution requires.
The supersaturated solution is in contact with solid solute starts depositing the solute so as to reach
the concentration of a saturated solution. Thus in a supersaturated solution the concentration of
the solute falls when come in contact with the solute.
The solution which is dissolving some more solute in it than the solute that can dissolve is called
super saturated solution.
Activity
To test whether a given solution of a substance is unsaturated, saturated, or supersaturated, at
a particular temperature. Prepare an unsaturated, saturated and a supersaturated solution of
common salt separately in three beakers. Now, add a small amount of the crystal of common
salt turn by turn to each of the solution.
1. If it dissolves and the concentration of the solution increases, it is unsaturated.
2. If it does not dissolve even on vigorous stirring and the concentration of the solution
remains the same, it is saturated.
3. If it grows in size and the concentration of the solution falls, it is supersaturated.
Scan for practical experiment
visit: csp.codes/c09e10
Blooming Science Book 9 155
Dilute and Concentrated Solutions
The quantity of solute dissolved in a certain quantity of solvent, by weight or volume denotes
the concentration of the solution. Accordingly two types of solutions are known dilute and
concentrated solution.
A solution containing relatively small amount of solute in a fixed amount of solvent or compared
to that of the solvent is a dilute solution. The solution made by mixing a teaspoon of sugar in a
cup of water.
Solution containing relatively more quantity or large amount of solute in the fixed amount of
solvent is a concentrated solution. The solution made by mixing three teaspoons of sugar in a
cup of water.
Activity
Take two beakers. Put equal amount of water in two beakers. Put a spoonful of sugar in
one beaker and three teaspoons in the other. Stir with the glass rod. Take these prepared
solutions separately. The solution containing three teaspoonful of sugar is more sweeter than
the solution with one teaspoon of sugar.
Solubility
When a substance dissolves in water or in other liquid (solvent), it is said to be soluble in it. The property
of a substance by which it tends to dissolve in water (or any other solvent) is called its solubility.
Activity
To determine the solubilities of potassium sulphate, sodium chloride and sodium nitrate at
room temperature:
1. Take 100gm of water in each of three separate beakers at room temperature.
2. Take the weights of all the three beakers with water separately.
3. Add potassium sulphate to the first, common salt to the second and potassium nitrate
to the third beaker.
4. Stir well to dissolve the required solute to make each solution saturated.
5. Take the weight of each beaker containing saturated solution.
6. Calculate the amount of solute added in each beaker containing 100 gms of water to
make it a saturated solution at room temperature.
Calculation In 100gm of water about 210gm ?DO
of sugar can dissolve in room
Weight of the solvent in each beaker = x gm temperature. At 100oC about 500gm You
sugar disolves in 100gm water. Know
Weight of the saturated solution in each
beaker = y gm
∴ The weight of the solute = (y - x)gm.
156 Blooming Science Book 9
In other words, the difference in the weight of the two is the amount of a solute required to form
a saturated solution at room temperature. If the above activity is done properly, following result
will be obtained.
Solute Amount of solvent (water) Amount of solute Solubility
Potassium sulphate 100 gms 12 gms 12
Sodium chloride 100 gms 35 gms 35
Potassium nitrate 100 gms 60 gms 60
At room temperature
The result of this activity shows that for different substances, the amount of solute needed to
make a saturated solution in the same solvent is different at a particular temperature.
The amount of solute (in gms) required to form a saturated solution in 100 gms of solvent at a
particular temperature is called the solubility of the substance at that temperature.
The above activity shows that solubility of sodium chloride is 35 at 20oC, that of potassium
nitrate is 60 at 20oC and the potassium sulphate is 12 at 20oC.
Here, Solubility = Wt. of solute (in gm) × 100, at a particular temperature
Wt. of solvent (in gm)
Different substances have different solubilities. Solubilities of some common substances in water
at 20oC and 30oC are given in table below.
Solubilities of Some Common Substances
Solute Solubility
Copper sulphate at 20oC at 30oC
Sodium nitrate 21 25
Sodium chloride 88 95
35 37
Above table shows clearly that the solubility of the solute increases with the rise in temperature,
but the solubility of sodium chloride increases slightly with the rise in temperature.
Solved Numerical Problems
1. The solubility of a solute at 30oC is 40. What amount of water is required to make
saturated solution of 80gm of a solute?
Solution:
Weight of a solute = 80 gm
Solubility at 30oC = 40
Weight of solvent (water) = ?
Blooming Science Book 9 157
According to the formula,
Solubility = Wt of solute (in gm) × 100, at 30oC
Wt. of solvent (in gm)
Weight of solvent (gm) = Wt of solute (in gm) × 100, at 30oC
Solubility
= 80 × 100 = 200 gm
40
∴ 80 gm of solute needs 200 gm of solvent (water) to form saturated solution at 30oC.
2. 7 gm of saturated solution of salt saturated at 60oC is evaporated to dryness; 2 gm of
white residue is left behind. What is the solubility of salt at that temperature?
Solution:
Weight of saturated solution = 7gm
Weight of solute (Salt) = 2gm
Solubility of salt at 60oC = ?
To find the weight of solvent (water),
Weight of Solution = Weight of Solute + Weight of Solvent
Weight of Solvent = Weight of Solution - Weight of Solute
Weight of Solvent = 7gm - 2gm = 5gm
To find solubility, according to the formula,
Solubility = Wt of solute (in gm) × 100, at 60oC
Wt. of solvent (in gm)
= 2 = 40
5
∴ The solubility of salt at 60oC is 40.
3. At 30oC, 7 gram of sugar dissolves in 5 gram of water to form a saturated solution.
Find the solubility of sugar?
Given,
Mass of solute (sugar) = 7 gm
Mass of solvent (water) = 5 gm
Solubility at 30oC = ?
According to the formula,
158 Blooming Science Book 9
Solubility = Wt of solute (in gm) × 100, at 30oC
Wt. of solvent (in gm)
= 7 gm × 100
5 gm
= 140
∴ The solubility of sugar at 30oC is 140.
4. How much copper sulphate will be precipitated out when 30 gram of its saturated
solution is cooled down from 30oC to 20oC? (Solubility of the salt at 30oC and 20oC is
35 and 21 respectively.
Solution,
Weight of saturated solution of copper sulphate = 30 gm
Solubility of copper sulphate at 30oC = 35
Solubility of copper sulphate at 20oC = 21
For 30oC
Amount of solute (copper sulphate) = x gm
Amount of solvent = (30 - x) gm
Solubility of copper sulphate = 35
According to the formula,
Solubility = Wt. of solute × 100, at 30oC
Wt. of solvent
or, 35 = x
(30 - x)
or, 35 (30 - x) = 100x
or, 1050 - 35x = 100x
or, 1050 = 135xv or, x = 1050 = 7.78 gm
135
Therefore, weight of solvent in 30 gm of saturated solution at 30oC
= Weight of saturated solution - Weight of solute (x)
= 30 gm - 7.78 gm = 22.22 gm
Again, for 20oC, (the weight of the solvent remains the same)
Blooming Science Book 9 159
Given,
Weight of solute = y gm.
Amount of solvent = 22.22 gm
According to the formula,
Solubility = Wt. of colute × 100, at 20oC
Wt. of solvent
or, 21 = y × 100
22.22
or, 21 x 22.22 = 100 y
or, 100 y = 466.67
or, y = 466.67
100
= 4.67
Therefore, the weight of copper sulphate that will precipitate out
= Weight of copper sulphate (x) at 30oC - Weight of copper sulpahte (y) at 20oC
= (7.78 - 4.67) gm = 3.11 gm
∴ 3.11 gm of copper sulphate will be precipitated out.
Effect of Temperature on Solubility
Molecules in liquid are less closely packed and the intermolecular forces of attraction between
the molecules are weaker than those in solids. The molecules of the liquids are constantly moving
in different directions with different speeds. Thus, on heating, the kinetic energy of the molecules
will increase more and move faster.
Molecules of the solid solute are held together by strong intermolecular force of attraction. So the
molecules in solids are closely packed. When the temperature is increased, the kinetic energy of
the molecules will increase which makes the molecules to move fast.
Similarly, on stirring the solution with a spoon or a glass rod helps to increase the kinetic energy
of the molecules of solvent. Thus, on heating or stirring the kinetic energy of the molecules of
the solute increases. The molecules of the solute strike each other and get separated from each
other. These separated solute molecules get mixed with the solvent molecules to form a solution.
Molecules of solvent
Molecules of solute
160 Blooming Science Book 9
When the molecules vibrate more at high temperature, the intermolecular force of attraction
between the molecules weakens and the spaces between them increases. That’s why warm water
dissolves more solute than cold water. Thus the solubility of the substances increase with the rise
of temperature.
Molecules moves Heat Molecules moves
slowly fast
Inter molecules space
Solution at low temperature Solution at high temperature
Thus, the solubility of any solid in a solvent depends on temperature. The solubilities of most
solutes like potassium nitrate, copper sulphate, ammonium chloride etc. increase with increase in
temperature. The solubilities of sodium chloride and potassium chloride etc increase slightly with
an increase in temperature. Solubilities of some substances in water such as calcium sulphate,
calcium hydroxide and sodium sulphate decrease with the rise of the temperature.
Solubility Curve
The variation in the solubility of any given substance with change of temperature is shown by
solubility curve. The curve line drawn in graph showing the relationship between temperature
and solubility of the substance at different temperature is called a solubility curve.
To draw solubility curves, temperature is represented along the X-axis and solubility along the
Y-axis. Various solubility points plotted are connected by a smooth curve which is a solubility
curve. The solubility of copper sulphate at different temperatures are given in table below.
Solubilities of Copper sulphate at different temperatures.
Temperature 0oC 0 10 20 30 40 50 60 70
Solubility 14 17 21 24 29 34 40 47
The solubility curve drawn for the above data is shown in the figure below.
Solubility (grams solute 50
per 100 gm of water) 45
40
35 Scan for practical experiment
30
25
20
15
10
0 10 20 30 40 50 60 70 visit: csp.codes/c09e11
Temperature oC
Blooming Science Book 9 161
The solubility curve of three typical substances are shown in figure below.
130
120
110
100
Solubility90
Potassium chloride80
70 Magnesium chloride
60
50
40 Sodium chloride
30
20
10
10 20 30 40 50 60 70 80 90 100
Temperature oC
The solubility of sodium chloride is seen to raise little with rise of temperature. The solubility of
lead nitrate, potassium nitrate and sodium nitrate increases rapidly with increase of temperature.
The solubility of calcium sulphate and calcium hydroxide decreases with rise of temperature.
Uses of Solubility Curves
The general shape of the curve indicates the rate of change in the solubility with a rise in
temperature. A steep curve e.g. that of potassium nitrate shows that solubility increases rapidly
with rise of temperature.Aflat curve, e.g., that of sodium chloride indicates that solubility increases
slowly with the rise of temperature. The solubility of given substance at a given temperature can
be determined from its solubility curve. Solubility curve can be used to determine the amount of
substance deposited when the solution is cooled. Solubilities of different substances at a particular
temperature can be determined. The substance to be crystallized out first from a solution of two
or more solution can be predicted.
Information obtained from Solubility Curves
After studying the solubility curve, the following information can be obtained,
1. The solubility of a substance at a particular temperature can be determined.
2. The solubility of a given substance at any temperature can be determined.
3. The solubility curve helps us to predict which substance will crystallize out first from a
solution containing two or more solutes.
4. The solubility curve helps us to compare the solubilities of different subtances at the same
temperature.
5. It brings the change in the composition of a solute subtance.
6. It gives a clear idea that solubility of substance changes with the temperature.
162 Blooming Science Book 9
Crystals and Crystallization
Let us spread some sugar or common salt on a paper and observe it through a magnifying glass
or naked eyes. Each piece of sugar or common salt can be seen in the form of regular geometrical
shape. This kind of a particle is a crystal. A crystal is a solid substance bounded by plane faces
and having sharp edges. It has regular and geometric shape. Crystals are formed by a process
called crystallizsation. Crystals are formed due to the systematic arrangement of atoms or
molecules. The person who studies about crystal is called a crystallographer. Crystals are formed
by different methods like cooling of saturated solution, sublimation and solidification of fused
substances. They can be found in different shapes like cubical, tetragonal, orthorhombic etc.
Crystallization is the best method of getting purest form of a substance.
A hot saturated solution contains more of a given solute than does a cold saturated solution. So
when a hot saturated solution is cooled, some of the dissolved solid is thrown out of the solution
in the form of crystals. The crystals thus obtained from the saturated solution are purer than
the initial substance dissolved. This is because the solution is saturated only with respect to the
principal substance and unsaturated with respect to the impurities present in small amount in
the solution. Hence ,the principal substance crystallizes out and the impurities are left behind in
the remaining solution. Thus, crystallization is an excellent method for purification of soluble
substances containing smaller amounts of soluble impurities.
The crystals are glistening solid particles having a definite geometrical shape, flat sides and sharp
edges. The process of deposition of crystals from a saturated solution is called crystallization.
Activity
Saturated Crystal
solution
Saturated
Heat solution
Some amount of water is taken in a beaker and copper sulphate is mixed with it. It is stirred
with a glass rod. At the beginning, the copper sulphate gets dissolved easily. Some more
copper sulphate is added in the beaker and the solution is stirred. If we go on adding sulphate
and stirred it, finally it is found that no more copper sulphate is dissolved in the solution. If
the solution is heated and more copper sulphate is added, the copper sulphate gets dissolved
again. Now, the solution is placed in a beaker containing cold water and let it to cool. After
a few minutes, the crystals of copper sulphate are found to be formed on the surface of the
water.
To get well crystals, (i) saturated solution should be cooled slowly, (ii) slow evaporation and
slow cooling should be done.
Blooming Science Book 9 163
Conditions of Crystallization
To obtain well defined crystals, following conditions of crystallization should be known.
1. Saturated solution should be cooled down slowly.
Rapid cooling of the hot saturated solution results in the deposition of fine crystals. Slow
cooling produces few but larger crystals.
2. A cold saturated solution should not be heated to dryness to obtain crystals. Because after
evaporation the residue obtained is an amorphous mass.
3. Slow evaporation of a saturated solution produces crystals so, cooling and evaporation
must be as slow as possible.
In short, well defined crystals are formed Glass rod
only when a hot saturated solution is Thread
allowed to cool or evaporated very slowly Beaker
without being shaken.
Copper sulphate crystals
Large crystals of green vitriol, potash and Copper sulphate solution
alum can be obtained in the similar way.
All substances do not form crystals. Thus, there are two types of solid substances. They are
crystalline and amorphous.
Crystalline Solids
These solids have typical geometrical shapes along with definite and rigid structure. These solids
have fixed melting and boiling points. Sodium chloride and copper sulphate are crystalline solids.
Amorphous Solids
These solids do not have any definite geometrical forms or shapes. Rubber and plastics are
amorphous solids.
Let’s Learn
1. When an unsaturated solution is cooled, it becomes saturated. It is because when the
solution is cooled, the intermolecular space decreases and less particles of solute can be
adjusted.
2. Temperature plays an important role in solubility because the solubility of substance is
different at different temperature.
3. Solution is called homogeneous mixture because the solute particles can not be separated
with our naked eyes.
164 Blooming Science Book 9
Main Points to Remember
1. Mixture is a combination of two or more than two substances combined in any proportions.
2. On the basic of combining particles, homogeneous and heterogeneous are two types of mixture.
3. On the basic of size of solute particles, mixture is classified as solution, colloids and suspension.
4. A solution in a homogeneous mixture of solute and solvent.
5. A solution in which excess solute does not dissolve is called a saturated solution at a given
temperature.
6. A solution in which excess solute can dissolve is called an unsaturated solution at a given
temperature.
7. A supersaturated solution at a given temperature is a solution with excess solute than
required to make saturated solution at that temperature.
8. The solubility of a substance at a given temperature is defined as the amount of solute
dissolved in 100gm of solvent to form a saturated solution at that temperature.
9. A solubility curve is a line obtained on a graph paper by plotting the solubility of a
substance at different temperatures.
10. A crystal is a solid substance bounded by plane faces and having sharp edges.
11. Crystals are formed by a process called crystallization.
PRO J ECTWORK
Prepare unsaturated and saturated solution of edible salt in two glasses with equal volume
of water and taste both solution. Compare the taste (amount salt) and discuss why does this
happen?
Exercise
A. Choose the correct answer from the given alternatives:
1. Organic solvent dissilves...................solute.
a. inorganic b. organic c. chimical d. all of them
2. In which solution crystals are formed?
a. unsaturated b. saturated c. concentrated d. supersatuated
3. What is the relation of solubility of substance which temperature?
a. directly b. Indirecty c. equal none of them
Blooming Science Book 9 165
4. The formula to calculate solubility of a substance is
a. wt. of solute × 100 b. wt. of solute × 100
wt. of solution wt. of solvent
c. wt. of solvent × 100 d. wt. of solvent × 100
wt. of solution wt. of solute
5. Blood is an example of ...........................
a. wlloids b. suspension c. solution d. mixture
B. Answer the following questions.
1. Define the following terms:
Mixture, solution, solvent, solute, supersaturated solution, solubility and solubility
curve.
2. Write any two differences between:
(i) Dilute solution and concentrated solution
(ii) Saturated solution and unsaturated solution
(iii) Solute and solvent
(iv) Colloid and suspension
3. What is saturated solution? How does temperature affect it? Explain with examples.
4. What information can be obtained from solubility curves?
5. Write the importance of solution.
6. You are given a beaker, salt, glass rod, water, a thermometer. How would you make
saturated solution. Explain in brief with a diagram.
7. What are crystals? How are they formed? How can you obtain big crystals?
8. In what respect does a mixture of sugar and water differ from a mixture of sand and
sugar?
9. How does a homogeneous mixture differ from a heterogeneous mixture?
10. When excess solute is mixed in solutions A, B and C, you will get the following results.
The excess salt disappears in A, excess salt remains same in B, and the size of excess salt
increase in size in C. Distinguish the solutions given in the figures below.
Salt C
AB
Water Water Water
Beaker Beaker Beaker
166 Blooming Science Book 9
11. Explain the types of solid substances.
12. Draw a solubility curve of ammonium chloride form the give data
Temperature (oC) 0 10 20 30 40 50 60 80 100
Solubility (gm) 24.4 32.8 37.3 41.3 46.2 50.6 55.0 64.0 72.8
Also find out the solubility at 25oC and 65oC.
13. Why temperature always written with solubility of a substance?
Numerical problems
(i) At 30oC, 7 gm of sugar dissolves in 5 gm of water to from a saturated solution find
solubility of sugar at 30oC. (140)
(ii) At 30oC, 60 g of ammonium nitrate is dissolved in 25 g of water to form a saturated
solution. Find out the solubility of ammonium nitrate. (240).
(iii) 560 g of common salt is dissolved in 1805 g of water at 10oC to form a saturated
solution. Calculate solubility. (31.02)
(iv) The solubility of substance at 30oC is 140 and weight of solute is 300 gm. Find the
weight of the solvent. (214.28 gm)
(v) The weight of saturated solution of a substance at 25oC is 2.4 kg and the weight of
solvent is 2 kg. Find out the solubility. (20)
(vi) At 10oC, solubility of lead nitrate is 55. Find the wt. of lead nitrate to dissolve in
200 g of water to form a saturated solution. (110)
(vii) The solubility of copper sulphate at 30oC is 25. What do you mean by it?
(viii) 25g of saturated solution of salt at 30oC is evaporated to dryness. 10g of residue is
left behind. What is the solubility of salt at 30oC? (66.7)
(ix) The solubility of a substance at 70oC is 50 and at 20oC is 30. 80g of water is needed
to form saturated solution with salt at 70oC is cooled to 20oC. Calculate the wt. of
the salt thrown out. (16 gm)
(x) 24g of water needed to form saturated solution of silver nitrate at 70oC is cooled to
15oC. Calculate the wt. of the salt thrown out. The solubility at 70oC and 15oC are
525 and 196 respectively. (78.96 gm)
Blooming Science Book 9 167
Chapter SOME GASES
11
Learning Outcomes Estimated Periods: 5+3
On the completion of this unit, the students will be able to:
• describe the laboratory preparation and properties of some gases like hydrogen,
oxygen and nitrogen.
• explain the utility of above gases.
There are eleven elements which are found in a gaseous state. Some of them are hydrogen,
nitrogen, oxygen, inert gases like helium, neon, argon, krypton, etc. Among these gases, hydrogen
is the lightest gas known. There are some compounds which are found in the gaseous state
in nature. Some of them are carbon dioxide, water vapour, ammonia, etc. Gases like nitrogen,
oxygen, inert gases and water vapour mix together in different volumes to form air. The thick
layer of air which surrounds the earth is called atmosphere. The main gases in the air, which
occupy about 90% volume of the air, are nitrogen and oxygen. The composition of different
gases in the air is tabled below:
Composition of Air Percentage by volume Nitrogen Oxygen
78.07 Composition of air
S.N. Gases in air 20.98
1. Nitrogen 0.03
2. Oxygen 0.85
3. Carbon dioxide 0.002
4. Argon 0.008
5. Neon 0.06
6. Other inert gases
7. Water vapour
Without oxygen no living thing can exist in nature. They breathe in oxygen and breathe out carbon
dioxide. Nitrogen is used by plants in the form of fertilizer. Plants take carbon dioxide and give out
oxygen during photosynthesis. Carbon dioxide is formed in the air also by the combustion of fuel.
In this way, these different gases are interchanged between air and living things and their content
remains in balanced form in the atomosphere. The content of these gases is always constant in the
air. But the amount of moisture and carbon dioxide may vary from place to place. In this class, we
will study the general method of preparation, properties and uses of some gases.
168 Blooming Science Book 9
Hydrogen
Symbol : H Atomic number: 1
Molecular formula: H2 Atomic weight: 1.008 a.m.u
Valency: 1
Molecular weight: 2.016
Position in periodic table: Group - I‘A’ Melting point: -259°C
Period - 1 Boiling point: - 253°C
Electronic configuration : 1 (1s1)
1p+ 1p+
1p+
Atomic model of hydrogen Molecular structure of hydrogen
Hydrogen was first proved to be distinct element by an eminent English scientist, Henery
Cavendish in 1766 AD and named it inflammable gas. In 1783 AD, Lavoisier proposed the name
hydrogen to it.
Occurrence
Hydrogen is a reactive element and thus does not occur much in free state. It is found in volcanic
gases and natural gases in trace amount. But, it occurs abundantly in combined state.
In combined form, it is an important constituents of water, acid, alkali and many organic
compounds of vegetables and animal products. The chief source of hydrogen is water in which
its amount is double than that of oxygen.
General Methods for the Preparation of Hydrogen Gas
Hydrogen gas can be prepared by the action of metals with acids, alkalis and water respectively.
1. From Acids: Metals like zinc, iron, magnesium, etc. are more electropositive than
hydrogen and react with acid to produce hydrogen gas.
Zn + dil.H2SO4 → ZnSO4 + H2↑
Metals like silver, copper, etc. are less electropositive than hydrogen. So these metals do
not produce hydrogen gas from acids.
2. From Alkalis: Hydrogen gas can be obtained from the action of metals like zinc,
aluminium, etc. with boiling caustic soda.
Zn + 2NaOH → Na2ZnO2 + H2↑
Sodium zincate
Blooming Science Book 9 169
3. From Water: At ordinary temperature, metals like sodium, potassium, calcium, etc. react
with water to liberate hydrogen gas. Hydrogen gas has 3 ?DO
2Na + 2H2O → 2NaOH + H2↑ isotopes, 1H1,1H2 and 1H2 is
Preparation of Hydrogen Gas in Laboratory You
normal hydrogen with one Know
proton and one electron.
Principle: When granulated zinc reacts with dilute 1H2 is called deuterium
sulphuric acid, they react together to form zinc with one proton and one
sulphate and hydrogen gas. The principle reaction neutron.
is as follows
Zinc + Sulphuric acid → Zinc Sulphate + 1H3 is called tritium with
one proton and 2 neutron.
Hydrogen
Zn + dil H2SO4 → ZnSO4 + H2↑
Thistle funnel
Dilute
Sulphuric Acid
Delivery Hydrogen Gas
tube Gas jar
Scan for practical experiment
Woulfe’s Pneumatic
Bottle trough
Bee-hive visit: csp.codes/c09e12
shelf
Granulated Zinc
FIg: Laboratory preparation of hydrogen gas
Take a few grains of granulated zinc in Woulfe’s bottle fitted with a thistle funnel and delivery
tube with corks. Put the next end of the delivery tube under water in a pneumatic trough having
beehive shelf. Invert a gas jar, completely filled with water over the beehive shelf and let the end
of the delivery tube into it. Now pour dilute sulphuric acid through the thistle funnel till it covers
the pieces of zinc and lower end of the funnel dips in the acid. The apparatus should be air tight.
A brisk action sets in and the hydrogen gas evolves. The first formed gas is contaminated with air
inside the Woulfe’s bottle and delivery tube and is allowed to escape. Then the evolved hydrogen
gas gets collected in the gas jar by the downward displacement of water. The zinc sulphate is left
in the Woulfe’s bottle in the form of solution.
Note: If the reaction between zinc and acid is very slow, we should add little copper sulphate
solution to accelerate the chemical reaction between them. Here, copper sulphate acts as a catalyst.
Precautions
1. Impure zinc should be used instead of pure zinc because the reaction between the pure
zinc and the dilute sulphuric acid is very steady whereas impurities present in the zinc
increase the rate of reaction.
170 Blooming Science Book 9
2. Concentrated sulphuric acid should not be used because it produces sulphur dioxide
instead of hydrogen gas.
3. The end of thistle funnel must be under the acid in the Woulfe’s bottle.
4. The apparatus should be made airtight.
5. The end of the delivery tube with beehive shelf should be under the level of water in the
trough.
6. The fitted apparatus for the experiment should be kept far away from fire after pouring acid in
the Woulfe’s bottle otherwise, hydrogen itself burns and may explode when mixed with air.
Test of Hydrogen
Hydrogen burns in air with a faint pale-blue flame. To check whether the produced gas hydrogen
or not, when a lighted splinter is introduced to a mouth of the gas jar, the gas burns itself
with a very faint pale-blue flame at the mouth of the jar with pop sound and the splinter gets
extinguished. Thus, we can conclude that the produced gas is hydrogen because it is combustible
but not a supporter of combustion.
Manufacture of Hydrogen
Hydrogen gas is used for many purposes. So, it is manufactured in large scale. Usually the
following two methods are used for the manufacture of hydrogen:
1. From Methane-Steam Process: When a mixture of steam and methane (paraffin
hydrocarbon) is passed over heated nickel catalyst at 1200oC, and compressed to 30
atmosphere, hydrogen gas is manufactured. Methane is obtained as by-product of
petroleum industry.
CH4 + H2 1200oC CO + 3H2↑
30 atm/Ni
Note: Catalyst is defined as a chemical substance which alters (increase or decrease)
the rate of speed of chemical reaction, itself remaining chemically unchanged
because it does not take part in the chemical reaction. The phenomenon is known
as catalysis. Water
2. By the Electrolysis of Water: Hydrogen Hydrogen Oxygen
gas is manufactured by the electrolysis of
water where the electrical energy is supplied.
For the electrolysis, small amount of dilute H2SO4 + water
acid or alkali is poured into a voltameter Anode
containing water to make strong electrolyte Cathode
as shown in the figure. In the electrolytic
cell or voltameter iron is used as cathode Asbestos
while the nickel-plated iron acts as anode. Diaphragm
An asbestos diaphragm separates these two
electrodes from each other. This diaphragm Electrolysis of water
prevents the mixing of hydrogen gas and oxygen gas. When electric current is passed, hydrogen
is collected at cathode and oxygen at anode.
When dilute H2SO4 is added to water for electrolysis, the ionization of water takes place
which is as follows:
Blooming Science Book 9 171
[H2O H+ + OH-] × 4
At Cathode
4H+ + 4e- → 4H → 2H2↑
At Anode
4OH- → 4OH + 4e-↑
4OH → 2H2O + O2↑
Nascent (Newly Born) Hydrogen
The atomic form of hydrogen produced at the time of chemical reaction is known as nascent
hydrogen. The name nascent has been derived from newly born. These nascent type of species
are very unstable because (i) they have high reactivity (ii) they are in the high pressure (iii) they
possess high internal energy. So, the produced nascent species immediately combine themselves
to produce molecular forms of hydrogen which is less reactive than the nascent species.
Zn + H2SO4 → ZnSO4 + 2H (nascent)
H + H → H2
Properties of Hydrogen
Physical Properties
1. Hydrogen is a colourless, tasteless and odourless gas.
2. It is very slightly soluble in water.
3. It is lightest gas known. 22.4 litres of this gas weight only 2 grams.
4. It is neutral to litmus.
5. It liquefies at -253oC and solidifies at -259oC.
Chemical Properties
1. Hydrogen burns in air or in oxygen with a blue flame and it produces water.
2H2 + O2 ∆ 2H2O
2. Metals like sodium, potassium or calcium burn in hydrogen forming corresponding
hydrides.
2Na + H2 30oC 2NaH
Sodium Hydride
Ca + H2 30oC CaH2
Calcium Hydride (hydrolith)
3. When dry hydrogen is passed over heated oxides of iron, copper, lead, etc. it reduces the
metallic oxides to their metals.
Fe3O4 + 4H2 ∆ 3Fe + 4H2O
CuO + H2 ∆ Cu + H2O
PbO + H2 ∆ Pb + H2O
172 Blooming Science Book 9
In the above reactions, the removal of oxygen from the compound is called reduction and
hydrogen which brings about reduction is called reducing agent. The compound from
which oxygen is removed is said to have reduced. But the oxides of calcium, zinc and
magnesium are not reduced by hydrogen.
4. Hydrogen reacts with chlorine, bromine, iodine and fluorine (halogens) in different
conditions to form their acids.
H2 + F2 in darkness 2HF
(Hydrofluoric Acid)
H2 + Br2 sunlight 2HBr
(Hydrobromic Acid)
H2 + I2 40oC 2HI
(Hydroidic Acid)
5. Hydrogen combines with nitrogen under 200-900 atmospheric pressure, at 500oC
temperature and in the presence of a catalyst iron with molybdenum to give ammonia.
500oC, Fe, Mo
N2 + 3H2 200 - 900 atm 2NH3
Uses of Hydrogen
1. It is used in manufacturing of ammonia. Ammonia is used for the manufacture of fertilizers.
2. Hydrogen gas is used in manufacturing of vansapati ghee. To manufacture the vanaspati
ghee, hydrogen gas is passed through vegetable oils, under 8-10 atmospheric pressure in
the presence of nickel powder at temperature of 200oC, oil changes into vegetable ghee or
solid fat known as vanaspati ghee. This process is called hydrogenation. Hydrogenation
is a process in which unsaturated compounds combine with hydrogen in the presence of
catalyst, and in other suitable conditions and are converted into saturated compounds.
200oC, Ni
Vegetable oil + H2 8 - 10 atm Solid fat Vanaspati ghee (Saturated)
3. Hydrogen is used as a reducing agent because it reduces metallic oxides into corresponding
metals.
4. The mixture of hydrogen and oxygen burns to produce a high temperature of about
3000oC. Hence, hydrogen is used for the production of oxy-hydrogen flame for welding
and cutting of metals.
5. It is used for filling balloons because it is the lightest gas. As hydrogen is highly
combustible, helium is mixed with it to fill balloon.
6. The liquid form of hydrogen is used as fuel in rockets.
Blooming Science Book 9 173
Let’s Learn
1. Zinc is mostly used for the laboratory preparation of hydrogen gas rather than other metals
because of the following facts:
a. Metals like sodium and potassium react violently with acid.
b. Calcium and magnesium are very expensive in comparison to zinc.
c. Aluminium forms a protective coating of Al2O3 whereas iron reacts with acid
very slowly and requires more heat.
2. Metals displace the hydrogen from acids because metals are more reactive or more
electropositive than hydrogen.
3. Hydrogen gas is not found in air because it is more reactive and the lightest gas known.
Main Points to Remember
1. Air is a mixture of different gases like nitrogen, oxygen, carbon dioxide, inert gases, water
vapour, etc.
2. Hydrogen is the lightest gas known, so it is found only in combined state and not in air.
3. Hydrogen gas can be obtained from the action of metals with acids, alkalis and water.
4. Hydrogen gas is prepared in the laboratory by treating zinc with dilute sulphuric acid.
Zn + dil H2SO4 ZnSO4 + H2↑
5. Hydrogen gas is collected by the downward displacement of water in the laboratory as it
is the lightest gas known.
6. Hydrogen is a combustible gas but not a supporter of combustion.
7. Hydrogen is a colourless, odourless, and tasteless gas. It can be liquefied or solidified. It
is insoluble in water and it is neutral to the indicators.
8. Hydrogen burns with oxygen to produce water. This property is known as combustion.
2H2 + O2 ∆ 2H2O
9. Hydrogen reacts with nitrogen at suitable conditions to produce ammonia.
10. Hydrogenation is a process in which unsaturated compounds combine with hydrogen in
the presence of catalyst and saturated compounds are produced.
Exercise iv. Neon
A. Choose the correct answer from the given alternatives:
1. Which is the lightest gas?
i. Hydrogen ii. Nitrogen iii. Helium
2. Which is the most abundant gas in atmosphere?
i. O2 ii. H2 iii. N2 iv. CO2
174 Blooming Science Book 9
3. Principle reaction of preparation of hydrogen gas in labrotory is :
i. 2H2O electricity 2H2↑+O2↑ ii. Zn+dil.H2So4 ZnSo4+H2↑
iii. Mg+H2O (steam) Mgo+H2↑ iv. None of above
4. Complete:
2K+H2O .................+H2↑
i. K2OH ii. KOH iii. 2KOH iv. K(OH)2
biocatalyst
6CO2+6H2O+.................
5. Complete: C6H12O6
i. Power ii. Energy iii. Light iv. Sound
B. Answer the following questions.
1. Write the composition of air.
2. Answer the following questions after observing the figure.
a. Which gas is going to be prepared in the figure?
b. What are the mistakes in the arrangement? Correct the figure.
c. What happens if Conc acid is used?
Thistle funnel
Dilute
Sulphuric Acid
Delivery Hydrogen Gas
tube Gas jar
Woulfe’s Pneumatic
Bottle trough
Granulated Zinc Bee-hive
shelf
3. Write correct balanced chemical equation and well labeled diagram for the laboratory
preparation of hydrogen from zinc and dilute sulphuric acid.
4. What would you observe when
a. a piece magnesium ribbon is added to dilute sulphuric acid.
b. a mixture of hydrogen and chlorine is made.
c. a mixture of hydrogen and oxygen is heated.
d. Hydrogen gas is passed over heated copper (II) oxide.. Explain briefly how
hydrogen is manufactured from steam on a large scale.
Blooming Science Book 9 175
6. What are the physical properties of hydrogen?
7. State one property of hydrogen
a. that made it useful in meteorological balloons.
b. that makes it a good fuel.
8. What is a catalyst? Name the catalyst used in the manufacture of hydrogen gas.
9. What is hydrogenation? Where is this process used? Name the catalyst used in the process.
10. What do you mean by reduction property of hydrogen? Explain.
11. Fill in the blanks in the following word equations by using correct formulae.
a. Zinc + ……..→ Zinc chloride + Hydrogen
b. ……….. + Water → Carbon monoxide + Hydrogen
c. Copper oxide + Hydrogen → …………… + Water
d. Lead oxide + ………….. → Lead + Water
12. Fitted apparatus for a gas preparation in the
laboratory is shown below. Study the diagram and
answer the following questions:
a) Which gas is being collected in a gas gar?
b) Write and equation of the chemical reaction
that takes place during gas preparation.
c) What happens when a burning match stick
is held near the gas jar?
d) What is mean by granulated zinc? What is its use?
e) What happens if hot concentrated sulphuric acid is used instead of dilute sulphuric
acid?
f) How is the gas collected? Explain why this method is used.
g) List two precautions to ensure safe preparation and collection of the gas.
h) How would you confirm that the gas collected is hydrogen?
i) Why is zinc used?
176 Blooming Science Book 9
Oxygen
Symbol: O Atomic Number: 8
Atomic Mass: 16
Molecular Formula: O2 Molecular weight: 32
Valency: 2
Position in periodic Table : Group: VIA, Period: 2, Block: P
Electronic Configuration: 2, 6 (1s2, 2s22p4)
Introduction
Oxygen gas was first prepared by British Chemist Joseph Priestely in 1823 AD. He prepared it
by the action of heat on red mercuric oxide. Later Lavoisier studied its properties in detail and
showed that one fifth of the air is oxygen. He named this gas as oxygen.
Oxygen is the most abundant of all element on the earth. It Ozone gas is made ?DOYou
constitutes 46.6% of the earth’s crust. It is found in nature both from three atoms of Know
free in air and in combined state. It is also found dissolved
in water. In combined state, it is found in water and in many oxygen. It is written
minerals on the earth’s crust. It is also an important constituent
of all living matter. It is very important in respiration and as O3
combustion processes. About 72% of our body is composed
of oxygen compounds. It makes up about 21% by volume of
the air.
In free state, oxygen exists as a diatomic molecule. Each molecule contains a double covalent
bon (O = O).
8p+ 8p+ 8p+
8n 8n 8n
Oxygen molecule
Laboratory Preparation of Oxygen
By H eating
In the laboratory, oxygen is prepared by heating a mixture of potassium chlorate and manganese
dioxide in the ratio 4:1 in a hard glass test tube. Manganese dioxide acts as a catalyst.
2KClO3 360oC - 370oC 2KCl + 3O2 Scan for practical experiment
2KClO3 2KCl + 3O2
240oC - 250oC
MnO2
visit: csp.codes/c09e13
Blooming Science Book 9 177
Hard glass test tube Delivery Oxygen gas
Mixture of potassium chlorate tube Gas jar
and manganese dioxide
Bunsen burner
Trough
Water
Beehive shelf
Fig: Laboratory Preparation of Oxygen Gas
When Potassium Chlorate is heated, first it melts and then at about 360oC - 370oC, it decomposes
to liberate oxygen. If however, a little manganese dioxide is added, potassium chlorate gives off
oxygen steadily at a much lower temperature (240oC - 250oC).
Procedure
Potassium Chlorate is a white crystalline solid. It is soluble in water. Manganese dioxide is a
black solid. Powdered Potassium Chlorate and manganese dioxide are taken in the ratio of 4:1
mass in a hard glass test tube. The hard glass test tube is clamped with sliding down on mouth
side. The free end of the delivery tube is introduced into a bee-hive shelf placed under water in a
trough. The mixture in the tube is heated for about two minutes to displace the air inside the test
tube and then put a gas jar filled with water over the bee-hive shelf. The gas jar is filled to the
brim with water. It is covered with a well greased lid and lowered upside down into the trough
containing water before the lid is removed and placed over the bee-hive shelf. There should be
no air bubbles inside the water filled jar.
Continuous heating will produce the bubbles of oxygen gas rising up and displace the water in
the gas jar by downward displacement of water. When all the water has been pushed out, the gas
jar is raised, its mouth is covered with the gas plate while it is still under water and is taken out.
Precautions
a) The end of the delivery tube should be removed from the trough before heating is stopped.
(It is done to prevent the entry of water into the test tube.)
b) The test tube should be sloping down on the mouth side so that if moisture is condensed it
should not move to the hot part of the test tube. This will prevent the breaking of the tube.
Without Using Heat
The most convenient and quickest method of preparing oxygen gas in the laboratory is by adding
hydrogen peroxide solution to powdered manganese dioxide.
2H2O2 + MnO2 → 2H2O + O2 + MnO2
178 Blooming Science Book 9
Hydrogen Delivery tube Oxygen
peroxide Gas jar
Thistle Water
funnel
Conical
flask
Manganese dioxide
Fig: Laboratory Preparation of Oxygen Gas
Manganese dioxide is placed in a conical flask and the apparatus are fitted as shown in figure
above, on dropping the hydrogen peroxide from the thistle funnel over manganese dioxide,
hydrogen peroxide decomposes and releases oxygen gas. Manganese dioxide acts as a catalyst.
The gas is collected by the downward displacement of water in the gas jar since the gas is only
slightly soluble in water.
When the gas jar is filled with the gas, it is covered with a wall greased glass lid beneath the
surface of water in the trough and is removed.
Test: A burning splinter is brought close to the gas jar containing oxygen. When splinter is
inserted inside the gas jar, the splinter burns brightly which shows the presence of oxygen in gas
jar
To retard the decomposition of H2O2 by light, it is kept in brown coloured bottle. It is used
for bleaching silk, wool, leather, paper etc. It bleaches hair to a golden colour.
Manufacture of Oxygen
The cheapest raw materials for the production of oxygen are air and water.
From Liquid Air
The air freed from moisturised carbon dioxide is liquefied by compression, cooling and sudden
expansion. The liquid air is then allowed to vaporize. Nitrogen having a lower boiling point
(-195oC) escapes first. The residual liquid, which is nearly pure oxygen is then allowed to
vaporsie, boiling point (-183oC) collected and stored in steel cylinders.
By the Electrolysis of Acidulated Water
In the electrolysis of water, a direct current is passed through water to which a little sulphuric
acid is added. Hydrogen is collected at the cathode and the oxygen at the anode, their volumes
are in the ratio of 2:1. Both the gases are then stored under pressure in steel cylinders.
Blooming Science Book 9 179
Properties of Oxygen
Physical Properties
1. It is a colourless, odourless an tasteless gas.
2. It is slightly soluble in water (about 3%).
3. Liquid oxygen is pale blue in colour.
4. It is a neutral gas. It does not affect litmus and other indicators.
5. Combustion or burning is the combination of substance with oxygen. Oxygen does not
burn but helps things to burn i.e. supporter of combustion.
6. The gas liquefies at –183ºC and solidifies at –219ºC.
Chemical Properties
1. Reaction with Non-Metals
Oxygen also reacts with most non-metals to form covalent molecules. For example, glowing
carbon, burning sulphur and heated phosphorus burn brightly in oxygen forming their respective
oxides.
C (s) + O2 (g) ∆ CO2 (g)
S (s) + O2 (g) ∆ SO2 (g)
4P (s) + 5O2 (g) ∆ 2P2O5 (g)
Non-metals oxides are acidic oxides. Their solutions in water turn blue litmus red.
CO2 (g) + H2O (l) → H2CO3 (aq)
Oxygen combines with nitrogen at very high temperature during lightning discharge and form
nitric oxide.
Nitrogen + Oxygen heat Nitric oxide
N2 (g) + O2 2000oC - 3000oC 2NO (g)
The nitric oxide gas dissolves in rain water. Plants absorb it along with water and use it to make protein.
2. Reaction with Metals
Many metals, among them sodium, calcium and magnesium when strongly heated, burn in
oxygen forming their respective oxides.
[metal + oxygen = metallic oxide]
4Na (s) + O2 (g) ∆ 2Na2O
4K (s) + O2 (g) heat 2K2O
2Ca (s) + O2 (g) heat 2CaO (s)
2 Mg (s) + O2 (g) heat 2 MgO (s)
180 Blooming Science Book 9
With magnesium, it burns with a very bright dazzling light, leaving behind a white residue of
magnesium oxide. The oxides of the metals are basic oxides, because their solutions in water are
basic i.e. turns red litmus blue. These reactions are called oxidation reactions.
3. Reaction with Iron
At ordinary temperature, iron combines with oxygen and water very slowly forming hydrated
ferric oxide (Fe2O3, XH2O). It is formed as reddish brown scales on the surface of iron and is
called rusting of iron. Rusting is slow oxidation of iron when it is exposed to the moist air.
Iron + Oxygen + Water + → Hydrated ferric oxide
4Fe (s) + 3O2 (s) + XH2O → 2Fe2O3 . XH2O (s)
4. Combination with Organic Compounds (Combustion)
Many organic compounds like carbohydrates, ethyl alcohol, oil, petrol, wax, etc react with
oxygen to form carbon dioxide.
CH4 (g) + 2O2 (g) ∆ CO2 (g) + 2H2O (l)
C6H12O6 (s) + 6O2 (g) ∆ 6CO2 (g) + 6H2O (l)
[Fuel + O2 2H2O + CO2]
Uses
1. Oxygen present in air is supporter of life. It helps in oxidation of food to release energy in
living beings.
2. It is extensively used in hospitals for the artificial respiration of pneumatic patients.
3. Miners, mountaineers and deep-sea divers carry oxygen cylinders with them.
4. Liquid O2 is used as part of the fuel in rocket and jet planes to make the burning of fuel
more vigorous. Liquid O2 is preferred to gaseous O2 because it has much less volume
(greater density) for the same mass of oxygen.
5. The mixture of O2 and acetylene burns producing a very high temperature i.e. oxy-
acetylene flame. This flame is used in welding and cutting of metals.
6. Oxygen is also used in steel making to remove impurities.
Some Reasonable Facts
1. Oxygen is called supporter of combustion because it helps other substances to burn.
2. Oxygen helps in oxidation because it reacts with metals to form respective oxides
2Pb + O2 → 2PbO
3. Mountaineers carry oxygen cylinders because at high altitude there is not sufficient
oxygen and they use cylindrical oxygen for artificial respiration.
4. Manganese dioxide is used in preparation of oxygen because it is a catalyst which
reduces temperature and makes reaction faster.
Blooming Science Book 9 181
Main Points to Remember
1. Oxygen gas was discovered by Joseph Priestley and Lavoisier gave the name oxygen.
2. Atmospheric air contains about 20.95% of oxygen by volume.
3. Oxygen gas is prepared in the laboratory by heating the mixture of potassium chlorate and
manganese dioxide (catalyst) in the ratio 4:1.
4. Oxygen gas is also prepared in the laboratory without heating by using hydrogen peroxide
in the presence of manganese dioxide (catalyst). Oxygen gas is collected by downward
displacement of water.
5. Oxygen gas is colourless, odourless and tasteless gas.
6. It shows oxidation reaction.
Exercise
1. What is the percentage of oxygen in air? Why is this gas important?
2. What is the role of manganese dioxide in the laboratory preparation of oxygen?
3. State three main physical properties of oxygen gas.
4. State the importance of the dissolved oxygen in water.
5. What happens when the following elements burn in O2? Explain with chemical equation.
a) Magnesium b) Calcium
c) Carbon d) Iron
6. Complete the following chemical equations-
a) H2O2 (aq) → …………….. + O2 (g)
b) ……………. + O2 (g) → Fe2O3 (s)
c) KClO3 → ………….. + O2
d) CH4 + ……………….. → CO2 + H2O
7. What happens when?
a) potassium chlorate is heated.
b) burning charcoal is inserted into the gas jar containing oxygen.
c) potassium permanganate is heated.
d) iron nail is exposed to moisture.
e) mercuric oxide is heated.
9. What is oxy-acetylene flame?
182 Blooming Science Book 9
10. Draw a labeled diagram for the preparation of oxygen gas from the mixture of potassium chlorate
and manganese dioxide in the laboratory and give the chemical equation for the reaction.
11. Oxygen gas in prepared in the laboratory by catalytic decomposition of hydrogen peroxide.
a) What is meant by a catalyst? Give one catalytic reaction.
b) Name the catalyst usually used in the above gas preparation.
c) Give the balanced chemical equation for the above reaction.
d) Write two uses of oxygen gas.
e) State how oxygen is collected?
f) Draw a labeled diagram of the fitted apparatus of this experiment.
13. The diagram is an arrangement of an apparatus used for the preparation of oxygen gas in
the laboratory. Answer the following questions:
a) Name the substances A and B in the mixture.
b) Name an apparatus labeled X and Y.
c) Write the balanced chemical equation of the reaction that occurs.
d) Why is the hard glass test tube containing the mixture kept inclined?
e) How is the gas collected? Explain with reason.
X
Mixture of A and B
Y
14. Oxygen is called supporter of combustion. What does it mean?
Blooming Science Book 9 183
Nitrogen Gas
Symbol :N
Atomic No :7
Atomic mass : 14 a.m.u. 7p+ 7p+
Valency : 3, 5 7n0 7n0
Electronic configuration : 2, 5 (1s22s22p3) Atomic structure of nitrogen
Molecular formula : N2
Molecular mass : 28 a.m.u.
Nitrogen gas was discovered by a Scottist physician and chemist Daniel Rutherford in 1772 AD.
He named it mephitic (poisonous) air. In 1776 AD, Lavoisier studied its properties and named it
as Azota (a=no; zoo - life) as it does not support combustion and respiration. The name nitrogen
was derived from nitre (NaNO2) as it constitutes nitrogen.
Nitrogen gas is found both in free state and combined state. In free state it is found in atomospheric
air. It constitutes about 78.08% of air. In combined state it is found in the form of ammonia,
ammonium salts [NH4Cl, (NH4)2SO4], nitrate salts etc.
Laboratory Preparation of Nitrogen Gas
In the laboratory, nitrogen gas is prepared by heating a mixture of sodium nitrite and ammonium
chloride solutions and the gas prepared is collected by downward displacement of water.
NH4Cl + NaNO2 ∆ NH4NO2 + NaCl.
NH4NO2 N2 + 2H2O.
Delivery tube
Stand Gas jar
N2
N2 R.B. Flask
Burner
Trough
Beehive shelf
Fig: Laboratory Preparation of Nitrogen Gas
Test: It gives yellow ash of magnesium nitride, when a burning magnesium ribbon is introduced
in the gas jar.
Precautions: i) Apparatus must be air tight.
ii) Chemicals must be used in solution form.
iii) Gentle heat must be applied uniformly.
184 Blooming Science Book 9
Manufacture of Nitrogen Gas
It can be manufactured from two major sources (a) air and (b) nitrogenous compounds.
(a) From Atmospheric Air: Air is a mixture of different gases in which nitrogen is in higher
percentage by volume. Thus, nitrogen can be manufactured from atmospheric air. For
this, air free from carbon dioxide and moisture is cooled and liquefied under suitable
conditions (high pressure and low temperature). Then liquid air is subjected to fractional
distillation. Liquid nitrogen has lower boiling point (-196oC) than liquid oxygen (-185oC).
Therefore it distills out first leaving behind liquid oxygen. Nitrogen vapour is collected in
cylinders.
b) By burning phosphorus in air:
This gas can also be prepared by burning phosphourous in air in a closed vessel.
Phosphorous on burning in air combines with oxygen forming phosphorous pentoxide
and nitrogen remains behind in the vessel.
4P + 5O2 → 2P2O5 Crucible Gas jar
Nitrogen
Experiment: Take a crucible with a piece of phosphorous
and place it on a cork floating on water. Ignite the Phosphorus
phosphorous and immediately cover it with a gas jar. Cork
Phosphorus combines with oxygen forming phosphorous
pentoxide which gets dissolved in water and water level
rises in the gas jar. The gas left behind in the jar is nitrogen
gas.
(c) From Ammonia: Ammonia when reacts with Trough
chlorine gives nitrogen gas. Water
8NH3 + 3Cl2 → N2 + 6NH4Cl
Properties of Nitrogen Gas
Physical Properties
1. It is a colourless, odourless and tasteless gas.
2. It is slightly lighter than air. Its density is 14 while that of air is 14.4
3. It is slightly soluble in water.
4. It is neutral. It does not have any effect to indicators.
5. It is non-combustible gas and does not support combustion.
6. It can be liquefied to a colourless liquid which boils at -196oC and freezes at -269oC.
Chemical properties
1. Though it is non-supporter of combustion, some of the active metals like magnesium,
calcium, aluminium etc. continue to burn in atmosphere of nitrogen forming their
respective nitrides.
3Mg + N2 → Mg3N2 (magnesium nitride)
3Ca + N2 → Ca3 N2 (Calcium nitride)
2Al + N2 → 2Al N (Aluminium nitride)
Blooming Science Book 9 185
2. Nitrogen combines with hydrogen at 450 - 500oC under 200 atmospheric pressure in
presence of iron catalyst and molybdenum promoter. This method of preparation of
ammonia from nitrogen and hydrogen is called Haber’s process. It is a method for the
manufacture of ammonia.
N2 + 3H2 500oC 2NH3
200 atm
3. At very high temperature of 2000 - 3000OC, nitrogen combines with oxygen forming
nitric oxide. Nitric oxide is also formed during lightning.
N2 + O2 2000/3000ºC 2NO
Uses of Nitrogen Gas
1. It is used for manufacturing of ammonia, nitric acid, cyanamides, nitrides etc.
2. Being inactive, it is used in filling electric bulbs and high temperature thermometer.
3. It is used to replace fuels in fuel tanks of aeroplane to prevent possible explosion which
may occur due to the formation of explosive mixture of air and fuel.
4. It is used in manufacture of fertilizers such as calcium cyanamide, ammonium nitrate,
calcium nitrate etc.
5. It is used in filling air space of compounds to retain their flavour and colour better.
Let’s Learn
1. Nitrogen is filled in an electric bulb because it is an inactive gas and does not expand on
heating.
2. The major volume of atmosphere is covered by nitrogen because it is an important
components for the preparation of nutrients for plants.
3. Nitrogen is prepared from atmosphere because there is 78% nitrogen in air and it is
collected by fractional distillation of liquid air.
Main Points to Remember
1. Nitrogen gas is found both in free state and combined state.
2. Nitrogen gas can be prepared in the laboratory by heating a mixture of ammonium chloride
and sodium nitrite.
3. Nitrogen gas is manufactured in large scale by liquefaction of air and by the reaction
between ammonia and chlorine.
4. This gas is non-combustible and non-supporter of combustion.
5. Some metals like aluminium, magnesium and calcium burn in atmosphere of nitrogen
forming their respective nitrides.
6. Nitrogen combines with hydrogen forming ammonia by Haber’s process. Thus, ammonia
obtained is used to manufacture chemical fertilizers.
7. Nitrogen gas is used to fill electric bulbs and fuel tank of rocket.
186 Blooming Science Book 9
PRO J ECTWORK
Draw the diagram to show the laboratory. Preparation of above discussed gases in a chart
paper and keep it in your study room/classroom.
Exercise
1. Explain the laboratory preparation of nitrogen gas with a labeled diagram. Give also a
chemical reaction involved in it.
2. Write the physical properties of nitrogen.
3. Write three chemical properties of nitrogen with suitable chemical reactions.
4. Write the uses of nitrogen.
5. Explain sure test for nitrogen in brief.
6. Explain the manufacture of nitrogen gas.
a. by air b. by nitrogenous compounds like NH3
7. What happens when:
a) Ammonia is reacted with chlorine.
b) Calcium carbide (CaC2) is reacted with nitrogen.
c) A burning magnesium ribbon is inserted into a gas jar containing nitrogen.
8. What is the percentage of nitrogen in the air?
9. Give reason:
a) Nitrogen is filled in electric bulb and fuel tanks of aeroplanes.
b) Nitrogen in the air can play important role for living beings.
10. Complete the following equations and balance them:
(a) Mg + ……………………… ∆ Mg3N2
(b) Mg3N2 + ………............... Mg(OH)2 + NH3
(c) NH4Cl + ………….. ∆ N2 + H2O + NaCl
(d) N2 + O2 3000oC …………….
(e) Mg + N2 ……………
(f) N2 + H2 ……………
Activity for Practice
1. Take a few grains of potassium permanganate in test tube and supply heat to it. During the
course of heating, gas is produced. Study the property of the produced gas.
2KMnO4 ∆ K2MnO4 + MnO2 + O2
2. Take a balloon filled with H2 gas and transfer it to the gas jar. Test the properties of
hydrogen.
Blooming Science Book 9 187
Chapter METALS
12
Learning Outcomes Estimated Periods: 5
On the completion of this unit, the students will be able to:
• describe role of metals in living organisms ( zinc as an enzyme, importance of Na+ and
K+, negative effects of mercury and lead in human health)
• describe general properties of metals.
Introduction
We have already discussed that elements are mainly grouped into two categories i.e. Metal and
Non-metal. There is no sharp line of demarcation between metal and non-metals in the periodic
table. Some elements exhibit closely to metals whereas some exhibit the properties of non-metals.
We have made an attempt to differentiate between metals and non-metals in following sections.
Physical Properties of Metals
The following are the general properties of metals:
1. Most of the metals are found in the solid state except a few like mercury and caesium.
2. Metals are generally hard.
3. They can produce metallic clink sound.
4. They posses metallic luster.
5. They can be beaten into thin plate. This property of metal is called malleability. The metals
that show this property are called malleable.
6. The metals can be drawn into wires. This property of metals is called ductility and the
metals are ductile.
7. Metals are the good conductors of heat and electricity.
8. Generally metals have high boiling and melting points.
Chemical Properties of Metal
1) Reaction with oxygen
(Metal + oxygen heat metallic oxide or basic oxide)
eg. 2Na + O2 heat 2Na2O
But some metal oxides like Cr2O3, Mn2O7, SnO2 etc are not basic oxides
2) Reaction with hydrogen
Some active metals like sodium, potassium, calcium etc. react with hydrogen to form
unstable hydride
i.e. (Metal + hydrogen heat metallic hydride)
(2Na + H2 heat 2NaH)
188 Blooming Science Book 9
(2K + H2 heat 2KH)
(Ca + H2 heat CaH2) (hydrolith) Iron when kept in air for some ?DO
3) Reaction with acids
You
Some metals react with acids to produce days, rust are formed (brown Know
salt and gas.
powder) on its surface. Do
eg. Zn + 2HCl ZnCl2+ H2 same thing happens in copper?
In copper, if left for some days
4) Reaction with salt solution of less in air, it forms green layer of
electropositive metals carbonate and hydroxide ie, CuCo3
Cu(OH)2
eg. Fe + CuSO4 FeSO4+ Cu
5) Reaction with conc.acids
eg. Cu + 2H2SO4 CuSO4+ SO2 + 2H2O
Activity
1. To find the hardness and brittleness of substances
Take pieces of wood copper wire, iron nail, carbon rod ( pencil carbon) rubber and aluminium
vessel in your hand turn by turn and check their hardness. Also try to break them. Can you
find which of them are hard and which of them can be broken into pieces?
Fill in the table from your finding.
Materials Hard Soft Brittle Non-Brittle
Wood
Copper wire
Iron nail
Carbon rod
Aluminium Vessel
We found that metals are hard and less-brittle where as non-metals are usually soft and
more-brittle.
2. To find metallic lustre
Take some substances like carbon rod ( from pencil), tin sheet copper wire, paper, gold ring and
wood. Draw a line or mark a ‘scratch’ surface of each substances one by one. Non examine which
of shine and which of them donot shine, then filled in the table below:
Substances Metallic lustre Non-metallic lustre
Carbon Blooming Science Book 9 189
Gold
Paper
Copper
Tin
Wood
It is found that usually metals shine and and possess metallic lustre but not metals( except
grophite) donot possess metallic lustre.
3. To show malleable property (or Malleability)
Take all the substances as in activity above and hammer them one by one. You can see that only
metals can be hammered into thin sheets non-metals can’t be hammered into thin sheet. Now
filled table according to your finding.
Substances can be beaten into thin sheet can’t be beaten into thin sheet
4. To show ductility of substances
Use the materials as in activity three and find out which of them can be turned into wire. Then
fill the table.
Substances can be made into wire can’t be changed into wire
5. To show Conductivity of substances
To demonstrate metals conduct electricity. Take metal coin, plastic,
rubber, tin, paper clip and carbon rod. Now use these substances turn
by turn to complete the circuit as in the diagram shown. What will you
find ? Does the lamp glow with all the material given to you?
No, only metals conducts electricity and bulb glows with +-
non-metal the currents doesnot flow and the bulb won’t glow. Now fill Coin
the table.
Objects that conducts electricity that does conduct electricity
Metal
coin
Plastic
6. To show metals are good conductor of heat
Take all the materials as in activity 5. Now burn a candle and heat their one end one by one and
catch the other end with your hand. Can you tell which of them conduct heat? The metals conduct
heat and non-metal does not conduct heat.
Objects that conducts heat that does conduct heat
Metal
coin
Plastic
190 Blooming Science Book 9
7. Melting and boiling point of substances
The metals usually are solids so, they have high melting and boiling points. But some metals
like sodium, potassium, Lithium have low melting and boiling points. Some metals with their
melting and boiling point are given below:
Metals M.P B.P Metals like sodium, potassium are ?DO
Mg 650 oc 1107 oc very reactive metal, they even react
Al 660 oc 2467 oc You
Know
Fe 1535 oc 2750 oc with water and get dissolved in it.
Cu 1085 oc 2750 oc These metals when kept in water
Na 97.79 oc 882.94 oc burns violently due to release of
hydrogen gas.
K 63.5 oc 759 oc
Uses of Metals
1. Metals like gold, silver and copper are used for making jewelry and medals.
2. They are used for making coins.
3. They are used for making various articles like bolts, pipes, chains vehicles and railway
track.
4. They are used in making agricultural appliances, weapons and many other tools.
5. They are used for making utensils, frames, electrical goods, etc.
6. They are used for making many useful alloys such as brass and bronze, etc.
Differences between Metals and Non-metals
Metals Non-metals
1. State: They are solid at ordinary 1. They exist in all three states. With the exception
temperature and usually volatile at high of Bromine, Carbon and Silicon, they are
temperature. (exception: Mercury) either gases or volatile at low temperature.
2. Metallic luster: They possess a metallic 2. Generally they possess no metallic luster
luster and take a high polish. (exception: graphite)
3. Density: Their specific gravity is generally 3. These possess a low specific gravity.
high.
4. Malleability and ductility: Metals are 4. These are neither malleable nor ductile.
malleable, ductile and tenacious (exception: They possess little tenacity.
As, Sb, Bi).
5. Conductivity: Metals are good conductors 5. They are poor conductors of heat and
of heat and electricity (exception: Pb). electricity (exceptions: graphite and carbon).
6. Alloy formation: They possess the power 6. With the exception of carbon, silicon and
of forming homogeneous mixture with phosphorous the non-metals possess little
other metals called alloys. power of alloy formation.
7. Atomicity: Their vapours are generally 7. They form poly-atomic molecules
monoatomic.
Blooming Science Book 9 191
8. Nature of oxides: Generally, they form 8. They produce acidic oxides.
basic oxides. Chromium and manganese
form acidic oxides in addition to the basic
oxides and Al, Zn and Sn form amphoteric
oxides.
9. Hydrides: Metals generally do not form 9. They form stable hydrides with hydrogen.
hydrides. In case they form hydrides, they
form unstable hydrides.
10. Electro-chemical nature: In aqueous solution, 10. They form anions in aqueous solution. Non-
metals form cations. Some metals such as metals are liberated at the anode during
chromium and manganese form cations. electrolysis and therefore, electronegative
During electrolysis, metals are deposited at (exception: Hydrogen).
anode and therefore, they are electropositive.
11. Solubility: Metals can be dissolved 11. Many non-metals dissolve without
generally by a chemical reaction. chemical change taking place.
Metalloids
From the above discussions, we come to the conclusion that there is no sharp line of demarcation
between the metals and non-metals. Most of the elements satisfy the above tests only partially.
There are again some elements, like arsenic and antimony, which exhibit the properties of metals
as well as non-metals. Elements that exhibit the properties of both metals and non-metals are
called metalloids. Unlike metal, they are neither malleable nor ductile nor tenacious but arsenic
is good conductor of electricity and antimony forms alloys with metals.
Alloys
An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal. It
possesses in general the properties of metals as a class and in particular the general properties
of the elements of which it is composed, but these are not necessarily intermediated. Examples,
(i) Brass is an alloy of copper and zinc.
(ii) Bronze is an alloy of copper and tin.
(iii) German silver is an alloy of copper, zinc and nickel.
(iv) Steel is an alloy of iron and carbon.
(v) Stainless steel is an alloy of iron, chromium and carbon.
Alloys are generally made to impart one of the following special properties to the elements.
(i) To increase hardness.
(ii) To increase the strength.
(iii) To improve the colour.
(iv) To lower the melting point of metals.
Amalgam
The alloys of metals with mercury are called amalgams, eg. sodium amalgam, zinc amalgam, etc.
192 Blooming Science Book 9
Role of Metals in Animal Life
Our body consists of mainly C,H,N,O,P and S elements. Besides these major elements, there are
some elements which are found in less amount but perform the important functions in our body.
They are NA, K, Ca, Zn, Mg, Fe, Cu. These elements usually are available in the compound from
in our body. They remain in the form of enzymes, protens, lipids etc. in our body. Here, we will
discuss the roles of some important elements in our body.
1. Zinc
Zinc is available in very less amount in our body. An healthy adult body contains 2-3g of zinc.
Zinc is used by our body to make some 300 energies and mostly zinc is available in muscles and
bones of our body.
Role of zinc in human body:
i. Zinc plays important role in cell division and wound healing etc.
ii. It helps in maintaining immune system of our body.
iii. It helps in normal growth and development of child.
iv. It increase the fertility power of man and woman.
v. It helps in vision as our eye retina contain high concentration of zinc. In old age, the
zinc content in retina decreases which makes the vision blur.
vi. It also plays important role in slowing ageing process.
vii. It’s deficiency causesd dwarfism, anorexia, night blindness, diarrhoea, delay sexual
maturation etc.
2. Sodium and Potassium ions and their importance
Sodium / Potassium is an essential biological element to the more advanced animals; the ratio
sodium/potassium concentrations in intercellular and extracellular fluids is responsible for the
transport ions through the cellular membranes, the regulation of the osmotic pressure inside the
cell, the transmission of nervous pulses and other electrophysiological functions.
Role of sodium / potassium in human body:
i. Sodium ions are primarily found inside the human cells such as nerve cells.
They regulate the flow of water across the cell membrance.
ii. They are needed for transport of sugars and amino acids into cells..
iii. Potassium ion is important in heart function and in skelaton and muscle
contruction.
iv. Potassium ion maintains the electricity balance in the body.
v. Sodium/potassium pump helps with responding to stimule, transmmitting
nerve impulses and regulating cellular volume.
A drop of sodium levels in blood plasma below a normal range is known as hypoatremia.
Hypoatremia leads to headache, nausea, coma, seizures etc. Low potassium levels leads to
hypertension.
Blooming Science Book 9 193
Harmful effects of Mercury (Hg) and Lead (Pd)
Mercury: Mercury is a heavy metal and is found in liquid state in room temperature. It is a toxic
metal. The harmful effects or poisoning effects of mercury is called hydrargyria or mercurialism.
Some effects are as follow:
i) The mercury compounds or mercury droplets are toxic to humans as they have serious
impacts on nervous system, kidneys, liver, etc.
ii) It interfers the development of foetus during pregnancy.
iii) It damages the hearing, listening and speaking ability of a person.
iv) The inhalation of elemental mercury vapours can cause neurological and behavioral
disorders, such as tremors, emotional instability, insomnia, memory loss, neuromuscular
changes and headaches.
Lead: Lead is also a heavy metal which is widly used in varieties of products. It is excessively
used in paints, plastic toys, etc.
Important sources of environmental contamination include mining, smelting, manufacturing
and recycling activities, and, in some countries, the continued use of leaded paint and leaded
gasoline. More than three quarters of global lead consumption is for the manufacture of lead-acid
batteries for motor vehicles. Lead is, however, also used in many other products, for example
pigments, paints, solder, stained glass, crystal vessels, ammunition, ceramic glazes, jewelry, toys
and in some cosmetics and traditional medicines. Drinking water delivered through lead pipes or
pipes joined with lead solder may contain lead.
Some effects of lead are as follows:
i) It affects our nervous system badly..
ii) It causes headache, stomach pain, anaemia, retardation in growth and development in
children.
iii) Lead usually deposites in bone which distrubs the formation of blood cells and also make
the bone weak by destroying calcium in bone.
Let’s Learn
1 Silicon is metalloids because it shows the property of both metals and non-metals.
2. Metals are extracted from metallic ores. It is because ores are impure forms of metal when
metals are digged from the interior of earth.
3. A moist iron piece left in air becomes reddish brown. It is because the moist iron combines
with atmospheric air and forms rust in the presence of iron.
4Fe + 2H2O + 3O2 → 2Fe2O3 + 2H2O
Main Points to Remember
1. Metals are generally hard, malleable, ductile and possess metallic lusture. They
have melting point and boiling point.
2. Hydrolysis is the general term that explains the chemical reactions effected with the
addition of water with the elements.
3. The elements that exhibit the properties of both metals and non-metals are called metalloids.
4. Important metals are: cobalt, copper, lead, iron, zinc, gold, silver, aluminium, etc.
194 Blooming Science Book 9
5. An ore is a mineral from which the metal can be profitably and economically extracted.
For example, bauxite is the chief ore of aluminium.
6. The metals can be found as: oxides, hydroxides, halides, carbonates, sulphides, sulphates,
nitrates, silicates, phosphates, etc.
7. Metallurgy is the process of obtaining a metal from its ore.
8. An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal. The
alloys of metals with mercury are called amalgams.
9. Sodium, Potassium and Zinc are necessary minerals in our body.
10. Lead and Mercury are toxic elements. They cause harm in our body.
PRO J ECTWORK
Collect/Observe some materials used in your kitchen and differentiate them into metals and
non-metals. Make a table and justify with one or two important properties present in them.
Exercise
1. Choose the correct answer from the given alternatives:
a. Which is the general property of metal?
i. Hard ii. Malleable iii. Sonorous iv. All of above
b. Which is harmful metal for humans?
i. iron ii. lead iii. copper iv. none of above
c. Which mineral is found in human blood?
i. Ca ii. Fe iii. P iv. Au
d. Which one is not an element?
i. Brass ii. lithium iii. silver iv. all of above
e. Which is necessary for human body?
i. lead ii. mercury iii. potassium iv. none of above
2. Answer the following questions:
a) What are the differences between metals and non-metals?
b) Explain the term hydrolysis? Is hydrolysis a chemical reaction?
c) Define metalloids with example.
d) What are the properties of metalloids?
e) Explain the usefulness of metals for the development of mankind.
f) What is metallurgy? In how many ways can you bring concentration on ore minerals?
g) Write any four alloys with their compositions..
Blooming Science Book 9 195
h) What are the roles of zinc in our body? Explain.
i) Explain the roles of sodium / potassium in our body.
j) Write the harmful effects of lead and mercury to human being.
k) Name the metals which are found in liquid state at the normal temperature.
l) Name the non-metal which is found in liquid state at the normal temperature.
m) Write any two metals which are not malleable.
n) Name any two oxides of metals which are not base.
3. Why are the melting and boiling points of copper higher than that of sodium?
4. What happens when, (write with a balanced chemical equation).
a) Copper is placed in dilute and concentrated Nitric acid?
b) Iron reacts with copper sulphate solution?
c) Iron reacts with dilute Hydrochloric acid?
d) Copper reacts with concentrated Sulphuric acid?
e) Aluminium reacts with concentrated Hydrochloric acid?
5. Give reasons:
a) Meltals are used to make cooking utensils.
b) Copper and alumimium are widely used for making conducting wires.
c) Alloys are made.
196 Blooming Science Book 9
Chapter
13 CARBON AND ITS COMPOUNDS
Learning Outcomes Estimated Periods: 5
On the completion of this unit, the students will be able to:
• identify and demonstrate the existence of carbon in some common substances like
wood, sugar, oil, etc.
• explain the physical and chemical properties of carbon.
• differentiate organic and inorganic compounds.
Introduction
We know far more compounds that contains carbon than those compounds that do not contain
carbon. Carbon is the basic element in the structure of the biosphere. The biosphere is the layer
of living matter that inhabits the earth’s sphere. The organic world, the world of living and
growing things is largely formed of compounds that have carbon as their principal element.
The importance of carbon compounds lies in the fact that these compounds often have very
large molecules that are chained like or exist in the network pattern. The properties of carbon
account for particular qualities of some rocks, toughness of wood, strength of silk and elasticity
of rubber. Carbon compounds form the basis of science of foods, dyes, drugs, perfumes, plastics
and innumerable new compounds.
Position of Carbon in the Periodic Table
Carbon element has four valence electrons in its atom. So, it is placed in group IV A of the
periodic table along with silicon (Si), germanium (Ge), Tin (Sn) and Lead (Pb). Carbon is the
first element of group IV A and is the most electronegative element in the group.
Since carbon has two shells in its atom. It is placed in second period of the periodic table. Its
sub-shell configuration is 1s2, 2s2, 2p2. It is a p - block element.
Occurrence of Carbon
Carbon is found in nature in the free as well as in the combined state. Carbon 6p+
ranks twelfth in abundance among the elements in the earth’s crust. In a free
state, it is found as coal, diamond and graphite. In a combined state, carbon 6n
occurs in the form of: (i) carbonates (limestone, marble and dolomite), (ii)
petroleum and natural gas, (iii) organic compounds such as proteins and fats,
(iv) carbon dioxide (in the air), etc. Carbon occurs in the form of carbonates
in rocks, petroleum, natural gas and in all living things.
Carbon is a non-metal that has been known since ancient time in the form of coal and charcoal.
The word carbon is derived from the Latin word ‘Carbo’ which means charcoal or soot. In
fact, carbon is found in nature in various forms like coal, charcoal, graphite, diamond. Carbon
is also found in combined state in the form of compounds like carbondioxide or air, protein,
carbohydrates and fat of living beings also.
Blooming Science Book 9 197
The symbol of carbon is C. Its atomic number is 6 and atomic mass is 12. Its electronic
configuration is 1s2, 2s2, 2p2. It has four valence electrons. Its valency is 4.
Carbon is one of the most widely distributed elements in nature. Carbon occurs in nature, both as
free element and in the combined state. In the free state, it occurs in the form of diamond, graphite
and coal. In the combined state, it occurs in the form of various inorganic and organic compounds.
Sources of Carbon
Some of the sources of carbon are given below:
i. Carbon occurs in the atmosphere as carbon-dioxide.
ii. Carbon in the form of compounds is found in all living beings, both plants and animals
as carbohydrates, fats, proteins and vitamins. These are the nutrients of living beings. So
carbon has very close relationship with our life.
iii. Carbon occurs in the carbonate and bicarbonate compounds.
iv. Carbon occurs in natural hydrocarbons like coal, petroleum and natural gas.
v. Carbon occurs in natural water as dissolved carbon-dioxide.
Carbon in Common Substances
Activity
Collect the following carbon containing substances obtained from plants and animals.
Substances Sources Wood Delivery
Wood Plants (Trees) pieces tube
Sugar Sugarcane
Kerosene Fossils Burner Clear lime
Petrol Fossils water
Edible oil Plants (seeds of mustard, Stand
soyabean, sunflower etc.)
Fats Animals
Take a sample of each of the substance mentioned above. Burn each substance as shown in
the figure above and observe. Each substance burns and changes into black substances like
smoke, coal, charcoal, etc. This shows that each of these substances contain carbon. Carbon
on burning releases carbon dioxide gas which when passed into lime water turns it milky
(lime water test for carbon dioxide)
198 Blooming Science Book 9
Activity
To Show the sugar or starch contains carbon
Method Sugar Spoon
Burner
1. Take a small amount of sugar in a metal spoon that has a
long handle. Place the spoon over the spirit lamp and heat
it gently in a flame.
2. Pass the gas produced on burning the sugar charcoal into
lime water.
Observation
The sugar melts (m.p. 186oC) slowly and converts into thick, yellow liquid. Further heating
to about 200oC, converts the yellow liquid into brown coloured solid called Caramel. It on
further heating changes into a black substance called sugar charcoal. The sugar charcoal on
burning produces gas which turns lime water milky.
Conclusion
Sugar burns in air and forms carbon dioxide. It is the carbon dioxide that turns lime water milky.
Nature of carbon
The outermost orbit of a carbon atom consists of 4 electrons. It needs four other electrons to
attain stable state. One carbon atom can bond with four other carbon atoms or other elements to
form a compound. This property of carbon is responsible for the formation of a large number of
carbon compounds. Carbon has main two properties which are explained below:
a) Catenation:
The property of carbon due to which it can form long chain of carbon by linking with other
carbon atoms with covalent bond is called catenation. eg C C C C (linear chain)
covalent bond
CCCC
C
(branched chain)
C
Due to this property of carbon, long chain oragnic compound like polyethene, PVC, etc. are
made.
b) Tetra valency:
The valency of carbon is four. It can form covalent compounds with other elements by sharing
four electrons. The carbon can form covalent compounds with oxygen, hydrogen, nitrogen, chlo-
rine, etc. The catenation and tetravalency properties of carbon make possible for the existence of
large varieties of organic compounds.
Blooming Science Book 9 199
Allotropes of Carbons ?DO
The existence of an element in more than one form in Graphene is an allotrope You
the same physical state is called allotropy. The various of a carbon consisting of a Know
such forms of the element are called allotropes of that sigle layer of carbon atoms
element. Allotropes always exhibit different physical arranged in an hexagonal
properties but possess same chemical properties. latti. It is the strongest
The allotropes of carbon are of two types - the material ever tested. It is 200
crystalline and amorphous carbon. Diamond and times stronger than steel but
graphite are crystalline carbon. Charcoal, coke, lamp lighter than paper.
black soot, etc are amorphous or non-crystalline carbon.
Allotropes of an element consist of only one kind of atoms with their different arrangement
having different physical properties but same chemical properties.
Allotropes of Carbon
Crystalline Amorphous
Diamond Graphite Graphene Fullerence Coke Coal Charcoal Lamp black
(Bucky ball)
soot
Crystalline Carbon
Diamond and graphite are crystalline allotropic forms of carbon. Both of them consist of each
atoms only. When diamond and graphite are strongly heated in air, they burn completely to give
carbon dioxide. This shows that they are made up of carbon atoms and are chemically identical.
But they differ in the arrangement of carbon atoms in their crystals. This leads to the difference
in their appearance and many other physical properties.
Diamond
Diamond is the purest form of carbon. A diamond is a transparent substance.
It is one of the naturally occurring hardest substances. It has an extra-ordinary
brilliance. It is the purest form of carbon.
Structure of Diamond
A diamond crystal possess a large number of carbon atoms. Each carbon
atoms in the diamond crystal is linked to four other carbon atoms by
strong covalent bonds. Hence, it is difficult to break diamond.
Uses of Diamond
When people think of diamonds, they tend to think of jewelry, such as diamond rings, earrings,
or necklaces. However, as the hardest naturally occurring substance on the planet, diamonds are
used for many other, commonly unknown, purposes as well. Some unusual uses of diamonds.
200 Blooming Science Book 9
The words you are searching are inside this book. To get more targeted content, please make full-text search by clicking here.
Blooming Science-9-2077 final final for press
Discover the best professional documents and content resources in AnyFlip Document Base.
Search