Chap 1: EXPERIMENTAL CHEMISTRY
I). Measurements:
The study of chemistry requires the measurement of various quantities, such as mass and
volume.
Mass:
The mass of a substance is the amount of matter it contains. Its unit is the gram (g) for
small amounts or the kilogram (kg) for larger amounts. It is measured using a beam
balance or an electronic balance.
Fig 1.1. Beam Balance
Fig 1.2. Electronic balance.
Time:
The S.I unit for time is seconds (s). Other units are minutes (min) and hours (h). A
stopwatch or a stopclock is used to measure time.
Fig 1.3 A digital stopwatch
Temperature:
The S.I unit for temp is Kelvin (K). Another unit is the degree Celsius (°C).
Temperature can be measured using a thermometer. The mercury-in-glass and the
alcohol-in-glass thermometers are commonly used in labs, with ranges from -10 to
110°C.
Volume:
The volume of a substance is the amount of space it occupies. It is measured in cubic
centimeters (cm³), cubic decimeters (dm³) or liters (l). There are many lab apparatuses
used for measuring the volume of a fluid.
Some of these are as follows:
Fig 1.4. A beaker is used to measure approximate volumes.
Fig 1.5. A measuring cylinder is accurate to the nearest cm³.
Fig 1.6. A burette is accurate to the nearest 0.1 cm³. The volume of liquid required is run
off from the bottom through a tap.
Fig 1.7. A pipette measures fixed volumes (25cm³, 50cm³, 10cm³, etc.) very accurately.
II). Collection of Gases:
Speed of Chemical Reactions: In a chemical reaction where a gas is given off, we can
measure the speed of the reaction by recording the volume of gas given off in a certain
time interval, which can be measured using a gas syringe, or by recording the progressive
loss in mass as the gas escapes into the air, which can be measured using a balance.
To collect a gas, the following properties may be taken into consideration:
1. Its density with respect to air.
2.
its solubility in water
If the gas [to be collected] is denser than air, it can be collected using the downward
delivery method, in which the air is displaced upward.
If the gas is lighter than air, it can be collected using the upward deliver method, in which
the air is displaced downward, and the gas is collected in an upside-down container.
For gases that are insoluble or slightly soluble in water, the method of displacement of
water may be used, in which the gas is passed through water into an upside down
container; the gas collects on top and the water gets displaced downwards.
Chap 2: The Particulate Nature of Matter
I). States of Matter:
A matter is any substance that has mass and occupies space. There are three states of
matter: solid, liquid and gas. Most of the matters can exist in all three states. Matter can
change from one state to another due to changes in temperature or pressure. A common
example is water: on room temperature it is liquid, on freezing it becomes solid (ice), and
on boiling it becomes gas (water vapour).
Property
Shape
Volume
Compressibility
Solid
Liquid
Fixed
Not fixed
Fixed
Fixed
Cannot be
Cannot be
compressed
compressed
Table 2.1. Properties of solids, liquids and gases.
Gas
Not fixed
Not fixed
Can be compressed
II). Kinetic Particle Theory:
The kinetic particle theory states that all matter is made up of tiny particles, and that these
particles are in constant, random motion. However, the movement of these particles
varies according to the state of matter.
The Solid State:
The forces of attraction between the particles of solid are very strong, so they are held
very close together and cannot move about freely. They have very less kinetic energy,
enough to enable them to vibrate about their fixed positions. This is why a solid has a
fixed shape.
As these particles are already very closely packed, a solid cannot be further compressed,
which is why it has a fixed volume.
The Liquid State:
The forces of attraction between the particles of a liquid are weaker than those in a solid.
These particles are not held in fixed positions, but instead arranged in a disorderly
manner, and, due to higher kinetic energy as compared to solid, can move freely by
sliding over each other. This is why a liquid does not have a fixed shape.
The particles of a liquid are further away from each other as compared to those in a solid,
but are still packed quite closely together. Thus a liquid cannot be compressed and has a
fixed volume.
The Gaseous State:
The particles of a gas have high kinetic energy and very low forces of attraction,
therefore they are not held in fixed positions. They can move about rapidly in any
direction, therefore a gas has no fixed shape.
A gas takes up as much free space as provided to it, which is why there are large
distances between particles. These spaces make it possible for the gas to be compressed,
or forced into coming closer. Therefore, a gas has no fixed volume.
III). Changes in State:
Solid to Liquid: Melting
Liquid to Solid: Freezing
Liquid to Gas: Evaporation or Boiling
Gas to Liquid: Condensation
Solid to Gas: Sublimation
Gas to Solid: Solidification
Melting and boiling points are important in identifying a substance and testing its purity.
Pure substances have definite melting and boiling points. The presence of impurities will
cause boiling points to rise and melting points to fall.
● The constant temperature at which a pure solid changes into a liquid is called its
melting point.
● The constant temperature at which a pure liquid changes into a gas is called its
boiling point.
Heating curve of a pure substance:
Cooling curve of a pure substance:
Note: the parts in both graphs that do not show any temperature change are the parts in
which the state changes occur. In them, both states occur together.
IV). Diffusion:
Diffusion is the movement of particles from a region of higher concentration to a region
of lower concentration. As it involves a movement of particles, it is only possible in
liquids and gases. Gases diffuse faster than liquids as they are generally lighter.
Gases take up whatever space is available to them. This is because of diffusion. This
phenomenon is shown in the figure below.
Chap 3: Methods of Purification
A pure substance is one that contains only one type of substance, i.e. only one type of
atom or molecule.
A mixture contains two or more different substances physically combined together. They
can be separated by different methods of purification.
Dissolving, Filtering and Evaporating:
This method is used when a mixture of solids (both soluble and insoluble) are dissolved
in a particular liquid.
For example, take a mixture of salt, sand and water. On stirring, the salt dissolves in
water, whereas the sand remains as it is. If we filter this mixture, the sand will be
collected in the filter paper as the residue, while the dissolved salt and water will pass
through the filter paper as filtrate.
To separate the salt from water, we put the solution in an evaporating dish and heat
gently. The water evaporates, leaving behind salt crystals. The slower the evaporation,
the larger the crystals. This process of evaporating some of the solvent and letting the
remainder dissolved solute to cool and turn into crystals is known as crystallization.
Separating Funnel:
This method is used to separate two immiscible liquids (liquids that do not mix together).
For example, we could take the mixture of oil and water and place them in a separating
funnel. The lighter liquid (oil) collects on the top and the heavier liquid (water) collects at
the bottom. When the tap is opened, the water at the bottom starts to flow out. It can be
collected in a container placed below the tap.
Fig 3.1. A separating funnel.
Simple Distillation:
This method is used to separate a pure liquid from a solution.
For example, take salt water, with the objective of obtaining the pure water from it and
removing the impurities. The flask containing the mixture is heated and when the solution
boils, steam is given off. This is cooled down and condensed in a condenser, which
consists of a jacket of cold water with the coldest water entering at the bottom of the
jacket and circulating out through the top. This way, it makes sure that the coldest part of
the condenser is just before the vapour escapes. The condensed water is called the
distillate and is collected in a receiver.
The water which collects in the flask is very pure as all its impurities are left behind in
the flask. It is called distilled water. A thermometer at the top of the flask shows the
temperature at which the vapour distills over (100°C).
Fig 3.2. Simple Distillation Apparatus.
Fractional Distillation:
This technique is used to separate two mixtures which dissolve in each other. These
liquids are known as miscible liquids, and their separation depends on their different
boiling points.
For example, let’s take a mixture of alcohol and water, having boiling points of 78°C and
100°C respectively. A fractionating column is used for this distillation. The fractionating
column is packed with glass beads, which provide a large surface area for condensation.
Method:
When the flask is heated, the vapour coming off from the mixture will contain both
ethanol and water molecules, but will be richer in ethanol molecules due to their lower
boiling point. At first the vapours will condense back, but as the coloumn warms up the
ethanol molecules will reach the top of the coloumn and will be distilled over, while the
water molecules, due to their higher boiling point, will continue to condense and fall back
in the flask. While the ethanol is being distilled over, the thermometer at the top shows a
constant temperature of 78°C, the boiling point of ethanol. The ethanol passes through
the condenser, in which the cool running water cools and condenses the hot ethanol
vapour, and the liquid ethanol then flows down the inner tube of the condenser into the
receiver. This ethanol is now the distillate.
Fig 3.3. Fractional Distillation Apparatus.
Paper Chromatography:
The technique of using a solvent to separate a mixture into its components is called
chromatography. A drop of the mixture is placed on a chromatography paper, and one
end of the paper is dipped in a solvent. The solvent travels up, separating the different
components of the mixture and spreading them on the chromatography paper. This
technique is useful in separating dyes in ink, pigments in plants, etc, and to detect traces
of banned substances in food.
Testing the Purity of substances:
● A pure substance has a fixed and exact melting point.
● A pure substance has a fixed and exact boiling point.
● A pure substance shows only one spot on a chromatogram.
Chap 4: Elements, Compounds and Mixtures
Element:
An element is a substance that cannot be split into two or more simpler substances by chemical
means or by electricity.
Elements can be classified into metals and non-metals.
Metals
Shiny appearance (lustrous)
Mostly solids at r.t.p*.
Malleable
Sonorous
Ductile
High melting points and boiling points.
●
●
●
Non-metals
Dull appearance (non-lustrous)
Either gases, volatile liquids or solids with low
melting points at r.t.p.
Brittle (if solid)
Generally low melting points and boiling
points.
Generally poor conductors of heat.
Generally poor conductors of electricity.
Good conductors of heat.
Good conductors of electricity in all states of
matter.
Table 4.1. Differences between Metals and Non-metals.
*r.t.p: room temperature and pressure.
Atom:
An atom is the smallest possible particle of an element that can take part in a chemical change.
Compound:
A compound is a pure substance which contains only one type of molecule made up of atoms of
more than one element. It is formed when two or more different elements are chemically
combined together.
Heat can be used to form compounds. Heat can also be used to break down compounds into
elements or simpler compounds. Such a chemical reaction is called thermal decomposition.
Molecule:
A molecule is the smallest particle of a compound and is made up of a group of two or more
atoms chemically combined together. The atoms that join to make one molecule can either be
same or different.
To show the atoms present in the molecule of a compound, we use its chemical formula, e.g., the
chemical formula of water is H2O. It contains the elements Hydrogen (symbol H) and Oxygen
(symbol O). The subscript number 2 shows the number of atoms of the element preceding it.
Mixture:
A mixture is not a pure substance, as its components are not chemically combined. It is formed
when two or more substances are physically combined.
The differences between mixtures and compounds are listed below:
Mixture
1. Component substances can be separated by
physical means.
2. Its physical properties (colour, density, etc)
are an average of those of the substances in it.
3. Normally little or no energy is given out or
taken in when a mixture is formed.
Compound
Component substances cannot be separated by
physical means.
Its physical properties are individual and not the
result of its elements.
Energy is usually taken in or given out when a
compound is formed. This is because a
chemical reaction takes place here.
4. A mixture’s composition can vary.
A compound’s composition cannot vary.
Constituent elements are present in a fixed
proportion by mass.
5. Its chemical properties are the result of the Its chemical properties are quite different from
substances in the mixture.
those of its elements.
Table 4.2. Differences between mixtures and compounds.
Chap 5: Atomic Structure
Dear students, this lecture is about the atom and its structure. Understanding this structure in chemistry is
essential to further understand things such as chemical properties, bonding, etc.
The atom consists of a ‘nucleus’, and there are orbits around it called ‘shells’. You can understand the
shape better if you think of it as a miniature solar system: the sun is the nucleus and the orbits on which
the planets rotate are the shells. The nucleus of an atom consists of sub atomic particles called ‘neutrons’
and ‘protons’. The protons are positively charged and the neutrons have no charge, which means they are
neutral. So in general, the nucleus is positively charged. The shells have subatomic particles called
‘electrons’ on them, which rotate around the nucleus, staying on their orbits (you can think of them as the
planets). These electrons are negatively charged.
The charge on an overall atom is zero, which means that there are an equal number of protons and
electrons present in an atom. The number of neutrons does not affect the charge on the atom, so there is
no fixed trend as to how many neutrons each atom contains.
The following table shows some of the properties of the three subatomic particles:
Particle
Symbol
Relative Mass
Relative Charge
Proton
p
1
+1
Neutron
n
1
0
Electron
e
1/1840
-1
(negligible)
Atomic Number:
The atomic number is the number by which each type of element is known. In all the atoms of a single
element, the atomic number will be the same. Also, this atomic number will not be shared by an atom of
any other type of element. This means that every element has a unique atomic number.
This atomic number is basically the number of protons in each element. For example, each atom of the
element Hydrogen contains 1 proton, so Hydrogen has the atomic number ‘1’. This atomic number is
shared by no other element, and neither does any element have only 1 proton. So basically, the number of
protons of an atom decides what element it belongs to. All these numbers are given in the periodic table
of elements. The atomic number is also called the proton number of an element. Its symbol is Z.
Nucleon Number:
The nucleon number of an element is the total number of protons and neutrons in its atom, i.e. the number
of subatomic particles in its nucleus. As we saw earlier that both the particles neutron and proton have the
relative mass of ‘1’, while the electron has negligible mass, therefore the sole mass of the atom depends
on how many subatomic particles are present in the nucleus. We ignore the electrons when taking the
mass of an atom. Therefore, the nucleon number is also the mass number of the element. Its symbol is A.
This number for every element is also given in the periodic table.
If we subtract the nucleon number (A) from the proton number (Z), we can get the number of neutrons in
each atom.
Isotopes:
Isotopes are atoms of the same element with different number of neutrons.
As I mentioned earlier, the ‘type’ of element depends on the number of protons it has. Therefore if the
number of neutrons varies in the atoms of a single element, it does not mean that the atoms with the
different nucleon numbers belong to different elements.
For example, let’s take Hydrogen. It has 1 proton and no neutron (nucleon number = 1), but it has 2
naturally occurring isotopes. These are ‘deuterium’, with 1 proton and 1 neutron, and tritium, with 1
proton and 2 neutrons. As you can see, the proton number does not change, so we can say that deuterium
and tritium are actually the element hydrogen, but as their nucleon number is different, we call them
‘isotopes of hydrogen’. They are called this because their chemical properties (which depend on proton
number) are the same as hydrogen, but their physical properties are slightly different.
Arrangement of Electrons in an Atom (The Electronic Configuration):
As we have already seen, the electrons revolve around the nucleus on orbits called shells. But how are
these electrons arranged on these shells? They are divided between these shells in fixed quantities, that is,
the first shell (which is always filled first) can hold a maximum of 2 electrons. The second, third and
subsequent shells can generally hold up to 8 electrons. At this level, u need not worry about shells further
than the third, but learn that all shells after the first tend to complete their ‘octet’, i.e. they each hold up to
8 electrons. E.g., let’s take Calcium (proton number 20). It has 20 electrons, so its ‘configuration’
(division of electrons) will be:
First shell: 2
Second shell: 8
Third shell: 8
Fourth shell: 2
This configuration can be written as: 2,8,8,2.
Fig 5.1. The electronic configuration of Calcium (2,8,8,2).
In the figure above, only the shells and the electrons are shown. The nucleus (not shown) lies in the
middle.
You must learn the electronic configuration of the first 20 elements.
These configurations are very important, as the chemical properties of an element depend upon the
number of valence electrons in its atoms. The valence shell is the outermost shell of an atom, and the
electrons in this shell are known as the valence electrons.
Another importance of these configurations is that they are responsible for the arrangement of the
elements in the periodic table. You will learn more about this in the chapter ‘The Periodic Table’, but just
for a general idea you can learn that the elements in the periodic table are arranged according to their
increasing proton number, i.e. 1, 2, 3, and so on. Also, the number of shells decides which period does an
element belong to, while the number of valence electrons decide which group does it fall in.
Chap 6: The Periodic Table of Elements
The Periodic Table is an arrangement of elements in order of their increasing proton (atomic) number.
Groups:
The vertical columns of elements in the Periodic Table are called groups. There are eight groups in the
periodic table (there are more than eight vertical columns, but the block of elements between group II and
group III are not classified according to groups, so that we do not count the columns in that bock when
counting groups). The group number of an element is determined by its number of valence electrons. For
example, Sodium (Na) has 1 valence electron, so it falls in group I. The group number is generally written
in roman numbers. The elements in the same group have similar chemical properties, as the chemical
properties majorly depend upon the number of valence electrons. Their physical properties, however, are
very different. We will learn about all this as the chapter proceeds.
Periods:
The horizontal rows of elements in the Periodic Table are called periods. There are seven periods in the
Periodic Table. The period number indicates the number of shells each atom of an element contains. All
the elements in one period have the same number of shells. For example, Potassium (K) and Calcium (Ca)
both have 4 shells, so they both are placed in period 4.
From left to right across a period, there is a decrease in metallic properties and an increase in non-metallic
properties.
Fig 6.1. The Periodic Table.
As you can see in the periodic table, different parts of the table are coloured differently. Let’s learn these
classifications in detail.
Transition Metals:
The block of metals between group II and group III is known as transition metals. These are typical
metals. They are strong and hard, good conductors of heat and electricity, and have high melting points.
The section of the periodic table in which the transition metals are placed does not have any group
numbers. There are some properties typical of all transition metals, and the other metals (the ones placed
in groups) do not generally have these. E.g., the transition metals form coloured compounds. They are
used as catalysts to speed up chemical reactions. They have variable oxidation states. You will learn more
about oxidation states in the chapters of bonding, but I will give you a general idea.
You see, atoms of most elements can either loose or gain electrons, giving them a positive or negative
charge. This charged atom is known as an ion. The oxidation state is the charge an atom of an element
would have if it existed as an ion in a compound. For example, oxygen in its ion form, i.e. O2-, has the
oxidation state -2. This state is fixed for many elements, but not for transition metals.
An example of transition metals is Iron (Fe). It forms coloured compounds. It can exist in form of two
ions, Iron (II) and Iron (III). The number in the brackets refers to its oxidation state. Iron (II) compounds
are green, and Iron (III) compounds are reddish-brown. Iron is used as a catalyst in the Haber Process
(you will learn more about this process later). Other typical properties of metals are also true for this
metal.
Groups I and II:
These two groups consist of metals. These are reactive metals, and their reactivity increases as you go
down the group. Group I metals are also known as Alkali Metals, while Group II metals are called
Alkaline Earth Metals. They are called this because they react with water to form an alkali and hydrogen
gas. Alkali Metals are very reactive, even reacting with air and cold water. They are soft, and can be cut
easily. When freshly cut, each of these elements shows a shiny and silvery surface that rapidly tarnishes
in air. They also have low melting and boiling points, and low densities. Why then are they called metals,
you will learn so in a later chapter. Examples of Group I metals are Sodium (Na) and Potassium (K).
Examples of Group II metals are Magnesium (Mg) and Calcium (Ca).
Note: Hydrogen does NOT belong to Group I. It has 1 valence electron so it is generally placed in group
I, but it does not have the properties typical of Group I elements.
Group VII:
These elements are called Halogens. The Halogens have low melting points (M.P) and boiling points
(B.P). They are also coloured. The M.Ps and B.Ps of the Halogens increase as we go down the group, and
their colour gets darker. Halogens react with most metals to form halides. Their ions are called halide
ions. In a solution, a more reactive halogen replaces a less reactive halogen. You will learn more about
this later, and can relate all later learnings back to this.
Examples of Halogens are Fluorine (F) and Chlorine (Cl). They, however, always exist as F2 and Cl2. This
is a property typical of all halogens, i.e. they are diatomic (refer to the chapter of bonding).
Group VIII (also called Group 0):
These elements are called Noble Gases. They are also referred to as inert gases because they are
unreactive. The noble gases:
● Are monatomic elements
● Are all colourless gases at room temperature
Have low M.Ps and B.Ps that increase on going down the group
Are insoluble in water.
Their atoms have full outer shells of electrons, therefore they do not need to lose, gain or share electrons,
hence they are unreactive. They do not form ions or molecules.
Namely, the Noble Gases are Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon
(Rn).
●
●
Chap 7: Ionic Bonding
Dear students, from here the tricky part of chemistry starts. So as not to confuse you, I have divided the
chapter of ‘bonding’ into two parts: ionic bonding, and covalent and metallic bonding. This lecture deals
only with ionic bonding.
Before we begin, you all must understand the need for bonding. You see, atoms of many elements are
unstable, i.e. their valence (outermost) shell is incomplete. We have already learnt that atoms try to
complete their ‘duplet’ or ‘octet’, and for this many of them either gain or lose electrons. Note that I say
‘many’, not ‘all’. This transfer of electrons causes atoms to become charged, either positively or
negatively; these charged atoms are called ‘ions’. Anyway, we all know opposite charges attract each
other, and that’s the whole phenomenon behind ionic bonding. There is a transfer of electrons between
atoms, and as they become oppositely charged, they attract each other and a strong bond is formed
between them. This is known as ionic bonding.
Ionic bonding occurs generally between metals of group I and II, and non-metals of groups VI and VII.
This is because of their configurations. Let us learn why.
Formation:
As I mentioned earlier, atoms try to complete their duplet or octet, in order to achieve stability. But who
gains electrons and who loses them?
Let us take the element sodium, which has the configuration 2,8,1. You can see that it has 1 electron in its
outermost shell. If it loses that electron, it will achieve stability. However, it could also gain 7 electrons
and achieve that goal. Compare losing the one electron to gaining SEVEN electrons; what would be
easier? Obviously losing the one it already has. So a metal (Group I and II) will always go for losing its
valence electrons, gaining a positive charge. A positively charged ion is known as a cation. Remember,
metals always form cations.
Now let us consider Chlorine. It has the configuration 2,8,7. It needs to gain either 1 electron to complete
its octet, or needs to lose 7 of its own electrons. Again, the easier solution would be to gain the one
instead of losing the seven, so it will gain, unlike metals. So, non-metals generally gain electrons to
achieve stability, forming negatively charged ions (anions).
You see in the above examples that Sodium loses 1 electron, while Chlorine gains 1. Where does the lost
electron go, and where does the gained electron come from? You see, the sodium atom gives an electron
to anyone willing to take it, and if we are reacting sodium and chlorine together, the chlorine accepts the
atom that sodium is giving. There is an exchange of electrons, and opposite charges are formed (+1 on
sodium and -1 on chlorine), causing ‘electrostatic attraction’ between them that holds them together, and
thus an ionic bond is formed.
This can be shown by a dot and cross diagram.
Fig 7.1. Exchange of electrons between Na and Cl
In diagram 7.1, you can see that electrons of Na are represented by crosses, and electrons of Cl by dots.
The electron that Cl gains from Na is also represented by a cross, to show that a transfer took place. The
bond is then formed due to the attraction between opposite charges.
In O levels, ionic bonds are usually shown in this way.
Fig 7.2. Ionic Bond between Na and Cl.
This new compound is totally different from its products. It is called Sodium Chloride (NaCl), also
known as common salt in daily language. Compounds formed by ionic bonding are called ionic
compounds, therefore NaCl is an ionic compound.
Note: NaCl was just an example to explain the main concept of ionic bonds. There are many more
examples other than this. Also, Na and Cl do not always bond together.
Other Examples:
I mentioned earlier that Group I, II, VI and VII elements take part in ionic bonding. Aluminium (Group
III) forms some ionic compounds as well. Let us look at some more examples of Ionic Compounds.
Fig 7.3. Ionic Bonding in MgCl2.
In the above example, you can see that Magnesium (Group II) can lose two electrons, so two chlorine
atoms are needed to complete this bond. A +2 charge is formed on Mg atom, while a -1 charge comes on
both the Cl atoms, as each gains one electron.
Fig 7.4. Ionic Bonding in MgO.
In the above example, you can see that Magnesium must lose two electrons, while Oxygen must gain two.
So Magnesium gives both its electrons (gaining a +2 charge) to Oxygen, which accepts (gaining a -2
charge). So, Magnesium Oxide (MgO) is formed.
Structure:
Ionic Compounds form giant ionic structures, called giant lattice structure or crystal lattice. Let us, again,
take the example of NaCl, which forms a cubic lattice.
Fig 7.5. Cubic Lattice of NaCl.
In a cubic lattice, each Na+ is bonded to six Cl-, and each Cl- is in turn bonded to six Na+. No two Na+ ions
are placed together, and no two Cl- ions are placed together. This is why the structure is so strong; there is
attraction from every side which makes it difficult for ions to break apart.
Chemical Formulae:
The formula of an ionic compound is constructed by balancing the charges on the positive ions with those
on the negative ions. The positive charges must equal the negative charges. For example, in Sodium
Chloride, Sodium (whose symbol is Na) has charge +1, while Chlorine (whose symbol is Cl) has charge 1. So 1 Na needs 1 Cl, therefore the chemical formula “NaCl”.
Now, let’s consider Magnesium Oxide. Mg has +2 charge, while O has -2 charge. There is 1 to 1 ratio
between the ions, so the chemical formula is MgO.
But what happens when there isn’t a 1 to 1 ratio? Let us take Aluminium Oxide. Al has the charge +3,
while O has the charge -2. The following way is adopted then.
Fig 7.6. Determining Chemical Formula by cross multiplying charges.
The subscript after each symbol represents the number of atoms of that element in each bond of a
compound. Here, we can see that in Al2O3, two Al3+ ions join with three O2- ions to form a complete bond
The above strategy may be adopted for determining the chemical formula of any ionic compound.
Remember, in 1-to-1 ratio, the numbers cancel out. Like in MgO, it should be actually Mg2O2, but the 2s
cancel out, and we don’t write 1 in any chemical formula.
Physical Properties:
●
●
●
Most ionic compounds have high melting and boiling points.
They are usually soluble in water.
They are usually insoluble in organic solvents.
●
They do not conduct electricity in their solid state. However, they do conduct electricity in their
aqueous or molten state. This is because ions are free to move around when in aqueous or molten
states.
Chap 8: Covalent and Metallic Bonding
I). Covalent Bonds:
The bond formed between atoms that share electrons is called a covalent bond. After bonding, each
electron attains the electronic configuration of a noble gas, or attains stability. This bond is formed
between a non-metal and a non-metal.
A molecule is a group of two or more atoms held together by covalent bonds. Covalent compounds do not
have chemical formulae; instead, they have molecular formulae.
Formation and Arrangement:
Many non-metallic elements exist as diatomic molecules. This means that they always exist in the
molecular state, with two identical atoms joined together through covalent bonding. For example,
hydrogen gas always exists as H2, which is its diatomic form.
Fig 8.1. Covalent bond in Hydrogen molecule.
Note: always show the electrons of different members with different signs, i.e. one with a dot and the
other with a cross.
There are other ways of representing a hydrogen molecule as well:
Molecular Formula: H2
Dot & cross diagram: H·×H
Structural Formula: H
_
H (The single line shows a single covalent bond.)
The sharing of two electrons, or a ‘pair of electrons’, forms a single covalent bond.
Between each molecule of a covalent compound, there are weak van der Waal’s forces holding the
molecules together.
Other Examples:
_
Fig 8.3. Cl2 molecule, Cl Cl
Fig 8.2. O2 molecule, O=O
Covalent Compounds:
Molecules made from two or more different types of atoms linked together by covalent bonding are called
molecular compounds or covalent compounds.
A common example is H2O, water. It is made up of two single covalent bonds.
Fig 8.3. H2O
molecule,
Another example is CH4, methane. It has four single bonds in it.
Fig 8.4. CH4 molecule,
Carbon dioxide, CO2, has two double bonds.
Fig 8.5. CO2 molecule, O=C=O
Physical Properties:
●
●
●
●
They have low melting and boiling points.
They are mostly gases or volatile liquids at room temperature.
Most are insoluble in water, and soluble in organic compounds.
Most do not conduct electricity in either solid, liquid or gaseous state.
II). Metallic Bonding:
Metal atoms are held strongly to each other by metallic bonding. In the metal lattice, the atoms lose their
valence electrons and become positively charged. The valence electrons no longer belong to any metal
atom and are said to be ‘delocalised’. They move freely between the metal ions like a cloud of negative
charge. Hence, this lattice structure is described as a lattice of positive ions surrounded by a ‘sea of
mobile electrons’. Therefore a metallic bond can be defined as: ‘the force of attraction between positive
metal ions and the sea of delocalised electrons.’
Physical Properties of Metals:
●
●
●
●
They are good conductors of electricity.
They are good conductors of heat.
They are malleable (can be hammered into different shapes).
They are ductile (can be drawn into wires without breaking).
Chap 9: The Mole
(Part 1)
Understanding the concept of mole is extremely important in chemistry, as it will be there in
almost everything related to this subject. From paper point of view, it will be important in MCQs
and theory paper, as well as practical or A.T.P paper. To make it simpler, I have decided to teach
this chapter in two parts. This is the first part of the lecture.
Now, don’t get confused with the amount of things you must learn and understand; trust me, it
gets easier once you get the hang of it, and if you practice it enough. For this, we will be doing
some questions together, too.
A mole is a certain amount of substance. It is a general term to describe an amount of atoms, ions
or molecules, and it enables chemists to count these particles by weighing. A mole is defined as
the amount of substance which contains the Avogadro Number of particles. The Avogadro
Number (or Avogadro Constant) is defined as the number of atoms in 12g of the carbon-12
isotope, and its value is equal to 6.02 x 10²³.
Relative Atomic Mass:
The mass of one mole of atoms is its ‘relative atomic mass’ in grams. In other words, the relative
atomic mass of any atom is the number of times the mass of one atom of an element is greater
than 1/12 of the mass of one carbon-12 atom. Its symbol in chemistry is Ar.
Relative Molecular Mass:
The mass of one mole of molecules is its ‘relative molecular mass’ in grams.
Or you could say that the relative molecular mass in grams of any compound contains the same
number of molecules, equal to the Avogadro Number. Its symbol is Mr.
The relative molecular mass of ionic compounds is more accurately known as the ‘relative
formula mass’, as ionic compounds do not have molecules and thus using the word molecular
would be wrong. The symbol for this is also Mr. In all places that Mr is mentioned, it may mean
either.
To calculate the Mr of a compound, you must first know its molecular formula. First note down
the masses of all the elements present in the compound (you can take this information from the
periodic table), and then multiply the mass of each element with the number of atoms in one
molecule (the number in the subscript). Add the resulting numbers. This is your Mr.
Molar Gas Volume:
One mole of any gas at room temperature pressure occupies a volume of 24 dm³, or 24000 cm³ (1
dm³ = 1000cm³). This is sometimes known as the ‘molar gas volume’ as it contains the Avogadro
Number of particles. In more general terms, Avogadro’s Law states that equal volumes of all
gases at the same temperature and pressure contain the same number of particles.
Calculations with Moles:
Memorize the following formulae:
Number of Moles =
mass in grams .
Relative atomic mass
Or:
Number of Moles = mass in grams .
Relative molecular mass
These two formulas (which are basically the same; one has Ar and the other Mr) will help you as
long as you are studying chemistry.
Let’s do some questions on moles together. Remember, in the paper you will be provided the
periodic table from where you can see the atomic numbers and masses of respective elements, but
it is always good to memorize the repeated ones.
Q. How many moles are there in 54g of Water?
A. (keeping the above formulae in mind)
Mr of H2O = (2 x 1) + 16 = 18
Number of Moles = 54/18 = 3 moles.
Another type of question could be:
Q. How many grams are there in 7 moles of Carbon dioxide?
(In this question, we reverse the formula to obtain what we need.)
A. Mr of CO2 = 12 + (2 x 16) = 44
Mass in grams = 7 x 44 = 308 g
Now, let’s deal with volumes of gases. Memorize the following formulae:
1.
Volume of gas = mass of gas
Mr of gas
x molar gas volume
2.
Vol of gas = number of moles x molar gas volume
These two formulae, in this form or another, are going to be used repeatedly by you. Let’s do
some questions for practice.
Q. What volume (at r.t.p) would 0.2 moles of oxygen gas occupy?
A. Vol of O2 = 0.2 x 24 = 48 dm³.
Q. What volume would 32 g of methane gas occupy (at r.t.p)?
A. Mr of CH4 = 12 + (4 x 1) = 16
Vol of CH4 = 32/16 x 24 = 48 dm³.
Q. What mass in grams would 3 dm3 of Hydrogen gas occupy?
A. Mr of H2 = 2 x 1 = 2
Mass of H2 = 3/24 x 2 = 0.25 g
These were basic questions. Let’s now go to some trickier ones: ones that involve
equations as well as calculations.
Q. How many ions are there in 20 g of Magnesium Oxide?
A. Mr of MgO = 24 + 16 = 40
Number of moles of MgO = 20/40 = 0.5 mol
MgO
Mg2+ + O2-
From the equation we can see that 1 mol of MgO contains 1 mol of ions and 1 mol of ions. 0.5
mol of MgO will contain 0.5 mol of ions and 0.5 mol of ions.
Hence, 0.5 mol of MgO contains 1 mol of ions.
Number of ions = 1 x 6 x 10²³ = 6 x 10²³ ions.
Chap 10: Metals
The physical properties of any substance are determined by the way its particles are arranged or
packed. In a metal, atoms are packed tightly in layers and are held together by strong metallic
bonds. Due to these strong metallic bonds, metals usually have high densities, melting points and
boiling points. Due to the packing in regular layers, it is easy for the layers of atoms to slide over
each other when force is applied. This makes metals soft, ductile and malleable. The structure of
metals can also be described as positive metal ions surrounded by a sea of mobile electrons. The
mobile electrons allow metals to conduct electricity when they are connected to an electric
source. This also makes them good conductors of electricity.
Alloys:
Pure metals have many useful properties, but they are not widely used due to their softness and
high reactiveness. Most metallic substances used nowadays are alloys.
An alloy is a mixture of a metal with one or a few other elements. For example, bronze is an alloy
of copper and tin, brass is an alloy of copper and zinc, and steel is an alloy of iron, chromium,
nickel, and carbon.
Alloys are made by mixing the molten elements in the right proportions and then allowing them
to solidify. The alloys produced have more useful physical properties than the pure metals, such
as:
● Alloys are generally harder and stronger than their constituents. This is because when a
pure metal is alloyed, a different metal is added to the pure metal. Atoms of the added
element have a different size than those of the pure metal. This breaks up the regular
arrangement of atoms in the pure metal, making the sliding of layers over each other
difficult. Therefore alloys are harder and less malleable.
● Alloying can also be used to improve the appearance of metals.
● Alloys are also more resistant to corrosion than pure metals.
● Alloying is used to lower the melting points of metals.
The Reactivity Series:
Metals not only have same physical properties, they also undergo many similar chemical
reactions. However, one metal may react more or less vigorously with a substance than another
metal. The metal that reacts more vigorously is said to be more reactive than the other metal.
In the reactivity series, the metals are arranged in order of decreasing reactivity, i.e. at the top is
the most reactive and at the bottom is the least reactive. The reactivity series is as follows:
Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Carbon (C)
Zinc (Zn)
Iron (Fe)
Tin (Sn)
Lead (Pb)
Hydrogen (H)
Copper (Cu)
Mercury (Hg)
Silver (Ag)
Gold (Au)
Platinum (Pt)
Note how I included non-metals like hydrogen and carbon in the series. You will learn why later.
You should learn this reactivity series as it is. It is essential to know it during the paper, especially
when dealing with reactions of metals, and displacements. Let us learn more about these.
Reactions:
Metals with acids:
Many metals react with dilute acids to produce a salt and hydrogen. For example,
Mg (s) + HCl
(aq) MgCl + H2
In the above reaction, the salt magnesium chloride is formed. In all reactions of metals with
hydrochloric acid, a metal chloride and hydrogen gas are formed.
Note how I say ‘many’ metals react with dilute acid, not all. Only metals that are more reactive
than hydrogen react with dilute acids to produce hydrogen gas. Metals like copper and silver will
not react with the acids.
Decomposition:
The more reactive a metal is, the more difficult it is to decompose its oxides, i.e. reduce the oxide
to the metal.
Oxides of copper, as well as oxides of those metals above copper in the reactivity series, will
need the help of reducing agents such as carbon, along with the heat required for decomposition,
to turn them into their respective metal oxides. Below copper, there is no need for the reducing
agent, for the oxides of those metals simply decompose on heating. However, metal oxides of
magnesium, which is a highly reactive metal, will not decompose on heating, even in the presence
of a reducing agent. Same goes for the oxides of metals above magnesium in the reactivity series.
Other than carbon, hydrogen is also at times used as a reducing agent during decomposition.
However, it only reduces the oxides of iron and the metals below it in the reactivity series, while
the oxides of metals more reactive than iron will stay unreacted.
Let us now learn why some metals react while others don’t.
The Displacement Reactions of Metals:
More reactive metals can displace less reactive metals from their salt solutions, oxides, etc. for
example, solid iron displaces copper ions from a solution of copper (II) sulphate. Due to the
action of iron, copper metal is precipitated out of the solution as a pink solid.
In other words, atoms of the more reactive metal become ions and form compounds while ions of
the less reactive metal change back to atoms.
We see this behavior of metals in the reactions I told u about before this, too. In the reaction of
metal with acids, only the more reactive ones were able to displace hydrogen and form a salt,
while the less reactive metals like copper stayed unreacted.
In the decomposition reaction, we needed the reducing agent for more reactive metals. But metals
that were more reactive than carbon stayed unreacted even when the reducing agent carbon was
present. In case of the reducing agent hydrogen, it could not displace those that were more
reactive, so the metal oxides of metals more reactive than hydrogen stayed as they were, and the
reaction did not alter them.
For oxides of very reactive metals that do not get displaced and cannot be obtained as only the
metal through combustion, a process called electrolysis is used. You will learn more about
electrolysis in the next chapter.
Chap 11: Electrolysis
Electrolysis is the process of using electricity to break down or decompose a compound. It takes a
place in an electrolytic cell, which works like an electric circuit and has three main components: a
battery, electrodes and an electrolyte.
The plates which carry the electricity into the liquid are called electrodes. The cathode is the
electrode connected to the negative terminal of a cell. The anode is the electrode connected to the
positive terminal of a cell.
Electrolytes are compounds which when molten or dissolved in water conduct electric current and
are decomposed in the process. Here, electricity is passed through the liquid by the movement of
ions between the electrodes. Examples of electrolytes are acids, alkalis or salts dissolved in water
or molten salts. All these are ionic substances. A non-electrolyte is a liquid which does not allow
the passage of electricity, and its examples are water and all covalent substances which do not
contain ions.
Process:
When electricity is passed through an electrolyte, it undergoes chemical decomposition. The
electrolyte splits up. The ions migrate towards the oppositely charged electrode. At the anode,
negative ions (anions) lose their electrons to the anode, which is very ready to accept electron as
it is positively charged, and has a lack of electrons. At the cathode, positive ions (cations) gain
electrons from the cathode, which has an excess of electrons and therefore an overall negative
charge. This release of ions at the electrode results in the chemical decomposition of the
electrolyte. It also allows electrons to travel from the cathode to the anode, and the movement of
ions during electrolysis allows conduction of electricity.
The process of losing or gaining electrons at the electrodes is called discharge. When ions are
discharged at the electrodes, they form atoms or molecules.
Factors Affecting Electrolysis:
In electrolysis, when more than one type of cation or anion is present in a solution, only one
cation or one anion are preferentially discharged. This is known as selective discharge of ions.
Position in the Reactivity Series (for cations):
The cations of an element lower in the reactivity series are discharged at the cathode in preference
to other cations in the solution. This is because cations of a less reactive element accept electrons
more readily.
Relative Ease of Discharge (anions):
The hydroxide ions give up electrons most easily, followed by iodide, bromide and chloride ions.
Remember that sulphate and nitrate ions will not be discharged during electrolysis.
Concentration:
If the concentration of a particular ion is high, then this can alter the preferential discharge. The
higher concentration makes it more likely to be given off.
Type of Electrode:
Let us take as an example the electrolysis of aqueous copper (II) sulphate solution. If we use
carbon electrodes, they will be inert electrodes and will not affect the electrolysis. However, if we
use copper electrodes, these are active electrodes and do affect the electrolysis. At the anode, the
copper electrode will dissolve into the solution, while at the cathode, the copper ions will be
deposited as pink copper metal.
Industrial Application of Electrolysis:
Electrolysis has many varied industrial applications. The extraction of metals from their ores, in
particular aluminium, and the purification of metals, especially copper, are important industrial
uses.
Another large scale use of electrolysis is in the manufacture of the important alkali, sodium
hydroxide. This is produced by the electrolysis of concentrated sea water. The chlorine and
hydrogen gases which are produced during this electrolysis are both commercially useful.
Finally, electrolysis is used for electroplating, which is the forming of a thin protective coating of
a metal on the surface of another which is likely to corrode.
Chemical to Electrical Energy:
A device which converts chemical energy into electrical energy is called a cell, or battery (a
collection of cells). It consists of a pair of dissimilar metals in an electrolyte. The more reactive
metal dissolves and turns into ions, thereby producing electrons. These electrons then travel to the
less reactive metal electrode, and bubbles of hydrogen are produced at this electrode.
This production or movement of electrons is electricity, so electrical energy has been generated
and the bulb that completes the circuit and is connected to the electrodes lights up.
How bright the bulb is depends on the difference in the reactivities of the two metals. If the
metals are far apart in the reactivity series, the bulb is bright, but does not stay alight for a very
long time. If the metals are closer in the reactivity series, the bulb is less bright, but stays alighted
for a longer time.
If the electrolyte in the cells is a paste, as opposed to a liquid, it is called a dry cell. These are the
sort of cells most commonly found in the home. They are a convenient, portable energy source
and are reasonably cheap to buy.
Chap 12: Speed of Reaction
The speed at which a chemical reaction proceeds can vary, and is called the rate of reaction. We
will study about this in detail in this chapter.
I). Measuring Speed of Reaction:
The speed of reaction can be found by measuring these quantities at regular time intervals, when
the chemical reaction is in progress.
● The volume of a gas produced by the reaction.
● The mass of the reactant that remains.
Either of the mentioned is measured, and a graph of that against time is plotted. The speed of
reaction at any instant is measured by drawing a tangent against that point on the graph, and
calculating its gradient.
II). Factors Affecting Speed of Reaction:
There are various factors that affect the speed of a chemical reaction, the major ones being:
● The temperature at which the reaction is occurring
● The concentration of the substances used (reactants)
● The pressure on the reaction (if the reactants are gaseous)
● The particle size or surface area of the reactants.
For the reaction to occur between two particles, the reacting particles must collide with each
other, and they must collide with a certain minimum amount of energy, known as activation
energy.
The more the collisions of particles, the faster the rate of reaction. In other words – as we are
talking about the factors that affect the speed of reaction – when any factor increases the rate of
effective collisions between reacting particles, it will also increase the speed of reaction.
Temperature:
The higher the temperature, the faster the rate of reaction. You see, when the temperature is
raised, the reactant particles have a greater heat energy, causing them to move about more and
with a greater kinetic energy. They, therefore, stand a better chance of colliding into another
reactant molecule with sufficient energy to convert into product molecules.
Concentration:
The more concentrated the reactants, the faster the speed of a chemical reaction. This is because
at a higher concentration, there is a greater likelihood that reacting molecules will collide with
one another with sufficient energy to form products.
Pressure:
In the reactions that involve gases, the speed of reaction increases if the pressure is increased.
Again, this is really the effect of concentration as higher pressures force the particles close
together and so their concentration within a certain volume increases. More collisions therefore
occur, and the speed of reaction increases.
Particle Size:
The greater the surface area, the faster the rate of a chemical reaction. Smaller particles like
powders have a much greater surface area than larger lumps or crystal. With a greater surface
area, the other reactant can attack it more easily, and thus the speed of reaction increases.
III). Catalysts and Enzymes:
A catalyst is a substance which increases the rate of a chemical reaction by lowering its activation
energy, without itself being chemically changed at the end of the reaction. It works by providing
an alternate, more direct route from reactants to products.
Enzymes are biological catalysts; they are present in all living organisms and are responsible for
breakdown of foodstuffs.
As I mentioned, catalyst speed up chemical reactions. So that is another factor that affects the rate
of chemical reactions: the presence of catalysts.
Chap 13: Redox Reactions
Redox reactions are reactions that involve oxidation and reduction. There are several definitions
of oxidation and reduction, but we will go through a general idea of them.
Oxidation:
Oxidation is a gain in oxygen, or a loss of hydrogen. Any combustion process can be classified as
oxidation, as it involves a gain in oxygen. Oxidation can also be defined as a loss of electrons, or
an increase in the oxidation state. Substances which help oxidation to take place are called
oxidising agents. Oxygen itself is the obvious oxidising agent, but substances with lots of oxygen
in their molecules can also act as oxidizing agents. An oxidising agent itself, however, becomes
reduced.
Reduction:
Reduction is a loss of oxygen, or a gain in hydrogen. It is also a gain in electrons, or a decrease in
the oxidation state. Substances which help reduction to take place are called reducing agents.
These agents themselves become oxidized to reduce.
For example, let us consider this ionic equation:
Cu2+ (s) + H2 (g) Cu (s) + 2H+ (aq)
This is a redox reaction. In it, Copper’s oxidation state changes from +2 to 0. As it is a decrease
in oxidation state, it means that in this reaction copper is being reduced.
On the other hand, Hydrogen’s oxidation state changes from 0 to +1. As it is an increase in
oxidation state, it means that hydrogen is being oxidised.
Oxidising agents:
These are the electron acceptors. Some examples are:
● Non-metals (chlorine, bromine, etc). This is because non-metals are good at accepting
electrons.
● Acidified Potassium Manganate (VII)
● Acidified Potassium Dichromate (VI)
Reducing Agents:
These are the electron donors. Some examples are:
● Metals (potassium, sodium, etc). This is because they are good at donating electrons.
● Potassium Iodide
● Carbon
● Hydrogen
● Carbon monoxide
Color Changes
It is often helpful to identify redox reactions by the color changes involved. Some of the color
changes that you should remember are:
● Acidified Potassium Manganate (VII): this is a powerful oxidising agent. It works by
reducing its Manganate (VII) ion to Manganate (II) ion. The colour changes from purple
to colourless.
●
●
Acidified Potassium Dichromate (VI): this is also a powerful oxidising agent. It works by
reducing its dichromate (VI) ion to chromium (III) ion. The colour changes from orange
to green.
Aqueous Potassium Iodide: this is a powerful reducing agent. Its iodide ion is oxidised to
iodine. The colour changes from colourless to brown.
These colour changes can also be used as tests. Test for oxidising agent by adding potassium
iodide which turns from colourless to brown, if present. Test foe reducing agent by adding
acidified potassium dichromate (VI) which turns from orange to green, if present.
Let us do a couple of questions together now.
Q1. Which underlined agent is ‘not’ an oxidising agent?
A. Cl2 (g) + 2Br-
2Cl- (aq) + Br2 (l)
(aq)
B. PbS (s) + 4H2O2
(aq)
C. SO2 (g) +
2H2S(g)
D. Mg (s) + Cl2
(g)
PbSO4 (s) + 4H2O (l)
2H2O (l) + 3S (s)
MgCl2 (s)
For answer, let us consider all options one by one.
In option A, Cl2 is underlined. Its oxidation state changes from 0 to -1, which means it is being
reduced. We learnt that oxidising agents themselves get reduced. However, an alternate and a
more accurate method would be to see weather or not Br is getting oxidised. Its oxidation state
changes from -1 to 0, which means that it is getting oxidised. Therefore Cl2 here is acting as an
oxidising agent, so it is eliminated as the answer.
In option B, H2O2 is underlined. So let’s take a look at PbS. There can be two ways. Either check
the oxidation state, which changes from 0 to +8 (increases), or you can see that it is gaining
Oxygen. It both ways, you can see that PbS is getting oxidised. Therefore, H2O2 is acting as an
oxidising agent, so we eliminate option B as well.
In option C, SO2 is underlined. So let’s look at H2S. The hydrogen gains oxygen, while the
Sulphur’s oxidation state changes from -2 to 0 (increases). Therefore SO2 is an oxidising agent.
So we eliminate option C as well.
In option D, Mg is underlined. So we take a look at Cl2. its oxidation state changes from 0 to -2,
which is a decrease. Therefore, we can say that Cl2 is getting reduced, so Mg is acting as a
reducing agent.
So, option D is the right answer.
Q2. For the following equations, state which substances have been oxidised, and which have been
reduced, giving your reasons.
a). C (s) + H2O
(g)
CO (g) + H2 (g)
Ans. Oxidation: Carbon is being oxidised in this reaction. Its oxidation state increases from 0 to
+2. It also gains oxygen, changing from C to CO.
Reduction: Hydrogen is being reduced. Its oxidation state decreases from +2 to 0. It also loses
oxygen.
b). Fe2O3 (s) + 3CO
(g)
2Fe (s) + 2CO2 (g)
Ans. Oxidation: CO is getting oxidised in this reaction. Its oxidation state increases from 0 to +2,
and it also gains oxygen.
Reduction: Fe becomes reduced. It loses oxygen. Its oxidation state decreases from +3 to 0.
Chap 14: Extraction of Metals
Dear students, this is a simple and short chapter. We will learn in this lecture how and
why must we extract metals from their natural ores.
We will start off with ores. Ores are natural compounds of metals found in the ground,
often oxides, sulphides or carbonates. Obtaining metals from their ores is a reduction
process as the metal ion has to gain electrons. Ores of metals high in reactivity series are
difficult to extract and electrolysis is used. Middle order metals like iron are extracted by
reducing with carbon in the blast furnace.
Blast Furnace:
The main ore of iron is called haematite. We extract iron from this.
The furnace is charged from the top, with our three raw materials: haematite, coke and limestone.
It is heated by blowing hot air in at the base of the furnace, through pipes. What happens now is
that the coke present in the furnace burns, and as it is a very exothermic reaction, we get all the
heat we need. The coke also burns to provide carbon monoxide, which is the reducing agent for
the ore. The limestone present in the furnace removes the main impurity in the ore, which is sand
(silicon dioxide). The iron produced in the blast furnace is called pig iron and contains about 4%
carbon.
Steel:
It is produced by purifying the pig iron from the blast furnace by burning off impurities with
oxygen gas. Controlled amounts of carbon are then added (0.5% carbon is mild steel, 1.5%
carbon is high carbon steel). Mild steel is an example of steel that is low in carbon content, and is
soft and easy to mould. It is used in car bodies and machinery.
Rusting:
It is the reaction of iron with air (oxygen) and water to form rust (hydrated iron(III) oxide).
Rusting is a redox process. Many a times, this rusting needs to be prevented, and for that some
methods have been devised, for example, rusting can be prevented by painting, coating with oil or
grease, metal planting or electroplating, and also by sacrificial protection.
Stainless steel is an example of an alloy of iron that has been made specifically resistant to
corrosion.
Aluminium Extraction:
The main ore of aluminium is bauxite (aluminium oxide). After it is mined, it is first purified by
adding it in sodium hydroxide. The impurities do not dissolve and are filtered off. The dissolved
aluminium oxide is then precipitated out as aluminium hydroxide by diluting with water. This is
then heated to form pure white aluminium oxide, or alumina. The aluminium is then extracted by
electrolysis.
The alumina is added to and electrolytic cell and melted. A compound called cryolite (another
aluminium compound) is added to lower the melting point of the alumina. The electrodes are
made of graphite. Oxygen gas collects at the anode. The cathode is very interesting; it is the
lining of the electrolytic cell that is made the cathode. Molten aluminium collects here.
At cathode: Al3+ (l) + 3e-
Al (l)
At anode: 2O2- (l)
O2 (g) + 4e-
Aluminium is a more reactive metal than its apparent lack of reactivity is explained by a thin
protective oxide coating. This coating (unlike iron oxide) does not flake off and makes aluminium
corrosion resistant.
Aluminium metal is relatively light, corrosion resistant, and a good conductor of heat and
electricity. When alloyed, it becomes quite resistant.
Aluminium can be used in making kitchen foils due to its malleability and heat conducting
ability, in overhead power cables due to its light weight and electrical conducting abilities, and in
aircraft construction due to its lightness and resistance to corrosion.
Chap 15: Acids and Bases
Acids:
Acid is a substance which produces hydrogen ions, H+, when it is dissolved in water.
All acids contain hydrogen ions, but not all compounds that contain hydrogen are acids, e.g. NH3.
This is because NH3 does not produce hydrogen ions in water. It is the hydrogen ions produced
that are responsible for the properties of acids.
Properties of Acids:
1.
2.
3.
4.
5.
6.
Acids have a sour taste.
They have a pH below 7.
They are generally corrosive in nature. The lower the pH, the more corrosive they are.
They turn blue litmus paper red.
They dissolve in water to form solutions that conduct electricity.
Acids react with reactive metals to form hydrogen and a salt.
metal + acid
salt + hydrogen
7.
Acids react with carbonates to form a salt, carbon dioxide and water.
carbonate + acid
salt + carbon dioxide + water
8.
Acids react with metal oxides and hydroxides to form a salt and water.
metal oxide + acid
salt + hydrogen
Acids only show the properties of acids when they are dissolved in water. This is because acids
dissociate in water to produce the hydrogen ions which are responsible for the acidic properties.
Bases:
A base is any metal oxide or hydroxide. This means that a base contains either oxide ions O2- or
hydroxide ions OH-.
We can also define a base as a substance that reacts with an acid to give a salt and water only.
base + acid
salt + water
For example,
CuO (s) + H2SO4 (aq)
CuSO4 (aq) + H2O (l)
In this case, copper (II) sulphate CuSO4 is the salt produced.
Another example,
NaOH (aq) + HCl (aq)
NaCl (aq) + H2O (l)
In this case, sodium chloride NaCl is the salt produced.
Note in the above cases, the oxide ions or the hydroxide ions from the bases react with the
hydrogen ions from the acids to form water. This reaction is known as neutralization, and it is
represented by the ionic equation below.
OH- (aq) + H+ (aq)
H2O (l)
It occurs between alkalis and acids.
Alkalis:
An alkali is a base that is soluble in water. In an aqueous solution, an alkali produces hydroxide
ions OH-.
Most bases are insoluble in water. They are not considered as alkalis. Some common examples of
alkalis include sodium hydroxide NaOH, potassium hydroxide KOH, calcium hydroxide Ca(OH)2
and aqueous ammonia NH3.
Properties of Alkalis:
Alkalis have a bitter taste and a soapy feel.
They have a pH above 7.
The higher the PH, the more corrosive they are.
They turn red litmus paper blue.
All alkalis produce hydroxide ions when dissolved in water.
All alkalis can react with water to form salt and water only. This is known as
neutralization. We have already discussed this.
7. Alkalis heated with ammonia salts give off ammonia gas.
alkali + ammonium salt
ammonia + salt + water
1.
2.
3.
4.
5.
6.
8.
Alkalis can react with a solution of one metal salt to give a metal hydroxide and another
metal salt.
alkali + salt (of metal A)
metal hydroxide + salt (of metal B)
Uses of Acids:
Sulphuric acid:
● In making detergents
● In making fertilizers
● In car batteries
Hydrochloric acid:
● In leather processing
● For cleaning metals
Ethanoic acid:
● (In vinegar) To preserve food
● In making adhesives such as glue
Uses of Bases:
Ammonia solution:
● In window cleaning solutions.
● In fertilizers.
Calcium oxide:
● In neutralizing acidic soil
Sodium hydroxide:
● In making soaps and detergents
Concentration and Strength:
The term concentration tells us how much a substance is dissolved in 1 dm3 of the solution. This
is changeable, of course. By adding more of the substance, we increase its concentration. By
adding more of the solvent, we decrease the concentration.
The term strength refers to how easily an acid or an alkali dissociates when dissolved in water.
This is a property that is not changeable.
Types of Oxides:
An oxide is a compound of oxygen and another element. Most oxides can be grouped into four
types: acidic oxides, basic oxides, amphoteric oxides, and neutral oxides.
Acidic Oxides:
Non-metals may form acidic oxides. Most acidic oxides dissolve in water to form an acid. These
oxides do not react with acids. However, they react with alkalis to form salt and water.
An example of acidic oxide is sulphur trioxide SO3. it is a gas which, when dissolved in water,
forms Sulphuric acid H2SO4.
Basic Oxides:
The oxides of metals are basic oxides. Most basic oxides are insoluble in water. Those that are
soluble are called alkalis.
Basic oxides are solids at room temperature. They react with acids to form a salt and water. They
have no reaction with bases or alkalis.
An example of basic oxide is calcium oxide. It dissolves in water to form calcium hydroxide,
while it reacts with hydrochloric acid or any other acid to form calcium chloride or any other salt,
in addition with water.
Amphoteric Oxides:
Amphoteric oxides are metallic oxides that react with both acids and bases to form salts and
water. Meaning, amphoteric oxides are oxides that contain both the abilities of acids and alkalis.
If an amphoteric oxide is reacted with an acid, it will show the properties of an alkali and a
neutralization reaction occurs. Similarly, if it is reacted with an alkali, it will show acidic
properties and here, too, a neutralization reaction will occour.
Examples of amphoteric oxides are aluminium oxide, zinc oxide, lead (II) oxide, etc.
Neutral Oxides:
Some non-metals form oxides that show neither basic nor acidic properties. These oxides are
called neutral oxides. They are insoluble in water. Some examples are carbon monoxide and nitric
oxide. Water itself is also a neutral oxide.
Chap 16: Salts
All salts are ionic compounds. A salt is formed when a metallic ion or an ammonium ion replaces
one or more hydrogen ions of an acid. We will study about this substance in detail in this chapter.
Solubility:
All nitrates, and all ammonium, sodium and potassium salts are soluble in water. Most chlorides
and sulphates are also soluble, while the majority of carbonates are insoluble.
Learning the solubility of salts is important before learning how the salts are formed. The
solubilities in detail are:
Nitrates: All soluble
Chlorides: All soluble, EXCEPT silver chloride AgCl, lead (II) chloride PbCL2, and mercury (II)
chloride HgCl2.
Sulphates: All soluble, EXCEPT calcium sulphate CaSO4, barium sulphate BaSO4 and lead (II)
sulphate PbSO4.
Carbonates: All insoluble, except group I carbonates and ammonium carbonate.
Preparation:
We will learn three methods to prepare salts. Before deciding which method to use, we must
know the answer to these two questions.
1. Is the salt to be prepared soluble in water?
2. Are the starting materials soluble in water?
Method 1:
This method is used if the salt to be prepared is soluble in water, while the starting materials are
insoluble.
For example, let’s see how to prepare zinc sulphate. We learnt that it is a soluble salt. Let’s think
of some starting materials for it now. For the ‘sulphate’ part, we could use Sulphuric acid H2SO4.
For the ‘zinc’ part, we could use the metal zinc itself.
Zn(s) + H2SO4
ZnSO4 + H2
In this method, we react the acid with an excess of the substance. Zinc is an insoluble metal. All
the acid will be used up as it reacts with the zinc, while the excess zinc can be filtered out.
What we do now is that we take the Sulphuric acid in a beaker and add zinc into the acid while
stirring. We keep adding the metal until no more effervescence is observed. Now we have in the
beaker zinc sulphate solution and unreacted zinc. We filter the mixture to get rid of the excess
metal, and collect the filtrate. This filtrate is only the zinc sulphate solution. Now we heat this
filtrate to concentrate the solution, and then leave it to cool so that it can crystallize. These
crystals are zinc sulphate crystals. Wash these crystals with distilled water to remove impurities,
and dry between sheets of filter paper. These are now pure zinc sulphate crystals.
Method 2:
This method, known as titration, is suitable for preparing salts that are soluble, and their starting
materials are also soluble.
For example, let’s see how to make sodium chloride. It is a soluble salt. We can make it with the
alkali sodium hydroxide, and the acid hydrochloric acid.
NaOH (aq) + HCl (aq)
NaCl (aq) + H2O
First, we add the sodium hydroxide in a beaker, while the hydrochloric acid is added to the
burette, which is clamped to a stand and is placed vertically above the beaker. We add some
indicator ‘phenolphthalein’ to the beaker. Phenolphthalein is colorless in acid and neutral
solutions while pink in alkali, so it turns pink in the beaker. We open the burette tap to slowly let
the acid flow into the burette and watch closely for a color change in the beaker. As soon as it
turns colorless, we stop the burette tap and record the reading. In the beaker, we have now sodium
chloride salt, as well as the impurity phenolphthalein. As we want only the pure salt, we repeat
this whole experiment again, this time without the indicator. We add the same amount of alkali
into the beaker, and put in the same volume of the acid that we noted in the first experiment.
Now that we have the salt, we repeat the same processes that were used in the first method. We
concentrate the salt by heating it, and then let it cool and dry so that it can crystallize. We filter
and collect the crystals, wash them with distilled water, and then dry between sheets of filter
paper. We now have pure crystals of sodium chloride.
Method 3:
This method is known as ‘precipitation’. This method is used to prepare insoluble salts. All
insoluble salts can be prepared using this method.
For example, let’s learn how to make lead (II) sulphate, PbSO4.
We will first mix two solutions: one that contains the positive ions, and the other that contains the
negative ions of the salt to be prepared. We can use lead nitrate for the lead part, and Sulphuric
acid for the sulphate part.
This method, you will learn, is the simplest method. We add the lead (II) nitrate solution to a
beaker, and add Sulphuric acid (in excess) and stir until no more precipitate forms. Then, filter to
collect the precipitate, which is actually our salt lead (II) sulphate. To remove impurities, wash
with distilled water, and then dry between filter papers.
Qualitative Analysis:
The process of identification of cations and anions is called qualitative analysis or salt analysis.
We identify which cations (positive ions) are present in a salt by using aqueous sodium
hydroxide, and aqueous ammonia. For cation identification, learn the following table.
Sodium Hydroxide solution, NaOH (aq)
Aqueous Ammonia, NH3 (aq)
Observations on adding
Observations on adding
a. a few drops of NaOH
a. a few drops of NH3
b. excess NaOH
b. excess NH3
3+
Al
a. a white precipitate is formed
a. a white precipitate is formed
b. the precipitate dissolves in excess to give a
b. precipitate insoluble in excess
colourless solution
Ca2+
a. white ppt formed
a. no ppt
b. ppt insoluble in excess
b. no ppt
Zn2+
a. white ppt formed
a. white ppt formed
b. ppt dissolves in excess to give a colorless
b. ppt dissolves in excess to give
solution
a colorless solution
2+
Cu
a. light blue ppt
a. light blue ppt
b. ppt insoluble
b. ppt dissolves in excess to give
a deep blue solution
Fe2+
a. green ppt
a. green ppt
b. ppt insoluble
b. ppt insoluble
Fe3+
a. reddish-brown ppt
a. reddish-brown ppt
b. ppt insoluble
b. ppt insoluble
+
NH4
a. no ppt is formed
No reaction
b. On heating, ammonia gas is given off.
Ammonia turns moist red litmus paper
blue.
To identify anions, learn these reactions:
CO32- (carbonate):
Test: add dilute hydrochloric acid. Pass the gas given off into limewater.
Observation: effervescence is observed. Gas given off forms a white ppt with lime water. Carbon
dioxide gas is given off.
Cl- (chloride):
Test: add dilute nitric acid, then add silver nitrate solution.
Observation: a white ppt of silver chloride is formed.
I- (iodide):
Test: add dilute nitric acid, then add silver nitrate solution
Observation: a yellow ppt of silver iodide is formed.
NO3- (nitrate):
Test: add dilute sodium hydroxide. Then add a piece of aluminium foil. Warm the mixture. Test
the gas given off with a piece of moist red litmus paper.
Observation: the moist red litmus paper turns blue. Ammonia gas is given off.
SO42- (sulphate):
Test: add dilute nitric acid, then add barium nitrate solution.
Observation: a white ppt of barium sulphate is formed.
Chap 17: Energy from Chemicals
Exothermic and Endothermic Changes:
An exothermic reaction is one where energy (heat) is given out causing a temperature rise in the
surroundings. The temperature of the reaction mixture rises. The container feels hot. Examples of
exothermic reactions include:
● Combustion of fuels
● Rusting of iron
● Corrosion of metals
● Respiration
● Neutralization reaction
An endothermic reaction is one where energy (heat) is taken in causing a temperature drop in the
surroundings. The temperature of the reaction mixture falls. The container feels cold. Examples
of endothermic reactions include:
● Photosynthesis
● Thermal decomposition
The amount of energy involved in a reaction is known as he heat change or the enthalpy change
of the reaction. ΔH for an exothermic reaction is negative. This is because chemicals have lost
energy to the surrounding. ΔH for an endothermic reaction is positive. This is because the
chemicals gain energy from the surroundings.
Making and Breaking Bonds:
An exothermic change occurs when chemical bonds are made.
An endothermic change occurs when chemical bonds are broken.
The overall energy change of a reaction is the difference between the energy given out when
bonds are made and the energy taken in when bonds are broken.
Activation Energy:
All reactions need energy in order to get started, weather it is an endothermic reaction or an
exothermic one. The minimum energy that reacting particles must possess in order for a chemical
reaction to occur is called the activation energy, EA.
Energy Profile Diagrams:
To show the activation energy of a reaction, energy profile diagrams are used. This diagram is a
way of representing the energy changes that occur during a chemical reaction. The energy
difference between the products and reactants represents the enthalpy change of the reaction. So,
an energy profile diagram shows the activation energy required and the enthalpy change for a
reaction.
The energy profile diagram for an exothermic reaction would be:
The energy profile diagram for an endothermic reaction would be:
Chap 18: Ammonia and its Uses
Nitrogen:
Nitrogen is a gas. It is the main constituent of air (78%). It plays an important role in the
formation of animal and plant protein. It is also used in the manufacture of ammonia gas, an
important industrial gas, through the Haber Process. We will learn how in this chapter.
Haber Process:
Ammonia is the most important compound of nitrogen. It is produced industrially by the Haber
Process.
The raw materials for the process are hydrogen and nitrogen. Hydrogen is produced industrially
by cracking oil, and nitrogen from liquefaction of air. These two gases, when obtained, are
combined directly in a ratio of 3:1, and are passed over an iron catalyst at a temperature of 450°C
and a pressure of 200 atm.
The reaction between hydrogen and nitrogen to form ammonia is a reversible process. This means
that some of the ammonia formed may revert back to nitrogen and hydrogen. So to achieve the
maximum yield of ammonia at the minimum cost, the reaction conditions are very carefully
controlled.
● The yield of ammonia is increased under high pressures. High pressure also increases the
rate of reaction. However, maintaining high pressure is costly because expensive
equipment is required.
● A lower temperature increases the yield of ammonia and reduces the decomposition of
hydrogen to nitrogen. However, a lower temperature also results in a slow reaction. At
industry level, we cannot afford slow reactions. Therefore the high temperature of 450°C
is a compromise for a faster reaction.
● Even the high temperature and high pressure used to react the nitrogen and hydrogen
result in a slow reaction, so an iron catalyst is used to speed up the reaction.
Displacement of Ammonia from its Salts:
Another way of obtaining ammonia gas is through the displacement of ammonia from its salt. For
example,
NH4Cl (s) + NaOH (aq)
NH3 (g) + H2O (l) + NaCl (aq)
In this reaction, you can see that ammonia gas is produced on heating. Whenever an ammonium
salt is heated with an alkali, ammonia is displaced from the salt. Lets take the example of another
reaction.
2NH4Cl (s) + Ca(OH)2 (aq)
2NH3 (g) + 2H2O (l) + CaCl2 (aq)
Testing for Ammonia Gas:
Ammonia is the only common alkaline gas, so it can be identified by moist red litmus paper
turning blue. However, we will learn a more specific chemical test for it now. In this test, we hold
a glass rod dipped into some concentrated hydrochloric acid close to the suspected ammonia. This
will give off fumes of hydrogen chloride gas which, in the presence of ammonia, form a dense,
white ‘smoke’ of ammonium chloride.
In the same way, mixing a gas jar of hydrogen chloride and ammonia gas produces the same
dense, white smoke, which is actually ammonium chloride.
Uses of Ammonia:
The production of ammonia is done on large scale and large quantities are produced because it
has many important uses.
Manufacture of Fertilizers:
The main use of ammonia is in the manufacture of fertilizers. Approximately 75% of all ammonia
produced is converted into various ammonium compounds like ammonium sulphate (NH4)2SO4,
ammonium nitrate NH4NO3, and urea NH2CONH2. These compounds are called nitrogenous
fertilizers. They are solids for easy handling, and water soluble so that they seep into the soil to be
absorbed by the roots of the plant. These fertilizers provide nitrogen to the plant, which is an
important element for healthy plant growth, as it helps in making the proteins which are needed
for healthy growth of stems and leaves.
There are other important elements for healthy plant growth, too, such as phosphorus, which is
necessary for good root growth, and potassium, which is important in the development of flowers
and food. These important elements come from other sorts of fertilizers. If a fertilizer contains
only one of the important elements, it is called a straight fertilizer. If these fertilizers are mixtures
of the important elements, they are called compound fertilizers.
Beside these elements, there are also trace elements needed for the healthy growth of a plant, for
example calcium, magnesium, sodium, sulphur, and tiny amounts of copper, zinc, boron,
manganese and iron.
But what you basically must keep in mind in this chapter is that how ammonia is used to
manufacture fertilizers. These are artificial fertilizers. The nitrogen content of such a fertilizer can
be calculated by percentage composition.
The use of artificial fertilizers, however, has its disadvantages. Large scale use of these causes a
pollution problem eutrophication: the high solubility of the fertilizer allows it to be leached from
the soil and washed into streams, and when it finally settles in still water, it causes algae to grow
and removes oxygen from the water.
Manufacture of Nitric Acid:
Nitric acid is made by the catalytic oxidation of ammonia over heated platinum. Oxidising
ammonia produces oxides of nitrogen which can then be dissolved in water to produce nitric acid.
Most of the nitric acid made is used to make the all-important fertilizers. Other uses of nitric acid
include making explosives and making dyes.
Use as a Solvent:
Aqueous ammonia is used as a degreasing agent, as it is a good solvent of grease and fat.
Let us go through some questions now, so that we can get a clearer idea of the whole chapter.
Q1. Which of these compounds does not always contain nitrogen?
A. Protein
B. Fertilizer
C. Nitrate
D. Nitrite
The answer to that question, as most of you might have already deduced, is B, Fertilizers. I
mentioned this fact earlier in my lecture that fertilizers are generally made up of ammonia,
potassium, or phosphorus, not necessarily ammonia.
Now solve this question yourself.
Q2. Which of the following reactants will not produce ammonia on heating?
A. Ammonium chloride and potassium hydroxide
B. Ammonium sulphate and calcium oxide
C. Potassium hydroxide and sodium nitrate
D. Sodium hydroxide and ammonium nitrate
Chap 19: Sulphuric Acid
Hello students. In this chapter, we learn about how Sulphuric acid is manufactured by a process
called contact process. But before we begin, let’s learn a little about the oxides of sulphur:
sulphur dioxide and sulphur trioxide.
Both oxides of sulphur are acidic in aqueous solution: sulphur dioxide dissolves to form
sulphurous acid and sulphur trioxide dissolves to form Sulphuric acid.
Sulphur dioxide is much more common than the higher oxide and can easily be obtained by the
combustion of sulphur in air. It has several uses. Its main use is in the manufacture of Sulphuric
acid. Other than that it is used as a bleach, especially in bleaching wood pulp for making paper,
and as a food preservative.
Sulphur trioxide is the higher oxide of sulphur and is more correctly called sulphur (VI) oxide. It
is much more difficult to prepare as it involves the catalytic oxidation of sulphur dioxide.
Industrially, it has great significance, and is used in the manufacture of Sulphuric acid.
The Contact Process:
The sulphuric acid has great industrial importance, therefore it has to be prepared on large-scale.
It is manufactured industrially by the contact process. Let us now learn this process.
In the first step, we burn sulphur in air (oxygen). (Sulphur and air (oxygen) are our raw
materials). This way, sulphur dioxide SO2 is formed.
In step two, we purify the sulphur dioxide, by removing impurities like arsenic compounds which
would otherwise poison the catalyst. It is then passed through an electrostatic dust precipitator,
which charges dust particles that are then removed by being attracted to oppositely charged
plates.
The third step is then performed. In this step, sulphur dioxide and air and washed and passed over
the catalyst I mentioned earlier: vanadium (V) oxide. This is done at a temperature of 450C and at
pressures 2-3 atm. This reaction is reversible and exothermic. The sulphur dioxide and oxygen in
air react to produce sulphur trioxide.
The next step is to dissolve the sulphur trioxide produced in concentrated sulphuric acid, to form
oleum, or fuming sulphuric acid.
In step five, we dilute this oleum with water to the required strength of acid.
Although we can achieve sulphuric acid directly by dissolving sulphur trioxide in water, we
cannot practically dissolve these two together, as the resulting reaction is too violent.
Uses of Sulphuric Acid:
There are many uses of sulphuric acid industrially. Some of them are as follows:
● Sulphuric acid is used in the production of fertilizers, e.g., ammonium sulphate,
potassium sulphate, etc.
● It is used in the manufacture of non-soapy, organic detergents.
● It is used as an electrolyte inside car batteries.
● It is used in the cleaning of metals by removing the surface oxide coating. This is called
pickling.
● It is used in the making of artificial silks.
These are enough for you to learn at this stage. Thank you students.
Chap 20: Environmental Chemistry
In this chapter, we study about air, its pollutants, oxygen, hydrogen, water and more
environmental issues and facts.
Air and Oxygen:
We all know that oxygen is important for human and animal life. It is even important for plant
life. The process by which living organisms produce energy from their food is called respiration,
and oxygen is essential for this process.
Oxygen is one of the main constituent of air. 20% of air is oxygen. The nitrogen in air is almost
inert at room temperature, and after that the highest composition is of oxygen. So when any
reaction says ‘react with air’, it actually means react with O2. However, we have to study other
things about oxygen and the air today.
Combustion:
This is the process of burning. When anything burns, it requires fuel, heat and oxygen.
Combustion actually means burning in the presence of oxygen.
Oxides:
When oxygen reacts with elements (metals or non-metals), the oxide of that element is formed.
Oxygen is very reactive, and reacts with most metals.
For example, when oxygen O2 reacts with sodium Na, it forms sodium oxide Na2O.
When oxygen reacts with carbon C, it forms carbon’s oxide which is known as carbon dioxide,
CO2.
Air Pollution:
The addition of poisonous gases in the environment is called air pollution. These gases include
carbon monoxide CO, sulphur dioxide SO2, etc, and are called pollutant gases. Some pollutants
like the oxides of sulphur and nitrogen are also acidic oxides and water soluble, so they cause
acid rain and thus cause plants to die and buildings to be eaten away.
The main source of these pollutant gases entering the air is from the burning of fuels. All oils and
fuels contain sulphur, and when they are burnt the sulphur reacts with the oxygen and forms
sulphur dioxide gas.
When fuels like petrol and diesel are burnt in an internal combustion engine, the amount of
oxygen present is limited. The carbon in these fuels reacts with the limited oxygen to form carbon
monoxide, which is poisonously fatal. Had there been an excess of the oxygen, carbon dioxide
would have been formed, which is a lot harmless compared to carbon monoxide. Also, the oxides
of nitrogen that are released from car exhausts cause major pollution problems.
Don’t worry though; you won’t have to stop using your cars in order to save the environment!
Engineers have already thought of this problem and figured out a way to reduce the pollution that
our cars produce. They have made catalytic converters that are fit to car exhausts. Inside the
converter is a special metal, mostly platinum, which acts as a catalyst. The function of this
converter is to turn the poisonous exhaust gases we discussed earlier into harmless gases. Carbon
monoxide is converted into carbon dioxide, while all oxides of nitrogen are converted into
nitrogen gas. It does this by transferring oxygen atoms from the oxides of nitrogen to the carbon
dioxide.
Ozone:
Ozone gas is actually three atoms of oxygen combined in form of a molecule. Even though it is
useful in the atmosphere high up, it is a pollutant and harmful at ground level. It is formed when
an electric spark passes through oxygen. Levels of ozone are high near railway tracks and
electrified wires. It is also produced in car engines.
Ozone is harmful to the plants, and an irritant for nose and throat for humans. It reacts with
ultraviolet radiation in sunlight to produce a photochemical smog, which is not fit for humans.
However, up in the atmosphere it is beneficial as it helps keep out ultraviolet radiation from the
sun.
Lead Compounds:
Lead compounds are usually added to petrol to make it heavier and thus more efficient. When this
petrol burns, the lead particles stay as they are. These particles are then emitted out of the vehicle,
and into the air. Such particles, when breathed in, can build up inside the body, and are toxic and
poisonous. So these lead particles are also pollutants.
Hydrogen:
We all know by now that hydrogen is the first element of the periodic table. It is the simplest of
all atoms. It just has one proton. It is the only atom that has no neutron.
The hydrogen atom has only one electron. It combines with another hydrogen atom to form a
complete duplet, and thus makes a molecule of hydrogen H2. Hydrogen in its original state exists
always as a diatomic molecule, i.e. it is always in the form of H2.
If we need to test for hydrogen, we take a lighted splint near the mouth of the suspected
hydrogen. If there is a loud ‘pop’ sound, it means that the gas is hydrogen.
Water:
Water is the most common of all liquids. Life is not possible without water. It is essential for
human and animal life, and also for plant growth.
Water is an extraordinary solvent, and dissolves many things. This has its advantages as well as
disadvantages. Beneficial dissolved substances in water are mineral salts and oxygen. However,
here are also some water pollutants that may get dissolved in this water. These include metal
compounds, sewage, nitrates from fertilizers, phosphates from detergents, and crude oil from
spillage. These water pollutants must be removed before consumption.
Due to these and other water pollutants, there is a need for water purification, as we cannot
consume dirty water. The place where dirty contaminated water is treated and purified is called
waterworks.
There are few main stages in the purification of water. Let us go through these.
In the first stage, the easily removable impurities are removed. Large impurities like leaves,
sticks, etc are removed through screening, while a settling agent like aluminium sulphate is used
for smaller particle impurities, so that they form small clumps and settle at the bottom, so that
screening them out is easier. The water containing these clumps is left for sometime in
sedimentation tanks so that the large solid particles may settle down. Next, the water is filtered
through layers of sand and gravel to remove any fine solid particles. Sometimes the water is
passed over beds of activated charcoal to remove tastes and odors.
After this, the water is chlorinated. The purpose of chlorine is to kill any harmful bacteria that
might be present. However, it must be added carefully and in a controlled quantity so that it
cannot be detected and is thus harmless if the water is consumed. Sometimes even fluoride salts
are added to the water to prevent tooth decay.
Finally, the purified water is stored in places from where it can be easily pumped or gravity-fed to
the taps of the consumer.
But this method is not practiced everywhere. So the water from seas and oceans is also used in
many countries. But as you know, this water is salty and cannot be used in that state; therefore it
must be desalinated first. To desalinate it, fractional distillation is used. We have already learnt
how that is carried out.
Uses for Water:
I’ll list them up for you.
● Essential for life
● Cooling in power stations and producing steam for generators.
● Cooking and cleaning in the home
● Manufacture of food and drinks
● Building and construction industry
That is pretty much all for this chapter. I thank you students for your concentration. Good day!
Chap 21: Organic Chemistry
Organic chemistry is the chemistry of the compounds of carbon. Actually, it is the study of
‘hydrocarbons’. Hydrocarbons are compounds containing hydrogen and carbon and no other
element. There are many types of hydrocarbons. They are classified into several types according
to their structure.
Homologous Series:
A homologous series is a family of organic compounds with similar chemical properties.
Examples of homologous series are the alkanes, the alkenes, the alcohols, and the carboxylic
acids. Compounds of the same homologous series contain the same functional group. A
functional group is an atom or a group of atoms that gives a molecule its characteristic properties.
Let us study the different examples of homologous series now.
Alkanes:
The members of this group of hydrocarbons are distinguished by possessing the general
molecular formula CnH2n+2, where n is a real number, for successive members of the group. For
example, the first member of the group is methane CH4, in which n=1. The second member is
ethane C2H6, in which n=2. And so on.
Note that in all the homologous series, whenever n=1, the compound’s name will start with methand the next part of the word will indicate the group it falls in. Like in methane, the ‘ane’ shows
that it belongs to the alkane group.
We will learn the names and formulas of the first ten members of the alkane group.
Methane CH4
Ethane C2H6
Propane C3H8
Butane C4H10
Pentane C5H12
Hexane C6H14
Heptane C7H16
Octane C8H18
Nonane C9H20
Decane C10H22
Structure:
Each carbon atom is any molecule can make four bonds. It may bond with less than four elements
by making double and triple bonds, but the bonds must equal four. In alkanes, each carbon bonds
with four hydrogen atoms. The molecular structures of the first three alkanes are:
Methane CH4
Ethane C2H6
Propane C3H8
And so on. I am sure you get the general idea.
The alkanes are called unsaturated compounds. It is because they cannot take part in addition
reactions, as they do not have any double bonds that can be easily broken. They can, however,
take part in substitution reactions under the right conditions, which means that any other element
or radical can take the place of any hydrogen in the compound to form a new compound.
Alkenes:
The alkenes are members of a homologous series of general molecular formula CnH2n. For
example, ethene is C2H4. Every carbon in an alkene makes bonds with four other elements, i.e.
one carbon and other hydrogens. However, two of the carbons in each alkene molecule are linked
by a double bond. It is shown in these molecular structures:
Ethene, C2H4
Propene, C3H6
You see, all carbons are still making four bonds, but there exists one double carbon to carbon
bond in each alkene molecule. This double bond is the ‘functional group’ of alkenes that I
mentioned earlier. This double bond exists in every alkene and is responsible for the chemical
properties of the alkenes.
Alkenes are said to be unsaturated because of the double bond in their structures. This means that
one molecule of it can combine with two hydrogen atoms (or their equivalent) in addition
reactions.
For example,
+ H2
Polymerization:
This will be treated in more detail in the next chapter about macromolecules. However,
unsaturated molecules can, under the right conditions, join together to form giant molecules
called polymers. All plastics and man-made fibers are polymers. Each polymer is made up of
thousands of identical units called monomers. For example, the polymer polyethene is made up of
a very large number (n) of ethene monomer molecules.
This is a monomer of ethene. You can see that the double bond is broken, leaving two free bonds
to bond with any other element. So they form chains and each monomer in the chain bonds with
the next monomer. It is then called a polymer.
This is the polymer. The zigzag lines show that this continues.
Alkynes:
The alkynes are those hydrocarbons that contain one triple bond. The triple bond is their
functional group and decides their properties.
Alcohols:
These are compounds containing not only the elements carbon and hydrogen, but also oxygen.
Like hydrocarbons, they form a homologous series, with the general formula CnH2n+1OH. The OH
group is its functional group, and determines its properties. Do not confuse these with hydroxides
of the alkalis, as alcohols (or alkonols) are neutral.
Alcohols can take part in the reactions combustion and oxidation. In the combustion reaction,
alcohols burn in air to produce carbon dioxide and water. If we oxidise alcohols, they become
alcohols. For example, if we oxidise ethanol, it becomes ethanoic acid.
The molecular structures of alcohols are like this:
Methanol CH3OH
Ethanol C2H5OH
Alcohols are used mainly in making alcoholic drinks, as solvents, and as fuels.
Carboxylic Acids:
The carboxylic acids are a homologous series of organic compounds with the carboxyl functional
group COOH. They have the general formula CnH2n+1COOH. Carboxylic acids are weak acids.
They are a class of organic acids.
The carboxylic acids are all soluble. The more their carbon content, the more is their boiling
point.
Let’s take a look at the molecular structure of ethanoic acid.
Ethanoic acid CH3COOH
Note that there are two carbons, but n=1. the other carbon is a part f the functional group of the
acid.
Esters:
Esters are formed when we react acids with carboxylic acids. For example, if we add ethanol and
ethanoic acid together, we get water and an ester, ethyl ethanoate, CH3COOC2H5. This reaction of
alcohols with carboxylic acid is called esterification. This is a reversible reaction, and we can
gain the acid and alcohol back by boiling the ester with sodium hydroxide. This is in fact a
reaction with water, and is called hydrolysis.
Esters have a sweet smell, so they are used in making perfumes. They are also used in flavorings,
as solvents, etc.
Organic chemistry is not finished, but we will end this lecture here so that it does not get too
complicated. Thank you for your attention students.
Chap 22: Macromolecules
Hello students. This is a tricky chapter so I suggest you all pay attention. In this chapter we study
about the macromolecules that the organic compounds we studied about in the previous chapter
form.
You will recall that we learnt a little about polymerization. It is the joining together of small
molecules, referred to as monomers, to form giant macromolecules called polymers. Let us learn
more about these polymers in this chapter.
Synthetic Polymers:
Synthetic or man-made polymers can be made by two types of polymerizations: addition
polymerization and condensation polymerization.
Addition polymerization is the linking together of unsaturated monomers to form a polymer,
without losing any atoms or molecules. Each polymer is only made of one kind of monomer
molecule. This type of polymerization is suitable with alkenes and alkynes.
In this type of polymerization, the double bond or the triple bond of the original molecule breaks
and forms the monomer that then combines with other monomers to form a long chain. The
conditions for this polymerization have to be very exact, and are different for every type of
polymer to be made. The temperature and pressure usually has to be very high, and a catalyst is
used.
An example of synthetic polymer made by addition polymerization is polyethene, also known as
plastic. Addition polymers are good insulators of heat and electricity, and are resistant to
chemical attack. For example, Perspex is a polymer made to use car windscreens, because it is
transparent and less easy to break compared to glass. Polyvinyl chloride (PVC), another polymer,
is used to make pipes, raincoats, thin gloves and flooring mats.
Condensation polymerization is the linking together of monomers with the elimination of a
simple molecule like water. Each polymer may contain two kinds of monomer molecule.
There are two main groups of condensation polymers: the polyamides and the polyesters.
Nylon is an example of a synthetic polyamide. It is made from the monomers dicarboxylic acid
and diamine.
Besides nylon, Terylene is another example of a condensation polymer. Terylene is made from
the monomers benzene-1, 4-dicarboxylic acid, and ethane-1,2-diol.
Clothes made from synthetic fibres are shrink-proof and crease proof. They are also easier to
wash and dry. Examples of items made from nylon and Terylene are curtains, parachutes, fishing
lines and sleeping bags.
Synthetic polymers, or plastics, are now used in place of natural materials such as wood, metal,
cotton and leather because they are relatively cheaper, easily molded into various shapes, light,
tough, waterproof, and resistant to decay and rusting. However, there are also many
disadvantages of these plastics in respect to the environment. When plastics burn, fires can spread
very quickly and poisonous gases are produced, causing air pollution. Plastics are nonbiodegradable; they cannot be decomposed by bacteria in the soil. As we cannot burn them either,
we must dump them somewhere, so they take up place and cause land pollution.
Natural Macromolecules:
The three main classes of food are proteins, fats and carbohydrates. All of these are natural
polymers or macromolecules, where small monomer units are linked to form a giant molecule.
Food which contains only carbon, hydrogen and oxygen are either carbohydrates or fats. Food
which contains carbon, hydrogen, oxygen and sulphur or nitrogen are proteins.
All natural macromolecules are biodegradable and are broken down during digestion using
biological catalysts called enzymes. These catalyse the reaction of protein, fats and carbohydrates
with water. This reaction is called hydrolysis.
Proteins: the monomers in proteins are linked by amide or peptide linkage. This is in fact the
same amide linkage that nylon possesses but proteins are formed from monomer units of amino
acids. To break proteins up, the enzyme pepsin is used in acidic medium. Proteins break up to
form amino acids.
Fats: they are linked by the same ester linkage that is in Terylene, but with different monomer
units. To break fats up, we boil them with sodium hydroxide, and the enzyme lipase is used. They
break up to form fatty acids and glycerol.
Carbohydrates: the monomers in carbohydrates are linked by oxygen linkage. To break these
up, we boil them with dilute hydrochloric acid, and the enzyme amylase is used. They break up to
form simple sugars, for example glucose, maltose, etc.
This is all for this chapter. Thank you students.