1. science
2. physics

CHM111 COURSE COMPACT
Course
Course code: CHM 111
Course title and Credit unit: GENERAL PHYSICAL CHEMISTRY (3 Unit)
Course status: Compulsory
Course Duration
Three hours per week for 15 weeks (45 hours)
Time of Lecture: Monday (1pm – 2pm) & Thursday (8am – 10am)
Lecture Location:
New College Building, LT1
Colleges Taking CHM 111: College of Science and Engineering (CSE) and College of Agricultural
Sciences (CAS)
Lecturers’ Data:
1. Name of the Lecturer: Dr Bello, O.S (Adjunct)
Qualifications obtained: BSc, MSc, Ph.D.
College: College of Science and Engineering
Department: Physical Sciences
Unit: Industrial Chemistry
E-mail: [email protected]
2. Name of the Lecturer: DADA, Oluwasogo Adewumi
Qualifications obtained: B.Sc (Hons), M.Sc (Chemistry), MCSN, MSAN
College: College of Science and Engineering
Department: Physical Sciences
Unit: Industrial Chemistry
E-mail: [email protected]
Office Location: 1st College Building, 1st Floor, A132.
Consultation Hours: Monday – Wednesday: 2 – 4pm,Thursday – Friday: 12 – 2pm
3. Name of Lecturer: Dr D.F. Latona,
Qualifications obtained: BSc (B.U.K), MSc (Unilorin), PhD (O.A.U)
College: College of Science and Engineering
Department: Physical Sciences
1
Unit: Industrial Chemistry
E-mail: [email protected]
Office Location: 1st College Building, 2nd Floor.
Consultation Hours: Monday – Wednesday: 2 – 4pm,Thursday – Friday: 12 – 2pm
Office Location: 1st College Building, 1st Floor, A206.
COURSE CONTENT
 Brief Overview of Course
Historical development of atoms, Fundamental of particles atoms and atomic structure.
Modern electronic theory of atoms; electronic configuration of elements, Periodicity of
Elements, Radioactivity, Stoichiometry, States of matter, Chemical Energetics, Chemical
Kinetics, Chemical Equilibrium, and Electrochemistry.
 Course Objectives/Goals
1. The physical chemistry course is often the first opportunity that a student has to synthesize
descriptive, theoretical, and mathematical knowledge about chemistry into a coherent
whole. To facilitate this, CHM 111 is constructed about the idea of Historical development
of atoms, Fundamental of particles atoms and atomic structure. Modern electronic theory
of atoms; electronic configuration of elements, Periodicity of Elements, Radioactivity,
Stoichiometry, States of matter, Chemical Energetics, Chemical Kinetics, Chemical
Equilibrium, and Electrochemistry.
2. It is to provide strong foundation in Physical Chemistry that would be a basic tool for their
relevance in the industries.
 Method of Lecture Delivery/Teaching Aids: Provision of detailed explanation on different
experiments that will be carried out for proper understanding of the subject matter. Provision
of electronic lecture notes with detailed examples and practice questions. Two classes will be
run which will be taken concurrently in room 106 and 402 each time the lecture is taken.
 Course Requirement:
In order to maximally participate, benefit and get the desired grade for this course, each student is
required to:
1. Regularly and actively attend the class and be sited latest 10 minutes to the commencement of
the lecture
2. Since one of the modes or methods of delivering lecture will be through the use of power point
presentation, students are required to come with their computer systems.
3. Constantly attempt the class assignment which will be one of the means to get the student
2
familiar with the pattern which test or examination could be tested.
4. Timely submit the term paper that will be given at the beginning of the course
5. Properly prepare for an individual presentation that would be given based on the discretion of
the course lecturer. The individual power point presentation, term paper, class test and
examination will be the basis for the grading of the student’s score.
Method of Grading
S/N
Grading
Score (%)
1.
Assignment
5
2.
Test
10
3.
Presentation
5
4.
Term Paper
10
5.
Final Examination
70
Total
100
 Course Delivery Strategies:
Core teaching methods via power point presentations will be adopted. The class will be more of
interactions and discussion of case study as an application practical tool of getting the student
familiar with the course.
 Tutorials: One hour tutorial per week as demanded.
 Topics for Assignment: Students will be given assignment based on topics covered.
 Alignment with Goals and Vision of Landmark University: To build a total man and intelligent
student who will break a new ground in the field of Agriculture, Science and Engineering.
 Contemporary Issues/ Industry Relevance:
Most of the topics covered under the course would provide strong foundation for the students
in areas of Agriculture, Science, and Engineering.
It will provide knowledge on how to handle stoichiometric task useful in quantitative / Quality
control unit in the industry
It will also afford the students the opportunity to develop their quantitative and analytical skills
in kinetics and thermodynamics
 Recommended Readings:
Physical Chemistry by Robert G. Mortimer
Physical Chemistry by K.K. Sharma Sharma
Physical Chemistry by Thomas Engel
3
 Course Outline:
Course coordinator: Mr DADA, O.A.
Module 1
 Week 1 – 3: Taken by DADA, A.O
 Description:
- Historical development of atoms, Fundamental of particles atoms and atomic structure.
Modern electronic theory of atoms; electronic configuration of elements by
- Periodicity of Elements and Radioactivity
 Study Question:
1. Which of the following quantum numbers gives information on types of orbital? (A).
Principal (B) Magnetic (C) Azimuthal (D) Spin
2. The wavelength of first spectral line in the Balmer series is 6561 Å units. Calculate the
wavelength of the second spectral line in Balmer series. (A) 4870 Å (B) 4860 Å (C) 2058 Å
(D) 6560 Å
3. What is the relationship between 32Ce76 and 34Se76?. They are (A) Isotopes (B) Isotones (C)
Isomers (D) Isobars
4. Which orbital is specified by n=3, l=2 (A) 3s (B) 3f (C) 3d (D) 3p
5. Which of the following scientist was the first to investigate the existence of Proton? (A) J.J.
Thompson, (B) Eugene Goldstein (c) Ernest Rutherford (d) Neil Bohr
6. In a typical experiment in Chemistry laboratory at LMU, 23.41g of ethanoic acid was
reacted with 40.00g of methanol to give 26.54g of methyl ethanoate. Calculate the
percentage yield of the ester.
[ C=12, O=16, H=1] (A) 26.54% (B) 91 % (C) 40% (D) 60%
7. In the historical development of atom, which of the scientist developed a mathematical
model for the distribution of electron in atoms? (a) Neil Bohr (b) Ernest Rutherford (c)
Erwin Schrodinger (d) Max Planck
Module 2
Week 4 – 6:
- Chemical Energetics by Dr Latona, D.F.
- Stoichiometry by DADA, A.O
Week 7: Mid-Semester examination.
Module 3
4
Week 8 – 10
- Chemical Kinetics and Chemical Equilibrium by Dr Bello, O.S.
- States of matter by DADA, A.O.
Module 4:
Week 11
- Electrochemistry by Dr Latona, D.F.
Week 12, 13 and 14: Thorough Revision.
 Objectives:
1. To fully equip the students for the examination
2. To get the student familiar with the course
3. To examine different technical questions.
4. To build more confidence in the mind
Week 15
Topic: Examination
Objectives:
To examine the students on all that has been taught during the semester.
Lecture Note
DADA OLUWASOGO’S ASPECT
CHM 111(3) GENERAL PHYSICAL CHEMISTRY
HISTORICAL DEVELOPMENT OF ATOMS
Postulates of Dalton's atomic theory
John Dalton, a British school teacher, published his theory about atoms in the year 1808. His findings
were based on experiments and also from laws of chemical combination.
Main assumptions or postulates of Dalton






All matter consists of indivisible particles called atoms.
Atoms of the same element are similar in shape and mass, but differ from the atoms of other
elements.
Atoms cannot be created or destroyed.
Atoms of different elements may combine with each other in a fixed, simple, whole number
ratio to form compound atoms.
Atoms of same element can combine in more than one ratio to form two or more compounds.
Atoms are the smallest unit of matter that can take part in a chemical reaction.
5
Drawbacks of Dalton's atomic theory of matter
 The indivisibility of an atom was proved wrong, for, an atom can be further subdivided into
protons, neutrons and electrons. However an atom is the smallest particle, which takes part in
chemical reactions.
 According to Dalton, the atoms of same element are similar in all respects. This is wrong
because atoms of some elements vary in their mass and density. Such atoms of the same
element having different masses are called isotopes. For example, chlorine has two isotopes
having mass numbers 35 a.m.u and 37 a.m.u.
 Dalton also said atoms of different elements are different in all respects. This has been proved
wrong in certain cases like argon and calcium atoms, which have the same atomic mass of 40.
Such atoms of different elements that have the same atomic mass are called isobar.
 According to Dalton atoms of different elements combine in simple whole number ratio to form
compounds. This is not seen in complex organic compounds like sugar C12H22O11.
 The theory completely fails to explain the existence of allotropes. The difference in properties
of charcoal, graphite, diamond went unexplained in spite of being made up of same kind of
atoms.
Merits of Dalton's atomic theory


It has enabled us to explain the laws of chemical combination.
Dalton was the first person to recognize a workable distinction between the ultimate particle of
an element (atom) and that of a compound (molecule).
Discovery of electron
PRODUCTION OF CATHODE RAYS
In 1885, Sir William Crookes carried out a series of investigations into the behaviour of metals heated
in a vacuum. The experiment of Crookes and others showed that a heated cathode produced a stream
of radiation, which could cause gases at low pressure to glow and which, make other substances emit
light too. The radiation emitted from the cathode was given the name 'Cathode rays'. By mid-nineties it
was known that these rays could be deflected by a magnetic field and they carried a negative charge.
Some scientists felt that these rays were waves and others were inclined to think they were particles.
Perrin tube
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In 1897, J.J Thomson showed that the stream of particles were indeed electrons. He conducted the
famous discharge tube experiment by passing electricity at high voltage through a gas at low pressure.
A common discharge tube is a long glass tube having two metal plates, sealed at its two ends as
electrodes. It has a side tube through which air can be pumped out by using a vacuum pump, so that
experiments can be performed at low pressure.
Production of Cathode rays
When the pressure of air in the discharge tube is reduced to .001 mm of mercury and a high voltage is
applied to the electrodes, the emission of light by air stops. But the phenomenon of fluorescence is
observed in which the walls of the discharge tube at the end opposite to the cathode begin to glow
with a greenish light. It is now deduced that some invisible rays were formed at the cathode, which on
striking the glass tube emitted a green light. Since they are formed at the cathode they are known as
cathode rays.
PROPERTIES (EXPERIMENTAL OBSERVATIONS)
1. Travel in straight lines
When an opaque object like a metal cross is placed in the path of cathode rays in a discharge tube, a
shadow of the metal cross is formed at the end opposite to the cathode.
Cathode rays cast shadows of the objects placed in their path
2. Produce mechanical effects
On placing a light paddle in the path of cathode rays in a discharge tube the blades of the paddle wheel
rotate. This shows that cathode rays are a beam of particles having mass and possessing kinetic energy.
7
Cathode rays can rotate a high paddle wheel placed in their path
3. Are negatively charged
When an electric field is applied in the path of cathode rays, they are deflected towards the positive
plate of the electric field, which shows cathode rays are made up of negatively charged particles.
4. Effect of electric field
Effect of electric field on cathode rays

The nature of cathode rays does not depend on the nature of gas taken in the discharge tube or
material of the cathode.

The ratio of the charge to mass (e/m ratio) of cathode ray particles obtained from different gases
was found to be exactly the same.
Conclusion
Since all gases form cathode rays, it means that all atoms contain electrons. The cathode ray particles
being negatively charged an electron is negatively charged.
Characteristics of an electron
1. The mass of an electron is 1/1840 of the mass of a hydrogen atom. Since the mass of hydrogen
atom is 1 amu., the relative mass of an electron is 1/1840 amu. The absolute mass is 9 x 10-28
gram.
2. Charge of an electron
An electron is found to carry 1.6 x 10-19 coulomb of negative charge. Since this is the smallest negative
charge carried by any particle, it is taken as unit negative charge.
Discovery of proton
PRODUCTION OF ANODE RAYS
8
Since the atom is electrically neutral there must be positively charged particles present in the atom to
neutralize the negative charges of the electrons. Goldstein experimentally proved the existence of
protons in the atom.
Experiment
Production of anode rays
In this a discharge tube having a perforated cathode is used. When a high voltage of about 10,000 volts
is applied to a discharge tube having a perforated cathode and containing air at very low pressure of
about .001 mm of mercury, a faint red glow is observed behind the cathode. These rays are formed at
the anode and when these rays strike the walls of the discharge tube they produce a faint red light.
Since they are formed at the anode (positive electrode) they are known as anode rays or positive rays.
Properties of anode rays

Travel in straight lines: They cast a shadow of the objects placed in their way.

Produce mechanical effect: A paddle wheel placed in their path starts rotating.

Rays are positively charged: Anode rays are deflected towards the negative plate of an electric field.
The nature of the anode rays depends upon the gas taken in the discharge tube. Different gases
give different types of positive rays, which contain particles having different masses and different
charges. Therefore the e/m ratio is not constant for positive ray particles obtained from different
gases.
In the case of hydrogen the e/m ratio is the highest as the positive particles obtained from hydrogen
are the lightest. The positive particles obtained from hydrogen gas are called 'protons'. It comes from
the Greek word 'Proteios' meaning 'of first importance'.

Formation of positive rays
When high electrical voltage is applied to a gas, its atoms break up into negatively charged particles
(electrons) and positively charged particles. These positively charged particles formed by the removal
of electrons from the gas atoms are called positive rays.
CHARACTERISTICS OF A PROTON
1. Mass
A proton is actually a hydrogen atom, which has lost its electron. Since the mass of an electron is small,
the mass of a proton is equal to the mass of a hydrogen atom. As the mass of hydrogen atom is 1
9
a.m.u., the relative mass of a proton is 1 a.m.u. The mass of a proton is 1840 times that of an electron.
The absolute mass of a proton is 1.6 x 10-24 gram.
2. Charge
The proton is equal and opposite to the charge of an electron. So the absolute charge of a proton is 1.6
x 10-19 coulomb of positive charge. This being the smallest positive charge carried by any particle, it is
taken as 1 unit positive charge. The relative charge of a proton is +1 (plus one).
Effect of low pressure in the discharge tube
When the gas atoms in the discharge tube are at atmospheric pressure they collide with the electrons
preventing them from reaching the anode. As no electrons reach the anode no current flows through
the discharge tube. When the gas pressure is very low there are few gas atoms in the discharge tube.
As such there is no hindrance to the movement of electrons and the gas conducts electricity.
NB:
Thompson from his discharge tube experiment concluded that atoms consisted of positively and
negatively charged particles. The neutrality of atoms is due to equal numbers of negative and
positive charged ions.
Failure of J.J Thompson’s Model: On the basis of his model, he was able to explain
1. Photoelectric effects
2. Kinetics Theory of gases
3. Emission of Electromagnetic waves
Failure of J.J Thompson’s Model
1. He failed to explain the scattering of α particles
2. He failed to explain emission spectra of the elements
Discovery of neutron
Although Rutherford's model was enormously successful at explaining the scattering of alpha particles,
there was a problem when the experimental data were used to calculate the mass of the nucleus.
From the results of the scattering experiments, it was possible to calculate the charge on the
nucleus. However the mass of these protons was only about half the overall mass of the nucleus. In
1920, William Draper Harkins an American physicist suggested that the missing mass could be
accounted for if the nucleus contained other particles with mass similar to that of a proton but no
charge. He named this particle neutron. James Chadwick finally discovered it in 1932. He bombarded
the element beryllium with alpha-particles. He observed the emission of a radiation with the
following properties:

The radiation was highly penetrating

The radiation remained unaffected in an electric or magnetic field i.e., the
radiation was neutral

The particles constituting the radiation had the same mass as the proton. Thus the relative mass of
such a particle = 1 amu and the absolute mass = 1.6 x 10-24g. Because of their electrical neutrality,
these particles were called neutrons
We are now in a position of explain why the atomic mass of carbon is 12. It is now known that the
carbon atom contains 6 protons and 6 neutrons each having a mass of 1 amu.
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Atomic mass of carbon = Mass of 6 protons + Mass of 6 neutrons
6 x 1 + 6 x 1 = 12
Concept of atomic number, mass number, fractional atomic mass, isotopes, isobars
The nuclei of atoms is made up of protons and neutrons. These two components of the nucleus are
referred to as nucleons. The electrons occupy the space outside the nucleus. Since an atom is
electrically neutral, the number of protons in the nucleus is exactly equal to the number of electrons.
This number is the atomic number given by the symbol Z.
Atomic number represents the number of protons in an atom. As atoms are electrically neutral, an
atom contains as many electrons as it has protons.
ATOMIC MASS, MASS NUMBER OR NUCLEON NUMBER
The total number of protons and neutrons present in one atom of an element is known as its mass
number.
Mass number = number of protons + number of neutrons
It can also be said that:
Mass number = Atomic number + number of neutrons
Atomic number and mass number can also be represented on the symbol of an element as follows:
Let's take an example:
Represent the atom of sodium whose atomic mass is 23 and atomic number is 11. Calculate the
number of protons, electrons and neutrons.
23
Na11
Atomic number Z =11
Atomic mass A = 23
No. of protons = Z = 11
No. of electrons = 11
No. of neutrons = A - Z
23 - 11 = 12

Study the table carefully
Lithium - 7 Silicon - 28 Copper - 65 Dysprosium - 164 Uranium - 238
Proton Number Z
3
14
29
66
92
Nucleon number A
7
28
65
164
238
Number of neutrons N (A-Z)
4
14
36
98
146
Symbol
7 Li 3
28 Si 14
65 Cu 29
164 Dy 66
238 U 92
Attempt the following:

The atomic number of an element is 12. How many protons and electrons are there in the atom?
Solution:
11
Atomic number = 12
Number of protons = Number of electrons = Atomic number = 12

The nucleus of an atom of an element contains 11 protons and 12 neutrons. Determine the atomic
number and mass number of the element.
Solution:
Atomic number = number of protons = 11
Mass number = number of protons + number of neutrons
= 11 + 12 = 23
Fractional atomic mass
It is interesting to note that atoms of a given atomic number can have different number of neutrons.
Some examples are listed below:
Hydrogen
Hydrogen atom (Z=1) has no neutrons.
Number of protons = 1
Number of electrons = 1
Number of neutrons = 0
It has been reported that the hydrogen element has atoms with mass number 2 and 3 also i.e.,
Atoms of elements having the same atomic number with different mass numbers are called isotopes.
12
Protium Deuterium Tritium
At.mass = 1 At.mass = 2 At.mass = 3
Nuclear composition of isotopes of chlorine:
Nuclear composition of isotopes of carbon:
Characteristics of isotopes
13

All isotopes of an element have the same number of valence electrons thus have identical chemical
properties.

The physical properties of the isotopes are different due to the difference in the number of
neutrons in their nuclei. The densities, melting points and boiling points etc., are slightly different.
REASON FOR FRACTIONAL ATOMIC MASSES OF ELEMENTS
Atomic masses of many elements are in fractions not in whole numbers.
Example
Cl - 35.5
Cu = 63.5
The fractional atomic masses of elements are due to the existence of isotopes having different masses.
Example
Natural chlorine consists of two isotopes:
Calculate the average atomic mass of chlorine.
TRY IT OUT!

A naturally occurring sample of Lithium contains 7.42% of
of
is 6.015 and that of
lithium.
and 92.58% of
. The relative mass
is 7.016. Calculate the atomic mass of a naturally occurring sample of
Solution

Which two of the following nuclei are isotopes of each other?
The two isotopes are:
ISOTOPES
These are the elements having same atomic number but different mass number. They have the same
atomic number because the number of protons inside their nuclei remains the same. The difference in
their mass number is due to the difference in their number of neutrons.
Since they are neutral isotopes are elements having same number of electrons, which make them to
possess identical chemical properties. Let us see some examples 1H1, 1H2, 1H3 are all isotopes of
hydrogen. They all have their atomic number to be unity but the number of neutrons are 0, 1, 2 and z
respectively. 17Cl37, 17Cl35 are isotopes of chlorine. They have 17 protons in the nucleus but have
number of neutrons equal to 20 and 18 respectively. Practically every element consists of a mixture of
several isotopes. The relative abundance of different isotopes differs from element to element. For
example chlorine is composed of two isotopes of masses 34.98U and 36.98U, which are nearly integral
multiples of the mass of hydrogen atom. Their relative abundances are 75.4 and 24.6 percent
respectively. Mass of natural chlorine atom can be found as
14
= 35.47.
The isotope can occur either naturally or can be produced artificially in the laboratory.
ISOBARS
Isotopes are chemically same and physically different. But the converse is true in isobars. That is
isobars are elements, which are chemically different but physically same. So, isobars are atoms of
different elements having the same atomic mass but different atomic number. Since their number of
electrons is different, their chemical properties are different. The light nuclei have unstable isobars.
Heavy nuclei have stable isobars and these occur in pairs. Suppose the number of protons of one
isobar matches with that of another they are called as mirror-nuclides of each other.
Examples of isobars are
Since isobars are different elements they appear in different places in the periodic table.
ISOTONES
Isotones are elements having the same number of neutrons. Examples of isotones are Chlorine - 37
and Potassium - 39. Both have 20 neutrons in their nuclei
Rutherford's α ray scattering experiment and nuclear model of atom; limitation
After the discovery of the radioactive particles, Rutherford performed an experiment where he
bombarded a thin sheet of gold with a particles obtained from a radioactive substance. On striking the
gold foil some of the a particles scattered and produced flashes on a zinc sulphide (ZnS) screen placed
at the back of the gold foil. These tiny flashes were observed by a movable microscope. The
observations made in this scattering experiment were as follows:
15

Most of the a particles pass through the metal and are undeflected.

Some of the a particles are deflected through small angles.
Only a very few of them are deflected through as much as 90o or even larger angles.
The results of the scattering experiment could not explain the Thomson's atomic model. Hence,
Rutherford
concluded that:


As most of the a particles passed undeflected, most of the space inside the atom is empty or
hollow.

Some that deflected with large angles show that heavy positively charged body must be present
inside the body of the atom, which repelled the like charge of the a particle. This heavy, positively
charged body was named as nucleus.

The number of heavy positively charged particles that undergo deflection is very small. The volume
occupied by the nucleus is very small compared to the volume of the atom.

When heavy particles like the a get deflected, the nucleus of the atom must also have an
appreciable mass.
Rutherford's nuclear model
On the basis of the scattering experiment, Rutherford described the structure of the atom as:

An atom consists of a positively charged nucleus surrounded by electrons that move around it. The
positive charge of the nucleus is due to the protons.

Electrons and neutrons are held together by coulombic force of attraction.

The effective volume of the nucleus is extremely small as compared to the effective volume of the
atom. The volume occupied by the nucleus is about 10-12 times the volume of the atom.

The entire mass of the atom is concentrated at the nucleus.

Since each atom is electrically neutral, the number of positive charges in the nucleus of an atom is
equal to the number of electrons in it.
Fig: 3.12 - Rutherford's model of an atom
Limitations
1.
Rutherford's model suffered some drawback, as it could not explain the stability of the atom inspite of the revolving electrons around the nucleus.
2.
Bohr's model of atom and explanation of hydrogen spectra
In 1913, Neils Bohr proposed a model of an atom based on the Planck's quantum theory of radiation.
The basic postulates of Bohr's theory are:

An atom consists of a small, heavily positively charged nucleus around which electrons revolve in
definite circular paths called orbits.
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
These orbits are associated with definite energies called energy shells/energy levels. They are
designated as K, L, M, N, …. etc. shells or numbered as 1, 2, 3, 4, …..etc. from the nucleus.

As long as the electron remains in a particular orbit /energy shell its energy remains constant. This
accounts for the stability of an atom.

Only those orbits are permitted in which angular momentum of the electron is a whole number
multiple of
where h is Plancks constant. Any moving body taking a circular orbit has an angular
momentum equal to the product of its mass (m), velocity of movement (v) and radius of orbit (r). In
other words the angular momentum of an electron
Thus,
This postulate introduces the concept of quantization of angular momentum.
Electrons can either lose or absorb energy abruptly, when they jump from one energy level to
another. For instance when an electron moves from the 'normal or ground state - E1' of an atom
i.e., the state of lowest energy as required by its 'n' and 'l' values, to a higher level, it causes the
atom to be in its 'excited state - E2' i.e., where electrons in an atom occupy energy levels higher
than those permitted by its 'n' and 'l' values. The reverse is also true and the change in energy isDE,
DE = E2 - E1 = hn

Fig: 3.13 - Energy changes in an electron jump
Bohr's atomic model explained successfully:

The stability of an atom. Bohr postulated that as long an electron remains in a particular orbit it
does not emit radiation i.e. lose energy. Hence it does not become unstable.

The atomic spectrum of hydrogen was explained due to the concept of definite energy levels. The
one electron of hydrogen being closest to the nucleus is in its lowest energy shell (n =1) or normal
ground state. It can absorb a definite amount of energy and jump to a higher energy state. This
excited state being unstable, the electron comes back to a lower energy level.
17
When the energy emitted during transition, strikes a photographic plate, it gives its impression in the
form of a line. This difference is also the energy of photon expressed as E2 - E1 = hn.
The frequency of the emitted radiation is:
Since E2 and E1 have only definite values and are characteristic of energy levels of atoms, the values of
'n' will also be definite and characteristic of the atoms. Thus each transition will produce a light of
definite wavelength, which is observed as a line in the spectrum.
For example, if the electron jumps down from the third to the first energy level having energies E3 and
E1 respectively, then the wavelength of the spectral line would be
Similarly, when the electron jumps down from the fourth to the first energy level having energies E4
and E1 respectively or from the fifth to the second i.e., E5 and E2, then we have
These will give different lines in the spectrum of the atom corresponding to different transitions having
definite wavelengths.

The sample of hydrogen gas contains a large number of atoms and when energy is supplied, the
electrons in different hydrogen atoms absorb different amounts of energies. These are raised to
different energy states. For example, the electrons in some atoms may jump to second energy level
(L), while in others it may be to the third (M), fourth (N) or fifth (O) and so on. These electrons
come back from the higher energy levels to the ground state in one or more jumps emitting
different amount of energies.
Fig: 3.14 - Different routes to the ground state from n = 4
Different lines depending upon the difference in energies of the levels concerned can be summarized
in the form of series named after the scientists who have discovered them.
Lyman series from n = 2, 3, 4, 5……to n = 1
Balmer series from n = 3, 4, 5, 6……to n = 2
Paschen series from n = 4, 5, 6, 7……to n = 3
Brackett series from n = 5, 6, 7, 8……to n = 4
Pfund series from n = 6, 7, 8, 9……to n = 5.

The energy of the electron in a particular orbit of hydrogen atom could be calculated by Bohr's
theory. The energy of the electron in the 'nth' orbit has been found to be
18
where 'm' is the mass and 'e' is the charge of the electron. The energy expression for hydrogen like ions
such as He, Li can be written as:
where 'Z' is the nuclear charge, which is equal to atomic number.
Problems
7. If the energy difference between the electronic states of hydrogen atom is 214.68 kJ mol-1, what will
be the frequency of light emitted when the electron jumps from the higher to the lower energy state?
(Planck's constant = 39.79 x 10-14 kJ mol-1)
Solution
The frequency (n) of emitted light is related to the energy difference of two levels (DE) as
E = 214.68 kJ mol-1, h = 39.79 x 10-14 kJ mol-1
= 5.39 x 1014 s-1
8. The wavelength of first spectral line in the Balmer series is 6561 Å units. Calculate the wavelength of
the second spectral line in Balmer series.
Solution
According to Rydberg equation:
For the first line in Balmer series, n1 = 2, n2 = 3
For the second line in Balmer series, n1 = 2, n2 = 4
Dividing equations (i) by (ii)
A spectrum is an assembly of energy levels in the form of radiations emitted by an atom in its excited
state. Every atom gives discontinuous line spectra. Each line in the spectra corresponds to a specific
wavelength and it is unique to a given element. No two elements give same pattern of lines in their
spectra.
Atomic emission spectra
When a substance is heated to a high temperature, the atoms in the vapours get energized. These
energized atoms then return to the ground state by emitting electromagnetic radiations of certain
definite wavelength. A series of bright lines separated from each other by dark spaces is obtained and
this is called atomic emission spectra.
19
Fig: 3.9 - Mechanism of emission spectra
Atomic absorption spectra
When the atomic vapours from a sample are placed in the path of white light from an arc lamp, it
absorbs the light of certain characteristic wavelengths and the light of other wavelengths get
transmitted. In such conditions a series of dark lines on a white background are formed. This is called
an absorption spectrum.
The dark lines in the absorption spectrum and the bright lines in the emission spectrum of a given
element appear at the same wavelength.
Fig: 3.10 - Comparison of Absorption and emission spectra of sodium vapour
Since each element gives a definite pattern of lines at certain definite frequencies or wavelengths, the
atomic spectra is used in chemical analysis to identify and estimate the elements present in any
sample.
Atomic spectra of hydrogen atom
Hydrogen is the simplest element with its atom having only one electron. Hence, the atomic spectrum
of hydrogen has played a significant role in the development of atomic structure. In the emission
spectrum of hydrogen, when an electric discharge is passed through hydrogen gas, the molecules of
hydrogen break into atoms. The hydrogen atoms get energized and go into an excited state. The
excited atoms then return to the ground state by emitting light. Hydrogen atoms emit bluish light. On
passing this light through a prism, a discontinuous line spectrum consisting of several sharp lines is
obtained. This is the line spectrum of hydrogen.
Four sharp coloured lines were observed in the visible region of this spectrum by Balmer, in the ultra
violet region by Lyman, in the infrared region by Paschen, Brackett and Pfund. These series of lines are
named after these scientists who discovered them. Balmer expressed these lines in terms of inverse of
their wavelength ( ) by a mathematical relation, which was later modified by Rydberg.
where 'RH' is the Rydberg's constant and 'n1', 'n2' are integers with values equal to or greater than 3
and 'l' is the wavelength.
20
Fig: 3.11 - Line spectrum of hydrogen atom
Problem
6. Calculate the wavelength of the spectral line when the electron in the hydrogen atom undergoes a
transition from 4th energy level to 2nd energy level. What is the colour of the radiation?
Solution
According to Rydberg's equation
here, n1 = 2, n2 = 4 and R = 109678 cm-1
l = 486 x 10-9 m = 486 nm
Limitation of Bohr's model of atom

It could not explain the line spectrum of multi electron atoms.

This model failed to explain the effect of magnetic field on the spectra of atoms (Zeeman effect).

The effect of electric field on the spectra could not be explained by Bohr's model (Stark effect).

The shapes of molecules arising out of directional bonding could not be explained.

The dual nature of electrons (both as wave and particle) and the path of motion of the electron in
well defined orbits were not correct.


Elementary idea of quantum mechanical model
In 1924, de Broglie's suggested that all material objects including an electron have a dual
character; they behave as particles as well as waves. The wavelength associated with a particle
of mass 'm', moving with velocity 'v' is given by de Broglie's relation as:

21

The discovery of the wave like character of the electron helped in the making of the modern
electron microscope.
Dual nature of electron (de-Broglie equation)
De-Broglie noted that according to the theory of relativity, the role played by momentum 'P' and the
energy 'E' of the particle is done by angular frequency 'w' and the propagation vector 'k' of a wave.
According to Einstein relation
E = mc2
Where 'm' is the mass of the particle and 'c' = is the velocity of light and according to quantum theory.
where υ is the frequency
E=hυ
hυ = mc2
(Since c = ul)
(Since mc = p)
This relation ship tells us that if there is a particle of mass 'm' moving with a velocity if manifests itself
in the form of a wave, its wavelength would be h/p. This is de Broglie wave equation.
De-Broglie's hypothesis attributing a dual particle-wave character to matter appears very strange at
first sight. Since it is very much against the direct evidence of one's senses at the macroscopic level.
But at the level of the atom, the behavior of matter has been already found to be unconventional. It
was such evidence of non-classical behavior, which gave scope for De Broglie's proposal. The precise
form of the hypothesis was determined however by the speculative postulate that at the most
fundamental level, matter and radiation which form the basic constituents of the physical should be
similar in nature. Confirmation that nature does exhibit such an aesthetically pleasing symmetry
between matter and radiation came from the experiments by Davison, Germer, and G.P. Thompson.
These experiments demonstrated the existence of 'De Broglie Waves' associated with electrons in a
very direct fashion.
Heisenberg's uncertainty principle
Heisenberg, in 1927 pointed out that it is not possible to measure simultaneously both the momentum
(or velocity) and the position of a microscopic particle with absolute accuracy. Mathematically this may
be expressed
22
Dp = uncertainity in momentum
The constant on the right side of the equation (the product of the two uncertainties) tells us that the
two uncertainties are inversely related. If the momentum of the particle is measured with more
accuracy there will be a large uncertainity in its position and vice versa.
Uncertainity is not due to the lack of refined techniques available, but because we cannot observe
microscopic bodies without disturbing them. [Observations made as result of the impact of light
suffer a change in the position or velocity of these microscopic objects]. This does not hold good for
large objects of daily light, as the changes that o
Probability concept
According to Heisenberg's uncertainty principle, it is impossible to describe the exact position of an
electron at a given moment in terms of position, we can speak of most probable regions where the
probability of finding an electron in the space around the nucleus of an atom is high. The electron does
not always remain at a fixed distance from a nucleus. It keeps moving in the whole space around the
nucleus but tends to remain most of the time within a small volume around the nucleus, where the
probability of locating the electron is maximum.
A new atomic model, was needed to explain

Wave nature (dual character) of atoms.

The idea of uncertainity in the position of electrons in a atom.
Concept of fixed energy states.
Schrodinger put the wave model or quantum mechanical model of atom forward. The behavior of an
electron is defined by the mathematical representation:

where,
ψ = (psi) is a wave function of space coordinates 'x', 'y', 'z' and represents the amplitude of the electron
wave.
m = mass of the electron
E = the total permissible energy level, which the electron can have.
V = potential energy of the electron given by ze2/r.
h = Planck's constant having the value 6.626 x 10-34 J s.
d= (delta)stands for infinitesimal change.
The wave length function y (psi) describes a number of possible states of an electron in an atom. Since
a large number of solutions are possible, four quantum numbers were introduced, which describe
meaningful permissible values of energy and location with respect to its nucleus
Shape of atomic orbital (s and p orbitals only)
Shape of s orbital
s orbitals are non-directional and spherically symmetrical, This means that the probability of finding
the electron is same in all directions at a particular distance from the nucleus, The 1s orbital is shown
in the figure 1.3.
23
fig 1.3 - Shapes of 1s and 2s orbitals
It is observed that density of charge cloud is maximum at the nucleus and decreases with increase in
distance from the nucleus. The 2s orbital is also non-directional and spherically symmetrical. In this
case, the density is maximum at the nucleus and becomes small at large distances. However, the
effective volume or size of 2s is larger than 1s orbital. An important feature of 2s orbital is that there is
a spherical shell within 2s orbital (region without dots) where the probability of finding the electron is
practically zero, This is called a node or a nodal surface. Thus, a 2s orbital differs from 1s orbital in
being larger in size and having a nodal surface. The higher s orbitals have also spherical shapes and
number of nodal surfaces in s orbital for any given energy level is n - 1, where n represents energy
level.
Shape of p orbital:
For p-orbitals (l=1), there are three possible orientations corresponding to m = -1, 0, +1 values. This
means that there are three p - orbitals in each p-subshell. These are designated as px, py and pz; For
e.g., 2px, 2py and 2pz.
fig 1.4 - (a) Shapes of three 2p orbitals
24
These three orbitals are equal in energy but differ in their orientations. Each orbital consists of two
lobes symmetrical about a particular axis. Depending upon the orientation of the lobes, these are
designated as 2px, 2py and 2pz, as they are symmetrical about x, y and z-axes respectively. That is,
2pxorbital has two lobes symmetrical around x-axis and 2py orbital has two lobes symmetrical around
y-axis while the lobes of 2pz orbital are symmetrical around z-axis. The shape of the orbital is called
dumb bell shape. The two lobes of each orbital are separated by a plane having zero electron density.
This plane is known as nodal plane. For 2pz orbital, the nodal plane (in XY plane) is shown in below
figure.
fig 1.4 - (b) Nodal plane in 2pz orbital
2px and 2py orbitals have also similar nodal planes. It should be noted that the probability of finding the
electron in a particular p orbital is equal in both the lobes. The p orbitals of higher energy levels (n = 3,
4, 5… etc.) have similar shapes although their sizes are bigger.
NB: be familiar with the shape of d – orbitals
Quantum numbers
Orbitals of electrons in atoms differ in size, shape and orientation. Definite energies and angular
movements characterize atomic orbitals. The state of an electron in any atom is defined by certain
permissible values of energy and angular momentum, which describe its location with respect to its
nucleus and its energy level. These permissible states are called orbitals and are expressed by a set of
four numbers 'n', 'l', 'm' and 's' called quantum numbers. These numbers serve as the signature of the
electrons, uniquely describing its position in the atom. The 'n', 'l' and 'm' indicate the spatial
distribution while 's' indicates the spin orientation of the electrons.
1. Principal quantum number
This quantum number determines the main energy shell or energy level in which the electron is
present. The principal quantum number gives the average distance of the electron from the nucleus
and energy associated with it.
It is denoted by the letter 'n' that can take whole number values starting from 1, 2, 3, 4, ….. . The shell
with n = 1 is called first shell or 'K' shell. The shell with n = 2 is the 'L' shell and so on. The first shell is
closest to the nucleus. As the value of 'n' increases, the distance from the nucleus as well as the energy
of the electrons increases.
2. Azimuthal quantum number or angular quantum number
The Azimuthal quantum number determines the angular momentum of the electron, denoted by the
letter 'l'. The value of 'l' gives the sub level or sub shell in a given principal energy shell to which the
electron belongs. It can have only positive integral values from zero to (n-1) where 'n' is the principal
quantum number. The various sub shell values of l are also designated by the letters s, p, d, f,…… For
any main energy level, the energies of the sub shell follow the order s > p > d > f.
25
The different sub shells are represented by first writing the value of 'n' and then the letter designated
for the value of 'l'.
To illustrate,
n = 1 l = 0 one sub shell 1s
n = 2 l = 0,1 two sub shells 2s, 2p
n = 3 l = 0,1,2 three sub shells 3s, 3p, 3d
n =4 l = 0,1,2,3 four sub shells 4s, 4p, 4d, 4f
Thus for each value of 'n' there are 'n' values of 'l'.
The value of azimuthal quantum number gives the shape of the sub shell or orbital. So it is also called
as orbital quantum number.
3. Magnetic quantum number
The magnetic quantum number describes the behaviour of electron in a magnetic field. In the absence
of external magnetic field electrons/orbitals having same values of 'n' and 'l' but different values on 'm'
have the same energies. They are called degenerate orbitals. However, in the presence of an external
magnetic field the orbitals vary in their energies slightly. This is because the preferred orientation of
the orbital in space is a result of interaction of its own magnetic field with that of the external magnetic
field.
It is denoted by the letter 'm' the values of which depends on 'l'. This quantum number can have all
integral values from '-l' to '+l' including 0. Thus for given 'l' value there are (2l + 1) values of 'm'. Two
orbitals in the same shell can have identical 'n' and 'l' values but they must have different fixed values
of 'm'.
The number of orbitals in each sub shell are given below:
s sub shell l = 0 m = 0 only one orientation one orbital
p sub shell l = 1 m = +1,0, -1 three orientations three orbitals
d sub shell l = 2 m = +2,+1,0,-1,-2 five orientations five orbitals
4. Spin quantum numbers
The orientation of spin of an electron is designated by its spin quantum number 's'. The spin
orientation is an intrinsic characteristic of the electron connected more with its magnetic behaviour
rather than rotation of an electron about its own axis. This number can have only two values
corresponding to clockwise and anticlockwise spins i.e., +½ and -½. The clockwise spin is represented
by an arrow (h) pointing upwards. The anti clockwise spin is represented by an arrow (i) pointing
downwards. Each orbital can accommodate a maximum of two electrons provided they have opposite
spins.
26
Fig: 3.15 - Number of subshells and orbitals in the K,L and M shells
Question/Problem
1.(a) What are the permissible values for l and m when n = 3? (b) Which orbital is
specified by l = 2 and n = 3?
Solution
(a) For n = 3, the permissible values for 'l' and 'm' are: l = 0, 1, 2
For 'l' = 0 m = 0 (s-orbital)
For 'l' = 1 m = +1, 0, -1 (p-orbital)
For 'l' = 2 m = +2, +1, 0, -1, -2 (d-orbital)
(b) For 'n' = 3 and 'l' = 2:
'l' = 2 means 'd' orbitals
The given orbital is '3d'.
PAULI'S EXCLUSTION PRINCIPLE
This principle was proposed by Pauli in 1952. It states that no two electrons in an atom can have same
values for all the four quantum numbers. Thus, in the same atom, two electrons may have the same
values for three quantum numbers but the fourth must be different. Electrons having the same value
of n, l and m are said to belong to the same orbital. For instance, consider K shell, i.e. n =1. The
electron will have only one value of (l) which is l = 0 and one value of m, which is m = 0 but it can have
two values of s, either + ½ or - ½. This means that although n, l and m are the same for the two
electrons but their spin quantum numbers are different. Thus, an orbital can have maximum of two
electrons. Moreover, if an orbital has two electrons, they must be of opposite spin. For one s = + ½ and
for other s = - ½ . For e.g., the set of four numbers for the two electrons present in 1s-orbital are:
Thus, the Pauli's exclusion principle may also be stated as "An orbital can have maximum of two
electrons and these two must have opposite spin". The maximum number of electrons in each main
energy level can be predicted by extending this principle. The tabular column given below the various
combinations of quantum numbers for electrons in n = 1 to n =3 levels.
Distribution of electrons in terms of permitted values of quantum numbers
27
Hund's rule of maximum multiplicity
According to this rule, electron pairing will not take place in orbitals of same energy (same sub-shell)
until each orbital is first singly filled with parallel spin. In other words in a set of orbitals having same
energy (degenerate orbitals), the electrons distribute themselves to occupy separate orbitals with
same spin as far as possible. This rule can be illustrated by considering the example of carbon. The
atomic number of carbon is 6 and it contains two electrons in 2p subsheIl and these can be distributed
in the following three ways:
28
Since all the three 2p orbitals have same energy, therefore, it does not take any difference as to which
of the three orbitals contain electrons. In state (a) both the electrons are in the same orbital. In state
(b), the two electrons are present in different orbitals but with opposite spins while in state (c), the
electrons are present in different orbitals with same spins. Now, the electrons are charged particles
and repel one another. The electron-electron repulsions are minimum when the electrons are as far
apart as possible with parallel spins. Thus, state (c) has minimum repulsions and corresponds to lower
energy (stable state). This is in accordance with Hund's rule. This principle is very important in guiding
the filling of p, d and f subshells, which have more than one type of orbitals.
Aufbau principle and Bohr Bury rule
The distribution of electrons in different orbitals is known as its electronic configuration. This
characterizes each electron in an atom. The electronic configuration is expressed by indicating the
principal quantum number and its respective orbital along with the number of electrons present in it.
For example the notation 3px1 indicates that in the third principal shell there is one electron in the 'px'
orbital.
Sometimes the electronic configuration is also described by box notation form i.e., putting an arrow for
single electron in a box or a pair of arrows for two electrons in a box. The direction of the arrows gives
the orientation of its spin.
Further the box is labelled on top by writing the symbol of the orbital.
Rules-for-Filling-the-Orbitals">
Rules for Filling the Orbitals
Aufbau principle
The principle states that the electron in an atom are so arranged that they occupy orbitals in the
order of their increasing energy. Since the energy of a 'n' orbital in the absence of any magnetic field
depends on the 'n' and 'l' quantum number values, the order of filling orbitals with electrons may be
obtained from the (n + l) rule of Bohr Bury's rule.
According to this principle the orbital with the lowest energy will be filled first. The orbital having lower
(n + l) value has lower energy. However for orbitals whose (n + l) values are equal, the orbital having
lower value of 'n' has lower energy. It is important to remember that because of this rule, this
sequence of energy levels pertains to energy level up to '3p' and thereafter, '4s' orbitals comes first
instead of '3d'. Thus, the orbitals should be filled in the order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s
29
Fig: 3.19 - A simple way to determine the relative energies of different orbitals
Bohr Bury rule
Electrons revolve around the nucleus in different energy levels or shells and each shell is associated
with definite energy. The energy of the K shell is the least while those of L, M, N and O shells increases
progressively. We also know that any system that has least energy is the most stable.
1st energy level is K shell
2nd energy level is L shell
3rd energy level is M shell
4th energy level is N shell and so on.
Electronic configuration of an element
The arrangement of electrons in the various shells/orbits/energy levels of an atom of the element is
known as electronic configuration.
Bohr and Bury Scheme - Important Rules

Maximum number of electrons that can be accommodated in a shell is given by 2n2 where n=shell
number
30
For 1st energy level, n = 1
Maximum number of electrons in 1st energy level = 2n2
2 x (1) 2 = 2

For 2nd energy level n=2
Maximum number of electrons in the 2nd energy level = 2n2
2 x 22 = 2 x 4 = 8

For 3rd energy level n=3
Maximum number of electrons in the 3rd energy level = 2n2
= 2x(3)2
= 2x9=18

For 4th energy level n=4
Maximum no.of electrons in the 4th energy level = 2n2
= 2x(4)2
= 2x16=32

Sl No.
Electron Shell
Maximum capacity
1.
K Shell
2 electrons
2.
L Shell
8 electrons
3.
M Shell
18 electrons
4.
N Shell
32 electrons
The outermost shell of an atom cannot accommodate more than 8 electrons, even if it has a
capacity to accommodate more electrons. This is a very important rule and is also called the Octet
rule. The presence of 8 electrons in the outermost shell makes the atom very stable.
Keeping these points in mind let us write the electronic configuration of elements.

Electronic configurations of some important elements
Element
Symbol
Atomic
number
Electronic configuration (or Electron
arrangement) KLMN
Hydrogen
H
1
1
Helium
He
2
2
Lithium
Li
3
2,1
Beryllium
Be
4
2,2
Boron
B
5
2,3
Carbon
C
6
2,4
Nitrogen
N
7
2,5
Oxygen
O
8
2,6
Fluorine
F
9
2,7
Neon
Ne
10
2,8
31
Sodium
Na
11
2,8,1
Magnesium Mg
12
2,8,2
Aluminium Al
13
2,8,3
Silicon
Si
14
2,8,4
Phosphorus P
15
2,8,5
Sulphur
S
16
2,8,6
Chlorine
Cl
17
2,8,7
Argon
Ar
18
2,8,8
Potassium
K
19
2,8,8,1
Calcium
Ca
20
2,8,8,2
Geometric Representation of Atomic Structure
Example:
Steps:

The first 2 electrons will go to the 1st shell = K Shell (2n2)

The next shell L takes a maximum of 8 electrons (2n2).
In this way 2 + 8 = 10 electrons have been accommodated. The next 2 electrons go to the M Shell.
KLM
2,8,2

Example
32
Special case of potassium and calcium elements
Atomic number of potassium is 19 and its electronic configuration is
KLMN
1881
Atomic number of calcium is 20 and its electronic configuration is
KLMN
2882
This abnormal behaviour can be explained as follows:
It is found that shells have sub shells. The smaller sub shells are termed s, p, d and f. The maximum
number of electrons that can go into these are 2, 5, 10 and 14 respectively. These sub shells can
overlap, resulting in energies that may differ from that predicted purely on the basis of n=1, 2, 3 etc.
Therefore when electrons start filling, they may go to a new outer shell even before the inner shell is
filled to capacity.
Electronic configuration of the atoms and ions
Element
Symbol
Atomic
number
Electronic configuration (or Electron
arrangement) KLMN
Hydrogen
H
1
1
Helium
He
2
2
Lithium
Li
3
2,1
Beryllium
Be
4
2,2
Boron
B
5
2,3
Carbon
C
6
2,4
Nitrogen
N
7
2,5
Oxygen
O
8
2,6
Fluorine
F
9
2,7
Neon
Ne
10
2,8
Sodium
Na
11
2,8,1
33
Magnesium Mg
12
2,8,2
Aluminium Al
13
2,8,3
Silicon
Si
14
2,8,4
Phosphorus P
15
2,8,5
Sulphur
S
16
2,8,6
Chlorine
Cl
17
2,8,7
Argon
Ar
18
2,8,8
Potassium
K
19
2,8,8,1
Calcium
Ca
20
2,8,8,2
Periodic Table
Periodic Classification of Elements
At the turn of the nineteenth century, about 30 elements were known. 50 years later by the 1850s,
scientists had discovered sixty three chemical elements and the numbers kept increasing. With the
discovery and study of more and more elements and their compounds, the various data about them
also increased. It became progressively difficult to organize all that was known about the elements and
scientists made several attempts to look for some pattern in their properties, on the basis of which
they could organize study such a large number of elements with ease.
Initially scientists had classified elements into metal and non-metals. However, some elements
possessed properties which could neither be classified as metals nor non-metals called metalloids. This
classification was found to be insufficient for scientific study. Later on, a number of chemists
attempted to make a rational and systematic classification of the physical and chemical properties of
elements and tabulate the results in the form of a table.
Mendeleev's periodic law and periodic table
Mendeleev arranged the elements into 8 groups and 7 periods. He did this in increasing order of their
atomic masses. Mendeleev stated the law of chemical periodicity as: "The physical and chemical
properties of elements are periodic functions of their atomic mass."
Contributions of Mendeleev's Periodic Table
(i) Systematic study of elements
Mendeleev's Periodic table simplified the study of elements. It became useful in studying and
remembering the properties of a large number of elements, in a simpler way. This is because the
elements showing similar properties belonged to the same group.
(ii) Prediction of new elements
While arranging the elements, in increasing order of atomic mass, Mendeleev left three blanks for
elements that were not discovered at that time. He was able to predict the properties of these
unknown elements more or less accurately. He named them eka-boron, eka-aluminium and eka-silicon.
He named them so, as they were just below boron, aluminium and silicon in the respective sub-groups.
Eka-boron was later named as scandium, eka-aluminium as gallium and eka-silicon as germanium.
A comparative study of the properties of the elements predicted and later
34
discovered
Property
Atomic weight
Oxide
Specific gravity
Sulphate
Property
Eka-boron
Scandoum
44
Eb2O3
3.5
Eb2(SO4)3
43.79
Sc2O3
3.864
Sc2(SO4)3
Eka-aluminium
Atomic weight
Specific gravity
Melting point
Formula of oxide
Solubility in acid
and alkali
Gallium
58
5.9
Low
Ea203
Dissolves slowly in both
acid and alkali
Property
69.9
5.94
303. 15°K
Ga203
Dissolves slowly in
both acid and alkali
Eka-silicon
Atomic weight
Specific Gravity
Melting point
Valency
Reaction with acid
and alkali
Germanium
72.32
5.47
958°C
4
Dissolves neither by
hydrochloric acid nor
sodium hydroxide
72
5.5
High
4
Slightly attacked by
acids, resists attack by alkali
(iii) Correction of atomic masses
Mendeleev's periodic table helped in correcting the atomic masses of some of the elements, based on
their positions in the periodic table. For e.g., atomic mass of beryllium was corrected from 13.5 to 9.
Atomic masses of indium, gold, platinum were also corrected.
Mendeleev's Periodic Table
Grp:1
Grp:2
Grp:3
Grp:4
Grp:5
Grp:6
Grp:7
Grp:8
H=1
-
-
-
-
-
-
-
Li=7
Be=9.4
B=11
C=12
N=14
O=16
F=19
-
Na=23
Mg=24
Al=27.3
Si=28
P=31
S=32
Cl=35.5
-
K=39
Ca=40
-
Ti=48
V=51
Cr=52
Mn=55
Fe=56
Co=58.9
Ni=58.7
cu=63
-
Zn=65
-
-
As=75
se=78
Br=80
-
Rb=85
Sr=87
Yt=88
Zr=90
Nb=94
Mo=96
-
Ru=104
Rh=104
pb=106
Ag=108
-
cd=102
In=113
Sn=118
Sb=122
Te=127.6
I=126.9
-
35
Grp:1
Cs=133
Grp:2
Ba=137
Grp:3
Di=138
Grp:4
Ce=140
Grp:5
-
Grp:6
-
Grp:7
Grp:8
-
-
-
-
Er=178
La=180
Ta=82
W=184
-
Os=195
Ir=197
Pt=198
Au=199
-
Hg=200
Tl=204
pb=207
Bi=208
-
-
-
-
-
-
Th=231
-
U=240
-
-
Anamolies of Mendeleev's periodic table
(i) Position of hydrogen
The position of hydrogen is not correctly defined. It is still not certain whether to place hydrogen in
group I A or VII A.
(ii) Anomalous pair
In certain pairs of elements like, Ar (40) and K (39); Co (58.9) and Ni (58.6); Te (127.6) and I (126.9) the
arrangement was not justified. For e.g., argon was placed before potassium whereas its atomic mass is
more than potassium.
(iii) Position of isotopes
Isotopes are atoms of the same element having different atomic mass but same atomic number. For
e.g., there are three isotopes of hydrogen with atomic mass 1, 2, and 3. According to Mendeleev's
periodic table these should be placed at three separate places. However isotopes have not been given
separate places in the periodic table.
(iv) Grouping of elements
Some similar elements are separated and dissimilar elements are grouped together e.g., copper and
mercury have similar properties but are placed in different groups. On the other hand elements of
group IA such as lithium, sodium and potassium were grouped with dissimilar elements such as copper,
silver and gold.
(v) Cause of periodicity
Mendeleev's table was unable to explain the cause of periodicity among elements.
(vi) Position of lanthanides and actinides
Fourteen elements that follow lanthanum called lanthanides and fourteen elements following actinium
called actinides were not given proper places in Mendeleev's periodic table.
Modern periodic law, and modern periodic table
Later, Henry Gywn-Jeffreys Moseley, an English Physicist,showed that the atomic number of an
element is numerically equal to the number of electrons round the nucleus.He suggested the modern
periodic law which states that the properties of the elements are the periodic functions of their atomic
numbers.
From his modern periodic law, periodicity derived its meaning which is ‘‘The recurrence of similar
properties of the elements when they are arranged in the order of increasing atomic number, after
certain regular intervals’’.
36
Advantages of modern Periodic table
1. The modern periodic table is divided into two main categories:
(i) Vertical columns called Groups and
(ii) Horizontal columns called Periods.
37
2. There are 18 vertical columns or groups. These are further sub-divided into A and B (groups I to VII),
VIII group and zero group.
3. Member of the same group have similar electronic configuration of the valence shell and thus show
same valency.
4. Elements of groups IA to VIIA are called groups of typical elements, representative elements or
normal elements.
5. Groups IA and IIA are strongly metallic and are called group of 'alkali metals and alkaline earth
metals' respectively.
6. Groups IB to VII B and VIII lie in the middle of the table between IIA and IIIA groups and are called
groups of transition elements. They consist of metals.
7. The Zero group consists of 'Noble gases'.
8. There are 7 horizontal rows in the periodic table. These are called the periods.
9. In a period, the number of valence shell remains the same for all elements. However, the number of
electrons in the valence shell increases from left to right .
10. The number of elements present in each period is given in the following table.
Period
Valence shell
Type of Period
Number of elements
Atomic Number of elements
1st Period
n=1
Short period
2
Atomic number 1 and 2
2nd Period
n=2
Short period
8
Atomic number 3 to 10
3rd Period
n=3
Long period
8
Atomic number 11 to 18
4th Period
n= 4
Long period
18
Atomic number 19 to 36
5th Period
n=5
Long period
18
Atomic number 37 to 54
*6th Period
n=6
Long period
32
Atomic number 55 to 86
*7th Period
n=7
Incomplete
23
Atomic number 87 to 109
38
11. The 6 th period consists of elements that have atomic numbers 58 to 71. They are called
Lanthanides. The 7th period consists of elements that have atomic numbers 90 to 105. They are called
Actinides. Both of them are called inner transition elements.
12. The 7 th period is an incomplete period as it has only 23 elements.
13. Lanthanides and actinides are not accommodated in the main body of the periodic table but are
placed separately at the bottom of the table.
14. The position of hydrogen is not certain. Thus it can be placed in both group IA and group VIIA.
15. Group VIIA elements are called halogens or salt producers. Representative periodic table for eight
groups up to calcium (atomic number 20) with their electronic configuration is given in the table.
Representative periodic table for elements with atomic number from 1 to 20 with their electronic
configuration
Group -->
1A
2A
3A
4A
5A
6A
7A
0
Period
Period 1
Period 2
Period 3
Period 4
1H
2He
1
2
3Li
4Be
5B
6C
7N
8O
9F
10Ne
2,1
2,2
2,3
2,4
2,5
2,6
2,7
2,8
11Na
12Mg
13Al
14Si
15P
16S
17Cl
18Ar
2,8,1
2,8,2
2,8,3
2,8,4
2,8,5
2,8,6
2,8,7
2,8,8
19K
2,8,8,1
20Ca
2,8,8,2
Division of elements into s,p, d and f blocks
Periodic Table
KEY
ELEMENT CLASSIFICATION
Non-Metals
Rare Earth Elements
39
Semimetallics
Alkali Metals
Transition Metals
Alkaline-earth Metals
Other Metals
Radioactive Rare Earth Elements
Halogens
Inert Gases
Periodicity of physical properties
The recurrence of similar properties of the elements when they are arranged in the order of increasing
atomic number, after certain regular intervals, is called periodicity.
Cause of periodicity of elements
The modern periodic table is based on the electronic configuration of the elements. The properties of
an element are determined largely by the electrons in its outermost or valence shell. Valence electrons
interact with other atoms and take part in all chemical reactions, while inner shell electrons have little
influence on the properties of elements.
When elements are placed in the order of their increasing atomic number, the elements having the
same number of valence shell electron is repeated in such a way, so as to fall under the same group.
Since, the electronic configuration of the valence shell electrons is same they show similar properties.
Periodicity of valency
The electrons present in the outer most shell are called valence electrons. The valency of an element is
defined as its 'combining capacity'.
Variation along the period
The number of valence electrons increase from left to right in a period. So, the valency of the elements
first increase from 1 to 4 and then decreases to zero.
Variation down the group
On moving down the group the number of valence electrons remain the same and therefore all the
elements in the group exhibit the same valency.
Periodicity of atomic radii
When the elements are arranged in the order of increasing atomic numbers there is a recurrence of
similar properties after certain regular intervals. This regularity is called the periodicity in properties.
The distribution of electrons in the various shells determines their physical and chemical character. It
has been observed that the properties of the elements depend more on the arrangements of the
electrons in the outermost shell (valence shell) and not on the inner shells. This repetition of similar
electronic configuration in the outermost or valence shell after certain regular intervals causes
periodicity in the properties of elements.
Atomic radii
The size of the atom is significant in governing its property. If the atom is assumed to be spherical, then
the radius of the sphere gives the atomic radius. But it is difficult to exactly determine the radius of the
atom because

The probability of finding the electron is never zero even at large distances from the nucleus and so
the atom does not have a well defined boundary.
40

It is not possible to isolate an atom and measure its radius.
The size of the atom changes in going from one set of environment to another and from one
bonded state to another.
So, one can arbitrarily define atomic radius as the effective size which is the distance of closest
approach of one atom to another atom in a given bonding situation. This approximate radius can be
determined by measuring the inter-nuclear distance between the two centres of the neighbouring
atoms in a covalent molecule. This is usually done by diffraction and spectroscopic techniques.

Fig: 4.3 - Calculation of atomic radius
The inter-nuclear distance corresponds to the diameter of the atom and therefore half of this distance
gives the atomic radius for a homonuclear molecule like Cl-Cl or Br-Br.
Hence, it may be defined as one half of the distance between the centers of the nuclei of two similar
atoms bonded by a single covalent bond. This is also called as covalent radius.
For a hetero-nuclear molecule, the covalent radius is the distance between the center of the nucleus of
the atom and the mean position of the shared paired of electrons between the bonded atoms.
The covalent radii are smaller than the atomic radii in the uncombined atoms because the overlap
region between atomic orbitals of two atoms becomes common in a covalent bond.
The forces of attraction (Van der Waals forces) existing between non-bonded atoms and molecules are
weak and the atoms are held at larger inter-nuclear distances. Thus these radii known as Van der
Waals radii are always larger than covalent radii.
Van der Waals radius is defined as one half of the inter-nuclear distance between two adjacent atoms
belonging to the two nearest neighbouring molecules of the substance in the solid state.
Variation of Atomic Radii
Variation in a period
Atomic radii in general, decrease with increase in atomic number, going from left to right in a period.
This is explained on the basis of increasing nuclear charge along a period. The nuclear charge increases
progressively by one unit while the corresponding addition of one electron takes place in the same
principal shell. As the electrons in the same shell do not screen each other from the nucleus, the
nuclear charge is not neutralized by the extra valence electron. Consequently the electrons are pulled
closer to the nucleus by the increased effective nuclear charge resulting in the decrease in the size of
the atom. In this way the atomic size goes on decreasing across the period.
41
Fig: 4.4 - Variation of atomic radius with atomic number in a period
The atomic radius abruptly increases in the case of noble gas element Neon as it does not form
covalent bonds. So the value of Neon radius is Van der Waals radius which is considerably higher than
the value of other covalent radii.
Variation in a group
The atomic radii of elements increases from top to bottom in a group because the nuclear charge
increases with increasing atomic number. Although, there is an increase in the principal quantum
number from one atom to another, the number of electrons in the valence shell remain the same. The
effect of increase in the size of the electron cloud out weighs the effect of increased nuclear charge
and so the distance of the valence electron from the nucleus increases down the group. Thus the size
of the atom goes on increasing down the group in spite of increasing nuclear charge.
Fig: 4.5 - Variation of atomic radius with atomic number in a group
Problem
4.The atomic mass of germanium is 72.6 and its density is 5.47g cm-3. What is the atomic volume of
germanium?
Solution
42
Periodicity of ionic radii
These are radii of ions in ionic crystals. Ionic radius may be defined as the effective distance from the
center of nucleus of an ion up to which it has an influence on its electron cloud. In ionic compounds the
inter nuclear distance may be taken as equal to the sum of the ionic radii of the two ions. The inter
nuclear distance in ionic crystals are obtained from X-ray studies.
Fig: 4.6 - Internuclear distance and ionic radii
Radius of the cation
A cation is formed by the loss of one or more electrons from the gaseous atom. Thus, the whole of the
outer most shell of electrons is removed resulting in the smaller size in the cation.
For example, in lithium atom, there is only one electron in the outermost '2s' shell. As the lithium atom
changes to Li+ ion the outer most '2s' shell disappears completely. This disappearance results in the
decrease in size.
With the removal of electrons from an atom the magnitude of the nuclear charge remains the same
while the number of electrons decreases. As a result the nuclear charge acts on less number of
electrons. The effective nuclear charge per electron increases and the electrons are more strongly
attracted and pulled towards the nucleus. This causes a decrease in the size of the ion.
Radius of the anion
The negative atom is formed by the gain of one or more electrons in the neutral atom. The number of
electrons increases while the magnitude of nuclear charge remains the same. The same nuclear charge
acts on larger number of electrons than were present in the neutral atom. The effective nuclear charge
per electron is reduced and the electron cloud is held less tightly by the nucleus. This causes an
increase in the size of the ion. Thus anions are larger in size than the corresponding atom.
Variation of ionic radii in a group
The ionic radii in a particular group increases in moving from top to bottom because of the increase in
the principal quantum number though the number of electrons in the valence shell remains the same.
In isoelectronic series of ions, as the nuclear charge increases the electrons are pulled more and more
strongly and the size decreases.
Problem
5. Out of Na+ and Na which has the smaller size and why?
43
Solution
Na+ has a smaller size than Na. Na+ is formed by the removal of one electron from Na. However both of
these posses the same nuclear charge. Therefore electrons in Na + are more tightly held than in Na. The
removal of one electron from Na also leads to complete removal of the third shell so that in Na +, the
outermost shell is second. Hence Na+ has a smaller size than Na.
Periodicity of ionization energy
The amount of energy required to remove the most loosely bound electron from an isolated gaseous
atom is called ionization energy (IE).
Ionization energy is also called as ionization potential because it is measured as the minimum potential
required to remove the most loosely held electron from the rest of the atom. It is measured in the
units of electron volts (eV) per atom or kilo joules per mole of atoms (kJ mol-1).
1 eV per atom = 96.64 kJ mol-1 = 23.05 k cal mol-1
Thus, the ionization energy gives the ease with which the electron can be removed from an atom. The
smaller the value of the ionization energy, the easier it is to remove the electron from the atom.
Factors Governing Ionization Energy
Size of the atom
As the size of the atom increases the outermost electrons are held less tightly by the nucleus
(attractive force between the electron and the nucleus is inversely proportional to the distance). As a
result it becomes easier to remove the electron and therefore the ionization energy decreases with the
increase in atomic size.
Charge on the nucleus
The attractive force between the nucleus and the electron increases with the increase in nuclear
charge making it more difficult to remove an electron. The ionization energy thus increases with the
increase in the nuclear charge.
Screening effect
In multielectron atoms, the outermost electrons are shielded or screened from the nucleus by the
inner electrons. This is known as shielding or screening effect. The outer most electrons do not feel the
complete charge of the nucleus and the actual charge felt is called the effective nuclear charge. When
the inner electrons are more, the screening effect will be large, the nuclear attraction will be less. Thus
when the inner electrons increase the ionization energy will decrease.
Penetration effect
The 's' electrons are more penetrating (maximum probability of finding near the nucleus) towards the
nucleus than the p electrons. The order of penetration power in a given shell is s > p > d > f
If the penetration power of the electron is more, it will be closer to the nucleus and will be held more
firmly. Thus ionization energy will increase with the increase in the penetration power of the electrons.
For the same sub shell the ionization energy would be more to remove an 's' electron than to remove a
'p' electron which in turn will be more than that for removing a 'd' electron.
Electronic arrangement
Certain electronic configuration like half-filled and completely-filled shells have extra stability. It is
more difficult to remove electron from these stable configuration and the ionization energy is very
44
high. For example, the noble gases have the most stable configuration and so have high ionization
energy; elements like Be and Mg have completely filled orbitals while N and P have exactly half-filled
sub shells. Thus, their ionization energies are high. The more stable the electronic configuration, the
higher is the ionization energy.
Variation along a period
The ionization energy increases with increasing atomic number in a period. This is because

The nuclear charge increases on moving across a period from left to right.
The atomic size decreases along a period though the main energy level remains the same.
Due to the increased nuclear charge and simultaneous decrease in atomic size, the valence electrons
are more tightly held by the nucleus. Therefore more energy is needed to remove the electron and
hence ionization energy keeps increasing. However some irregularities have been noticed due to the
extra stability of the half filled and completely filled configurations.
For example, the nuclear charge on Boron is more than Beryllium, yet there is slight decrease in
ionization energy from Be to B. This is because, in boron the last electron goes to '2p' orbital which is at
a slightly higher energy than '2s' orbital. Also, the electronic configuration of B is less stable than that
of Be (has completely filled orbitals). Hence the ionization energy is less than that of Be. Similarly,
nitrogen, which has half filled '2p' orbitals, is more stable than oxygen. Therefore the ionization energy
of nitrogen is more than that of oxygen.

Fig: 4.7 - Variation of ionization energy with atomic number
45
Variation down a group
The ionization energy gradually decreases in moving from top to bottom in a group. This is due to the
fact that:

The nuclear charge increases in going from top to bottom in a group.

An increase in the atomic size due to an additional energy shell (level) 'n'.
Due to the increase in the number of inner electrons there is an increase in the shielding effect on
the outer most electron. The effect of increase in atomic size and the shielding effect is much more
than the effect of increased nuclear charge.
As a result , the electron becomes less firmly held to the nucleus and so the ionization energy
decreases as we move down the group.

Fig: 4.8 - The variation of ionization energy with atomic number
Successive ionization energies
The energies required to remove subsequent electrons from a gaseous atom is called as successive
ionization energies. They are termed as first ,second, third …… ionization energy depending on the
removal of the first, second, third electron respectively.
The second ionization energies are higher than the first due to the fact that after the removal of the
first electron the atom changes into a monovalent positive ion. In this ion, the number of electrons
decreases but the nuclear charge remains same and so the remaining electrons are held more tightly
by the nucleus and it becomes difficult to remove the second electron. Hence the value of the second
ionization energy (IE2) is higher than the first (IE1). In the same way the removal of the second electron
will result in the formation of di-positive ion making the attraction between the nucleus and the
remaining electrons stronger. This results in higher value of third ionization energy (IE3).
Problems
6. Calcium (Z = 20) loses electrons successively to form Ca+, Ca2+ and Ca3+ ions. Which step will have
highest ionization energy?
Solution
The step, which involves the formation of Ca3+ from Ca2+� would have highest ionization energy.
46
7. Consider the ground state electronic configurations given below:
(A) 1s2 2s2 2p6 (B) 1s2 2s2 2p4 (C) 1s2 2s2 2p6 3s2 (D) 1s2 2s2 2p6 3s1 (E) 1s2 2s2 2p5.
Which of the above configuration is associated with the lowest and which is associated with highest
ionization energy?
Solution
Lowest ionization energy = D
Highest ionization energy = A.
8. Why is the ionization energy of nitrogen unexpectedly high?
Solution
The electronic configuration of nitrogen (1s22s22p3) shows that it possesses exactly half-filled 'p'
orbitals as its outershell configurations. Since, half-filled orbitals are extraordinarily stable, it is more
difficult to remove an electron from this atom. Thus nitrogen has unexpectedly high ionization energy.
Periodicity of electron affinity
Electron affinity is the amount of energy released when an electron is added to an isolated gaseous
atom.
Electron affinity is the ability of an atom to hold an additional electron. If the atom has more tendency
to accept an electron then the energy released will be large and consequently the electron affinity will
be high. Electron affinities can be positive or negative. It is taken as positive when an electron is added
to an atom. It is expressed as electron volts per atom (eV per atom) or kilo joules per mole.
Factors affecting electron affinity

When the nuclear charge is high there is greater attraction for the incoming electron. Therefore
electron affinity increases as the nuclear charge increases.

With the increase in the size of the atom the electron affinity decreases because the distance
between the nucleus and the incoming electron increases.

Electron affinities are low or almost zero in elements having stable electronic configurations (half
filled and completely filled valence subshells) because of the small tendency to accept additional
electron.
Variation along a period
The size of an atom decreases and the nuclear charge increases on moving across a period. This results
in greater attraction for the incoming electron. Hence the electron affinity increases in a period from
left to right.
Variation down a group
As we move down a group the atomic size and nuclear size increases. As the effect of increase in
atomic size is more pronounced the additional electron feels less attracted by the large atom.
Consequently the electron affinity decreases.
However there are some irregularities observed in the above general trend. They are:
The halogens have highest electron affinity because they have only one electron less than the stable
noble gas configuration.They have a strong tendency to accept an additional electron. This makes their
electron affinity values high.
47
The electron affinity value of noble gases are zero because of the stable electronic configuration of
ns2np6 which has no tendency to take in additional electron. No energy is released and their electron
affinities are zero.
The electron affinity values for Be, Mg, N and P are almost zero because of the extra stability of
completely filled orbitals in Be and Mg and half filled orbitals in N and P.
Electron affinity of fluorine is unexpectedly less than that of chlorine. The low electron affinity value of
F is due to the very small size of F atom. This small size results in strong inter electronic repulsions in
the relatively compact 2p subshell of fluorine: thus the incoming electron does not feel much
attraction. Electron affinity for some third period elements (e.g., P, S, Cl) are greater than
corresponding second period elements (e.g., N, O, F) because of the smaller atom size of second period
elements, which produces larger electronic repulsions for the additional electron.
Succesive electron affinities
The second electron is added to the negatively charged ion and the addition is opposed by coulombic
repulsions. The energy has to be supplied to force the second electron into the anion.
First electron affinity
Second electron affinity
The second electron affinities in which energy is absorbed have negative values while the first electron
affinity have positive values as energy is released.
Problem
9. Arrange the following in the decreasing order of electron affinity: B, C, N, O.
Solution
All these elements belong to the same period. The size of an atom decreases and the nuclear charge
increases on moving across a period. This results in greater attraction for the incoming electron. Hence
the electron affinity increases in a period from left to right O, C, B, N.
Periodicity of electronegativity
Electronegativity is the tendency of an atom to attract electrons towards itself in a molecule of a
compound. The value of electronegativity of an element describes the ability of its atom to compete
for electrons with the other atom to which it is bonded. Electronegativity is however not the property
of an isolated atom.
Electronegativity increases from left to right in each period ending at group VII.
In the 3rd period, electronegativity increases from sodium to chlorine i.e., chlorine can accept electrons
most easily in that period followed backwards by sulphur, phosphorus, silicon, aluminium, magnesium
and sodium. All the atoms of the above mentioned elements have three shells but chlorine has the
smallest atomic radii. Hence chlorine experiences more positive charge from the nucleus than all other
atoms in that period. So, if one electron is available, chlorine can attract it most easily.
48
Types of Electronegativity

When the molecule is formed by transfer of electrons (ionic bonding) the transfer takes place from
electropositive atom to electronegative atom (Fig. 2.4). In the example below, Na is electropositive
and Cl is electronegative.

If the molecule is formed by sharing of electrons (covalent bond) the bonded pair of electrons shift
towards more electronegative atom resulting in the formation of polar molecule. In the example
below, chlorine atom is more electronegative as compared to hydrogen atom, resulting in a
covalent bond where the shared pair of electron shifts towards the more electronegative atom.
This results in polar molecules (Fig. 2.5).
The electron pair is more closer to the chlorine atom and so the molecule gets polarized i.e., the
chlorine atom gets a negative charge while the hydrogen atom gets a positive charge.
Remember :

Fluorine is the most electronegative element.
A summary of periodic properties and their variation in groups and periods is given below:
Variation along the period
49
Due to the increased nuclear charge across the period, the electronegative or non-metallic character
increases in going from left to right in a period.
Variation down the group
Due to the increased atomic size the non-metallic or electronegative character decreases as we go
down the group.
Electronegativity is however not the property of an isolated atom. Electronegativity is measured on a
number of scale levels, the most commonly used are of Pauling or Mulliken.
Factors Affecting Electronegativity
1. Atomic size
As the size of the atom decreases it has greater tendency to attract the bonding electrons towards
itself. Therefore smaller atoms have higher electronegativity values than the larger ones.
2. Ionisation energy and electron affinity
Higher Ionisation energy and electron affinity lead to higher electronegativity.
3. Number and nature of atoms
The electronegativity depends on the number and nature of atoms bonded to it.
4. Type of hybridization
The electronegativity increases with the increase in 's' character in the hybrid orbital. This is because
the 's' orbitals being more near to the nucleus have greater tendency to attract the shared pair of
electron.
5. Charge on the ion
A cation has high electronegativity while an anion has less electronegativity than its parent atom. A
cation with a higher positive charge is more electronegative.
NB: Further illustrations will be done in the Class
RADIOACTIVITY
Concept of radioactivity
The discovery of the electron towards the end of the nineteenth century was the starting point of new
avenues of research in science, which were to give physicists an insight into the structure and nature of
the atoms of matter.
In 1896, Henri Becquerel discovered that when a photographic plate wrapped in black paper was
placed near a salt of uranium, it got affected. The same observation was repeated with other salts of
uranium. This led him to conclude that the uranium salt emitted some deep penetrating radiation.
Further experiments showed that the intensity of the emitted radiation depended directly on the
concentration of uranium in its various salts. The emission of radiation was completely unaffected by
any change in the physical and chemical conditions of the system. Becquerel concluded that the origin
of the radiation was somehow rooted in the nucleus of the uranium atom. The radiation discovered by
Becquerel was initially named Becquerel rays.
Later on, Madam Curie and Pierre Curie discovered other substances i.e., polonium, radium, etc.,
which were more active than uranium.
These substances, which have the property of spontaneous emission of radiation are called radioactive
substances and the process of spontaneous emission of radiation is called Radioactivity.
50
It was later found that all elements, which have an atomic number greater than 82 are naturally
radioactive.
Radioactivity is a nuclear phenomenon. It is the spontaneous emission of radiation from the nucleus.
Radioactive rays (alpha ray, beta ray & gamma ray)
In 1899, the study of radioactivity was taken up by Ernest Rutherford. He placed a little radium at the
bottom of a small lead box and subjected the rays that emerged from it to the action of a very strong
magnetic field at right angles to their direction. He found that the rays separated into three distinct
constituents as shown in the figure below.
Lead box containing radium
For convenience, Rutherford called the three types of radiation alpha (a), beta (b) and gamma (g) rays.
The a-rays were deflected in a direction opposite to that of b-rays. This showed that the a-rays carried
a positive charge, b-rays carried a negative charge and those which passed undeviated were neutral or
uncharged were g-rays.
Similarly, if the radiations given out by a radioactive substance are subjected to an electric field
perpendicular to their path, they separate into three constituents. Those which turn towards the
negative plate are the positively charged alpha particles. Those, which turn towards the positive plate,
are the negatively charged beta particles. Those, which pass undeviated, are the uncharged gamma
radiations.
Further investigation has shown that an alpha ray is a stream of helium nuclei, a beta ray is a stream of
electrons and a gamma ray is an electromagnetic radiation whose frequency is higher than that of Xrays.
Properties of alpha particles

They are deflected by electric and magnetic fields. The deflection is less compared to that of
bparticles, because a particles are heavy.

They affect a photographic plate and cause fluorescence on striking a fluorescent material.

They ionise the gas through which they pass.

The mass of an alpha particle is 6.643 x 10-27 kg, which is roughly four times the mass of a proton.
Its charge is +3.2 x 10-19 C, which is two times the charge of a proton. Thus its specific charges i.e.,
q/m value is 4.826 x 107 C/kg.
51

An alpha particle consists of two protons and two neutrons.

The speed of a-particle is of the order of 107 m/s.

Its penetrating power is very small.

They have large kinetic energy.

They destroy living cells and cause biological damage.

They get scattered while passing through thin mica or gold foils.
Properties of Beta particles

They get deflected by electric and magnetic fields. The deflection is large since a beta particle is
lighter than an a-particle.

They affect photographic plates.

They ionise the gas through which they pass.

The mass is 9.1 x 10-31 kg and its charge is 1.6x10-19 C. Thus its specific charge q/m value is 1.76 x
1011 C/kg.

They are fast moving electrons emitted from the nucleus of the atom.

Its speed is of the order of 108 m/s.

The penetrating power of b particles is more than that of

They cause fluorescence on striking a fluorescent material.

They produce X-rays when they are stopped by metals of high atomic number and high melting
point such as tungsten.

They cause greater radiation damage as they can easily pass through the skin of the body.
particles.
Properties of gamma radiation

They are not deflected by electric and magnetic fields.

They affect photographic plates.

The ionising power is very low compared to alpha-particles and beta-particles.

They are electromagnetic waves like X-rays and light rays. The wavelength of these rays are shorter
than that of X-rays.

The speed is the same as the speed of light.

The penetrating power is high.

They cause fluorescence when they strike a fluorescent material.

They are diffracted by crystals.

Although X-rays and g-rays have similar properties, their origin is different. X-rays originate in the
electron cloud outside the nucleus, where as gamma rays originate in the nucleus.

They can easily pass through the human body and cause immense biological damage.

They are very useful in the treatment of cancer.
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Distinction among the properties of α, β and γ radiation
Property
α-particle
β-particle
Nature
Stream of positively
charged particles i.e,
helium nuclei,
Stream of negatively charged
Electromagnetic
particles i.e., energetic
waves.
electrons.
Speed
Nearly 107 ms-1
About 90% of the speed of
light or 2.7 x 108 ms-1
Rest mass
4 times the mass of
Massof electron i.e,9.1X10-27
proton i.e., 6.64 X 10 kg 31kg
No mass
Charge
Positive charge +3.2x1019
C (+2e)
Negative charge —1.6x1019
C(-e)
No charge
1.76X1011C kg-1
-
Specific charge[q/m] 4.83X107C kg-1
γ-radiation
3 x lO8 ms-1 (in
vacuum)
10-13m or 10-3A0
Wavelength
Effect of electric and
Less deflected
magnetic fields
More deflected in direction
opposite to alpha particles.
Unaffected
Minimum
Ionising power
Maximum (10,000 times
of γ)
less than alpha(100 times of
γ)
Penetrating power
Small (3-8 cm in air)
Large (About 1 mm of lead or Very large (About 30
about 5 mm of Al)
cm of iron)
Biological damage
Cause some damage
Cause more damage
Cause immense
damage
Meaning of natural and artificial radioactivity
Radioactivity is broadly classified into two categories:
a) Natural radioactivity and
b) Artificial or Induced radioactivity.
If a substance emits radiations by itself, it is said to possess natural radioactivity.
If a substance does not possess radioactivity but starts emitting radiations on exposure to rays from a
natural radioactive substance, it is said to possess induced or artificial radioactivity.
The first artificial transmutation was caused out by Rutherford in 1919 who bombarded nitrogen gas
with alpha particles and obtained hydrogen and oxygen.
The isotope
are stable and no further disintegration takes place. Nowadays other
bombarding particles (i.e., projectiles) like protons (1H1), deutrons1H2 electrons can be accelerated to
very high speeds by fluctuating electric and magnetic fields in machines such as cyclotron, synchrotron
etc. These high speed particles are more efficient in causing nucleus to disintegrate on impact.
The typical transmutations involving various particles are summarised below:
(i) a particle induced reactions
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Since a particle is used and a neutron is produced, the reaction is termed as (a, n) reaction.
The first example of radioactivity produced by artificial means is given below:
The isotope produces is itself radioactive.
undergoes decay by positron (b+) emission.
(ii) Deutron induced reactions
(iii) Proton - induced reactions
(iv) Neutron - induced reactions
Thus this phenomenon in which artificial transmutation of a stable non - radioactive nucleus leads to
the formation of radioactive isotope is called artificial radioactivity or induced radioactivity.
Some of the isotopes produced as a result of neutron bombardment find applications in different
areas.
The features of artificial radioactive isotopes are given below:
i) Artificial radioactive isotopes exhibit behaviour similar to natural radioactive elements and follow
some rate law with regard to their disintegration.
ii) Natural radioactivity is shown by elements with high atomic numbers (> 83) whereas artificial
radioactivity can be induced in elements with low atomic numbers.
iii) Artificial radioactive isotopes generally have short half-life periods.
iv) They are very rarely found in nature because they decay off as soon as they are formed due to short
half-life periods.
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Nuclear reactions, Nuclear energy
Nuclear reactions
Two consequences of nuclear reactions are the phenomena of nuclear fission and nuclear fusion. Both
are important sources of nuclear energy, which may be used for peaceful or destructive purposes.
When a nucleus is bombarded with some sub-atomic particles such as a - particles, neutrons, protons
etc, these particles are captured by the target nucleus, which then disintegrates. The new element
formed has mass either slightly greater or slightly smaller than the parent element.
The process of splitting of a heavier nucleus (like that of U235) into a number of fragments of much
smaller mass, by suitable bombardment with sub-atomic particles is called nuclear fission.
Of the three natural isotopes of uranium
nucleus undergoes nuclear fission when bombarded with slow neutrons. U236 is formed which being
unstable, further breaks up in several different ways.
The tremendous amount of energy released during nuclear fission is because of the loss in mass. The
sum of the masses of the fragments produced and neutrons released as a result of fission is less than
the sum of the masses of target 235U and bombarding neutron. The loss in mass gets converted into
energy according to Einstein equation
E = mc2.
- nuclide splits up into 144Ba and 90Kr along with
For e.g., we can calculus the loss of mass when
the release of two neutrons.
Dm the mass defect or the mass converted into energy is given by
Dm = 236.127 - 235.846 = 0.281 amu
1 amu = 931.48 MeV
\ Energy released = 0.281 amu = 931.48 x 0.281
= 261.75 MeV
The neutrons emitted from the fission of first uranium atom hit other uranium nuclei and cause their
fission resulting in the release of more neutrons, which further continue the fission process. In this
way, a nuclear chain reaction sets up releasing tremendous amount of energy.
If the chain reaction takes place as mentioned, the energy released would be very high. This is the
principle underlying the nuclear or atom bomb. However some of the released neutrons escape from
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the surface unused and do not involve in the chain process. It is found that the net amount of energy
released from 1 kg of uranium 235 is equivalent to that available from 2 x 104 kg of coal.
It has been found that a branching chain reaction is possible only with a quantity of
larger than a
certain critical amount (so that only a few neutrons escape).
The minimum amount of the fissionable material required so as to continue the chain reaction process
under a given set of conditions is called the critical mass. If the mass of the material is more than the
critical mass, it is referred to as super-critical mass whereas if the mass of the fissionable material is
smaller than the critical mass, it is called sub-critical mass. The critical mass of U-235 is between 1 kg to
100 kg.
It must be noted that the naturally occurring uranium contains most U-238 isotope (about 99.3%)
which is not fissionable with slow neutrons.
Nuclear energy
Nuclear energy is the energy released when certain changes take place in the nucleus of an atom.
Nuclear energy is partly renewable and partly non-renewable source of energy.
Nuclear fission and nuclear fusion are the two types of reactions that release nuclear energy:
a) Nuclear fusion - is renewable because hydrogen needed for this process is available in plenty in
nature.
b) Nuclear fission - is non-renewable because uranium needed for this process is radioactive and has
only limited existence.
Nuclear isotopes and uses
Naturally occurring radioactive substances have high nucleon number. It is possible to make artificial
radioactive substances by bombarding lighter nuclides with - particles, protons or neutrons. The
radioactive substances produced in this manner are known as radioisotopes.
A nuclide is any species of atom of which each atom has an identical proton number and also an
identical nucleon number. Different nuclides, which have the same proton number (but different
nucleon numbers) are called isotopes (isotopic nuclides). The first radioisotope was an unstable
isotope of phosphorus. It was produced in 1934 by bombarding aluminium with - particles, i.e.,
Phosphorus - 30 was produced, together with a neutron. Notice that on each side of the equation the
sum of the nucleon number is 31 and the sum of the proton number is 15. Phosphorus - 30 decays by
ejecting a positron and has a half-life of about 3 minutes.
The positron has not been mentioned before because it does not occurring natural radioactivity. It has
a mass equal to that of the electron, and a positive electron. It is denoted by
When magnesium is bombarded by neutrons a radioisotope of sodium is formed. The reaction is
The sodium decays with the emission of a beta- particle.
OR
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The important point is that it is now possible to produce any radioisotope. Most of those produced
have short half-life periods. This is very important because the activity and hence the danger from
radioactive emissions does not last very long.
Uses of Radioactive Isotopes
All isotopes of a substance have the same chemical properties and behave in an identical manner. The
advantage of a radioisotope is that its position can be detected very easily by the radiation which it
emits. It has wide applications in various fields like:
1
In Medicine
Radio isotopes are used in detection of diseases and also in radio therapy:

The rays from radium is used in the treatment of skin diseases.

Radiation from Co60 ( - rays) is used to diagnose and treat thyroid disorders.

Radio iodine (I131) is used to diagnose and treat thyroid disorders.

Radio phosphorus (P32) is used in the treatment of leukaemia and tumours.

Radio sodium (Na24) in the form NaCl is used to study circulation of blood.
2
In Agriculture

Radioactive phosphorus (P32) is used in the study of metabolism of plants.

Radioactive sulphur (S35) helps to study advantages and disadvantages of
fungicides.

Pests and insects on crops can be killed by - radiations.,

- rays are used for preservation of milk, potatoes etc.

3.
Yield of crops like carrot, root, apples, grapes can be increased by irradiation with radioisotopes.
In Industry

In is used in the manufacture of paper, plastic and metal sheets. (Used to control the thickness of
the sheets.)

Radioisotopes can be used to estimate the amount of wear in bearings.

Leaks in pipes may be traced by introducing a small quantity of radioisotopes into the fluid in the
pipe.

It is also used to detect the cracks in the welding, casting etc.
Natural Radioactive Isotopes - Radiocarbon and Carbon 14 Dating
There are a small number of radioisotopes of low proton number, which occur naturally. They are
produced by bombardment by radiation from outer space (cosmic rays). The most well known of these
is radioactive carbon-14, which is produced when nitrogen is bombarded by neutrons.
Carbon-14 decays with the emission of alpha and beta - particle, and reverts to nitrogen
In nature, radiocarbon is formed when high energy atomic particles called cosmic rays break down the
atoms in the atmosphere into electrons, protons, neutrons and other particles. Some of the neutrons
strike the nuclei of nitrogen atoms in the atmosphere get converted into radiocarbon atoms.
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Carbon-14 has a long half-life of about 5600 years. It is reasonable to assume that equilibrium has been
reached between the rate at which carbon - 14 forms in the atmosphere and the rate at which it
decays, and that the amount of it in the atmosphere is constant. When plants photosynthesise they
take in CO2 from the air. The carbon 14 atoms in these molecules slowly decay to Nitrogen. Human
beings and other animals take in radiocarbon chiefly from the food provided by plants. Thus, all living
things contain radiocarbon. Plants are utilised as cotton or linen; or might form coal. But whatever
happens the C - 14 in it gradually decays. When the plant or animal dies fresh carbon is no longer taken
in and the C - 14, which is present, decays. Thus, the length of time a specimen has been dead may be
determined by the activity of the C - 14, which remains in it. Carbon - dating has therefore become an
important tool for archaeologists and anthropologists.
Radiocarbon or C -14, is a radioactive isotope of carbon. It is used to determine the age of fossils and
other ancient organic matters.
So by finding out how much carbon 14 is there in an object, we can approximate the age of the sample.
The less the C - 14 compared with carbon 12, the older the sample is. C - 14 dating is widely used in
Archaeology to determine the age of archaeological samples like tools, ornaments, paintings, furniture
etc.
Carbon Dating
The age of fossils of plant or animal origin can be determined by carbon dating technique developed by
Willard Libby in 1949. The radioactive isotope of carbon (6C14) is used to determine the date at which
an animal or plant had died.
Carbon Dating
The method of measuring the age of archaeological materials that contain matter of living origin using
the radioactive isotope 6C14 is called carbon dating. It is continuously formed in the upper strata of the
atmosphere by the action of neutrons in the cosmic rays on 7N14.
By photosynthesis, plants take up CO2 from the atmosphere which contains small amount of radio
isotope 6C14. It is used by plants to build carbohydrates which are then consumed by living animals.
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Because of the natural plant - animal carbon cycle, an equilibrium will be set up and all living matter
will contain a constant equilibrium concentration of C - 14, if the intensity of cosmic rays reaching the
earth remains constant over a long period of time.
Once the plant or animal dies, the process of incorporation of 6C14 stops and 6C14 already present
begins to decay.
(t1/2 = 5770 years)
Thus, by knowing the equilibrium concentration of 6C14 in a living plant and the concentration of 6C14 in
a dead piece of organic matter at a particular time, the age of the material can be determined.
NB:
-
Remaining aspects of the CHM 111 can be assess on the University E-learning Site
Further illustration will be done in the Class
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