Atomic structure

Atomic
Structure
By- AtandritChatterjee

In This Presentation
 Introduction
 Electron
 Millikan’s oil drop experiment
 Proton
 Thomson's Atomic Model
 Rutherford’s Experiment- Discovery of Nucleus
 Rutherford’s Nuclear Model of Atom
 Planck’s quantum theory
 Bohr’s model of an Atom
 Neutron
 Atomic Number and Mass Number
 Isotopes, Isobars and Isotones
 Electronic Configuration
 Valency

Introduction
The concept of an atom is originated from
Greek philosophers like Democritus and
John
Dalton.
Democritus studied the nature of matter
and the constituents of all the substances.
In 1808 John Dalton put forward atomic
theory to explain the laws of chemical
combination. According to him, an atom is
the smallest unit of matter which takes part
in
a chemical reaction. He considered that
atoms are indivisible particles
At the end of 18th and early 20th centuries
modern concept an atom developed by
scientists like J.J Thomson, Goldstein,

Electron
Discovery
Towards the end of the 19th century J.J.Thomson (1856-
1940) was studying electric discharges at the well-known
Cavendish laboratory in Cambridge, England. Several people
had been studying the intriguing effects in electric
discharge tubes before him. Spectacular glows could be
observed when a high voltage was applied in a gas volume at
low pressure. It was known that the discharge and the glow
in the gas were due to something coming from the cathode,
the negative pole of the applied high voltage. Thomson
made a series of experiments to study the properties of
the rays coming from the cathode. He observed that the
cathode rays were deflected by both electric and magnetic
fields - they were obviously electrically charged.

Next, J.J. Thomson put a negatively charged metal
plate to determine if the charge carried by the
cathode rays was negative or positive. When the
charged metal plates were introduced he found that
the cathode rays bent toward the positive plate and
away from the negative plate. This showed that the
charge carried by the cathode rays was negative. The
particle that J.J.Thomson discovered in 1897, the
electron, is a constituent of all the matter we are
surrounded by. He received the Nobel Prize in 1906
for the discovery of the electron, the first elementary
particle.

Properties of Cathode Rays
1.They travel in straight lines from the
negative pole (cathode).
2.They produce fluorescence in the glass walls
of the discharge tube.
3.They cast shadows if some target is placed
in their path.
4.They can produce mechanical motion, e.g.,
they cause a light pedal wheel placed in their
path to rotate.
5.They possess heating effect and can heat
thin metal filaments to incandescence.
6.They are deflected from their rectilinear
path by electrostatic and magnetic fields and
behave in the manner of a stream of
negatively charged particles.
7.They can impart negative charge to objects
in their paths.
8.They can cause ionization in gases.

Millikan’s Oil Drop Experiment
In 1909, Robert Millikan
and Harvey Fletcher
conducted the oil drop
experiment to determine
the charge of an electron.
They suspended tiny
charged droplets of oil
between two metal
electrodes by balancing
downward gravitational
force with upward drag and
electric forces. The density
of the oil was known, so
Millikan and Fletcher could
determine the droplets'
masses from their observed
radii (since from the radii
they could calculate the

Using the known electric field and the values
of gravity and mass, Millikan and Fletcher
determined the charge on oil droplets in
mechanical equilibrium. By repeating the
experiment, they confirmed that the charges
were all multiples of some fundamental value.
They calculated this value to be 1.5924 × 10−19
Coulombs (C), which is within 1% of the
currently accepted value of 1.602176487 ×
10−19 C. They proposed that this was the
charge of a single electron.
Q is the charge of an electron, E is the
electric field, m is mass of the droplet, and g
is gravity.
Q⋅ E = m⋅ g
Q = m⋅ g /E
One can see how Millikan calculated the
charge of an electron. Millikan found that all

Proton
Discovery
• Canal Ray experiment is the experiment
performed by German scientist Eugen
Goldstein in 1886 that led to the
discovery of proton. The discovery of
proton which happened after the
discovery of electron further
strengthened the structure of atom. In
the experiment, Goldstein applied high
voltage across a discharge tube which had
a perforated cathode. A faint luminous
ray was seen extending from the holes in
the back of the cathode.
• The apparatus of the experiment
incorporates same apparatus as of

containing two pieces of metals at the different end which
acts as electrode. The two metal pieces are connected
with external voltage. The pressure of gas inside the tube
is lowered by evacuating the air.
• Procedure of the experiment
• Apparatus is setup by providing a high voltage source and
evacuating the air to maintain low pressure inside the
tube.
• High voltage is passed to the two metal pieces so as to
ionize air and make it conductor of electricity.
• The electricity starts flowing as circuit was complete.
• When the voltage was increased to several thousand volts,
a faint luminous ray was seen extending from the holes in
the back of the cathode.
• These rays were moving in the opposite direction of
cathode rays and was named canal rays.

Properties of Canal Rays
1. They are the streams of positive ions of the gas
enclosed in the discharge tube. The mass of each
ion is nearly equal to the mass of the atom.
2. They are deflected by electric and magnetic
fields. Their deflection is opposite to that of
cathode rays.
3. They travel in straight lines.
4. The velocity of canal rays is much smaller than
the velocity of cathode rays.
5. They affect photographic plates.
6. These rays can produce fluorescence.
7. They ionize the gas through which they pass.

Thomson’s Atomic Model
J. J. Thomson, who discovered the electron in 1897,
proposed the plum pudding model, also known as
Thomson's atomic model, of the atom in 1904
before the discovery of the atomic nucleus in
order to include the electron in the atomic model.
In Thomson's model, the atom is composed of
electrons surrounded by a soup of positive charge
to balance the electrons' negative charges, like
negatively charged "plums" surrounded by
positively charged "pudding" . The electrons (as
we know them today) were thought to be
positioned throughout the atom in rotating rings.
In this model the atom was also sometimes
described to have a "cloud" of positive charge.
Thomson’s model could successfully explain the
electrical neutrality of atom. However, it failed to
explain how the positively charged particles are
shielded from the negatively charged electrons
Thomson’s Atomic Model

Rutherford’s Experiment-Discovery Of
Nucleus
Ernest Rutherford was interested in
knowing how the electrons are arranged
within an atom. Rutherford designed an
experiment which led to the discovery of
nucleus. In this experiment, fast moving
alpha (α)-particles were made to fall on a
thin gold foil. He selected a gold foil
because he wanted as thin a layer as
possible. This gold foil was about 1000
atoms thick. α-particles are doubly-
charged helium ions. Since they have a
mass of 4µ, the fast-moving α-particles
have a considerable amount of energy. It
was expected that α-particles would be

Since the α-particles were much heavier
than the protons, he did not expect to
see large deflections. But, the α-
particle scattering experiment gave
totally unexpected results .
• Observations of Rutherford's
scattering experiment:
• Most of the fast moving α-particles
passed straight through the gold foil.
• Some of the α-particles were deflected
by the foil by small angles.
• Surprisingly one out of every 12,000
alpha particles appeared to rebound.
Conclusion of Rutherford's
scattering experiment:
• Most of the space inside the atom is

through the gold foil without getting
deflected.
• Very few particles were deflected from
their path, indicating that the positive
charge of the atom occupies very little
space.
• A very small fraction of α-particles were
deflected by very large angles, indicating
that all the positive charge and mass of
the gold atom were concentrated in a very
small volume at the centre of the atom
and this centre of positive charge is
known as the nucleus of an atom.
From the data he also calculated that the
radius of the nucleus is about 105 times
less than the radius of the atom.

Rutherford’s Nuclear Model of
Atom
On the basis of his experiment, Rutherford
put forward the model of an atom, which
had the following features:
• There is a positively charged centre in an
atom called the nucleus. Nearly all the
mass of an atom resides in the nucleus.
• The electrons revolve around the nucleus
in well-defined orbits.
• The size of the nucleus is very small as
compared to the size of the atom.
Rutherford's alpha particle scattering
experiment shows the presence of nucleus
in the atom. It also gives the following
important information about the nucleus

• Nucleus of an atom is positively charged and
contains all the protons and neutrons.
• Nucleus of an atom is very dense and hard.
• Nucleus of an atom is very small as compared to
the size of the atom as a whole.
• An atom is electrically neutral. Therefore, the
number of electrons in an atom is equal to the
number of protons in it.
• The electrostatic attraction between protons and
electrons holds the atom together.
Rutherford model of atom is also called Nuclear
model of atom.

• Drawback
It does not obey the Maxwell theory of
electrodynamics, according to it “A small
charged particle moving around an
oppositely charged centre continuously
loses its energy”. If an electron does so,
it should also continuously lose its
energy and should set up spiral motion
ultimately falling into the nucleus. If
this were so, the atom should be highly
unstable and hence matter would not
exist in the form that we know. We
know that atoms are quite stable.

Planck’sQuantum Theory
In 1901, Max Planck concluded from his
experimental observation that energy can be
absorbed or radiated by a body in the form of
small packets of energy called quanta (sing.-
quantum) . In case of light, a quantum of energy
is called a 'photon'. Planck postulated that the
energy of a particular quantum of radiant energy
could be described by the equation:-
E œ v
E = h.v
Where h= Planck’s constant=6.626×10-34J⋅s
V = frequency of the radiation
Now,
c = v.w
Where c = speed of light = 3.0 x 108 m/sec
W = wavelength in meters

v = c/w
So, from the above equation, we
get:-
E = h.c/w
The total amount of energy
emitted or absorbed by a
body will be some whole
number quanta. Hence
E = nh.v
Where n = number of photons

Bohr’s ModelOf An Atom
In 1913 Bohr proposed his quantized shell model of
the atom to explain how electrons can have stable
orbits around the nucleus. The motion of the
electrons in the Rutherford model was unstable
because, according to classical mechanics and
electromagnetic theory, any charged particle
moving on a curved path emits electromagnetic
radiation; thus, the electrons would lose energy
and spiral into the nucleus. To remedy the
stability problem, Bohr modified the Rutherford
model by requiring that the electrons move in
orbits or shells. Each orbit or shell is associated
with a definite amount of energy. Hence these
are also called energy levels and are designated
as K, L, M, N respectively. The energy of an
electron depends on the size of the orbit and is
lower for smaller orbits. As long as the electron

also called stationary orbits and the electrons are
said to be in stationary energy state. Radiation
can occur only when the electron jumps from one
orbit to another. If an electron jumps from a
lower energy level to a higher energy level, it
absorbs energy. Whereas when it jumps from
higher to lower energy level, energy is emitted
in the form of electromagnetic radiations. The
energy emitted or absorbed (∆E) is an integral
multiple of h.v. The atom will be completely
stable in the state with the smallest orbit, since
there is no orbit of lower energy into which the
electron can jump.

The electron can revolve only in the orbit
in which the angular momentum of the
electron (m.v.r) is quantised, i.e. m.v.r is
a whole number multiple of h/2π. This is
known as principle of quantization of
angular momentum. The angular
momentum is written as-
m.v.r = nh/2π
Where, n is an integer (n = 1,2,3,4.........)
and is called principal quantum number
m = mass of the electron
v = velocity of the electron in its orbit
r = distance of the electron from the
nucleus
h= Planck’s constant

By applying the concept of quantization of energy,
Bohr calculated the radii and energy of the nth
orbit of single electron species:-
rn = (n2 .h2)/(4π2.m.e2Z)
=( 0.529.n2)/Z Å
En=(-2π2.m.e4.Z2)/(n2h2)
=(- 13.6.Z2)/n2 eV
Where,
m = mass of the electron
e = charge of the electron
Z = atomic number or nuclear charge
h= Planck’s constant
Velocity of electron in an orbit:-
vn=Z.e2/2nh
With the help of these expressions, Bohr gave a
satisfactory explanation for the spectra of
hydrogen and hydrogen like species.

• Drawback
• It violates the Heisenberg Uncertainty
Principle . The Bohr atomic model theory
considers electrons to have both a known
radius and orbit i.e. known position and
momentum at the same time, which is
impossible according to Heisenberg.
• The Bohr atomic model theory made
correct predictions for smaller sized
atoms like hydrogen, but poor spectral
predictions are obtained when larger
atoms are considered.
• It failed to explain the Zeeman effect
when the spectral line is split into several
components in the presence of a magnetic
field.
• It failed to explain the Stark effect when

Neutron
Discovery
In 1930 it was discovered that Beryllium,
when bombarded by alpha particles,
emitted a very energetic stream of
radiation. This stream was originally
thought to be gamma radiation. However,
further investigations into the
properties of the radiation revealed
contradictory results. Like gamma rays,
these rays were extremely penetrating
and since they were not deflected upon
passing through a magnetic field,
neutral. However, unlike gamma rays,
these rays did not discharge charged
electroscopes (the photoelectric
effect). Irene Curie and her husband

paraffin, protons were knocked loose which could
be easily detected by a Geiger counter.
In 1932, Chadwick proposed that this particle was
Rutherford's neutron. In 1935, he was awarded
the Nobel Prize for his discovery. Using
kinematics, Chadwick was able to determine the
velocity of the protons. Then through
conservation of momentum techniques, he was
able to determine that the mass of the neutral
radiation was almost exactly the same as that of
a proton. This is Chadwick's equation:-

• Properties of Neutron
1. The neutron is a charge-less
subatomic particle
2. Its mass is almost equal to that
of a proton.
3. Treating it semi classically, it has
an energy given by:-
E=m.v2/2

Atomic Number and MassNumber
• Atomic Number
The number of protons in the nucleus of an
atom is known as atomic number. It is
characteristic of a chemical element and
determines its place in the periodic
table. Its symbol is Z. Ex:- One atom of
sodium has 11 protons in it, so the atomic
number of sodium is 11.
• Mass Number
The total number of protons and neutrons
in a nucleus is known as mass number. Its
symbol is A. Therefore,
Mass Number = Atomic number + Number
of neutrons

Isotopes, Isobars andIsotones
• Isotope
Isotope is each of two or more forms of the same
element that contain equal numbers of protons
but different numbers of neutrons in their
nuclei, and hence differ in relative atomic mass
but not in chemical properties; in particular, a
radioactive form of an element.
• Isobar
Isobar is each of two or more isotopes of different
elements, with the same atomic weight.
• Isotone
Two nuclides are isotones if they have the same
neutron number N, but different proton number
Z. Ex:- Boron-12 and carbon-13 nuclei both
contain 7 neutrons, and so are isotones .

Electronic Configuration
Electron configuration was first conceived
of under the Bohr model of the atom. An
electron shell is the set of allowed states
that share the same principal quantum
number, n that electrons may occupy. An
atom's nth electron shell can
accommodate 2n2 electrons.
K
L
M
N
O
Shell
2
8
18
32
50
2n2

The factor of two arises because the allowed states are
doubled due to electron spin—each atomic orbital admits
up to two otherwise identical electrons with opposite
spin, one with a spin +1/2 (usually denoted by an up-
arrow) and one with a spin −1/2 (with a down-arrow).
A subshell is the set of states defined by a common
azimuthal quantum number, ℓ, within a shell. The values ℓ
= 0, 1, 2, 3 correspond to the s, p, d, and f labels,
respectively. For example the 3d subshell has n = 3 and ℓ
= 2. The maximum number of electrons that can be
placed in a subshell is given by 2(2ℓ+1). This gives two
electrons in an s subshell, six electrons in a p subshell,
ten electrons in a d subshell and fourteen electrons in an
f subshell. It is arranged in the order shown in the
picture. Ex:- Argon, Ar (18) - 1s22s22p63s23p6.

The numbers of electrons that
can occupy each shell and
each subshell arise from the
equations of quantum
mechanics, in particular the
Pauli Exclusion Principle, which
states that no two electrons
in the same atom can have the
same values of the four
quantum numbers.

Valency
When atoms of one element combine
with the atoms of another element
to form formula units, they do so in
fixed numbers depending upon the
capacities of the atoms to form
bonds. Valency of an element is a
measure of the combining capacity
of its atom to form chemical bonds.
Ex:- Valency of sodium (Na) is one
as it doesn’t have its outermost
shell completely filled and has only
one electron.

Atomic structure

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