Atomic
Structure
Introduction
About the structure of atom a theory was put on by John Dalton in 1808. According to this theory matter was made from small indivisible particles called atoms.
About the structure of atom a theory was put on by John Dalton in 1808. According to this theory matter was made from small indivisible particles called atoms.
But after several experiments many
particles have been discovered with in the atom which are electrons, protons,
neutrons, positrons etc. For the discovery of these fundamental particles the
experiments are as follows.
1. Faraday’s experiment indicates the existence of electron.
1. Faraday’s experiment indicates the existence of electron.
2. Crook’s tube experiment explains
the discovery of electron and proton.
3. Radioactivity also confirms the
presence of electrons and protons.
4. Chadwick’s experiment shows the
presence of neutrons.
The details of these experiments are
given below.
Faraday’s
Experiment
Passage of Electricity Through Solution
In this experiment Faraday passed the
electricity through an electrolytic solution. He observed that when two metal
plates called electrodes are placed in an electrolytic solution and electricity
is passed through his solution the ions present in the solution are moves
towards their respective electrodes. In other words these ions are moves
towards the oppositely charge electrodes to give up their charge and liberated
as a neutral particles.
Faraday also determined the charges of different ions and the amount of elements liberated from the electrolytic solution. Due to this experiment presence of charge particles in the structure of atoms is discovered. The basic unit of electric charge was later named as electron by Stoney in 1891.
Diagram Coming Soon
Faraday also determined the charges of different ions and the amount of elements liberated from the electrolytic solution. Due to this experiment presence of charge particles in the structure of atoms is discovered. The basic unit of electric charge was later named as electron by Stoney in 1891.
Diagram Coming Soon
Crook’s Tube
Or Discharge Tube Experiment
Passage of Electricity Through Gases Under
Low Pressure
Introduction
The first of the subatomic particles to be discovered was electron. The knowledge about the electron was derived as a result of the study of the electric discharge in the discharge tube by J.J. Thomson in 1896. This work was later extended by W. Crooke
The first of the subatomic particles to be discovered was electron. The knowledge about the electron was derived as a result of the study of the electric discharge in the discharge tube by J.J. Thomson in 1896. This work was later extended by W. Crooke
Working of Discharge Tube
When a very high voltage about 10,000
volts is applied between the two electrodes, no electric discharge occurs until
the part of the air has been pumped out of the tube. When the pressure of the
gas inside the tube is less than 1 mm, a dark space appears near the cathode
and thread like lines are observed in the rest of 0.01 mm Hg it fills the whole
tube. The electric discharge passes between the electrodes and the residual gas
in the tube begins to glow. These rays which proceed from the cathode and move
away from it at right angle in straight lines are called cathode rays.
Properties of Cathode Rays
1. They travel in straight lines away
from the cathode and produce shadow of the object placed in their path.
2. The rays carry a negative charge.
3. These rays can also be easily
deflected by an electrostatic field.
4. The rays can exert mechanical
pressure showing that these consist of
material particle which are moving
with kinetic energy.
5. The produce fluorescence when they
strike the glass wall of the discharge tube.
6. Cathode rays produce x-rays when they strike a metallic plate.
6. Cathode rays produce x-rays when they strike a metallic plate.
7. These rays consists of material
particle whose e/m resembles with electron.
8. These rays emerge normally from the
cathode and can be focused by
using a concave cathode.
Positive Rays
In 1890 Goldstein used a discharge
tube with a hole in the cathode. He observed that while cathode rays were
emitting away from the cathode, there were coloured rays produced
simultaneously which passed through the perforated cathode and caused a glow on
the wall opposite to the anode. Thomson studied these rays and showed that they
consisted of particles carrying a positive charge. He called them positive
rays.
Properties of Positive Rays
1. These rays travel in a straight
line in a direction opposite to the cathode.
2. These are deflected by electric as well as magnetic field in the way indicating that they are positively charged.
2. These are deflected by electric as well as magnetic field in the way indicating that they are positively charged.
3. The charge to mass ratio (e/m) of
positive particles varies with the nature of the gas placed in the discharge
tube.
4. Positive rays are produced from
the ionization of gas and not from anode electrode.
5. Positive rays are deflected in electric field. This deflection shows that these are positively charged so these are named as protons.
5. Positive rays are deflected in electric field. This deflection shows that these are positively charged so these are named as protons.
The Information Obtained From Discharge Tube
Experiment
The negatively charge particles
electrons and the positively charge particles protons are the fundamental
particle of every atom.
Radioactivity
In 1895, Henry Becqueral observed
that uranium and its compounds spontaneously emitted certain type of radiation which
affected a
photographic plate in the dark and
were able to penetrate solid matter. He called these rays as radioactivity rays
and a substance which possessed the property of emitting these radioactivity
rays was said to be radioactivity element and the phenomenon was called
radioactivity.
On further investigation by Maric Curic, it was found that the radiation emitted from the element uranium as well as its salts is independent of temperature and the source of the mineral but depend upon the mineral but depend upon the quantity of uranium present e.g. Pitchblende U3O8 was found to be about four times more radioactive than uranium.
On further investigation by Maric Curic, it was found that the radiation emitted from the element uranium as well as its salts is independent of temperature and the source of the mineral but depend upon the mineral but depend upon the quantity of uranium present e.g. Pitchblende U3O8 was found to be about four times more radioactive than uranium.
Radioactive Rays
Soon after the discovery of radium it
was suspected that the rays given out by radium and other radioactive substance
were not of one kind. Rutherford in 1902 devised an ingenious method for
separating these rays from each other by passing them between two oppositely
charged plate. It was observed that the radioactive rays were of three kinds,
the one bending towards the negative plate obviously carrying positive charge
were called α-rays and those deflected to the positive plate and carrying -ve
charge were named as β-rays. The third type gamma rays, pass unaffected and
carry no charge.
Properties of α – RAYS
1. These rays consists of positively
charged particles.
2. These particles are fast moving
helium nuclei.
3. The velocity of α-particles is
approximately equal to 1/10th of the velocity
of light.
4. Being relatively large in size,
the penetrating power of α-rays is very low.
5. They ionize air and their
ionization power is high.
Properties of β – RAYS
1. These rays consists of negatively
charged particles.
2. These particles are fast moving
electron.
3. The velocity of β-particles is
approximately equal to the velocity of light.
4. The penetrating power of β-rays is
much greater than α-rays.
5. These rays ionizes gases to lesser
extent.
Properties of γ – RAYS
1. Gamma rays do not consist of particles.
These are electromagnetic
radiations.
2. They carry no charge so they are not deflected by electric or magnetic
2. They carry no charge so they are not deflected by electric or magnetic
field.
3. Their speed is equal to that of light.
3. Their speed is equal to that of light.
4. These are weak ionizer of gases.
5. Due to high speed and non-material
nature they have great power of
penetration.
Chadwick
Experiment (Discovery of Neutron)
When a light element is bombarded by
α-particles, these α-particles leaves the nucleus in an unstable disturbed
state which on settling down to stable condition sends out radioactivity rays.
The phenomenon is known as “Artificial Radioactivity”.
In 1933, Chadwick identified a new
particle obtained from the bombardment of beryllium by α-particles. It had a
unit mass and carried no charge. It was named “Neutron”.
Spectroscopic
Experiment
After the discovery of fundamental
particles which are electrons, protons & neutron, the next question
concerned with electronic structure of atom.
The electronic structure of the atom was explained by the spectroscopic studies. In this connection Plank’s Quantum theory has great impact on the development of the theory of structure of atom.
The electronic structure of the atom was explained by the spectroscopic studies. In this connection Plank’s Quantum theory has great impact on the development of the theory of structure of atom.
Planck’s
Quantum Theory
In 1900, Max Planck studied the
spectral lines obtained from hot body radiations at different temperatures.
According to him,
When atoms or molecules absorb or emit radiant energy, they do so in separate units of waves called Quanta or Photons.
When atoms or molecules absorb or emit radiant energy, they do so in separate units of waves called Quanta or Photons.
Thus light radiations obtained from
excited atoms consists of a stream of
photons and not continuous waves.
The energy E of a quantum or photon
is given by the relation
E = h v
Where v is the frequency of the emitted
radiation and h the Planck’s
constant. The value of h = 6.62 x
10(-27) erg. sec.
The main point of this theory is that
the amount of energy gained or lost is
quantized which means that energy change
occurs in small packets or
multiple of those packets, hv, 2 hv,
3 hv and so on.
Spectra
A spectrum is energy of waves or
particles spread out according to the increasing or decreasing of some
property. E.g. when a beam of light is allowed to pass through a prism it
splits into seven colours. This phenomenon is called dispersion and the band of
colours is called spectrum. This spectrum is also known as emission spectrum.
Emission spectra are of two types.
1. Continuous Spectrum
2. Line Spectrum
1. Continuous
Spectrum
When a beam of white light is passed
through a prism, different wave lengths are refracted through different angles.
When received on a screen these form a continuous series of colours bands:
violet, indigo, blue, green, yellow and red (VIBGYOR). The colours of this
spectrum are so mixed up that there is no line of demarcation between different
colours. This series of bands that form a continuous rainbow of colours is
called continuous spectrum.
Diagram Coming Soon
2. Line
Spectrum
When light emitted from a gas source
passes through a prism a different kind of spectrum may be obtained.
If the emitted from the discharge
tube is allowed to pass through a prism some discrete sharp lines on a
completely dark back ground are obtained. Such spectrum is known as line
spectrum. In this spectrum each line corresponds to a definite wave length.
Diagram Coming Soon
Identification of Element By Spectrum
Each element produces a
characteristics set of lines, so line spectra came to serve as “finger prints”
for the identification of element. It is possible because same element always
emit the same wave length of radiation. Under normal condition only certain
wave lengths are emitted by an
element.
Rutherford’s
Atomic Model
Evidence for Nucleus and Arrangement of
Particles
Having known that atom contain electrons
and a positive ion, Rutherford and Marsden performed their historic “Alpha
particle scattering experiment” in 1909 to know how and where these fundamental
particles were located in the structure of atom.
Rutherford took a thin of gold with
thickness 0.0004 cm and bombarded in with α-particles. He observed that most of
the α-particles passed straight through the gold foil and thus produced a flash
on the screen behind it. This indicated that old atoms had a structure with
plenty of empty space but some flashes were also seen on portion of the screen.
This showed that gold atoms deflected or scattered α-particles through large
angles so much so that some of these bounced back to the source.
Based on these observations
Rutherford proposed a model of the atom which is known as Rutherford’s atomic
model.
Diagram Coming Soon
Assumption Drawn From the Model
1. Atom has a tiny dense central core
or the nucleus which contains practically the entire mass of the atom leaving
the rest of the atom almost empty.
2. The entire positive charge of the atom is located on the nucleus. While electrons were distributed in vacant space around it.
2. The entire positive charge of the atom is located on the nucleus. While electrons were distributed in vacant space around it.
3. The electrons were moving in
orbits or closed circular paths around the nucleus like planets around the sun.
4. The greater part of the atomic
volume comprises of empty space in which electrons revolve and spin.
Weakness of Rutherford Atomic Model
According to the classical
electromagnetic theory if a charged particle accelerate around an oppositely
charge particle it will radiate energy. If an electron radiates energy, its
speed will decrease and it will go into spiral motion finally falling into the
nucleus. Similarly if an electron moving through orbitals of ever decreasing
radii would give rise to radiations of all possible frequencies. In other words
it would given rise to a continuous spectrum. In actual practice, atom gives
discontinuous spectrum.
X-Rays and
Atomic Number
In 1895, W.Roentgen discovered that
when high energy electrons from cathode collide with the anode in the Crook’s
tube, very penetrating rays are produced. These rays were named as X-rays.
Explanation
When an electron coming from the cathode strike with the anode in the crook’s tube, it can remove an electron from the inner shell of the atom. Due to removal of t his electron the electronic configuration of this ion is unstable and an electron from an orbital of higher energy drops into the inner orbital by emitting energy in form of a photon. This photon corresponds to electromagnetic radiations in the x-rays region.
When an electron coming from the cathode strike with the anode in the crook’s tube, it can remove an electron from the inner shell of the atom. Due to removal of t his electron the electronic configuration of this ion is unstable and an electron from an orbital of higher energy drops into the inner orbital by emitting energy in form of a photon. This photon corresponds to electromagnetic radiations in the x-rays region.
Relationship
Between Wave Length and Nuclear Charge
In 1911, Mosley stablished a
relationship between the wave length and nuclear charge. He found that when
cathode rays struck elements used as anode targets in the discharge tube,
characteristic x-rays were emitted. The wave length of the x-rays emitted
decreases regularly with the increase of atomic mass. On careful examination of
his data Mosely found that the number of positive charges on the nucleus
increases from atom to atom by single electronic unit. He called the number of
positive charges as the atomic number.
Diagram Coming Soon
Bohr’s Theory
Rutherford’s model of atom fails to
explain the stability of atom and appearance of the line spectra. Bohr in 1913
was the first to present a simple model of the atom which explained the
appearance of line spectra.
Some of the postulates of Bohr’s theory are given below.
1. An atom has a number of stable orbits or stationary states in which an
Some of the postulates of Bohr’s theory are given below.
1. An atom has a number of stable orbits or stationary states in which an
electron can reside without emission
or absorption of energy.
2. An electron may pass from one of these non-radiating states to another of lower energy with the emission of radiations whose energy equals the energy difference between the initial and final states.
3. In any of these states the electrons move in a circular path about the
2. An electron may pass from one of these non-radiating states to another of lower energy with the emission of radiations whose energy equals the energy difference between the initial and final states.
3. In any of these states the electrons move in a circular path about the
nucleus.
4. The motion of the electron in these states is governed by the ordinary laws of mechanics and electrostatic provided its angular momentum is an integral multiple of h/2π
4. The motion of the electron in these states is governed by the ordinary laws of mechanics and electrostatic provided its angular momentum is an integral multiple of h/2π
It can be written as
mvr = nh / 2π
Here mvr becomes the angular momentum
of the electron. Thus Bohr’s first condition defining the stationary states
could be stated as
“Only those orbits were possible in
which the angular momentum of the electrons would be an integral multiple of
h/2π”. These stationary states correspond to energy levels in the atom.
Calculation of Radius of Orbits
Consider an electrons of charge e
revolving.
Atomic number and e the charge on a
proton.
Let m be the mass of the electro, r
the radius of the orbit and v the
tangential velocity of the revolving
electron.
The electrostatic force of attraction
between the nucleus and the electron
according to Coulomb’s law
= Z e x e / r2
Diagram Coming Soon
The centrifugal force acting on the
electron.
= mv2 / r
Bohr assumed that these two opposing
forces must be balanced each other exactly to keep the electron in an orbit.
Therefore
Ze2 / r2 = m v2 / r
Ze2 / r2 = m v2 / r
Multiply both sides by r
r x Ze2 / r2 = r x m v2 / r
Ze2 / r = m v2
OR
r = Ze2 / m v2 ……………… (1)
The Bohr’s postulate states that only
those orbits are possible in which
mvr = nh / 2π
Therefore,
V = nh / 2πmr
Substituting the value of V in eq (1)
r = Ze2 / m(nh/2πmr)2
or
r = Ze2 x 4π2 mr2/n2h2
or
1/r = 4π2mZe2/n2h2
Cr
r = n2h2 / 4π2mZe2 …………… (2)
This equation gives the radii of all
the possible stationary states. The values of constants present in this equation
are as follows.
H = 6.625 x 10(-27) ergs sec OR 6.625
x 10(-37) J.s
Me = 9.11 x 10(-28) gm OR 9.11 x
10(-31) kg
E = 4.802 x 10(-10) e.s.u OR 1.601 x
10(-19) C
By substituting these values we get
for first shell of H atom
r = 0.529 x 10(-8) m OR 0.529
The above equation may also be
written as
r = n2 (h2 / 4π2mZe2) x n2 a0
……………….. (3)
For the first orbit n = 1 and r =
0.529. This is the value of the terms in the
brackets sometimes written as a0
called Bohr’s Radius. For the second shell
n = 2 and for 3rd orbit n = 3 and so
on.
Hydrogen Atom
Spectrum
Balmer Series
The simplest element is hydrogen
which contain only one electron in its valence shell.
Balmer in 1885 studied the spectrum
of hydrogen. For this purpose he used hydrogen gas in the discharge tube. Balmer
observed that hydrogen atom spectrum consisted of a series of lines called
Balmer Series. Balmer determined the wave number of each of the lines in the
series and found that the series could be derived by a simple formula.
Lyman Series
Lyman series is obtained when the
electron returns to the ground state i.e. n = 1 from higher energy level n(2) =
2, 3, 4, 5, etc. This series of lines belongs to the ultraviolet region of
spectrum.
Paschen Series
Paschen series is obtained when the
electron returns to the 3rd shell i.e. n = 3 from the higher energy levels n2 =
4, 5, 6 etc. This series belongs to infrared region.
Bracket Series
This series is obtained when an
electron jumps from higher energy levels to 4th energy level.
Heisenberg
Uncertainty Principle
According to Bohr’s theory an
electron was considered to be a particle but electron also behaves as a wave
according to be Broglie.
Due to this dual nature of electron
in 1925 Heisenberg gave a principle known as Heisenberg Uncertainty Principle
which is stated as,
It is impossible to calculate the
position and momentum of a moving
electron simultaneously.
It means that if one was known
exactly it would be impossible to known the other exactly. Therefore if the
uncertainty in the determination of momentum is Δpx and the uncertainty in
position is Δx then according to this principle the product of these two
uncertainties may written as
Δpx . Δx ≈ h
Δpx . Δx ≈ h
So if one of these uncertainties is
known exactly then the uncertainty in its determination is zero and the other
uncertainty will become infinite which is according to the principle.
Energy Levels
and Sub-Levels
According to Bohr’s atomic theory,
electrons are revolving around the nucleus in circular orbits which are present
at definite distance from the nucleus. These orbits are associated with
definite energy of the electron increasing outwards from the nucleus, so these
orbits are referred as Energy Levels or Shells.
These shells or energy levels are
designated as 1, 2, 3, 4 etc K, L, M, N etc.
The spectral lines which correspond to the transition of an electron from one energy level to another consists of several separate close lying lines as doublets, triplets and so on. It indicates that some of the electrons of the given energy level have different energies or the electrons belonging to same energy level may differ in their energy. So the energy levels are accordingly divided into sub energy levels which are denoted by letters s, p, f (sharp, principle, diffuse & fundamental).
The spectral lines which correspond to the transition of an electron from one energy level to another consists of several separate close lying lines as doublets, triplets and so on. It indicates that some of the electrons of the given energy level have different energies or the electrons belonging to same energy level may differ in their energy. So the energy levels are accordingly divided into sub energy levels which are denoted by letters s, p, f (sharp, principle, diffuse & fundamental).
The number of sub levels in a given
energy level or shell is equal to its value of n.
e.g. in third shell where n = 3 three
sub levels s, p, d are possible.
Quantum
Numbers
There are four quantum numbers which
describe the electron in an atom.
1. Principle
Quantum Number
It is represented by “n” which
describe the size of orbital or energy level.
The energy level K, L, M, N, O etc correspond to n = 1, 2, 3, 4, 5 etc.
If
The energy level K, L, M, N, O etc correspond to n = 1, 2, 3, 4, 5 etc.
If
n = 1 the electron is in K shell
n = 2 the electron is in L shell
n = 3 the electron is in M shell
2. Azimuthal Quantum Number
This quantum number is represented by
“l” which describes the shape of the orbit. The value of Azimuthal Quantum
number may be calculated by a relation.
l = 0 —-> n – 1
l = 0 —-> n – 1
So for different shell the value of l
are as
n = 1 K Shell l = 0
n = 2 L Shell l = 0, 1
n = 3 M Shell l = 0, 1, 2
n = 4 N Shell l = 0, 1, 2, 3
when l = 0 the orbit is s
when l = 1 the orbit is p
when l = 2 the orbit is d
when l = 3 the orbit is f
3. Magnetic Quantum Number
It is represented by “m” and explains
the magnetic properties of an electron. The value of m depends upon the value
of l. It is given by
m = + l —-> 0 —-> l
when l = 1, m has three values (+1,
0, -1) which corresponds to p orbital.
Similarly when l = 2, m has five
values which corresponds to d orbital.
4. Spin Quantum Number
It is represented by “s” which
represents spin of a moving electron. This spin may be either clockwise or
anticlockwise so the values for s may be +1/2 or -1/2.
Pauli’s
Exclusion Principle
According to this principle
No two electrons in the same atom can
have the same four quantum number.
Consider an electron is present in 1s orbital. For this electron n = 1, l = 0, m = 0. Suppose the spin of this electron is s = +1/2 which will be indicated by an upward arrow ↑. Now if another electron is put in the same orbital (1s) for that electron n = 1, l = 0, m = 0. It can occupy this orbital only if the direction of its spin is opposite to that of the first electron so s = -1/2 which is symbolized by downward arrow ↓. From this example, we can observe the application of Pauli’s exclusion principle on the electronic structure of atom.
Consider an electron is present in 1s orbital. For this electron n = 1, l = 0, m = 0. Suppose the spin of this electron is s = +1/2 which will be indicated by an upward arrow ↑. Now if another electron is put in the same orbital (1s) for that electron n = 1, l = 0, m = 0. It can occupy this orbital only if the direction of its spin is opposite to that of the first electron so s = -1/2 which is symbolized by downward arrow ↓. From this example, we can observe the application of Pauli’s exclusion principle on the electronic structure of atom.
Electronic
Configuration
The distribution of electrons in the
available orbitals is proceeded according to these rules.
1. Pauli Exclusion Principle
2. Aufbau Principle
3. (n + l) Rule
4. Hund’s Rule
The detail of these rules and
principles is given below.
1. Aufbau Principle
It is states as
The orbitals are filled up with
electrons in the increasing order of their
energy.
It means that the orbitals are fulled
with the electrons according to their energy level. The orbitals of minimum
energy are filled up first and after it the orbitals of higher energy are
filled.
2. Hund’s Rule
If orbitals of equal energy are
provided to electron then electron will go to different orbitals and having
their parallel spin.
In other words we can say that electrons are distributed among the orbitals of a sub shell in such a way as to give the maximum number of unpaired electrons and have the same direction of spin.
In other words we can say that electrons are distributed among the orbitals of a sub shell in such a way as to give the maximum number of unpaired electrons and have the same direction of spin.
3. (n + l) Rule
According to this rule
The orbital with the lowest value of
(n + l) fills first but when the two orbitals have the same value of (n + l)
the orbital with the lower value of n fills first.
For the electronic configuration the
order of the orbital is as follows.
1s, 2s, 2p, 3s, 4s, 3d, 4p, 5s, 4d, 5p, 6s etc.
1s, 2s, 2p, 3s, 4s, 3d, 4p, 5s, 4d, 5p, 6s etc.
Atomic Radius
For homonuclear diatomic molecules
the atomic radius may be defined as
The half of the distance between the two nuclei present in a homonuclear diatomic molecules is called atomic radius.
The half of the distance between the two nuclei present in a homonuclear diatomic molecules is called atomic radius.
It may be shown as
In case of hetronuclear molecular
like AB, the bond length is calculated
which is (rA + rB) and if radii of any
one is known the other can be
calculated.
For the elements present in periodic
table the atomic radius decreases from left to right due to the more attraction
on the valence shell but it increases down the group with the increase of
number of shells.
Ionic Radius
Ionic radius is defined as
The distance between nucleus of an
ion and the point up to which nucleus has influence of its electron cloud.
When an electron is removed from a
neutral atom the atom is left with an excess of positive charge called a cation
e.g
Na —-> Na+ + c-
But when an electron is added in a
neutral atom a negative ion or anion is formed.
Cl + e- —-> Cl-
As the atomic radius, the ionic radii
are known from x-ray analysis. The value of ionic radius depends upon the ions
that surround it.
Ionic radii of cations have smaller radii than the neutral atom because when an electron is removed. The effective charge on the nucleus increases and pulls the remaining electrons with a greater force.
Ionic radii of anions have a large radii than the neutral atom because an excess of negative charge results in greater electron repulsion.
Radius of Na atom = 1.57
Radius of Na+ atom = 0.95 (smaller than neutral atom)
Ionic radii of cations have smaller radii than the neutral atom because when an electron is removed. The effective charge on the nucleus increases and pulls the remaining electrons with a greater force.
Ionic radii of anions have a large radii than the neutral atom because an excess of negative charge results in greater electron repulsion.
Radius of Na atom = 1.57
Radius of Na+ atom = 0.95 (smaller than neutral atom)
Radius of Cl atom = 0.99
Radius of Cl- atom = 1.81 (larger
than neutral atom)
Ionization
Potential
Definition
The amount of energy required to remove most loosely bounded electron from the outermost shell of an atom in its gaseous state is called is called ionization potential energy.
The amount of energy required to remove most loosely bounded electron from the outermost shell of an atom in its gaseous state is called is called ionization potential energy.
It is represented as
M(gas) —-> M+(gas) + e- ………………. ΔE
= I.P
The energy required to remove first
electron is called first I.P. The energy required to remove 2nd or 3rd electron
is called 2nd I.P or 3rd I.P
M(gas) —-> M+(gas) + e- ………………. ΔE
= 1st I.P
M+(gas) —-> M++(gas) + e- …………….ΔE
= 2nd I.P
M++(gas) —-> M+++(gas) + e- …………
ΔE = 3rd I.P
The units of I.P is kilo-Joule per
mole.
Factors on which I.P Depends
1. Size of
the Atom
If the size of an atom is bigger the
I.P of the atom is low, but if the size of the atom is small then the I.P will
be high, due to fact if we move down the group in the periodic table. The I.P
value decreases down the group.
2. Magnitude
of Nuclear Charge
If the nuclear charge of atom is
greater than the force of attraction on the valence electron is also greater so
the I.P value for the atom is high therefore as we move from left to right in
the periodic table the I.P is increased.
3. Screening
Effect
The shell present between the nucleus
and valence electrons also decreases the force of attraction due to which I.P
will be low for such elements.
Electron Affinity
Definition
The amount of energy liberated by an atom when an electron is added in it is called electron affinity.
The amount of energy liberated by an atom when an electron is added in it is called electron affinity.
It shows that this process is an
exothermic change which is represented as
Cl + e- —-> Cl- ………… ΔH = -348 kJ / mole
Cl + e- —-> Cl- ………… ΔH = -348 kJ / mole
Factors on which Electron Affinity Depends
1. Size of
the Atom
If the size of atom is small, the
force of attraction from the nucleus on the valence electron will be high and
hence the E.A for the element will also be high but if the size of the atoms is
larger the E.A for these atoms will be low.
2. Magnitude
of the Nuclear Charge
Due to greater nuclear charge the
force of attraction on the added electron is greater so the E.A of the atom is
also high.
3. Electronic
Configuration
The atoms with the stable
configuration has no tendency to gain an electron so the E.A of such elements
is zero. The stable configuration may exist in the following cases.
1. Inert gas configuration
2. Fully filled orbital
3. Half filled orbital
Electro
negativity
Definition
The force of attraction by which an atom attract a shared pair of
electrons is called electro negativity.
Application of Electro negativity
1. Nature of
Chemical Bond
If the difference of electro
negativity between the two combining atoms is more than 1.7 eV, the nature of
the bond between these atoms is ionic but if the difference of electro
negativity is less than 1.7 eV then the bond will be covalent.
2. Metallic
Character
If an element possesses high electro
negativity value then this element is a non-metal but if an element exists with
less electro negativity, it will be a metal.
Factors for Electro negativity
1. Size of
the Atom
If the size of the atom is greater
the electro negativity of the atom is low due to the large distance between the
nucleus and valence electron.
2. Number of
Valence Electrons
If the electrons present in the
valence shell are greater in number, the electro negativity of the element is
high.
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