Quality Education At
Every level
Chapter
#03
Atomic Structure
Prepared
by:
Lecturer.
Syed Fayyaz Hussain
Atomic Structure
Introduction:
Atom
is a smallest particle of an element which can take part in a chemical
reaction. Atom consist of subatomic particles such as electron, proton and
neutron etc. evidence for the presence of electron, proton and neutrons in the
atom is derived through many experiment.
FARADAY’S
EXPERIMENT:
The
substance which can dissociate into their ions can conduct electricity is
called electrolyte. This phenomenon was studied in great detail
By Faraday’s .He
observed that,
(a)
Electrolyte gave cations and anions on dissociate.
(b)
Cations migrate towards cathode.
(c)
Anions migrate toward anode.
(d)
These ions neutralized their charges by gaining or
losing electrons.
(e)
Finally, they get deposited at the electrodes.
CONCLUSIONS:
Accurate
determination showed that 96490 coulombs of electricity liberated one gram
equivalent of the substance at the electrode. Hence there is some elementary
unit of electric charges associated with these ions and it was found that its
value is
ē = 1.60 x 10-19
coulombs.
The ions were observed
to carry some integral multiply of this charge. The basic unit of electric
charge was letter named by stony as “ELECTRONIC
CHARGE”.
Discharge Tube Experiment:
All
gases air are bad conductors of electricity at normal pressure. But on high
voltage and low pressure, they become good conductor. The conduction of
electricity was first studied by WILLIAM COOKS. The apparatus used for this
purpose is called “DISCHARGE TUBE”.
Construction:
Discharge tube consists
of a cylindrical glass tube closed at both ends and fitted with two metallic
electrodes. These electrodes are connected to the positive and negative
terminals of a battery.
The discharge tube also
posses a side tube which is connected to a vacuum pump in order to remove the
gas or air from it. The removal of air or gas reduced pressure inside the tube.
Working:
When a
high voltage is applied to a discharge tube at normal pressure, no phenomenon
is observed. But when the vacuum pump is started and most of the gas inside the
discharge tube is pumped out and the pressure is reduced to 1 torr, the tube soon begins to emit a
soft glow. This gives an indication that the gas in the tube begins to conduct
electricity. As the pressure is further reduced, the glowing rays move towards
anode. Since these rays are produced at the surface of the cathode therefore
these are called “CATHODE
RAYS”.
At still lower pressure
about 0.01 torr, the flow from
within the tube fades out and dark spaces appears in the discharge tube. At
this stage the glass tube at the anode end begins to emit a greenish
fluorescence.
Properties of Cathode Rays:
The various experiments
performed by different scientists indicate that the cathode rays possess the
following characteristics:
1. These rays travel in a straight lines
perpendicular to the cathode surface.
2. These rays produce sharp shadows if an
opaque object is placed in their path.
3. These rays emerge from cathode and can
be focused by using a concave cathode.
4. These rays can penetrate small
thickness of meter e.g. Aluminum or Gold foils
without producing any perforation in it.
5. These rays can exert mechanical
pressure, showing they possess Kinetic Energy.
6. These rays are deflected in a magnetic
field. This behavior indicates that they
consist
of charged particles.
7. These rays are deflected towards
positively charge plate in an electrical field,
indicating that they consist of negatively charged
particles called ELECTRONS.
8. The charge to mass ratio (e/m) of these
rays is 1.76 x 1011 coulomb / Kg.
9. The charge and mass of these rays are
1.6 x 10-19 coulomb & 9.1 x 10-31 Kg
respectively.
10. The rays were seen neither to depend on
the material of which the electrodes were
made nor upon the gas which is filled in the tube.
Conclusion:
On the
basis of these properties, it was concluded that cathode rays are negatively
charged particles called “ELECTRONS”.
Since nature of the cathode rays does not change with the nature of the gas and
the cathode used in the tube, hence it can be safely that electrons are the fundamental
particles of all atoms.
DISCOVERY
OF PROTON:
During
the study of the passage of electricity through gases at low pressure, it was
observed by Goldstein that, in
addition to cathode rays there are also other rays traveling in opposite
direction to the cathode rays, and if the cathode is perforated, then these
rays can pass through the perforations (canals) and accumulate at the behind of
cathode. He named these rays as “Positive
Rays” or “Canal Rays”.
Properties of Positive Rays:
1.These rays travel in
a straight line.
2.These rays are deflected by electric and
magnetic fields opposite to the cathode rays.
Therefore
these are positively charged rays.
3.These rays are not
emitted from anode but are produced from the
ionization of gas as a
result of bombardment of electrons.
4.Their
positive charge was found to be equal to that of electron or simple multiple of it.
5.These
particles were found to be much heavier than electrons. The mass of these
particles were found to depend on the kind
of gas taken in the discharge tube
but it is
never less than that of an atom or
hydrogen.
6. Positive
rays unlike cathode rays have different value of e/m depending on the gas
present in the discharge tube.
Conclusion:
The charge
to mass ratio (e/m) of these rays
depends upon the natural of the gas in the discharge tube. The highest (e/m) ratio is obtained in case of
hydrogen gas. Therefore hydrogen without electron (H+) is the
smallest positive particles of matter. This particle was given the name “PROTON” by Goldstein in 1986. Later on Rutherford
showed that proton like electron is also a fundamental particle of matter.
Discovery of Neutron:
The
total mass of the protons and electrons in each atom are not sufficient to
account for the atomic masses of the different elements. For this reason,
scientists believed that there must be another heavy particle inside the atom.
In 1932, James Chadwick discovered the third
particle in the nucleus of atom by means of artificial radioactivity.
Chadwick bombarded the
nucleus of Beryllium atoms with a particles and found that it gave highly penetrating
radiation. Chadwick put forward the suggestion that these penetrating radiation
were due to material particles with mass comparable with that of an atom of
hydrogen but carrying no charge. These particles are called “Neutrons” The neutrons must have come
out from atoms on disintegration of the bombarded element. This is indicated by
the equation.
4Be9
+ 2He4 6
C12 + 0n1
Radioactivity:
The
natural phenomenon in which certain elements emits invisible radiations is
called “Radioactivity” and the
substances which give out these radiations are called radioactive substance.
If the radioactive rays
are passed through a strong magnetic field as shown in the diagram, they split
into three kinds of rays, which are as follows:
(a) Alpha a rays
(b) Beta b rays
(c) Gamma g rays
Properties of a -Rays:
1. These
rays consist of stream of doubly positively charged particles which are four
times as heavy as hydrogen atoms. So they
are He nuclei or Helium ion (He++).
2. The velocity of these particles is 
of the velocity of
light.
of the velocity of
light.
3. They are deflected by electrical and
magnetic field and moves towards –ve plate.
4. They produce intense ionization of the
gas through which they pass.
5. Due
to heavy mass and low velocity, these particles have low penetration power. Therefore, these
particles can hardly pass through thin layers of solids or a few centimeters of
gas.
6. These particles can effects
photographic plate.
Properties of b Rays:
1. They
consist of negatively charged particles which have same e/m ratio as the cathode rays, so they are actually electrons.
2. Their
velocity is about ten times greater than that of a rays. Hence it is nearly equal to the
velocity of light.
3. These
rays are deflected by electric & magnetic fields in the direction opposite to the a -
particles.
4. Due
to lesser mass, the ionization that it produces is small as compared to a -
particles.
5. These
rays due to lesser mass and greater velocity have more penetration power than that of a particles.
6. The
e/m ratio of these rays is equal to
that of an electron.
7. These rays can effects photographic
plate.
Properties of g- Rays:
1. Gamma
rays do not consist of particles, actually they are electromagnetic radiations.
2. They
travel with very high speed which is equal to the velocity of light.
3. They
are not deflected either by electric or by magnetic field. Hence they carry no
charge.
4. They
are weak ionizers of gases, due to their non – material nature.
5. Due to high
velocity and non – material nature, they have greater power of penetration through solid.
6. Their power to effects a photographic
plate is small.
7. These
are more energetic rays than X – rays due to their shorter wave – length.
Plank’s Quantum Theory:
German physicist, Max Plank proposed this theory in 1900
to describe the origin of the radiations from heated bodies. This theory is
stated as:
“Radiant energy is
emitted or absorbed by a body in the form of small packets, called QUANTA, instead of being continuously.
This QUATA of energy are often called PHOTON.
Also the amount of
energy given off or absorbed is directly related to the frequency of the light
emitted”.
i.e.
E a u
=> E = hu
where E = energy of quanta
u =
frequency of radiation
h = Plank’s constant
=
6.625 x 10-34 J.sec
Spectra:
When
an element absorb sufficient amount of energy from a flame or its source its
emits radiant energy. When these radiation are passed through a prism, they get
separated into different components of colours.
e.g. When a ray of
light is passed through the prism, it is dispersed into seven colours. The
group of these colours of components is called “SPECTRUM”. Hence.
“A band of rays of
different wave – lengths obtained from the decomposition of radiation is called
SPECTRUM”.
In a spectrum each
colour or component has its own frequency and energy. The light of single wave
– length is called “Monochromatic”.
There are two types of
spectra.
ü
Continuous spectrum
ü
Discontinuous spectrum
Continuous Spectrum:
“It is
spectrum in which different colours are different into each other. There are no
dark spaces separating these colours”.
The visible portion of
spectrum obtained from ordinary sunlight and lights from incandescent sources
are the examples of continuous spectrum.
Line Spectrum:
“In
the line spectrum, the bands of colours are separated by dark spaces”.
This type of spectrum
may be obtained when light emitted from a gas source pass through a prism. A
common way of doing this is to pass an electric current through the gas at low
pressure (Crook’s tube). The neon
lights used in advertisement make use of this method for producing light and so
do Sodium Vapour street
lights. If the light from the discharge tube is allowed to pass through the
prism, some discrete sharp lines on an otherwise complete dark background are
obtained such spectrum is called “LINE
SPECTRUM”.
X – Rays:
Professor William Roentgen in 1895
discovered that when cathode rays (electrons) collide with a metal anode, a
very penetrating radiation is produced, which were named “X – Rays”. They were
also called Roentgen rays”. These
rays have the following properties.
Ø
They are invisible rays.
Ø
They can penetrate paper, rubber, glass, metal
and human flash.
Ø
They have very short wave length.
Ø
They have high energy and are electromagnetic in
nature.
Ø
Different metals produce X – Rays of different
wave lengths.
Discovery of Atomic Number:
In 1913, Henry Moseley bombarded different
metal anodes with cathode rays in order to det ermine
the magnitude of positive charge in atoms. From the study of different
wavelengths of the metals, the concluded that wavelengths of the X – rays
emitted decreased regularly with the increase in atomic mass. On Careful
examination he found that this is due to a fundamental factor in the nucleus,
i.e. the number of positive charges in the nucleus. He called the number of
positive charges contained in the nucleus as “ATOMIC NUMBER (Z)”
Rutherford’s Model of Atom:
After the discovery of electron,
proton & neutron, the attempts were made to see how these particles are
arranged in an atom. On the basis of following experiments, Rutherford
in 1911 not only discovered nucleus of the atom but also proposed a model of
atom.
Experiment:
(i.e. He2+
ions) was obtained from radioactive element polonium. He bombarded thin foil of
gold with a
particles and observed the effect on a fluorescent screen.
Observations:
In his experiment, Rutherford
observed that most of the a particles passed through the foil without any
deflection, but a few particles were deflected to very great extent. Also, he
was found that if 0.0004 cm thick gold foil is used then only one out of 8000
particles suffered a deflection through an angle greater than 90o.
Conclusions:
On the basis of these observations, Rutherford put forward the following conclusions.
1. Since majority of the a particles, pass straight
and undeflected through foil, it indicates that most of the volume occupied by
atom is extraordinary empty.
2. The
deflection of only few a particles which are positively charged indicates that centre of atom has positive charge. Rutherford
call the centre of atom NUCLEUS.
3. The total mass of the atom is concentrated in the nucleus,
whose dimension is negligible when compared with radius of atom.
4. The atom as a whole is neutral, therefore it was concluded
that the number of positively
charged particles (protons) within the nucleus
must be equal to negatively charged electrons.
5. The electrons are not stationary but revolving around the
nucleus with a very high speed. The centrifugal forces resulting from the motion of electrons balances
the electrostatic attraction of positive nucleus and keeps the electros away
from nucleus.
Weakness of Model:
The Rutherford ’s
atomic model has two serious defects:
1. According to
Maxwell’s theory a revolving electron will emit energy continuously. Hence due
to decrease in energy, the orbit of
revolving electron will become smaller and smaller until it would drop into the
nucleus. But it never happen so.
2. Since revolving
electron according to Rutherford emits energy
continuously, so it should give continuous spectrum but in actual practice a
line spectrum is obtained.
Bohr’s Atomic Theory:
The weakness in the Rutherford ’s model and the formation of line spectrum
were improved by Neil Bohr, who proposed a new theory to explain the electronic
structure of the atom in 1913. This theory is based upon the following
assumptions.
Assumptions:
1. Electrons in an
atom revolve around the nucleus in fixed circular which he called orbits or
energy levels.
2. As long as an
electron revolves in a particular energy level it does not emit or absorb
energy.
3. When an electron absorbs energy, it
moves to a higher energy level, further away from the nucleus. When it loss
energy, it returns to a lower energy
level, closer to the nucleus and the energy
is emitted as light.
4. The electron loses a definite quantity
of energy called “Quantum”, when it
jumps from an orbit of higher energy level to lower energy level.
5. The energy is emitted in the form of
radiations. The frequency of the energy emitted is directly proportional to the
difference in energy between two
levels.
i.e. E2 – E1 a u
E2 – E1 = hu
D E =
hu
where, E1 = Energy of electron in 1st
orbit.
E2
= Energy of electron in 2nd orbit.
DE =
Energy difference b/w two levels.
u = Frequency of the energy emitted.
h = Plank’s Constant
6. The angular momentum (mvr) of an electron in any orbit is integral multiple of 
.
.
where,
m = mass of electron
v =
velocity of electron
r =
Radius of the orbit
n =
Quantum number
=
1, 2, 3, ……………
Application of Bohr Theory:
§ In Hydrogen Atom:
By applying the concept of the
quantum of energy and the laws of classical mechanics Bohr worked out a
mathematical expression for the derivation of radius, energy, frequency &
wave number of an electron moving with a certain orbit in case of Hydrogen atom
Radius of an Orbit:
Suppose an atom
of Hydrogen with atomic number ‘Z’ and electro with mass ‘m’ , charge ‘e’ is
revolving around its nucleus at the distance ‘r’.
The electrostatic force
(centripetal force) between the nucleus and the electron would be
The centrifugal force which keeps
the electron away from the nucleus would be 
. Since the electron is in equilibrium.
. Since the electron is in equilibrium.
\ centrifugal force =
centripetal force
\ 
--------------
(i)
--------------
(i)
According to postulate of Bohr’s
Theory:
Þ 
--------------
(ii)
--------------
(ii)
Put this value of ‘v2’
in equation (i)
we have
Þ 
Þ 
Þ 
Þ 
Þ 
--------------
(iii)
--------------
(iii)
Energy of Electron:
An electron
possess Kinetic Energy due to its motion around the nucleus and the potential
energy due to the coulombic attraction force of the nucleus.
According to equation
Þ 
Þ 
Þ 
--------------
(iv)
--------------
(iv)
Now P.E. = Fcoul X Distance
Þ 
Þ 
The total energy of electro will
be the sum of its K.E. & P.E.
\ E = K.E + P.E
Þ 
Þ 
--------------
(v)
--------------
(v)
We know that
\ 
Þ 
Þ 
--------------
(vi)
--------------
(vi)
Frequency:
Suppose and
electron with energy E2, jumps from higher energy state n2,
to a lower energy state n1, with energy E1. Then the
energy released DE
by the electron will be
DE = E2 – E1
The energies E2 and E1
of the electron will be
and
\ 
Þ 2
2
2
Þ 
-------------- (vii)
-------------- (vii)
According to Bohr’s Postulate,
DE = hu
Þ 
Þ 
--------------
(viii)
--------------
(viii)
Wave Number:
The number of
waves per unit distance is called “Wave Number” , if “C” is the velocity of
light then the wave number is given by
Putting this value in eq. (viii)
we have
Þ 
Þ 
-------------- (ix)
-------------- (ix)
But
is a constant term which is called Rydberg constant denoted
by RH and has value 109678 cm-1.
Þ 
--------------
(x)
--------------
(x)
The expression is for wave
number.
SPECTRUM OF
HYDROGEN ATOM:
Although
Hydrogen atoms contains only one electron, its spectrum gives a large number of
series. Balmer in 1885, studied the spectrum of Hydrogen . he found that , when
energy is supplied to the sample of hydrogen gas, individual atoms absorb
different amount of energy. The electrons in higher energy levels are unstable
& drop back to the lower energy levels &during this process energy is
emitted in the from of line spectrum containing various lines of particular
frequency & wave length.
Balmer observed that a series
of lines appeared in the visible region
when the electrons drop from 3rd , 4th ,----nth energy
levels to the second orbit . These spectral line are known as “BALMER SERIES”. He proposed an
empirical formula to find wave no.(u) of each line.
where n2 = 3,4,5,…………
Lyman later on discovered another series in uv-region . wave no. of
each line was sound by similar formula.
where n2 = 2,3,4,…………
Paschen discovered another such series in infrared region. Wave no.
of each line was given by:-
Þ 
where n2 = 4,5,6,……………
Bracket found another series infar-infrared region . Wave no. of
such lines was given by.
Þ 
where n2 = 5,6,7,……………
Pfund also found another series far-infrared region. Wave no. of
each line was given by:-
Þ 
where n2 = 6,7,8,……………
Heisenberg’s Uncertainty Principle:
Statement:
According to
this principle,
“The position & the momentum
of an electron cannot be determined accurately simultaneously. One is measured
more accurately, the other becomes equally more uncertain at any given
instant.”
Mathematical Expression:
If DPx
is the uncertainty in the determination of the momentum of a particle & Dx was
the uncertainly in the simultaneous determination of it’s position , then the
product of these two uncertainties is given by:
DPx . Dx @ h
Thus if one of the two i.e. Px
or X was known exactly , then the uncertainty in its other would become
infinite . It means that the certainty of determination of one property
introduces uncertainty
For the determination of other.
Explanation:
The uncertainty
arise due to the fact that a light with shorter wave length than the located
its position. But the momentum of photon (particle of light) increases with decreasing
its wave-length.
An electron is so small a
particle, that it will be disturbed from its position, on colliding with a high
momentum photon particle used to locate it, & hence it is not possible to say as to what actual
position of electron is,
Orbit:
“The fixed
circular path on electron around the
nucleus called “Shell”
or “Energy Level” or “Orbit”.
These orbit are designated as K,
L, M, N, etc. The maximum number of electron in one orbit is “2n2“,
where “n” is the number of orbit. These orbit possess a definite amount of
energy increasing outwards from the nucleus.
Orbitals:
If a atomic
spectrum is observed, the spectral lines are found to be consist of two or more
fine lines closely packed together. Thus the electron In the same orbit may
differ in energy by small amount. Thus energy level are further divided into Sub – Energy Levels or “Orbitals”
& can be defined as
“Orbitals are
the regions around
the nucleus in which the probability of finding the electron is
maximum.”
The Orbitals have been named as s,
p, d, f, etc. The maximum no. of electron in s, p, d, f, are 2, 6, 10 & 14
respectively. Each orbit has “Orbital” equal
to its quantum number ‘n’. Thus the first orbit contain 1, the second orbit has
2 (i.e. s & p) , the third orbit has 3 (i.e. s, p & d) & fourth has
4 (i.e. s, p, d & f ) Orbitals.
SHAPES OF ORBITAL
S – orbital:
All’s
Orbitals are spherical in shape with the nucleus at the centre. Therefore in an
‘S’ orbital, the probability of finding
the electron is uniformly distributed
around the nucleus. It has only one possible orientation in space,
because it spread over all the three
axes uniformly . It has no nodal plane.
P – orbital:
The P - orbital are dumb-bell
shaped
and they are oriented in the
space
along the three mutually
perpendicular
axis (x,y,z) and are called Px,
Py and Pz
orbitals. All the three P-
orbitals are
perpendicular to each other.
These are degenerated orbitals,
that are equal energy. Each P – orbital has two lopes, one of which is labeled
(+) and the other is labeled (-).
Quantum NumberS:
In 1926 Schrödinger was a mathematician who
calculate the probability of location of the electrons in an orbital. Quantum
Numbers are the constant used in Schrödinger.
Wave equation to describe the energy of an electron, the shape of orbitals an
orientation in space around the nucleus of atom. The four quantum numbers are:
ü
Principle quantum number
ü
Azimuthal quantum number
ü
Magnetic quantum number
ü
Spin quantum number
Principle Quantum NumberS:
Principle
quantum number describes the size of an orbit and is represent by ‘n’. They
have any integral value i.e. n=1,2,3,………………… it never zero. The size and the
energy of orbital increase with increase the value of ‘n’.
Azimuthal Quantum NumberS:
Azimuthal quantum
number describes the shape of an orbital and its represented by ‘l’. It can have integral values ranging
from zero to n-1 i.e. l=0,1,2,…………… (n-1)
If l = o, the orbital
is called s orbital
If l = 1, the orbital
is called p orbital
If l = 2, the orbital
is called d orbital
Magnetic Quantum Numbers:
The magnetic
quantum number describes the different orientation of an orbital in the space
in applied magnetic field and it is represented by ‘m’. The value of m =- l to + l through zero. i.e. , m = -
l ….0….. + l.
If l = 2
Then m = -2 , -1, 0 , +1, +2
i.e. d-orbital has the five
orientations.
Spin Quantum Numbers:
It describes the
direction of spin of an electron around the nucleus of an atom and is
represented by ‘s’. Its values are = + ½ and – ½.
+ ½ spin for clockwise direction
- ½ spin for anti clockwise
direction
Principle of Electronic Configuration:
The
distribution of electron in the orbitals of the atom is called electronic
configuration. Following rules are used in the filing of the electrons in any
orbital.
(i)
Pauli exclusion principle
(ii)
Aufbau principle
(iii)
(n+ l) rule
(iv)
Hund’s rule
Pauli Exclusion Principle:
According to
this principle “No two electron of the same atom will have the same value of
all its four quantum numbers”.
Therefore in an atom two
electrons may have a maximum of three same quantum numbers, same value but the
fourth would be different. Thus in an orbital, when the value of n, l and m are same , the two electrons
can occupy the same orbital only there spin are opposite.
e.g. for 1st K shell
n=1, l=0,
m=0, s=+ ½
n=1, l=0,
m=0, s=- ½
Aufbau principle:
According to
this rule the electron are fed in various orbitals in order of increase orbital
energy starting with the one ‘s’ orbital”.
Hence electronic configuration of
an atom can be built up by placing the electrons to the lowest available
orbital until the total number of electrons added is equal to the atomic number
‘Z’. The sequence of increasing orbital energy is below
1s, 2s,2p,3s,3p,4s,3d,4p,……………
(n+ l)
Rule:
According to
this rule “In builting up the electronic configuration the orbitals 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”.
Here ‘n’ and ‘l’ stands for principal and azimuthal
quantum numbers respectively. Actually this rule is used to determine the
energy of any orbital and then we can apply the Aufbau Principle. e.g. 3d –
orbital is filled later than 4s – orbital
\ 3d – orbital has
(n+l ) = 3+2 = 5
4s
– orbital has (n+l ) = 4+0 = 4
Similarly 4p – orbitals fills
before 5s – orbital because 4p orbital has lesser value of ‘n’.
\ 4p– orbital has
(n+l ) = 4+1 = 5
5s
– orbital has (n+l ) = 5+0 = 5
Hund’s Rule:
According to
this rule “When the orbitals of same energy levels are available, then electrons
are distributed in orbitals in such a way as to give the maximum numbers of
unpaired electrons . Only when the orbital are separately occupied then the
pairing of electrons commences”.
This rule explain the filling of
electrons in degenerate (having same energy) orbitals like p,d,f. In simple, it
state that the electrons remain unpaired as far as possible. i.e. If there are
available orbitals of equal energy to the electrons, the electrons would lie in
separate orbitals and have same spin rather than to lie in the same orbital and
have paired spin.
 |
 |
||||
 |
|||||
e.g. (N(Z=7) = 1s2, 2s2,2Px
, 2Py , 2Pz is true |
and
not 1s2, 2s2,2Px,
2Py, 2Pz
Atomic Radius:
“It is half
distance between two nuclei of two identical atoms covalently bounded
together”.
Atomic Radius is measured in
Angstrom (Ao) or ‘nm’ units
1 Ao = 10-8cm
= 10-10m
1nm = 10-9m
Trend in the Periodic Table:
The atomic radii
of elements regularly decrease from left to right along a period. Because
increasing atomic number along a period increase the positive charge in the
nucleus, resulting in attraction of the outermost shell electrons more and more
in word and thus the shell are contracted.
The atomic radii of elements
regularly increase from top to bottom in a group. Because as we move from 1 element
to the other down a group, a new shell is added. This results in the expansion,
of the size of the atom on one hand and the reduction in the attraction of the
nucleus for the electron on the other. Therefore the atom radii of the elements
increase down the group.
Ionic Radius:
Ions are formed
due to loss or gain of electron from the outermost shell. Positive ion is
called cation and negative ion is called anion. Ionic Radius is defined as “The
radius of spherical ion”.
The ionic radii of cations are
smaller from which they are formed because when an atom looses electron to form
a positive ion, the outermost shell vanishes away and the next shell becomes
the outermost shell, while the quantity of the charge in the nucleus remains
the same. As the result of it the attraction of the nucleus for the outermost
shell electrons increases causing decrease in the ionic radius.
On the other hand anions are
larger than the atoms from which they are formed because in the case of
negative ion the electrons are added in the same outer most shell. Therefore
the ions tend to expand with the increasing negative charge on them, resulting
in the increase in their ionic radii.
Ionization Potential:
“The minimum
amount of energy require to remove completely the outer most electrons from an
atom in gaseous state resulting in the formation of positive ions is called
Ionization Potential” (IP). Simply ionization energy.
When the first electron is
removed it is called first ionization energy when the second, third and fourth
electron are removed the energies used are called second, third, fourth
ionization energy value respectively and son on. The value of I.E. values of
element increase from first I.E. to second I.E. to third I.E. and so on.
Al 1st I.E. Al+ + e-, DHo = +518 Kj/moleAl+ 2nd t I.E. Al2+
+ e-, DHo
=+1817Kj/moleAl2+ 3rd
I.E.
Al3+ + e-,
DHo =+2745Kj/mole
Trends in the Periodic Table:
I.E. decrease
down a group:
As we go down a
group the addition of new shell increase the atomic radii as well as shielding
effect. This result in the decrease in I.E. of the elements down a group in the
periodic table.
I.E. increases along a period:
As we go from
left to right in the periodic table the atomic number (+charge in the nucleus) of
elements increase regularly. This increases the attraction force between the
outermost shell electron and the nuclei resulting in the decrease in the atomic
radii. Consequently the outermost electrons are held more strongly in the atom
and thus more energy is required from this removal. Therefore the ionization of
element increase along a period from left to right.
Electron Affinity:
“It is the
amount of energy released, when an electron is added to the outermost shell of
gaseous atom to form a negative ion”.
Electron Affinity (E.A.) is
effected by the factors such as atomic size, nuclear charge and shielding
effect of the electrons.
The E.A. increases from left to
right in a period due to decrease in the atomic size and increase in the
nuclear charge. But it decrease down a group from top to bottom because the
atomic size increase.
Electronegativity:
“Electronegativity
is the force of attraction of an atom to attract a shared pair of electron
towards itself” All elements have different values of electronegativities.
Fluorine is the most electronegative element and it used as standard for the
comparison of electronegativities of other element. Its value is fixed as 4.
Tend in the Periodic Table:
The
electronegativity values of the element increase from left to right in a period
because of increase in nuclear charge and the decrease in the size of atom.
On other hand electronegativity
values of the elements decrease from top to bottom in the group due to increase
in atomic size.
Nature of Chemical Bond:
The nature of a chemical bond can
be decided on the basis of electronegativity difference between the two bonded
atom i.e.
i. “When E.N. difference is less than 0.5
it will be a non polar covalent bond”.
ii. “When the E.N. difference is between
0.6 to 1.7, it will be Polar
– Covalent bond.
iii. “When E.N. difference is greater than
1.7 then it will be “ionic bond”.
e.g.
(a) C – H , is a non Polar Covalent bond
\ E.N.
difference = 2.2 – 2.1 = 0.4 < 0.7
(b) H – Cl, is a Polar Covalent bond
\ E.N.
difference = 3.5 – 2.1 = 1.4 > 0.7 but < 1.7
(c) Na – Cl, is a ionic bond
\ E.N.
difference = 3.5 – 0.9 = 2.4 > 1.7
Table 5.1 Highlights in the Development of Contemporary
Atomic Theory
Year
|
Name
of Scientists
|
Work
|
1803
|
John Dalton (English)
|
Proposed an atomic theory to explain the laws of definite
composition and multiple proportions.
|
1879
|
William Crookers (English)
|
Repeated the observations of earlier workers and
concluded that rays produced by an electric current in an evacuated tube
consist of a stream of charged particles produced at the cathode.
|
1885
|
Johann J. Balmer (Swiss)
|
Developed a simple empirical equation that could be used
to calculate the wavelength of lines in the spectrum of hydrogen atoms.
|
1886
|
Eugen Goldstein (Germen)
|
Characterized “canal rays” (positive ions) produced by an
electric current in an evacuated cube.
|
1895
|
Wilhelm Roentgen (Germen)
|
Discovered X rays produced by an electric current in an
evacuated cube.
|
1896
|
Antoine Becquerel (French)
|
Discovered natural radioactivity.
|
1890
|
Johannes Rydberg (German)
|
Modified Balmer’s equation to describe the frequencies of
the lines in the hydrogen spectrum.
|
1897
|
J. J. Thomson (British)
|
Showed that Crookes’s cathode rays are beams of
negativity charged particles (now known to be electrons) and calculated their
charge to-mass ratio.
|
1898
|
Marie Curie (Polish) and her Husband, Pierre (French)
|
Isolated polonium and radium from pitchblende (a lead
ore) by chemical processes.
|
1900
|
Max Planck (German)
|
Proposed a quantum theory that described the light
emitted from a hot object as composed of discrete units called quanta or photons.
|
1903
|
Hantaro Nagoka (Japanese)
|
Postulated a saturn-like atom, a positively charged
sphere surrounded by a halo of electrons.
|
1904
|
J. J. Thomson
|
Proposed a model of the atom with electrons embedded in a
sea of positive charges, the “plum-pudding model.”
|
1905
|
Albert Einstein (German)
|
Described the relationship of mass and energy.
|
1909
|
Robert Milliken (American)
|
Determined the charge on the electron in his oil drop
experiment.
|
1911
|
Ernest Rutherford (New Zealander)
|
Attributed the scattering of a particles by a thin gold
foil to an atomic structure in which an atom’s mass is concentrated in a
tiny, positively charged nucleus.
|
1912
|
J. J. Thomson
|
Detected the isotopes neon-20 and neon-22.
|
1913
|
Niels Bohr (Danish)
|
Showed that the hydrogen spectrum could be explained by a
previously unrecognized property of matter-the energy of electrons in atoms
is limited to certain values (quantized).
|
1913
|
Henry Moseley (English)
|
Observed that an element emits characteristics X rays
that depend on its nuclear charge (atomic number).
|
1920
|
Orme Masson (Australian), William Harkins (American),
|
Independently postulated the existence of an uncharged
particle with the mass of proton (the neutron).
|
1924
|
Louis de Broglie (French)
|
Combined equations from Einstein and Planck to suggest
that electrons have wave-like properties.
|
1925
|
Wolfgang Pauli (German)
|
Stated that no two electrons in the same atom can have
the same set of quantum numbers (Pauli’s exclusion principle).
|
1926
|
Erwin Schrödinger (Austrian)
|
Combined the particle nature of an electron, its wave
properties, and quantum restrictions and developed an equation that described
the energy and likely location of electron in a probability relationship.
|
1927
|
Werner Heisenberg (German)
|
Showed that we cannot determine simultaneously both the
exact position and the momentum of an election (Heisenberg’s uncertainty
principle).
|
1927
|
Frederick Hund (German)
|
Determined that electrons in subshells have maximum
unpairing and that unpaired electrons have the same spin (Hund’s rule).
|
1932
|
James Chadwick (British)
|
Characterized neutroms.
|
Ref#
Page no. 144 and 145 General Chemistry 10th edition by Robinson,
Odom, Holtzclaw
City Of Knowledge
(Science Campus)
Add: 24/2 Sheet # 10 Model Colony Near Model Colony Police Station Karachi ,Pakistan
Ph # +9221-35449369, +92312-5449369
No comments:
Post a Comment