The
term atom was introduced by Dalton.
·
In 1807, John Dalton proposed his famous atomic theory
Dalton’s
Atomic Theory :
i.
Matter is composed up extremely small particles called atoms.
ii.
Atoms are indivisible. They can neither be created nor be destroyed.
iii.
Atoms of the same element are alike in properties.
iv.
Atoms combine in small whole numbers to form compound atoms.
Discovery
of Subatomic Particles :
·
J.J. Thomas (1897) concluded that cathode rays consist of a stream of
fast moving negatively charged particles called electrons. He also
determined the velocity of the electron and their charge-mass (e/m)
ratio using different gases, and found the value 1.75875 ×
1011CKg-1.
·
In 1909 Millikan determined the charge and mass of the electron by
using his famous oil drop technique and found the values 1.60206 ×
10-19C and 9.1091 × 10-11Kg respectively.
·
Goldstein (1886) used perforated cathode in the discharge tube and
observed the emission of anode rays, which consist of positively
charged particles known as protons. The charge and mass of the proton
are found the values 1.6 × 10-19C and 1.672 × 10-27Kg
respectively.
·
In 1932, Chadwick discovered neutron, a neutral particle by the
bombardment of α-particles on beryllium or boron.
J.J.
Thomson’s Model :
According
to this model, atom is compared with watermelon. Seeds are analogous
to electrons and the edible part is the positive electrically.
Drawbacks
of Watermelon model :
i.
It cannot explain the stability of atom.
ii.
Hydrogen spectra could not be explained.
Rutherford’s
Experiment (Model) :
In
1911, Rutherford concluded that when α-particles struck on thin (4 ×
10-5cm thick) sheet of gold :
-Most
of the α-particle continued their straight path => Most part of
atom is empty.
-
Some α-particle deviated => The positive charge body of the atom
is concentrated only at the centre of the atom called neucleus.
-
Some α-particle (very few) bounced back => The positively charged
mass is occupying a very small space.
The
net outcome of Rutherford’s model gave the idea about nucleus
(diameter 10-15m). He is the discoverer of nucleus. Hence
the model is also known as nuclear model.
-
We know that, nucleus has a diameter of the order of 10-15m
while the atom has a diameter of the order of 10-10m.
Drawbacks
of Rutherford’s model :
-Stability
of atom couldn’t be explained. According to classical
electrodynamics, accelerating particle emits radiation continuously.
Therefore the revolving electron must loose energy continuously and
the electron will steadily drift towards the nucleus and ultimately
will fall into the nucleus collapsing the atom.
-Hydrogen
Spectra couldn’t be explained.
=The
atomic number is a fundamental property of the element and equal to
number of protons in the nucleus i.e. equal to unit positive on the
nucleus.
=The
sum of the number of proton (Z) and neutrons (N) in a called the mass
number (A) of the atom or A = N + Z. Hence proton, neutron and
electron are the fundamental particles of an atom.
Bohr’s
Model of Atom :
This
theory, proposed by Neil Bohr, is based on both classical and quantum
theory of planck. The important postulates of this model are:
i.
The first postulate give the idea about circular orbits. An electron
can revolve only in those orbits whose angular momentum (mvr) is an
integral multiple of h/2π. i.e., mvr = nh/2π.
Where,
m = mass of electron, v = velocity of electron, r = radius of orbit,
h = Planck’s constant and n = number of orbit in which electron is
present.
ii.
So long as the electrons revolve in particular orbit, it neither
gains nor looses energy. So the circular orbit is also known as
stationary orbit or energy level.
iii.
When electron jumps from one orbit to another, the difference in
energy is emitted as radiation given by ΔE = E2 – E1=
hv.
When
electron jumps from lower energy level, it gains energy but loses
when vice-versa.
Advantages
of Bohr’s model :
Bohr’s
theory satisfactorily explains the spectra of species having one
electron, viz hydrogen atom, He+,Le2+ etc.
Further his postulates can also be used for calculating :
i.
Radii of various orbits of hydrogen atom like species, r =
n2h2
4π2me2z
=
0.529 × n2 Å
z
where,
0.529 is the radius of the first orbit of hydrogen atom.
ii.
The velocity of electron in an orbit , v =
2πe2
nh
The
velocity of electron in the first orbit (i.e. when n = 1), also known
as Bohr’s velocity is,
V1
= 2.19 × 108 cms-1
Which
is 1/138 of velocity of light.
i.e.
Vn = c/138n
iii.
Energy of electron in different orbits En =
-2π2mz2e4
= -13.6 ev
n2h2
n2
-
Energy of an electron is directly proportional to the square of n.
i.e. En ∝ n2.
-
Although the energy of an electron increases with increase in the
value of n (orbit), yet the difference of energy between successive
orbits decreases.
Thus,
E2 – E1 > E3 – E2
> E4 – E3 etc.
iv.
It could also explain the hydrogen spectra with following spectral
lines :
a.
Lyman series :
Lies
in U-V region, Line appears in the atomic spectrum due to drops of
electrons from higher energy levels to the lowest energy level (n1
= 1)
b.
Balmer series :
Lies
in visible region, Lines appear due to drop of electrons from energy
levels 3, 4, 5, 6, .................. etc. to the energy level 2 (n1
= 2).
c.
Similarly, the Paschen, Brakett and Pfund series :
Correspond
to drop of electrons from higher energy levels 3, 4 and 5
respectively. These series lie in infrared region.
Limitations
of Bohr’s model :
i.
It does not explain the spectra of atoms having more than one
electron.
ii.
Bohr’s theory could not explain this multiple or fine structure of
spectral lines.
iii.
It does not explain the splitting of spectral lines into a group of
finer lines under the influence of magnetic field (Zeeman’s effect)
and electric field (Stark’s effect).
iv.
This model could not obey de-Broglies hypothesis and Heisenberg’s
uncertainty principle.
Bohr
Sommerfield’s model :
In
order to explain the fine spectrum sommerfield modified Bohr’s
model and gave few postulates :
i.
The path of electron is elliptical, circular path is the special case
of elliptical path.
ii.
Orbit is composed of sub – orbits ,
Number
of sub orbits = orbit number.
De-Broglie’s
Equation :
“the
momentum of a particle in motion is inversly proportional to the
wavelength of the waves associated with it”.
i.e.
λ ∝ (1/mv) ∴ λ = (h/mv)
thus
the de broglie’s equation points out that every thing in nature
posses the properties of particles as well as waves.
Heisenberg
Uncertainity Principle :
“It
is impossible to determine simultaneously the position and momentum
of microscopic particles with absolute certainity.” Mathematically,
Δx.Δp
≥ h/2π
Where,
Δx = uncertainity in position, Δp = uncertainity in momentum.
Quantum
Numbers :
The
“state” of an electron in an atom is completely defined by a set
of four numbers.
Here
state refers to position, energy and orientation of electrons.
Four
Quantum numbers are :
1.
Principal Quantum numbers (n) :
It
gives the idea about principal energy level to which electron belongs
and the average distance between the electrons and nucleus higher the
principal quantum number, greater is its size and also higher is its
energy.
Though
theoretically its value ranges from 1 to ∞ , only 1 to 7 are known,
beyond 7, attraction between electron and nucleus is very small that
atoms get ionized. They are distignated either as 1, 2, 3, 4, 5, 6
and 7 or K, L, M, N, O, P and Q respectively.
The
maximum number of electrons in n principal quantum number is given by
2n2.
2.
Azimuthal Quantum numbers (l) :
This
number denotes the sub-shell to which the electron belongs and
determines the shapes of the orbital and the energy associated with
the angular momentum of the electron. For a given value of principal
quantum number n, the azimuthal quantum number (l) may have all
integral values from 0 to (n-1) each representing different
sub-shell, denoted by s (sharp), p (principal), d (diffused) and f
(fundamental).
|
Symbol
of sub-shell :
|
s
|
p
|
d
|
F
|
|
Value
of l
|
0
|
1
|
2
|
3
|
3.
Magnetic Quanum numbers (m) :
4.
Spin Quantum numbers (s) :
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