Entry Test Preparation 2015
Chemistry
BOOK 1
Chapter # 4
ATOMIC STRUCTURE
Sub-atomic Particles of Atom:
Electronics, protons and neutrons, are called fundamental particles of
atom.
Electron:
Electron was discovered by Thomson during cathode ray experiments.
Electrons have and electric charge of -1.6022
x 10-19 coulomb, a mass of 9.11 x 10-31 kg based on
charge/mass measurements and a relativistic rest mass of about 0.511 MeV/c2.
The mass of the electron is
approximately 1/1836 of the mass of the proton. The common electron symbol is e-.
Cathode Rays:
·
Cathode rays are beams of
electrons.
·
e/m (electron to mass)
value of cathode rays is found to be a universal constant independent of the
nature of the gas. It is numerically equal to 1.7589 x 1011 CKg-1, C is the unit of charge
called coulomb.
Proton:
The proton is a subatomic
particle with and electric charge of one positive fundamental unit (1.602 x 10-19
coulomb, a diameter of about 1.5 x 10-15
m, and a mass of 938.3 MeV/c2 (1.6726 x 10-27 kg), 1.007
276 466 88(13) amu or about 1836 times the mass of an electron.
Neutron:
The neutron is a subatomic
particle with no net electric charge and a mass of 939.573 Mev/c2 (1.6749 x 10-27
kg, slightly more than a proton). Its spin is ½. Its antiparticle is called the
antineutron. The neutron, along with the proton, is a nucleon.
Discovery of Electrons
At ordinary pressure the air
inside the Discharge Tube do not conduct electricity even when the electrodes
were connected to source of very high potential of about 5000 volts.
If the pressure inside the tube
is reduced, and a high voltage of 5000-10000 volts is applied, an electric
discharge is produce through the gas producing a uniform glow inside the tube.
By reducing
the pressure in the tube still further to about 0.01 torr the original glow
disappears. Rays are produced which creates fluorescence on the glass wall
opposite to the cathode. These rays are called Cathode Rays.
Cathode rays are actually
negatively charged particles.
Properties of Cathode Rays and Discovery of Electrons:
Cathode rays are negatively
charged.
They produce a greenish
fluorescence on striking the walls of the glass tube.
Cathode rays cast shadow when a
opaque object is placed in their path. This proves that they travel in a
straight line perpendicular to the surface of cathode
These rays can drive a small
paddle wheel placed in their path. This shows that these rays possess momentum.
Cathode rays can produce X-rays
when they strike an anode particularly with large atomic mass.
Cathode rays can produce heat
when they fall on matter e.g. when cathode rays from a concave cathode are
focused on a platinum foil, it begins to glows.
Cathode rays can ionize gases.
These rays can pass through a
thin metal foil like aluminum or gold foil.
They can cause a chemical change
because they have reducing effect.
Discovery of Protons (Positive Rays):
The positive rays had first been
observed by Goldstein in 1886. He had cut small holes in the cathode, and seen
pencil-like rays streaming through the holes, which he called canal radiation.
The positive charge of the rays was not identified until sixteen years later.
They were much harder to deflect than cathode rays and were deflected in the
opposite direction, suggesting they could be massive, positively charged
particles.
Characteristic Properties of Positive Rays:
These rays travel in a straight
line and are deflected by electric and magnetic fields in a way that reveals
their positive charge.
The ratio of charge to mass,
(e/m) for these positively charged rays in considerably smaller than for
electrons.
The e/m ratio of the positive
rays depends upon the nature of the gas in the tube. The highest e/m ratio is
obtained if hydrogen gas is present.
Measurement of e/m value of electron:
In 1897, J.J. Thomson devised and
instrument to measure the e/m value of electrons. It came out to be 1.7588 x 1011
coulombs kg -1. This means that 1 kg of electrons have 1.7588 x 1011
coulombs of charge.
Measurement of Charge on Electron
Millikan’s Oil Drop Method:
In 1909, Millikan determined the
charge on electron. It was found to be 1.6022 x 10-19 coulombs. The
charge present on an electron is the smallest charge of electricity that has
been measured so far.
Discovery of Neutron
21684Po → 21282Pb + 42He
42He + 94Be → 126C + 1on
·
The positive charge in
spread over a sphere of radius 10-15m, called nucleus, and the outer-most
electron clouds are about 10-10m from the centre of the nucleus.
·
The positive charge of a
nucleus is due to protons but mass of the nucleus is due to both protons and
neutrons.
·
The total number of protons
and neutrons in the nucleus determines the nuclear mass.
·
Electrons revolve around
the nucleus in circular orbits in which the required centripetal force is
provided by the electrostatic force of attraction between nucleus (+Ze) and
electron, that is
Discovery of Nucleus:
Ernest Rutherford was observing
the effects of shooting a narrow beam of small alpha particles at a thin gold
foil (0.00004cm). Rutherford noticed when the
alpha particles struck the thin metal, some of them scattered instead of
continuing straight through. The number of alpha particles scattered by an
angle greater than 90 degrees was much more than predicted. Rutherford
explained that this was possible because any particle that ran into the nucleus
would rebound and get deflected, which would not happen if a particle has
simply traveled through an atom.
Planck’s Quantum theory
According to his revolutionary
theory, energy travels in a discontinuous manner and it is composed of large
number of tiny discrete units called quanta. The main points of his theory are:
Energy is not emitted or absorbed
continuously. Rather it is emitted or absorbed in a discontinuous manner and in
the form of wave packets. Each was packet or quantum is associated with a
definite amount of energy. In case of light, the quantum of energy is often
called photon.
The amount of energy associated
with a quantum of radiation is proportional to the frequency (v) of the
radiation.
A body can emit or absorb energy
only in terms integral multiples of a quantum.
E = n h r Where in = 1, 2, 3,……………
E = h c / 
………………..
Bohr’s Model of Atom:
·
Electrons revolve around
the nucleus in circular orbits and an electron does not radiate energy if it is
in its orbits.
·
Electron can revolve only
in those orbits whose angular momentum (mvr) is an integral multiple of a
factor 
,
,
mvr = n
·
An electron continues to
move in a particular orbit called stationary orbit or energy level, without
losing energy. This state is called the ground state.
·
If energy is supplied to an
electron, it may jump from a lower energy level to a higher energy level by absorbing
one or more quanta or photons of energy. This new state is called the excited
state. The energy absorbed in a jump is equal to the difference in energies of
the two energy levels. Same amount of energy is emitted during de-excitation.
Thus, energy absorbed or released in an electron jump (∆E) is given by
Where ∆E is the change between
energy, E1 of the electron in the lower energy orbit and energy E2
of the electron in the higher orbit, h is the plank’s constant v is the
frequency of the emitted radiation, c is the velocity of light, and
∆E = E2-E1
= hv (energy of photon) = 
is the wavelength of
the emitted radiation.
Radius of nth Bohr’s orbit:
Where angstrom (A) is a unit of
length.
A = 1.0 x 10-10 meters = 0.1 nm = 100 pm
Where ro is:
ro =εoh2 /πme2
Where
h = Planks constant = 6.6 x 10-34 Js
m = mass of electron
e = charge on electron
r2-r1<
r3-r2< r4-r3< ------------------
raduis of
second Bohr orbit is four times that of the first orbit.
r2
= 0.529 x (2)2 = 0.529 4 A
Energy of electron in nth Bohr orbit:
En = -me4/8 εo2h2
[1/n2] j/atom
En = -2.18 x 10-18 [1/n2]
En = -2.18 x 10-18 x 6.02 x 1023 / 1000
x n2
= -1313.315 / n2 Kj /mol
E2-E1> E3-E2> E4-E3 >----------------
En=
Where -13.6 eV is the energy of
electron in first Bohr orbit, that is, the ionization energy of hydrogen atom.
Speed of Electron in an Orbit:
Limitation of Bohr’s Atomic Model:
Bohr’s Model is true only for one
electron system, for example. H, He+, Li2+, Be3+
and so on. It fails for atoms having more than one electron.
A visual display or dispersion of
the components of white light, when it is passed through a prism is called a spectrum. Spectrum is of two types.
·
Continuous spectrum
·
Atomic or line spectrum
Continuous Spectrum:
In this type of spectrum, the
boundary line between the colours cannot be marked. The colours diffuse into
each other. The best example of continuous spectrum is rainbow. It is obtained
from the light emitted by the sun or incandescent (electric light) solids.
Atomic or line Spectrum:
The spectrum which is not
continuous and has breaks is called line spectrum. Atomic spectrum is usually
line spectrum. When an element or its compound is volatilized on a flame and
the light emitted is seen through a spectrometer, we see distinct lines
separated by dark spaces. This type of spectrum is called line spectrum or
atomic spectrum. There are two ways in which an atomic spectrum can be viewed.
·
Atomic emission spectrum
·
Atomic absorption spectrum
Atomic Emission Spectrum:
When solids are volatilized or
elements in their gaseous states are heated to high temperature or subjected to
an electrical discharge, radiations or certain wavelengths are emitted. The
spectrum of this radiation contained bright lines against dark background. This
is called atomic emission spectrum.
Atomic absorption spectrum:
When a beam of white light is
passed through a gaseous sample of an element, the element absorbs certain
wavelengths while the rest of the wavelengths pass through it. The spectrum of
such radiation is called an atomic absorption spectrum. While the rest of
wavelengths pass through it. The spectrum of these radiations is called an atomic
absorption spectrum. The wavelengths of the radiation that have been absorbed
by the element appear as dark lines and the background is bright.
Hydrogen Spectrum:
Hydrogen spectrum is an important
example of atomic spectrum. Hydrogen is filled in a discharge tube at a very
low pressure; on applying a very high voltage the electrodes, a bluish light is
emitted from the discharge tube. The light when viewed through a spectrometer
shows several isolate sharp lines. These are called spectral lines. The wavelengths
of these lines lie in the visible, ultra violet and infrared regions. These
spectral lines can be classified into five groups called spectral series. These
series are named after their discoverers as below:
(i) Lyman series (U.V region)
(ii) Balmer series (visible region)
(iii) Paschen series (I.R. region)
(iv) Brackett series (I.R. region)
(v) Pfund series (I.R. region)
X-rays
X-rays are produced when
electrons having high kinetic energy, collide with heavy metal (anode) in the
discharge tube. When the electrons are stopped on collision. Energy is released
in the form of electromagnetic waves. In the discharge tube, the electrons
produced by a heated tungsten filament are accelerated by high voltage. It
gives them sufficient kinetic energy to bring about the emission of X-rays on
striking the target metal. X-rays are emitted from the target in all
directions, by only a small portion of them is used for useful purposes through
windows.
The wavelength
of X-rays produced depends upon the nature of the target metal.
Every metal
has its own characteristic X-rays
Mosley’s Law:
This law states that the
frequency of a spectral line in X-rays spectrum varies as the square of atomic
number of an element emitting it.
This law leads
to the conclusion that it is the atomic number and not the atomic mass of the
element which determines its characteristic properties, both physical and
chemical.
……………………………
Wave Particle Nature of Matter:
Planck’s quantum theory of
radiation tells us that light shows a dual character. It behaves both as a
material particle and a wave.
According to de Broglie all
matter particles in motion have a dual character. It means that electrons,
protons, neutrons, atom and molecules possess the characteristics of both the
material particle and a wave.
This is called wave particle
duality in matter. De Broglie derived a mathematical equation which relates the
wavelength () of the electron to the momentum of electron.
) of the electron to the momentum of electron.
h = Planck’s
constant = 6.625 
10-34 Js
10-34 Js= de Broglie’s
wavelength
m = mass of the particle
v = velocity of electron
Schrödinger Theory of Wave Nature of Particle:
A German physicist Erwin
Schrodinger proposed that electrons in any atoms may be considered as standing
waves. The electron wave can be associated with a circlar standing wave about
the nucleus. The condition on the allowed wavelengths for the circular standing
wave is
This relation is called
uncertainty principle.
This equation shows that if ∆x is
small then ∆p will be large and vice versa. So, if one quantity is measured
accurately then the other becomes less accurate.
The uncertainty principle is
applicable only for small particles like electrons neutrons etc, and has no
significance for large particles, i.e., macroscopic particle.
Atomic Orbital:
The volume of space in which
there is 95% chance of finding an electron is called atomic orbital.
Quantization
To say that something is quantized
means that it comes in particular, well-defined packet of quantities that do
not change linearly form one value to another, rather the change Is in
multiples.
The quantum mechanical model of atom:
The important consequences of the
quantum mechanical description of atoms include the following:
·
The energies of electrons
in atoms are quantized.
·
The number of possible
energy levels for electrons of different elements is a direct consequence of
the wave like properties of electrons.
·
The position and momentum
of the electron cannot be determined simultaneously.
·
The region in space around
the nucleus in which the electron is most probably located is what can be
predicted in an atom.
·
Electrons of different
energies are likely to be found in different regions. The region in space round
the nucleus in which an electron with a specific energy Is most probably
located is called an atomic orbital.
Quantum Numbers:
·
The motion of an electron
in an atom is determined by four quantum numbers. The set of four quantum numbers
for a particular electron is unique in the atom. No two electrons can have same
set of quantum numbers.
Principal quantum number (n):
This is used to specify the
energy and position of an electron in an orbit or shell. This is denoted by n.
n = 1, 2, 3, ………..
(Only positive
integers)
(Only positive
integers)
·
Position of electron from
the nucleus in the nth orbit is given by 
·
Corresponding energy 
n specifies the orbit number of
the electron.
Azimuthal Quantum Number (l):
This is used to describe a
subshell. This is also called subsidiary quantum number.
I = 0, 1, 2, 3, ……….(n-1)
|
I
|
0
|
1
|
2
|
3
|
|
Subshell
|
s
|
p
|
d
|
F
|
|
s
= sharp spectral lines.
p
= Principle spectral lines.
d
= degenerate spectral lines.
f = fundamental spectral lines.
|
Magnetic Quantum Number (m):
This quantum number determines
the preferred orientations of orbitals in space.
m = -1 ………. 0………..+1
·
Total values of m = (2/ +
1) = number of orbitals in a subshell (l).
·
Thus s-subshell (m = 0)
have one orbital.
·
p-subshell (m=-1, 0, +1)
has three orbitals (Px, Py, Pz) These are degenerate orbitals.
·
d-sub-shell (m= -2, -1, 0,
+1, +2) has five orbitals (dxy, dyx, dzx, dx2 – y2,dz2).
These are also
degenerate orbitals.
·
Splitting of spectral lines
occurs when placed in a magnetic field (Zeemen’s effect) or in an electric
filed (Stark effect).
·
Total lines from a single
line in the normal spectrum = (2/+1)
·
Total number of orbitals in
nth shell = n2
Spin Quantum Number (s):
Electron spins on its axis, like
a top, in a clockwise and anticlockwise direction. Spin quantum number
describes the orientation of the electron in an orbital.
S = +1/2 or 
(spin up for
clockwise).
(spin up for
clockwise).
S = -1/2 or 
(spin down for anticlockwise).
(spin down for anticlockwise).
In an orbital, there are maximum
of two electrons such that there are +
and -
and -
Sun of spins
of the two electrons of an orbital is zero 
Shapes of Orbitals:
Depending upon the values of
azimuthal quantum number, there are four types of orbitals, these orbitals are
s, p, d and f having azimuthal quantum number values as l = 0, 1, 2, 3,
respectively.
Shape of s-orbitals:
There are three values of
magnetic quantum number for p-sub shell. So p sub-shell has three orientations
in space i.e., along x, y, and z-axes. All the three p-orbitals namely Px,
Py and Pz have dumb bell shapes.
Shape of d-orbitals:
For d sub-shell three are five
values of magnetic quantum number. So there are five space orientations
along x, y and z-axes. Thery are
designed as dxy, dyz, dxz, dx2, dx2.
Stationary state:
An electron in an atom whose
energy does not change with time is in a stationary state. Ground state:
This is the state of an atom or molecule
in which all the electrons have their lowest energies.
Excited state:
If electrons are given extra
energy and move from a lower to a higher state, then the atom or molecule is
said to be in an excited state.
Electronic distribution;
The following rules have been
adopted to distribute the electrons in sub-shells or orbitals.
Aufbau rule:
·
Electrons in various
sub-shells of a shell are filled in the increasing order of their energies.
·
Smaller the value of (n +
l), smaller the energy.
·
If two or more orbits have
same value of (n + l), subshell with lower value of n has lower energy.
·
Energy order of various
subshells is given below.
|
1s<2s<2p<3s<4s<3d<4s<4p<……
|
Hund’s Rule of Maximum
Multiplicity:
Pairing of electrons in various
orbitals (like Px, Py, Pz) of a sub-orbit take
place only after each orbital is half filled (singly occupied. For example,
electronic configuration of P(Z=15) is given below
15 P : 1s2
2s2 2p6 3s2 
Pauli’s Exclusion Principle:
It states that, no two electrons
in an atom can have the same set of quantum numbers. For example, two electrons
in 1s2 (helium) have same values of n, l, m but they differ in spin.
|
Quantum Number
|
n
|
l
|
m
|
s
|
|
1st
electron
|
1
|
0
|
0
|
+1/2
|
|
2nd
electron
|
1
|
0
|
0
|
-1/2
|
Electro negativity:
Its is the power of an atom to
attract the shared pair of electrons towards itself in a molecule.
Ionization potential:
The amount of energy required to
remove the most loosely bound electron from an atom in the gaseous state is
called Ionization Energy.
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