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ATOMIC STRUCTURE
The word atom is derived from the Greek word atom which means indivisible. The Greeks concluded
that matter could be broken down into particles to small to be seen. These particles were called
atoms. Dalton believed atom was a solid, indivisible, with different atoms for different elements. But
1
later many research works proved all atoms (except normal hydrogen 1H )are composed of three
fundamental particles, namely electrons, protons and neutrons.
Fundamental particles
(a) Electron: The nature and existence of electron was established by experiments on
conduction of electricity through gases, i.e., discovery of cathode rays. J.J. Thomson (1897)
8
determined specific charge (e/m) or the charge to mass ratio of the electron (-1.7588 x 10
coulomb/g) and proved that whatever gas be taken in the discharge tube (Discharge tube
consists of a glass tube with metal electrodes fused in the walls), the value of e/m is always
the same. From that he concluded that all atoms contained electrons. The name electron was
given by Irish physicist, Stoney
(b) Proton: A proton is a fundamental particle with a charge equal in magnitude but opposite in
sign to the charge on the electron, and having a much larger mass. The nature and existence
of proton was established by the discovery of Positive rays (Goldstein in 1886).e/m value of
the anode rays is dependent on the nature of the gas taken in the discharge tube, i.e., positive
particles are different in different gases. The specific charge (e/m) of the anode rays is found
to be maximum when gas present in the discharge tube is Hydrogen. The name proton was
given by Ruther ford.
Neutron: It is sub-atomic particle, which carries no charge. Although the existence of neutrons
was predicted in 1920, the prediction was verified by Chadwick in 1932.
The characteristics of fundamental particles are given below.
Particle
Symbol
Mass in
amu or u
Mass in kg.
Charge
esu
Electron
-1e0
0.000548
9.1091x10 –31
-4.803 x 10-10
1
in
-1.602 x 10-19
Relative
Charge
-1
-27
+4.803x 10
+1.602 x 10
+1
–27
0
0
0
Proton
+1P
1.00757
1.6725x10
Neutron
1
0n
1.00893
1.6748x10
-10
Charge
in
coulomb(C)
-19
Note: The esu, an abbreviation for electrostatic unit, is the cgs unit of charge. coulomb is the
MKS unit of charge.
THOMSON’S ATOMIC MODEL: J.J Thomson concluded that all atoms contain electrons from his
experiments on conduction of electricity through gases. He proposed that atoms were made up of
electrons embedded in a uniform matrix of protons. The total positive charge was balanced by the
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Atomic Structure
total negative charge. If electrons were removed, the remaining ion was left with excess positive
charge.
This became known as Thomson’s Plum Pudding model (or watermelon model).
The watermelon model was the accepted model for the structure of the atom until Ernest Rutherford’s
alpha scattering experiments in 1911.However his model has no experimental evidence.
RUTHERFORD’S ATOMIC MODEL: Rutherford (1911) bombarded a thin gold foil with
α – particles. It was observed that
(a) Most of the α – particles passed undeflected.
(b) Very few α – particles underwent small and large deflections.
(c) A very very few (1 out of 100,000) even returned on the original path.
Based on these observations Rutherford proposed his model of atom.
i) Atoms are spherical in shape and mostly hollow
(ii) The central part of the atom is positively charged. It is massive and extremely small in size. This
part is called the ‘nucleus’ of the atom.
(iii) Electrons revolve around the nucleus, like planets around the sun
Note:
-12
The diameter of the nucleus is of the order of 10
-8
-13
to 10
cm. It is very small as
-8
compared to the size of the atom( ≈ 10 cm) [ 1 x 10 cm = 1Å ]
Drawbacks of Rutherford Model:
1.
According to Classical laws of physics (or classical electrodynamics) a revolving electron should
lose energy continuously and fall into the nucleus through a spiral path and the atom must be
destroyed, however, the atom is stable.
2.
If the electron loses energy continuously, the atomic spectra should consist of continuous
bands. Experimentally, atomic spectra are made up of discrete spectral lines.
ATOMIC NUMBER (Z):
The number of protons or electrons in an atom is known as its atomic number.
The atom as a whole is neutral in nature i, e., total negative charge contributed by electrons is equal to
total positive charge contributed by protons,
i, e.,
Number of electrons = Number of protons
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MASS NUMBER (A):
Since electrons have negligible mass, the protons and neutrons are mass particles. The sum of
number of protons and neutrons in the nucleus of an atom is termed mass number.
General symbol for an atom of an element (E) indicating its mass number (A) and atomic number(Z)
A
ZE
;
C
Example : 6C12
A (mass number) = Z (Atomic number or Number of protons) + Number of neutrons
A-Z = Number of neutrons
Note: Mass number is always whole number but mass of proton or a neutron is not a whole number so
atomic mass (or) weight is not necessarily a whole number. Atomic mass of an element is not
equal to sum of masses of all the protons and neutrons simply. Most of the elements in nature
exist as mixture of isotopes. For these elements atomic mass is calculated on the bases of their
abundance in nature and atomic masses of individual members.
Isotopes: Atoms of same element which have same atomic number but different mass numbers are
called isotopes
Example: Cl atomic weight is 35.5 due to existence of Cl 35 and Cl 37 , which are present in the ratio of
3:1 in nature
Most of our information about the arrangement of electrons in atoms has come from studies of the
interaction of matter with light. To understand the nature of these interactions, we shall consider
electromagnetic radiations.
ELECTROMAGNETIC RADIATIONS
The theory of electromagnetic radiations(classical electromagnetism or classical electrodynamics) was
developed by James Clerk Maxwell. Electromagnetic radiations are usually treated as wave motions.
The electronic and magnetic fields oscillate in directions perpendicular to each other and to the
direction of motion of the wave. These waves travel as a continuous sequence of alternating crests
and troughs Ordinary light rays, X-rays,γ-rays are some examples of electromagnetic radiations. All
types of electromagnetic radiations travel through space with the same velocity i.e. 3 x 1010 cm sec –1
8
–1
or 3 x 10 m sec .
There are three fundamental characteristics associated with wave motion. They are i) wavelength (λ)
ii) frequency (v) and iii) velocity (c)
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Wavelength. The distance between two consecutive crests or troughs is known as wavelength
denoted by λ.
Frequency: The number of waves that pass through a given point in 1 second is called its frequency
(number of waves per sec) denoted by v.
Velocity: The distance traveled by wave in 1 second is its velocity c
Velocity= frequency X wavelength
c=v λ
Electromagnetic spectrum
The arrangement of the various types of electromagnetic radiations in order of their increasing(or
decreasing)wavelengths or(frequencies)is known as electromagnetic spectrum.
Wavelength increases →
Cosmic rays γ - rays
X-rays
UV-rays visible infrared(IR)
Micro waves
Radio waves
Frequency decreases →
SPECTRA OR SPECTRUMS
Spectrum is the impression produced on a photographic film when the radiations of particular
wavelengths are analysed with a spectroscope (a device in which a beam of light is passed through a
prism and received on a photograph). It is broadly of two types.
Emission spectrum. Spectrum produced by emitted radiation is known as emission spectrum. it is
again of two types.
i) Continuous spectrum: The spectrum consists of continuous bands of radiations corresponding to
different wavelengths.
Example: solar spectrum. [Spectrum of sunlight]
ii) Line or atomic spectrum: The spectrum consists of a series of sharp lines; each line corresponds
to a particular wavelength.
Example; if an electric discharge is passed through hydrogen gas taken in a discharge tube under low
pressure, the bright light is emitted and the emitted radiation is analysed with spectroscope. It is found
to consist of a series of sharp lines called hydrogen emission spectrum
Absorption spectrum. Spectrum produced by the absorbed radiations is called absorption spectrum.
QUANTUM THEORY
Quantum theory is a fundamental branch of theoretical physics that replaces classical
electromagnetism at the atomic and subatomic levels. The discovery(Max planck,1901) that waves
could be measured in particle-like small packets of energy called quanta and frequency of radiation(v)
is directly proportional to energy of radiation(E) (E=hv where h is called planck’s constant) led to the
branch of physics that deals with atomic and subatomic systems which we today call Quantum theory.
The foundations of quantum theory were established during the first half of the 20th century by Max
Planck, Albert Einstein, Niels Bohr, Werner Heisenberg, Erwin Schrödinger, Max Born, John von
Neumann, Paul Dirac, Wolfgang Pauli and others. Some fundamental aspects of the theory are still
actively studied
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BOHR’S ATOMIC MODEL
Bohr’s atomic model is based on quantum theory of radiation and the classical laws of
physics. The important points of Bohr model are:
(i)
Electrons revolve around the nucleus in closed circular paths called orbits or shells. As long
as the electron stays in a given orbit it does not radiate energy. Therefore these orbits
called “stationary orbits or stationary shells”.
(ii)
Each stationary orbit is associated with definite amount of energy. These orbits are
designated by K, L, M, N, O-------from the nucleus. The orbit close to nucleus has less
energy compared to the orbit away from the nucleus.
(iii)
Energy is emitted in the form of radiation when an electron jumps from outer orbit to inner
orbit. Energy is absorbed when electron jumps from lower orbit to higher orbit
Ehigher - Elower =hv
Where h is Planck’s constant and v is the frequency of radiation.
(iv)
The angular momentum of the revolving electron is an integral multiple of
momentum of the electron is quantized]
i, e.,
mvr = n h
1
–
where, n is integer (n =1,2,3,4…………)
+
m = mass of the electron
Electron in its most stable
‘ground state’ orbit, i.e.
principal quantum number 1
v = velocity of the electron
r = radius of the circular orbit
h = Planck’s constant
Advantages of Bohr’s theory :
Bohr successfully calculated the radii and energies of various orbits of hydrogen
2 2
n h __
4π2mZe2
r = 0.529 x 10-8 x n2 cm
For hydrogen atom,
2
= 0.529 x n Å
rn = r1 x n2 Å
or
/2π, [Angular
3 etc
2
2π
Radius of an orbit =
h
+
2+
3+
and for hydrogen-like mono electron species such as He ,Li , Be
2
rn = 0.529 x n Å
Z
Where Z = Atomic number of the species.
2
4
- 2 π me
Energy of an electron in the nth orbit =
n2h2
For hydrogen atom
-19
En = -21.79 x 10
J per atom
2
n
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= -13.6 eV per atom
2
n
= -313.6 kcal per mole
n2
2
2
and
En = energy of hydrogen first orbit x Z /n for Hydrogen – like mono electron species such as
+
2+
He , Li and Be3+.
Bohr’s model explains the stability of the atom. The frequencies of spectral lines calculated from
Bohr’s equation are in close agreement with the frequencies observed experimentally in hydrogen
spectrum. The spectrum of hydrogen-like ions can also be explained.
Defects of Bohr’s theory: (i) It fails to explain the spectra of multi – electron atoms.
(ii) It fails to explain fine spectrum (spectral lines consist of several closely packed lines) of hydrogen.
(iii) It does not provide an explanation why angular momentum of the electron should always be an
integral multiple of h/2 π.
(iv) It does not explain splitting of spectral lines into a group of finer lines under the influence of
magnetic field (Zeeman effect) and electric field (stark effect).
MODERN THEORY OF ATOMIC STRUCTURE
Louis de Broglie proposed the particle and wave nature (dual nature) of the electron. Based on the
dual nature of the electron, the quantization of its angular momentum is explained. According to the
uncertainity principle proposed by Heisenberg it is impossible to trace path (orbit) traversed by an
electron, the only possibility is maximum probability or relative chance of finding electron in space
around the nucleus(orbital). Erwin Schrodinger proposed the wave equation for the electron. From the
Schrodinger Wave Equation the basic information about quantum numbers and overall electron
behavior have been derived.
Atomic orbital is the space around the nucleus in which relative chance (probability) of finding
electron is maximum (95%).
QUANTUM NUMBERS :
Quantum numbers are used to locate a particular electron in an atom and there are four types of
quantum number which give a complete picture of an electron.
Principal Quantum Number
 It was proposed by Bohr, denoted by n.
 It will have any integer value except zero.
 It gives the size of the orbit and energy of the orbit.
 As the value of n is increasing the size and energy of the orbit increases.
 The number of electrons that can present in an orbit is equal to 2n2.
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Azimuthal Quantum Number
To explain fine spectrum of hydrogen Sommerfeld introduced the idea of elliptical orbits. The angular
momentum of the electron in an elliptical path is quantised
mvr=k h/2π, where k is Sommerfeld’s
integer called azimuthal quantum number.
K= 1,2,3,4…………….n.
Bohr quantized the size of the orbit (major axis of the path) and Sommerfeld quantised the shape of
the orbit (minor axis)
Both the principal quantum number n and azimuthal quantum number k are related to the ellipse by
n/k= length of the major axis/length of the minor axis.
Sommerfeld proposed that, for a given n value, a set of k values1 to n, are possible means for a given
stationary orbit or shell a set of “sub- orbits” or “sub-shells” are present.
Minor axis
Major axis
Note: From quantum mechanics it was found that the azimuthal quantum number can take up values
starting from 0 to n-1 but not 1 to n and azimuthal quantum number is represented by ‘l’ instead of k.
Azimuthal quantum number gives information regarding the shape of sub orbit or subshell and the
number of subshells in a shell. It can have values from 0 to (n-1), i, e., l = 0 (s- subshell), l = 1
(p-subshell), l = 2(d-subshell), l = 3(f-subshell).
(i)
Magnetic quantum number (m): proposed by Lande to explain Zeeman effect and stark
effect.It describes the orientations of the subshells. It can have values from –l to +l
including zero, i, e., total (2l+1) values. Each value corresponds to an orbital.
a) For l = 0 (s sub – shell) , m = value is 0. Hence there is only one orientation for the s
sub – shell.
b) For l = 1 (p sub – shell), m = values are 0,-1 and +1. Hence three orientations are
possible for the p sub – shell. The three corresponding orbitals are written as px, py
and pz.
For l = 2 (i.e. d sub – shell), m = values are –2, -1, 0, +1 and +2. Hence d sub – shell can have five
different orientations, and orbitals corresponding to these are dxy, dyz, dzx, d x2-y2 dz2
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Spin Quantum Number
 It was proposed by Uhlenbeck and Goudsmit.
 It is denoted by 's'.
 Electron moving in an orbital can spin on its own axis.
 The spin of the electron may be clockwise or anti clock wise.
 The clockwise spin is denoted by +1/2 or  and anticlockwise spin is denoted

by –1/2 or  .
Shapes of atomic orbitals
S-orbital
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Various Shells, sub-shells and their quantum numbers
Principal
energy
level (n)
1
Total number of sub
energy levels and their
azimuthal quantum
numbers or
designation (l)
The number of orbitals per
sub – level and their magnetic
quantum numbers (m)
One [0 or 1s]
Number of
orbitals per
energy sublevel
(per energy level,
2
n ).
One (0)
1] (1)
0 or 2s
1 or 2p
One(0)
Three (-1, 0,+1)
1
0 or 3s
1 or 3p
2 or 3d
One (0)
1
Three (-1, 0,+1)
3
Five (-2, -1, 0, +1, +2)
5
2
Two
3
(4)
3
Three
(9)
4
Four
0 or 4s
1 or 4p
2 or 4d
3 or 4f
1
One (0)
3
Three (-1, 0, +1)
5
Five (-2, -1, 0, +1,+2)
7
(16)
Seven (-3, -2, -1, 0, +1, +2, +3)
ELECTRONIC CONFIGURATION OF ATOMS.
Electronic configurations are written in nlx method where n is the principal quantum number, l is the
azimuthal quantum number and x is the number of electrons present in it .To describe the
arrangement and distribution of electrons, following selective principles are required.
Pauli’s Exlusion Principle:
No two electrons in an atom can have all four quantum numbers same:
e.g
correct for ns2
e.g
incorrect for ns
2
Following results can be inferred from Pauli’s exclusion principle.
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2
a). Maximum number of electrons in an orbit can be 2n
b). Maximum number of electrons in a sub-shell can be 2, 6, 10, 14 in s, p, d, f
respectively.
c). Maximum number of electrons in an orbital is 2 only.
Aufbau Principle: Aufbau is a German word meaning building up. This gives us a sequence in
which various subshells are filled up depending on the relative order of the energy of orbitals.
“Electrons are added progressively to the various orbitals in the order of increasing energy starting
with the orbital of lowest energy”
(n + l) rule:
a.
The orbital with lower values of (n + 1) possesses lower energy relative and should be filled
first.
1s2 2s2 2p6 3s2 3p6 3d1
e.g 19K
2
2
6
2
6
1s 2s 2p 3s 3p
wrong
1
4s correct
n + 1 of 4s = 4 + 0 = 4
n + 1 of 3d = 3 + 2 = 5
b.
If (n + 1) is same for two orbitals the orbital with lower values of ‘n’ possess
and should be filled first.
e.g 21Sc
1s2 2s2 2p6 3s2 3p6 4s2 4p1
2
2
6
2
6
2
1
1s 2s 2p 3s 3p 4s 3d
lower energy
wrong
correct
n + 1 of 4p = 4 + 1 = 5
n + 1 of 3d = 3 + 2 = 5
Thus, 3d should be filled first
‘n’ of 3d < ‘n’ of 4p
A sub-shell having nearly full filled or nearly half filled configuration tends to acquire exactly full filled or
exactly half filled nature in order to attain stability
Example L
e.g 24Cr
[Ar] 3d4 4s2 wrong
1
5
9
2
[Ar] 4s 3d
e. g 29Cu [Ar] 3d 4s
1
10
[Ar] 4s 3d
e. g 46Pd
9
2
[Kr] 4d 5s
1
10
[Kr] 5s 4d
Correct
wrong
Correct
wrong
Correct
Note: Most of the exceptions to the electron configuration predicted from the aufbau principle occur
among elements with atomic numbers larger than 40
Ground state and Excited state of atom: when all the electrons present in atom are in the lowest
possible energy states, it is called as ground state of atom. When the electrons jump from a lower
shell to higher, it comes in a higher energy state and is said to be in excited state.
Hund’s Rule: In an atom electron pairing takes place only, after all the available degenerate orbitals
are half-filled with electrons of parallel spin.
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WORKED EXAMPLES
1.
An oil drop has 6.39 x 10-19 C charge. Find out the number of electrons in this drop.
-19
charge on oil drop = 6.39 x 10
-19
1.602 x 10
.
C now we know that
C is charge on one electron
-19
. . 6.39 x 10
-19
C will be charge on = 6 .39 x 10
-19
1.602 x 10
= 4 electrons
2.
(a) Write the electronic configuration of elements with atomic numbers 19, 28 and 29.
(b) Calculate the atomic number and name the element that corresponds to each of the
following electronic configuration.
2
2
6
2
6
1
(i)
1s 2s 2p 3s 3p 4s
(ii)
1s2 2s2 2p6 3s2 3p6 4s1 3d5
(iii)
1s2 2s2 2p6 3s2 3p6 4s1 3d10
(a) Electronic configuration of elements with atomic number
19
1s2 2s2 2p6 3s2 3p6 4s1
28
1s2 2s2 2p6 3s2 3p6 4s2 3d8
29
1s2 2s22p6 3s2 3p6 4s1 3d10
(b) (i) Atomic number of the element is 2 + 2 + 6 + 2 + 6 + 1 = 19
Therefore, the element is potassium.
(ii) Atomic number of the element is 2 + 2 + 6 + 2 + 6 +5 + 1= 24
Therefore, the element is chromium.
(iii) Atomic number of element is 2 + 2 + 6 + 2 + 6 +10 + 1 = 29
Therefore, the element is copper.
3.
(a) An electron is in a 4f orbital. What possible values for the quantum numbers n, l,m
and can it have?
(b) What designation is given to an orbital having (i) n = 2, l = 1 and (ii) n = 4, l = 0?
(a) For an electron in a 4f orbital,
n = 4, l = 3, m = -3, -2, –1, 0, 1, +2, +3, s = +1/2 and –1/2 for each value of m.
(b) For an orbital having n = 2, l = 1 designation is 2p.
For an orbital having n = 4, l = 0 designation is 4s.
4.
A neutral atom has 2K, 8L, 5M electrons. Find out the following from the data: a) atomic
number (b) total number of s electrons (c) total number of p electrons (d) number of
protons in the nucleus.
(a) Atomic number = Number of protons = Number of electrons
Total number of electrons = 2 + 8 + 5 = 15
Hence atomic number = 15
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(b) Total number of s electrons. To find out it, we are to write electronic configuration of Atomic
number = 15.
1s2 2s2 2p6 3s2 3p3
 Total number of s electrons = 6
(c) Total number of p electrons = 9
(d) Number of protons in the nucleus = Number of electrons
 number of protons = 15
5.
Why chlorine has fractional atomic weight?
35
The fractional atomic weight (35.5) of chlorine is due to the fact that in the ordinary chlorine Cl
37
and Cl isotopes are present in the ratio of 3:1
 Average atomic weight of Cl =
3 x 35 + 1 x 37
4
= 35.5 amu. (approximately)
6.
When α- particles are sent through a thin metal foil, very few α- particles returned on the
original path explain.
The whole of the atomic mass is concentrated in the nucleus i.e. the central nucleus is rigid and
hence α- particles which strike on it are returned to the original path.
7.
An element has two naturally occurring isotopes of atomic masses 63 and 65 its atomic
weight is 63.55. Find out the percentage abundances for two isotopes?
Let the percentage abundance of the isotope of mass 63 be x. Therefore, the percentage
abundance of the other isotope is 100 – x. Thus we have
63 x x+ 65 (100 – x )
= 63.55
100
63x + 6500 – 65x = 6355
x = 72. 5
Thus, the percentage abundance of the isotope of atomic mass 63 is 72.5% and the percentage
abundance of the isotope of atomic mass 65 is 27.5 % .
8.
Calculate the radius of the 4th Bohr’s orbit in hydrogen atom?
From Bohr’s theory
rn = 0.529 x n2 Å
Given n = 4
9.
2
=
0.529 x 4 Å
=
8.464 Å
According to Sommerfeld how many elliptical orbits are possible for the fourth orbit
electron.
Given electron is present in 4th orbit
n=4
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For given ‘n’ k values are
k = 1 --------- n n = 4 , k = 1,2,3,4
n
length of major axis
We know
=
k
length of minor axis
if n = 4,k = 4
i.e, length of major axis = length of minor axis then electron path is circular.
In remaining cases for n = 4, k = 1, n = 4 k = 2 & n = 4 k = 3 path of the electron is elliptical
Total 3 elliptical paths are possible for the given electron.
10.
The first excited state refers to the electronic configuration with energy closest to but
higher than that of ground state. Write electronic configurations of the first excited state
of the following
(a) C
(b) Ne
(c) Li
For carbon z is 6
Ground state 1s2 2s2 2p2
First excited state 1s2 2s1 2p3
For Neon z is 10
Ground state 1s2 2s2 2p6
First excited state 1s2 2s2 2p5 3s1
For Lithium z is 3
Ground state 1s2 2s1
First excited state 1s1 2s2
11.
a)
What values are permitted for the orbital Azimuthal quantum number l for an
electron with principal quantum number n =3?
b)
How many different values for the magnetic quantum number are possible for an
electron with Azimuthal quantum number l = 3?
c)
How many electrons can be placed in each of the following subshells: s, p, d, f ?
d)
What is the lowest shell, which has an f subshell?
a) 2,1,0
b) seven (-3, -2, -1, 0, +1, +2, +3)
c) s: 2, p:6, d:10 ,f: 14.
d) Fourth
12.
Which shell would be the first to have a g subshell?
Since the azimuthal quantum number l must be less than the principal quantum number, n, and
since the g subshell designation represents an l value of 4, the minimum n value possible is 5.
In literal notation (K, L, M, N, O) the fifth shell is designed the O shell.
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LEVEL – I
1.
2.
The two electron occupying the same orbital are distinguished by:
a) principal quantum number
b) azimuthal quantum number
c) magnetic quantum number
d) spin quantum number
Filling 3d completely for copper, the next electron enters in
a) 4p – orbital
3.
b) 3d – orbital
b) 1 and 2
th
c) 4.77 Å
d) 2.12 Å
th
b) 1: 4: 9
c) 1: 4: 6
d) 1: 2: 3
th
According to Bohr’s theory, the angular momentum for an electron of 5 orbit is:
b) 2.5h/
c) 5/h
d) 25h/
The maximum number of electrons in a subshell is given by the Expression:
a) 4l + 2
8.
d) 6 and 1
The ratio of radii of 2 , 4 and 6 orbits of hydrogen atom is:
a) 5h/
7.
b) 1.06 Å
nd
a) 2: 4: 6
6.
c) 2 and 0
The Bohr orbit radius for the hydrogen atom (n = 1) is approximately 0.530 Å. The radius for the
first excited state (n = 2) orbit is:
a) 0.13 Å
5.
d) 4d – orbital
For the sixth electron the values of principal and orbital quantum numbers are respectively
a) 2 and 1
4.
c) 4s – orbital
b) 4l – 2
c) 2l + 1
d) 2n2
The combination of quantum numbers that are allowed for an electron in an atom is
a) n = 2, l = 2, m = 1, s = 1/2
b) n = 3, l = 1, m = 1, s = ½
c) n = 5, l = 1, m = 2, s = ½
d) n = 4, l = -1, m = 0, s = ½
9. Four electrons were shown with a set of quantum numbers each. Which set of quantum number
show an electron with highest energy?
10.
a) . n = 3, l = 2, m = +1, s = +1/2
b) n = 4, l = 1, m = 0, s = -1/2
c) n = 4, l = 2, m = -1, s = +1/2
d) n = 5, l = 0, m = 0, s = -1/2
The highest value of e/m of anode rays has been observed when the discharge tube is filled
with:
a) Nitrogen
11.
d) helium
b) –4.5eV
c) –6.8 e V
d) +6.8 e V
Which of the following properties of an element is a whole number?
a) atomic mass
13.
c) hydrogen
The energy of an electron in the first Bohr orbit of H atom is –13.6 eV. The possible energy
values (s) of the excited state (s) for electron is Bohr orbits of hydrogen is (are) :
a) –3.4 e V
12.
b) oxygen
b) atomic number
c) atomic radii
d) atomic volume
7
If the nitrogen atom had electronic configuration 1s , it would have energy lower than that of the
2
2
3
normal ground state configuration 1s 2s 2p , because the electrons would be closer to the
nucleus, yet 1s7 is not observed because it violates
a) Heisenberg uncertainty principal
b) Hund’s rule
c) Pauli exclusion principle
d) Bohr postulate of stationary orbit
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14.
For a particular value of azimuthal quantum number (1), the total number of magnetic quantum
number (m) is given by s. (SAT, 1998)
a) l = (m+1)/2
15.
Atomic Structure
b) l = (2m+1)/2
c) l = (m – 1)/2
d) m = (2l –1)/2
Two electrons A and B in an atom have the following set of quantum numbers :
A : 3,
2,
–2, +1/2
B : 3,
0,
0,
+1/2
Which statement is correct for A and B ?
a) A and B have same energy
c) B has more energy than A
16.
As we move away from the nucleus, the difference between the adjacent energy levels
a) Increases
17.
b) Remains constant
c) Decreases
d) None of these
Existence of fundamental particles violates :
a) Dalton’s atomic theory
c) Mendeleef’s theory
18.
b) A has more energy than B
d) A and B represents same electron
b) be Broglie’s theory
d) Newland’s theory
Which statement does not form part of Bohr’s model of the hydrogen atom ?
a) Energy of the electrons in the orbit is quantized
b) The electron in the orbit nearest the nucleus is in the lowest energy
c) Electrons revolve in different obits around the nucleus
d) The position and velocity of the electrons in the orbit cannot be determined simultaneously
19.
The ratio of specific charge (e/m) of an electron to that of a hydrogen ion is
a) 1 : 1
20.
b) 1840 : 1
c) 1 : 1840
d) 2 : 1
The specific charge for positive rays is much less than the specific charge for cathode rays. This
is because
a) Positive rays are positively charged
b) Charge on positive rays is less
c) Positive rays comprise ionised atoms whose mass is much higher
d) Experimental method for determination is wrong
21.
The configuration 1s2, 2s2 2p5, 3s1 shows
a) Ground state of fluorine
c) Excited state of neon atom
22.
Which are isoelectronic with each other?
a) Na+ and Ne
23.
b) K+ and O
c) Ne and O
d) Na+ and K+
Which atom has as many as s-electron as p-electrons?
a) H
24.
b) Excited state of fluorine
d) Excited state of O2– ion
b) Mg
c) N
d) Na
The number f orbitals that could be associated with a principal quantum number are
a) n2
b) 2n2
c) n2 + 2
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d) n2 – 2
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25.
The atomic weight of an element is 23 and atomic number is 11. The number of protons,
electrons and neutrons respectively present in the atom of the elements are
a) 11, 11, 12
26.
b) 12, 12, 11
b) 1  4
c) 4  3
b) 18
d) 3  2
c) 12
d) 20
The ratio of radius of III and IV Bohr’s orbits in hydrogen atom is
a) 3 : 4
29.
d) 12, 11, 12
Maximum number of electrons which can be accommodated in a g sub-shell is
a) 14
28.
c) 11, 12, 11
Which transition of electron in the hydrogen atom emits maximum energy?
a) 2  1
27.
Atomic Structure
b) 3 : 8
c) 9 : 16
2
2
6
1
d) 8 : 9
1
The electronic configuration 1s , 2s 2p , 3s , 3p correctly describes :
a) Ground state of Na
b) Ground state of Si
+
c) Excited state of Mgd) Excited state of Al3+
30.
Which represent the correct set up of the four quantum numbers of 4s-electron?
a) 4, 3, 2, +1/2
31.
b) 4, 2, 1, 0
b) 4p
c) 5g
b) – 5.40 eV
c) – 0.85 eV
d) – 2.72 eV
An electron has a spin quantum number + ½ and a magnetic quantum number – 1. It can not
be present in
a) d – orbital
35.
b) n = 2, l = 1, m = 1
d) n = 3, l = 0, m ,= 0
The energy of hydrogen atom in its ground states is – 13.6eV. The energy of the level
corresponding to the to the quantum number n = 5 is
a) – 0.54 eV
34.
d) 4d
The correct set of quantum numbers for the unpaired electron of chlorine atom is
a) n = 2, l = 1, m = 0
c) n = 3, l = 1 , m = 1
33.
d) 4, 0, 0, 1/2
The atomic orbital not allowed in quantum theory is
a) 3f
32.
c) 4, 3, –2, +1/2
b) f – orbital
c) p – orbital
d) s – orbital
Assertion A: The angular momentum of an electron in an atom is quantized.
Reason R: In an atom only those orbits are permitted in which angular momentum of the
electron is whole number multiple of h/2π.
1. Both A and R are true and R is the correct explanation of A
2. Both A and R are true and R is not the correct explanation of A
3. A is true R is false
4) A is true R is false
36. Ground state electronic configuration of nitrogen atom can be represented by










a)
b)
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Atomic Structure








c)

d)
37. Which of the following electronic configuration is against to Pauli’s exclusion principle & Aufbau
principle
a)
b)
c)
d)














38. Which of the following electronic configuration is against to Aufbau & Hund’s rule
a)
b)















c)
d)
LEVEL – II
2+
4+
1.
Write down the electronic configuration of Cr
present?
and Mn . How many unpaired electrons are
2.
An element X has two isotopes of atomic masses 16 and 18. Its overall atomic weight is 6.4.
The proportion of the isotope 16 in X ?
3.
Which transitions among the adjacent orbits in hydrogen spectrum show larger difference in
energy?
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Atomic Structure
4.
What values are assigned to quantum numbers n, l and m for
(i) 2s orbital (ii) 2p orbital (iii) 4d orbital
5.
Fourth shell (n = 4) may contain upto 32 electrons. Explain this fact with the help of quantum
numbers.
6.
Bohr’s atomic model is superior than the Rutherford’s atomic model, explain,
1997)
7.
What do you expect in between two electronic orbits of an atom?
8.
Differentiate orbit and orbital. (SAT, 2000)
(SAT,
KEY TO LEVEL – I
1–D
2–C
3–A
4–D
5–B
6–B
7–A
8–B
9–C
10 – C
11 – A
12 – B
13 – C
14 – C
15 – B
16– C
17 – A
18 – D
19 – B
20 – C
21 – C
22 – A
23 – B
24– A
25 – A
26 – A
27 – B
28 – C
29 – C
30 – D
31 – A
32 –C
33 – A
34 – D
35 – A
36 – A
37 – B
38 – B.
WORK SHEET
th
1.
Calculate the 10 Bohr’s orbit in hydrogen atom? Give the formula of rn  ?
2.
What are quantum numbers and explain it?
3.
Differentiate orbit and orbital?
4.
Calculate the wave number corresponding to the following wane lengths?
5.
Cal. The frequency and energy of a photan, wave length is 6200A0 (Plank’s constant
h  6.625 1027 )
6.
What is the relation between velocity, frewney and wave length?
7.
Cal. the frewney of a radiation whose wave length is 2000A
8.
What is the ratio between energies of two radiations whose wave lengths are 6000A0 and
0
2000A ?
9.
Write electronic configaration of cu?
10.
Write electronic configaration of cromivm?
11.
A liquid drop has 6  10
12.
What is 180 topes and give example of hydrogen and also names of hydrogen?
13.
Carbon occurs in nature of c
40
0
c charge find out the number of electrons in this drop?
12
12
percentatage abundance of c
13
and c
the average atomic mass of ‘c’ is 13.022. what is
13
and c
in nature?
14.
The atomic weight of an element 3 and atomic number is 1. the number of protons, electrons
neutrons are -------------------------------
15.
What is the mass of 1 mole of electrons?

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