APEX INSTITUTE FOR IIT-JEE /AIEEE / PMT, 0120-4901457, +919990495952, +919910817866 www.apexiit.co.in/ CONTENTS 2.1 Composition of atom 2.2 Atomic number, Mass number and Atomic species 2.3 Electromagnetic radiation’s 2.4 Atomic spectrum- Hydrogen spectrum 2.5 Thomson's model 2.6 Rutherford's nuclear model 2.7 Planck's Quantum theory and Photoelectric effect 2.8 Bohr’s atomic model 2.9 Bohr – Sommerfeld’s model 2.10 Dual nature of electron 2.11 Heisenberg’s uncertainty principle 2.12 Schrödinger wave equation 2.13 Quantum numbers and Shapes of orbitals 2.14 Electronic configuration principles 2.15 Electronic configurations of elements Assignment (Basic and Advance Level) Answer Sheet of Assignment Science has produced a microscopic structure of the atom, but it’s structure is so detailed and so subtle of something which is far removed from our immediate experience that it is difficult to see how many of its features were constructed. Yet among all the experiments used to form the theory of atomic structure, there stand a few which have been most in-fluential in shaping its major features. Erwin Schrödinger Chapter 2
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CONTENTS
2.1 Composition of atom
2.2 Atomic number, Mass number and Atomic species
2.3 Electromagnetic radiation’s
2.4 Atomic spectrum- Hydrogen spectrum
2.5 Thomson's model
2.6 Rutherford's nuclear model
2.7 Planck's Quantum theory and Photoelectric effect
2.8 Bohr’s atomic model
2.9 Bohr – Sommerfeld’s model
2.10 Dual nature of electron
2.11 Heisenberg’s uncertainty principle
2.12 Schrödinger wave equation
2.13 Quantum numbers and Shapes of orbitals
2.14 Electronic configuration principles
2.15 Electronic configurations of elements
Assignment (Basic and Advance Level)
Answer Sheet of Assignment
Science has produced a microscopic
structure of the atom, but it’s structure is so detailed and so subtle of something which is far removed from our immediate experience that it is difficult to see how many of its features were constructed. Yet among all the experiments used to form the theory of atomic structure, there stand a few which have been most in-fluential in shaping its major features.
Erwin Schrödinger
Chapter 2
John Dalton 1808, believed that matter is made up of extremely minute indivisible particles, called
atom which can takes part in chemical reactions. These can neither be created nor be destroyed. However,
modern researches have conclusively proved that atom is no longer an indivisible particle. Modern structure of
atom is based on Rutherford’s scattering experiment on atoms and on the concepts of quantization of energy.
2.1 Composition of atom.
The works of J.J. Thomson and Ernst Rutherford actually laid the foundation of the modern picture of the
atom. It is now believed that the atom consists of several sub-atomic particles like electron, proton, neutron,
positron, neutrino, meson etc. Out of these particles, the electron, proton and the neutron are called
fundamental subatomic particles and others are non-fundamental particles.
EElleeccttrroonn ((––11eeoo))
(1) It was discovered by J.J. Thomson (1897) and is negatively charged particle. Electron is a component particle of cathode rays.
(2) Cathode rays were discovered by William Crooke's & J.J. Thomson (1880) using a cylindrical
hard glass tube fitted with two metallic electrodes. The tube has a side tube with a stop cock. This tube was
known as discharge tube. They passed electricity (10,000V) through a discharge tube at very low pressure
( 210 − to )10 3 Hgmm− . Blue rays were emerged from the cathode. These rays were termed as Cathode rays.
(3) Properties of Cathode rays
(i) Cathode rays travel in straight line.
(ii) Cathode rays produce mechanical effect, as they can rotate the wheel placed in their path.
(iii)Cathode rays consist of negatively charged particles known as electron.
(iv) Cathode rays travel with high speed approaching that of light (ranging between 910 − to 1110 −
cm/sec)
(v) Cathode rays can cause fluorescence.
Gas at low pressure
Cathode rays
Cathode Anode
TC Vaccum pump
High voltage + –
Discharge tube experiment for production of cathode rays
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(vi) Cathode rays heat the object on which they fall due to transfer of kinetic energy to the object.
(vii) When cathode rays fall on solids such as −XCu, rays are produced.
(viii) Cathode rays possess ionizing power i.e., they ionize the gas through which they pass.
(ix) The cathode rays produce scintillation the photographic plates.
(x) They can penetrate through thin metallic sheets.
(xi) The nature of these rays does not depend upon the nature of gas or the cathode material used in discharge tube.
(xii) The e/m (charge to mass ratio) for cathode rays was found to be the same as that for an −e 81076.1( ×− coloumb per gm). Thus, the cathode rays are a stream of electrons.
Note : q When the gas pressure in the discharge tube is 1 atmosphere no electric current flows
through the tube. This is because the gases are poor conductor of electricity.
q The television picture tube is a cathode ray tube in which a picture is produced due to fluorescence
on the television screen coated with suitable material. Similarly, fluorescent light tubes are also
cathode rays tubes coated inside with suitable materials which produce visible light on being hit
with cathode rays.
(4) R.S. Mullikan measured the charge on an electron by oil drop experiment. The charge on each
electron is .10602.1 19 C−×−
(5) Name of electron was suggested by J.S. Stoney. The specific charge (e/m) on electron was first
determined by J.J. Thomson.
(6) Rest mass of electron is gm28101.9 −× 1837/1000549.0 == amu of the mass of hydrogen atom.
(7) According to Einstein’s theory of relativity, mass of electron in motion is, m ′ ])/(1[
)electron(mof mass Rest2cu−
=
Where u = velocity of electron, c= velocity of light.
When u=c than mass of moving electron =∞.
(8) Molar mass of electron = Mass of electron × Avogadro number = .10483.5 4−×
(9) 1.1 2710× electrons =1gram.
(10) 1 mole electron = 5483.0 mili gram.
(11) Energy of free electron is ≈ 0. The minus sign on the electron in an orbit, represents attraction
between the positively charged nucleus and negatively charged electron.
(12) Electron is universal component of matter and takes part in chemical combinations.
(13) The physical and chemical properties of an element depend upon the distribution of electrons in outer
shells.
(14) The radius of electron is .1028.4 12 cm−×
(15) The density of the electron is = mLg /1017.2 17−× .
Example : 1 The momentum of electron moving with 1/3rd velocity of light is (in g cm sec–1)
(5) Mass of proton = Mass of hydrogen atom= amu00728.1 gram2410673.1 −×= 1837= of the mass of
electron.
(6) Molar mass of proton = mass of proton × Avogadro number 008.1= (approx).
(7) Proton is ionized hydrogen atom )( +H i.e., hydrogen atom minus electron is proton.
(8) Proton is present in the nucleus of the atom and it's number is equal to the number of electron.
(9) Mass of 1 mole of protons is ≈ 1.007 gram.
(10) Charge on 1 mole of protons is ≈ 96500 coulombs.
(11) The volume of a proton (volume = 3
3
4rπ ) is ≈ 1.5 .10 338 cm−×
(12) Specific charge of a proton is 41058.9 × Coulomb/gram.
NNeeuuttrroonn ((oonn11,, NN))
Perforated cathode
Cathode rays TC Vaccum pump
High voltage + –
Perforated tube experiment for production of anode rays
Anode rays
(1) Neutron was discovered by James Chadwick (1932) according to the following nuclear reaction,
(2) The reason for the late discovery of neutron was its neutral nature.
(3) Neutron is slightly heavier (0.18%) than proton.
(4) Mass of neutron = gram2410675.1 −× = kg2710675.1 −× = ≈amu00899.1 mass of hydrogen atom.
(5) Specific charge of a neutron is zero.
(6) Density = ../105.1 14 ccgram−×
(7) 1 mole of neutrons is ≈ 1.008 gram.
(8) Neutron is heaviest among all the fundamental particles present in an atom.
(9) Neutron is an unstable particle. It decays as follows :
nutrino anti
00
electron
01
proton
11
neutron
10 ν++→ − eHn
(10) Neutron is fundamental particle of all the atomic nucleus, except hydrogen or protium.
Comparison of mass, charge and specific charge of electron, proton and neutron
Name of constant Unit Electron(e–) Proton(p+) Neutron(n)
Mass (m)
amu
kg
Relative
0.000546
9.109 × 10–31
1/1837
1.00728
1.673 × 10–27
1
1.00899
1.675 × 10–24
1
Charge(e)
Coulomb (C)
esu
Relative
– 1.602 × 10–19
– 4.8 × 10–10
– 1
+1.602 × 10–19
+4.8 × 10–10
+1
Zero
Zero
Zero
Specific charge (e/m) C/g 1.76 × 108 9.58 × 104 Zero
• The atomic mass unit (amu) is 1/12 of the mass of an individual atom of 126 C , i.e. kg2710660.1 −× .
Other non fundamental particles
Particle Symbol Nature Charge esu
×10–10
Mass
(amu)
Discovered by
Positron ++ β,1, 0ee + + 4.8029 0.0005486
Anderson (1932)
Neutrino ν 0 0 < 0.00002
Pauli (1933) and Fermi (1934)
Anti-proton −p – – 4.8029 1.00787 Chamberlain Sugri (1956) and
Weighland (1955)
Positive mu meson +µ + + 4.8029 0.1152 Yukawa (1935)
Negative mu meson −µ – – 4.8029 0.1152 Anderson (1937)
Positive pi meson +π + + 4.8029 0.1514
Powell (1947) Negative pi meson −π – – 4.8029 0.1514
Neutral pi meson 0π 0 0 0.1454
2.2 Atomic number, Mass number and Atomic species.
1126
42
94 nCHeBe o+→+ or 114
74
211
5 nNHeB o+→+
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(1) Atomic number or Nuclear charge
(i) The number of protons present in the nucleus of the atom is called atomic number (Z).
(ii) It was determined by Moseley as,
where, −= Xν rays frequency
Z= atomic number of the metal
ba & are constant.
(iii) Atomic number = Number of positive charge on nucleus = Number of protons in nucleus =
Number of electrons in nutral atom.
(iv) Two different elements can never have identical atomic number.
(2) Mass number
(i) The sum of proton and neutrons present in the nucleus is called mass number.
Mass number (A) = Number of protons + Number of neutrons or Atomic number (Z)
or Number of neutrons = A – Z .
(ii) Since mass of a proton or a neutron is not a whole number (on atomic weight scale), weight is not
necessarily a whole number.
(iii) The atom of an element X having mass number (A) and atomic number (Z) may be represented
by a symbol,
Note : q A part of an atom up to penultimate shell is a kernel or atomic core.
q Negative ion is formed by gaining electrons and positive ion by the loss of electrons.
q Number of lost or gained electrons in positive or negative ion =Number of protons ± charge on
ion.
(3) Different Types of Atomic Species
Atomic species Similarities Differences Examples
Isotopes
(Soddy)
(i) Atomic No. (Z)
(ii) No. of protons
(iii) No. of electrons
(iv) Electronic
configuration
(v) Chemical properties
(vi) Position in the periodic
table
(i) Mass No. (A)
(ii) No. of neutrons
(iii) Physical properties
(i) HHH 31
21
11 ,,
(ii) OOO 188
178
168 ,,
(iii) ClCl 3717
3517 ,
1−sν
Z
)( bZa −=ν or abaZ −
A
Z
X
Mass number
Atomic number
Element
e.g. 147
168
199 ,, NOF etc.
Isobars
(i) Mass No. (A)
(ii) No. of nucleons
(i) Atomic No. (Z)
(ii) No. of protons, electrons
and neutrons
(iii)Electronic configuration
(iv) Chemical properties
(v) Position in the perodic
table.
(i) CaKAr 4020
4019
4018 ,,
(ii) BaXeTe 13056
13054
13052 ,,
Isotones
No. of neutrons (i) Atomic No.
(ii) Mass No., protons and
electrons.
(iii) Electronic
configuration
(iv) Physical and chemical
properties
(v) Position in the periodic
table.
(i) SPSi 32
16
31
15
30
14 ,,
(ii) CaK 4020
3919 ,
(iii) HeH 42
31 ,
(iv) NC 147
136 ,
Isodiaphers
Isotopic No.
(N – Z) or (A – 2Z)
(i) At No., mass No.,
electrons, protons,
neutrons.
(ii) Physical and chemical
properties.
(i) 23190
23592 , ThU
(ii) 199
3919 , FK
(iii) 5524
6529 , CrCu
Isoelectronic
species
(i) No. of electrons
(ii) Electronic
configuration
At. No., mass No. (i) )22(,, 22−− eCNOCOON
(ii) )14(,, 2−− eNCNCO
(iii) )2(,,, 2 −++− eBeLiHeH
(iv)
)18(,,,, 223 −++−−− eCaandKArClSP
Isosters
(i) No. of atoms
(ii) No. of electrons
(iii) Same physical and
chemical properties.
(i) 2N and CO
(ii) 2CO and ON 2
(iii) HCl and 2F
(iv) CaO and MgS
(v) 66 HC and 633 HNB
Note : q In all the elements, tin has maximum number of stable isotopes (ten).
q Average atomic weight/ The average isotopic weight
100
isotope of 2nd mass relative isotope of 2nd % isotope 1stof mass relative isotope 1stof % ×+×=
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Example : 4 The characteristics X- ray wavelength for the lines of the αk series in elements X and Y are 9.87Å and
2.29Å respectively. If Moseley’s equation )75.0(109.4 7 −×= Zν is followed, the atomic numbers of X and
Y are
(a) 12, 24 (b) 10, 12 (c) 6, 12 (d) 8, 10
Solution : (a) λ
νc
=
8
10
8
105132.51087.9
103×=
×
×=
−xν
8
10
8
104457.111029.2
103×=
×
×=
−yν
using Moseley’s equation we get
)75.0(109.4105132.5 78 −×=×∴ xZ …..(i)
and )75.0(1090.4104457.11 78 −×=× yZ ….. (ii)
On solving equation (i) and (ii) .24,12 == yx ZZ
Example : 5 If the straight line is at an angle 45° with intercept, 1 on axis,−ν calculate frequency ν when atomic
number Z is 50.
(a) 20001−s (b) 2010
1−s (c) 24011−s (d) None
Solution : (c) a==°= 145tanν
ab=1
49150 =−=∴ ν
.2401 1−= sν
Example : 6 What is atomic number Z when 12500 −= sν ?
(a) 50 (b) 40 (c) 51 (d) 53
Solution : (c) .51,12500 =−== ZZν
Example : 7 Atomic weight of Ne is 20.2. Ne is a mixutre of 20Ne and 22Ne . Relative abundance of heavier isotope
is
(a) 90 (b) 20 (c) 40 (d) 10
Solution:(d) Average atomic weight/ The average isotopic weight
100
isotope of 2nd mass relative isotope of 2nd % isotope 1stof mass relative isotope 1stof % ×+×=
Examples based on Moseley equation
1−sν
Z
θ
a= θtan ab=intercept
Examples based on Atomic number, Mass number and Atomic species
ZXA
100
22)100(202.20
×−+×=∴
aa; 90=∴ a ; per cent of heavier isotope 1090100 =−=
Example : 8 The relative abundance of two isotopes of atomic weight 85 and 87 is 75% and 25% respectively. The
average atomic weight of element is
(a) 75.5 (b) 85.5 (c) 87.5 (d) 86.0
Solution:(b) Average atomic weight/ The average isotopic weight
100
isotope of 2nd mass relative isotope of 2nd % isotope 1stof mass relative isotope 1stof % ×+×=
5.85100
25877585=
×+×=
Example : 9 Nitrogen atom has an atomic number of 7 and oxygen has an atomic number of 8. The total number of
electrons in a nitrate ion is
(a) 30 (b) 35 (c) 32 (d) None
Solution : (c) Number of electrons in an element = Its atomic number
So number of electrons in N=7 and number of electrons in O=8.
Formula of nitrate ion is −3NO
So, in it number of electrons
×= 1 number of electrons of nitrogen ×+3 number of electrons of oxygen +1 (due to negative charge)
3218371 =+×+×=
Example :10 An atom of an element contains 11 electrons. Its nucleus has 13 neutrons. Find out the atomic
number and approximate atomic weight.
(a) 11, 25 (b) 12, 34 (c) 10, 25 (d) 11, 24
Solution : (d) Number of electrons =11
∴ Number of protons = Number of electron =11
Number of neutrons = 13
Atomic number of element = Number of proton = Number of electrons =11
Further, Atomic weight = Number of protons + Number of neutrons =11 + 13=24
Example : 11 How many protons, neutrons and electrons are present in Pa 3115)( Arb 40
18)( Agc 10847)( ?
Solution : The atomic number subscript gives the number of positive nuclear charges or protons. The neutral atom
contains an equal number of negative electrons. The remainder of the mass is supplied by neutrons.
Atom Protons Electrons Neutrons
P3115 15 15 31 – 15=16
Ar4018 18 18 40 – 18=22
Ag10847 47 47 108 – 47=61
Example :12 State the number of protons, neutrons and electrons in 12C and .14C
Solution : The atomic number of 12C is 6. So in it number of electrons = 6
Number of protons =6; Number of neutrons =12 – 6=6
The atomic number of 14C is 6. So in it number of electrons = 6
Number of protons = 6; Number of neutrons =14 – 6=8
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Wavelength Crest
Energy
Trough
Vibrating source
Example :13 Predict the number of electrons, protons and neutrons in the two isotopes of magnesium with
atomic number 12 and atomic weights 24 and 26.
Solution : Isotope of the atomic weight 24, i.e. .2412 Mg We know that
Number of protons = Number of electrons =12
Further, Number of neutrons = Atomic weight – Atomic number =24 – 12 =12
Similarly, In isotope of the atomic weight 26, i.e. 2612 Mg
Number of protons = Number of electrons =12
Number of neutrons = 26 – 12 = 14
2.3 Electromagnetic Radiations.
(1) Light and other forms of radiant energy propagate without any medium in the space in the form of
waves are known as electromagnetic radiations. These waves can be produced by a charged body moving in a
magnetic field or a magnet in a electric field. e.g. −α rays, −γ rays, cosmic rays, ordinary light rays etc.
(2) Characteristics : (i) All electromagnetic radiations travel with the velocity of light. (ii) These consist
of electric and magnetic fields components that oscillate in directions perpendicular to each other and
perpendicular to the direction in which the wave is travelling.
(3) A wave is always characterized by the following five characteristics:
(i) Wavelength : The distance between two
nearest crests or nearest troughs is called
the wavelength. It is denoted by
λ (lambda) and is measured is terms of
centimeter(cm), angstrom(Å), micron( µ )
or nanometre (nm).
mcmÅ 108 10101 −− ==
mcm 64 10101 −− ==µ
mcmnm 97 10101 −− ==
nmÅcm 748 1010101 === µ
(ii) Frequency : It is defined as the number of waves which pass through a point in one second. It is
denoted by the symbol ν (nu) and is expressed in terms of cycles (or waves) per second (cps) or hertz
(Hz).
=λν distance travelled in one second = velocity =c
λν
c=
(iii) Velocity : It is defined as the distance covered in one second by the wave. It is denoted by the
letter ‘c’. All electromagnetic waves travel with the same velocity, i.e., .sec/103 10 cm×
sec/103 10 cmc ×== λν
Thus, a wave of higher frequency has a shorter wavelength while a wave of lower frequency has a
longer wavelength.
(iv) Wave number : This is the reciprocal of wavelength, i.e., the number of wavelengths per
centimetre. It is denoted by the symbol ν (nu bar). It is expressed in 11 or −− mcm .
λν
1=
(v) Amplitude : It is defined as the height of the crest or depth of the trough of a wave. It is denoted by
the letter ‘A’. It determines the intensity of the radiation.
The arrangement of various types of electromagnetic radiations in the order of their increasing or decreasing wavelengths or frequencies is known as electromagnetic spectrum.
Name Wavelength (Å) Frequency (Hz) Source
Radio wave 714 103103 ×−× 95 101101 ×−× Alternating current of high
Cosmic Rays 0.01- zero −× 20103 infinity Outer space
2.4 Atomic spectrum - Hydrogen spectrum.
AAttoommiicc ssppeeccttrruumm
(1) Spectrum is the impression produced on a photographic film when the radiation (s) of particular
wavelength (s) is (are) analysed through a prism or diffraction grating. It is of two types, emission and absorption.
(2) Emission spectrum : A substance gets excited on heating at a very high temperature or by giving
energy and radiations are emitted. These radiations when analysed with the help of spectroscope, spectral lines are obtained. A substance may be excited, by heating at a higher temperature, by passing electric current at a very low pressure in a discharge tube filled with gas and passing electric current into metallic filament. Emission spectra is of two types,
(i) Continuous spectrum : When sunlight is passed through a prism, it gets dispersed into continuous
bands of different colours. If the light of an incandescent object resolved through prism or spectroscope, it also gives continuous spectrum of colours.
(ii) Line spectrum : If the radiations obtained by the excitation of a substance are analysed with help of a spectroscope a series of thin bright lines of specific colours are obtained. There is dark space in between two consecutive lines. This type of spectrum is called line spectrum or atomic
spectrum..
(3) Absorption spectrum : When the white light of an incandescent substance is passed through any
substance, this substance absorbs the radiations of certain wavelength from the white light. On analysing the transmitted light we obtain a spectrum in which dark lines of specific wavelengths are observed. These lines
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Positively charged sphere
Electron –
–
–
– –
–
+ +
+ +
+ +
+
Positive charge spreaded throughout the sphere
constitute the absorption spectrum. The wavelength of the dark lines correspond to the wavelength of light
absorbed.
HHyyddrrooggeenn ssppeeccttrruumm
(1) Hydrogen spectrum is an example of line emission spectrum or atomic emission spectrum.
(2) When an electric discharge is passed through hydrogen gas at low pressure, a bluish light is emitted.
(3) This light shows discontinuous line spectrum of several isolated sharp lines through prism.
(4) All these lines of H-spectrum have Lyman, Balmer, Paschen, Barckett, Pfund and Humphrey series.
These spectral series were named by the name of scientist discovered them.
(5) To evaluate wavelength of various H-lines Ritz introduced the following expression,
−===
2
2
2
1
111
nnR
c
ν
λν
Where R is universal constant known as Rydberg’s constant its value is 109, 678 1−cm .
2.5 Thomson's model.
(1) Thomson regarded atom to be composed of positively charged protons and negatively charged
electrons. The two types of particles are equal in number thereby making atom
electrically neutral.
(2) He regarded the atom as a positively charged sphere in which
negative electrons are uniformly distributed like the seeds in a water melon.
(3) This model failed to explain the line spectrum of an element and the
scattering experiment of Rutherford.
2.6 Rutherford's nuclear model.
(1) Rutherford carried out experiment on the bombardment of thin (10–4 mm) Au foil with high speed
positively charged −α particles emitted from Ra and gave the following observations, based on this experiment
:
(i) Most of the −α particles passed without any deflection.
(ii) Some of them were deflected away from their path.
(iii) Only a few (one in about 10,000) were returned back to their original direction of propagation.
(iv) The scattering of −α particles .
2sin
1
4
∝
θ
b
θ
r0 Nucleus
α-particle (energy E eV)
Scattering of α -particle
(2) From the above observations he concluded that, an atom consists of
(i) Nucleus which is small in size but carries the entire mass i.e. contains all the neutrons and
protons.
(ii) Extra nuclear part which contains electrons. This model was similar to the solar system.
(3) Properties of the Nucleus
(i) Nucleus is a small, heavy, positively charged portion of the atom and located at the centre of the
atom.
(ii) All the positive charge of atom (i.e. protons) are present in nucleus.
(iii) Nucleus contains neutrons and protons, and hence these particles collectively are also referred to as
nucleons.
(iv) The size of nucleus is measured in Fermi (1 Fermi = 10–13 cm).
(v) The radius of nucleus is of the order of .105.1 13 cm−× to .105.6 13 cm−× i.e. 5.1 to 5.6 Fermi.
Generally the radius of the nucleus ( )nr is given by the following relation,
This exhibited that nucleus is 510 − times small in size as compared to the total size of atom.
(vi) The Volume of the nucleus is about 3910 − 3cm and that of atom is ,10 324 cm− i.e., volume of the
nucleus is 1510 − times that of an atom.
(vii) The density of the nucleus is of the order of 31510 −cmg or 810 tonnes 3−cm or cckg /10 12 . If
nucleus is spherical than,
(4) Drawbacks of Rutherford's model
(i) It does not obey the Maxwell theory of electrodynamics, according to it “A small charged
particle moving around an oppositely charged centre continuously loses its energy”. If an electron
does so, it should also continuously lose its energy and should set up spiral motion ultimately
failing into the nucleus.
(ii) It could not explain the line spectra of −H atom and discontinuous spectrum nature.
Size of the nucleus = 1 Fermi = 10–15 m Size of the atom 1 Å = 10–10 m
Nucleus +
10–15 m
10–10 m
Planetry electron
–
3/113 )104.1( Acmrr on ××== −
Density = = nucleus of the volume
nucleus of the mass
323
3
410023.6
number mass
rπ××
Unstability of atom
e–
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Example:14 Assuming a spherical shape for fluorine nucleus, calculate the radius and the nuclear density of fluorine
nucleus of mass number 19.
Solution : We know that,
3/113 )104.1( Ar −×= 3/113 19104.1 ××= − cm131073.3 −×= (A for F=19)
Volume of a fluorine atom 3
3
4rπ= 313 )1073.3(14.3
3
4 −×××= 3371018.2 cm−×=
number sAvogadro'
nucleusof mol oneof Mass nucleus singleof Mass = g
2310023.6
19
×=
Thus nucleus singleof Volume
nucleus singleof Mass nucleusof Density =
3723 1018.2
1
10023.6
10−×
××
= 11310616.7 −== cmg
Example: 15 Atomic radius is the order of ,10 8 cm− and nuclear radius is the order of .10 13 cm− Calculate what
2.7 Planck's Quantum theory and Photoelectric effect.
PPllaanncckk''ss QQuuaannttuumm tthheeoorryy
(1) Max Planck (1900) to explain the phenomena of 'Black body radiation' and 'Photoelectric effect' gave quantum theory. This theory extended by Einstein (1905).
(2) If the substance being heated is a black body (which is a perfect absorber and perfect radiator of energy) the radiation emitted is called black body radiation.
(3) Main points
(i) The radiant energy which is emitted or absorbed by the black body is not continuous but
discontinuous in the form of small discrete packets of energy, each such packet of energy is called a 'quantum'. In case of light, the quantum of energy is called a 'photon'.
(ii) The energy of each quantum is directly proportional to the frequency (ν ) of the
radiation, i.e.
ν∝E or λ
νhc
hE ==
where, =h Planck's constant = 6.62×10–27 erg. sec. or .sec1062.6 34 Joules−×
(iii) The total amount of energy emitted or absorbed by a body will be some whole number quanta.
Hence ,νnhE = where n is an integer.
(iv) The greater the frequency (i.e. shorter the wavelength) the greater is the energy of the radiation.
thus, 1
2
2
1
2
1
λ
λ
ν
ν==
E
E
Examples based on Properties of the nucleus
+
(v) Also ,21 EEE += hence, 21 λλλ
hchchc+= or
21
111
λλλ+= .
Example: 16 Suppose J1710 − of energy is needed by the interior of human eye to see an object. How many photons of
green light )550( nm=λ are needed to generate this minimum amount of energy
(a) 14 (b) 28 (c) 39 (d) 42 Solution : (b) Let the number of photons required =n
1710 −=λ
hcn ; 286.27
10310626.6
105501010834
91717
==×××
××=
×=
−
−−−
hcn
λ photons
Example: 17 Assuming that a 25 watt bulb emits monochromatic yellow light of wave length .57.0 µ The rate of
(1) When radiations with certain minimum frequency )( 0ν strike the surface of a metal, the electrons are
ejected from the surface of the metal. This phenomenon is called photoelectric effect and the electrons emitted are called photo-electrons. The current constituted by photoelectrons is known as photoelectric current.
(2) The electrons are ejected only if the radiation striking the surface of the metal has at least a minimum
frequency )( 0ν called Threshold frequency. The minimum potential at which the plate photoelectric current
becomes zero is called stopping potential. (3)The velocity or kinetic energy of the electron ejected depend upon the frequency of the incident radiation and is independent of its intensity.
(4) The number of photoelectrons ejected is proportional to the intensity of incident radiation. (5) Einstein’s photoelectric effect equation : According to Einstein, Maximum kinetic energy of the ejected electron = absorbed energy – threshold energy
−=−=
0
0
2
max
11
2
1
λλνν hchhmv
where, 0ν and 0λ are threshold frequency and threshold wavelength.
Note : q Nearly all metals emit photoelectrons when exposed to u.v. light. But alkali metals like
lithium, sodium, potassium, rubidium and caesium emit photoelectrons even when exposed to visible light.
Examples based on Planck's Quantum theory
λ
hcE =
Visible light
Visible light
U.V. light Metal
Metal other than alkali
metals
Alkali metals
Photo electrons
No photo electrons
Photo electrons
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q Caesium (Cs) with lowest ionisation energy among alkali metals is used in photoelectric cell.
2.8 Bohr’s atomic model.
(1) This model was based on the quantum theory of radiation and the classical law of physics. It gave new
idea of atomic structure in order to explain the stability of the atom and emission of sharp spectral lines.
(2) Postulates of this theory are :
(i) The atom has a central massive core nucleus where all the protons and neutrons are present. The
size of the nucleus is very small.
(ii) The electron in an atom revolve around the nucleus in certain discrete orbits. Such orbits are
known as stable orbits or non – radiating or stationary orbits.
(iii) The force of attraction between the nucleus and the electron is equal to centrifugal force of the
moving electron.
Force of attraction towards nucleus = centrifugal force
(iv) An electron can move only in those permissive orbits in which the angular momentum (mvr) of the
electron is an integral multiple of .2/ πh Thus, π2
hmvr n=
Where, m = mass of the electron, r = radius of the electronic orbit, v = velocity of the electron in its
orbit.
(v) The angular momentum can be .2
,......2
3,
2
2,
2 ππππ
nhhhh This principal is known as quantization of
angular momentum. In the above equation ‘n’ is any integer which has been called as principal
quantum number. It can have the values n=1,2,3, ------- (from the nucleus). Various energy levels
are designed as K(n=1), L(n=2), M(n=3) ------- etc. Since the electron present in these orbits is
associated with some energy, these orbits are called energy levels.
(vi) The emission or absorption of radiation by the atom takes place when an electron jumps from one
stationary orbit to another.
(vii) The radiation is emitted or absorbed as a single quantum (photon) whose energy νh is equal to
the difference in energy E∆ of the electron in the two orbits involved. Thus, Eh ∆=ν
Where ‘h’ =Planck’s constant, =ν frequency of the radiant energy. Hence the spectrum of the atom
will have certain fixed frequency.
(viii) The lowest energy state (n=1) is called the ground state. When an electron absorbs energy, it gets
excited and jumps to an outer orbit. It has to fall back to a lower orbit with the release of energy.
(3) Advantages of Bohr’s theory
(i) Bohr’s theory satisfactorily explains the spectra of species having one electron, viz. hydrogen atom, ++ 2, LiHe etc.
E1
E2 E1 – E2 = hν
Emission
E1
E2 E1 – E2 = hν
Absorption
(ii) Calculation of radius of Bohr’s orbit : According to Bohr, radius of orbit in which electron moves
is
where, n =Orbit number, m =Mass number [ ],101.9 31 kg−× e =Charge on the electron [ ]19106.1 −×
Z =Atomic number of element, k = Coulombic constant [ ]229109 −× cNm
After putting the values of m,e,k,h, we get.
nmZ
nrorÅ
Z
nr nn 529.0529.0
22
×=×=
(a) For a particular system [e.g., H, He+ or Li+2]
(iv) Ionisation energy : The minimum energy required to relieve the electron from the binding of nucleus.
.6.132
2
eff.ionisation eV
n
ZEEE n +=−= ∞
(v) Ionisation potential : e
EV ionisation= ionisation
(vi) Separation energy : Energy required to excite an electron from excited state to infinity.
S.E. = .excitedEE −∞
(vii) Binding energy : Energy released in bringing the electron from infinite to any orbit is called its binding energy (B.E.).
Note : q Principal Quantum Number 'n' = .).(
6.13
EB.
(5) Spectral evidence for quantisation (Explanation for hydrogen spectrum on the basisof bohr atomic
model)
(i) The light absorbed or emitted as a result of an electron changing orbits produces characteristic
absorption or emission spectra which can be recorded on the photographic plates as a series of
lines, the optical spectrum of hydrogen consists of several series of lines called Lyman, Balmar,
Paschen, Brackett, Pfund and Humphrey. These spectral series were named by the name of
scientist who discovered them.
(ii) To evaluate wavelength of various H-lines Ritz introduced the following expression,
where, R is = =3
422
ch
meπ Rydberg's constant
It's theoritical value = 109,737 cm–1 and It's experimental value = 1581.677,109 −cm
This remarkable agreement between the theoretical and experimental value was great achievment
of the Bohr model.
(iii) Although H- atom consists only one electron yet it's spectra consist of many spectral lines as shown in fig.
(iv) Comparative study of important spectral series of Hydrogen
Lyman series
Pfund
series
Brackett
series
Paschen
series
Balmer
series
En
ergy
lev
el
n=7 n=6
n=5
n=4
n=3
n=2
n=1
n=8
Humphrey series
−===
2
2
2
1
111
nnR
c
ν
λν
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S.No. Spectral
series
Lies in
the region
Transitio
n
12 nn >
Rnn
nn
)( 21
22
22
21
max−
=λ R
n 21
min =λ 21
22
22
min
max
nn
n
−=
λ
λ
(1) Lymen
series
Ultraviolet
region
11 =n
∞= ....4,3,22n
2 and 1 21 == nn
R3
4max =λ
∞== 21 and 1 nn
R
1min =λ 3
4
(2) Balmer
series
Visible
region
21 =n
∞= ....5,4,32n
3 and 2 21 == nn
R5
36max =λ
∞== 21 and 2 nn
R
4min =λ
5
9
(3) Paschen
series
Infra red
region
n1 = 3
∞= ....6,5,42n
4 and 3 21 == nn
R7
144max =λ
∞== 21 and 3 nn
R
9min =λ 7
16
(4) Brackett
series
Infra red
region
41 =n
∞= ....7,6,52n
5 and 4 21 == nn
R9
2516max
×=λ
∞== 21 and 4 nn
R
16min =λ 9
25
(5) Pfund series Infra red
region
51 =n
∞= ....8,7,62n
6 and 5 21 == nn
R11
3625max
×=λ
∞== 21 and 5 nn
R
25min =λ 11
36
(6) Humphrey
series
Far
infrared
region
61 =n
∞= ....8,72n
7 and 6 21 == nn
R13
4936max
×=λ
∞== 21 and 6 nn
R
36min =λ 13
49
(v) If an electron from nth excited state comes to various energy states, the maximum spectral lines
obtained will be = .2
)1( −nn n= principal quantum number.
as n=6 than total number of spectral lines = .152
30
2
)16(6==
−
(vi) Thus, at least for the hydrogen atom, the Bohr theory accurately describes the origin of atomic spectral lines.
(6) Failure of Bohr Model
(i) Bohr theory was very successful in predicting and accounting the energies of line spectra of hydrogen i.e. one electron system. It could not explain the line spectra of atoms containing more than one electron.
(ii) This theory could not explain the presence of multiple spectral lines.
(iii) This theory could not explain the splitting of spectral lines in magnetic field (Zeeman effect) and in electric field (Stark effect). The intensity of these spectral lines was also not explained by the Bohr atomic model.
(iv) This theory was unable to explain of dual nature of matter as explained on the basis of De broglies concept.
(v) This theory could not explain uncertainty principle.
(vi) No conclusion was given for the concept of quantisation of energy.
Example: 18 If the radius of 2nd Bohr orbit of hydrogen atom is r2. The radius of third Bohr orbit will be
(a) 29
4r (b) 24r (c) 3
4
9r (d) 29r
Solution : (c) 22
22
4 mZe
hnr
π= ∴
2
2
3
2
3
2=
r
r ∴ 23
4
9rr =
Example: 19 Number of waves made by a Bohr electron in one complete revolution in 3rd orbit is
(a) 2 (b) 3 (c) 4 (d) 1
Solution : (b) Circumference of 3rd orbit = 32 rπ
According to Bohr angular momentum of electron in 3rd orbit is
mvr3 = π2
3h
or 3
2 3r
mv
h π=
by De-Broglie equation, mv
h=λ
3
2 3rπλ =∴ λπ 32 3 =∴ r
i.e. circumference of 3rd orbit is three times the wavelength of electron or number of waves made by Bohr electron in one complete revolution in 3rd orbit is three.
Example: 20 The degeneracy of the level of hydrogen atom that has energy 16
11R− is
(a) 16 (b) 4 (c) 2 (d) 1
Solution : (a) 2n
RE H
n −= ∴ 162
HH R
n
R−=−
i.e. for th4 sub-shell
i.e. 1+3+5+7=16, ∴ degeneracy is 16
Example: 21 The velocity of electron in the ground state hydrogen atom is 2.18 .10 18 −× ms Its velocity in the second orbit would be
(a) 181009.1 −× ms (b) 181038.4 −× ms (c) 15105.5 −× ms (d) 181076.8 −× ms
Solution : (a) We know that velocity of electron in nth Bohr's orbit is given by
smn
Zv /1018.2 6×=
for 1, =ZH
Q smv /1
1018.2 6
1
×=
Q smsmv /1009.1/2
1018.2 66
2 ×=×
=
Example: 22 The ionization energy of the ground state hydrogen atom is .1018.2 18 J−× The energy of an electron in its second orbit would be
+1 0 +1
three p
n=4
1=0
m=0
one s
1
five d
+2 –1 0 +1 +2
2
–3 –2 –1 0 +1 +2 +3 seven f
3
Examples based on Bohr’s atomic model and Hydrogen spectrum
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Solution : (d) For vvLi =+3 for 2zH × =15200 ×9= 1,36800 1−cm
Example: 24 The Bohr orbit radius for the hydrogen atom (n = 1) is approximately 0.530Å. The radius for the first excited state (n = 2) orbit is (in Å)
(a) 0.13 (b) 1.06 (c) 4.77 (d) 2.12
Solution : (d) The Bohr radius for hydrogen atom (n = 1) = 0.530Å
The radius of first excited state (n = 2) will be = ÅZ
n120.2
1
)2(530.0530.0
22
=×=×
Example: 25 How many chlorine atoms can you ionize in the process ,−+ +→ eClCl by the energy liberated from the
following process :
−− →+ CleCl for 23106 × atoms
Given electron affinity of ,61.3 eVCl = and IP of eVCl 422.17=
(a) 1.24 2310× atoms (b) 201082.9 × atoms (c) 151002.2 × atoms (d) None of these
Solution : (a) Energy released in conversion of 23106 × atoms of −Cl ions = 23106 × × electron affinity
= 6× 2310 2410166.261.3 ×=× eV.
Let x Cl atoms are converted to +Cl ion
Energy absorbed ×= x ionization energy
2410166.2422.17 ×=×x ; 2310243.1 ×=x atoms
Example: 26 The binding energy of an electron in the ground state of the He atom is equal to 24eV. The energy required to remove both the electrons from the atom will be
(a) 59eV (b) 81eV (c) 79eV (d) None of these
Solution : (c) Ionization energy of He 6.132
2
×=n
Z 6.13
1
22
2
×= eV4.54=
Energy required to remove both the electrons
= binding energy + ionization energy
= 4.546.24 + = 79eV
Example: 27 The wave number of the shortest wavelength transition in Balmer series of atomic hydrogen will be
(a) 4215 Å (b) 1437Å (c) 3942Å (d) 3647Å
Solution : (d)
−=
22
21
2
shortest
111
nnRZ
λ
∞−××=
22
2 1
2
11109678
cm510647.3 −×=λ Å3647=
Example: 28 If the speed of electron in the Bohr's first orbit of hydrogen atom is x, the speed of the electron in the third Bohr's orbit is
(a) x/9 (b) x/3 (c) 3x (d) 9x
Solution : (b) According to Bohr's model for hydrogen and hydrogen like atoms the velocity of an electron in an atom is
quantised and is given by nh
Zev
22π∝ so
nv
1∝ in this cass 3=n
Example: 29 Of the following transitions in hydrogen atom, the one which gives an absorption line of lowest frequency is
(a) n=1 to n=2 (b) 3=n to 8=n (c) 2=n to 1=n (d) 8=n to 3=n
Solution : (b) Absorption line in the spectra arise when energy is absorbed i.e., electron shifts from lower to higher orbit, out of a & b, b will have the lowest frequency as this falls in the Paschen series.
Example: 30 The frequency of the line in the emission spectrum of hydrogen when the atoms of the gas contain electrons in the third energy level are
(a) Hz1410268.1 × and Hz1610864.2 × (b) Hz1010214.3 × and Hz1210124.1 ×
(c) Hz1210806.1 × and Hz1510204.6 × (d) Hz1410568.4 × and Hz1510924.2 ×
Solution : (d) If an electron is in 3rd orbit, two spectral lines are possible
(a) When it falls from 3rd orbit to 2nd orbit.
In equation
−×=
22
21
15 1110289.3
nnν
289.31 =ν36
510289.3
3
1
2
110 15
22
15 ××=
−× = Hz1414568.4 ×
(b) When it falls from 3rd orbit to 1st orbit :
Hz1515
2
152 10924.2
9
810289.3
3
1
1
110289.3 ×=××=
−××=ν
Example: 31 If the first ionisation energy of hydrogen is J1810179.2 −× per atom, the second ionisation energy of helium per atom is
Solution : (a) For Bohrs systems : energy of the electron 2
2
n
Z∝
Ionisation energy is the difference of energies of an electron ),( ∞E when taken to infinite distance and rE
when present in any Bohr orbit and αE is taken as zero so ionisation energy becomes equal to the energy
of electron in any Bohr orbit.
2
2
H
HH
n
ZE ∝ ;
2
2
He
HeHe
n
ZE ∝ or
22
1
×=
He
H
E
E [as ]1,1,2,1 ==== HeHHeH nnZZ
or 1818 10716.8410179.24 −− ×=××=×= HHe EE Joule per atom.
Example: 32 The ionization energy of hydrogen atom is 13.6eV. What will be the ionization energy of +He (a) 13.6eV (b) 54.4eV (c) 122.4eV (d) Zero
Solution : (b) I.E. of 26.13 ZeVHe ×=+ eVeV 4.5446.13 =×
Example: 33 The ionization energy of +He is 18106.19 −× J atom–1. Calculate the energy of the first stationary state of 2+Li
(a) -118 atom106.19 J−× (b) J181041.4 −× atom–1
(c) -119 atom106.19 J−× (d) 117 atom1041.4 −−× J
Solution : (d) I.E. of ZEHe (2 2×=+ for )2=He
I.E. of 32 3×=+ ELi (Z for Li=3)
9
4
).(.
).(.2
=∴+
+
LiEI
HeEI or I.E. ).(.
4
9)( 2 ++ ×= HeEILi 18106.19
4
9 −××= = 171041.4 −× J atom–1
2.9 Bohr – Sommerfeld’s model.
(1) In 1915, Sommerfield introduced a new atomic model to explain the fine spectrum of hydrogen atom.
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(2) He gave concept that electron revolve round the nucleus in elliptical orbit. Circular orbits are formed in special conditions only when major axis and minor axis of orbit are equal.
(3) For circular orbit, the angular momentum = π2
nh where n= principal quantum number only one
component i.e. only angle changes.
(4) For elliptical orbit, angular momentum = vector sum of 2 components. In elliptical orbit two components are,
(i) Radial component (along the radius) = π2
hnr
Where, n r = radial quantum number
(ii) Azimuthal component = n φπ2
h
Where, n φ = azimuthal quantum number
So angular momentum of elliptical orbit =ππ
φ22
hn
hnr +
Angular momentum = π
φ2
)(h
nnr +
(5) Shape of elliptical orbit depends on,
φn
n=
axisminorofLength
axismajorofLength
φ
φ
n
nnr +=
(6) n φ can take all integral values from l to ‘n’ values of n r depend on the value of n φ . For n = 3,
n φ can have values 1,2,3 and n r can have (n –1) to zero i.e. 2,1 and zero respectively.
Thus for n = 3, we have 3 paths
n n φ n r Nature of path
3 1 3 elliptical
2 1 elliptical
3 0 circular
The possible orbits for n = 3 are shown in figure.
Thus Sommerfield showed that Bohr’s each major level was composed of several sub-levels. therefore it provides the basis for existance of subshells in Bohr's shells (orbits).
(7) Limitation of Bohr sommerfield model :
(i) This model could not account for, why electrons does not absorb or emit energy when they are moving in stationary orbits.
(ii) When electron jumps from inner orbit to outer orbit or vice –versa, then electron run entire distance but absorption or emission of energy is discontinuous.
(iii) It could not explain the attainment of expression of π2
nh for angular momentum. This model could
not explain Zeeman effect and Stark effect.
• Nuclear
K= 1
K= 2
K= 3
φ = change
r = change
r
φ1 r φ2 φ2 r
φ1 r
φ = change
r = constant
φ1 r
r
2.10 Dual nature of electron.
(1) In 1924, the french physicist, Louis de Broglie suggested that if light has both particle and wave like nature, the similar duality must be true for matter. Thus an electron, behaves both as a material particle and as a wave. (2) This presented a new wave mechanical theory of matter. According to this theory, small particles like electrons when in motion possess wave properties. (3) According to de-broglie, the wavelength associated with a particle of mass ,m moving with velocity v
is given by the relation
,mv
h=λ where h = Planck’s constant.
(4) This can be derived as follows according to Planck’s equation, λ
νch
hE.
==
=
λν
cQ
energy of photon (on the basis of Einstein’s mass energy relationship), 2mcE =
equating both mc
hormc
hc== λ
λ2 which is same as de-Broglie relation. ( )pmc =Q
(5) This was experimentally verified by Davisson and Germer by observing diffraction effects with an electron beam. Let the electron is accelerated with a potential of V than the Kinetic energy is
eVmv =2
2
1; eVmvm 222 =
PeVmmv == 2 ; eVm
h
2=λ
(6) If Bohr’s theory is associated with de-Broglie’s equation then wave length of an electron can be determined in bohr’s orbit and relate it with circumference and multiply with a whole number
n
rornr
πλλπ
22 ==
From de-Broglie equation, mv
h=λ . Thus
n
r
mv
h π2= or
π2
nhmvr =
Note : q For a proton, electron and an α -particle moving with the same velocity have de-broglie
wavelength in the following order : Electron > Proton > α - particle.
(7) The de-Broglie equation is applicable to all material objects but it has significance only in case of microscopic particles. Since, we come across macroscopic objects in our everyday life, de-broglie relationship has no significance in everyday life.
Example: 34 An electron is moving with a kinetic energy of 4.55 × 10 25− J. What will be de-Broglie
(1) One of the important consequences of the dual nature of an electron is the uncertainty principle,
developed by Warner Heisenberg.
(2) According to uncertainty principle “It is impossible to specify at any given moment both the position
and momentum (velocity) of an electron”.
Mathematically it is represented as , π4
.h
px ≥∆∆
Where =∆x uncertainty is position of the particle, =∆p uncertainty in the momentum of the particle
Now since vmp ∆=∆
So equation becomes, π4
.h
vmx ≥∆∆ or m
hvx
π4≥∆×∆
The sign ≥ means that the product of x∆ and p∆ (or of x∆ and v∆ ) can be greater than, or equal to but
never smaller than .4π
hIf x∆ is made small, p∆ increases and vice versa.
(3) In terms of uncertainty in energy, E∆ and uncertainty in time ,t∆ this principle is written as,
π4
.h
tE ≥∆∆
Note :q Heisenberg’s uncertainty principle cannot we apply to a stationary electron because its
velocity is 0 and position can be measured accurately.
Example: 36 What is the maximum precision with which the momentum of an electron can be known if the
uncertainty in the position of electron is ?001.0 ű Will there be any problem in describing the
momentum if it has a value of ,2 0a
h
πwhere 0a is Bohr’s radius of first orbit, i.e., 0.529Å?
Solution : π4
.h
px =∆∆
Q mÅx 1310001.0 −==∆
∴ 22
13
34
1027.51014.34
10625.6 −
−
−
×=××
×=∆p
Example: 37 Calculate the uncertainty in velocity of an electron if the uncertainty in its position is of the order
of a 1Å.
Solution : According to Heisenberg’s uncertainty principle
Examples based on uncertainty principle
π4.
hpx ≥∆∆
m
hxv
π4. ≈∆∆
xm
hv
∆≈∆
.4π
1031
34
1010108.97
224
10625.6
−−
−
××××
×= 15 sec108.5 −×= m
Example: 38 A dust particle having mass equal to ,10 11 g− diameter of cm410 − and velocity .sec10 14 −− cm The
error in measurement of velocity is 0.1%. Calculate uncertainty in its positions. Comment on the result .
Solution : 174
sec101100
101.0 −−−
×=×
=∆ cmv
Q m
hxv
π4. =∆∆
∴ cmx 10
711
27
1027.51011014.34
10625.6 −
−−
−
×=××××
×=∆
The uncertainty in position as compared to particle size.
cmdiameter
x 6
4
10
1027.510
1027.5 −
−
−
×=×
=∆
=
The factor being small and almost being negligible for microscope particles.
2.12 Schrödinger wave equation.
(1) Schrodinger wave equation is given by Erwin Schrödinger in 1926 and based on dual nature of electron.
(2) In it electron is described as a three dimensional wave in the electric field of a positively charged nucleus.
(3) The probability of finding an electron at any point around the nucleus can be determined by the help of Schrodinger wave equation which is,
0)(8
2
2
2
2
2
2
2
2
=Ψ−+∂
Ψ∂+
∂
Ψ∂+
∂
Ψ∂VE
h
m
zyx
π
Where yx, and z are the 3 space co-ordinates, m = mass of electron, h = Planck’s constant,
E = Total energy, V = potential energy of electron, Ψ = amplitude of wave also called as wave function.
∂ = stands for an infinitesimal change.
(4) The Schrodinger wave equation can also be written as :
0)(8
2
22 =Ψ−+Ψ∇ VE
h
mπ
Where ∇ = laplacian operator.
(5) Physical Significance of Ψ and 2Ψ
(i) The wave function Ψ represents the amplitude of the electron wave. The amplitude Ψ is thus a
function of space co-ordinates and time i.e. )......,,( timeszyxΨ=Ψ
(ii) For a single particle, the square of the wave function )( 2Ψ at any point is proportional to the
probability of finding the particle at that point.
(iii) If 2Ψ is maximum than probability of finding −e is maximum around nucleus. And the place
where probability of finding −e is maximum is called electron density, electron cloud or an atomic orbital. It is different from the Bohr’s orbit.
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(iv) The solution of this equation provides a set of number called quantum numbers which describe specific or definite energy state of the electron in atom and information about the shapes and orientations of the most probable distribution of electrons around the nucleus.
2.13 Quantum numbers and Shapes of orbitals.
QQuuaannttuumm nnuummbbeerrss
(1) Each orbital in an atom is specified by a set of three quantum numbers (n, l, m) and each electron is
designated by a set of four quantum numbers (n, l, m and s).
(2) Principle quantum number (n)
(i) It was proposed by Bohr’s and denoted by ‘n’.
(ii) It determines the average distance between electron and nucleus, means it is denoted the size of
atom.
ÅZ
nr 529.0
2
×=
(iii) It determine the energy of the electron in an orbit where electron is present.
moleperKcaln
ZE 3.313
2
2
×−=
(iv) The maximum number of an electron in an orbit represented by this quantum number as .2 2n No
energy shell in atoms of known elements possess more than 32 electrons.
(v) It gives the information of orbit K, L, M, N------------.
(vi) The value of energy increases with the increasing value of n.
(vii) It represents the major energy shell or orbit to which the electron belongs.
(viii) Angular momentum can also be calculated using principle quantum number
π2
nhmvr =
(3) Azimuthal quantum number (l)
(i) Azimuthal quantum number is also known as angular quantum number. Proposed by Sommerfield
and denoted by ‘l’.
(ii) It determines the number of sub shells or sublevels to which the electron belongs.
(iii) It tells about the shape of subshells.
(iv) It also expresses the energies of subshells fdps <<< (increasing energy).
(v) The value of )1( −= nl always where ‘n’ is the number of principle shell.
(vi) Value of l = 0 1 2 3………..(n-1)
Name of subshell = s p d f
Shape of subshell = Spherical Dumbbell Double
dumbbell
Complex
(vii) It represent the orbital angular momentum. Which is equal to )1(2
+llh
π
(viii) The maximum number of electrons in subshell )12(2 += l
electrons2subshell →−s electrons10subshell →−d
electrons6subshell →−p electrons.14subshell →−f
(ix) For a given value of ‘n’ the total value of ‘l’ is always equal to the value of ‘n’.
(x) The energy of any electron is depend on the value of n & l because total energy = (n + l). The
electron enters in that sub orbit whose (n + l) value or the value of energy is less.
(4) Magnetic quantum number (m)
(i) It was proposed by Zeeman and denoted by ‘m’.
(ii) It gives the number of permitted orientation of subshells.
(iii) The value of m varies from –l to +l through zero.
(iv) It tells about the splitting of spectral lines in the magnetic field i.e. this quantum number proved
the Zeeman effect.
(v) For a given value of ‘n’ the total value of ’m’ is equal to .2n
(vi) For a given value of ‘l’ the total value of ‘m’ is equal to ).12( +l
(vii) Degenerate orbitals : Orbitals having the same energy are known as degenerate orbitals. e.g. for p
subshell zyx ppp
(viii) The number of degenerate orbitals of s subshell =0.
(5) Spin quantum numbers (s)
(i) It was proposed by Goldshmidt & Ulen Back and denoted by the symbol of ‘s’.
(ii) The value of 1/2,-and1/2 is'' +s which is signifies the spin or rotation or direction of electron on
it’s axis during movement.
(iii) The spin may be clockwise or anticlockwise.
(iv) It represents the value of spin angular momentum is equal to .)1(2
+ssh
π
(v) Maximum spin of an atom ×= 2/1 number of unpaired electron.
(vi) This quantum number is not the result of solution of schrodinger equation as solved for H-atom.
Distribution of electrons among the quantum levels
n l m s Designation of
orbitals
Electrons
present
Total no. of
electrons
1 (K shell) 0 0 +1/2, –1/2 1s 2 2
N
+1/2
S
–1/2
Magnetic field
N S
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2 (L shell) 0
1
0
+1
0
–1 2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
−+
−+
−+
−+
2s
2p
6
2
8
3 (M shell)
4(N shell)
0
1
2
0
1
2
3
0
+1
0
–1
+2
+1
0
–1
–2
0
+1
0
–1
+2
+1
0
–1
–2
+3
+2
+1
+0
–1
–2
–3
2/1,2/1
2/1,2/1
2/1,2/1
−+
−+
−+
2/1,2/1 −+
−+
−+
−+
−+
−+
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
−+
−+
−+
−+
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
−+
−+
−+
−+
−+
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
−+
−+
−+
−+
−+
−+
−+
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
2/1,2/1
3s
3p
3d
4s
4p
4d
4f
10
6
2
14
10
6
2
18
32
Z
Y
X
Nucleus
SShhaappee ooff oorrbbiittaallss
(1) Shape of ‘s’ orbital
(i) For ‘s’ orbital l=0 & m=0 so ‘s’ orbital have only one unidirectional
orientation i.e. the probability of finding the electrons is same in all
directions.
(ii) The size and energy of ‘s’ orbital with increasing ‘n’ will be
.4321 ssss <<<
(iii)It does not possess any directional property. s orbital has spherical
shape.
(2) Shape of ‘p’ orbitals
(i) For ‘p’ orbital l=1, & m=+1,0,–1 means there are three ‘p’ orbitals, which is symbolised as .,, zyx ppp
(ii) Shape of ‘p’ orbital is dumb bell in which the two lobes on opposite side separated by the nodal
plane.
(iii) p-orbital has directional properties.
(3) Shape of ‘d’ orbital
(i) For the ‘d’ orbital l =2 then the values of ‘m’ are –2,–1,0,+1,+2. It shows that the ‘d’ orbitals has five
orbitals as .222 ,,,,
zyxzxyzxy ddddd−
(ii) Each ‘d’ orbital identical in shape, size and energy.
(iii) The shape of d orbital is double dumb bell .
(iv) It has directional properties.
(4) Shape of ‘f’ orbital
(i) For the ‘f’ orbital l=3 then the values of ‘m’ are –3, –2, –1,0,+1,+2,+3. It shows that the ‘f’ orbitals
have seven orientation as .and,,,, 233222222 ),()()( xzyzzxyzyxzyxyyxxfffffff
−−−
(ii) The ‘f’ orbital is complicated in shape.
Y
X
dZ2
Z
Y
X
dX2–Y
2
Z Y
X
dYZ
Z Y
X
dXY
Z Y
X
dZX
Nodal Plane
Z Y
X
Pz orbital Nodal Plane
Nodal Plane
Z Y
X
Py orbital
Nodal Plane
Nodal Plane
Z Y
Px orbital
X
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2.14 Electronic configuration principles.
The distribution of electrons in different orbitals of atom is known as electronic configuration of the
atoms.
Filling up of orbitals in the ground state of atom is governed by the following rules:
(1) Aufbau principle
(i) Auf bau is a German word, meaning ‘building up’.
(ii) According to this principle, “In the ground state, the atomic orbitals are filled in order of
increasing energies i.e. in the ground state the electrons first occupy the lowest energy orbitals
available”.
(iii)In fact the energy of an orbital is determined by the quantum number n and l with the help of (n+l)
rule or Bohr Bury rule.
(iv) According to this rule
(a) Lower the value of n + l, lower is the energy of the orbital and such an orbital will be filled up
first.
(b) When two orbitals have same value of (n+l) the orbital having lower value of “n” has lower
energy and such an orbital will be filled up first .
Thus, order of filling up of orbitals is as follows:
dfspdspspspss 5665454433221 <<<<<<<<<<<<
(2) Pauli’s exclusion principle
(i) According to this principle, “No two electrons in an atom can have same set of all the four
quantum numbers n, l, m and s .
(ii) In an atom any two electrons may have three quantum numbers identical but fourth quantum
number must be different.
(iii)Since this principle excludes certain possible combinations of quantum numbers for any two
electrons in an atom, it was given the name exclusion principle. Its results are as follows :
(a) The maximum capacity of a main energy shell is equal to 22n electron.
(b) The maximum capacity of a subshell is equal to 2(2l+1) electron.
(c) Number of sub-shells in a main energy shell is equal to the value of n.
(d) Number of orbitals in a main energy shell is equal to .2n
(e) One orbital cannot have more than two electrons.
(iv) According to this principle an orbital can accomodate at the most two electrons with spins
opposite to each other. It means that an orbital can have 0, 1, or 2 electron.
(v) If an orbital has two electrons they must be of opposite spin.
Correct Incorrect
(3) Hund’s Rule of maximum multiplicity
(i) This rule provides the basis for filling up of degenerate orbitals of the same sub-shell.
(ii) According to this rule “Electron filling will not take place in orbitals of same energy until all the
available orbitals of a given subshell contain one electron each with parallel spin”.
(iii)This implies that electron pairing begins with fourth, sixth and eighth electron in p, d and f
orbitals of the same subshell respectively.
(iv) The reason behind this rule is related to repulsion between identical charged electron present in
the same orbital.
(v) They can minimise the repulsive force between them serves by occupying different orbitals.
(vi) Moreover, according to this principle, the electron entering the different orbitals of subshell have
parallel spins. This keep them farther apart and lowers the energy through electron exchange or
resonance.
(vii) The term maximum multiplicity means that the total spin of unpaired −e is maximum in case of
correct filling of orbitals as per this rule.
Energy level diagram
The representation of relative energy levels of various atomic orbital is made in the terms of energy level
diagrams.
One electron system : In this system 21s level and all orbital of same principal quantum number have
same energy, which is independent of (l). In this system l only determines the shape of the orbital.
Multiple electron system : The energy levels of such system not only depend upon the nuclear charge
but also upon the another electron present in them.
Diagram of multi-electron atoms reveals the following points :
Energy level diagram of one electron system
En
erg
y
1s
5
4
3
2
2s 2p
3s 3p 3d
4s 4p 4d 4f 6 p 5 d 4 f 6 s 5 p 4 d 5 s
4 p 3 d 4 s 3 p
3 s 2 p
2 s 1 s
En
erg
y
Energy level diagram of multiple electron system
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(i) As the distance of the shell increases from the nucleus, the energy level increases. For example energy
level of 2 > 1.
(ii) The different sub shells have different energy levels which possess definite energy. For a definite shell,
the subshell having higher value of l possesses higher energy level. For example in 4th shell.
Energy level order 4f > 4d > 4p > 4s
l= 3 l = 2 l = 1 l= 0
(iii)The relative energy of sub shells of different energy shell can be explained in the terms of the (n+l)
rule.
(a) The sub-shell with lower values of (n + l) possess lower energy.
For 3d n = 3 l= 2 ∴ n + l = 5
For 4s n = 4 l = 0 n + l = 4
(b) If the value of (n + l) for two orbitals is same, one with lower values of ‘n’ possess lower energy level.
Extra stability of half filled and completely filled orbitals
Half-filled and completely filled sub-shell have extra stability due to the following reasons :
(i) Symmetry of orbitals
(a) It is a well kown fact that symmetry leads to stability.
(b) Thus, if the shift of an electron from one orbital to another orbital differing slightly in energy
results in the symmetrical electronic configuration. It becomes more stable.
(c) For example 753 ,, fdp configurations are more stable than their near ones.
(ii) Exchange energy
(a) The electron in various subshells can exchange their positions, since electron in the same subshell have equal energies.
(b) The energy is released during the exchange process with in the same subshell.
(c) In case of half filled and completely filled orbitals, the exchange energy is maximum and is greater
than the loss of orbital energy due to the transfer of electron from a higher to a lower sublevel e.g. from 4s to 3d orbitals in case of Cu and Cr .
(d) The greater the number of possible exchanges between the electrons of parallel spins present in
the degenerate orbitals, the higher would be the amount of energy released and more will be the stability.
(e) Let us count the number of exchange that are possible in 4d and 5d configuraton among electrons
with parallel spins.
To number of possible exchanges = 3 + 2 + 1 =6
To number of possible exchanges = 4 + 3 + 2 +1 = 10
3 exchanges by 1st e– 2 exchanges by 2nd e– Only 1 exchange by 3rd e–
d4 (1) (2) (3)
2 exchange by 3rd e– 3 exchanges by 2nd e– 4 exchanges by 1st e–
(2) (3) d5 (1)
1 exchange by 4th e–
(4)
2.15 Electronic configurations of Elements.
(1) On the basis of the elecronic configuration priciples the electronic configuration of various elements are given in the following table :
Electronic Configuration (E.C.) of Elements Z=1 to 36
Element Atomic
Number 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f
H 1 1
He 2 2
Li 3 2 1
Be 4 2 2
B 5 2 2 1
C 6 2 2 2
N 7 2 2 3
O 8 2 2 4
F 9 2 2 5
Ne 10 2 2 6
Na 11 2 2 6 1
Mg 12 2
Al 13 2 1
Si 14 10
electrons
2 2
P 15 2 3
S 16 2 4
Cl 17 2 5
Ar 18 2 2 6 2 6
K 19 2 2 6 2 6 1
Ca 20 2
Sc 21 1 2
Ti 22 2 2
V 23 3 2
Cr 24 5 1
Mn 25 5 2
Fe 26 6 2
Co 27 18
electrons
7 2
Ni 28 8 2
Cu 29 10 1
Zn 30 10 2
Ga 31 10 2 1
Ge 32 10 2 2
As 33 10 2 3
Se 34 10 2 4
Br 35 10 2 5
Kr 36 2 2 6 2 6 10 2 6
(2) The above method of writing the electronic configurations is quite cumbersome. Hence, usually the electronic configuration of the atom of any element is simply represented by the notation.
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(3) (i) Elements with atomic number 24(Cr), 42(Mo) and 74(W) have 51 )1( dnns − configuration and not 42 )1( dnns − due to extra stability of these atoms.
(ii) Elements with atomic number 29(Cu), 47(Ag) and 79(Au) have 101 )1( dnns − configuration instead
of 92 )1( dnns − due to extra stability of these atoms.
(4) In the formation of ion, electrons of the outer most orbit are lost. Hence, whenever you are required to write electronic configuration of the ion, first write electronic configuration of its atom and take electron from
outermost orbit. If we write electronic configuration of Fe ),24,26(2 −+ = eZ it will not be similar to Cr (with −e24 ) but quite different.
[ ][ ]
+ 62
62
34
34
dsArFe
dsArFeo
outer most orbit is 4th shell hence, electrons from 4s have been removed to make +2Fe .
(5) Ion/atom will be paramagnetic if there are unpaired electrons. Magnetic moment (spin only) is
)2( += nnµ BM (Bohr Magneton). )/1027.91( 24 TJBM −×= where n is the number of unpaired electrons.
(6) Ion with unpaired electron in d or f orbital will be coloured. Thus, +Cu with electronic configuration
[ ] 103dAr is colourless and +2Cu with electronic configuration [ ] 93dAr (one unpaired electron in 3d) is
coloured (blue).
(7) Position of the element in periodic table on the basis of electronic configuration can be determined as, (i) If last electron enters into s-subshell, p-subshell, penultimate d-subshell and anti penultimate f-
subshell then the element belongs to s, p, d and f – block respectively. (ii) Principle quantum number (n) of outermost shell gives the number of period of the element.
(iii)If the last shell contains 1 or 2 electrons (i.e. for s-block elements having the configuration 21−ns ), the group number is 1 in the first case and 2 in the second case.
(iv) If the last shell contains 3 or more than 3 electrons (i.e. for p-block elements having the
configuration 612 −npns ), the group number is the total number of electrons in the last shell plus
10. (v) If the electrons are present in the (n –1)d orbital in addition to those in the ns orbital (i.e. for d-
block elements having the configuration (n –1) 21101 −− nsd ), the group number is equal to the total
number of electrons present in the (n –1)d orbital and ns orbital.
nl
NUMBER OF PRINCIPAL SHELL
NUMBER OF ELECTRONS PRESENT
SYMBOL OF SUBSHELL
x
e.g. 1s2 means 2 electrons are present in the s- subshell of the 1st main shell.
3d5
4s1
Cr (24) [Ar] 3d10
4s1
Cu (29) [Ar]
1. The fundamental particles present in the nucleus of an atom are [CPMT 1983, 84]
(a) Alpha particles and electrons (b) Neutrons and protons
(c) Neutrons and electrons (d) Electrons, neutrons and protons
2. Cathode rays were discovered by
(a) William Crookes (b) J. Stoney (c) Rutherford (d) None of these
58. The electronic configuration of a dipositive metal +2M is 2, 8, 14 and its atomic weight is 56 a.m.u. The number of neutrons in its nuclei would be [MNR 1984, 89; Kerala PMT 1999]
(a) 30 (b) 32 (c) 34 (d) 42
59. The total number of unpaired electrons in d-orbitals of atoms of element of atomic number 29 is [CPMT 1983]
(a) 10 (b) 1 (c) 0 (d) 5
60. Chlorine atom differs from chloride ion in the number of [NCERT 1972; MP PMT 1995]
(a) Proton (b) Neutron (c) Electrons (d) Protons and electrons
61. The number of electrons in one molecule of 2CO are [IIT 1979; MP PMT 1994; Rajasthan PMT 1999]
(a) 22 (b) 44 (c) 66 (d) 88
62. The nitrogen atom has 7 protons and 7 electrons, the nitride ion )( 3−N will have [NCERT 1977]
(a) 7 protons and 10 electrons (b) 4 protons and 7 electrons
(c) 4 protons and 10 electrons (d) 10 protons and 7 electrons
63. The total number of neutrons in dipositive zinc ions with mass number 70 is [IIT 1979; Bihar MEE 1997]
(a) 34 (b) 40 (c) 36 (d) 38
64. If W is atomic weight and N is the atomic number of an element, then [CPMT 1971, 80, 89]
(a) Number of NWe −=−1 (b) Number of NWn −=10
(c) Number of NWH −=11 (d) Number of Nn =1
0
65. The number of electrons in the atom which has 20 protons in the nucleus is [CPMT 1981, 93; CBSE 1989]
(a) 20 (b) 10 (c) 30 (d) 40
66. Six protons are found in the nucleus of [CPMT 1977, 80, 81; NCERT 1975, 78]
(a) Boron (b) Lithium (c) Carbon (d) Helium
67. A sodium cation has different number of electrons from
(a) −2O (b) −F (c) +Li (d) +++Al
68. An atom which has lost one electron would be [CPMT 1986]
85. The number of electrons and neutrons of an element is 18 and 20 respectively. Its mass number is [CPMT 1997; Pb. PMT 1999; MP PMT 1999]
(a) 17 (b) 37 (c) 2 (d) 38
86. The number of electrons in 14019 ][ −K is [CPMT 1997; AFMC 1999]
(a) 19 (b) 20 (c) 18 (d) 40
87. In the nucleus of 4020 Ca there are [CPMT 1990; EAMCET 1991]
(a) 40 protons and 20 electrons (b) 20 protons and 40 electrons
(c) 20 protons and 20 neutrons (d) 20 protons and 40 neutrons
88. The atomic weight of an element is 39. The number of neutrons in its nucleus is one more than the number of protons. The number of protons, neutrons and electrons respectively in its atom would be [MP PMT 1997]
111. 2CO is isostructural with [IIT 1986; MP PMT 1986, 94, 95]
(a) 2SnCl (b) 2SO (c) 2HgCl (d) All the above
112. In an X-ray experiment, different metals are used as the target. In each case, the frequency (ν) of the radiation produced is
measured. If Z= atomic number, which of the following plots will be a straight line
Advance LLeevveell
(a) ν against Z (b) ν
1 against Z (c) ν against Z (d) ν against Z
113. In Moseley's equation )([ bZa −=ν ], which was derived from the observations made during the bombardment of metal targets
with X-rays,
(a) a is independent but b depends on the metal (b) Both a and b depend on the metal (c) Both a and b are independent of the metal and are constant (d) b is independent but a depends on the metal 114. If molecular mass and atomic mass of sulphur are 256 and 32 respectively, its atomicity is [Rajasthan PET 2000]
(a) 2 (b) 8 (c) 4 (d) 16 115. Assertion (A) : The atoms of different elements having same mass number but different atomic number are known as isobars
Reason (R) : The sum of protons and neutrons, in the isobars is always different [AIIMS 2000] (a) Both A and R are true and R is a correct explanation of A (b) Both A and R are true but R is not a correct explanation of A (c) A is true but the R is false (d) A is false but R is true
116. The mass number of an anion, 3−X , is 14. If there are ten electrons in the anion, the number of neutrons in the nucleus of atom,
2X of the element will be [MP PMT 1999]
(a) 10 (b) 14 (c) 7 (d) 5 117. Atoms consists of protons, neutrons and electrons. If the mass of neutrons and electrons were made half and two times
respectively to their actual masses, then the atomic mass of 126 C [NCERT 1982]
(a) Will remain approximately the same (b) Will become approximately two times (c) Will remain approximately half (d) Will be reduced by 25%
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118. A neutral atom (Atomic no. > 1) consists of [CPMT 1982] (a) Only protons (b) Neutrons + protons (c) Neutrons + electrons (d) Neutrons +proton + electron 119. Compared with an atom of atomic weight 12 and atomic number 6, the atom of atomic weight 13 and atomic number 6
[NCERT 1971] (a) Contains more neutrons (b) Contains more electrons (c) Contains more protons (d) Is a different element
120. Assertion (A) : Nuclide 1330 Al is less stable than 20
40 Ca
Reason (R): Nuclides having odd number of protons and neutrons are generally unstable [IIT 1998] (a) Both A and R are correct and R is the correct explanation of A (b) Both A and R are correct but R is not the correct explanation of A (c) A is correct but R is incorrect (d) A is incorrect but R is correct
121. Which of the following are iso-electronic species 3423 IV,III,II,I NHNHNHCH −−−− +−+ [CPMT 1999]
(a) I, II, III (b) II, III, IV (c) I, II, IV (d) I and II
122. The charge on the atom containing 17 protons, 18 neutrons and 18 electrons is [AIIMS 1996]
(a) + 1 (b) – 2 (c) –1 (d) Zero
123. Rutherford’s α-particle scattering experiment proved that atom has [MP PMT 2001]
181. The energy of a photon is calculated by [Pb. PMT 2000]
(a) νhE = (b) νEh = (c) V
Eh = (d)
V
hE =
182. Which of the following is true for Thomson's model of the atom
(a) The radius of an electron can be calculated using Thomson's model.
(b) In an undisturbed atom, the electrons will be at their equilibrium positions, where the attraction between the cloud of positive charge and the electrons balances their mutual repulsion
(c) When the electrons are disturbed by collision, they will vibrate around their equilibrium positions and emit electromagnetic radiation whose frequency is of the order of magnitude of the frequency of electromagnetic radiation of a vibrating electron.
(d) It can explain the existence of protons.
183. When a gold sheet is bombarded by a beam of α–particles, only a few of them get deflected whereas most go straight,
undeflected. This is because
(a) The force of attraction exerted on the α–particles by the oppositely charged electrons is not sufficient.
(b) A nucleus has a much smaller volume than that of an atom.
(c) The force of repulsion acting on the fast moving α–particles is very small.
(d) The neutrons in the nucleus do not have any effect on the α–particles.
184. From the α–particle scattering experiment, Rutherford concluded that
(a) α–particles can come within a distance of the order of 10–14m of the nucleus
(b) The radius of the nucleus is less than 10–14m
(c) Scattering follows Coulomb's law
AAddvvaannccee Level
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(d) The positively charged parts of the atom move with extremely high velocities.
185. Rutherford's scattering formula fails for very small scattering angles because
(a) The full nuclear charge of the target atom is partially screened by its electron
(b) The impact parameter between the α–particle source and the nucleus of the target is very large compared to the size of the
nucleus
(c) The kinetic energy of the α–particles is large
187. The nucleus of an atom can be assumed to be spherical. The radius of the nucleus of mass number A is given by
cmA 3/1131025.1 ×× − . Radius of atom is one Å. If the mass number is 64, then the fraction of the atomic volume that is occupied by the nucleus is [NCERT 1983]
189. The radius of first Bohr’s orbit for hydrogen is 0.53 Å. The radius of third Bohr’s orbit would be [MP PET 1994]
(a) 0.79 Å (b) 1.59 Å (c) 3.18 Å (d) 4.77 Å
190. The energy of an electron in the first Bohr orbit of H atom is ––13.6 eV. The possible energy value (s) of the excited state (s) for electrons in Bohr orbits to hydrogen is (are) [IIT 1998]
(a) – 3.4 eV (b) – 4.2 eV (c) – 6.8 eV (d) + 6.8 eV
191. Energy of electron of hydrogen atom in second Bohr orbit is [MP PMT 2000]
192. 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
[CBSE 1998; BHU 1999]
(a) 0.13 Å (b) 1.06 Å (c) 4.77 Å (d) 2.12 Å
193. The energy of an electron in nth orbit of hydrogen atom is [MP PET 1999]
(a) eVn4
6.13 (b) eV
n3
6.13 (c) eV
n2
6.13 (d) eV
n
6.13
194. As electron moves away from the nucleus, its potential energy [UPSEAT 2003]
(a) Increases (b) Decreases (c) Remains constant (d) None of these
195. In which one of the following pairs of experimental observations and phenomenon does the experimental observation correctly account for phenomenon [AIIMS 1983]
Experimental observation Phenomenon
(a) X- ray spectra (a) Charge on the nucleus
(b)
α-particle scattering (b)
Quantized electron orbit
(c) Emission spectra (c) The quantization of energy
(d)
The photoelectric effect (d)
The nuclear atom
196. When an electron jumps from ‘L’ level to ‘M’ level, there occurs [EAMCET 1979]
(a) Emission of energy (b) Emission of X-rays (c) Absorption of energy (d) Emission of γ-rays
197. In Balmer series of hydrogen atom spectrum which electronic transition causes third line [MP PMT 2000]
(a) Fifth Bohr orbit to second one (b) Fifth Bohr orbit to first one
(c) Fourth Bohr orbit to second one (d) Fourth Bohr orbit to first one
198. In which of the following transitions will the wavelength be minimum
(a) n = 6 to n = 4 (b) n = 4 to n = 2 (c) n = 3 to n = 1 (d) n = 2 to n = 1
199. The frequency of one of the lines in Paschen series of hydrogen atom is Hz1410340.2 × . The quantum number 2n which
produces this transition is [Delhi PMT 2001]
(a) 6 (b) 5 (c) 4 (d) 3
200. Positronium consists of an electron and a positron (a particle which has the same mass as an electron, but opposite charge) orbiting round their common centre of mass. Calculate the value of the Rydberg constant for this system.
(a) 4/∞R (b) 2/∞R (c) ∞R2 (d) ∞R
201. What are the average distance and the most probable distance of an electron from the nucleus in the 1s orbital of a hydrogen
atom [ =0a the radius of the first Bohr orbit]
(a) 05.1 a and 0a (b) 0a and 05a (c) 05.1 a and 05.0 a (d) 0a and 05.0 a
202. Choose the correct relations on the basis of Bohr theory
(a) Velocity of electron n
1∝ (b) Frequency of revolution
3
1
n∝
(c) Radius of orbit Zn2∝ (d) Force on electron 4
1
n∝
203. For a hydrogen atom, what is the orbital degeneracy of the level that has energy 9
∞−=
hcR, where ∞R is the Rydberg constant for
the hydrogen atom
(a) 1 (b) 9 (c) 36 (d) 3
204. In a hydrogen atom, if energy of an electron in ground state is 13.6 eV, then that in the 2nd excited state is [AIEEE 2002]
(a) – 1.51 eV (b) – 3.4 eV (c) – 6.04 eV (d) – 13.6 eV
205. The ionization energy of hydrogen atom is – 13.6 eV. The energy required to excite the electron in a hydrogen atom from the
ground state to the first excited state is (Avogadro’s constant )10022.6 23×= [BHU 1999]
206. The value of the energy for the first excited state of hydrogen atom will be [MP PET 2002]
(a) – 13.6 eV (b) – 3.40 eV (c) – 1.51 eV (d) – 0.85 eV
207. An atom has 2 electrons in K shell, 8 electrons in L shell and 6 electrons in M shell. The number of s-electrons present in that element is [CPMT 1989]
(a) 6 (b) 5 (c) 7 (d) 10
208. In Bohr series of lines of hydrogen spectrum, the third line from the red end corresponds to which one of the following inter-orbit jumps of the electron for Bohr orbits in an atom of hydrogen [AIEEE 2003]
(a) 3 → 2 (b) 5 → 2 (c) 4 → 1 (d) 2 → 5
209. If electron falls from n = 3 to n = 2, then emitted energy is [AFMC 1997; MP PET 2003]
(a) 10.2 eV (b) 12.09 eV (c) 1.9 eV (d) 0.65 eV
210. The emission spectrum of hydrogen is found to satisfy the expression for the energy change. ∆E (in joules) such that
Jnn
E
−×=∆
22
21
111018.2 where ....3,2,11 =n and ....4,3,22 =n The spectral lines correspond to Paschen series to [UPSEAT 2002]
(a) 11 =n and 4,3,22 =n (b) 31 =n and 6,5,42 =n (c) 11 =n and 5,4,32 =n (d) 11 =n and =2n infinity
211. The energy required to dislodge electron from excited isolated H-atom., eVIE 6.131 = is [DCE 2000]
(a) = 13.6 eV (b) > 13.6 eV (c) < 13.6 and > 3.4 eV (d) eV4.3≤
212. If change in energy, sJhJE −×=×=∆ −− 348 1064.6,103)( and ,/103 8 smc ×= then wavelength of the light is
[CBSE 2000]
(a) 31036.6 × Å (b) 51036.6 × Å (c) 181064.6 −× Å (d) 181036.6 × Å
213. The value of Planck’s constant is Js341063.6 −× . The velocity of light is 18100.3 −× ms . Which value is closest to the wavelength
in nanometres of a quantum of light with frequency of 115108 −× s [CBSE 2003]
(a) 7103 × (b) 25102 −× (c) 18105 −× (d) 1104 ×
214. The wavelength of a spectral line for an electronic transition is inversely related to [IIT 1988]
(a) The number of electrons undergoing the transition
(b) The nuclear charge of the atom
(c) The difference in the energy of the energy levels involved in the transition
(d) The velocity of the electron undergoing the transition
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215. If wavelength of photon is m11102.2 −× , Jsh 34106.6 −×= , then momentum of photon is [MP PET 1999]
252. If uncertainty in the position of an electron is zero, the uncertainty in its momentum would be [CPMT 1988]
(a) Zero (b) < h / 2λ (c) > h / 2λ (d) Infinite
253. The position of both an electron and a helium atom is known within 1.0 nm and the momentum of the electron is known within 1261050 −−× mskg . The minimum uncertainty in the measurement of the momentum of the helium atom is
254. Assertion (A): The position of an electron can be determined exactly with the help of an electron microscope.
Reason (R): The product of uncertainty in the measurement of its momentum and the uncertainty in the measurement of the position cannot be less than a finite limit. [NDA 1999]
(a) Both A and R are true and R is the correct explanation of A
(b) Both A and R are true but R is not the correct explanation of A
(c) A is true but R is false
(d) A is false but R is true
AAddvvaannccee Level
255. The uncertainty in the position of an electron (mass = )101.9 28 g−× moving with a velocity of 14100.3 −× scm accurate upto
0.001% will be (Use π4
h in the uncertainty expression, where sergh −×= −2710626.6 ) [CBSE 1995]
(a) 51092.1 −× cm (b) 7.68 cm (c) 5.76 cm (d) 3.84 cm
256. The uncertainty in the position of a moving bullet of mass 10 gm is m510 − . Calculate the uncertainty in its velocity [DCE 1999]
292. If n = 3, then the value of ‘l’ which is incorrect [CPMT 1994]
(a) 0 (b) 1 (c) 2 (d) 3
293. The angular momentum of an electron depends on [BHU 1978, NCERT 1981]
(a) Principal quantum number (b) Azimuthal quantum number
(c) Magnetic quantum number (d) All of these
294. The shape of an orbital is given by the quantum number [NCERT 1984; MP PMT 1996]
(a) n (b) l (c) m (d) s
295. Which of the following set of quantum number is not valid [AIIMS 2001]
(a) n = 1, l = 2 (b) n = 2, m = 1 (c) n = 3, l = 0 (d) n = 2, l = 0
296. 2p orbital have [NCERT 1981; MP PMT 1993, 97]
(a) n = 1, l = 2 (b) n = 1, l = 0 (c) n = 2, l = 1 (d) n = 2, l = 0
297. The maximum number of electrons which each sub-shell can occupy is [PU CET 1989]
(a) 22n (b) 2n (c) 2 ( 2l + 1) (d) 2l + 1
298. Which of the following represent the correct sets of the four quantum numbers of a 4d electron [MLNR 1992; UPSEAT 2001]
(a) 4, 3, 2, 2
1+ (b) 4, 2, 1, 0 (c) 4, 3, – 2,
2
1+ (d) 4, 2, 1,
2
1−
299. For the n = 2 energy level, how many orbitals of all kinds are possible [Bihar CEE 1995]
(a) 2 (b) 3 (c) 4 (d) 5
300. The total number of orbitals in an energy level designated by principal quantum number n, is equal to [AIIMS 1997]
(a) 2n (b) 22 n (c) n (d) 2n
301. The quantum numbers for the outermost electron of an element are given below as n = 2, l = 0, m = 0,2
1+=s . The atoms is
[EAMCET 1978]
(a) Lithium (b) Beryllium (c) Hydrogen (d) Boron
302. The maximum number of electrons in an atom with l = 2 and n = 3 is [MP PET / PMT 1998]
(a) 2 (b) 6 (c) 12 (d) 10
303. Correct set of four quantum numbers for valence electron of rubidium (Z = 37) is [IIT 1984; JIPMER 1999; UPSEAT 2003]
(a) 5, 0, 0, 2
1+ (b) 5, 1, 0,
2
1+ (c) 5, 1, 1,
2
1+ (d) 6, 0, 0,
2
1+
304. If n + l = 6, then total possible number of sub-shells would be [Rajasthan PMT 1997]
(a) 3 (b) 4 (c) 2 (d) 5
305. If value of azimuthal quantum number l is 2, then total possible values of magnetic quantum number will be
(a) 7 (b) 5 (c) 3 (d) 2
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306. Orbital angular momentum for a d-electron is [MP PET 2003]
(a) π2
6h (b)
π2
6h (c)
π2
12h (d)
π2
12h
307. The number of quantum numbers required to describe an electron in an atom completely is [CET Pune 1998]
(a) 1 (b) 2 (c) 3 (d) 4
308. An −e has magnetic quantum number as – 3, what is its principal quantum number [BHU 1998]
(a) 1 (b) 2 (c) 3 (d) 4
309. The quantum number which is designated by letters s, p, d and f instead of number is [BHU 1980]
(a) n (b) l (c) lm (d) sm
310. An electron having the quantum numbers n = 4 , l = 3, m = 0, s = – 1/2 would be in the orbital [Orissa JEE 1997]
(a) 3s (b) 3p (c) 4d (d) 4f
311. The magnetic quantum number of valence electron of sodium (Na) is [Rajasthan PMT 2002]
(a) 3 (b) 2 (c) 1 (d) 0
312. How many electrons can be fit into the orbitals that comprise the 3rd quantum shell n = 3 [MP PMT 1986, 87; Orissa JEE 1997]
(a) 2 (b) 8 (c) 18 (d) 32
313. A sub-shell l = 2 can take how many electrons [NCERT 1973, 78]
(a) 3 (b) 10 (c) 5 (d) 6
314. What is the maximum number of electrons which can be accommodated in an atom in which the highest principal quantum number value is 4 [MP PMT 2000]
(a) 10 (b) 18 (c) 32 (d) 54
315. How many electrons can be accommodated in a sub-shell for which n = 3, l = 1 [CBSE 1990]
(a) 8 (b) 6 (c) 18 (d) 32
316. If an electron has spin quantum number of 2
1+ and a magnetic quantum number of – 1, it cannot be presented in an
389. The atomic number of an element is 35. What is the total number of electrons present in all the p-orbitals of the ground state atom of that element [EAMCET (Engg.) 2003]
(a) 6 (b) 11 (c) 17 (d) 23
390. After np orbitals are filled, the next orbital filled will be
(a) sn )1( + (b) pn )2( + (c) dn )1( + (d) sn )2( +
391. The number of unpaired electrons in an 2O molecule is [MNR 1983]
(a) 0 (b) 1 (c) 2 (d) 3
392. How many unpaired electrons are present in +2Ni (atomic number = 28) cation
[IIT 1981; MNR 1984; MP PAT 1993; MP PMT 1995; Kerala PMT 2003]
(a) 0 (b) 2 (c) 4 (d) 6
393. The electronic configuration of calcium ion )( 2+Ca is [CMC Vellore 1991]
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403. The number of orbitals in 2p sub-shell is [NCERT 1973; MP PMT 1996]
(a) 6 (b) 2 (c) 3 (d) 4
404. The statements [AIIMS 1982]
(i) In filling a group of orbitals of equal energy, it is energetically preferable to assign electrons to empty orbitals rather than pair them into a particular orbital.
(ii) When two electrons are placed in two different orbitals, energy is lower if the spins are parallel are valid for
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423. For a given value of quantum number l, the number of allowed values of m is given by
(a) l + 2 (b) 2l + 2 (c) 2l + 1 (d) l + 1
424. The set of quantum numbers not applicable for an electron in an atom is [MLNR 1994]
(a) 1=n , 1=l , 1=lm , 2
1+=sm (b) 1=n , 0=l , 0=lm ,
2
1+=sm
(c) 1=n , 0=l , 0=lm , 2
1−=sm (d) 2=n , 0=l , 0=lm ,
2
1+=sm
425. Which of the following statements is not correct for an electron that has the quantum numbers n = 4 and m = 2 [MLNR 1993]
(a) The electron may have the quantum number 2
1+=s (b) The electron may have the quantum number l = 2
(c) The electron may have the quantum number l = 3 (d) The electron may have the quantum number l = 0, 1, 2, 3
426. The electrons identified by quantum numbers n and l (i) n = 4, l = 1 (ii) n = 4, l = 0 (iii) n = 3, l = 2 (iv) n = 3, l = 1 can be placed in order of increasing energy from the lowest to highest, as [IIT 1999]
(a) (iv) < (ii) < (iii) < (i) (b) (ii) < (iv) < (i) < (iii) (c) (i) < (iii) < (ii) < (iv) (d) (iii) < (i) < (iv) < (ii)
427. Which of the following sets of quantum numbers is not allowed [Orissa JEE 1997]
(a) n = 1, l = 0, m = 0, s = + 1/2 (b) n = 1, l = 1, m = 0, s = – 1/2
(c) n = 2, l = 1, m = 1, s = + 1/2 (d) n = 2, l = 0, m =0, s = – ½
428. What are the values of the orbital angular momentum of an electron in the orbitals 1s, 3s, 3d and 2p
450. When the value of the principal quantum number n is 3, the permitted values of the azimuthal quantum number l and the magnetic quantum numbers m, are
l m
(a) 0 0
1 + 1, 0, – 1
2 + 2,+ 1, 0, – 1, – 2
(b) 1 1
2 + 2, 1, – 2
3 + 3,+ 2, 1, – 2, – 3
(c) 0 0
1 1, 2, 3
2 + 3, + 2, 1, – 2, – 3
(d) 1 0, 1
2 0, 1, 2
3 0, 1, 2, 3
451. Five valence electrons of 15 P are labelled as
If the spin quantum of B and Z is + 1 /2, the group of electrons with three of the quantum number same are [JIPMER 1997]
(a) AB, XYZ, BY (b) AB (c) XYZ, AZ (d) AB, XYZ
452. The quantum numbers + 1/ 2 and – 1/ 2 for the electron spin represent [IIT Screening 2001]
(a) Rotation of the electron in clockwise and anticlockwise direction respectively
(b) Rotation of the electron in anticlockwise and clockwise direction respectively
(c) Magnetic moment of the electron pointing up and down respectively
(d) Two quantum mechanical spin states with have no classical analogue
453. Which of the following violates the Pauli exclusion principle
(a) (b) (c) (d)
454. Which of the following violates the Aufbau principle
(a) (b) (c) (d)
455. Which of the following electronic configurations have the highest exchange energy
(a) (b)
(c) (d)
456. Which of the following set of quantum numbers is permissible [AIIMS 2001]
(a) n = 3; l = 2; m = 2 and 2
1+=s (b) n = 3; l = 4; m = 0; and
2
1−=s
(c) n = 4; l = 0; m = 2; and 2
1+=s (d) n = 4; l = 4; m = 3; and
2
1+=s
457. Which of the following sets of quantum number is not possible [MP PET 2001]
(a) n = 3; l = + 2; m = 0;2
1+=s (b) n = 3; l = 0; m = 0;
2
1−=s
AB X Y Z 3s 3p
↑↓ ↑↓ ↑↓ ↑ ↓ ↑ ↑↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑
2s 2p
↑↓ ↑↓
2s 2p
↑↓ ↑↓ ↑↓ ↑
2s 2p
↑↑ ↑↓ ↑↓ ↑
2s 2p
↑ ↑↓ ↑ ↑
3d 4s
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
3d 4s
↑ ↑ ↑ ↑ ↑ ↑ 3d 4s
↑ ↑ ↑ ↑
3d 4s
↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑
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(c) n = 3; l = 0; m = – 1;2
1+=s (d) n = 3; l = 1; m = 0;
2
1−=s
458. Which of the following set of quantum numbers belong to highest energy [CPMT 1999]
(a) n = 4, l = 0, m = 0,2
1+=s (b) n = 3, l = 0, m = 0,
2
1+=s
(c) n = 3, l = 1, m = 1,2
1+=s (d) n = 3, l = 2, m = 1,
2
1+=s
459. Assertion (A) : The cation energy of an electron is largely determined by its principal quantum number
Reason (R) : The principal quantum number n is a measure of the most probable distance of finding the electron around the nucleus
[AIIMS 1996]
(a) Both A and R are true statements and R is the correct explanation of A
(b) Both A and R are true statements and R is not the correct explanation of A
(c) A is true but R is a false statement
(d) Both A and R are false statements
460. Which of the following set of quantum number is not possible [Pb. PMT 2002]
n l 1m 2m
(a) 3 2 1 + 1/2
(b) 3 2 1 – 1/2
(c) 3 2 1 0
(d) 5 2 – 1 + 1/2
461. For the energy levels in an atom, which one of the following statements is correct [AIIMS 1983]
(a) There are seven principal electron energy levels
(b) The second principal energy level can have four orbitals and contains a maximum of eight electrons
(c) The M energy level can have maximum of 32 electrons
(d) The 4s sub-energy level is at a higher energy than the 3d sub-energy level
462. The orbital diagram in which the Aufbau’s principle is violated is [IIT 1988; AMU 1999]
2s xp2 yp2 zp2
(a) ↑↓ ↑↓ ↑
(b) ↑ ↑↓ ↑ ↑
(c) ↑↓ ↑ ↑ ↑
(d) ↑↓ ↑↓ ↑↓ ↑
463. The maximum probability of finding an electron in the xyd orbital is [MP PET 1999]
(a) Along the x-axis (b) Along the y-axis
(c) At an angle of 450 from the x and y-axes (d) At an angle of 900 from the x and y-axes
464. Krypton )(36 Kr has the electronic configuration 610218 434)( pdsAr . The 37th electron will go into which one of the
following sub-levels [CBSE 1989; CPMT 1989; EAMCET 1991]
(a) 4f (b) 4d (c) 3p (d) 5s
465. Which one is in the ground state [Delhi PMT 1996]
(a) (b)
(c) (d)
↑
↑↓
↑ ↑
↑
↑↓
↑ ↑ ↑↓
↑
↑ ↑ ↑
↑ ↑
↑↓
↑ ↑ ↑
↑
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468. Which of the following has maximum energy [AIIMS 2002]
(a) (b)
(c) (d)
469. Elements upto atomic number 103 have been synthesized and studied. If a newly discovered element is found to have an atomic number 106, its electronic configuration will be [AIIMS 1980]
(a) Isotopes (b) Fissionable (c) Not fissionable (d) None of these
489. The valence electron in the carbon atom are [MLNR 1982]
(a) 0 (b) 2 (c) 4 (d) 6
490. The atomic number of an element is 35 and mass number is 81. The number of electrons in the outer most shell is [UPSEAT 2001]
(a) 7 (b) 6 (c) 5 (d) 3
491. The atomic weight of an element is double its atomic number. If there are four electrons in 2p orbital, the element is [AMU 1983]
(a) C (b) N (c) O (d) Ca
492. If electron, hydrogen, helium and neon nuclei are all moving with the velocity of light, then the wavelengths associated with these particles are in the order [MP PET 1993]
(a) Electron > hydrogen > helium > neon (b) Electron > helium > hydrogen > neon
(c) Electron < hydrogen < helium < neon (d) Neon < hydrogen < helium < electron
493. When atoms are bombarded with alpha particles, only a few in million suffer deflection, others pass out undeflected. This is because
[MNR 1979; NCERT 1980; AFMC 1995]
(a) The force of repulsion on the moving alpha particle is small
(b) The force of attraction on the alpha particle to the oppositely charged electrons is very small
(c) There is only one nucleus and large number of electrons
(d) The nucleus occupies much smaller volume compared to the volume of the atom
494. The total number of valence electrons in 4.2 gm of −3N ion is ( AN is the Avogadro’s number) [CBSE 1994]
AAddvvaannccee Level
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(a) AN6.1 (b) AN2.3 (c) AN1.2 (d) AN2.4
495. In neutral atom, which particles are equivalent [Rajasthan PMT 1997]
(a) ++ ep , (b) +− ee , (c) +− pe , (d) 0, np+
496. An element have atomic weight 40 and it’s electronic configuration is 62622 33221 pspss . Then its atomic number and number of
neutrons will be [Rajasthan PMT 2002]
(a) 18 and 22 (b) 22 and 18 (c) 26 and 20 (d) 40 and 18
497. Which phrase would be incorrect to use [AMU (Engg.) 1997]
(a) A molecule of a compound (b)A molecule of an element (c)An atom of an element (d) None of these
498. Splitting of signals is caused by [Pb. PMT 2000]
(a) Proton (b) Neutron (c) Positron (d) Electron
499. Choose the correct statement
(a) A node is a point in space where the wave function )(Ψ has zero amplitude.
(b) The number of peaks in radial distribution is n – l
(c) Radial probability density )(4)( 2,
2, rRrrp lnln π= .
(d) 2Ψ represents the atomic orbital
(e) All the above
500. Which of the following electronic configurations have zero spin multiplicity