Chapter 22 Atoms free study material by TEACHING CARE online tuition and coaching classes

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Important Atomic Models.

(1) Thomson’s model

J.J. Thomson gave the first idea regarding structure of atom. According to this model.

(i) An atom is a solid sphere in which entire and positive charge and it’s mass is uniformly distributed and in which negative charge (i.e. electron) are embedded like seeds in watermelon.

Success and failure

Explained successfully the phenomenon of thermionic emission, photoelectric emission and ionization.

The model fail to explain the scattering of a– particles and it cannot explain the origin of spectral lines observed in the spectrum of hydrogen and other atoms.

(2) Rutherford’s model

Rutherford’s a-particle scattering experiment

Rutherford performed experiments on the scattering of alpha particles by extremely thin gold foils and made the following observations

Number of scattered particles :

Most of the a-particles pass through the foil straight away
Some of them are deflected through small
A few a-particles (1 in 1000) are deflected through the angle more than 90^o.
A few a -particles (very few) returned back e. deflected by 180^o.
Distance of closest approach (Nuclear dimension)

The minimum distance from the nucleus up to which the a-particle approach, is called the distance of closest

approach (r ). From figure r = 1 . Ze2 ; E = 1 mv² =

K.E. of a-particle

0 0 4pe E 2

Impact parameter (b) : The perpendicular distance of the velocity vector ( v ) of the a-particle from the centre of the nucleus when it is far away from the nucleus is known as impact It is given as

b = Ze ² cot(q / 2)

Þ b µ cot(q / 2)

4pe

æ 1 mv2 ö

₀ ç 2 ÷

Note : @If t is the thickness of the foil and N is the number of a-particles scattered in a particular direction

(q = constant), it was observed that

N = constant Þ

N1 = t1 .

N 2 t 2

After Rutherford’s scattering of a-particles experiment, following conclusions were made as regard as atomic structure :

Most of the mass and all of the charge of an atom concentrated in a very

small region is called atomic nucleus.

Nucleus is positively charged and it’s size is of the order of 10^–15 m » 1 Fermi.
In an atom there is maximum empty space and the electrons revolve around the nucleus in the same way as the planets revolve around the

10–10 m

leus

Size of the nucleus = 1 Fermi = 10^–15 m

Size of the atom 1 Å = 10^–10 m

Draw backs

Stability of atom : It could not explain stability of atom because according to classical electrodynamic theory an accelerated charged particle should continuously radiate Thus an electron moving in an circular path around the nucleus should also radiate energy and thus move into smaller and

smaller orbits of gradually decreasing radius and it should ultimately fall into nucleus.

According to this model the spectrum of atom must be continuous where as practically it is a line
It did not explain the distribution of electrons outside the

(3) Bohr’s model

Bohr proposed a model for hydrogen atom which is also applicable for some lighter atoms in which a single electron revolves around a stationary nucleus of positive charge Ze (called hydrogen like atom)

Bohr’s model is based on the following postulates.

The electron can revolve only in certain discrete non-radiating orbits, called stationary orbits, for which total

angular momentum of the revolving electrons is an integral multiple of h (= h)

ç ÷

i.e. L = n æ h ö = mvr;

where n = 1, 2, 3,……. = Principal quantum number

è ø

The radiation of energy occurs only when an electron jumps from one permitted orbit to

When electron jumps from higher energy orbit (E₁) to lower energy orbit (E₂) then difference of energies of these orbits i.e. E₁ – E₂ emits in the form of photon. But if electron goes from E₂ to E₁ it absorbs the same amount of

energy.

E₁

E₂

Emission

E₁ – E₂ = hn

E₁

E₁ – E₂ = hn

E₂

Absorption

Note : @According to Bohr theory the momentum of an e ^–

revolving in second orbit of H atom will be h

@ For an electron in the n^th orbit of hydrogen atom in Bohr model, circumference of orbit

= nl ; where l = de-Broglie wavelength.

Bohr’s Orbits (For Hydrogen and H₂-Like Atoms).

(1) Radius of orbit

For an electron around a stationary nucleus the electrostatics force of attraction provides the necessary centripetal force

i.e. 1 (Ze)e = mv²

……. (i) also mvr = nh

…….(ii)

4pe ₀ r ² r 2p

From equation (i) and (ii) radius of n^th orbit

rn =

n2h2

n²h²e n²

2 = ⁰ = 0.53 Å

éwhere k = 1 ù

4p kZme

pmZe² Z

ë 4pe ₀ û

Þ r µ n

Note : @The radius of the innermost orbit (n = 1) hydrogen atom (z = 1) is called Bohr’s radius a₀ i.e.

a₀ = 0.53 Å .

(2) Speed of electron

From the above relations, speed of electron in n^th orbit can be calculated as

2pkZe ²

Ze ²

æ c ö Z ₆ Z ^v

vn =

nh 2e

nh = ç

÷. = 2.2 ´ 10

n m / sec

₀ è 137 ø n

where (c = speed of light 3 ´ 10⁸ m/s)

Note : @The ratio of speed of an electron in ground state in Bohr’s first orbit of hydrogen atom to velocity of

light in air is equal to

e ² = 1

(where c = speed of light in air)

(3) Some other quantities

2e ₀ ch 137

For the revolution of electron in n^th orbit, some other quantities are given in the following table

(4) Energy

Potential energy : An electron possesses some potential energy because it is found in the field of nucleus

potential energy of electron in n^th orbit of radius r

is given by U = k. (Ze)(-e) = – kZe 2

rn rn

Kinetic energy : Electron posses kinetic energy because of it’s Closer orbits have greater kinetic energy than outer ones.

As we know

mv²

= k.(Ze)(e)

Þ Kinetic energy

K = kZe ²

= |U |

r_n r ² 2r_n 2

Total energy : Total energy (E) is the sum of potential energy and kinetic energy e. E = K + U

Þ E = –

kZe ²

also rn =

n²h²e ₀

. Hence

E = – ç

me ⁴ ö z ²

2 2 2

= – ç

me ⁴

ö 2

÷ ch

= –Rch Z 2

= -13.6

Z 2 eV

2rn

pmze

ç 8e ₀ h ÷

ç 8e ₀ ch ÷ n n n

where

R = me ⁴

8e ²ch³

= Rydberg’s constant = 1.09 ´ 10⁷ per metre

Note : @Each Bohr orbit has a definite energy

@ For hydrogen atom (Z = 1) Þ

E = – 13.6 eV

n n2

@ The state with n = 1 has the lowest (most negative) energy. For hydrogen atom it is E₁ = – 13.6 eV.

@ Rch = Rydberg’s energy ~– 2.17 ´ 10 ^-18 J ~– 31.6 eV .

@ E = –K = U .

Ionisation energy and potential : The energy required to ionise an atom is called ionisation It is the energy required to make the electron jump from the present orbit to the infinite orbit.

æ Z ² ö 13.6Z ²

Hence Eionisation = E¥ – En = 0 – ç- 13.6 n2 ÷ = + n2 eV

For H₂-atom in the ground state

Eionisation

= + 13.6(1)² = 13.6 eV

The potential through which an electron need to be accelerated so that it acquires energy equal to the

ionisation energy is called ionisation potential.

Vionisation

= Eionisation

Excitation energy and potential : When the electron is given energy from external source, it jumps to higher energy This phenomenon is called excitation.

The minimum energy required to excite an atom is called excitation energy of the particular excited state and corresponding potential is called exciting potential.

EExcitation

= EFinal

EInitial

and

VExcitation

= Eexcitation

Binding energy (B.E.) : Binding energy of a system is defined as the energy released when it’s constituents are brought from infinity to form the system. It may also be defined as the energy needed to separate it’s constituents to large If an electron and a proton are initially at rest and brought from large distances to form a hydrogen atom, 13.6 eV energy will be released. The binding energy of a hydrogen atom is therefore 13.6 eV.

Note : @For hydrogen atom principle quantum number n = .

(5) Energy level diagram

The diagrammatic description of the energy of the electron in different orbits around the nucleus is called energy level diagram.

Energy level diagram of hydrogen/hydrogen like atom

	n = ¥	Infinite	Infinite	E_¥ = 0 eV	0 eV	0 eV
n = 4		Fourth	Third	E₄ = – 0.85 eV	– 0.85 Z²	+ 0.85 eV
n = 3		Third	Second	E₃ = – 1.51 eV	– 1.51 Z²	+ 1.51 eV
n = 2		Second	First	E₂ = – 3.4 eV	– 3.4 Z²	+ 3.4 eV
n = 1		First	Ground	E₁ = – 13.6 eV	– 13.6 Z²	+ 13.6 eV
Principle		Orbit	Excited	Energy for H₂	Energy for H₂	Ionisation energy
quantum			state	– atom	– like atom	from this level (for
number						H₂ – atom)

Note : @In hydrogen atom excitation energy to excite electron from ground state to first excited state will be

– 3.4 – (-13.6) = 10.2 eV .

and from ground state to second excited state it is [ – 1.51 – (-13.6) = 12.09 eV ].

@ In an

H ₂ atom when

e ^– makes a transition from an excited state to the ground state it’s

kinetic energy increases while potential and total energy decreases.

(6) Transition of electron

When an electron makes transition from higher energy level having energy E₂(n₂) to a lower energy level having energy E₁ (n₁) then a photon of frequency n is emitted

(i) Energy of emitted radiation

Rc h Z ² æ Rch Z ² ö

E₂

₂ æ 1 1 ö

n₂

DE,n,l

DE = E2 – E1 = – ç-

÷ = 13.6Z ç – ÷

n2 è

n1 ø

ç n1

n2 ÷

Emission

(ii) Frequency of emitted radiation

–

DE E₂ – E₁ ₂ æ 1 1 ö

DE = hn Þ n = =

h = Rc Z

ç n2

n2 ÷

(iii)

–

Wave number/wavelength

è 1 2 ø

1 n 1

2 æ 1

–

1 ö 13.6Z ² æ 1 1 ö

Wave number is the number of waves in unit length n

= l = c

Þ l = RZ

ç n2

n2 ÷

hc ç n²

n2 ÷

è 1 2 ø è 1 2 ø

Number of spectral lines : If an electron jumps from higher energy orbit to lower energy orbit it emits raidations with various spectral

If electron falls from orbit n₂ to n₁ then the number of spectral lines emitted is given by

N = (n2 – n1 + 1)(n2 – n1 )

E 2

If electron falls from n^th orbit to ground state (i.e. n₂

= n and n₁ = 1) then number of spectral lines emitted

N = n(n – 1)

E 2

Note : @Absorption spectrum is obtained only for the transition from lowest energy level to higher energy levels. Hence the number of absorption spectral lines will be (n – 1).

Recoiling of an atom : Due to the transition of electron, photon is emitted and the atom is recoiled

–

h ₂ æ 1 1 ö

Recoil momentum of atom = momentum of photon = l

= hRZ

ç n2

n2 ÷

è 1 2 ø

Also recoil energy of atom = p 2

= h2

2ml²

(where m = mass of recoil atom)

(7) Draw backs of Bohr’s atomic model

It is valid only for one electron atoms, g. : H, He⁺, Li⁺², Na⁺¹ etc.
Orbits were taken as circular but according to Sommerfield these are
Intensity of spectral lines could not be
Nucleus was taken as stationary but it also rotates on its own
It could not be explained the minute structure in spectrum
This does not explain the Zeeman effect (splitting up of spectral lines in magnetic field) and Stark effect (splitting up in electric field)
This does not explain the doublets in the spectrum of some of the atoms like sodium (5890Å & 5896Å)

Hydrogen Spectrum and Spectral Series.

When hydrogen atom is excited, it returns to its normal unexcited (or ground state) state by emitting the energy it had absorbed earlier. This energy is given out by the atom in the form of radiations of different wavelengths as the electron jumps down from a higher to a lower orbit. Transition from different orbits cause

different wavelengths, these constitute spectral series which are characteristic of the atom emitting them. When observed through a spectroscope, these radiations are imaged as sharp and straight vertical lines of a single colour.

Spectral series

The spectral lines arising from the transition of electron forms a spectra series.

Mainly there are five series and each series is named after it’s discover as Lymen series, Balmer series, Paschen series, Bracket series and Pfund
According to the Bohr’s theory the wavelength of the radiations emitted from hydrogen atom is given by

1 = R é 1 – 1 ù

l ê n² n² ú

ë 1 2 û

where n₂ = outer orbit (electron jumps from this orbit), n₁ = inner orbit (electron falls in this orbit)

First line of the series is called first member, for this line wavelength is maximum (l_max)

Last line of the series (n₂ = ¥) is called series limit, for this line wavelength is minimum (l_min)

Quantum Numbers.

An atom contains large number of shells and subshells. These are distinguished from one another on the basis of their size, shape and orientation (direction) in space. The parameters are expressed in terms of different numbers called quantum number.

Quantum numbers may be defined as a set of four number with the help of which we can get complete information about all the electrons in an atom. It tells us the address of the electron i.e. location, energy, the type of orbital occupied and orientation of that orbital.

Principal Quantum number (n) : This quantum number determines the main energy level or shell in which the electron is The average distance of the electron from the nucleus and the energy of the electron

depends on it.

En µ n2

and

rn µ n²

(in H-atom)

The principal quantum number takes whole number values, n = 1, 2, 3, 4,….. ¥

(2) Orbital quantum number (l) or azimuthal quantum number (l)

This represents the number of subshells present in the main shell. These subsidiary orbits within a shell will be denoted as 1, 2, 3, 4 … or s, p, d, f … This tells the shape of the subshells.

The orbital angular momentum of the electron is given as L =

h (for a particular value of n).

For a given value of n the possible values of l are l = 0, 1, 2, ….. upto (n – 1)

Magnetic quantum number (m_l) : An electron due to it’s angular motion around the nucleus generates an electric field. This electric field is expected to produce a magnetic field. Under the influence of external magnetic field, the electrons of a subshell can orient themselves in certain preferred regions of space around the nucleus called

The magnetic quantum number determines the number of preferred orientations of the electron present in a subshell.

The angular momentum quantum number m can assume all integral value between – l to +l including zero.

Thus m_l can be – 1, 0, + 1 for l = 1. Total values of m_l associated with a particular value of l is given by (2l + 1).

Spin (magnetic) quantum number (m_s) : An electron in atom not only revolves around the nucleus but also spins about its own axis. Since an electron can spin either in clockwise direction or in anticlockwise Therefore for any particular value of magnetic quantum number, spin quantum number can have two

values, i.e.

m = 1

s 2

(Spin up) or

m = – 1

s 2

(Spin down)

This quantum number helps to explain the magnetic properties of the substance.

Electronic Configurations of Atoms.

The distribution of electrons in different orbitals of an atom is called the electronic configuration of the atom.

The filling of electrons in orbitals is governed by the following rules.

(1) Pauli’s exclusion principle

“It states that no two electrons in an atom can have all the four quantum number (n, l, m_l and m_s) the same.”

It means each quantum state of an electron must have a different set of quantum numbers n, l, m_l and m_s. This principle sets an upper limit on the number of electrons that can occupy a shell.

N max

in one shell = 2n²; Thus N_max in K, L, M, N …. shells are 2, 8, 18, 32,

Note : @ The maximum number of electrons in a subshell with orbital quantum number l is 2(2l + 1).

(2) Aufbau principle

Electrons enter the orbitals of lowest energy first.

As a general rule, a new electron enters an empty orbital for which (n + l ) is minimum. In case the value

(n + l) is equal for two orbitals, the one with lower value of n is filled first.

Thus the electrons are filled in subshells in the following order (memorize) 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, ……

(3) Hund’s Rule

When electrons are added to a subshell where more than one orbital of the same energy is available, their spins remain parallel. They occupy different orbitals until each one of them has at least one electron. Pairing starts only when all orbitals are filled up.

Pairing takes place only after filling 3, 5 and 7 electrons in p, d and f orbitals, respectively.

Example: 1 The ratio of areas within the electron orbits for the first excited state to the ground state for hydrogen atom is [BCECE 2004]

(a) 16 : 1 (b) 18 : 1 (c) 4 : 1 (d) 2 : 1

Solution : (a) For a hydrogen atom

2 r 2 n4

pr ² n⁴

A n⁴ 2⁴

A 16

Radius r µ n

Þ 1 = 1 Þ 1 = 1

Þ 1 = 1 =

= 16

Þ 1 =

2 4 2 4

2 2 2

2 4 14

A2 1

Example: 2 The electric potential between a proton and an electron is given by

V = V0

ln r , where

r₀ is a constant.

Assuming Bohr’s model to be applicable, write variation of rn with n, n being the principal quantum number

[IIT-JEE (Screening) 2003]

(a)

rn µ n

(b)

rn µ 1 / n

(c)

rn µ n²

rn µ 1 / n²

Solution : (a) Potential energy U = eV = eV0 ln r

\ Force

F = – dU

= eV0 . The force will provide the necessary centripetal force. Hence

mv ² = eV₀

Þ v =

…..(i) and

mvr = nh

…..(ii)

r r 2p

Dividing equation (ii) by (i) we have

mr = æ nh ö

ç ÷

or r µ n

è ø

Example: 3 The innermost orbit of the hydrogen atom has a diameter 1.06 Å. The diameter of tenth orbit is

(a) 5.3 Å (b) 10.6 Å (c) 53 Å (d) 106 Å

[UPSEAT 2002]

r æ n ö ²

d æ n ö ²

d æ 10 ö ²

Solution : (d) Using r µ n² Þ ² = ç ² ÷

or ² = ç ² ÷

Þ ² = ç ÷ Þ d = 106 Å

r1 è n1 ø

d1 è n1 ø

1.06

è 1 ø

Example: 4 Energy of the electron in n^th orbit of hydrogen atom is given by En

= – 13.6 eV . The amount of energy needed

to transfer electron from first orbit to third orbit is [MH CET 2002; Kerala PMT 2002]

(a) 13.6 eV (b) 3.4 eV (c) 12.09 eV (d) 1.51 eV

Solution : (c) Using

E = – 13.6 eV

For n = 1 , E1 = –13.6 = -13.6 eV

and for n = 3 E3 = – 13.6 = -1.51 eV

So required energy = E₃ – E₁ = -1.51 – (-13.6) = 12.09 eV

Example: 5 If the binding energy of the electron in a hydrogen atom is 13.6 eV, the energy required to remove the

electron from the first excited state of Li ⁺ ⁺ is [AIEEE 2003]

(a) 122.4 eV (b) 30.6 eV (c) 13.6 eV (d) 3.4 eV

13.6 ´ Z ²

Solution : (b) Using

En = – eV

For first excited state n = 2 and for Li ⁺⁺ , Z = 3

\ E = – 13.6 ´ 3² = – 13.6 ´ 9 = -30.6 eV . Hence, remove the electron from the first excited state of Li ⁺⁺ be 30.6 eV

22 4

Example: 6 The ratio of the wavelengths for 2 ® 1 transition in Li⁺⁺, He⁺ and H is [UPSEAT 2003]

(a) 1 : 2 : 3 (b) 1 : 4 : 9 (c) 4 : 9 : 36 (d) 3 : 2 : 1

1 ₂ æ 1 1 ö 1

1 1 1

Solution : (c) Using

= RZ ç –

÷ Þ µ

2 Þ lLi : lHe+

: l_H =

: : = 4 : 9 : 36

l ç n1

n2 ÷ Z

9 4 1

Example: 7 Energy E of a hydrogen atom with principal quantum number n is given by

E = –13.6 eV . The energy of a

Solution : (a)

photon ejected when the electron jumps n = 3 state to n = 2 state of hydrogen is approximately

[CBSE PMT/PDT Screening 2004]

DE 13 æ 1 1 ö 5

(a) 1.9 eV (b) 1.5 eV (c) 0.85 eV (d) 3.4 eV

= .6 – = 13.6 ´ = 1.9 eV

ç 22 32 ÷ 36

Example: 8 In the Bohr model of the hydrogen atom, let R, v and E represent the radius of the orbit, the speed of electron and the total energy of the electron respectively. Which of the following quantity is proportional to the quantum number n [KCET 2002]

R/E (b) E/v (c) RE (d) vR

e ₀n²h²

Solution : (d) Rydberg constant R = pmZe 2

Velocity v =

Ze ²

2e ₀nh

and energy

E = –

mZ ²e ⁴

8e ²n²h²

Now, it is clear from above expressions R.v µ n

Example: 9 The energy of hydrogen atom in nth orbit is E_n, then the energy in nth orbit of singly ionised helium atom will be [CBSE PMT 2001]

(a) 4E_n (b) E_n/4 (c) 2E_n (d) E_n/2

13.6 Z ²

E æ Z ö ²

æ 1 ö ²

Solution : (a) By using

E = –

Þ H = ç H ÷

= ç ÷

Þ EHe = 4 En .

n2 EHe è ZHe ø

è 2 ø

Example: 10 The wavelength of radiation emitted is l₀

when an electron jumps from the third to the second orbit of

hydrogen atom. For the electron jump from the fourth to the second orbit of the hydrogen atom, the wavelength of radiation emitted will be [SCRA 1998; MP PET 2001]

(a)

16 l (b)

20 l (c)

27 l (d)

25 l

25 ⁰

27 ⁰

20 ⁰

16 ⁰

Solution : (b) Wavelength of radiation in hydrogen atom is given by

1 = Ré 1 – 1 ù Þ 1 = Ré 1 – 1 ù = Ré 1 – 1 ù = 5 R

…..(i)

l ê n2 n2 ú l

êë 2²

3² úû

êë 4

9 úû 36

ëê 1

₂ úû 0

and

1 = Ré 1 –

1 ù = Ré 1 –

1 ù = 3R

…..(ii)

l ¢ ëê 22

4 ² úû êë 4

16 úû 16

From equation (i) and (ii)

l ¢ = 5R ´ 16 = 20

Þ l ¢ = 20 l

l 36 3R 27

27 ⁰

Example: 11 If scattering particles are 56 for 90 ^o

angle then this will be at 60 ^o

angle [RPMT 2000]

(a) 224 (b) 256 (c) 98 (d) 108

Solution : (a) Using Scattering formula

é æ q

ö ù 4

é æ 90° ö ù 4

ê sinç ¹ ÷ ú êsinç ÷ ú 4

N µ 1 Þ

N 2 = ê è 2 ø ú

Þ N 2 = ê è 2 ø ú

= é sin 45°ù

= 4 Þ N

= 4 N

= 4 ´ 56 = 224

sin ⁴ (q / 2)

N₁ ê æ q ₂ ö ú

N₁ ê æ 60° ö ú

êë sin 30° úû 2 1

ê sinç

2 ÷ ú

êsinç 2 ÷ ú

ëê è

ø úû

ë è ø û

Example: 12 When an electron in hydrogen atom is excited, from its 4^th to 5^th stationary orbit, the change in angular momentum of electron is (Planck’s constant: h = 6.6 ´ 10 ^–³⁴ J–s ) [AFMC 2000]

(a)

4.16 ´ 10 ^–³⁴ J– s

(b)

3.32 ´ 10 ^–³⁴ J–s

(c)

1.05 ´ 10 ^–³⁴ J–s

(d)

2.08 ´ 10 ^–³⁴ J–s

Solution : (c) Change in angular momentum

DL = L

L = n2 h – n1h

Þ DL = h (n – n ) = 6.6 ´ 10 -34 (5 – 4) = 1.05 ´ 10 ^–³⁴ J-s

² ¹ 2p 2p

2p 2 1

2 ´ 3.14

Example: 13 In hydrogen atom, if the difference in the energy of the electron in n = 2 and n = 3 orbits is E, the ionization energy of hydrogen atom is [EAMCET (Med.) 2000]

(a) 13.2 E (b) 7.2 E (c) 5.6 E (d) 3.2 E

Solution : (b) Energy difference between n = 2 and n = 3;

E = Kæ 1

– 1 ö = Kæ 1 – 1 ö = 5 K

…..(i)

ç 22

32 ÷ ç

÷ 36

Ionization energy of hydrogen atom n

= 1 and n

= ¥ ;

E¢ = Kæ 1 – 1 ö = K

…..(ii)

From equation (i) and (ii)

1 2

E¢ = 36 E = 7.2 E

ç 12

¥ 2 ÷

Example: 14 In Bohr model of hydrogen atom, the ratio of periods of revolution of an electron in n = 2 and n = 1 orbits is

[EAMCET (Engg.) 2000]

(a) 2 : 1 (b) 4 : 1 (c) 8 : 1 (d) 16 : 1

3 Þ T2

= n3 = 23 = 8

Solution : (c) According to Bohr model time period of electron T µ n

Þ T2 = 8T1 .

1 3 13 1

Example: 15 A double charged lithium atom is equivalent to hydrogen whose atomic number is 3. The wavelength of

required radiation for emitting electron from first to third Bohr orbit in

Li++

will be (Ionisation energy of

hydrogen atom is 13.6 eV) [IIT-JEE 1985; UPSEAT 1999]

(a) 182.51 Å (b) 177.17 Å (c) 142.25 Å (d) 113.74 Å

Solution : (d) Energy of a electron in nth orbit of a hydrogen like atom is given by

E = -13.6 Z 2 eV , and Z = 3 for Li

n n2

Required energy for said transition

DE = E

– E = 13.6Z ² æ 1 –

1 ö = 13.6 ´ 3² é 8 ù = 108.8 eV = 108.8 ´ 1.6 ´ 10 ^–¹⁹ J

3 1 ç 12

32 ÷

êë 9 úû

Now using DE = hc

Þ l = hc

Þ l = 6.6 ´ 10 ^–³⁴ ´ 3 ´ 10⁸

108.8 ´ 1.6 ´ 10 ^–¹⁹

= 0.11374 ´ 10 ^–⁷ m Þ l = 113.74 Å

Example: 16 The absorption transition between two energy states of hydrogen atom are 3. The emission transitions between these states will be [MP PET 1999]

(a) 3 (b) 4 (c) 5 (d) 6

Solution : (d) Number of absorption lines = (n – 1) Þ 3 = (n – 1) Þ n = 4

Hence number of emitted lines = n(n – 1) = 4(4 – 1) = 6

2 2

Example: 17 The energy levels of a certain atom for 1st, 2nd and 3rd levels are E, 4E/3 and 2E respectively. A photon of wavelength l is emitted for a transition 3 ® 1. What will be the wavelength of emissions for transition 2 ® 1

[CPMT 1996]

(a) l/3 (b) 4l/3 (c) 3l/4 (d) 3l

Solution : (d) For transition 3 ® 1

DE = 2E – E = hc

Þ E = hc

…..(i)

For transition 2 ® 1

4 E – E = hc Þ E = 3hc

…..(ii)

From equation (i) and (ii)

3 l ¢ l ¢

l ¢ = 3l

Example: 18 Hydrogen atom emits blue light when it changes from n = 4 energy level to n = 2 level. Which colour of light would the atom emit when it changes from n = 5 level to n = 2 level [KCET 1993]

(a) Red (b) Yellow (c) Green (d) Violet

Solution : (d) In the transition from orbits

5 ® 2

more energy will be liberated as compared to transition from 4 ® 2. So

emitted photon would be of violet light.

Example: 19 A single electron orbits a stationary nucleus of charge +Ze, where Z is a constant. It requires 47.2 eV to excited electron from second Bohr orbit to third Bohr orbit. Find the value of Z [IIT-JEE 1981]

(a) 2 (b) 5 (c) 3 (d) 4

Solution : (b) Excitation energy of hydrogen like atom for n₂ ® n₁

DE = 13.6Z

Þ Z = 5

ç –

2 æ

÷ eV Þ 47

₂ æ

1 – 1

÷ = 13.6 ´

2 2

47.2 ´ 36

22 32

13.6 ´ 5

1 2

.2 = 13.6Z ç

Z Þ Z =

= 24.98 ~- 25

Example: 20 The first member of the Paschen series in hydrogen spectrum is of wavelength 18,800 Å. The short wavelength limit of Paschen series is [EAMCET (Med.) 2000]

(a) 1215 Å (b) 6560 Å (c) 8225 Å (d) 12850 Å

Solution : (c) First member of Paschen series mean it’s l_max = 144

Short wavelength of Paschen series means l_min = 9

Hence

lmax = 16 Þ l

= 7 ´ l

= 7 ´ 18,800 = 8225 Å .

lmin 7

min 16

max 16

Example: 21 Ratio of the wavelengths of first line of Lyman series and first line of Balmer series is

[EAMCET (Engg.) 1995; MP PMT 1997]

(a) 1 : 3 (b) 27 : 5 (c) 5 : 27 (d) 4 : 9

Solution : (c) For Lyman series

1 = R é 1 –

1 ù = 3R

…..(i)

l êë1² 2² úû 4

For Balmer series

1 = R é 1

– 1 ù = 5R

…..(ii)

l êë 2²

32 úû 36

From equation (i) and (ii)

lL1 = 5 .

l 27

Example: 22 The third line of Balmer series of an ion equivalent to hydrogen atom has wavelength of 108.5 nm. The ground state energy of an electron of this ion will be [RPET 1997]

(a) 3.4 eV (b) 13.6 eV (c) 54.4 eV (d) 122.4 eV

1 ₂ æ 1 1 ö 1

₇ ₂ æ 1 1 ö

Solution : (c) Using

= RZ ç – ÷ Þ

-9 = 1.1 ´ 10

´ Z ç 2 – 2 ÷

l ç n1 n2 ÷

108.5 ´ 10

è 2 5 ø

Þ 1

108.5 ´ 10 ^–⁹

= 1.1 ´ 10⁷ ´ Z ² ´

100

Þ Z ² =

100 = 4

108.5 ´ 10 ^–⁹ ´ 1.1 ´ 10 ^–⁷ ´ 21

Þ Z = 2

Now Energy in ground state

E = -13.6Z ² eV = -13.6 ´ 2² eV = -54.4 eV

Example: 23 Hydrogen (H), deuterium (D), singly ionized helium

(He⁺)

and doubly ionized lithium

(Li ⁺⁺ )

all have one

electron around the nucleus. Consider

n = 2 to

n = 1

transition. The wavelengths of emitted radiations are

l₁, l₂, l₃ and l₄ respectively. Then approximately [KCET 1994]

(a)

l₁ = l₂ = 4l₃ = 9l₄

(b)

4l₁ = 2l₂ = 2l₃ = l₄

(c)

l₁ = 2l₂ = 2

2l₃ = 3

2l₄ (d)

l₁ = l₂ = 2l₃ = 3l₄

Solution : (a) Using DE µ Z ²

(∵ n₁ and n₂ are same)

Þ hc µ Z ²

Þ lZ ² = constant Þ l₁

Z ² = l₂

Z ² = l₃

Z ² = l₄

Z ⁴ Þ l₁

´ 1 = l₂

´ 1² = l₃

´ 2² = l₄

´ 3³

Þ l₁ = l₂ = 4l₃ = 9l₄ .

Example: 24 Hydrogen atom in its ground state is excited by radiation of wavelength 975 Å. How many lines will be there in the emission spectrum [RPMT 2002]

(a) 2 (b) 4 (c) 6 (d) 8

Solution : (c) Using

1 = Ré 1 – 1 ù

Þ 1 = 1.097 ´ 10⁷ æ 1 – 1 ö

Þ n = 4

l ê n2

n2 ú

975 ´ 10 ^–¹⁰

ç 12

n 2 ÷

ëê 1

₂ úû è ø

Now number of spectral lines

N = n(n – 1) = 4(4 – 1) = 6 .

2 2

Example: 25 A photon of energy 12.4 eV is completely absorbed by a hydrogen atom initially in the ground state so that it is excited. The quantum number of the excited state is [UPSEAT 2000]

(a) n =1 (b) n= 3 (c) n = 4 (d) n = ¥

Solution : (c) Let electron absorbing the photon energy reaches to the excited state n. Then using energy conservation

Þ – 13.6 = -13.6 + 12.4 Þ – 13.6 = -1.2 Þ n ² = 13.6 = 12 Þ n = 3.46 ≃ 4

n 2 n 2

1.2

Example: 26 The wave number of the energy emitted when electron comes from fourth orbit to second orbit in hydrogen is

20,397 cm^–1. The wave number of the energy for the same transition in He⁺ is [Haryana PMT 2000]

(a) 5,099 cm^–1 (b) 20,497 cm^–1 (c) 40,994 cm^–1 (d) 81,998 cm^–1

1 æ 1 1 ö

n æ Z ö ² æ Z ö 2

Solution : (d) Using

= n = RZ ² ç – ÷ Þ n

µ Z ² Þ ² = ç 2 ÷ = ç ÷ = 4

Þ n ₂ = n ´ 4 = 81588 cm^–¹ .

l ç n2 n2 ÷

n 1 è Z1 ø è 1 ø

è 1 2 ø

Example: 27 In an atom, the two electrons move round the nucleus in circular orbits of radii R and 4R. the ratio of the time taken by them to complete one revolution is

(a) 1/4 (b) 4/1 (c) 8/1 (d) 1/8

Solution : (d) Time period T µ Z 2

For a given atom (Z = constant) So T µ n³ …..(i) and radius R µ n²

…..(ii)

T æ R

ö 3 / 2

æ R ö ³ ^/ ² 1

\ From equation (i) and (ii) T µ R ³ ^/ ² Þ ¹ = ç 1 ÷ = ç ÷ = .

T2 è R2 ø

è 4 R ø 8

Example: 28 Ionisation energy for hydrogen atom in the ground state is E. What is the ionisation energy of Li⁺⁺ atom in the 2^nd excited state

E (b) 3E (c) 6E (d) 9E

Z 2

Solution : (a) Ionisation energy of atom in nth state En =

For hydrogen atom in ground state (n = 1) and Z = 1 Þ E = E0

…..(i)

For

Li ⁺⁺ atom in 2^nd excited state n = 3 and Z = 3, hence

E¢ = E0 ´ 3² = E

32 0

…..(ii)

From equation (i) and (ii) E¢ = E .

Example: 29 An electron jumps from n = 4 to n = 1 state in H-atom. The recoil momentum of H-atom (in eV/C) is (a) 12.75 (b) 6.75 (c) 14.45 (d) 0.85

Solution : (a) The H-atom before the transition was at rest. Therefore from conservation of momentum

Photon momentum = Recoil momentum of H-atom or P

= hn

= E4 – E1

= -0.85eV -(-13.6 eV) = 12.75 eV

recoil c c c c

Example: 30 If elements with principal quantum number n > 4 were not allowed in nature, the number of possible elements would be

[IIT-JEE 1983; CBSE PMT 1991, 93; MP PET 1999; RPET 1993, 2001; RPMT 1999, 2003; J & K CET 2004]

(a) 60 (b) 32 (c) 4 (d) 64

Solution : (a) Maximum value of n = 4

So possible (maximum) no. of elements

N = 2 ´ 1² + 2 ´ 2² + 2 ´ 3² + 2 ´ 4 ² = 2 + 8 + 18 + 32 = 60 .

Tricky example: 1

If the atom 100 Fm²⁵⁷ follows the Bohr model and the radius of 100 Fm²⁵⁷ is n times the Bohr radius, then find n

[IIT-JEE (Screening) 2003]

(a) 100 (b) 200 (c) 4 (d) ¼

Solution : (d) (r ) = æ m2 ö(0.53 Å) = (n ´ 0.53 Å) Þ m2 = n

_m ç Z ÷ Z

ç ÷

è ø

m = 5 for 100 Fm ²⁵⁷ (the outermost shell) and z = 100

\ n = (5)² = 1

100 4

Tricky example: 2

An energy of 24.6 eV is required to remove one of the electrons from a neutral helium atom. The energy (in eV) required to remove both the electrons from a neutral helium atom is [IIT-JEE 1995]

(a) 79.0 (b) 51.8 (c) 49.2 (d) 38.2

Solution : (a) After the removal of first electron remaining atom will be hydrogen like atom.

So energy required to remove second electron from the atom E = 13.6 ´ 22 = 54.4 eV

\ Total energy required = 24.6 + 54.4 = 79 eV

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Important Atomic Models.

(1) Thomson’s model

Success and failure

(2) Rutherford’s model

Draw backs

(3) Bohr’s model

Bohr’s Orbits (For Hydrogen and H2-Like Atoms).

(1) Radius of orbit

(2) Speed of electron

(3) Some other quantities

(4) Energy

(5) Energy level diagram

Energy level diagram of hydrogen/hydrogen like atom

(6) Transition of electron

(i) Energy of emitted radiation

(ii) Frequency of emitted radiation

(iii)

Wave number/wavelength

(7) Draw backs of Bohr’s atomic model

Hydrogen Spectrum and Spectral Series.

Spectral series

Quantum Numbers.

(2) Orbital quantum number (l) or azimuthal quantum number (l)

Electronic Configurations of Atoms.

(1) Pauli’s exclusion principle

(2) Aufbau principle

(3) Hund’s Rule

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Bohr’s Orbits (For Hydrogen and H₂-Like Atoms).