Study

1. RUTHERFORD’S MODEL
Observation:
A. Most of the α - particles (nearly 99%)continued with their straight path.
B. Some of the α - particles passed very close to the centre of the atom
and deflected by small angles.
C. Very few particles thrown back (180о).

2. Conclusion
A. Most of the α- particles were continued their straight path that means
most of the space inside the atom is empty.
B. The centre of an atom has a positively charged body called nucleus
which repel positively charged α - particles and thus explained the
scattering phenomenon.
C. Almost all mass of an atom is concentrated in its nucleus.
D. The size and volume of the nucleus is very small as compared to the
total size and volume of atom.

3. Planck’s Quantum Theory
According to this theory atoms or molecules could emit or absorb energy
only in discrete quantities (small packets) and not in any arbitrary amount.
Planck gave the name quantum to the smallest quantity of energy that can
be emitted in the form of E.M. radiation.
Quantum of light is known as PHOTON

4. Planck’s Quantum Theory
The energy of Quantum 𝖺 ν of wave E
=hν
or
The energy of a photon is proportional to its frequency and is given by
E =hν where h =6.626 x 10-34Jsec,

5. Planck’s Quantum Theory
A body can emit or absorb energy only in terms of the integral multiples of
quantum, i.e.
E=n.hν, where n =1, 2, 3, .......
i.e. a body can emit or absorb energy as hν, 2hν ...... but it can not emit or
absorb energy in fractional values of hν such as 1.5 hν, 2.5 hν.
The energy absorbed/released by a substance
=n.hν
=n.hc/λ
n =number of Quantum

6. Planck’s Quantum Theory
The energy of a single quantum is really small thus a new unit is used (eV)
1eV =1.6x10-19 C.V =1.6x10-19J
⇒ 1eV =1.6 x 10-19J

7. Planck’s Quantum Theory
E(ev) =1240/λ(nm) or 12400/λ(Å)

8. Photoelectric Effect
Emission of electrons from a metal surface when exposed to light radiations
of appropriate wavelength is called photoelectric effect. The emitted
electrons are called photoelectrons.

9. Photoelectric Effect
Emission of electrons from a metal surface when exposed to light radiations
of appropriate wavelength is called photoelectric effect. The emitted
electrons are called photoelectrons.

10. Intensity
I =energy of Radiation/Time x surface area
I =nhν/t.A
A ⟶ Fixed
ν ⟶ Fixed
High intensity =More number of protons

11. Threshold Energy
Work function or threshold energy may be defined as the minimum amount
of energy required to eject electrons from a metal surface.
According to Einstein,
Maximum kinetic energy of the ejected electron=absorbed energy-work
function
where, ν0, and λ0, are threshold frequency and threshold wavelength
respectively

12. Bohr’s Atomic Model
Applicable only for mono electric system H, He+, Li2+, Be3+
● The atom has a central massive core nucleus where all the
protons and neutrons are present. The size of the nucleus is' cry
small.
● 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.

13. Bohr’s Atomic Model
An electron can move only in those permissive orbits in which the
angular momentum (mvr) of the electron is an integral multiple of
h/2π Thus,
where, m =mass of the electron, r =radius of the electronic orbit,
v = velocity of the electron in its orbit.

14. rn =0.529 n2 Å
Z
v =2.188 x 106 x Z m/sec
n
Bohr’s Atomic Model

15. Thus, TE of e- =KE +PE
n
=E =-13.6
Z2
n2
eV/atom
n
=E =-21.8 x 10-19
Z2
n2
J /atom

16. KE =-TE
&
PE =2TE
or TE =-KE =PE/2

17. For Hydrogen,
n
E =-13.6
Z2
n2
eV/atom
E1 =-13.6eV
E2 =-3.4eV
E3 =-1.51eV
E4=-0.85eV

18. (En) any atom =(En)H x Z2

19. n =4 (3rd excited state)
n =3 (2nd excited state)
n =2 (1st excited state)
n =1(ground state)

20. Wavelength of Photon
Eph =Eho - Elo RH =Rydberg’s constant
=109677 cm-1

21. Wavelength of Photon
=912Å
1
RH

22. Ionization of monoelectric species.
The minimum energy required to remove the ground state e- out of the atom
nlo=
1⟶ nho =∞
Epn =Eho - Elo
=E∞- E1 =0-E1
=IE =-(E1)
For H,
IE=
13.6ev

24. ● When mono electric mono atomic H(g) is heated, all gas and state e-s
will get excited to higher state
● Thus, after some time, these es will be to come back to ground state IN
EVERY POSSIBLE WAY
● Because of this, so many photons of diff. λ will be emitted.
Line spectrum of Hydrogen

25. Electron transitions for the Hydrogen atom

26. 1. It could not explain the line spectra of atoms containing more than
one electron.
2. This theory could not explain the presence of multiple spectral lines.
3. 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.
4. This theory could not explain uncertainty principle.
Failure of Bohr Model

27. It suggests wave-like nature of particle.
In 1924, de Broglie proposed that an electron, like light, behaves both as
material particle and as a wave.
The electrons, protons and even atoms, when in motion, possess wave
properties.
This proposal gave a new theory, known as wave mechanical theory of
matter.
De Broglie Hypothesis

30. If a charged particle present at rest is accelerated by V volts, then,
λ =
Inc. in KE =q.v
=(KE)particle =0 +qv
=qv
h
√2m q.v
q= charge particle (c)
only mag.
V=voltage (v)

32. Heisenberg’s Uncertainty Principle
For a MOVING MICROSCOPIC particle, it is impossible to determine exact
value of both potion & momentum SIMULTANEOUSLY.
While treating e- as a wave it is not possible to ascertain simultaneously the
exact position and velocity of the e- more precisely at a given instant

33. Δx =error in position (m)
Δp =error in momentum (kg m/s)
Δx. Δp ≥h/4π
⇒ (Δx)min .(Δp)min =h/4π

35. Wave mechanical model
Quantum mechanical model
- Proposed by heisenberg and schrodinger.
- Exact pos. of e- can’t be determined inside the atom also
- Only the probability of finding the e- can be determined.

36. Principal quantum number (n):
- It was proposed by Bohr and denoted by ‘n’
- It determines the average distance between electron and
nucleus,
- It determine the energy of the electron in an orbit where
electron is present.
- It gives the information of orbit K, L, M, N,....
- The value of energy increases with the increasing value of n.
- It represents the major energy shell from which the electron
belongs.

37. Azimuthal quantum number or angular quantum
number (l):
- It was proposed by sommerfield and denoted by ‘l’.
- It determines the number of subshells or sublevels to which
the electron belongs.
- The value of l is integral values upto (n - 1), starting from zero
where ‘n’ is the number of principle shell.
- It tells about the shape of subshells.

38. Magnetic quantum number (m) :

39. Magnetic quantum number (m) :
- It was proposed by linde and denoted by ‘m’
- It gives the number of permitted orientation of subshells.
- The value of m varies from -l to +l through zero.

40. n =
1
l = 0 m = 0
1 s
n =
2
l = 0 m = 0
l = 1 m = -1, 0, 1
n =
3
l = 0
l = 1
m = 0
m = -1 , 0, 1
l = 2 m = -2 , -1, 0,
1, 2

41. Spin quantum number (s):
It was proposed by goldschmidt and uhlenbeck and denoted by the
symbol of ‘s’
The value of ‘s’ is +½or - ½,which is signified as the spin or rotation
or direction of electron on it’s axis during the movement.

42. The spin may be clockwise or anticlockwise.
It represents the value of spin angular momentum is equal to
h/2π √s(s+1).
Maximum spin of an atom =½x number of unpaired electron.

43. Nodes
Radial
Or spherical
Nodes
n - l - 1
Angular nodes
or nodal planes
or cones
(l)

44. Nodes
Radial
Or spherical
Nodes
n - l - 1
Angular nodes
or nodal planes
or cones
(l)

45. Atomic Structure
- Radial modes =n - l - 1
- Angular nodes =l
- Total nodes =n - 1

46. Rules for R2(r)v/sr graph
- At origin, nucleus is present.
- For s orbitals, the curve starts from a max value
- For p, d, f orbitals, the curve starts from origin.
- The pt. Where the curve cuts x - axis represents radial nodes.

47. Rules for 4πr2 R2(r)dr v/sgraph
- Nucleus is present at origin
- For every orbital , the curve starts from origin
- ∵ if r =o P(r) =O
- All curves are in 1st quadrant only [for all orbitals]
- Each touch on x - axis is a radial node

48. Rules to write e.conf.
1. Aufbau’s Rule
For mono electronic system,
Energy is dependent on value of n only
1s <2s =2p <3s =3p =3d <4s =4p =4d =4f

49. Rules to write e.conf
1. Aufbau’s Rule
for multi electronic system,
Energy is dependent on value of n+l
- Subshell having small (n+l) has lesser energy
- If for 2 subshells, n+l is same
Subshell with smaller n is of less energy
⇒ 1S α 2S <2p <3S <3p <4S <3d <

50. Periodic Classification

51. Hund’s rule of maximum multiplicity
- The subshells are first filled with the e-s of some spin and then
the pairing occurs.
- maximum multiplicity conf is the most stable.

52. Exchange energy
In a subshell with degenerate orbitals, the e-s having same spin can
exchange their pos,
⇒ energy is released ⇒ exchange energy
Since energy is released ⇒ stability inc.

53. Pauli’s exclusion principle
No two e-s in an atom can have same value of all the 4 Q.No.
or
In an- orbital, e-s must be present with opp. spin

54. Periodic Classification
Sc Ti V G Mn Fe Co Ni Cu
21 22 23 24 25 26 27 28 29
4S23d1 4S23d2 4S23d3 4S23d5 4S23d6 4S23d7 4S23d8

55. d - Block exceptions

56. Periodic Classification
21 22 23 24 25 26 27 28 29
Sc Ti V Cr Mn Fe Co Ni Cu
Y Zr Nb Mo Tc Ru Rh Pd Ag
Hf Ta W Re Os Ir Pt Au

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B.
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58. Configuration of Ions

59. A.
B.
C.
D.

60. Periodic Classification
Magnetic Moment:
μ =√n(n +2) B.M
Measure of paramagnetic Nature
n = 1 μ = 1.73 BM
n = 2 μ = 2 BM
n = 3 μ = 3 BM

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