1) Experiments with cathode ray tubes led to the discovery of the electron as a negatively charged fundamental particle. 2) Further experiments showed that atoms are mostly empty space and contain a small, dense nucleus made up of protons and neutrons, around which electrons orbit. 3) The photoelectric effect showed that light behaves as a particle (photon) rather than just a wave, transferring its energy in discrete quantized amounts to electrons and ejecting them from metal surfaces.

1. Structure of Atom
1
GRADE 11

2. DISCOVERY OF AN ELECTRON
 An electron was discovered by cathode ray discharge tubes experiment.
 A cathode ray tube is made of glass containing two thin pieces of metal
called electrodes, sealed in it. The electrical discharge through the gases
could be observed only at very low pressures and at very high voltages .
The pressure of different gases could be adjusted by evacuation. When
sufficiently high voltage is applied across the electrodes , the current
starts to flow through a stream of particles moving in the tube from the
negative electrode to the positive one . These rays were called the
cathode rays or cathode ray particles.
 The flow of current from cathode to anode was further checked by
making a hole in the anode and coasting the tube behind anode with
phosphorescent material called zinc sulphide coating , a bright spot on
the coating is developed .
2

3. Results of the experiment
 The cathode rays start from cathode and move towards the
anode.
 These rays were not visible but their behaviour could be
observed with a certain kind of material called FLOUROSCENT
OR PHOSPHORESCENT MATERIALS.
 In the absence of electrical or magnetic field these rays travel in
straight lines.
 In the presence of electrical or magnetic field , the behaviour of
cathode rays are similar to that expected from negatively
charged particles suggesting that the cathode rays consist of
negatively charged particles called electrons.
 The characteristics of cathode rays do not depend upon the
material of electrodes nature of the gas present in the tube .
3

4. CHARGE TO MASS RATIO OF
ELECTRON
 The measure of mass ratio of an electrical charge (e) to the mass
of an electron (me) by using the cathode rays discharge tube and
applying electrical and magnetic field perpendicular to each
other as well as to the path of electrons .
 The amount of deviation of the particles from their path in the
presence of electrical and magnetic field depend upon:
1) The magnitude of the negative charge on the particle, greater
the magnitude of the charge on the particle , greater is the
interaction with the electric and magnetic field and thus
greater is the deflection.
2) The mass of the particle: lighter the particle , greater the
deflection and vice versa.
3) The strength of the electrical and magnetic field-the deflection
of electrons from its original path increases with the increase
in the voltage across the path electrodes or the strength of the
magnetic field.
4

5.  When only electric field is applied, the
electrons deviate from their path and hit the
cathode ray tube at point A.
 When only magnetic field is applied electron
strikes the cathode ray tube at point C.
 When electrons deviate from their path,
then both electrical and magnetic field is
applied , it is possible to bring them back
and electrons go and hit the screen at point
B.
 = 1.758820 x 1011 C kg-1
5

6. Discovery of protons and neutrons
 Electrical discharge carried out in the modified cathode rays
tube led to the discovery of particles carrying positive charge
also known as canal rays.
 The characteristics of these rays are:
⇒ Unlike cathode rays, the positively charges particles depend
upon the nature of gas present in the cathode ray tube.These
gases are simple positively charged ions.
⇒ The charge to mass ratio of the particles is found to depend on
the gas from which they originate.
⇒ Some of the +vely charged particles carry a multiple of the
fundamental unit of electrical charge.
⇒ The behaviour of these particles in the magnetic or electrical
field is opposite to that observed for electron or cathode rays.
6

7.  The smallest and lightest positive ion was
obtained from hydrogen and was called the
proton.
 Discovery of neutrons: Chadwick felt that by
bombarding a thin sheet of beryllium by α-
particles. When electrically neutral particles
having a mass slightly greater than that of the
protons was emitted. He named these neutral
particles as neutrons. Thus this discovery Was a
very important discovery in the history of
chemistry.
7

8. Thomsons model of an atom
 According to Thomson, atom was in a
spherical in shape which had positive
charged particle sand negative charged
particles equally distribution and hence it
was electrically neutral.
 Its observation could be called as a plum
pudding model or a watermelon.
8

9. Rutherford’s model of atom
 Famous experiment of Rutherford was the α-
particle scattering experiment.
 A stream of high energy α particles from a
radioactive source was directed at a thin foil of a
gold metal. The thin foil had a circular fluorescent
zinc sulphide screen around it .Whenever α-
particles struck the screen, a tiny flash of light
was produced at the point.
9

10.  The results of this experiment were unexpected.
 (1) most of the α-particles passed through the gold
foil undeflected.
 (2)a small fraction of α-particles was deflected by
small angles
 (3) a very few α-particles (-1 in 20,000) bounced
back , that is were deflected by nearly 180 degree.
Observations:
 Most of the space in the atom is empty as most of the
α-particles passed through the foil undeflected.
 A few +vely charged α-particles were deflected.
 The deflection must be due to enormous repulsive
force showing the positive charge in the atom.
 The positive charge has to be concentrated in a very
small volume that repelled and deflected the
positively charged α-particles . 10

11.  Conclusions:
 The positive charge and most of the massof the
atom was densely concentrated in extremely
small region. This concentrated region was called
nucleus.
 The nucleus was surrounded by electrons moving
in a very high speed in circular paths called
orbits.
 Electrons and the nucleus are held together by
the electrostatic forces of attraction.
11

12. Atomic number and mass number
 Atomic number(Z)=number of protons
present in the nucleus = number of
electrons in the neutral atom.
 Electrons and protons together ina nucleus
are / were called nucleons.
 Mass number (A)= number of protons
(z)=number of neutrons(n)
12

13. Isobars and isotopes
 Isobars are elements having the same mass
number but different atomic number.
 Whereas isotopes are elements having same
atomic number but a different mass number.
 Hydrogen has 3 isotopes:
protium , deuterium and tritium.
Chemical properties of atoms are controlled by
the number of protons in the nucleus therefore
they show similar chemical properties and
similar chemical behaviour
13

14. Draw backs of Rutherford’s model
of atom.
 It could not explain the gravitational force in
nature.
 It could not explain planetary motion under the
influence of gravity.
 It could not explain Maxwell's electromagnetic
radiation property.
 It could not explain quantum mechanics as a
whole.
14

15. The Electromagnetic Spectrum
 Visible light is a small portion of
the electromagnetic radiation
spectrum detected by our eyes.
 Electromagnetic radiation
includes radio waves,
microwaves and X-rays.
 Described as a wave traveling
through space.
 There are two components to
electromagnetic radiation, an
electric field and magnetic field.
15

16. The Wave Nature of Light
 Wavelength, , is the
distance between two
corresponding points
on a wave.
 Amplitude is the size or
“height” of a wave.
 Frequency, , is the
number of cycles of the
wave passing a given
point per second,
usually expressed in Hz.

17. Planck’s Quantum Theory
 Higher T = shorter λ (higher E) maximum.
 Couldn’t explain with classical physics
17

18. Electromagnetic Radiation
 In 1900 Max Planck studied black body
radiation and realized that to explain the
energy spectrum he had to assume that:
1. energy is quantized
2. light has particle character
 Planck’s equation is
18
sJ10x6.626constantsPlanck’h
hc
or EhE
34-





19. The Wave Nature of Light
 The fourth variable of light is velocity.
 All light has the same speed in a vacuum.
 c = 2.99792458 x 108 m/s
 The product of the frequency and wavelength is the
speed of light.
 Frequency is inversely proportional to wavelength.
19
c  

20. The Wave Nature of Light
 Electromagnetic radiation can be categorized in terms of
wavelength or frequency.
 Visible light is a small portion of the entire electromagnetic
spectrum.
20

21. The Particulate Nature of Light
 Photoelectric effect: light striking a metal
surface generates photoelectrons.
 The light’s energy is transferred to electrons in
metal.
 With sufficient energy, electrons “break free” of the
metal.
 Electrons given more energy move faster (have
higher kinetic energy) when they leave the metal.
21

22. The Particulate Nature of Light
 Photoelectric effect is
used in photocathodes.
 Light strikes the
cathode at frequency .
Electrons are ejected if
 exceeds the threshold
value 0.
 Electrons are collected
at the anode.
 Current flow is used to
monitor light intensity.

23. Example Problem 6.1
 Neon lights emit an orange-red colored
glow. This light has a wavelength of 670 nm.
What is the frequency of this light?
 We know that c=ʋλ
Therefore c/λ=ʋ and solve.
23

24. The Particulate Nature of Light
 The photoelectric effect is not explained using a wave
description but is explained by modeling light as a particle.
 Wave-particle duality - depending on the situation, light is best
described as a wave or a particle.
 Light is best described as a particle when light is imparting
energy to another object.
 Particles of light are called photons.
 Neither waves nor particles provide an accurate description
of all the properties of light. Use the model that best
describes the properties being examined.
24

25. Photoelectric Experiments
a) For > 0, the number of electrons
emitted is independent of frequency.
Value of 0 depends on metal used.
b) As light intensity increases, the
number of photoelectrons increases.
c) As frequency increases, kinetic energy
of emitted electrons increases linearly.
d) The kinetic energy of emitted electrons
is independent of light intensity.

26. The Particulate Nature of Light
 The energy of a photon (E) is proportional
to the frequency ().
 and is inversely proportional to the wavelength
().
 h = Planck’s constant = 6.626 x 10-34 J s
26
E  h 
hc


27. The Particulate Nature of Light
 Binding Energy - energy holding an electron to a
metal.
 Threshold frequency, o - minimum frequency of light
needed to emit an electron.
 For frequencies below the threshold frequency, no
electrons are emitted.
 For frequencies above the threshold frequency, extra
energy is imparted to the electrons as kinetic energy.
 Ephoton = Binding E + Kinetic E
 This explains the photoelectric effect.
27

28. Example Problem 6.2
 The laser in a standard laser printer emits
light with a wavelength of 780.0 nm. What is
the energy of a photon of this light?
28

29. The Wave Nature of Light
 Refraction is the bending of light when it passes
from one medium to another of different density.
 Speed of light changes.
 Light bends at an angle depending on its wavelength.
 Light separates into its component colors.
29

30. Example Problem 6.3
 In a photoelectric experiment, ultraviolet
light with a wavelength of 337 nm was
directed at the surface of a piece of
potassium metal. The kinetic energy of the
ejected electrons was measured as 2.30 x
10-19 J. What is the electron binding energy
for potassium?
30

31. Atomic Spectra
 Atomic Spectra: the particular pattern of
wavelengths absorbed and emitted by an element.
 Wavelengths are well separated or discrete.
 Wavelengths vary from one element to the next.
 Atoms can only exist in a few states with very
specific energies.
 When light is emitted, the atom goes from a higher
energy state to a lower energy state.
31

32. Atomic Spectra
 Electrical current dissociates molecular
H2 into excited atoms which emit light
with 4 wavelengths.

33. Atomic Spectra and the Bohr Atom
 Every element has a unique spectrum.
 Thus we can use spectra to identify
elements.
 This can be done in the lab, stars, fireworks, etc.
33

34. Example Problem 6.4
 When a hydrogen atom undergoes a transition from E3 to
E1, it emits a photon with  = 102.6 nm. Similarly, if the
atom undergoes a transition from E3 to E2, it emits a
photon with  = 656.3 nm. Find the wavelength of light
emitted by an atom making a transition from E2 to E1.
34

35. The Bohr Atom
 Bohr model - electrons orbit the
nucleus in stable orbits.
Although not a completely
accurate model, it can be used to
explain absorption and emission.
 Electrons move from low
energy to higher energy
orbits by absorbing energy.
 Electrons move from high
energy to lower energy orbits
by emitting energy.
 Lower energy orbits are
closer to the nucleus due to
electrostatics.
35

36. The Bohr Atom
 Ground state: the lowest state
(orbital)where the electrons are initially
present is called the ground state.
 Excited state: the state where the electrons
on gaining energy I subjected to go to a
higher energy level is called the excited
state.
Atoms return to the ground state by emitting
energy as light.
36

37. Atomic Spectra and the Bohr Atom
 The Rydberg
equation is an
empirical equation
that relates the
wavelengths of the
lines in the
hydrogen
spectrum.
37
hydrogenofspectrumemission
in thelevelsenergytheof
numberstherefer tosn’
nn
m101.097R
constantRydbergtheisR
n
1
n
1
R
1
21
1-7
2
2
2
1











38. Atomic Spectra and the Bohr Atom
Example 4-8. What is the
wavelength of light
emitted when the
hydrogen atom’s energy
changes from n = 4 to n =
2?
38
 
 
2 1
2 2
1 2
7 -1
2 2
7 -1
7 -1
7 -1
6 -1
n 4 and n 2
1 1 1
R
n n
1 1 1
1.097 10 m
2 4
1 1 1
1.097 10 m
4 16
1
1.097 10 m 0.250 0.0625
1
1.097 10 m 0.1875
1
2.057 10 m






 
 
  
 
 
   
 
 
   
 
  
 
 

39. Atomic Spectra and the Bohr Atom
 In 1913 Neils Bohr incorporated Planck’s
quantum theory into the hydrogen
spectrum explanation.
 Here are the postulates of Bohr’s theory.
39
1. Atom has a number of definite and discrete
energy levels (orbits) in which an electron
may exist without emitting or absorbing
electromagnetic radiation.
As the orbital radius increases so does the energy
1<2<3<4<5......

40. Atomic Spectra and the Bohr Atom
3. An electron moves in a circular orbit about the
nucleus and it motion is governed by the
ordinary laws of mechanics and electrostatics,
with the restriction that the angular
momentum of the electron is quantized (can
only have certain discrete values).
angular momentum = mvr = nh/2
h = Planck’s constant n = 1,2,3,4,...(energy levels)
v = velocity of electron m = mass of electron
r = radius of orbit
40

41. Atomic Spectra and the Bohr Atom
 Light of a characteristic wavelength (and
frequency) is emitted when electrons move from
higher E (orbit, n = 4) to lower E (orbit, n = 1).
 This is the origin of emission spectra.
 Light of a characteristic wavelength (and
frequency) is absorbed when electrons jump from
lower E (orbit, n = 2) to higher E (orbit, n= 4)
 This is the origin of absorption spectra.
41

42. 42

43. Atomic Spectra and the Bohr Atom
 Bohr’s theory correctly explains the H
emission spectrum.
 The theory fails for all other elements
because it is not an adequate theory.
43

44. The Wave Nature of the Electron
In 1925 Louis de Broglie published his Ph.D.
dissertation.
 A crucial element of his dissertation is that electrons
have wave-like properties.
 The electron wavelengths are described by the de
Broglie relationship.
44
particleofvelocityv
particleofmassm
constantsPlanck’h
mv
h





45. The Wave Nature of the Electron
Example 4.9. Determine the wavelength, in m, of an electron,
with mass 9.11 x 10-31 kg, having a velocity of 5.65 x 107 m/s.
 Remember Planck’s constant is 6.626 x 10-34 Js which is also equal to
6.626 x 10-34 kg m2/s.
45
 
  
m1029.1
m/s1065.5kg109.11
ss/mkg10626.6
mv
h
11
731-
2234











46. The Quantum Mechanical
Picture of the Atom
 Werner Heisenberg in 1927 developed the
concept of the Uncertainty Principle.
 It is impossible to determine simultaneously
both the position and momentum of an
electron (or any other small particle).
 Detecting an electron requires the use of
electromagnetic radiation which displaces the
electron!
 Electron microscopes use this phenomenon
46

47. The Quantum Mechanical Model of the Atom
 Quantum mechanical model replaced the Bohr
model of the atom.
 Bohr model depicted electrons as particles in circular
orbits of fixed radius.
 Quantum mechanical model depicts electrons as waves
spread through a region of space (delocalized) called an
orbital.
 The energy of the orbitals is quantized like the Bohr
model.
47

48. The Quantum Mechanical Model of the Atom
 Diffraction of electrons shown in 1927.
 Electrons exhibit wave-like behavior.
 Wave behavior described using a wave
function, the Schrödinger equation.
 H is an operator, E is energy and  is the wave
function.
48
H  E

49. Potential Energy and Orbitals
 Total energy for electrons includes potential and
kinetic energies.
 Potential energy more important in describing atomic
structure; associated with coulombic attraction
between positive nucleus and negative electrons.
 Multiple solutions exist for the wave function for
any given potential interaction.
 n is the index that labels the different solutions.
49

50.  On solving the Schrödinger equation we get
three set of number which are called the
quantum numbers which give us the location
space and the 3-D spatial orientation of the
orbital.
 The probability density gives us the
probability of an electron in the region that is
where the electrons in most probable to be.
50

51. Potential Energy and Orbitals
 n can be written in terms of two
components.
 Radial component, depends on the distance
from the nucleus.
 Angular component, depends on the direction
or orientation of electron with respect to the
nucleus.
51

52. Potential Energy and Orbitals
 The wave function may have positive and
negative signs in different regions.
 Square of the wave function, 2, is always positive
and gives probability of finding an electron at any
particular point.
 Each solution of the wave function defines an
orbital.
 Each solution labeled by a letter and number
combination: 1s, 2s, 2p, 3s, 3p, 3d, etc.
 An orbital in quantum mechanical terms is actually
a region of space rather than a particular point.
52

53. Quantum Numbers
 Quantum numbers - solutions to the
functions used to solve the wave equation.
 Quantum numbers used to name atomic
orbitals.
 Vibrating string fixed at both ends can be used
to illustrate a function of the wave equation.
53

54. Quantum Numbers
 When solving the Schrödinger equation,
three quantum numbers are used.
 Principal quantum number, n (n = 1, 2, 3, 4, 5,
…)
 Secondary quantum number, l
 Magnetic quantum number, ml
54

55. Quantum Numbers
 The principal quantum number, n, defines the
shell in which a particular orbital is found.
 n must be a positive integer
 n = 1 is the first shell, n = 2 is the second shell, etc.
 Each shell has different energies.
55

56. Quantum Numbers
 The secondary quantum number, l, indexes energy
differences between orbitals in the same shell in an atom.
 l has integral values from 0 to n-1.
 l specifies subshell
 Each shell contains as many l values as its value of n.
56

57. Quantum Numbers
 The energies of orbitals are specified
completely using only the n and l quantum
numbers.
 In magnetic fields, some emission lines split
into three, five, or seven components.
 A third quantum number describes splitting.
57

58. Quantum Numbers
 The third quantum number is the magnetic
quantum number, ml.
 ml has integer values.
 ml may be either positive or negative.
 ml’s absolute value must be less than or equal to l.
 For l = 1, ml = -1, 0, +1
58

59. Quantum Numbers
 Note the relationship between number
of orbitals within s, p, d, and f and ml.

60. Electron spin-s
 The 3 quantum numbers labeling an atomic
orbital can be equally used to well define its
energy, shape and its orientation.
 A new quantum number was introduced called
the electron spin quantum numbers-ms .
 It had its own axis,, these had angular momentum
and two orientation or spins that is ½ and -½
and were normally represented by two arrows ↑
and ↓showing that these have opposite spins.
60

61. Summing up
 n defines the shell ,determines the size of the orbital and also to a large extent
the energy of the orbital.
 There are n subshells in the nth shell. l identifies the sub shell and determines
the shape of the orbital there are (2l+1)orbitals of each type in a subshells
that is one s orbital (l=0) three p orbitals(l=1) and 5 d orbitals(l= 2) per
subshells. To some extent l also determines the energy of the orbital in a multi
electron atom.
designates the orientation of the electron of the orbital . For a given value of l
or has (2l+1) values , the same as the number of orbitals per subshells. it
means that the number of orbitals is equal to the number of ways in which they
are oriented
refers to orientation of the spin of the electron 61

62. Shapes of atomic orbitals
 According to max born the square of the wave
function at a point gives the probability density
of the electron at that point.
 For 1s orbital the probability density is maximum
at the nucleus and its decreases sharply as we
move away from it
 For a 2s orbital the probability density first
decreases sharply to zero and then again starts
increases.
 After reaching a small maxima it decreases again
and approaches zero as the value of r increases
further. 62

63.  The region where this probability of finding an
electron reduces to zero is called nodal
surfaces or simply nodes.
 In general, for an ns orbital , an ns orbital has
(n-1)nodes, that is number of nodes increases
with the increase of quantum number
n.therefore for 2s it will be 1 and for 3s it will
be 2.
 Boundary surface diagrams of constant
probability density for different orbitals give a
fairly good representation of the shapes of the
orbitals . In this representation, a boundary
surface or a contour surface is drawn in space
for an orbital on which the value of probability
density is constant.
63

64.  In the boundary surface diagrams the nucleus is
taken to be at the origin or rather it is . Here, ,
diagrams are not spherical like the s-orbital. Here
the p-orbital consists of two section s called
lobes. That are on either sides of the plane where
the two lobes touch each other. The size, shape
and energy of the three orbitals are identical .
 Since the lobes are considered along the x, y and
the z axis they are designated as the above:-
 It should be understood that there is no relation
between first magnetic quantum number and x, y,
z directions 64

65. 65
Boundary surface diagrams for 2p orbitals

66. d- orbitals
 The 5 d-orbitals are designated as:-
The shapes of the first four d-orbitals are similar to
each other , where as the fifth one is different form
others, but all 5 have 3d- orbitals and are equivalent
in energy.
The d- orbitals for which n is greater than 3 also have
shapes similar to 3d orbital , but differ in energy.
When two nodal planes pass through the same origin
and bisecting the xy-pane and z-plane these nodes
are called angular nodes.
66

67.  Angular nodes are denoted by ’l’.
There are one angular node for the p-orbitals and 2
angular nodes for d-orbitals
The total number of nodes are given by (n-1)i.e the
sum of l angular nodes are (n-l-1) radial nodes.
Energies of orbitals:-
Energy increasing order in the orbitals is given as
follows:-
1s<2s=2p<3s=3p=3d<4s=4p=4d=4f<
As 2p and 2s orbitals are different, an electron has
the same energy as it is present in the 2p or the
2s orbital.
67

68.  The orbitals which have the same energy are
called degenerate. The 1s orbital in a hydrogen
atom corresponds to the most stable condition
and is called the ground state. And an electron
residing in this orbital is strongly held by the
nucleus
 The electrons residing in the 2s , 2p or higher
orbitals in the hydrogen atom are said to be in
the excited state.
 The attractive interactions of an electron
increases with the increase of the positive charge
(Ze) on the nucleus.
68

69.  Due to the presence of electrons in the inner
shells, the electron in the outer shells will not
experience full positive charge of the nucleus.
 The effect will be lowered due to the partial
screening of positive charge on the nucleus by the
inner shell electrons. This is known as shielding
of outer electrons from the nucleus by the
inner shell electrons and the net positive charge
experienced by the outer electrons is known as
effective nuclear charge.( )
 In other word the energy of interaction b/w the
nucleus and electron decreases with the increase
of atomic number Z. 69

70. Aufbau principle
 Aufbau principle deals with filling up of electrons.
The principle states:- In the ground state of the
atoms , the orbitals are filled in order of the
increasing energies.
 In other words, electrons first occupy the lowest
energy orbitals available tot hem and then enter
into higher energy orbital only after lower energy
orbital is filled.
 Order of increasing order of energies in the
orbital is as follows:-
1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,4f,5d,6p,7s……..70

71. 71
Order of filling
up of electrons

72. Pauli exclusion principle
 According to this principle :- no two electrons
in this atom have the same set of four
quantum numbers.
Or
It can also be stated otherwise as only two
electrons may exist in the same orbital and
these orbital must have opposite spins
The maximum number of electrons which can be
accommodated in the shell with the quantum
number n is according to the 2n2 rule. 72

73. Hund’s rule of maximum multiplicity
 This rule deals filling of electrons in the orbitals
belonging to the same subshells of equal energy
called degenerate orbitals.
 It states that pairing of electrons in the
orbitals belonging to the same sub shell (p,d,
or f) does not take place until each orbital of
that sub shell gets one electron that is singly
occupied.
 Some of the orbitals acquire extra stability due to
their symmetry. 73

74. Electronic configuration
 The distribution of electrons into orbitals of an
atom is called its electronic configuration.
 Electronic configuration can be represented in
two ways:-
(a)Normal notation and
(b) orbital diagram
As given in the textbook.
The electron in the completely filled electronic shell
with the highest principal quantum number are
called valence electrons.
74