2. Discovery of electron
Cathode-ray discharge tube experiment:
• Very low pressures and very high voltages required.
• Current starts flowing through a stream of particles moving in the tube from the
negative electrode (cathode) to the positive electrode (anode). These are called
cathode rays or cathode ray particles.
• Phosphorescent material (ZnS) showed bright spots when the cathode rays hit it.
3. RESULTS of Cathode-ray discharge tube experiment:
• Cathode ray start from cathode (-) and move towards the anode (+).
• Cathode ray is not visible. So, phosphorescent/fluorescent materials are
used.
• Travels straight (when no electrical or magnetic field is there).
• Behaves as negatively charged particles in the presence of electrical or
magnetic field.
• The characteristics of cathode rays (electrons) do not depend upon the
material of electrodes and the nature of the gas present in the cathode ray
tube.
Conclusion:
• Cathode ray consists of negatively charged electrons.
• Electrons are basic constituent of all the atoms.
5. Amount of deviation of the particles from their path depends upon:
Greater the magnitude of negative charge on the particle, greater is the
deflection.
lighter the particle, greater the deflection.
the deflection of electrons increases with the increase in the voltage or the
strength of magnetic field.
6. Charge on the Electron
• R. A. Millikan’s oil-drop experiment gave the charge on an electron (e).
7. Discovery of Protons
• Modified cathode ray tube: Discovery of canal rays carrying positively charged
particles.
Features of canal rays:
• Mass of positively charged particles depends upon the nature of gas.
• Charge to mass ratio of the particles depends on the gas from which these
originate.
• Some of the positively 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.
8. • The smallest and lightest positive ion was obtained from
hydrogen and was called proton.
9. Discovery of neutron
• Discovered by James Chadwick
• Bombarding a thin sheet of Be by a-particles.
• These are electrically neutral particles having mass lightly greater
than protons.
11. 1. Thomson Model of Atom
• J. J. Thomson, in 1898
• Atom possesses a spherical shape (radius approximately
10–10
m).
• The positive charge is uniformly distributed. The
electrons are embedded into it in such a manner as to
give the most stable electrostatic arrangement.
• Also called: plum pudding, raisin pudding or watermelon.
• Mass of the atom is assumed to be uniformly distributed
over the atom.
• Explained the overall neutrality of the atom
13. 2. Rutherford’s Nuclear Model of Atom
α–particle scattering experiment:
• bombard very thin gold foil with α–particles.
14. EXPECTATION:
• Particles would slow down and change directions only by a small
angles due to uniform distribution of mass.
OBSERVATION:
1. Most of the α–particles passed through the gold foil undeflected.
2. A small fraction of the α–particles was deflected by small angles.
3. A very few α–particles (~1 in 20,000) were deflected by nearly 180°.
15. CONCLUSION:
•Most of the space in the atom is empty.
•Positive charge of the atom is not spread throughout
the atom. It is concentrated in a very small volume
•Volume occupied by the nucleus is negligibly small as
compared to the total volume of the atom.
The radius of the atom is about 10–10
m, while that of
nucleus is 10–15
m.
16. Nuclear model:
•The positive charge and most of the mass of the atom
was densely concentrated in extremely small region
(nucleus).
•Electrons move around the nucleus with a very high
speed in circular paths called orbits.
•Electrons and the nucleus are held together by
electrostatic forces of attraction.
17. Drawbacks of Rutherford’s model
• When a body (electron) is moving in an orbit, it undergoes acceleration even if it
is moving with a constant speed in an orbit because of changing direction.
Therefore, an electron in an orbit will emit radiation, and the orbit will thus
continue to shrink.
Thus, he could not explain the stability of an atom.
• If the electrons were stationary, electrostatic attraction between the dense
nucleus and the electrons would pull the electrons toward the nucleus.
• This model does not explain the distribution of the electrons around the
nucleus and the energies of these electrons
18. Atomic Number and Mass Number
• Nucleons: Protons and neutrons in a nucleus
19. Isobars and Isotopes
• Isobars: Atoms with same mass number but different atomic number.
• Isotopes: Atoms with same atomic number but different mass number.
20. THINKLINE
Q) Why do all the isotopes of a given element
show same chemical behaviour?
22. Isoelectronic species
• Species having the same number of electrons.
• Example: Na+
, Mg2+
, F-
, Ne
Think of another example!
24. Development leading to Bohr’s model
• Dual nature of electromagnetic radiations
• Atomic spectra
25. Features of electromagnetic waves
• Electrical and magnetic fields are mutually perpendicular to each other.
• Both are perpendicular to direction of wave.
• Do not need medium to travel.
• Different types that vary in frequency/wavelength
26. Relation between frequency and wavelength
• Unit of frequency and wavelength?
• Wave number: number of wavelengths per unit length. (m-1
)
33. Particle-nature of electromagnetic radiation
• Black-body radiation
• Photoelectric effect
• variation of heat capacity of solids as a function of temperature
• Line spectra of atoms
System can take energy only in discrete amounts. All possible
energies cannot be taken up or radiated.
34. Black body radiation
• Black body is an ideal body that emits and absorbs all frequencies. (C black)
• The radiation emitted by black body is called black body radiation.
• Hot objects emit electromagnetic radiations over a wide range of wavelengths.
• Intensities of radiations of different wavelengths emitted by hot body depend upon its
temperature.
35. Planck’s Quantum theory
• Atoms and molecules can emit (or absorb) energy only in discrete
quantities and not in a continuous manner.
• Quantum: The smallest quantity of energy that can be emitted (or
absorbed) as electromagnetic radiation.
E = hν
• Quanta: Total amount of energy emitted or absorbed.
E = nhν; n = integer
37. Photoelectric effect
• The phenomenon in which electrons (or electric current) were ejected when certain metals (K,
Rb, Cs etc.) were exposed to a beam of light.
Important features:
no time lag between striking of light & ejection of electrons from the metal.
number of electrons ejected is proportional to intensity (brightness) of light (more photons).
For each metal, there is a characteristic minimum frequency (ν0) (threshold frequency) below
which photoelectric effect is not observed.
At a frequency ν>ν0, the ejected electrons come out with certain kinetic energy.
39. Dual Nature of electromagnetic radiation
• Particle-like behaviour: when radiation interacts with matter.
• Wave like behaviour: when radiation propagates.
44. Spectra
• Atoms and molecules may absorb
energy and reach to a higher energy
state.
• When they return to their normal state
from the excited state, they emit
radiation.
• Line Spectra is unique to each element
(like fingerprint).
47. Line Spectrum of Hydrogen
Rydberg constant for hydrogen: 109677 cm–1
48. BOHR’S MODEL FOR HYDROGEN ATOM
POSTULATES:
1. The electron in the hydrogen atom move around the nucleus in a circular path of
fixed radius and energy, called orbits (stationary states/allowed energy states).
2. The energy of an electron in the orbit does not change with time. However,
electron can make transition by absorbing or emitting energy.
3. Frequency of radiation absorbed or emitted when transition occurs:
4. The electron can revolve in only those orbits whose angular momentum (mvr) is
an integral multiple of h/2π.
49. Radius of orbit (Rn)
Here, a0: first orbit = 52.9 pm
Z: atomic number
50. Energy of orbit (En)
Here, RH: Rydberg constant = 2.1810–18
J
• Energy of electron free from influence of nucleus, the energy is taken as ZERO.
• As electron approaches nucleus, its energy decreases (and thus, negative value).
53. Limitation of Bohr’s theory
• Could not explain spectra of atoms having more than one electron.
• Does not satisfactorily explain the distribution of electrons in an atom.
• Could not explain Zeeman effect (splitting of spectral lines in magnetic field).
• Could not explain Stark effect (splitting of spectral lines in electrical field).
• Could not explain doublet in H-atom spectrum.
• It could not explain the ability of atoms to form molecules by chemical bonds.
56. Dual Behaviour of Matter
• Matter, like radiation, should also exhibit dual behaviour i.e., both
particle and wavelike properties.
• Electrons should also have momentum as well as wavelength.
• de Broglie equation:
Every moving object has a wave character. The wavelengths of macroscopic
objects are so short (due of large masses) that their wave properties cannot
be detected.
The wavelengths of electrons and other subatomic particles (small mass) can
be detected experimentally.
61. Heisenberg’s Uncertainty Principle
• It is impossible to determine simultaneously, the exact position and
exact momentum (or velocity) of an electron.
62. Significance of Uncertainty Principle
• It rules out existence of definite paths or trajectories of electrons and
other similar particles.
• Is significant only for motion of microscopic objects and is negligible
for that of macroscopic objects.
66. Bohr’s model & dual nature of electron
Bohr model of the hydrogen atom ignores:
Dual behaviour of matter
Heisenberg uncertainty principle
68. Quantum numbers & Orbitals
• Orbital is the space around nucleus where the probability of finding an electron is
maximum.
• A set of 4 quantum numbers gives complete information about an electron in an
atom.
Principal quantum number (n)
Azimuthal quantum number (l)
Magnetic quantum number (ml)
Electron spin quantum number (ms)
69. Principal Quantum Number (n)
For nth
shell:
No. of subshell = n
No. of orbitals = n²
Maximum no. of electrons = 2n²
Q) How many atomic orbitals are there
in the third shell?
70. Azimuthal Quantum Number (l)
For a given subshell (l):
No. of orbital = 2l + 1
Maximum no. of electrons = 2(2l + 1)
72. Electron spin quantum number (ms)
Energy of orbital = (n + l)
If (n + l) are equal for two orbital, then the one with greater n has
higher energy.
76. Shape of orbitals
• Shape of orbital is given by total
probability density psi2
.
• For 1s orbital, the probability density is
maximum at the nucleus and decreases
sharply as we move away from it.
• For 2s orbital, the probability density
first decreases sharply to zero and then
again starts increasing. After reaching a
maxima, it again starts decreasing.
• NODE: The region where the probability
density function reduces to zero.
77. NODE
• For principal quantum number (n):
Radial nodes = n-l-1
Angular node = l
TOTAL NODES = n-1
78. s-orbital
• Spherical shape
• Probability of finding electron is same in all
directions at a distance from the nucleus.
79. p-orbital
• Dumb-bell shape
• Probability of finding electron is maximum in the
two lobes on opposite sides of nucleus.
• p-orbital has ONE NODAL PLANE passing
through nucleus.
80. d-orbital
• Clover-leaf shape(dz2
= doughnut shape)
• Probability of finding electron is same in all
directions at a distance from the nucleus.
• d-orbital has TWO NODAL PLANE passing
through nucleus.
81. Energy of orbitals
• Energy of orbital = (n + l)
• 1s = 1
• 2s = 2
• 2p = 3
• 3s = 3
• 3p = 4
• 3d = 5
• 4s = 4
• 4p = 5
• 4d = 6
• 5s = 5
If two orbitals have the same (n + l) value, then the orbital
with greater n value will have more energy.
Energy of the same orbital decrease with increase in
atomic number. {So, 1s of H > 1s of Li}
Q) Now, arrange the given orbitals in increasing order of their energies.
83. Ground state and excited state
• For H-atom, 1s is the ground state. However, all other orbitals are
excited state for H-atom.
84. RULES for filling electrons in atom
1. Aufbau principle
2. Pauli exclusion principle
3. Hund’s rule of maximum multiplicity
85. Aufbau principle
• In the ground state of the atoms,
the orbitals are filled in order of
their increasing energies.
• Exceptions are possible.
Q) Fill 11 electrons in the orbitals.
Q) Fill 21 electrons in the orbitals.
86. Pauli exclusion principle
• No two electrons in an atom can have the same set of four quantum numbers.
• Only two electrons may exist in the same orbital and these electrons must have
opposite spin.
Q) Fill 11 electrons in orbital diagram. (show the electrons as arrows)
87. Hund’s rule of maximum multiplicity
• Pairing of electrons in the orbitals belonging to the same subshell (p,
d or f) does not take place until each orbital belonging to that subshell
has got one electron each i.e., it is singly occupied.
88. Electronic configuration of an atom
• The distribution of electrons into orbitals of an atom is called its
electronic configuration.
• Use the 3 rules to write the electronic configuration.
• Another common way: [Nobel gas]sx
py
dz
Q) Write the electronic configuration of Ca (Z = 20).
Q) Write the electronic configuration of Zn (Z = 30).
Q) Write the electronic configuration of Cu (Z = 29).
89. Stability of half-filled and fully-filled
subshell
• Half-filled and fully-filled subshell are highly stable. Thus, wherever
possible, this arrangement will be preferred.
• Thus, Cu = [Ar]3d10
4s1
instead of [Ar]3d9
4s2
.
Q) Now, write the electronic configuration of Cr (Z = 24)
90. Reason for exceptions
Symmetry (leads to stability)
Exchange energy:
Electrons with the same spin in the degenerate orbitals of a subshell can
exchange their position. The energy released during this exchange is called
exchange energy. Half-filled and fully filled subshells show maximum number of
exchanges.
![Electronic configuration of an atom
• The distribution of electrons into orbitals of an atom is called its
electronic configuration.
• Use the 3 rules to write the electronic configuration.
• Another common way: [Nobel gas]sx
py
dz
Q) Write the electronic configuration of Ca (Z = 20).
Q) Write the electronic configuration of Zn (Z = 30).
Q) Write the electronic configuration of Cu (Z = 29).](https://image.slidesharecdn.com/structureofatomppt-250722103156-e99d36b3/85/Structure-of-atom-ppt_ClassXI_CBSE_NCERT-88-320.jpg)
![Stability of half-filled and fully-filled
subshell
• Half-filled and fully-filled subshell are highly stable. Thus, wherever
possible, this arrangement will be preferred.
• Thus, Cu = [Ar]3d10
4s1
instead of [Ar]3d9
4s2
.
Q) Now, write the electronic configuration of Cr (Z = 24)](https://image.slidesharecdn.com/structureofatomppt-250722103156-e99d36b3/85/Structure-of-atom-ppt_ClassXI_CBSE_NCERT-89-320.jpg)
