MKVV SCHOOL

Structure of atom
Discovery of sub-atomic particles –

Electron – The cathode ray tube experiment majorly contributed in the discovery of electron. A
cathode ray tube is vacuum tube made up of glass that contains 2 thin metal electrodes. When a
certain amount of high voltage is applied across the electrodes the particles starts moving from
cathode to anode. This beam of particles was called as the cathode ray or cathode ray particles.

Summarised results of this experiment –

1. The cathode ray moves towards anode from cathode.

2. These ray travel in straight line when electrical or magnetic field is absent.
3. When electrical or magnetic field is present the behaviour of cathode ray is similar to that
expected from negatively charged particles. This suggests that cathode ray consists of
negatively charged particles called as “electrons”.

Proton – In cathode ray tube, due to electrical discharge “canal rays” were introduced which were
positively charged. The behaviour of these particles was opposite than the cathode ray. These
particles were named as “Proton”.

Neutron – These particles were discovered by James Chadwick (1932). Neutron is an electrically
neutral particle that have greater mass than the proton.

Atomic models –

1. Thomson atomic model –

In 1998, Sir J.J. Thomson proposed that atom is a spherical structure with having a radius of
10–10 m. According to this model the positive charge is uniformly distributed in the whole
sphere & negatively charged particles (electrons) are present in the scattered form, same as
the seeds present in watermelon, hence this model was also called as the “watermelon
model” or the “plum pudding” model.
electrons

Thomson’s model of atom

Positive
sphere

This model was only able to explain the overall neutral behaviour of the atom. But was not
consistent with the results after experiments.

2. Rutherford’s nuclear model of atom –

The Rutherford’s nuclear model of atom was based on an experiment called as “Rutherford’s
alpha particle experiment”. Rutherford shot a stream of high energy particles (α particles)
into a gold foil due to which following observations were noted –
a. Most of the α particles passed straight through the gold foil.
b. A small number of α particles was deflected by small angles.
c. 1 in 20,000 α particles returned back on their way.
 These observations lead to following conclusions –

i. Most of the space in atom is empty as the α particles passed through gold foil
undeflected.
ii. A few positively charged particles were deflected due to enormous repulsive force
shows that the positive charge is not spread throughout the atom.
iii. On the basis of Rutherford’s calculation the volume occupied by nucleus is negligibly
small & surrounded by electrons which orbit the nucleus.

Structure of atom according to above conclusions –

• The very small portion where atom is densely concentrated in centre is called as the
“Nucleus”.
• The nucleus is surrounded by the electrons that move around with high speed in circular
path that is called as “Orbit”.

Isotopes – The atoms with having same number of protons but different number of neutrons are
known as the “Isotopes”. The Isotopes shares almost same chemical properties but differs in the
physical properties.

• Isotopes of hydrogen –
1. Protium (Hydrogen)
2. Deuterium
3. Tritium

Electromagnetic radiations –

Electromagnetic radiation refers to the propagation of energy through space in the form of
oscillating electric and magnetic fields. It encompasses a broad spectrum of wavelengths and
frequencies, ranging from radio waves with long wavelengths and low frequencies to gamma rays
with short wavelengths and high frequencies.

characteristic of electromagnetic radiation include:

1. Wavelength and Frequency: Electromagnetic radiation exhibits varying wavelengths and

frequencies. Wavelength is the distance between successive peaks or troughs of a wave,
while frequency is the number of oscillations per unit of time.
2. Speed: Electromagnetic radiation travels at the speed of light, which is approximately
299,792,458 meters per second in a vacuum. This speed is constant for all electromagnetic
waves, regardless of their frequency or wavelength.
Wavelength is denoted by “λ”. The speed of wavelength can be denoted by “c” & the frequency of it
can be denoted by v.

Black body radiation –

Black body radiation refers to the electromagnetic radiation emitted by a perfect absorber and
emitter of energy, known as a black body. A black body is an idealized physical body that absorbs all
radiation incident upon it and emits radiation at all frequencies.

characteristics of black body radiation:

1. Continuous Spectrum: A black body emits radiation across all wavelengths. The spectrum of
this radiation depends only on the temperature of the body and not on its material
properties.
2. Intensity Distribution: The intensity of radiation emitted by a black body at different
wavelengths follows a specific distribution known as Planck's law, named after physicist Max
Planck, who first derived it in 1900. Planck's law describes how the intensity of radiation
varies with wavelength and temperature.
3. Peak Wavelength: The wavelength at which the intensity of radiation emitted by a black
body is maximum depends on its temperature. As temperature increases, the peak
wavelength shifts to shorter (bluer) wavelengths according to Wien's displacement law.
4. Temperature Dependence: The total power radiated by a black body per unit area (also
known as its radiative flux) is proportional to the fourth power of its temperature, according
to the Stefan-Boltzmann law.

Photoelectric effect –

The photoelectric effect refers to the phenomenon where electrons are ejected from a material
when it's exposed to electromagnetic radiation, typically light. It was first observed by Heinrich Hertz
in 1887, and its explanation by Albert Einstein in 1905 was a significant contribution to the
development of quantum mechanics.
Quantum number –

Quantum numbers are values used to describe the unique quantum state of an electron within an
atom. They define various properties of an electron such as its energy, angular momentum,
orientation in space, and spin. There are four quantum numbers:

1. Principal Quantum Number (n): Denoted by 'n', it determines the energy level of an
electron and its distance from the nucleus. The higher the value of 'n', the higher the energy
level and the farther the electron is from the nucleus.
2. Angular Momentum Quantum Number (l): Denoted by 'l', it determines the shape of the
orbital in which the electron resides. The values of 'l' range from 0 to (n-1) and represent
different sublevels:
o l = 0: s orbital (spherical)
o l = 1: p orbital (dumbbell-shaped)
o l = 2: d orbital (complex shapes)
o l = 3: f orbital (even more complex shapes), and so on.
3. Magnetic Quantum Number: Denoted by ml’, it specifies the orientation of the orbital in
space. The values of ml’ range from -l to +l, including zero. For example, for an l value of 2,
ml’ can be -2, -1, 0, 1, or 2, indicating the five different orientations in space for the d orbital.
4. Spin Quantum Number: Denoted by ms, it describes the intrinsic angular momentum or spin
of the electron. It has only two possible values: +1/2 (spin-up) or -1/2 (spin-down).

These quantum numbers provide a comprehensive description of the electron's state within an atom
and are crucial in understanding atomic structure and predicting electron behavior.

Introduction to orbitals –

Orbitals are regions in an atom where electrons are likely to be found. They are described by
mathematical equations called wave functions. There are several types of orbitals, each with its own
characteristic shape:

1. s Orbital: Spherical in shape, with a single lobe. An s orbital can hold a maximum of 2
electrons. There are different s orbitals corresponding to different energy levels (n values).
2. p Orbital: Consists of three dumbbell-shaped lobes oriented along the x, y, and z axes. There
are three p orbitals per energy level (labeled as px, py, and pz), and each can hold a
maximum of 2 electrons, totaling 6 electrons for the p orbitals in an energy level.
3. d Orbital: There are five d orbitals per energy level, each with a complex shape. These
shapes include cloverleaf-shaped lobes and dumbbell shapes with a ring around the middle.
Each d orbital can hold a maximum of 2 electrons, totaling 10 electrons for the d orbitals in
an energy level.
4. f Orbital: Even more complex in shape than d orbitals, with seven f orbitals per energy level.
The shapes of f orbitals are difficult to describe succinctly, but they have multiple nodes and
lobes. Each f orbital can hold a maximum of 2 electrons, totaling 14 electrons for the f
orbitals in an energy level.

These shapes describe the probability distribution of finding an electron in a particular region of
space around the nucleus. Each orbital type corresponds to a specific quantum number, which
determines its size, shape, and orientation within an atom.
Difference between orbit & orbital –

Orbit Orbital
1. It is a circular path where the electrons
1. It is a region around the nucleus of the
revolve around the nucleus.
atom where the probability of finding
2. It represents the planar motion of
electron is maximum.
electron.
2. It represents 3D motion of an electron.
3. Orbits are circular in shape.
3. Orbitals are different in shape.

Aufbau principle –

The Aufbau principle is a fundamental concept in chemistry that describes the order in which
electrons fill atomic orbitals in an atom. It's based on the idea that electrons occupy the lowest
energy orbitals available before filling higher energy ones. Here's how it works:

1. Lowest Energy Levels First: Electrons fill the lowest energy orbitals available within an atom
before moving to higher energy ones. The energy of an orbital generally increases with
increasing principal quantum number (n).

Electronic configuration
example of carbon
electron
 Pauli Exclusion Principle: No two electrons in an atom can
have the same set of four quantum numbers. This means that
each orbital can hold a maximum of two electrons, and these
electrons must have opposite spins (one "up" and one
"down").

Hund's Rule: When filling orbitals with the same energy (degenerate orbitals), electrons occupy
them singly with parallel spins before pairing up. This results in the maximum number of unpaired
electrons and minimizes electron-electron repulsion, stabilizing the atom.