Structure of Atom: Complete Chapter Revision for Class 11 NEET by PW

The Structure of Atom chapter explains the composition of an atom and how its electrons are arranged around the nucleus. It begins with the discovery of subatomic particles and the development of atomic models before moving to electromagnetic radiation, quantum theory, Bohr’s model, and the quantum mechanical description of electrons.

PW Structure of Atom Class 11 notes help you revise important experiments, formulas, atomic models, quantum numbers, and electronic configuration rules. The concepts of hydrogen spectra, de Broglie waves, Heisenberg’s uncertainty principle, and orbital shapes are also covered for Class 11 Chemistry and NEET preparation.

Subatomic Particles In An Atom

An atom contains three fundamental subatomic particles:

The electron has a mass of approximately 9.1 x 10^-31 kg, while the proton has a mass of approximately 1.67 x 10^-27 kg. The charge of an electron is -1.6 x 10^-19 coulomb, and the proton has an equal positive charge.

Discovery Of The Electron

Cathode Ray Experiment

J. J. Thomson studied cathode rays using an evacuated discharge tube containing gas at very low pressure. When a high potential difference was applied between the cathode and anode, rays travelled from the cathode towards the anode. These were called cathode rays.

The important observations were:

• They travel in straight lines.

• They produce fluorescence on a suitable screen.

• They are deflected by electric and magnetic fields.

• They bend towards the positive plate, showing that they carry negative charge.

• They have mass and very high speed.

• Their properties are independent of the gas and electrode material.

• They have the same charge-to-mass ratio for different gases.

• They are present in all atoms.

These observations showed that atoms contain smaller charged particles. The particles present in cathode rays were identified as electrons.

Thomson determined the specific charge of the electron: e/m = 1.75 × 10¹¹ C kg⁻¹. He could not determine the separate values of electron charge and mass.

Millikan’s Oil-Drop Experiment

R. A. Millikan determined the charge of an electron by observing charged oil droplets between oppositely charged plates.

The charge of an electron is: e = −1.6 × 10⁻¹⁹ C

The experiment also established the quantisation of charge: q = ±ne

where n is an integer. Therefore, charge exists as an integral multiple of the elementary charge.

Discovery Of The Proton And Neutron

Anode Or Canal Rays

Goldstein studied positively charged rays using a perforated cathode. These rays travelled from the anode toward the cathode and passed through the holes in the cathode.

Anode rays:

• Travel in straight lines.

• Produce fluorescence.

• Are deflected by electric and magnetic fields.

• Contain positively charged particles.

• Have greater mass than electrons.

• Depend on the nature of the gas used.

The positively charged particle associated with these rays was identified as the proton.

Discovery Of The Neutron

James Chadwick discovered the neutron by bombarding beryllium with alpha particles. The new particle had no charge, was not deflected by electric or magnetic fields, and had a mass close to that of a proton. It was present in the nucleus and was named the neutron.

Development Of Atomic Models

Thomson’s Atomic Model

Thomson proposed that the atom is a positively charged sphere containing embedded electrons. The positive charge was uniformly distributed, and the atom as a whole was neutral.

The model is also called the plum-pudding or watermelon model.

The model correctly suggested that the atom is spherical and electrically neutral. However, it failed to explain the actual arrangement of charge and was rejected after Rutherford’s experiment.

Rutherford’s Nuclear Model

Rutherford directed alpha particles at a thin gold foil surrounded by a fluorescent screen.

His observations were:

1. Most alpha particles passed through without deflection.

2. Some particles were slightly deflected.

3. A very small number were deflected through large angles.

4. About one particle in 20,000 rebounded nearly backward.

Rutherford concluded that:

• Most of the atom is empty space.

• Positive charge is concentrated in a very small, dense nucleus.

• Electrons are present outside the nucleus.

• The nucleus is much smaller than the atom.

The approximate dimensions are:

Atomic radius ≈ 10⁻¹⁰ m

Nuclear radius ≈ 10⁻¹⁵ m

Rutherford’s model is also called the planetary or solar model.

Limitations of Rutherford’s Model

According to classical electromagnetic theory, a revolving electron should continuously lose energy by radiation and eventually fall into the nucleus. This would make the atom unstable.

Rutherford’s model also could not explain the arrangement of electrons, quantised energy levels or the discontinuous spectrum of hydrogen.

Electromagnetic Radiation and Quantum Theory

Electromagnetic waves consist of mutually perpendicular electric and magnetic fields. They do not require a material medium and travel through vacuum at the speed of light.

The basic relation is: c = νλ

where:

• c = speed of light

• ν = frequency

• λ = wavelength

Thus, frequency and wavelength are inversely proportional.

The electromagnetic spectrum includes:

Radio waves → Microwaves → Infrared → Visible light → Ultraviolet → X-rays → Gamma rays → Cosmic rays

Along this sequence, frequency increases while wavelength decreases.

Planck’s Quantum Theory

Max Planck proposed that energy is emitted or absorbed in small packets called quanta. The quantum of electromagnetic radiation is called a photon.

• Energy of one photon: E = hν

• Since ν = c/λ: E = hc/λ

• For n photons: E = nhν

Photoelectric Effect

The photoelectric effect is the emission of electrons from a metal surface when suitable electromagnetic radiation falls on it. The emitted electrons are called photoelectrons.

Important observations include:

• Photoelectric emission is immediate.

• The number of photoelectrons increases with light intensity.

• Every metal has a threshold frequency.

• No photoelectric emission occurs below the threshold frequency.

• Increasing the frequency increases the maximum kinetic energy of the photoelectrons.

Einstein explained the photoelectric effect using:

Photon energy = Work function + Maximum kinetic energy

Therefore: hν = hν₀ + K.E.ₘₐₓ

And: K.E.ₘₐₓ = h(ν − ν₀)

Here, ν₀ is the threshold frequency.

Thus, light intensity primarily affects the number of emitted electrons, while frequency determines their maximum kinetic energy.

Bohr’s Atomic Model

Bohr’s model is applicable to single-electron species. According to the model, electrons move in fixed stationary orbits having definite energies and do not continuously emit radiation while occupying a stationary orbit.

The angular momentum of an electron is quantised: mvr = nh/2π

where n is a positive integer.

For hydrogen-like species:

• Radius of orbit ∝ n²/Z

• Velocity ∝ Z/n

• Energy = −13.6Z²/n² eV

The negative sign in the energy expression indicates that the electron is bound to the nucleus.

When an electron moves to a higher energy level, it absorbs energy. When it moves to a lower energy level, it emits energy.

The wavelength of emitted radiation is given by the Rydberg relation: 1/λ = RZ²(1/n₁² − 1/n₂²)

where n₂ > n₁.

Bohr’s model successfully explains the hydrogen spectrum but does not adequately explain multi-electron atoms and detailed spectral splitting.

Hydrogen Spectrum

The hydrogen spectrum contains different series based on the final energy level of the electron.

The first line of a series corresponds to the transition from the immediately higher energy level to the final level. It has minimum energy and maximum wavelength within that series.

The series limit corresponds to a transition from infinity to the final level. It has maximum energy and minimum wavelength.

Quantum Mechanical Model of Atom

Louis de Broglie proposed that moving particles have wave nature. The wavelength associated with a moving particle is: λ = h/p

Since p = mv: λ = h/mv

This is known as the de Broglie equation.

Heisenberg’s Uncertainty Principle

Heisenberg’s uncertainty principle states that the exact position and exact momentum of a particle cannot be determined simultaneously.

Δx × Δp ≥ h/4π

Therefore, electrons cannot be described as moving in fixed circular paths around the nucleus.

Schrödinger Wave Equation

The Schrödinger equation provides a mathematical description of the wave nature of electrons. Its solution gives a wave function (ψ).

The square of the wave function, ψ², represents the probability density of finding an electron at a particular location.

The quantum mechanical model is the accepted model for describing electrons and is applicable to multi-electron systems.

Quantum Numbers and Orbitals

Four quantum numbers are used to describe an electron in an atom.

The possible values are:

• n: 1, 2, 3, ...

• l: 0 to n − 1

• mₗ: −l to +l

• mₛ: +1/2 or −1/2

The number of orbitals in a subshell is: 2l + 1

The maximum number of electrons in a subshell is: 4l + 2

The maximum number of electrons in a shell is: 2n²

Shapes of Orbitals

• s orbital: spherical

• p orbital: dumbbell-shaped

• d orbitals: mainly four-lobed, except dᶻ²

The number of nodes is given by:

• Radial nodes = n − l − 1

• Angular nodes = l

• Total nodes = n − 1

Electronic Configuration Rules

The arrangement of electrons in different shells and subshells follows specific principles.

Aufbau Principle

According to the Aufbau principle, electrons occupy lower-energy subshells before filling higher-energy subshells. The relative filling order is determined using the (n + l) rule.

Hund’s Rule

According to Hund’s rule, electrons occupy degenerate orbitals singly with parallel spins before pairing begins.

Pauli Exclusion Principle

According to the Pauli exclusion principle, no two electrons in an atom can have the same set of four quantum numbers.

Therefore, two electrons occupying the same orbital must have opposite spins.

Paramagnetic and Diamagnetic Species

A species with all electrons paired is diamagnetic. A species with one or more unpaired electrons is paramagnetic.

The Structure of Atom chapter connects the discovery of subatomic particles with the development of modern atomic theory. These PW Structure of Atom Class 11 notes bring together atomic models, electromagnetic radiation, hydrogen spectra, quantum numbers, orbitals and electronic configuration rules in one place. 

NEET Resources You Might Like: