Atoms Class 12th Notes - Free NCERT Class 12 Physics Chapter 12 Notes - Download PDF

Chapter 12 of the NCERT Class 12 Physics book's Chapter 12 on Atoms provides a brief introduction to the basic concepts of atomic structure. The Atomic Theory of Dalton, Thomson's Atom Model, Rutherford's Nuclear Model, and Atomic Spectra are the major topics this chapter discusses. Spectral series, the Bohr Model of the Hydrogen Atom, energy levels of the electrons, and the limitations of Bohr's theory are some of the other topics it clarifies. The concept of the De Broglie Hypothesis is discussed in greater detail in this chapter.

This chapter plays an important role for CBSE board exams and also for the JEE and NEET as it gives the understanding of the structure of the atom and modern physics.

Also, students can refer,

• NCERT Notes Class 12 Physics
• NCERT Solutions for Class 12 Physics Chapter 12 Atoms
• NCERT Exemplar Class 12 Physics Chapter 12 Solutions Atoms

NCERT Class 12 Physics Chapter 12 Notes

This chapter is a vital part of modern physics and is frequently tested in CBSE board exams and entrance exams like JEE Main and NEET. It builds a strong conceptual base for atomic structure, quantum physics, and light-matter interaction, all of which are critical for higher studies in physics and chemistry.

1. Dalton’s Atomic Theory

• Proposed by John Dalton in the early 19th century.

• Main Postulates:

  • Matter is made up of indivisible atoms.

  • All atoms of an element are identical in mass and properties.

  • Atoms combine in fixed ratios to form compounds.

  • Chemical reactions involve the rearrangement of atoms.

• Limitations: Could not explain subatomic particles or atomic structure.

2. Thomson’s Atomic Model (Plum Pudding Model)

• Proposed by J.J. Thomson after discovering the electron.

• Model Description:

  • An atom is a positively charged sphere in which negatively charged electrons are embedded.

  • Like plums in a pudding (hence the name).

• Drawback: Failed to explain results from the alpha-particle scattering experiment.

3. Rutherford’s Nuclear Model of the Atom

• Based on the Gold Foil Experiment (Alpha Scattering Experiment).

• Alpha particles were directed at a thin gold foil. Most passed through, some deflected, and a few bounced back.

• Observations:

  • Most of the atom is empty space.

  • An atom has a small, dense, positively charged nucleus.

  • Electrons revolve around the nucleus like planets revolve around the Sun.

• Limitations:

  • Could not explain stability of atom (electrons should spiral into nucleus).

  • Did not explain discrete spectral lines of hydrogen.

• The distribution of scattered α-particles was analyzed and plotted as a graph as a function of scattering angle (θ).

• During the alpha particle scattering experiment, it was observed that as the scattering angle (θ) increased, the number of alpha particles that deflected decreased significantly. In contrast, the majority of alpha particles passed straight through the gold foil without any deflection, and only a very small fraction were deflected at large angles.

  These observations led to the following conclusions:

• Most of the atom is empty space—approximately 99.999999999996% %—since most alpha particles passed through undisturbed.

• The positive charge and mass of the atom are concentrated in a very small, dense region, which repelled the few alpha particles that were deflected. This region was named the nucleus.

• Rutherford proposed a model where:

• The nucleus is at the center of the atom, similar to the Sun in the solar system.

• This model is known as Rutherford’s Nuclear Model of the Atom, and Rutherford was the first to establish the concept of the atomic nucleus.

• Electrons revolve around the nucleus in circular orbits at very high speeds, like planets orbiting the Sun.

• The atom remains electrically neutral as the positive charge of protons is balanced by the negative charge of electrons.

• The electrostatic force of attraction between electrons and protons provides the centripetal force needed to keep the electrons in their orbits.

$\begin{array}{rl}& F_c=F_e\\ & \frac{m{v}^2}{r}=\frac{1}{4\pi {\epsilon }_0}\frac{Z{e}^2}{r^2}\end{array}$

Here r=radius of orbit,

v = velocity of orbiting electron,

e = charge of an electron,

m = mass of an electron,

Z = atomic mass of the atom,

ε₀= permittivity of free space

On solving, we get:

$r=\frac{Z{e}^2}{4\pi {\epsilon }_{0}m{v}^2}$

• Kinetic Energy(K):

putting the value of mv², we get:

$K=\frac{1}{2}m{v}^2=\frac{Z{e}^2}{8\pi {\epsilon }_{0}r}$

• Potential Energy(U):

Using the electrostatic potential between 2 charged bodies, we get:

$U=\frac{1}{4\pi {\epsilon }_0}\frac{-eX+Ze}{r}=\frac{-Z{e}^2}{4\pi {\epsilon }_{0}r}$

Here, negative signs show that there is a force of attraction, and energy has to be given to the system to overcome this force of attraction.

• Total Energy(T):

T = U + K

Some important relations to note:

Kinetic energy (K) = -(1/2)Potential energy(U)

Kinetic energy (K) = -Total energy(T)

Potential energy(U) = 2×Total energy(T)

Drawbacks of Rutherford’s Model:

1. Atomic Instability:
   Couldn’t explain why electrons don’t spiral into the nucleus while continuously radiating energy.

2. No Explanation of Spectral Lines:
   Failed to explain the line spectra of atoms, especially hydrogen.

3. No Concept of Energy Levels:
   Did not introduce quantized orbits; couldn’t explain why electrons stay in fixed paths.

4. Not Applicable to Multi-Electron Atoms:
   Could not describe atoms with more than one electron.

ATOMIC SPECTRA:

• When electrons in an atom absorb energy, they jump to higher energy levels (excited state).

• When they return to lower energy levels (ground state), they emit energy in the form of light.

• This emitted light, when passed through a prism, shows discrete lines known as the line spectrum or atomic spectrum.

• Each element has a unique atomic spectrum.
• These spectral lines correspond to specific wavelengths (energies) of emitted radiation.
• The atomic spectrum of hydrogen is divided into series based on the final energy level.

There are three main types of atomic spectra:

Emission Spectrum

• Happens when atoms gain energy (by heating or light) and get excited—electrons jump to higher energy levels.

• When electrons fall back to their original (ground) state, they release energy as light.

• The light emitted forms a bright-line spectrum (specific wavelengths).

• Every element has its own unique emission spectrum (like a fingerprint!).

Absorption Spectrum

• This is just the opposite of emission.

• White light (full spectrum) is passed through a substance.

• The substance absorbs certain wavelengths, and the rest passes through.

• The result is a spectrum with dark lines where absorption happened.

• The dark lines match the bright lines of the emission spectrum for the same element.

Continuous Spectrum

• Formed when white light passes through a prism or water droplets.

• The light spreads out into a continuous range of colors (like a rainbow).

• No gaps or dark lines—all wavelengths of visible light are present.

• Example: sunlight forming a rainbow is a continuous spectrum.

• Different colors bend differently:

  • Red bends the least (longest wavelength),

  • Violet bends the most (shortest wavelength).

SPECTRAL SERIES:

When an electron in a hydrogen atom jumps from a higher energy level to a lower one, it emits energy in the form of light. This light appears in specific wavelengths, forming what is called the hydrogen spectral series. The series are named based on the final energy level (n₁) the electron falls into.

• The Balmer Series is the only one visible to the human eye.

• All other series are in the infrared or ultraviolet region.

• The spectral lines get closer together as n2 increases.

TYPES OF SPECTRAL SERIES:

Balmer Series:

• Balmer was the first scientist to discover a spectral series of hydrogen atoms.

• It contains a visible radiation spectrum.

• It was experimentally found that these spectral lines could be expressed mathematically in the form of wavelength.

$\frac{1}{\lambda }=R(\frac{1}{2^2}-\frac{1}{n^2})$

Where

R = Rydberg constant = 109677cm^-1,

n=3, 4, 5…… (higher discrete energy state from which electron jumps to 2^nd energy state thus emitting radiation)

λ = wavelength of radiation in (cm)

• For maximum wavelength(λ_max) in the Balmer series, n=3 (has to be minimum):

$\frac{1}{{\lambda }_{max}}=R(\frac{1}{2^2}-\frac{1}{3^2})$

${\lambda }_{max}=\frac{36}{5R}$

• For minimum wavelength (λ_min) in the Balmer series

$\frac{1}{{\lambda }_{min}}=R(\frac{1}{2^2}-\frac{1}{{\mathrm{\infty }}^2})$

${\lambda }_{min}=\frac{4}{R}$

Lyman Series:

• These are the spectral series when radiation emitted is a result of jumping electrons from higher energy states to ground states.

Mathematically, it is expressed as

$\frac{1}{\lambda }=R(\frac{1}{1^2}-\frac{1}{n^2})$

Where $n=2,3,4\dots$.
- For the maximum wavelength in the Lyman series, $\mathrm{n}=2$ (has to be minimum):

${\lambda }_{max}=\frac{4}{3R}$

• For minimum wavelength in the Lyman series

${\lambda }_{min}=\frac{1}{R}$

Similarly, other series can also be expressed as:

Paschen Series:

Mathematically, it is expressed as

$\frac{1}{\lambda }=R(\frac{1}{3^2}-\frac{1}{n^2})$

where n=4, 5, 6....

Bracket Series:

Mathematically, it is expressed as

$\frac{1}{\lambda }=R(\frac{1}{4^2}-\frac{1}{n^2})$

where n=5, 6, 7.....

Pfund Series:

Mathematically, it is expressed as

$\frac{1}{\lambda }=R(\frac{1}{5^2}-\frac{1}{n^2})$

where n=6, 7, 8....

Atomic spectra play a key role in understanding the electronic structure of atoms, molecules, and ions.
Each element has its own unique spectral series, which acts like a fingerprint and helps in the identification of unknown elements.

In fact, many elements such as Helium (He), Rubidium (Rb), Caesium (Cs), Gallium (Ga), Thallium (Tl), and Scandium (Sc) were discovered using spectroscopic techniques.

BOHR’S MODEL

Proposed by Niels Bohr in 1913, this model explained the stability of the atom and the line spectra of hydrogen, improving upon Rutherford’s nuclear model.

Main Postulates of Bohr’s Model

1. Stationary Orbits:

   • Electrons revolve around the nucleus in fixed circular orbits without radiating energy.

   • These orbits are called stationary states or energy levels.

$L_N=M{V}_N{R}_N=\frac{NH}{2\pi }$

Where,

h= Planck’s constant

L_N= angular momentum of the electron in orbit

v_N= velocity of the electron in n^th orbit

R_N= radius of n^th orbit

n= permitted orbits on which electrons revolve

• Third Postulate: When electrons jump from a higher (initial) energy state to a lower (final) energy state, emit a photon of energy which is equal to the energy difference between the 2 energy states. Its frequency is given by:

$hv=E_i-E_f$

Bohr’s Radius: The radius on which electron moves around the nucleus in the orbit is described by Bohr’s model. It is known as Bohr’s radius.

Using the second postulate and Rutherford’s model

$mvr=\frac{nh}{2\pi }$

Using the value of $v^2$ from both the equations, we get

$\frac{n^2{h}^2}{4{\pi }^2{m}^2{r}^2}=\frac{1}{4\pi {\epsilon }_0}\frac{Z{e}^2}{mr}$

On solving (and putting Z=21 for the hydrogen atom) we get:

$r=\frac{n^2{h}^2{\epsilon }_0}{\pi m{e}^2}$

For the radius of the innermost orbit, putting n=1

The Velocity of Electrons in the Orbit:

$v=\frac{e^2}{2nh{\epsilon }_0}$

Energy of Orbits:

The orbital energy possessed by orbiting electrons in the discrete energy levels in Bohr’s model is known as the energy of orbits.

We also know from Rutherford’s Model that total energy (T) is given by

$T=\frac{-Z{e}^2}{8\pi {\epsilon }_{0}r}$

By putting the value of Bohr’s radius, we get

$T=\frac{-m{e}^2}{8\pi {\epsilon }_0^2{n}^2{h}^2}$

By putting the values of electron mass (m), charge (e), the permittivity of free space (ε₀), and Planck’s constant (h), we get,

$T=\frac{-13.6}{n^2}eV$

DRAWBACKS OF BOHR’S MODEL:

1. Only for Hydrogen-like Atoms:

   • The model works well only for hydrogen and other single-electron systems (like He⁺, Li²⁺).

   • It fails for multi-electron atoms (like helium, lithium, etc.).

2. No Explanation for Fine Spectral Lines:

   • Could not explain the fine structure or small splitting of spectral lines seen in high-resolution spectra.

3. No Zeeman and Stark Effects:

   • Failed to explain:

     • Zeeman effect – splitting of spectral lines in a magnetic field.

     • Stark effect – splitting in an electric field.

4. Electron Treated as a Particle Only:

   • Did not consider the wave nature of electrons, which was later explained by de Broglie’s hypothesis.

5. No Uncertainty Principle:

   • The model violates Heisenberg’s uncertainty principle, which says the exact position and momentum of an electron cannot be known simultaneously.

DE-BROGLIE’S HYPOTHESIS:

• De Broglie’s Hypothesis showed the wave-particle nature of matter.

• It explained that, like photons, electrons should even have mass or momentum and wavelength.

$p=mv=\frac{h}{\lambda }$

• It is valid just for the subatomic (microscopic) particles like electrons, protons, etc. Here mass is incredibly little, so wavelengths are massive enough to be discernible through an experiment.

• It is not valid for the macroscopic particles since their mass is very large, resulting in the wavelength being too small to be experimentally observable.

Significance of NCERT Class 12 Physics Chapter 12 Notes

These NCERT Class 12 Physics chapter 12 notes are very useful in covering the main topics of the Class 12 CBSE Physics Syllabus .

1. Foundation of Modern Physics:
   This chapter introduces key concepts like atomic models, quantum theory, and electron transitions, which are essential for understanding modern physics.

2. Helps in Board Exam Preparation:
   The chapter is a part of the CBSE Class 12 Physics syllabus and frequently appears in exams through conceptual and numerical questions.

3. Important for Competitive Exams:
   Concepts like Bohr’s model, hydrogen spectra, and de Broglie's hypothesis are often asked in JEE, NEET, and other entrance exams.

4. Connects Classical and Quantum Physics:
   It shows the shift from classical atomic models (Rutherford) to quantum models (Bohr and beyond), helping students grasp the evolution of scientific ideas.

5. Useful in Chemistry also
   Understanding atomic structure here helps in learning about electron configurations, periodic properties, and bonding in Chemistry.

6. Supports Further Studies:
   These concepts are foundational for higher studies in physics, engineering, and physical sciences.

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