Atomic Models: Class 11 Chemistry Notes on Thomson, Rutherford and Bohr

The Plum Pudding Model by J.J. Thomson shows the atom as a uniform sphere of positive charge with negatively charged electrons embedded within. The electron distribution resembles plums in a pudding, or much like seeds in a watermelon.

Thomson in 1904 suggested that atoms are positively charged spheres with electrons distributed evenly. This was based on his earlier discovery of the electron via cathode rays in 1897.

Merits of Thomson's Plum Pudding Model

• The equal amounts of positive and negative charge result in electrical neutrality. 
• It was the first atomic model to incorporate electrons as fundamental particles.
• It could explain why atoms are normally electrically neutral.

Key Postulates of Thomson’s Plum Pudding Model

• Electrons are continuously dispersed or spread out in a positive matrix.
• It accounts for the atom's neutrality in uncharged states.
• The mass is uniformly distributed, too. 

Limitations of Thomson’s Plum Pudding Model

• This model cannot explain the discrete line spectra observed in elements, as it offers no mechanism for quantised electron energies.
• It was later disproved by Rutherford's α-particle scattering experiment in 1909, which indicated a concentrated nucleus. If Thomson's atom were true, α-particles would pass through with minimal deflection due to gentle, distributed repulsion. Therefore, no α-particle should bounce back. But, Rutherford observed that some α-particles did bounce back. This proved Thomson wrong, with the existence of a concentrated positive charge that could cause large deflections.

 

Rutherford's nuclear model describes the atom as a tiny, dense, positively charged nucleus surrounded by electrons orbiting in a largely empty space. This model was closely understood as analogous to the planetary system. 

Based on alpha particle scattering experiments (1909 to 1911), Rutherford, with Hans Geiger and Ernest Marsden, concluded that atoms have a small, heavy nucleus containing protons (and later neutrons), while the electrons are in an extranuclear region. Most alpha particles passed through a gold foil undeflected, but some were scattered at large angles. 

Experimental Basis of Rutherford’s Nuclear Model

Alpha particles bombarded a thin gold foil, that is approximately $\left({10}^{-4}\mathrm{m}\mathrm{m}\right)$ .

Observations of the Rutherford Model

• Most particles passed straight and indicated a vast space.
• Some deflected at large angles, suggesting a dense, positive core.
• Few (1 in 10,000) retraced paths, confirming a compact nucleus.

Merits of the Rutherford Model

• Rutherford was the first to establish the nuclear structure of atoms.
• The Rutherford Model could explain the results of alpha particle scattering experiments.
• It introduced the concept of atomic number (number of protons).

Key Features of the Rutherford Model

• Nucleus holds nearly all atomic mass, with protons and neutrons (nucleons).
• Electrons orbit at a distance, occupying most of the atom's volume.
• Nucleus size: $1.5-6.5$ Fermi $\left({10}^{-13}\mathrm{c}\mathrm{m}\right),{10}^{-5}$ times the atom's size.
• Nucleus density: ${10}^{15}\mathrm{g}/{\mathrm{c}\mathrm{m}}^3$ , given by: $\text{Density}=\frac{\text{mass number}}{6.023\times {10}^{23}\times \frac{4}{3}\pi r_n^3},\mathrm{}{r}_n\approx 1.4\times {10}^{-13}{A}^{1/3}\mathrm{c}\mathrm{m}$

Limitations of the Rutherford Model

• According to classical physics, revolving electrons should emit radiation, lose energy, and spiral into the nucleus within ~10⁻⁸ seconds, for which the term atomic collapse was used. This would make all atoms unstable and matter as we know it impossible to exist.
• Rutherford had no explanation for the stability of atoms or the arrangement of electrons. The model could not predict why certain electron configurations are preferred or how chemical bonding occurs.
  
  
  

After Rutherford established that the positive charge of an atom is concentrated in a nucleus containing protons, scientists realised that the number of protons defines the atomic number (Z). It was also possible to determine that protons and neutrons together account for the atomic mass number (A) and almost all the atom's mass. Also note that we use a collective name for protons and neutrons together, naming them as nucleons. This term is used as they both reside in the nucleus.

To understand atomic number (Z) and atomic mass (A), let's consider hydrogen and helium atoms. 

• Hydrogen has 1 proton (atomic number = 1) and helium has 2 protons (atomic number = 2). Since atoms are electrically neutral, hydrogen has 1 electron and helium has 2 electrons to balance the positive charge.\
• Helium has 2 protons and 2 neutrons, so its atomic mass number is 4. Hydrogen typically has 1 proton and 0 neutrons, giving it an atomic mass number of 1.

To remember better, you can look into the NCERT's section 2.2.3, which highlights

"Atomic number (Z) = number of protons in the nucleus of an atom = number of electrons in a neutral atom"

"Mass number (A) = number of protons (Z ) + number of neutrons (n)"

Let's simplify another aspect of the Rutherford model, often necessary to be conceptually clear when preparing for exams. 

Isotopes

Isotopes are atoms of the same element with the same atomic number (Z) but different mass numbers (A). They have the same number of protons and electrons, but a different number of neutrons.
What you should remember, they share the same chemical properties. The atomic number determines the chemical behaviour. That is for electrons. But they have different physical properties due to differences in mass.

Example of Isotopes

Hydrogen isotopes

• Protium: 1/1 H 1 has 1 proton, 0 neutrons
• Deuterium: 2/1 D has 1 proton, 1 neutron
• Tritium: 3/1 T has 1 proton, 2 neutrons

Isobars

Isobars are atoms of different elements that have the same mass number (A). But these atoms have a different atomic number (Z).

Remember, they have different numbers of protons, neutrons, and electrons, which means their chemical properties are different. Only their nuclear mass is the same.

Example of Isobars

14/6 C and 14/7 N. Both have a mass number (A) of 14 but different atomic numbers (Z) of 6 and 7. That makes them different elements.

Isotopes vs Isobars

Property	Isotopes	Isobars
Atomic number (Z)	Same	Different
Mass number (A)	Different	Same
Number of protons	Same	Different
Number of neutrons	Different	Different
Element identity	Same element	Different elements
Chemical properties	Same (same electronic configuration)	Different (different electronic structure)
Physical properties	Different (due to mass difference)	Generally different

Bohr's atomic model introduces quantised electron orbits, where electrons revolve around the nucleus in fixed energy levels without radiating energy. Together, they explain atomic stability and spectral lines. 

You’ll learn that Niels Bohr (1913) proposed that electrons occupy stationary orbits with quantised angular momentum.  These could absorb or emit energy during transitions, accurately predicting the hydrogen's line spectra. 

Merits of Bohr's Atomic Model

• Bohr's Atomic Model successfully explains the line spectra of hydrogen and hydrogen-like ions (He⁺, Li²⁺).
• This model accurately calculates values for atomic radii and ionisation energies, with the orbital radius could be calculated as $r_n=\frac{n^2}{Z}\times 0.529\text{Å}$.
• Determines electron velocity: $v_n=\frac{2.188\times {10}^8\mathrm{Z}}{n}\mathrm{c}\mathrm{m}/\mathrm{s}$
• Bohr is the one who introduced the concept of quantised energy levels in atoms.
• Provides the theoretical foundation for understanding atomic stability.

Postulates of Bohr’s Atomic Model

• Electrons move in discrete orbits without energy loss. 

• Angular momentum is quantised using the equation $mvr=\frac{nh}{2\pi },\mathrm{}n=1,\mathrm{2,3},\dots$ where $n$ is the principal quantum number, $h$ is Planck's constant. Refer to the Planck's equation article.  
• Energy change during orbit transitions: $\mathrm{\Delta }E=E_{n_2}-E_{n_1}$
• Electron energy in orbit $n$ : $E_n=-\frac{2{\pi }^{2}m{Z}^2{e}^4{k}^2}{n^2{h}^2}=-13.6\frac{Z^2}{n^2}\mathrm{e}\mathrm{V}$ where $Z$ is atomic number, $m$ is electron mass, $e$ is the electron charge, $k$ is Coulomb's constant. 

• Energy emitted during the transition from a higher $\left(n_2\right)$ to a lower $\left(n_1\right)$ orbit is explained by the equation $\mathrm{\Delta }E=hv=13.6Z^2\left(\frac{1}{n_1^2}-\frac{1}{n_2^2}\right)\mathrm{e}\mathrm{V}$ .  Electron energy in the nth orbit: Eₙ = -13.6Z²/n² eV (for hydrogen-like atoms) is derived using Coulomb's Law for electrostatic attraction between nucleus and electron.
• The wave number of emitted radiation $\bar{v}=\frac{1}{\lambda }=R{Z}^2\left(\frac{1}{n_1^2}-\frac{1}{n_2^2}\right),\mathrm{}R=109678{\mathrm{c}\mathrm{m}}^{-1}$

Hydrogen spectral series in Bohr's Atomic Model

• • Lyman ( $n_1=1,\mathrm{U}\mathrm{V}$ ): ${\lambda }_{\text{max}}=\frac{4}{3R},{\lambda }_{\text{min}}=\frac{1}{R}$ .
  • Balmer ( $n_1=2$ , visible): ${\lambda }_{\text{max}}=\frac{36}{5R},{\lambda }_{\text{min}}=\frac{4}{R}$ .
  • Paschen ( $n_1=3,\mathrm{I}\mathrm{R}$ ): ${\lambda }_{\mathrm{m}\mathrm{a}\mathrm{x}}=\frac{144}{7R},{\lambda }_{\mathrm{m}\mathrm{i}\mathrm{n}}=\frac{9}{R}$ .

Limitations of Bohr's Atomic Model

• Bohr's atomic model does not apply to multielectron atoms due to electron-electron interactions
• This model of the atom cannot account for the Zeeman effect (the splitting of spectral lines in magnetic fields) or the Stark effect (splitting in electric fields). These effects reveal additional quantum mechanical complexities beyond Bohr's simple orbital model.
• Bohr ignores electron wave-particle duality and Heisenberg's uncertainty principle. It treats electrons as classical particles with definite positions and velocities, which quantum mechanics shows is impossible.

This section contrasts Thomson's and Rutherford's models with Bohr's revolutionary approach. Bohr's atomic model, which earned him the Nobel Prize in 1922, represented a major breakthrough in seeing atomic structure in a new light. We recommend you refer to our article on Developments Leading to the Bohr Model of Atom. 

Here is a comparison table, highlighting how atomic models were understood up until Bohr. 

Feature / Aspect	Thomson's Model (1898)	Rutherford's Model (1911)	Bohr's Model (1913)
Basic Idea	Atom is a uniform, positively charged sphere with embedded electrons (plum pudding model).	Atom has a tiny, dense, positively charged nucleus with electrons revolving around it (like planets around the Sun).	Electrons revolve in fixed quantised orbits with defined energies with no radiation in these stable orbits.
Charge Distribution	Positive charge uniformly spread. Electrons embedded to neutralise.	Positive charge concentrated entirely in the nucleus; electrons outside it.	Same as Rutherford's positive charge in nucleus, electrons outside, but only allowed in certain quantised paths.
Mass Distribution	Mass evenly spread throughout the atom.	Nearly all mass in the nucleus; electrons negligible in mass.	Same as Rutherford.
Electron Arrangement	Electrons are fixed or embedded within a sphere of uniform positive charge	Electrons move in circular, planet-like orbits around a central, positive nucleus. A major flaw was that, according to classical physics, these electrons should lose energy and spiral into the nucleus.	Electrons move in specific, stationary orbits (or energy levels) where their angular momentum is quantized, given by the formula: mvr = nh/(2π).
Empty Space in Atom	Atom thought to be solid with no space.	Atom mostly empty space, while nucleus is extremely small.	Yes, also mostly empty space.
Explanation of Spectral Lines	Cannot explain discrete spectra.	Cannot explain discrete spectra.	Can explain discrete line spectrum of hydrogen & hydrogen-like species using quantised energy levels.
Atomic Stability	Could not explain stability, and nothing prevents collapse.	Could not explain stability, where electrons should radiate energy and spiral inwards.	Stability explained: electrons in fixed orbits do not radiate energy; radiation occurs only during transitions.
Formulas / Results	None for size or energy levels.	Estimated atom & nucleus size from alpha scattering.	Derived formulas for orbital radius, electron velocity, and energy
Limitations	Failed for scattering and spectra.	Failed for stability and spectra.	Failed for multi-electron atoms; could not explain Zeeman/Stark effects or electron wave nature.

 

Go through the chapters in NCERT Class 11 Chemistry, and check out other topics, regularly updated to align with your learning needs. 

Chemistry Chemical Equilibrium

Chemistry Structure of Atom

Chemistry Redox Reactions

Chemistry Some Basic Concepts of Chemistry

Chemistry Organic Chemistry

Chemistry Classification of Elements and Periodicity in Properties

Chemistry Chemical Bonding and Molecular Structure

Chemistry Hydrocarbon