Atomic Structure: Subatomic Particles, Models, Spectra & Quantum Numbers

What is an atom?

An atom is the smallest particle of an element that can take part in a chemical reaction. It consists of three main subatomic particles:

• Protons (positively charged)
• Neutrons (no charge)
• Electrons (negatively charged)

Atomic number (Z): Number of protons in the nucleus (defines the element). Mass number (A): Total number of protons + neutrons. Atomic mass: Average mass of all isotopes of an element (measured in atomic mass units, u).

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Discovery of Subatomic Particles

1. Electron (J.J. Thomson, 1897)

• Thomson discovered electrons using a cathode ray tube experiment.
• He found that cathode rays were made of negatively charged particles (electrons).
• He proposed the "plum pudding model" where electrons were embedded in a positively charged "pudding" of matter.

2. Proton (E. Goldstein, 1886)

• Goldstein observed positively charged rays (canal rays) in a discharge tube.
• These rays were made of protons (positively charged particles).

3. Neutron (James Chadwick, 1932)

• Chadwick discovered neutrons using alpha particle bombardment on beryllium.
• Neutrons have no charge but contribute to the mass of the nucleus.

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Rutherford’s Nuclear Model (1911)

• Rutherford conducted the gold foil experiment where alpha particles were fired at a thin gold foil.
• Most particles passed through, but some bounced back.
• This led to the nuclear model:
  • Most of the atom is empty space.
  • The nucleus (protons + neutrons) is small, dense, and positively charged.
  • Electrons revolve around the nucleus like planets around the sun.

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Bohr’s Model of the Atom (1913)

• Bohr improved Rutherford’s model by introducing quantized energy levels.
• Electrons can only exist in specific orbits (shells) with fixed energy.
• When an electron jumps from a higher to a lower orbit, it emits energy as light (photons).
• When an electron absorbs energy, it jumps to a higher orbit.

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Wave Nature of Electrons (De Broglie, 1924)

• De Broglie proposed that electrons have wave-like properties (wave-particle duality).
• The wavelength of an electron is given by: $$\lambda =\frac{h}{mv}$$ where:
  • $\lambda$ = wavelength
  • $h$ = Planck’s constant ($6.626\times 10^{-34}$ Js)
  • $m$ = mass of electron
  • $v$ = velocity of electron

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Quantum Mechanical Model (Schrödinger, 1926)

• Electrons do not move in fixed orbits but exist in orbitals (probability clouds).
• Orbitals are regions where there is a high probability of finding an electron.
• The wave function (ψ) describes the electron’s behavior in an orbital.

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Quantum Numbers

Electrons in an atom are described by four quantum numbers:

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Electron Configuration

• Electrons fill orbitals following the Aufbau principle, Pauli exclusion principle, and Hund’s rule.
• The order of filling orbitals is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f ...

Example: Electron configuration of Magnesium (Z = 12)

• 1s² 2s² 2p⁶ 3s²

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Atomic Spectra

• When electrons jump between energy levels, they emit or absorb light of specific wavelengths.
• This creates atomic spectra (unique for each element).
• Hydrogen spectrum shows four main series:
  • Lyman series (UV region)
  • Balmer series (visible region)
  • Paschen series (infrared region)
  • Brackett series (infrared region)

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Isotopes, Isobars, and Isotones

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Solved Examples

Example 1: Calculate the number of neutrons in ${}_{26}^{56}Fe$

• Atomic number (Z) = 26 (number of protons)
• Mass number (A) = 56
• Number of neutrons = A - Z = 56 - 26 = 30

Example 2: Write the electron configuration of Chlorine (Z = 17)

• Follow the Aufbau principle: 1s² 2s² 2p⁶ 3s² 3p⁵

Example 3: Identify the element with electron configuration 2, 8, 2

• Total electrons = 2 + 8 + 2 = 12
• Atomic number (Z) = 12 → Magnesium (Mg)

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NEB Board-Style Questions

Short Answer Questions

1. What are the main postulates of Bohr’s atomic model?
2. Explain the significance of Rutherford’s gold foil experiment.
3. Define isotopes. Give an example.
4. What is the difference between an orbital and an orbit?
5. Write the electron configuration of Sodium (Z = 11).

Long Answer Questions

1. Describe the discovery of the electron, proton, and neutron. How did these discoveries lead to the development of atomic models?
2. Explain the quantum mechanical model of the atom. How does it differ from Bohr’s model?
3. What are atomic spectra? Describe the hydrogen spectrum and its significance.
4. Define isotopes, isobars, and isotones. Give one example of each.

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Exam Tip

• Memorize the order of filling orbitals (Aufbau principle).
• Understand the differences between Bohr’s model and the quantum mechanical model.
• Practice writing electron configurations for given atomic numbers.
• Know the significance of atomic spectra in identifying elements.
• Compare isotopes, isobars, and isotones clearly in tables or examples.
• Draw diagrams of atomic models (Thomson, Rutherford, Bohr) and orbitals (s, p, d, f) in exams.

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