Module 1 :Atomic Structure and Periodicity of Elements Chemistry Semester 1 Notes

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Atomic Structure – Introduction

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1. Rutherford’s Atomic Model (1911)

Experiment:

• Gold Foil Experiment (Geiger–Marsden).

• Alpha (α) particles shot at thin gold foil.

Observations:

1. Most α-particles passed straight → atom mostly empty space.

2. Few deflected → presence of dense positive center (nucleus).

3. Very few bounced back → nucleus is very small.

Model:

• Atom has small, dense, positively charged nucleus containing protons (and later neutrons).

• Electrons revolve around nucleus in circular paths.

• Most of atom is empty space.

Limitations:

• Could not explain why electrons don’t spiral into nucleus.

• Could not explain line spectra of atoms.

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2. Failure of Classical Physics

Classical theory predicted:

• Accelerating electrons lose energy continuously as radiation.

• Atom would collapse → not observed.

• Spectra should be continuous, but experiments showed line spectra.

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3. Blackbody Radiation Problem

• Blackbody: Perfect absorber and emitter of radiation.

• Observed: Energy distribution curve → intensity rises with wavelength, peaks, then falls.

• Classical Rayleigh–Jeans law failed at short wavelengths (predicted infinite energy → UV catastrophe).

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4. Planck’s Quantum Hypothesis (1900)

• Energy is not continuous; it is emitted or absorbed in small packets called quanta.

• Formula:

$$E = n h \nu$$

where:
$n$ = integer,
$h$ = Planck’s constant = $6.626 \times 10^{-34} \ \text{J·s}$,
$\nu$ = frequency.

• Explained blackbody radiation successfully.

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5. Photoelectric Effect (Einstein, 1905)

Observation:

• When light of certain frequency hits a metal surface, electrons are emitted instantly.

Laws:

1. No electrons if $\nu < \nu_0$ (threshold frequency).

2. Kinetic energy of electrons depends on frequency, not intensity.

3. Number of electrons ∝ light intensity (for $\nu > \nu_0$).

Einstein’s Equation:

$$h\nu = h\nu_0 + \frac{1}{2}mv_{\text{max}}^2$$

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6. Compton Effect (1923)

• X-rays scattered by electrons have greater wavelength than incident rays.

• Explained by photon concept of light.

• Wavelength change:

$$\Delta\lambda = \frac{h}{m_e c} (1 - \cos\theta)$$

where $m_e$ = electron mass, $\theta$ = scattering angle.

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7. Bohr’s Theory of Atom (1913)

Postulates:

1. Electrons revolve in fixed circular orbits (stationary states) without radiating energy.

2. Angular momentum is quantized:

$$mvr = \frac{nh}{2\pi}, \quad n = 1, 2, 3...$$

3. Radiation emitted/absorbed only when electron jumps between orbits:

$$\Delta E = h\nu$$

Energy of nth orbit (Hydrogen atom):

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

Rydberg Formula (spectral lines):

$$\frac{1}{\lambda} = R_H \left( \frac{1}{n_1^2} - \frac{1}{n_2^2} \right)$$

where $R_H = 1.097 \times 10^7 \ \mathrm{m^{-1}}$.

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8. Limitations of Bohr’s Theory

• Works only for hydrogen-like atoms.

• Cannot explain:

  • Fine structure of spectral lines.

  • Zeeman effect (in magnetic field).

  • Stark effect (in electric field).

  • Intensities of lines.

  • Wave nature of electrons.

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9. de Broglie’s Matter Waves (1924)

• Particles have wave-like properties:

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

• Supported Bohr’s quantization — electron waves fit perfectly in allowed orbits.

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10. Heisenberg’s Uncertainty Principle (1927)

• Impossible to know exact position ($\Delta x$) and exact momentum ($\Delta p$) of a particle at the same time:

$$\Delta x \cdot \Delta p \ge \frac{h}{4\pi}$$

• For electrons: Only probability distribution can be known → led to concept of orbitals instead of fixed paths.

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All points are included, but in a condensed, easy-to-revise format so you can score full marks.

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📘 MODULE – 2 : ATOMIC STRUCTURE & CHEMICAL BONDING (Short Notes)

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🔹 Bohr’s Atom Model

Postulates:

1. Electrons revolve around nucleus in fixed circular orbits (stationary states).

2. Energy of electron is constant in a given orbit → called energy level.

3. Allowed orbits satisfy:
   $mvr = \dfrac{nh}{2\pi}$
   (n = principal quantum number).

4. Energy ↑ when orbit is farther from nucleus.

5. Electron jumps between orbits:

   • Low → High: absorbs energy.

   • High → Low: emits energy.

   • Frequency: $\nu = \dfrac{E_2 - E_1}{h}$ (Bohr frequency rule).

6. n = 1,2,3,4… = K, L, M, N shells.

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🔹 Hydrogen Spectrum

• Heated atoms/electric discharge → line spectrum (discontinuous).

• Series:

Series	n₁	n₂	Region
Lyman	1	2,3,4…	UV
Balmer	2	3,4,5…	Visible
Paschen	3	4,5,6…	IR
Brackett	4	5,6,7…	IR
Pfund	5	6,7,8…	IR

• Formula (Rydberg):
  $\dfrac{1}{\lambda} = R \left( \dfrac{1}{n_1^2} - \dfrac{1}{n_2^2} \right)$
  $R = 1.097 \times 10^7 \, m^{-1}$.

• Bohr explains H and H-like ions (He⁺, Li²⁺).

Limitations:
❌ Cannot explain fine spectrum, multi-electron atoms, bonding & shapes, Zeeman/Stark effects, against de Broglie & Heisenberg.

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🔹 de Broglie Equation

• Matter has wave-particle duality.

• Wavelength:
  $\lambda = \dfrac{h}{mv} = \dfrac{h}{p}$

• Significance:

  • Macroscopic objects → λ negligible.

  • Microscopic particles (electrons) → λ large (wave nature significant).

Derivation (short):
Photon energy: $E = h\nu = hc/\lambda$.
Einstein: $E = mc^2$.
→ $\lambda = h/mc$.
For matter: replace c with v → $\lambda = h/mv$.

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🔹 Heisenberg’s Uncertainty Principle

• Impossible to know position & momentum of electron simultaneously.

• $\Delta x \cdot \Delta p_x \geq \dfrac{h}{4\pi}$

• Rules out definite electron paths (unlike Bohr).

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🔹 Schrödinger Wave Equation

• Describes electron as a standing wave.

• Solution = wave function (Ψ).

• Ψ → no meaning; Ψ² → probability of finding electron.

• Energy of electron → quantized.

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🔹 Orbitals

• Region in space where probability of finding electron is maximum.

• s orbital → spherical.

• p, d, f → directional.

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🔹 Quantum Numbers

1. Principal (n): main shell, distance from nucleus (n=1,2,3…).

2. Azimuthal (l): subshell (0 to n-1).

   • l=0 (s), l=1 (p), l=2 (d), l=3 (f).

3. Magnetic (m): orientation (values: -l to +l).

4. Spin (s): spin ±½ (↑ or ↓).

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🔹 Electron Arrangement

1. Pauli Exclusion: one orbital max 2 e⁻ with opposite spins.

2. Aufbau Principle: electrons fill in increasing energy order:
   1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p …

3. Hund’s Rule:

   • Each orbital gets 1 electron before pairing.

   • Unpaired electrons → parallel spins.

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🔹 Stability of Configurations

• Half-filled & fully filled orbitals = extra stable.

• Reasons:
  i) Symmetry.
  ii) Exchange energy (parallel spin exchange releases energy).

• Examples:

  • Cr (24): [Ar] 4s¹ 3d⁵

  • Cu (29): [Ar] 4s¹ 3d¹⁰

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🔹 Chemical Bonding

Ionic (Electrovalent) Bond

• Electrostatic attraction between +ve and -ve ions (formed by transfer of electrons).

• Examples: NaCl, MgCl₂.

Favourable conditions:

1. Low ionization energy of donor atom.

2. High electron affinity of acceptor atom.

3. High (negative) lattice energy.

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✅ This is your entire Module-2 in short form, all points covered.