This lesson explains how electrons occupy energy levels, how they absorb and release energy, and how this leads to the formation of atomic spectra.

1. Core Concepts (Short Notes)

3.1 Energy Levels in Atoms

Electrons in an atom are arranged in discrete (quantized) energy levels or shells.

• Lower energy levels are closer to the nucleus.

• Higher energy levels are farther from the nucleus.

• Electrons prefer the lowest possible energy level (ground state).

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3.2 Excitation & Relaxation

Electrons can absorb energy and move to higher levels.

• Excitation: Electron jumps from lower → higher level.

• Relaxation: Electron falls back to lower level, releasing energy.

The released energy appears as light of specific wavelengths.

3.3 Line Spectra

Atoms emit or absorb radiation only at specific wavelengths, producing:

• Emission spectra: Bright colored lines on a dark background.

• Absorption spectra: Dark lines on a colored background.

These lines act like a fingerprint for each element.

3.4 Hydrogen Spectrum

The hydrogen atom produces several series of spectral lines:

• Lyman series: UV region (transitions to n = 1)

• Balmer series: Visible region (transitions to n = 2)

• Paschen series: Infrared region (transitions to n = 3)

Each line corresponds to an electron transitioning between specific energy levels.

3.5 Quantized Energy

Electron energies are quantized; they cannot exist between energy levels. The energy difference between levels determines the color/wavelength of emitted or absorbed light.

2. Key Formulas to Memorize

Energy of a Photon

E = hν

Energy–Wavelength Relationship

E = hc / λ

Energy Level Transitions (Hydrogen)

ΔE = E₂ - E₁

If ΔE is:

• Positive: Absorption takes place.

• Negative: Emission occurs.

  
  

Rydberg Equation for Hydrogen Lines

Used to calculate wavelength: 1/λ = R (1/n₁² – 1/n₂²) Where:

• R = Rydberg constant

• n₁ < n₂

3. Tips & Tricks for Exams

• When an electron falls to a lower level → it emits energy.

• Bigger energy gap → higher frequency and shorter wavelength.

• Balmer series lines appear in the visible region—often tested.

• UV lines correspond to high energy transitions (ending at n = 1).

• Use diagrams to visualize transitions for clarity.

• Remember: absorption lines occur because electrons absorb energy to jump up.

4. Important Points to Remember

• Electrons occupy quantized energy levels.

• The spectrum of hydrogen gives strong evidence for quantized energy.

• Emission occurs when electrons lose energy.

• Each element has a unique spectrum based on its electronic structure.

• Short wavelength → high frequency → high energy.

Atomic spectra are crucial for identifying elements in stars and distant galaxies.