Energy Levels in Atoms and Line Spectra

Overview

This lesson links atomic structure to the line spectra observed from gases. You will use discrete electron energy levels to explain why atoms emit and absorb only particular photon energies and wavelengths.

What You Need to Know

• Electrons in isolated atoms can occupy only discrete energy levels.
• An emission line is produced when an electron moves to a lower energy level and emits a photon.
• An absorption line is produced when an electron absorbs a photon with exactly the energy needed to move to a higher level.
• The photon energy equals the energy difference between the two levels.
• Different atoms have different sets of energy levels, so their spectra can act like fingerprints.

How to Work Through It

1. Start by reading an energy-level diagram and identifying possible transitions.
2. Decide whether each transition would produce emission or absorption.
3. Calculate photon energies from differences between energy levels.
4. Link larger energy differences to higher frequencies and shorter wavelengths.

Check Your Understanding

• Why are atomic spectra made of lines rather than a continuous spread of colours?
• Which transition on an energy-level diagram gives the highest-frequency photon?
• What condition must be met for a photon to be absorbed by an atom?
• How can line spectra identify the elements in a gas?

Common Mistakes

• Treating energy levels as continuous rather than discrete.
• Forgetting that emission is a downward transition and absorption is an upward transition.
• Using the wrong sign for an energy change; photon energy should be a positive value.
• Assuming any photon can be absorbed if it is bright enough.

Next Steps

• Practise moving between energy differences, photon frequency, and wavelength.
• Use the photon model again when comparing wave and particle evidence.