Bohr’s Theory

1. Electrons revolve around the nucleus in fixed paths called orbits/energy levels

2. An electron has a fixed/quantised amount of energy and it usually occupies the lowest energy level available

3. As long as the electrons remains in one energy level it neither loses or gains energy

4. When an electron absorbs energy it jumps from a lower energy level (E₁) to a higher energy level (E₂)and is now in an excited state

5. The electron is now unstable and the excess energy is released as light as the electron falls from (E₁) to (E₂)

6. Light with a definite amount of energy lost → light of a definite frequency is given off

7. The frequency of light emitted depends on the energy difference between E₁ and E₂ → E₂ - E₁ = hf

   • h: Plank’s constant

   • f: Frequency of light emitted

8. Each definite amount of energy emitted gives rise to a particles line in the emission spectrum

9. Hence electrons must occupy definite energy level

10. A each element has a different number of electrons with a unique electronic configuration, there will be different numbers and types of electron transitions for each element which give rise to unique line emission spectra

• Paschen series: infrared light → invisible

• Balmer series: visible → emission line spectrum

• Lyman series: ultra violet light → invisible