In quantum physics, the energy of a photon—the fundamental particle of light—is directly proportional to its frequency. This relationship is expressed by the famous equation:
E = hf
Where:
- E is the energy of the photon
- h is Planck’s constant (a fundamental constant of nature)
- f is the frequency of the light wave
What This Means:
- Higher frequency light (like ultraviolet or X-rays) consists of photons with more energy.
- Lower frequency light (like radio or infrared) has less energetic photons.
- This explains why X-rays can penetrate tissues, while radio waves cannot—X-ray photons have much more energy.
Key Implications:
- Quantum Nature of Light: This idea, introduced by Max Planck and expanded by Albert Einstein, helped establish that light comes in discrete energy packets—photons—not just continuous waves.
- Photoelectric Effect: It explains why light of a certain frequency can eject electrons from a metal surface—even if the intensity is low—as long as each photon has enough energy.
- Spectroscopy: The energy-frequency link allows scientists to analyze atoms and molecules by studying the light they emit or absorb.