Electromagnetic radiation exchanges energy with matter in discrete packets called photons. A photon of frequency $\nu$ has energy

$$\boxed{E=h\nu},$$

where $h$ is Planck's constant. Because electromagnetic radiation in vacuum satisfies $c=\nu\lambda$,

$$\boxed{E=\frac{hc}{\lambda}}.$$

Shorter-wavelength photons therefore carry more energy than longer-wavelength photons.

For example, a $680,\mathrm{nm}$ red photon has energy

$$E=\frac{(6.626\times10^{-34},\mathrm{J,s})(2.998\times10^8,\mathrm{m/s})}{680\times10^{-9},\mathrm{m}}\approx2.92\times10^{-19},\mathrm{J}.$$

Multiplying by Avogadro's constant, $N_A\approx6.022\times10^{23},\mathrm{mol^{-1}}$, gives about

$$176,\mathrm{kJ/mol}$$

for one mole of such photons.

Absorption is selective

Atoms and molecules do not absorb every photon equally. Their electrons occupy allowed energy states. A photon can be absorbed efficiently when its energy matches an allowed transition between states and the electromagnetic field couples to that transition.

After absorption, the system is in an excited electronic state: its electronic energy is higher than before absorption.

An absorption spectrum records how strongly a substance absorbs radiation at different wavelengths. Different electronic structures therefore produce different absorption spectra.

Absorbing a photon does not guarantee that all of its energy becomes useful chemical work. An excited state may transfer excitation energy to another molecule, transfer an electron, emit light, or relax while dissipating energy into molecular motion.

The key bridge is

wavelength → photon energy → selective electronic excitation → possible energy or electron transfer

This principle underlies spectroscopy, vision, photochemistry and photosynthetic light capture.