Photons and the Photoelectric Effect showed that a photon's energy depends on its wavelength. Molecules can store energy in three ways: by spinning (rotation), by stretching and bending their bonds (vibration), and by moving electrons to higher orbitals. Each way needs photons of a different size, so each region of the spectrum does a different job.
Key idea
Microwaves make molecules rotate faster. Infrared light makes bonds vibrate. Visible and ultraviolet light move electrons to higher energy levels. Only photons with at least a bond's energy (most UV, and beyond) can break bonds.
Why it works
The energy steps for each kind of motion have very different sizes. Rotational levels are spaced by tiny amounts, well under $1\ \text{kJ/mol}$, which matches microwave photons. Vibrational levels are spaced by roughly 5 to 50 kJ/mol, which matches infrared photons. Electronic levels are hundreds of kJ/mol apart, which matches visible and ultraviolet photons. A molecule absorbs a photon only when the photon's energy fits one of its steps, the same rule you saw for hydrogen's levels. Compare bond energies: about 150 to 950 kJ/mol.
| Light | Wavelength | Energy per mole of photons | What it does |
|---|---|---|---|
| Microwave | 1 cm | 0.012 kJ/mol | rotation |
| Infrared | 10 μm | 12 kJ/mol | vibration |
| Visible | 500 nm | 239 kJ/mol | electronic excitation |
| Ultraviolet | 200 nm | 599 kJ/mol | electronic excitation; can break bonds |
Worked example: An infrared photon
Carbon dioxide absorbs infrared light near $15\ \text{μm}$. Find the energy per mole of these photons, and say what they do to the molecule.
- $$E = \dfrac{hcN_\text{A}}{\lambda} = \dfrac{(6.626 \times 10^{-34})(3.00 \times 10^{8})(6.022 \times 10^{23})}{15 \times 10^{-6}\ \text{m}} = 7.98 \times 10^{3}\ \text{J/mol}$$
- That is about $8.0\ \text{kJ/mol}$: far too little to break a bond (C=O is about 800 kJ/mol) or lift an electron, but the right size for a vibration. These photons make the molecule bend.
Answer: About $8.0\ \text{kJ/mol}$; they set the molecule bending.
Worked example: Enough to break a bond?
A germicidal lamp emits $254\ \text{nm}$ ultraviolet light. Does one photon carry enough energy to break a C–C bond ($347\ \text{kJ/mol}$)?
- Per mole of photons: $$E = \dfrac{(6.626 \times 10^{-34})(3.00 \times 10^{8})(6.022 \times 10^{23})}{2.54 \times 10^{-7}} = 4.71 \times 10^{5}\ \text{J/mol} = 471\ \text{kJ/mol}$$
- $471 > 347$, so yes: each photon carries more than one bond's worth of energy.
Answer: Yes: $471\ \text{kJ/mol} > 347\ \text{kJ/mol}$.
Common mistake
Saying infrared light makes electrons jump between levels, or that any light can break any bond. An infrared photon carries only a few to tens of kJ/mol, enough to set bonds vibrating, which we feel as warmth; electronic jumps and bond breaking need hundreds. Check: work out $hcN_\text{A}/\lambda$ and compare it with the energy of the step you have in mind.
This is the physics behind the greenhouse effect: the ground gives off infrared light, and $\mathrm{CO_2}$ and $\mathrm{H_2O}$ absorb some of it through their vibrations. A microwave oven works because its rapidly changing field twists polar water molecules back and forth, and their jostling heats the food.