How can absorption

Said in another way, electrons absorb only the photons that give them exactly the right energy they need to jump levels. Remember when we said that photons only carry very specific amounts of energy, and that their energy corresponds to their wavelength? The absorption spectrum of hydrogen shows the results of this interaction.

In the visible part of the spectrum, hydrogen absorbs light with wavelengths of nm violet , nm blue , nm blue-green , and nm red. Each of the absorption lines corresponds to a specific electron jump. Electrons can also lose energy and drop down to lower energy levels. When an electron drops down between levels, it emits photons with the same amount of energy—the same wavelength—that it would need to absorb in order to move up between those same levels.

The highest energy and shortest wavelength light is given off by the electrons that fall the farthest. Different elements have different spectra because they have different numbers of protons, and different numbers and arrangements of electrons.

The differences in spectra reflect the differences in the amount of energy that the atoms absorb or give off when their electrons move between energy levels. Molecules, like water, carbon dioxide, and methane, also have distinct spectra. All electrons vibrate at a specific frequency, which is known as their "natural" frequency. When light interacts with an atom of the same frequency, the electrons of the atom become excited and start vibrating.

During this vibration, the electrons of the atom interact with neighboring atoms and convert this vibrational energy into thermal energy. Consequently, the light energy is not to be seen again, that is why absorption differs from reflection and transmission. And since different atoms and molecules have different natural frequencies of vibration, they selectively absorb different frequencies of visible light.

As was mentioned above, everything is capable of absorbing light. For example, organic molecules are good at absorbing light.

If an organic molecule has electrons that have a high natural frequency then they absorb the light which has a high frequency as well. The longer the conjugated system conjugated system is a system of connected pi-orbitals with delocalized electrons , the longer the wavelength of the light absorbed.

Another example. Let's imagine that we are walking around a park with a lot of grass and plenty of beautiful flowers. These wavelengths are not detected by our eyes. The other wavelengths are reflected, and these are detected by our eyes. For example, grass appears green in white light:. Waves can also be transmitted at the boundary between two different materials.

When waves are transmitted, they continues through the material. Air, glass and water are common materials that are very good at transmitting light. They are transparent because light is transmitted with very little absorption.

An absorption spectrometer works in a range from about nm in the near ultra-violet to about nm in the very near infra-red. Only a limited number of the possible electron jumps absorb light in that region. Look again at the possible jumps.

This time, the important jumps are shown in black, and a less important one in grey. The grey dotted arrows show jumps which absorb light outside the region of the spectrum we are working in. Remember that bigger jumps need more energy and so absorb light with a shorter wavelength.

The jumps shown with grey dotted arrows absorb UV light of wavelength less that nm. The important jumps are:. That means that in order to absorb light in the region from - nm which is where the spectra are measured , the molecule must contain either pi bonds or atoms with non-bonding orbitals. Remember that a non-bonding orbital is a lone pair on, say, oxygen, nitrogen or a halogen. The diagram below shows a simple UV-visible absorption spectrum for buta-1,3-diene - a molecule we will talk more about later.

Absorbance on the vertical axis is just a measure of the amount of light absorbed. The higher the value, the more of a particular wavelength is being absorbed. You will see that absorption peaks at a value of nm. This is in the ultra-violet and so there would be no visible sign of any light being absorbed - buta-1,3-diene is colorless. You read the symbol on the graph as "lambda-max". That means that the only electron jumps taking place within the range that the spectrometer can measure are from pi bonding to pi anti-bonding orbitals.

A chromophore such as the carbon-oxygen double bond in ethanal, for example, obviously has pi electrons as a part of the double bond, but also has lone pairs on the oxygen atom. That means that both of the important absorptions from the last energy diagram are possible. You can get an electron excited from a pi bonding to a pi anti-bonding orbital, or you can get one excited from an oxygen lone pair a non-bonding orbital into a pi anti-bonding orbital.

The non-bonding orbital has a higher energy than a pi bonding orbital. That means that the jump from an oxygen lone pair into a pi anti-bonding orbital needs less energy. That means it absorbs light of a lower frequency and therefore a higher wavelength.

Ethanal can therefore absorb light of two different wavelengths:. Both of these absorptions are in the ultra-violet, but most spectrometers won't pick up the one at nm because they work in the range from - nm. Ethene contains a simple isolated carbon-carbon double bond, but the other two have conjugated double bonds. In these cases, there is delocalization of the pi bonding orbitals over the whole molecule.

Now look at the wavelengths of the light which each of these molecules absorbs. All of the molecules give similar UV-visible absorption spectra - the only difference being that the absorptions move to longer and longer wavelengths as the amount of delocalization in the molecule increases.

Compare ethene with buta-1,3-diene. In ethene, there is one pi bonding orbital and one pi anti-bonding orbital. In buta-1,3-diene, there are two pi bonding orbitals and two pi anti-bonding orbitals. This is all discussed in detail on the introductory page that you should have read. The highest occupied molecular orbital is often referred to as the HOMO - in these cases, it is a pi bonding orbital. The lowest unoccupied molecular orbital the LUMO is a pi anti-bonding orbital. Notice that the gap between these has fallen.

It takes less energy to excite an electron in the buta-1,3-diene case than with ethene. If you extend this to compounds with really massive delocalisation, the wavelength absorbed will eventually be high enough to be in the visible region of the spectrum, and the compound will then be seen as colored. A good example of this is the orange plant pigment, beta-carotene - present in carrots, for example. Beta-carotene has the sort of delocalization that we've just been looking at, but on a much greater scale with 11 carbon-carbon double bonds conjugated together.

The diagram shows the structure of beta-carotene with the alternating double and single bonds shown in red. The more delocalization there is, the smaller the gap between the highest energy pi bonding orbital and the lowest energy pi anti-bonding orbital. To promote an electron therefore takes less energy in beta-carotene than in the cases we've looked at so far - because the gap between the levels is less.

Remember that less energy means a lower frequency of light gets absorbed - and that's equivalent to a longer wavelength. Beta-carotene absorbs throughout the ultra-violet region into the violet - but particularly strongly in the visible region between about and nm with a peak about nm.