How can things around us have different colours if they have specific emission spectra?

Question

Objects appear in different colours because they absorb some colours (wavelengths) and reflect or transmit other colours. The colours we see are the wavelengths that are reflected or transmitted.

As far as I understand, when an atom absorbs a photon, one of its electrons gets excited (= unstable). So it drops back to the ground state, emitting that energy in the form of photons of a specific colour.

This means that we should find objects to be in certain colours. But, for example, hydrogen gas is colourless.

How is this possible? Shouldn't Hydrogen gas have a specific colour related to its emission spectra?

I was reading an article on this topic today (https://www.khanacademy.org/science/chemistry/electronic-structure-of-atoms/bohr-model-hydrogen/a/spectroscopy-interaction-of-light-and-matter), and this came into my mind.

Can anyone please help in clearing my understanding?

Answer 1

At room temperature, hydrogen gas exists as diatomic molecules that do not significantly absorb light in the visible spectrum, i.e., 400 nm to 700 nm. So hydrogen gas at room temperature is colorless. In a hydrogen gas discharge tube, that is not energized by an appropriate high voltage source, the hydrogen gas is colorless, as shown in my photograph below:

Application of high voltage, via an appropriate high voltage power supply, results in a mixture of excited hydrogen atoms and molecules. Both exist in the discharge and both emit light as they continually cycle around their respective excitation and de-excitation pathways.

This figure shows my photographs of the energized hydrogen discharge tube.

On the left is the raw output. To my eyes, it is quite noticeably red hued. On the right, a filter has been used to attenuate light at wavelengths longer than 600 nm. The light transmitted through the filter is, to my eyes, cyan hued.

Others may see the colors differently, but no matter: we have spectrometers, spectrographs, and the like. So, collecting light from the energized hydrogen discharge tube, and dispersing it using one of my homemade echelle spectrographs results in the following two dimensional spectrum (an echellogram):

I have annotated the echellogram to show grating orders, the Balmer line wavelengths, in angstroms, and a little about the acquisition of the echellogram. N.B. 1 angstrom = 0.1 nm.

As noted above, the echellogram shows the four Balmer series lines in the visible spectrum. Each is present twice, in consecutive echelle grating orders, as a consequence of how echelle spectrographs may operate. The Balmer lines are from excited hydrogen atoms emitting light as they de-excite. The other light is from other excited species, e.g., excited hydrogen molecules.

Note particularly the Fulcher alpha bands in the longer wavelength region of the echellogram. These have been know for many years and are useful for various purposes, e.g., in plasma diagnostics. A spectrum in conventional format, i.e., intensity versus wavelength, is shown below:

So it is a bit more complex than just those nice Balmer lines. The neat thing about spectroscopy is that the more you look, the more you see.

Answer 2

H absorbs very specific wavelengths, leaving almost all wavelengths undisturbed. Also the excited atoms then decay, producing the same wavelengths that were just absorbed.

Those wavelengths are the same that are produced in a gas discharge lamp, which is not at all colorless. The reason is that the lamp produces lots of light at those wavelengths. The excess is very noticeable.

Answer 3

Atomic hydrogen gas does glow with a particular colour, as long as it is at sufficient temperature to excite some of its atoms into excited levels. Here is an example:

First an RGB image taken with broadband filters. The reddish colours in this image are due to the Balmer alpha emission of hydrogen gas at 656 nm.

To show this, here is a similar image but this time the red is taken from a narrow-band filter centred on the Balmer alpha wavelength, clearly showing that the "redness" originates from that particular wavelength. [Images taken from Utah desert remote observatories].

Of course molecular (H$_2$) gas does not exhibit these colours because it is a molecule and transitions between molecular ro-vibrational levels are at infrared wavelengths.

As for why more everyday items are coloured - it has little to do with the wavelengths of atomic transitions. It is due to the solid state electron band structure that occurs when you pack atoms together in a solid material and which leads to a variety of reflectance and absorption properties.

Answer 4

As far as I understand, when an atom absorbs a photon, one of its electrons gets excited (and unstable). So it drops back to ground state, emitting that energy in the form of photons of a specific colour.

Other answers did not respond to this bit, but it’s important to realize that it is false in a lot of cases.

The electron does get excited and it is true that a system containing excited species is in a metastable state. The system wants to return to stable. Emitting a photon at the same wavelength as the absorbed photon is one way to do that. Note that this would make the material colorless, but also opaque. You would not be able to see the difference with diffuse reflection under typical viewing conditions.

However, the excess energy may also be dissipated in other ways; a common way is to dissipate it as heat.

For many materials that’s the whole story: absorption simply heats the material. Depending on which wavelengths are absorbed, you perceive a certain color.

It is reasonably common that some energy is dissipated as heat, and the remainder is emitted as a photon. The you observe a combination of the things described above; note that the dissipation may take a long time, such as in the case of fluorescent materials.