I'm not sure if this is more of a chemistry-type question, but my question focuses more on light spectroscopy than the chemical elements;
Why do some metals appear coloured in flame tests while they don't appear coloured at all at room temperature when in the presence of light? Is it just that flame provides far more energy than light could?
My understanding of colouring was that when a substance is exposed to wavelengths across the visible spectrum, it absorbs specific wavelengths for excitation and thus appears as the colours that are allowed to reflect/transmit, and the electrons de-excite via irradiative processes.
All the textbooks explain the flame tests as the emission of a specific wavelength during de-excitation, but as far as I was aware that results in a 'fluorescence' phenomena and not the colour that we observe.
Or is it just that some of the metal atoms vaporise and thus behave as gases, with emission spectrums being radiated in random directions and the absorption spectrum being negligible?
My thoughts right now;
- the flame provides far more energy than light, which is why it doesn't appear coloured at room temperature
- the specific energy absorbed doesn't change the spectrum incident in a non-negligible way, but the emission photons are much more intense relatively speaking, and thus the flame would appear as the emission colour
Any help would be appreciated, let me know if this is too chemistry-esque for this forum. Thanks!