Why do metals appear as coloured in flame tests, while at room temperature they don't? 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!
 A: 
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?

Yes. In a chunk of metal the energy states of the electrons are very different from in individual metal atoms. In a metal the atoms are all close together and they interact to form energy bands. The electrons in the uppermost band (the conduction band) behave in a similar way to free electrons and as a result they reflect all wavelengths of light and don't have any particular colour.
In a flame there is enough energy to dislodge metal atoms, and those individual atoms then undergo electronic transitions that emit light. For more on this see Possible colors of fire?
A: When a substance is put into a flame, the substance heats up. Some sources online say that 'electrons are excited by heat energy,' but this wasn't acceptable enough to me considering that by the quantum model of energy, electrons won't be excited gradually as excitation occurs as a one-quantum-one-electron interaction.
So, my conclusion (that satisfies me for the purpose of this question) is that in a heated environment (say ~$1000°$C for the outer regions of a bunsen flame)  atoms/electrons will have a much larger average kinetic energy - of a vibrational/rotational degree -  and thus the quantised energy to excite will be lower (as much more atoms have required 'activation energy' from collision theory) and can then excite by collisions with fast-moving gaseous particles (perhaps electron deflection in these gases?), which then de-excite via radiation of visible light photons (of course, depending on the electron transition difference).
This explanation would explain the immediate colouring shown as collision and then de-excitation would occur in a negligible time period at such a high kinetic energy. Also, the 'glow' that comes along with the colour would be the result of black-body radiation from the oscillation of vibrating electrons.
As to why it doesn't excite at room temperature and visible light, the circumstances required (high KE and particle collisions) aren't there, so it remains the shiny lattice as it would at room temp.
Anything I've gotten wrong here, please let me know!
