Why does the flame color change with chemicals?

Answer 1

The different colors of a fire represent different temperatures. Every material has a slightly different color spectrum corresponding to their respective flames when they're burned, but the most common spectrum is the one due to the consumption of oxygen in air, which is the familiar blueish flame at the bottom, orangeish flame in the middle, and yellowish-whiteish flame at the top. Therefore, that's the example I'll use in the rest of this answer.

The first thing to remember for the rest of this answer to make sense is that an electron in an atom is only allowed to have discrete, or quantized, amounts of energy. The second thing to remember is that an atom's electrons are always trying to reach the lowest possible value of these discrete energy states that they can, because that stabilizes them. The last thing to remember is that energy is conserved, always.

That being said, in a fire, the part of the flame that's closest to the source of what's being burned is naturally going to be the hottest part of the flame, since at that point, no thermal energy has yet been released. However, as we continue examining the flames of a fire at further and further distances away from the source, we notice that they get cooler and cooler. This is because more and more energy is continuously being emitted from the flames but no energy is getting absorbed by them, resulting in a net loss of energy. And, since temperature is nothing more than the average kinetic energy of all of the atoms within a system, a loss of energy is the same thing as a temperature drop.

So, what does all of this have to do with color? As stated above, the answer lies in the quantization of energy, the electron's desire to be in lower-energy states of higher stability, and the conservation of energy.

OK, follow along. Energy is conserved. When something gets burned, a chemical reaction takes place that releases energy. That energy can't just magically disappear, because energy is conserved. Where does it go? In this case, most of it is radiated away as infrared radiation, but some of it is absorbed by atomic electrons.

OK, so the electrons now have some of the energy, what next? Well, since energy is conserved, the electrons suddenly find themselves placed in one of those quantized energy levels I was talking about earlier, except these levels have way to much energy for the electrons to feel stable. Therefore, as stated before, twice in fact, the electrons will keep trying to drop down to lower and more stable energy states, and they will eventually succeed.

But, what about the friggin' colors?! I'm not sure if I mentioned this yet of not, but energy is always conserved. So, if an electron drops down to a lower energy state, which means that it loses some energy, that energy has to go somewhere, but where?!?!

And now for the moment you've all been waiting for...the energy that's lost from the electron ends up as an electromagnetic wave having a specific energy that's exactly equal to the energy lost by the electron. And, you guessed it, many of those specific values of electromagnetic energy fall within the visible light spectrum, hence COLOR!

But, why blue at the bottom, then orange, then yellow, then white? I knew I was going to ask that. Remember those discrete, quantized energy levels that the electron could be in? I hope so, because we were just talking about them. Well, it turns out that the values for the energies of those levels get closer and closer and closer together numerically as the electron gains more and more and more energy. For example, let's say that, hypothetically, the first energy level has a value of 10. Then the higher energy values would increase in a way kind of like this:18, then 24, 28, 30, 31, and so on. As you can see, the values themselves are increasing, but the difference between the values is decreasing, and that's what we're concerned with, because, you guessed it, energy is always conserved. Therefore, the electromagnetic wave that is emitted when the electrons lose energy is equal to the difference in energy between the initial and final state the electron was in.

So, you may have made the color connection by now, but in case not, here's two more useful little pieces of information:

1) Only one electron can be in one place at one time, or in this case energy state, or quantum state.

2) The greater the drop in energy of the electron, the greater the energy of the emitted electromagnetic wave, meaning the higher the frequency of that wave.

Keeping this information in mind, when these atomic electrons try to start plummeting down energy levels, because remember, they all want to do that, a problem occurs: there's a whole bunch of other electrons in their way, taking up their energy spot! So, since only one electron can occupy any particular spot at any given time, all the electrons with the higher energy values have to patiently wait for all of the electrons with the lower energy values to get out of their way first (some of them cheat and just cut in front of them though). This means that most of the electrons that initially lose their energy are those having the lowerenergy values, but the greater energy differences. Now's the time to remember that greater energy differences correspond to higher frequency wave emissions, which in turn correspond to, finally, the color blue, which is indeed the color of the flame closest to the source.

Once those electrons have moved out of the way, the ones energetically "above" them can fall into their vacated spots, emitting waves of slightly lower frequencies in the process, because, remember, the differences between energy levels gets smaller as the values for the energies get higher. Hence, waves corresponding to colors having less energy than blue, like green, yellow orange...WHOA, I don't remember there being any green mentioned in the fire!

You're right, it wasn't mentioned, because waves, unlike most particles, can be in the same place at the same time, a phenomena called superposition. Now, who remembers from elementary school what color light is that is a superposition of all the visible wavelengths? That's right, it's white! So what's happening in fire isn't that the colors are going from blue to red, but rather that they're going from blue to whiteas more and more waves are emitted with slightly lower and lower frequencies, all being superposed on top of each other.

Answer 2

Because, different atoms start off with electrons in different energy states and when they are burned the electrons get excited for a short period of time and then relax to their original state by giving off a color. The colors will change based on the different energy levels the electrons are at.

Answer 3

When a metal or metal salt is burned, the input of thermal energy raises the electrons in the metal atom to a higher energy state. These electrons cannot remain in this excited state for too long and will emit energy in the form of light to return to the more stable, grounded state. It is this light we see when a metal atom is burned in a flame.

Answer 4

The wavelength of light (and therefore the colour) emitted depends on the temperature. The shorter the wavelength, the hotter the object. Blue has a shorter wavelength than red, and therefore a blue flame is hotter than a red flame. We are all emitting radiation, but the wavelength is too long for us to see (infrared). If you heat something up enough, it will eventually emit light that you can see, it will be 'red hot'. So to answer your question, if adding a chemical increases the temperature of the flame, it will change the colour.

The peak wavelength can be determined by Wiens Law:

Wavelength(in nanometres) = (2.9 X 10^6)/Temperature(in Kelvin)

Answer 5

The reason why the color changes in the flame is because the flame will appear a different color dependent upon the chemical additives.

Answer 6

it gets hotter or cooler