Color Me Crazy

Courtesy of zazzle.com

In the coming week, with St. Patrick's day right around the corner, everything will start to turn green. From candy, to beverages, to shirts galore! But green isn't the only artificial color in the world around us. In fact, nearly everything around us is colored using dyes of some sort! Before we can dive into the structure of dyes, lets first discuss the phenomenon of color in a more general sense. 

Courtesy of wikipedia.com

Visible light is a form of electromagnetic radiation (waves that have both an electric and magnetic component) which has a wavelength in the range of about 380nm to about 740nm. Light above and below these wavelengths has the same properties as visible light, but as humans we lack the ability to see it. Other animals, such as bees, however, can see these different wavelengths. For a little more detail on how light perception and vision works, check out the previous blogpost here. The specific color of light is dependent on the wavelength its wavelength, 700nm corresponds to red, 380nm corresponds to violet, and all the other colors fall between the two. Scientists typically remember this range using the  acronym ROYGBIV for Red, Orange, Yellow, Green, Blue, Indigo, Violet. White light is a combination of all these visible wavelengths. 



The perception of color is where things get a little tricky. When the human eye perceives a particular color, it can mean one of two things, either the incident light is only of that specific wavelength, or it is the entire visible spectrum of wavelength minus the complimentary color of the one that is observed. For instance, when you see a LASER that's red, the reason the light you see looks red is because the LASER is sending out a frequency of light around 700nm. Conversely, if you see a red shirt it's because molecules in the shirt are absorbing blue-green wavelengths and reflecting all the rest of the visible spectrum. Let's look at this idea in a little more detail. 

Courtesy of stainsfile.info

Electrons within a molecule are said to occupy energy levels known as ground states. Above these ground states are a number of excited states which the electrons may transition to. The electrons can only reach there excited states, however, by absorbing very specific wavelengths of light. The unique wavelengh(s) absorbed by a given molecule are dependent on its structure. Below there are a few examples given. 

When white light hits a molecule (1), the electrons absorb certain wavelengths and are promoted to their excited states(2). The remaining light is then reflected at the observer who sees the remaining wavelengths that were not absorbed (3) (see figure above). 

So when you don your green in the coming weekend, or take a sip of an artificially flavored drink, take a second to think of all those excited electrons jumping up and down, and relish in the intricate processes that bring us the beautiful world of color!

More on the history of dyes and organic chemistry:

• http://stoltz.caltech.edu/litmtg/2002/trend-lit-8_22_02.pdf

References:

• http://stainsfile.info/StainsFile/dyes/dyecolor.htm
• http://www.rsc.org/Membership/Networking/InterestGroups/OrganicDivision/organic-chemistry-case-studies/organic-chemistry-colour-dyes.asp
• http://chemistry.umeche.maine.edu/~amar/spring2010/OrganicDyes.pdf
• http://monographs.iarc.fr/ENG/Monographs/vol99/mono99-7.pdf