The ABCs of Chemistry: Part 5
By Lydia from SLN More Blogs by This AuthorFrom the Science Bits in a World of Bytes Blog Series
A chemistry professor said that top reason people go into chemistry is explosions. Number two is color changes. I can’t remember why I chose it. I suppose it was more of a gradual realization during my junior year of high school that Hey! This is the area of science for me! The idea that something we weren’t able to image until recently combines in weird in and wonderful ways to make all that we see inspires awe.
O is for Ozone
Ozone is another form (allotrope) of oxygen but is less stable than normal oxygen, which has two atoms bound. It is three oxygens bonded in a bent configuration. “Ozone” comes from the Greek “ozein,” which means “to smell” in apparently an intransitive sense. Yes, it is named for its smell, which can be detected at concentrations as low as 10 ppb. At standard conditions, it is a pale blue gas. This color darkens as ozone is cooled; when solidified, it has a violet-black color. It is produced by the action of UV light in the upper atmosphere or by lightning strikes.
Ozone has use in industry as a powerful oxidant. It will oxidize all metals except gold, platinum, and iridium to their highest oxidation state (large positive charge). When it reacts with a double bond, alcohols, ketones, aldehydes, or carboxylic acids can be formed, so it has been useful in pharmaceutical production. Ozone can detoxify cyanide (-CN) by changing it to cyanate (-CNO-). White blood cells produce ozone and other oxygen species to defend the body. However, this oxidative power has a dark side. Ozone can cause irritation and destruction of body tissues. People who live in areas with large amounts of ground-level ozone have greater incidence of lung diseases. Exposure to ozone constricts the blood vessels, raising the likelihood of cardiovascular events.
However, ozone benefits us greatly when confined to the stratosphere. It reacts with UV rays , shielding us from some of them. However, the use of chlorofluorocarbons (CFCs) has made ozone's reactivity a bane. A CFC is emitted into the atmosphere. UV radiation breaks off a chlorine atom, which slams into an ozone molecule. It binds to an oxygen molecule. Another ozone molecule is broken apart by UV radiation; the free atom slams into the chlorine monoxide molecule previously made, freeing the chlorine atom to repeat the process. If the ozone reacts with CFCs, it will not be available to react with the UV rays.
P is for Particle-in-a-Box Model
This concept comes from quantum chemistry and explains colors produced by certain molecules. It’s a bit odd, so bear with me. It’s based on the wave-particle duality of matter, which says that solid objects have wavelike characteristics. Yes, we are included, but due to our size, our wavelength is too short to be measured, about 10-30 to 10-40 m. To give you an idea of how insignificant that is, consider this: atomic and subatomic lengths are measured on the scale of 10-10 to 10-12 m. Only submicroscopic particles like electrons have a wavelength of any importance.
This model uses a box that is infinitely high, so that the particle can be found only in the box. This particle, for example, an electron, has wavelength due to its dual nature. The wave travels in the box. Its energy correlates inversely to the size of the box. A bigger box means that the wave is more spread out and has a lower energy (frequency). Picture yourself as an intra-building courier for Wee Widget, Inc. The plant has one floor that’s fairly small because your production is not high yet. You can go from the president’s office on end to the factory floor on the other easily and often. As the company expands, you buy more land and add on to the factory. Now the distance from the president’s office to the factory is twice as long and, due to increased volume, you have to make the trip more often. You become tired and feel spread-out.
A similar occurrence happens to electrons in molecules. As the molecule gets bigger, the electron’s wave gets stretched so that its frequency, which is proportional to its energy, decreases. There is one thing you should know: certain types of molecules promote this behavior more than others. Consider beta-carotene, the structure of which is at the beginning of this entry. It gives carrots their vivid orange color. It has a long chain with alternating single and double bonds. This state is called “conjugation” and is continued into the ring. The electrons’ waves are so stretched out that the molecule emits light on the lower-energy portion of visible light, mostly in the red and orange segments.
Q is for Quinine
Quinine was, for much of history, the antimalarial drug of choice. It was first found in the bark of the cinchona tree; indigineous South American tribes used it as muscle relaxant. They imbibed it by mixing the bitter bark with sweet water to make tonic water. During the 1600s, A Jesuit priest living in Lima, Peru observed the natives using it for malarial treatment and sent a sample of bark to Europe. In the years following, the bark became a popular commodity.
Quinine works by attacking the bacterium that causes malaria, Plasmodium Falciparum. The exact mechanism is unknown, but studies of a related drug suggest that quinine impairs the bacterium’s metabolism of blood, allowing iron-containing compounds called hemes to build up to toxic levels.
The widespread use of quinine and related drugs like chloroquine has helped a lot of people, but it also resulted in bacterial resistance to their action. A 1998 paper reported that chloroquine, an antimalarial drug with fewer adverse side effects than quinine, failed to treat nearly 100 percent of cases in certain areas of Brazil, and that effective treatment consisted of a combination of quinine and tetracycline. The paper reported that the resistance was traced to less accumulation of the drug within the cell’s digestive vacuole. This knowledge, however, allows for design of better antimalarial agents.
Fun fact: Henry Perkins was trying to synthesize quinine in the lab. He failed, but he found something else useful. He had discovered dye molecules and started the age of the modern dye.
R is for Retinal
Guess what. Retinal helps you see! This molecule is a terpene, which is a linking of multiple isoprene units. Retinal is made by lopping off part of the long chain of beta carotene and then doubly bonding an oxygen to the carbon on the end of the chain. The picture you see above is close to retinal but is not quite it; it’s vitamin A. The oxygen has a single bond to the carbon and has a hydrogen attached, while retinal has a double bond between oxygen and carbon and no hydrogen on the oxygen. That, however, is the only difference. It can flip back and forth between forms at various complexes in the body.
As stated, retinal is involved in seeing. It is attached to a protein called an opsin, and the whole complex is called rhodopsin. The starting form found in the eye has a cis double bond in the chain, giving it a kink. When a photon smacks into the molecule, it changes the cis bond to a trans one, straightening out the molecule so it no longer fits its place in the protein. The retinal molecule moves around to try to fit, but it eventually breaks free. This triggers a change in the opsin’s confirmation, which starts a cascade of chemicals that result in vision. A multi-step process regenerates the cis-retinal to reattach to the opsin and restart the process.
The molecular level grabs and holds my attention. During my classes, I learned what goes on at the small level to give epoxy glues their staying power or DNA its order. I pushed electrons all through organic to help memorize the reaction mechanism of, say, the bromination of a ketone. In biochemistry, there were binding pockets for proteins, and cooperativity in hemoglobin to keep the oxygen flowing. All this happens on time scales crazily short from our point of view. Molecular craziness works macroscopic wonders.
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