How fluorescent lights work: quantum mechanics in the home


We have a tendency to think that “quantum mechanics” is synonymous with “out of the ordinary.” I do that, too, since there’s so much strange to talk about: the blurring of particles and waves, the apparent randomness that drove Einstein crazy, and so forth. It’s easy to forget that quantum mechanics also is an everyday matter. The odds are pretty good you’re reading this post on a computer screen (as opposed to a printout), and possibly the light you’re using is fluorescent.


The three major types of lights you can buy are incandescent bulbs, fluorescent lights (including compact fluorescent lights), and light-emitting diodes. Incandescent bulbs are the “normal” type (though they are becoming less so): They light up when an electric current runs through a thin wire made of tungsten, which heats up. The wattage of an incandescent is a measure of how much power it consumes, and most of that power goes to heat, not light, which is why you can burn your hand if you touch a bulb that’s been on any length of time. Because of the wasteful nature of that kind of bulb, a lot of people have made the switch to compact fluorescent lights (CFLs), which don’t run hot and use a lot less power for the same amount of light. And they work by using quantum mechanics!


Of course even incandescent bulbs are quantum-mechanical underneath: after all, everything Continue reading

Making Light in Electronics

By DXS Physics Editor Matthew Francis 

A while back, I wrote about one of the most common ways of making electric light: fluorescent bulbs. Understanding fluorescent lights requires quantum mechanics! While a lot of quantum physics seems pretty removed from our daily lives, it’s essential to most of our modern technology. In fact, reading what I’m writing requires quantum mechanics, since you are using a computer (maybe a handheld computer like an iPad or smart phone, but it’s still a computer) or a printout from a computer.

Modern electronics, including computers and phones, depend on semiconductors. Conductors (like the copper wire in power cords) let electricity flow easily, but semiconductors conduct electricity more reluctantly—but that very reluctance lets us control the flow. While they can’t sustain large currents like conductors can, we can tinker with the chemistry of semiconductors to make them conduct electricity in very precise ways. One of those ways lets semiconductor devices make light: those are known as light-emitting diodes, or LEDs.

You likely have many LEDs in your home: they’re common as indicator lights on appliances, and you might even have LED light bulbs. While they’re pretty expensive right now, the price of LED lights is getting lower all the time, and they have major advantages over both incandescent (old-style) light bulbs and fluorescents. They won’t burn out even as quickly as fluorescent lights (themselves longer-lived than incandescents), and consume less energy. Since they are based on solids rather than gases, they’re not going to break easily, either! But how do they work?

The Electrons in the Band

When I described fluorescent lights in my earlier post, I described how atoms have distinct energy levels inside them, and light is produced when electrons move between those energy levels. Fluorescent lights use gases (generally mercury vapor), so the atoms are relatively widely separated. In solids, including semiconductors, atoms are tightly packed together, forming bonds that don’t break without high pressures or temperatures. In fact, they may also share electrons with each other; a particularly dramatic example of this is in metals, where the electrons in the highest energy levels of the atoms all form a gas that surrounds the atoms. That’s why metals are such good conductors—a little push from a battery or other power source makes those electrons flow in one direction (on average at least), much as a fan creates currents in the air.

Semiconductors are a bit more complicated: their electrons are loosely bound, but still stuck to their host atoms. The way physicists understand this is something known as the band model: just like atoms have energy levels, solids have energy bands. Low energies correspond to electrons stuck to their atoms, which can’t leave; we call these closed shell electrons (for reasons that aren’t important for this particular post). Moderate energies are known as valence electrons, which stay put ordinarily, but can be persuaded to move if given the right incentive. Finally, high energies are conduction electrons, which aren’t tied to a particular atom at all; as their name suggests, they are the ones that carry electric current.

Whether a solid conducts electricity depends on its band structure, and the size of the energy barrier in between the bands, which is called a gap. Large gaps require large energies for electrons to jump them, while smaller gaps are more easily jumped. Conductors have negligible gaps between their valence and conduction bands, while insulators have huge gaps. Semiconductors lie in between; adding extra atoms to a semiconductor can make the gap smaller (a process known as “doping”, which sometimes makes describing it unintentionally funny).

Cars and Roads and Electrons

At low temperatures, semiconductors may not conduct electricity at all, since no electrons can jump the gap into the conduction band. Either warming them up a bit or applying an external electric current gives the electrons the energy they need to move into the conduction band.

I was pondering analogies about band structures to help us understand them, and thought of this one based on cars and roads. Think of closed shells as like parking spaces along a road: cars (which stand in for electrons) are stationary. Valence bands are the slow lane, which is clogged with traffic, so the cars technically can move, but don’t. The conduction bands are fast lanes: cars can really zip, but there’s a traffic barrier between the slow lane and fast lane. (That barrier is the weakest part of my analogy, so remember that we should be thinking of a barrier as something that can be traversed under some conditions but not others.)

One more complication: there are two types of semiconductors, known as n-type and p-type. In n-type, just a few electrons (cars) have access to the conduction band (fast lane) at a time, but in p-type, enough electrons get in to leave holes in the valence band. Applying a current to the semiconductor shifts another valence electron into the hole, but that leaves another hole, and so forth…so it looks like the hole is moving! In fact, physicists refer to this as “hole conduction”, which also sounds odd if you’re not used to it.

Now we’re finally ready to understand LEDs. If you join an n-type semiconductor to a p-type semiconductor, you make something known as a diode. (The prefix di- refers to the number two. If you join three semiconductors, you get a transistor of either the pnp or npn types, depending on the order you use.) The bands (lanes) don’t line up perfectly at the junction: the conduction band in the n-type is generally only slightly higher than the valence band of the p-type, so just a little nudge is needed to move electrons across. This means when they reach the junction between the materials, electrons from the n-type semiconductor can fill the holes on the p-type, which is a decrease in energy. Just as in individual atoms, moving from a higher energy level to a lower energy level makes a photon—and that’s where the LE in the D comes from!

LEDs tend to produce very pure colors, rather than the mixture of colors our eyes perceive as white light. To create LED light bulbs, generally blue LEDs are coated with a phosphorescent material, much like the kind used in fluorescent bulbs. Unlike fluorescents, though, there’s no gas involved, and less heat is lost (though there is still a little bit). Together these factors make LED light bulbs longer-lasting and more efficient even than fluorescents, though currently they are far more expensive.

Despite how common LEDs and other semiconductors are, they’re considered fairly advanced physics. But guess what: if I did my job right, you should understand LED physics now! What is often thought of as “advanced” is really everyday science, and it’s a part of how quantum mechanics (with all its electrons and fascinating interactions on the microscopic level) has helped create our modern world.

Biology Explainer: The big 4 building blocks of life–carbohydrates, fats, proteins, and nucleic acids

The short version
  • The four basic categories of molecules for building life are carbohydrates, lipids, proteins, and nucleic acids.
  • Carbohydrates serve many purposes, from energy to structure to chemical communication, as monomers or polymers.
  • Lipids, which are hydrophobic, also have different purposes, including energy storage, structure, and signaling.
  • Proteins, made of amino acids in up to four structural levels, are involved in just about every process of life.                                                                                                      
  • The nucleic acids DNA and RNA consist of four nucleotide building blocks, and each has different purposes.
The longer version
Life is so diverse and unwieldy, it may surprise you to learn that we can break it down into four basic categories of molecules. Possibly even more implausible is the fact that two of these categories of large molecules themselves break down into a surprisingly small number of building blocks. The proteins that make up all of the living things on this planet and ensure their appropriate structure and smooth function consist of only 20 different kinds of building blocks. Nucleic acids, specifically DNA, are even more basic: only four different kinds of molecules provide the materials to build the countless different genetic codes that translate into all the different walking, swimming, crawling, oozing, and/or photosynthesizing organisms that populate the third rock from the Sun.

                                                  

Big Molecules with Small Building Blocks

The functional groups, assembled into building blocks on backbones of carbon atoms, can be bonded together to yield large molecules that we classify into four basic categories. These molecules, in many different permutations, are the basis for the diversity that we see among living things. They can consist of thousands of atoms, but only a handful of different kinds of atoms form them. It’s like building apartment buildings using a small selection of different materials: bricks, mortar, iron, glass, and wood. Arranged in different ways, these few materials can yield a huge variety of structures.

We encountered functional groups and the SPHONC in Chapter 3. These components form the four categories of molecules of life. These Big Four biological molecules are carbohydrates, lipids, proteins, and nucleic acids. They can have many roles, from giving an organism structure to being involved in one of the millions of processes of living. Let’s meet each category individually and discover the basic roles of each in the structure and function of life.
Carbohydrates

You have met carbohydrates before, whether you know it or not. We refer to them casually as “sugars,” molecules made of carbon, hydrogen, and oxygen. A sugar molecule has a carbon backbone, usually five or six carbons in the ones we’ll discuss here, but it can be as few as three. Sugar molecules can link together in pairs or in chains or branching “trees,” either for structure or energy storage.

When you look on a nutrition label, you’ll see reference to “sugars.” That term includes carbohydrates that provide energy, which we get from breaking the chemical bonds in a sugar called glucose. The “sugars” on a nutrition label also include those that give structure to a plant, which we call fiber. Both are important nutrients for people.

Sugars serve many purposes. They give crunch to the cell walls of a plant or the exoskeleton of a beetle and chemical energy to the marathon runner. When attached to other molecules, like proteins or fats, they aid in communication between cells. But before we get any further into their uses, let’s talk structure.

The sugars we encounter most in basic biology have their five or six carbons linked together in a ring. There’s no need to dive deep into organic chemistry, but there are a couple of essential things to know to interpret the standard representations of these molecules.

Check out the sugars depicted in the figure. The top-left molecule, glucose, has six carbons, which have been numbered. The sugar to its right is the same glucose, with all but one “C” removed. The other five carbons are still there but are inferred using the conventions of organic chemistry: Anywhere there is a corner, there’s a carbon unless otherwise indicated. It might be a good exercise for you to add in a “C” over each corner so that you gain a good understanding of this convention. You should end up adding in five carbon symbols; the sixth is already given because that is conventionally included when it occurs outside of the ring.

On the left is a glucose with all of its carbons indicated. They’re also numbered, which is important to understand now for information that comes later. On the right is the same molecule, glucose, without the carbons indicated (except for the sixth one). Wherever there is a corner, there is a carbon, unless otherwise indicated (as with the oxygen). On the bottom left is ribose, the sugar found in RNA. The sugar on the bottom right is deoxyribose. Note that at carbon 2 (*), the ribose and deoxyribose differ by a single oxygen.

The lower left sugar in the figure is a ribose. In this depiction, the carbons, except the one outside of the ring, have not been drawn in, and they are not numbered. This is the standard way sugars are presented in texts. Can you tell how many carbons there are in this sugar? Count the corners and don’t forget the one that’s already indicated!

If you said “five,” you are right. Ribose is a pentose (pent = five) and happens to be the sugar present in ribonucleic acid, or RNA. Think to yourself what the sugar might be in deoxyribonucleic acid, or DNA. If you thought, deoxyribose, you’d be right.

The fourth sugar given in the figure is a deoxyribose. In organic chemistry, it’s not enough to know that corners indicate carbons. Each carbon also has a specific number, which becomes important in discussions of nucleic acids. Luckily, we get to keep our carbon counting pretty simple in basic biology. To count carbons, you start with the carbon to the right of the non-carbon corner of the molecule. The deoxyribose or ribose always looks to me like a little cupcake with a cherry on top. The “cherry” is an oxygen. To the right of that oxygen, we start counting carbons, so that corner to the right of the “cherry” is the first carbon. Now, keep counting. Here’s a little test: What is hanging down from carbon 2 of the deoxyribose?

If you said a hydrogen (H), you are right! Now, compare the deoxyribose to the ribose. Do you see the difference in what hangs off of the carbon 2 of each sugar? You’ll see that the carbon 2 of ribose has an –OH, rather than an H. The reason the deoxyribose is called that is because the O on the second carbon of the ribose has been removed, leaving a “deoxyed” ribose. This tiny distinction between the sugars used in DNA and RNA is significant enough in biology that we use it to distinguish the two nucleic acids.

In fact, these subtle differences in sugars mean big differences for many biological molecules. Below, you’ll find a couple of ways that apparently small changes in a sugar molecule can mean big changes in what it does. These little changes make the difference between a delicious sugar cookie and the crunchy exoskeleton of a dung beetle.

Sugar and Fuel

A marathon runner keeps fuel on hand in the form of “carbs,” or sugars. These fuels provide the marathoner’s straining body with the energy it needs to keep the muscles pumping. When we take in sugar like this, it often comes in the form of glucose molecules attached together in a polymer called starch. We are especially equipped to start breaking off individual glucose molecules the minute we start chewing on a starch.

Double X Extra: A monomer is a building block (mono = one) and a polymer is a chain of monomers. With a few dozen monomers or building blocks, we get millions of different polymers. That may sound nutty until you think of the infinity of values that can be built using only the numbers 0 through 9 as building blocks or the intricate programming that is done using only a binary code of zeros and ones in different combinations.

Our bodies then can rapidly take the single molecules, or monomers, into cells and crack open the chemical bonds to transform the energy for use. The bonds of a sugar are packed with chemical energy that we capture to build a different kind of energy-containing molecule that our muscles access easily. Most species rely on this process of capturing energy from sugars and transforming it for specific purposes.

Polysaccharides: Fuel and Form

Plants use the Sun’s energy to make their own glucose, and starch is actually a plant’s way of storing up that sugar. Potatoes, for example, are quite good at packing away tons of glucose molecules and are known to dieticians as a “starchy” vegetable. The glucose molecules in starch are packed fairly closely together. A string of sugar molecules bonded together through dehydration synthesis, as they are in starch, is a polymer called a polysaccharide (poly = many; saccharide = sugar). When the monomers of the polysaccharide are released, as when our bodies break them up, the reaction that releases them is called hydrolysis.

Double X Extra: The specific reaction that hooks one monomer to another in a covalent bond is called dehydration synthesis because in making the bond–synthesizing the larger molecule–a molecule of water is removed (dehydration). The reverse is hydrolysis (hydro = water; lysis = breaking), which breaks the covalent bond by the addition of a molecule of water.

Although plants make their own glucose and animals acquire it by eating the plants, animals can also package away the glucose they eat for later use. Animals, including humans, store glucose in a polysaccharide called glycogen, which is more branched than starch. In us, we build this energy reserve primarily in the liver and access it when our glucose levels drop.

Whether starch or glycogen, the glucose molecules that are stored are bonded together so that all of the molecules are oriented the same way. If you view the sixth carbon of the glucose to be a “carbon flag,” you’ll see in the figure that all of the glucose molecules in starch are oriented with their carbon flags on the upper left.

The orientation of monomers of glucose in polysaccharides can make a big difference in the use of the polymer. The glucoses in the molecule on the top are all oriented “up” and form starch. The glucoses in the molecule on the bottom alternate orientation to form cellulose, which is quite different in its function from starch.

Storing up sugars for fuel and using them as fuel isn’t the end of the uses of sugar. In fact, sugars serve as structural molecules in a huge variety of organisms, including fungi, bacteria, plants, and insects.

The primary structural role of a sugar is as a component of the cell wall, giving the organism support against gravity. In plants, the familiar old glucose molecule serves as one building block of the plant cell wall, but with a catch: The molecules are oriented in an alternating up-down fashion. The resulting structural sugar is called cellulose.

That simple difference in orientation means the difference between a polysaccharide as fuel for us and a polysaccharide as structure. Insects take it step further with the polysaccharide that makes up their exoskeleton, or outer shell. Once again, the building block is glucose, arranged as it is in cellulose, in an alternating conformation. But in insects, each glucose has a little extra added on, a chemical group called an N-acetyl group. This addition of a single functional group alters the use of cellulose and turns it into a structural molecule that gives bugs that special crunchy sound when you accidentally…ahem…step on them.

These variations on the simple theme of a basic carbon-ring-as-building-block occur again and again in biological systems. In addition to serving roles in structure and as fuel, sugars also play a role in function. The attachment of subtly different sugar molecules to a protein or a lipid is one way cells communicate chemically with one another in refined, regulated interactions. It’s as though the cells talk with each other using a specialized, sugar-based vocabulary. Typically, cells display these sugary messages to the outside world, making them available to other cells that can recognize the molecular language.

Lipids: The Fatty Trifecta

Starch makes for good, accessible fuel, something that we immediately attack chemically and break up for quick energy. But fats are energy that we are supposed to bank away for a good long time and break out in times of deprivation. Like sugars, fats serve several purposes, including as a dense source of energy and as a universal structural component of cell membranes everywhere.

Fats: the Good, the Bad, the Neutral

Turn again to a nutrition label, and you’ll see a few references to fats, also known as lipids. (Fats are slightly less confusing that sugars in that they have only two names.) The label may break down fats into categories, including trans fats, saturated fats, unsaturated fats, and cholesterol. You may have learned that trans fats are “bad” and that there is good cholesterol and bad cholesterol, but what does it all mean?

Let’s start with what we mean when we say saturated fat. The question is, saturated with what? There is a specific kind of dietary fat call the triglyceride. As its name implies, it has a structural motif in which something is repeated three times. That something is a chain of carbons and hydrogens, hanging off in triplicate from a head made of glycerol, as the figure shows.  Those three carbon-hydrogen chains, or fatty acids, are the “tri” in a triglyceride. Chains like this can be many carbons long.

Double X Extra: We call a fatty acid a fatty acid because it’s got a carboxylic acid attached to a fatty tail. A triglyceride consists of three of these fatty acids attached to a molecule called glycerol. Our dietary fat primarily consists of these triglycerides.

Triglycerides come in several forms. You may recall that carbon can form several different kinds of bonds, including single bonds, as with hydrogen, and double bonds, as with itself. A chain of carbon and hydrogens can have every single available carbon bond taken by a hydrogen in single covalent bond. This scenario of hydrogen saturation yields a saturated fat. The fat is saturated to its fullest with every covalent bond taken by hydrogens single bonded to the carbons.

Saturated fats have predictable characteristics. They lie flat easily and stick to each other, meaning that at room temperature, they form a dense solid. You will realize this if you find a little bit of fat on you to pinch. Does it feel pretty solid? That’s because animal fat is saturated fat. The fat on a steak is also solid at room temperature, and in fact, it takes a pretty high heat to loosen it up enough to become liquid. Animals are not the only organisms that produce saturated fat–avocados and coconuts also are known for their saturated fat content.

The top graphic above depicts a triglyceride with the glycerol, acid, and three hydrocarbon tails. The tails of this saturated fat, with every possible hydrogen space occupied, lie comparatively flat on one another, and this kind of fat is solid at room temperature. The fat on the bottom, however, is unsaturated, with bends or kinks wherever two carbons have double bonded, booting a couple of hydrogens and making this fat unsaturated, or lacking some hydrogens. Because of the space between the bumps, this fat is probably not solid at room temperature, but liquid.

You can probably now guess what an unsaturated fat is–one that has one or more hydrogens missing. Instead of single bonding with hydrogens at every available space, two or more carbons in an unsaturated fat chain will form a double bond with carbon, leaving no space for a hydrogen. Because some carbons in the chain share two pairs of electrons, they physically draw closer to one another than they do in a single bond. This tighter bonding result in a “kink” in the fatty acid chain.

In a fat with these kinks, the three fatty acids don’t lie as densely packed with each other as they do in a saturated fat. The kinks leave spaces between them. Thus, unsaturated fats are less dense than saturated fats and often will be liquid at room temperature. A good example of a liquid unsaturated fat at room temperature is canola oil.

A few decades ago, food scientists discovered that unsaturated fats could be resaturated or hydrogenated to behave more like saturated fats and have a longer shelf life. The process of hydrogenation–adding in hydrogens–yields trans fat. This kind of processed fat is now frowned upon and is being removed from many foods because of its associations with adverse health effects. If you check a food label and it lists among the ingredients “partially hydrogenated” oils, that can mean that the food contains trans fat.

Double X Extra: A triglyceride can have up to three different fatty acids attached to it. Canola oil, for example, consists primarily of oleic acid, linoleic acid, and linolenic acid, all of which are unsaturated fatty acids with 18 carbons in their chains.

Why do we take in fat anyway? Fat is a necessary nutrient for everything from our nervous systems to our circulatory health. It also, under appropriate conditions, is an excellent way to store up densely packaged energy for the times when stores are running low. We really can’t live very well without it.

Phospholipids: An Abundant Fat

You may have heard that oil and water don’t mix, and indeed, it is something you can observe for yourself. Drop a pat of butter–pure saturated fat–into a bowl of water and watch it just sit there. Even if you try mixing it with a spoon, it will just sit there. Now, drop a spoon of salt into the water and stir it a bit. The salt seems to vanish. You’ve just illustrated the difference between a water-fearing (hydrophobic) and a water-loving (hydrophilic) substance.

Generally speaking, compounds that have an unequal sharing of electrons (like ions or anything with a covalent bond between oxygen and hydrogen or nitrogen and hydrogen) will be hydrophilic. The reason is that a charge or an unequal electron sharing gives the molecule polarity that allows it to interact with water through hydrogen bonds. A fat, however, consists largely of hydrogen and carbon in those long chains. Carbon and hydrogen have roughly equivalent electronegativities, and their electron-sharing relationship is relatively nonpolar. Fat, lacking in polarity, doesn’t interact with water. As the butter demonstrated, it just sits there.

There is one exception to that little maxim about fat and water, and that exception is the phospholipid. This lipid has a special structure that makes it just right for the job it does: forming the membranes of cells. A phospholipid consists of a polar phosphate head–P and O don’t share equally–and a couple of nonpolar hydrocarbon tails, as the figure shows. If you look at the figure, you’ll see that one of the two tails has a little kick in it, thanks to a double bond between the two carbons there.

Phospholipids form a double layer and are the major structural components of cell membranes. Their bend, or kick, in one of the hydrocarbon tails helps ensure fluidity of the cell membrane. The molecules are bipolar, with hydrophilic heads for interacting with the internal and external watery environments of the cell and hydrophobic tails that help cell membranes behave as general security guards.

The kick and the bipolar (hydrophobic and hydrophilic) nature of the phospholipid make it the perfect molecule for building a cell membrane. A cell needs a watery outside to survive. It also needs a watery inside to survive. Thus, it must face the inside and outside worlds with something that interacts well with water. But it also must protect itself against unwanted intruders, providing a barrier that keeps unwanted things out and keeps necessary molecules in.

Phospholipids achieve it all. They assemble into a double layer around a cell but orient to allow interaction with the watery external and internal environments. On the layer facing the inside of the cell, the phospholipids orient their polar, hydrophilic heads to the watery inner environment and their tails away from it. On the layer to the outside of the cell, they do the same.
As the figure shows, the result is a double layer of phospholipids with each layer facing a polar, hydrophilic head to the watery environments. The tails of each layer face one another. They form a hydrophobic, fatty moat around a cell that serves as a general gatekeeper, much in the way that your skin does for you. Charged particles cannot simply slip across this fatty moat because they can’t interact with it. And to keep the fat fluid, one tail of each phospholipid has that little kick, giving the cell membrane a fluid, liquidy flow and keeping it from being solid and unforgiving at temperatures in which cells thrive.

Steroids: Here to Pump You Up?

Our final molecule in the lipid fatty trifecta is cholesterol. As you may have heard, there are a few different kinds of cholesterol, some of which we consider to be “good” and some of which is “bad.” The good cholesterol, high-density lipoprotein, or HDL, in part helps us out because it removes the bad cholesterol, low-density lipoprotein or LDL, from our blood. The presence of LDL is associated with inflammation of the lining of the blood vessels, which can lead to a variety of health problems.

But cholesterol has some other reasons for existing. One of its roles is in the maintenance of cell membrane fluidity. Cholesterol is inserted throughout the lipid bilayer and serves as a block to the fatty tails that might otherwise stick together and become a bit too solid.

Cholesterol’s other starring role as a lipid is as the starting molecule for a class of hormones we called steroids or steroid hormones. With a few snips here and additions there, cholesterol can be changed into the steroid hormones progesterone, testosterone, or estrogen. These molecules look quite similar, but they play very different roles in organisms. Testosterone, for example, generally masculinizes vertebrates (animals with backbones), while progesterone and estrogen play a role in regulating the ovulatory cycle.

Double X Extra: A hormone is a blood-borne signaling molecule. It can be lipid based, like testosterone, or short protein, like insulin.

Proteins

As you progress through learning biology, one thing will become more and more clear: Most cells function primarily as protein factories. It may surprise you to learn that proteins, which we often talk about in terms of food intake, are the fundamental molecule of many of life’s processes. Enzymes, for example, form a single broad category of proteins, but there are millions of them, each one governing a small step in the molecular pathways that are required for living.

Levels of Structure

Amino acids are the building blocks of proteins. A few amino acids strung together is called a peptide, while many many peptides linked together form a polypeptide. When many amino acids strung together interact with each other to form a properly folded molecule, we call that molecule a protein.

For a string of amino acids to ultimately fold up into an active protein, they must first be assembled in the correct order. The code for their assembly lies in the DNA, but once that code has been read and the amino acid chain built, we call that simple, unfolded chain the primary structure of the protein.

This chain can consist of hundreds of amino acids that interact all along the sequence. Some amino acids are hydrophobic and some are hydrophilic. In this context, like interacts best with like, so the hydrophobic amino acids will interact with one another, and the hydrophilic amino acids will interact together. As these contacts occur along the string of molecules, different conformations will arise in different parts of the chain. We call these different conformations along the amino acid chain the protein’s secondary structure.

Once those interactions have occurred, the protein can fold into its final, or tertiary structure and be ready to serve as an active participant in cellular processes. To achieve the tertiary structure, the amino acid chain’s secondary interactions must usually be ongoing, and the pH, temperature, and salt balance must be just right to facilitate the folding. This tertiary folding takes place through interactions of the secondary structures along the different parts of the amino acid chain.

The final product is a properly folded protein. If we could see it with the naked eye, it might look a lot like a wadded up string of pearls, but that “wadded up” look is misleading. Protein folding is a carefully regulated process that is determined at its core by the amino acids in the chain: their hydrophobicity and hydrophilicity and how they interact together.

In many instances, however, a complete protein consists of more than one amino acid chain, and the complete protein has two or more interacting strings of amino acids. A good example is hemoglobin in red blood cells. Its job is to grab oxygen and deliver it to the body’s tissues. A complete hemoglobin protein consists of four separate amino acid chains all properly folded into their tertiary structures and interacting as a single unit. In cases like this involving two or more interacting amino acid chains, we say that the final protein has a quaternary structure. Some proteins can consist of as many as a dozen interacting chains, behaving as a single protein unit.

A Plethora of Purposes

What does a protein do? Let us count the ways. Really, that’s almost impossible because proteins do just about everything. Some of them tag things. Some of them destroy things. Some of them protect. Some mark cells as “self.” Some serve as structural materials, while others are highways or motors. They aid in communication, they operate as signaling molecules, they transfer molecules and cut them up, they interact with each other in complex, interrelated pathways to build things up and break things down. They regulate genes and package DNA, and they regulate and package each other.

As described above, proteins are the final folded arrangement of a string of amino acids. One way we obtain these building blocks for the millions of proteins our bodies make is through our diet. You may hear about foods that are high in protein or people eating high-protein diets to build muscle. When we take in those proteins, we can break them apart and use the amino acids that make them up to build proteins of our own.

Nucleic Acids

How does a cell know which proteins to make? It has a code for building them, one that is especially guarded in a cellular vault in our cells called the nucleus. This code is deoxyribonucleic acid, or DNA. The cell makes a copy of this code and send it out to specialized structures that read it and build proteins based on what they read. As with any code, a typo–a mutation–can result in a message that doesn’t make as much sense. When the code gets changed, sometimes, the protein that the cell builds using that code will be changed, too.

Biohazard!The names associated with nucleic acids can be confusing because they all start with nucle-. It may seem obvious or easy now, but a brain freeze on a test could mix you up. You need to fix in your mind that the shorter term (10 letters, four syllables), nucleotide, refers to the smaller molecule, the three-part building block. The longer term (12 characters, including the space, and five syllables), nucleic acid, which is inherent in the names DNA and RNA, designates the big, long molecule.

DNA vs. RNA: A Matter of Structure

DNA and its nucleic acid cousin, ribonucleic acid, or RNA, are both made of the same kinds of building blocks. These building blocks are called nucleotides. Each nucleotide consists of three parts: a sugar (ribose for RNA and deoxyribose for DNA), a phosphate, and a nitrogenous base. In DNA, every nucleotide has identical sugars and phosphates, and in RNA, the sugar and phosphate are also the same for every nucleotide.

So what’s different? The nitrogenous bases. DNA has a set of four to use as its coding alphabet. These are the purines, adenine and guanine, and the pyrimidines, thymine and cytosine. The nucleotides are abbreviated by their initial letters as A, G, T, and C. From variations in the arrangement and number of these four molecules, all of the diversity of life arises. Just four different types of the nucleotide building blocks, and we have you, bacteria, wombats, and blue whales.

RNA is also basic at its core, consisting of only four different nucleotides. In fact, it uses three of the same nitrogenous bases as DNA–A, G, and C–but it substitutes a base called uracil (U) where DNA uses thymine. Uracil is a pyrimidine.

DNA vs. RNA: Function Wars

An interesting thing about the nitrogenous bases of the nucleotides is that they pair with each other, using hydrogen bonds, in a predictable way. An adenine will almost always bond with a thymine in DNA or a uracil in RNA, and cytosine and guanine will almost always bond with each other. This pairing capacity allows the cell to use a sequence of DNA and build either a new DNA sequence, using the old one as a template, or build an RNA sequence to make a copy of the DNA.

These two different uses of A-T/U and C-G base pairing serve two different purposes. DNA is copied into DNA usually when a cell is preparing to divide and needs two complete sets of DNA for the new cells. DNA is copied into RNA when the cell needs to send the code out of the vault so proteins can be built. The DNA stays safely where it belongs.

RNA is really a nucleic acid jack-of-all-trades. It not only serves as the copy of the DNA but also is the main component of the two types of cellular workers that read that copy and build proteins from it. At one point in this process, the three types of RNA come together in protein assembly to make sure the job is done right.


 By Emily Willingham, DXS managing editor 
This material originally appeared in similar form in Emily Willingham’s Complete Idiot’s Guide to College Biology

Why is the sky pink?


On Mars, the sky is pink during the day, shading to blue at sunset. What planet did you think I was talking about?

On Earth, the sky is blue during daytime, turning red at as the sun sinks toward night.

Scattering light

Well, it’s not quite as simple as that: if you ignore your dear sainted mother’s warning and look at the Sun, you’ll see that the sky immediately around the Sun is white, and the sky right at the horizon (if you live in a place where you can get an unobstructed view) is much paler. In between the Sun and the horizon, the sky gradually changes hue, as well as varying through the day. That’s a good clue to help us answer the question every child has asked: why is the sky blue? Or as a Martian child might ask: why is the sky pink?

First of all, light isn’t being absorbed. If you wear a blue shirt, that means the dye in the cotton (or whatever it’s made of) absorbs other colors in light, so only blue is reflected back to your eye. That’s not what’s happening in the air! Instead, light is being bounced off air molecules, a process known as scattering. Air on Earth is about 80% nitrogen, with almost all of the rest being oxygen, so those are the main molecules for us to think about.

As I discussed in my earlier article on fluorescent lights, atoms and molecules can only absorb light of certain colors, based on the laws of quantum mechanics. While oxygen and nitrogen do absorb some of the colors in sunlight, they turn right around and re-emit that light. (I’m oversimplifying slightly, but the main thing is that photons aren’t lost to the world!) However, other colors don’t just pass through atoms as though they aren’t there: they can still interact, and the way we determine how that happens is again the color.

The color of light is determined by its wavelength: how far a wave travels before it repeats itself. Wavelength is also connected to energy: short wavelengths (blue and violet light) have high energy, while long wavelengths (red light) have lower energy. When a photon (a particle of light) hits a nitrogen or oxygen molecule, it might hit one of the electrons inside the molecule. Unless the wavelength is exactly right, the photon doesn’t get absorbed and the electron doesn’t move, so all the photon can do is bounce off, like a pool ball off the rail on a billiards table. Low-energy red photons don’t change direction much after bouncing–they hit the electron too gently for that. Higher-energy blue and violet photons, on the other hand, scatter by quite a bit: they end up moving in a very different direction after hitting an electron than they moving before. This whole process is known technically as Rayleigh scattering, for the physicist John Strutt, Lord Rayleigh.

The blue color of the sky

Not every photon will hit a molecule as it passes through the atmosphere, and light from the Sun contains all the colors mixed together into white light. That means if you look directly at the Sun or the sky right around the Sun during broad daylight, what you see is mostly unscattered light, the photons that pass through the air unmolested, making both Sun and sky look white. (By the way, your body is pretty good at making sure you won’t damage your vision: your reflexes will usually twitch your eyes away before any injury happens. I still don’t recommend looking at the Sun directly for any length of time, especially with sunglasses, which can fool your reflexes into thinking everything is safer than it really is.) In other parts of the sky away from the Sun, scattering is going to be more significant.

The Sun is a long way away, so unlike a light bulb in a house, the light we get from it comes in parallel beams. If you look at a part of the sky away from the Sun, in other words, you’re seeing scattered light! Red light doesn’t get scattered much, so not much of that comes to you, but blue light does, meaning the sky appears blue to our eyes. Bingo! Since there is some green and other colors mixed in as well, the apparent color of the sky is more a blue-white than a pure blue.

(The Sun’s light doesn’t contain as much violet light as it does blue or red, so we won’t see a purple sky. It also helps that our eyes don’t respond strongly to violet light. The cone cells in our retinas are tuned to respond to blue, green, and red, so the other colors are perceived by triggering combinations of the primary cone cells.)

At sunset, light is traveling through a lot more air than it does at noon. That means every ray of light has more of a chance to scatter, removing the blue light before it reaches our eyes. What’s left is red light, making the sky at the horizon near the Sun appear red. In fact, you see more gradations of color too: moving your vision higher in the sky, you’ll note red shades into orange into yellow and so forth, but each color is less intense.

So finally: why is the Martian sky pink? The answer is dust: the surface of Mars is covered in a fine powder, more like talcum than sand. During the frequent windstorms that sweep across the planet, this dust is blown high into the air, where light (yes) scatters off of it. Since the grains are larger than air molecules, the kind of scattering is different, and tends to make the light appear red. (Actually, the sky’s “true” color is very hard to determine, since there is a lot more variation than on Earth.) When there is less dust in the atmosphere, the Martian sky is a deep blue, when the Sun’s light scatters off the carbon dioxide molecules in the air.

By DXS Physics Editor Matthew Francis

Modern Chemists

Our next installment of notable women in science brings us to chemists. Many of these women were born in the early part of the 20thcentury and forged their paths in tough times. All are still inspiring others today. Presented in no particular order:

Catherine Clarke Fenselau is a pioneer in mass spectrometryBorn in 1939, her interested in science was apparent before her 10th grade. She was encouraged to attend a women’s college, which at the time gave what she called “a special opportunity for serious-minded young women.” She graduated from Bryn Mawr with her A.B. in chemistry in 1961. Her graduate work at Stanford introduced her to the technology she would become known for, receiving her Ph.D. in analytical chemistry in 1965. Dr. Fenselau and her husband took positions at the Johns Hopkins University Medical School, at which time she had two sons. Johns Hopkins was under a mandate to accept female students and have female faculty at the time. Dr. Fenselau was made aware of the disparity of the treatment of male and female faculty, when in the 1970s the equal opportunity laws came into effect and she received an unexplained 25% raise. Her research resided in mass spectrometry, specifically in its use in biology. She became known as an anti-cancer researcher. Dr. Fenselau spoke often to chemists about feminism and goals, such as equal pay, opening closed career opportunities to women, and achieving the bonuses often only awarded to men. She has worked as an editor on several scientific journals. Some of her awards include the Garvan Medal, Maryland Chemist Award, and NIH Merit Award. Having  proper help at work and at home, and having supportive mentors and spouse has helped her achieve her success.

Elizabeth Amy Kreiser Weisburger is considered a real-lifemedical sleuth. Born in 1924, Dr. Weisburger was one of 10 children and schooled at home for her early education. She received her B.S. in chemistry, cum laude, Phi Alpha Epsilon from Lebanon Valley College. She received her Ph.D. in organic chemistry in 1947 from the University of Cincinnati. She married and had three children. Her research has caused her to be proclaimed a pioneer in the field of chemical carcinogenesis. She balanced her busy life of working at the NCI, committee work, giving lectures, attending meetings, writing and reviewing papers while caring for children with the aid of housekeepers and nursery childcare. Some of her awards include the Garvan Medal and the HillebrandPrize. Her life philosophy is summed up with “Don’t take life so seriously; you’ll never get out of it alive.”

Helen M. Free, photo from the ACS
Helen M. Free is a major contributor to science and science education. Born in 1923, Ms. Free attended the College of Wooster, graduating with honors and a B.S. in 1944. In 1978, she earned a M.A. from Central Michigan University. In the meantime, she worked as a chemist at Miles Laboratories. She developed clinical effective and easy to use laboratory tests. She worked her way up through the company and also held an adjunct professor position at Indiana University, South Bend. Ms. Free has used her time to be active in professional societies and has served as president for the American Association for Clinical Chemistry and the American Chemical Society. Her awards include the Garvan Medal, a Distinguished Alumni Award from Wooster, and is the first recipient ofthe Public Outreach Award bearing her name.

Jeanette Grasselli Brown is an industry researcher and director. Born in 1929, she graduated summa cum laudewith her B.S. from Ohio University in 1950 and received her M.S. in 1958 from Western Reserve University. She worked at Standard Oil of Ohio (now BP of America), and became the first woman director of corporate research there. She has received numerous awards including the Garvan Medal, Ohio Women’s Hall of Fame, and the Fisher Award in Analytical Chemistry. She has published 75 papers in scientific journals, written 9 books, and received 7 honorary Doctorate of Science degrees. She is an activist for the future of women in science.

Jean’ne Marie Shreeve is an important fluorine chemist. Born in 1933, she encountered sexism through her mother’s inability to be employed despite her training as a schoolteacher. Dr. Shreeve graduated with a B.A. from Montana State University in 1953, followed by an M.S. in 1956 from the University of Minnesota, and a Ph.D. in inorganic chemistry in 1961 from the University of Washington. After graduating, she worked her way through the professorial ranks at the University of Idaho. Besides her own research, Dr. Shreeve has devoted herself to educating other chemists. Some of her awards include U.S. Ramsey Fellow, Alfred P. Sloan Fellow, and Garvan Medal.

Joyce Jacobon Kaufman by Smithsonian Institution 
Joyce Jacobson Kaufman is distinguished in many fields. Born in 1929, she was reading before the age of 2 and was a voracious reader as a child. This led to her reading the biography of Marie Curie, which inspired her to be a chemist. Dr. Kaufman received her B.S., M.A., and Ph.D. in physical chemistry from Johns Hopkins University in 1949, 1959, and 1960, respectively. She married and had a daughter. Her research in the application of quantum mechanics to chemistry, biology, and medicine led to her renown in several fields. She has also spent much time in service positions. Her awards include the Martin Company Gold Medal for Outstanding Scientific Accomplishments (received 3 times), the Garvan Medal, and honored as one of ten Outstanding Women in the State of Maryland.

Madeleine M. Joullie is known for elegant research and inspirational teachingBorn in 1927, her early life in Brazil was overly-protective, so her father encouraged her to attend school in the U.S.A. She received her B.Sc. from Simmons College in 1949, and her M.Sc. and Ph.D. in chemistry in 1950 and 1953, respectively, from the University of Pennsylvania. She then worked her way through the professorial ranks at the University of Pennsylvania. Initially, only the women graduate students would work with her, and they were few and far between. She has explored many research avenues over the course of her career. Her awards include the Garvan Medal, the American Cyanamid Faculty Award, the Henry HillAward, and the Lindback Award for Distinguished Teaching.

Marjorie Caserio is a researcher, educator, author, andacademic administrator. Born in 1929, she entered university with the goal of becoming a podiatrist in order to generic income. She received several rejections from colleges due to her gender, and eventually was accepted to be the only woman in her class. She received her B.S. from Chelsea College, University of London in 1950 and an M.A. and Ph.D from Bryn Mawr in 1951 and 1956. Dr. Caserio is co-author of one of the most popular organic chemistry textbooks in the chemistry during the 1960s and 1970s. Her awards include the Garvan Medal and John S. Guggenheim Foundation Fellow.

Mary Lowe Good has won several awards and is a public servant. Born in 1931, she was supported in her aspirations by her parents. She received her B.S. in 1950 from the University of Central Arkansas, which was then the Arkansas State Teachers College. She went on to receive her M.S. and Ph.D. in inorganic and radiochemistry from the University of Arkansas in 1953 and 1955. Her career began in academic, but an appointment to the National Science Foundation by President Carter changed the course of her career. She served the International Union of Pure and Applied Chemistry, and president of the American Chemical Society and Zonta International Foundation. Some of her awards include Garvan Medal, CharlesLathrop Parsons Award, and 18 honorary doctorates.

Ruth Mary Roan Benerito is an academic and government scientistBorn in 1916, she began college at the age of 15 at Sophie Newcomb College, the women’s college of Tulane and received her B.S. in 1935. She received her M.S. from Tulane University in 1938, which she worked half-time while working another job at the same time. She taught at Tulane and its colleges before going to the University of Chicago to get her Ph.D. in 1948 in physical chemistry, again working on a part-time basis. Her career oscillated between academia and industry, earning her a large number of awards, including the Federal Women’s Award, the Southern Chemist Award, and inducted as a Fellow into the American Institute of Chemists and Iota Sigma Pi.  

Awards

The Garvan Medal is an award from the American Chemical Society to recognize distinguished service to chemistry by women chemists.

The Maryland Chemist Award recognizes and honors its members for outstanding achievement in the fields of chemistry.

The NIH Merit Award is a symbol of scientific achievement in the research community.

The Hillebrand Prize is awarded for original contributions to the science of chemistry.

The Distinguished Alumni Award from Wooster is presented annually to alumni who have distinguished themselves in one of more of the following area: professional career; service to humanity; and service to Wooster.

Helen M. Free Award recognizes outstanding achievements in the field of public outreach. 

Ohio Women’s Hall of Fame provides public recognition of contributions made to the growth and progress of Ohio and the nation.
The Fisher Award in Analytical Chemistry recognizes outstanding contributions to the field of analytical chemistry.

U.S. Ramsey Fellow is no longer offered.

Alfred P. Sloan Fellow is awarded to scientists and scholars of outstanding promise.

Outstanding Women in the State of Maryland awards women under the age of 40 for their achievements already made in an early career. 

The American Cyanamid Faculty Award  

The Henry Hill Award recognizes distinguished service to professionalism. 


John S. Guggenheim Foundation Fellow is awarded for demonstrating outstanding scholarship.

Charles Lathrop Parsons Award recognizes outstanding public service. 



The American Institute of Chemists advances the chemical sciences by establishing high professional standards of practice and to emphasize the professional, ethical, economic, and social status of its members for the benefit of society as a whole.

Iota Sigma Pi is a national honor society for women in chemistry.

Much of the information for this post came from the book Notable Women in the Physical Sciences: A Biographical Dictionary edited by Benjamin F. Shearer and Barbara S. Shearer. 

Adrienne M Roehrich, Double X Science Chemistry Editor

From alchemist to chemist: What kind of chemistry is that?


Figure 1: The Alchemist Discovering Phosphorus

What does the word chemistry  mean to you? For many, it was a class in high school or college to get through. In these introductory courses, called general chemistry, one gets a mix of all the flavors of chemistry – but the flavors are very different. To those who hear the calling of chemistry, it isn’t just any chemistry that will do. Some courses are more interesting to them than others. 

Many instructors start their general chemistry course with a history, introducing alchemy. Alchemy is considered to be the process by which to turn [name item of your choice] into gold. Alchemists were chemists by accident in that they performed many chemical reactions in their quests, discovering a number of elements in the process - embodied by Hennig Brandt’s discovery of phosphorus from the refinement of urine.
Alchemy relates to all the fields of chemistry. In perhaps the most famous of alchemy pictures, that by Joseph Wright of Derby entitled “The Alchemist Discovering Phosphorus,” the alchemist is kneeling by a very large round bottom flask. For many in modern chemistry, the round bottom flask signifies hours in the organic chemistrylaboratory mixing chemicals together to create something new.


Organic chemistry is the “branch of chemistry that deals with the structure, properties, and reactions of compounds that contain carbon” according to the American Chemical Association (ACS). Organic chemistry is the largest of chemistry fields in terms of number of people working in it. Organic chemists strive to make new compounds, usually to improve upon an existing one for a purpose and the field is often thought of in terms of synthesis applications.


The actual process of converting urine to phosphorus generally falls along the lines of inorganic chemical reactions. The form of phosphorus in urine is in the chemical sodium phosphate (Na3PO43-). Heating phosphates along with the organic products also in urine will form carbon monoxide (CO) and elemental phosphorus (P). The sodium phosphate, carbon monoxide, and elemental phosphate are all inorganic chemicals, falling under the field of inorganic chemistry.


Inorganic chemistry is “concerned with the properties and reactivity of all chemical elements,” according to UC-Davis chemwiki. While organic chemistry requires the presence of carbon in a specific type of bond, inorganic chemistry involves all the elements present in the periodic table. Inorganic chemistry delves into theories surrounding the bonding of metals to molecules and the shapes of molecules themselves.


Figure 2: Components of Urine
While the process of collecting phosphorus from urine requires organic and inorganic chemical reactions, the process of making the products in urine is biochemistry. Note in figure 2 that the primary product in urine is urea.


For students of biochemistry, images of the urea cycle (aka the Krebs cycle) are well known. According to the ACS, biochemistry is “the study of the structure, composition, and chemical reactions of substances in living systems.” Besides the chemical cycles to produce and use up necessary chemicals in biology, biochemistry encompasses protein structure and function (including enzymes), nucleic acids such as DNA, and biosynthesis.


As the alchemist turned urine to phosphorus, he added heat. The addition of heat to a reaction involves thermodynamics, a subsection of physical chemistry. If heat hadn’t been added, the reaction products would have been kinetic, which is another subsection of physical chemistry.


In a suite of physical chemistry courses, a student would also take quantum mechanics, rounding out the aim of physical chemists, which is to “develop a fundamental understanding at the molecular and atomic level of how materials behave and how chemical reactions occur,” according to the ACS. Physical chemists work by applying physics and math to the problems that chemists, biologists, and engineers study.


The alchemists who took exact measurements of their reactants and products, using quantitative methods, employed analytical chemistry. Presumably, the alchemists did this because every ounce of gold was precious, and they wanted to know how much substance they started with to produce the coveted metal.


Analytical chemistry focuses on obtaining and processing information about the composition and structure of matter. There are so-called wet lab ways to determine these quantities that often been employed. However, most analytical labs consist of the precision instrumentation that you may have seen on forensic crime shows, such as a mass spec, short for mass spectrometer, a frequent player on CSI.


While the alchemists were only trying to produce a substance to enrich pockets, they ultimately led to a rich science with several subfields, each with a trail leading from the practice of alchemy.

Adrienne M Roehrich, Double X Science Chemistry Editor
References:

(1734-1797), Joseph Wright of Derby. “English: The Alchemist Discovering Phosphorus or the Alchemist in Search of the Philosophers Stone.” Derby Museum and Art Gallery, Derby, U.K., 1771.

Lawton, Graham. “Pee-Cycling.” New Scientist, 20 December 2006 2006.

Weeks, Mary Elvira. “The Discovery of the Elements. Xxi. Supplementary Note on the Discovery of Phosphorus.” Journal of Chemical Education 10, no. 5 (1933/05/01 1933): 302.