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

Einstein's most famous equation, sort of. This is the transcription of the chalkboard from a public talk Einstein gave in Pittsburgh in 1934. (Credit: Dwight Vincent and David Topper)

Did Einstein write his most famous equation? Does it matter?

Why all the fuss about E = m c2?

By Matthew R. Francis

Albert Einstein in Pittsburgh, 1934. (Credit: Pittsburgh Sun-Telegraph/Dwight Vincent and David Topper)

The association is strong in our minds: Albert Einstein. Genius. Crazy hair. E = m c2. Maybe many people don’t know what else Einstein did, but they know about the hair and that equation. They may think he flunked math in school (wrong, though he did have conflicts with some teachers), that he was a ladies’ man (true, he had numerous affairs during both of his marriages), and that he was the smartest man who ever lived (debatable, though he certainly is one of the central figures in 20th century physics). Rarely, people will remember that he was a passionate antiracist and advocate for world government as a way of bringing peace.

Obviously whole books have been written about Einstein and E = m c2, but a blog post at io9 caught my attention recently. The post (by George Dvorsky) itself looked back to a scholarly paper by David Topper and Dwight Vincent [1], which reconstructed a public lecture Einstein gave in 1934. (All numbers in square brackets [#] are citations to the references at the end of this post.) This lecture was one of many Einstein presented over the decades, but as Topper and Vincent wrote, “As far as we know [the photograph] is the only extant picture with Einstein and his famous equation.”

Well, kind of. The photograph is really blurry, and the authors had to reconstruct what was written because you can’t actually see any of the equations clearly. Even in the reconstructed version (reproduced below)…there’s no E = m c2. Instead, as I highlighted in the image, the equation is E0 = m. Einstein set the speed of light – usually written as a very large number like 300 million meters per second, or 186,000 miles per second – equal to 1 in his chalkboard talk.

Einstein's most famous equation, sort of. This is the transcription of the chalkboard from a public talk Einstein gave in Pittsburgh in 1934. (Credit: Dwight Vincent and David Topper)

Einstein’s most famous equation, sort of. This is the transcription of the chalkboard from a public talk Einstein gave in Pittsburgh in 1934. (Credit: Dwight Vincent and David Topper)

What’s the meaning of this?

It is customary to express the equivalence of mass and energy (though somewhat inexactly) by the formula E = mc2, in which c represents the velocity of light, about 186,000 miles per second. E is the energy that is contained in a stationary body; m is its mass. The energy that belongs to the mass m is equal to this mass, multiplied by the square of the enormous speed of light – which is to say, a vast amount of energy for every unit of mass. –Albert Einstein [2]

Before I explain why it isn’t a big deal to modify an equation the way Einstein did, it’s good to remember what E = m c2 means. The symbols are simple, but they encode some deep knowledge. E is energy; while colloquially that term gets used for a lot of different things, in physics it’s a measure of the ability of a system to do things. High energy means fast motion, or the ability to make things move fast, or the ability to punch through barriers. Mass m, on the other hand, is a measure of inertia: how hard it is to change an object’s motion. If you kick a rock on the Moon, it will fly farther than it would on Earth, but it’ll hurt your foot just as much – it has the same mass and therefore inertia both places. Finally, c is the speed of light, a fundamental constant of nature. The speed of light is the same for an object of any mass, moving at any velocity.

Mass and energy aren’t independent, even without relativity involved. If you have a heavy car and a light car driving at the same speed, the more massive vehicle carries more energy, in addition to taking more oomph to start or stop it moving. However, E = m c2 means that even if a mass isn’t moving, it has an irreducible amount of energy. Because the speed of light is a big number, and the square of a big number is huge, even a small amount of mass possesses a lot of energy.

The implications of E = m c2 are far-reaching. When a particle of matter and its antimatter partner meet – say, an electron and a positron – they mutually annihilate, turning all of their mass into energy in the form of gamma rays. The process also works in reverse: under certain circumstances, if you have enough excess energy in a collision, you can create new particle-antiparticle pairs. For this reason, physicists often write the mass of a particle in units of energy: the minimum energy required to make it. That’s why we say the Higgs boson mass is 126 GeV – 126 billion electron-volts, where 1 electron-volt is the energy gained by an electron moved by 1 volt of electricity. For comparison, an electron’s mass is about 511 thousand electron-volts, and a proton is 938 million electron-volts.

In our ordinary units the velocity of light is not unity, and a rather artificial distinction between mass and energy is introduced. They are measured by different units, and energy E has a mass E/C2 where C is the velocity of light in the units used. But it seems very probable that mass and energy are two ways of measuring what is essentially the same thing, in the same sense that the parallax and distance of a star are two ways of expressing the same property of location. –Arthur Eddington [3]

Another side of the equation E = m c2 appears when we probe the structure of atomic nuclei. An atomic nucleus is built of protons and neutrons, but the total nuclear mass is different than the sum of the masses of the constituent particles: part of the mass is converted into binding energy to hold everything together. The case is even more dramatic for protons and neutrons themselves, which are made of smaller particles knowns as quarks – but the total mass of the quarks is much smaller than the proton or neutron mass. The extra mass comes from the strong nuclear force gluing the particles together. (In fact, the binding particles are known as gluons for that reason, but that’s a story for another day.)

A brief history of an idea

The E0 = m version of the equation Einstein used in his chalk-talk might seem like it’s a completely different thing. You might be surprised to know that he almost never used the famous form of his own discovery: He preferred either the chalkboard version or the form m = E/c2. In fact, in his first scientific paper on the subject (which was also his second paper on relativity), he wrote [4]:

If a body gives off the energy L in the form of radiation, its mass diminishes by L/c2. The fact that the energy withdrawn from the body becomes energy of radiation evidently makes no difference, so that we are led to the more general conclusion that … the mass of a body is a measure of its energy-content …

In other words, he originally used L for energy instead of E. However, it’s equally obvious that the meaning of E = m c2 is present in the paper. Equations, like sentences in English, can often be written in many different ways and still convey the same meaning. By 1911 (possibly earlier), Einstein was using E for energy [5], but we can use E or L or U for energy, as long as we make it clear that’s what we’re doing.

The same idea goes for setting c equal to one. Many of us are familiar with the concept of space-time: that time is joined with space (thanks to the fact that the speed of light is the same, no matter who measures it). We see the blurring of the boundary between space and time when astronomers speak of light-years: the distance light travels in one year. Because c – and therefore c2 – is a fixed number, it means the difference between mass and energy is more like the difference between pounds and kilograms: one is reachable from the other by a simple calculation. Many physicists, including me, love to use c = 1 because it makes equations much easier to write.

In fact, physicists (including Einstein) rarely use E = m c2 or even m = E/c2 directly. When you study relativity, you find those equations are specific forms of more general expressions and concepts. To wit: The energy of a particle is only proportional to its mass if you take the measurement while moving at the same speed as the particle. Physical quantities in relativity are measured relative to their state of motion – hence the name.

That’s the reason I don’t care that we don’t have a photo of Einstein with his most famous equation, or that he didn’t write it in its familiar form in the chalk-talk. The meaning of the equation doesn’t depend on its form; its usefulness doesn’t derive from Einstein’s way of writing it, or even from Einstein writing it.

A small representative sample of my relativity books, with my cats Pascal and Harriet for scale.

A small representative sample of my relativity books, with my cats Pascal and Harriet for scale.

Even more: Einstein is not the last authority on relativity, but the first. I counted 64 books on my shelves that deal with the theory of relativity somewhere in their pages, and it’s possible I missed a few. The earliest copyright is 1916 [6]; the most recent are 2012, more than 50 years after Einstein’s death. The level runs from popular science books (such as a couple of biographies) up to graduate-level textbooks. Admittedly, the discussion of relativity may not take up much space in many of those books – the astronomy and math books in particular – but the truth is that relativity permeates modern physics. Like vanilla in a cake, it flavors many branches of physics subtly; in its absence, things just aren’t the same.

References

  1. David Topper and Dwight Vincent, Einstein’s 1934 two-blackboard derivation of energy-mass equivalence. American Journal of Physics75 (2007), 978. DOI: 10.1119/1.2772277 . Also available freely in PDF format.
  2. Albert Einstein, E = mc2. Science Illustrated (April 1946). Republished in Ideas and Opinions (Bonanza, 1954).
  3. Arthur Eddington, Space, Time, and Gravitation (Cambridge University Press, 1920).
  4. Albert Einstein, Does the inertia of a body depend upon its energy-content? (translated from Ist die Trägheit eines Körpers von seinem Energiegehalt abhängig?). Annalen der Physic17 (1905). Republished in the collection of papers titled The Principle of Relativity (Dover Books, 1953).
  5. Albert Einstein, On the influence of gravitation on the propagation of light (translated from Über den Einfluss der Schwerkraft auf die Ausbreitung des Lichtes). Annalen der Physic35 (1911). Republished in The Principle of Relativity.
  6. Albert Einstein, Relativity: The Special and the General Theory (1916; English translation published by Crown Books, 1961).

Why Are Snowflakes Always Six-Sided?

(Today’s offering is a guest post by engineer Linda Gaines.)

It’s a well-known fact that all snowflakes have six sides. Or at least I thought it was. Why Google is unable to Google that fact and has on at least two occasions created a Doodle with an eight-sided snowflake is a mystery. What’s less mysterious is how scientists can be so sure that all snowflakes have six sides. Have we examined all snowflakes? No, of course not, but the explanation lies in two words: hydrogen bonding. Thanks to the intermolecular force of hydrogen bonding, all snowflakes have six sides, and hydrogen bonding also makes life as we know it possible. Now that’s an important bond.

You can’t really understand hydrogen bonding, though, without understanding why water molecules are arranged like they are. Water seems like a simple enough molecule. It consists of one oxygen atom with two hydrogen atoms bonded to it. The hydrogen atoms bond to the oxygen atom at a distance of exactly 104.5 degrees from each other (1). Why that particular angle?

An oxygen atom has a total of eight electrons. Two of them take up all the available spots in the shell closest to the atom’s nucleus. The remaining six electrons are relegated to the atom’s outermost (or valence) electronic shell. But this shell can actually hold eight electrons, so two spots are open. A hydrogen atom has one electron on its only electronic shell, and since that shell holds two electrons, it’s got room for one more.

Because oxygen has two available spaces and hydrogen has one, oxygen can share that space with two hydrogen atoms. Both hydrogen atoms share their single electron with the oxygen, and the oxygen shares an electrons with each of the hydrogen atoms. The remaining four of the oxygen’s electrons aren’t a part of this sharing arrangement, though. Electrons kick around in pairs, so these four non-sharing electrons form two pairs.

With these two pairs sitting alone and the other two electrons each sharing with a hydrogen, a water molecule has a tetrahedron (or three-sided pyramid) shape with four attachments emerging from the oxygen nucleus. Two of those attachments are electron clouds containing two electrons each (the pairs), and the other two attachments are hydrogen atoms with two electrons moving between the oxygen and hydrogen orbits. In a true tetrahedron, the attachments would all be 109.5 degrees from each other. With the water molecule, though, the hydrogen atoms are 104.5 degrees from each other because the two paired-electron clouds are grabby with space and force the electrons shared with the hydrogen atoms a little closer together.

So we’ve learned that the hydrogen and oxygen form a covalent bond, which means they share their electrons. What I haven’t told you is that the oxygen is very grabby with that electron, so the sharing isn’t exactly equal. The oxygen has a stronger hold on the electron and is pulling that negative charge closer to it and away from the hydrogens. What results is a slightly negative oxygen and slightly positive hydrogens. The oxygen actually has two areas of negativity, right across from where it’s bonded with each hydrogen. Water molecules can use these areas of slight charge to form a fairly strong bond with other molecules, a bond called a hydrogen bond. While not every molecule containing hydrogen can form this kind of bond with other molecules, molecules in which hydrogen is in this unequal sharing situation will be able to.

In the case of water molecules bonding to other water molecules, the two slightly negative areas of the oxygen can each bond with a slightly positive hydrogen from another water molecule. When all four slightly charged areas have each bonded with another water molecule via hydrogen bonding, the result is a tetrahedral (four-sided pyramid) shape.

These bonds make water an unusual substance. When the temperature drops and water starts to solidify, the hydrogen bonding becomes very important. The hydrogen bonding dictates the shape of the ice crystals. You’ve learned that each water molecule is linked to four other water molecules in a tetrahedral arrangement. 

As the water freezes, these tetrahedrons come closer together and crystallize into a six-ring or hexagonal structure. Look at the image to see how this happens. Each point on the hexagon is an oxygen atom, and each side is a hydrogen bonded to one oxygen. As the water approaches freezing temperature, the water molecules continue to crystallize in this tetrahedral arrangement.

But water does something unlike most substances. As it nears freezing, instead of continuing to contract, it expands slightly from about 4 degrees to 0 degrees Celsius as the motion of the molecules slows with the cold, and the hydrogen bonds extend the molecules to their fullest distance from each other. It’s like a ring of people holding hands, elbows bent, and then gradually straightening their arms to the fullest extension so that they’re at the greatest distance from each other. When water molecules do this, the hexagonal structure expands into a larger and larger hexagonal structure.

The snowflake, with its six sides, is what results from this process: It is a large, gorgeous ice crystal. Ice crystals are like mineral rock crystals. The macroscopic (large) shape you see is dictated by the microscopic, molecular crystalline structure. Ice has a hexagonal crystalline structure, so a snowflake has a hexagonal structure. Sodium chloride, aka table salt, has a cubic molecular structure, so the salt crystals you shake on your food have a cubic shape.

It’s interesting that hydrogen bonding causes snowflakes to be six sided (are you listening, Google?), but it carries far greater consequences than beautiful snowflakes. Breaking those hydrogen bonds apart so that water can transform from liquid to gas takes a lot of heat, so the boiling point of water is much higher than it is for other, similar molecules. Based on similar molecules, water’s boiling point should be about -80 degrees Celsius (-176 degrees Fahrenheit) (!) instead of the 100 degrees Celsius (212 degrees Fahrenheit) it really is (1).

And then there’s the fact that ice floats, which means that the solid form of water is less dense than the liquid form. It is highly unusual for the solid form of a substance to be less dense than its liquid. But because those hydrogen bonds force water into a pretty open, hexagonal crystalline structure as the temperature nears 0 degrees Celsius, molecules are not packed as closely together as they are at warmer temperatures.


Think of those people holding hands, stiff-arming each other as far apart as possible. If they all started slam dancing, their handholds would break, and they could get closer to one another. Water molecules are a bit like that when the temperature goes above 4 degrees Celsius. When ice melts, some of the hydrogen bonds break, and the water molecules can be closer together. The far-apart water molecules in ice form a less-dense substance than the close-together molecules of liquid water, so ice floats in liquid water.

This property of water is integral to life on Earth. When a freshwater lake starts to freeze, the ice floats on the top, insulating the water below and preventing it from freezing. The fish, plants, and other life in the lake remain alive beneath the protective and insulating icy layer. If ice sank instead, over periods of deep freeze during its 4.5 billion year existence, this blue planet would have developed an icy, inhospitable core. Instead, the fact that ice floats meant that Earth was a perfect incubator for life in its oceans. All because oxygen is just a little bit grabby with electrons.

(1) Petrucci, Ralph H. (1989) General Chemistry (Fifth Edition). New York: MacMillan.

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

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

Life and science challenges: flames, Hawkeye, the needle and the damage done

(source)

Of Heroin, Honorable Mentions, and Hawkeye: A day I will never forget

By Double X Science Biology Editor Jeanne Garbarino


“I look forward to seeing you in 3 months when you will be a whole person again.”

Those were my parting words to a special person in my life who was embarking on an undoubtedly difficult journey toward sobriety.  It was only 7:45am on Friday, June 1st, but already I had learned that the strings from a bikini top make a good tourniquet, and I actually held the syringe that, only moments before, contained a bolus of heroin.  I am still trying to believe that this really was the last time.  

As I attempted to wrap my head around what was happening, I remembered a description of a heroin high as told to me by a former addict.  According to this person, being on heroin feels like you’ve been swaddled in a warm blanket, and gently rocked by a loving mother, except the loving mother was actually the devil.  

Though I could never really understand what it feels like to be hooked on heroin, this helped me make some sense of it.  But, as much as I wanted to be sympathetic, I also wanted to grab my friend by the shoulders and scream.  “Why have you done this to yourself?  Why have you done this to us?”  It has truly been a difficult time, watching this person struggle.  And finding out that I can’t control any of it was probably the hardest lesson I’ve ever learned.

Still, life must go on.

I took a few deep breaths, which helped to quiet the tremble, and began to gather my thoughts.  What was it that I had to do today?  As if I flipped some switch, I began to plan out my day – renew my parking permit, finish that Western blot, read that thesis, and get that new post up on the site.  

Then, around 8:15am, I received an unexpected phone call.  It was Liz Bass from the Center of Communicating Science at Stony Brook University.  She was calling to see if I could make Alan Alda’s World Science Festival discussion about the Flame Challenge, which was to occur at 4pm that afternoon.  Not really knowing what was in store, I quickly accepted (um, hello, Alan Alda).  A second phone call about 20 minutes later informed me that I would be joining Alan on stage.  Was this really happening?  In about 30 minutes time, I went from despair to elation.  I also went to the store to buy a skirt since I was already in transit to my lab (and was dressed like a “scientist”).

As I sat on the train, I began to reflect.  Much of my free time during the month of March was dedicated to producing an entry to Alan Alda’s Flame Challenge contest, which, in an effort to raise science communication awareness, asked scientists from all over the world to define a flame to an 11 year old.  Because I enjoy working on a team, I asked my fellow scicommies, Deborah Berebichez and Perrin Ireland, to join me on this endeavor (three times the brain power!).  For several weeks, we worked on the script, and regularly discussed our progress during late night Google hangouts (which is a fantastic way to collaborate).  This was mostly due to the fact that we all have day jobs and obligations outside of work.  Luckily for me, Debbie and Perrin were willing to meet at a time that coincided after my children’s bedtime routine.

This experience was truly fun and rewarding.  Each of us has a certain set of strengths, which when combined, seemed to just synergize.  We literally examined every word in the script to make sure that it was clear, concise, and hopefully captivating.  Furthermore, we wanted to make sure that it was something an 11-year-old would both learn from and enjoy.

But, we did labor over one particular issue, and that was our use of the Bohr Model to represent an atom.  While this model might be commonplace in many classroom textbooks, scientists now know that electrons exist in orbitals, also known as electron clouds, and the calculations to determine the exact location(s) of an electron are based on probability.  Clearly, this was something very different than stating that electrons simply orbit around a nucleus.  

The analogy that electrons travel around the nucleus in the same way that planets travel around the sun is downright inaccurate.  However, this is an analogy that is still commonly used and is, in my opinion, a great example of how we sacrifice accuracy for simplicity.  I believe that this is the greatest challenge for a science communicator.  

As we talked through this issue, we tried to not lose site of the actual mission, which was to explain a flame to an 11 year old.  Would it help our story to break down the currently accepted atomic theory or would it detract from it?  In the end, we decided to keep our atomic structure simple, but noted that it was a simplified version of an atom.  We figured that by having this little disclaimer, it would inform our audience that there is more to it that what we showed, and maybe it would lead them down a road of scientific inquiry.  

Perhaps it was this attention to detail that landed our Flame Challenge video a spot in the top 15 entries (FYI there were close to 900 entries).   Or perhaps it was because our entry was cute and artistic.  Whatever the reason, we proudly accepted our honorable mention, and I was looking forward to discussing our video with the man himself.

Getting back to Friday, June 1st.   I arrived at the Paley Center for Media around 3:30pm (in a new skirt) and was immediately brought up to the 11th floor and into the green room of Alan Alda.  There, I met my fellow awardees (a combination of finalists and honorable mentions), and of course Alan Alda, who was fantastically charming and funny.  We all sat, around an old table, on which was a lovely array of cheese, nuts, banana chips, and get this, Swedish fish!  I don’t know what it was about the Swedish fish, but seeing this candy helped calm my nerves.

Alan helped us all to break the ice, and discussed his plans for the event.  Apparently we would be leading a panel discussion, and I would be on that panel.  On a stage.  In front of a very large audience.  And it was to be webcasted.  So I popped a few of those Swedish fish and told myself to not be nervous.

As my jaw worked to chew those sticky sweet candies, I couldn’t help but think about when I was a kid and how I used to sit with my dad and watch M*A*S*H.  I never would have believed you if you told me that I was going to be hanging out with Hawkeye when I was older.  But, there he was, telling us about the birth of the Flame Challenge.  I was tempted to ask him where Corporal Klinger is these days, but decided that my time would be better spent getting the plan for the panel firmed up in my brain.

After some quick chitchat, we were asked to make our way to the auditorium.  Seating was charted and mics were checked and around 4pm, it all began.  About an hour into it, we were asked to come on stage.  Each of our entries were highlighted, followed by a chance speak our piece.  Add in some Q&A from the audience and the panel discussion was complete.  A hearty round of applause later, I found myself getting whisked away for pictures.  

When the dust began to settle, I grabbed a beer and started to decompress.  I just couldn’t believe how this day turned out, especially given its start.  The stresses my family and I have been dealing with have certainly taken its toll on all of us, and I am grateful for that little dose of Hawkeye to help lighten things up.  I’m not sure if I will ever experience a day like that again, but that’s ok with me.  

Mariette DiChristina

Mariette DiChristina is editor in chief of Scientific American.

[Ed. note: This interview is the second installment in our new series, Double Xpression: Profiles of Women into Science. The focus of these profiles is how women in science express themselves in ways that aren’t necessarily scientific, how their ways of expression inform their scientific activities and vice-versa, and the reactions they encounter.]

Today’s profile is an interview with Mariette DiChristina, editor in chief, Scientific American, who answered our questions via email with DXS Biology Editor Jeanne Garbarino. Read on to find out what a Marx Brothers movie has to do with communicating science.

                         

DXS: First, can you give me a quick overview of what your scientific background is and your current connection to science?

MD: Like most kids, I was born a scientist. What I mean is, I wanted to know how everything worked, and I wanted to learn about it firsthand. At a tag sale, for instance, I remember buying a second-hand biology book called The Body along with my second-hand Barbie for 50 cents. “Are you sure your mom is going to be OK with you buying that?” asked the concerned neighbor, eyeing the biology book.

I memorized the names and orbital periods of the planets and of dinosaurs like some kids spout baseball stats (which I could also do as a kid, by the way). We didn’t have a lot of money, so I caught my own pet fish from a nearby pond by using my little finger as a pretend worm. I scooped up my fish with an old plastic container and put it on my nightstand. If it died, I buried it and dug it up later so I could look at the bones. My proudest birthday gifts were when I got a chemistry set and a microscope with 750x. A girlfriend and I got the idea to pick up a gerbil that had a bad habit of biting fingers, just so we could get blood to squeeze on a glass slide. (She was braver than I was about being the one to get bitten.)

In middle school, I was a proud member of the Alchemists—an after-school science club—so I could do extra labs and clean the beakers and put away Bunsen burners for fun. I knew I would be a scientist when I grew up.

But somewhere during my high school courses, I came to believe that being a scientist meant I’d have to pick one narrow discipline and stick to it. I felt that I liked everything too much to do that, however. As an undergraduate, I eventually figured out that what I really wanted was to be a student of many different things for life, and then share those things I learned with others. That led me to a journalism degree. It also means that, as far as knowledge about science goes, I fit the cliché of being “an inch deep and a mile wide.”

DXS: What ways do you express yourself creatively that may not have a single thing to do with science?

MD: This one is a tough one for me to answer because I am always trying to convince people that pretty much everything they care about in the headlines actually has to do with science! In my case, I’ve also always been interested in drawing and in visuals in general. I was a pretty serious art student in high school as well, although I later decided that I didn’t have enough passion for it to make that my career choice. My interest in art partly led me to work at magazines like Scientific American and Popular Science, where the ability to storyboard an informational graphic and otherwise think visually is very helpful.

When I’m home, I really enjoy making things with my two daughters, such as helping them with crafts or scrapbooks, although I definitely spend a lot more time on planning dinners and cooking for (and with) the family than anything else. I like the puzzle solving of setting up the meals for the week during the weekend, so it’s easier for my husband to get things ready weeknights. We’re big on eating dinner together as a family every night. I like gardening and mapping out planting beds. I’m better at planting than at keeping up with tending, however, because of my intense work schedule and travel. In short, if I have free time at all, I’m enjoying it with my family. And if we’re doing some creative expression while we’re at it, great!

DXS: Do you find that your connection to science informs your creativity, even though what you do may not specifically be scientific?

MD: My connection to science informs most things that I do in one way or another. When I’m making dinner, I sometimes find myself talking about the chemistry of cooking with the girls. Especially when our daughters were smaller, if one of them had a question, I’d try to come up with ways to make finding the answer together into a kind of science adventure or project.

I suppose that since I spend most of my waking hours thinking about how best to present science to the public, it’s just a mental routine, or a lens through which I tend to view the world.

DXS: Have you encountered situations in which your expression of yourself outside the bounds of science has led to people viewing you differently–either more positively or more negatively?

MD: It’s more the other way around. I get amusing reactions from people once they find out what I do. How could I seem so normal and yet work in a field that relates to…shudder…science? An attorney friend has sometimes kidded me, saying there’s no way he can understand what’s in Scientific American, so I must be incredibly smart. I don’t feel that way at all! Anybody who has a high school degree and an interest in the topic can understand a feature article in Scientific American. Science is for everyone. And science isn’t only for people who work in labs. It’s just a rational way of looking at life. I also believe science is the engine of human prosperity. And if I sound a little evangelistic about that, well, I am.
DXS: Have you found that your non-science expression of creativity/activity/etc. has in any way informed your understanding of science or how you may talk about it or present it to others?

MD: I think it’s helpful to look to non-science areas for ideas about ways to help make science appealing, especially for people who might be intimidated by the subject. My main job is to try to make a connection for people to the science we cover in Scientific American. I once had a boss at Popular Sciencewho made all us editors take an intensive, three-day screenwriting course that culminated in the showing and exposition, scene by scene, of the structure and writing techniques of Casablanca. When I came back, he gave me a big grin and said, “So, what did you think?” I got his point about bringing narrative techniques into feature articles. Like most people, I enjoy movies and plays; now I also look at them for storytelling tips. And there are lots of creative ways to tell science stories beyond words: pictures, slide shows, videos, songs. Digital media are so flexible.

DXS: How comfortable are you expressing your femininity and in what ways? How does this expression influence people’s perception of you in, say, a scientifically oriented context?

MD: I was the oldest of three daughters raised by a single dad (my mom died when I was 12) and I was always a tomboy, playing softball through college and so on. So I can’t say I’ve ever been terribly feminine, at least in the stereotypical ways. At the same time, I’m obviously a wife and a mother who, like most parents, tries not to talk about my kids so often that it’s irritating to friends and coworkers. I once was scolded in a letter from an irritated reader after I had mentioned my kids in a “From the Editor” column about education. He wrote that if I was so interested in science education and kids, I should go back home and “bake cookies.” I laughed pretty hard at that.

DXS: Do you think that the combination of your non-science creativity and scientific-related activity shifts people’s perspectives or ideas about what a scientist or science communicator is? If you’re aware of such an influence, in what way, if any, do you use it to (for example) reach a different corner of your audience or present science in a different sort of way?

MD: I’m sure that’s true. I think personality and approach also might shift perspectives. A girlfriend of mine once called me “the friendly face of science.” I guess I smile a lot, and I like to meet people and try to get to know them. That ability—being able to make a personal connection to different people—is important for every good editor. My job, essentially, is to understand your interests well enough to make sure Scientific American is something that you’ll enjoy each day, week, month.

Increasingly, also, the audiences are different in different media, so we need to understand how to flex the approach a bit to appeal to those different audiences. In print, for instance, according to the most recent data we have from MRI, the median age of Scientific American readers is 47, with 70 percent men and 30 percent women. The picture is quite different online, where, according to Nielsen, our median age is 40 and the male/female ratio is closer to half and half, with 56.5 percent men to 43.5 percent women. You need to bring a lot of creative thinking to the task of how to make one brand serve rather different sets of people.

Fortunately, I have terrific, creative staff! And another part of the way you do that, I think, is to invite your readers in to collaborate; we’ve done a bit of that in the past year on http://www.scientificamerican.com/, and I’m looking forward to experimenting further in the coming months. Ultimately, I’d like to turn Scientific American from a magazine with an amazing 166-year tradition of being a conduit of authoritative information about science and technology into a platform where curious minds can gather and share.

DXS: If you had something you could say to the younger you about the role of expression and creativity in your chosen career path, what would you say? 

MD: I was pretty determined to do something—whatever it was—that would let me satisfy my curiosity and passion about science. I would tell younger me, who, by the way, never intended to go into magazine management: It’s just as fun, rewarding and creative to be a science writer as you suspect it might be. I’d also tell the younger me something that didn’t occur to me early enough to pull it off—that a double major in journalism and science might be a good idea. And, I would add, it’s also a good idea to take some business classes, so you’ll be better armed for dealing with the working world.


Also on Double X Science

More about Mariette DiChristina

Mariette DiChristina oversees Scientific American Continue reading