Carbon: Four Times Cooler Than Most of the Other Molecules You Know

In a first-season episode of the original Star Trek series, the Enterprise encounters a strange creature that is killing colonists on a nearby planet.  Unable to locate or electronically sense this creature, Spock posits that it may be a life form heretofore unknown to mankind: it is a silicon-based life form.  The idea of silicon-based life has permeated popular science for decades, partly because a quick analysis seems to favor this theory as possible.  This is largely due to the fact that silicon has the same binding chemistry as carbon, the molecule on which all life is based.  In reality, however, silicon just can’t compete with all of the tools carbon has at its disposal.

So why is carbon ideal for serving as the basis of life in the universe?  It comes down to how many bonds it can form with other molecules.  To understand this concept, let’s take a quick step back.  Atoms such as those listed on the periodic table have a nucleus filled with protons and neutrons.  Electrons orbit around the nucleus.  Since protons are positively charged and electrons are negatively charged, they need to balance each other in order for the atom to be stable.  To achieve balance, the natural form of atoms typically has an equal number of protons and electrons.  Helium, which has two protons, also has two electrons.  (The balancing works like this: a proton is (+1) and an electron is (-1) so helium has two protons plus two electrons: (+2) + (-2) = 0.)  Neutrons don’t have a charge; they serve as a buffer (or “glue”) to hold the positive charges in the nucleus together.


Notice that the two negatively charged electrons are (intentionally) depicted as far apart from each other as possible in this diagram.  That’s because negative charges repel – think about when you try to push two negative ends of magnets together, and how they push away from each other.  It’s easy to avoid each other when there are only two electrons, but what happens when you add more?  Carbon, for example, has six protons and six electrons.  Would its electrons be happy all in one orbit, as posited below?


The answer is definitely not – that orbit near the nucleus is pretty crowded for a bunch of negative particles trying to stay as far away from each other as possible.  In fact, only two electrons can stably live on the small orbit closest to the nucleus, and the remaining four electrons move further away.  So carbon, alone, actually exists more like this:

Carbon This second orbit is bigger than the two-electron ring near the nucleus, and it has enough space for up to eight electrons.  In fact, the second orbit is most stable when it is “full,” and has eight electrons.  That means that molecules without a full outer ring won’t like to be alone, because they aren’t very happy (stable) alone.  In order to be happy, carbon will bond up to four times with other atoms such as hydrogen, oxygen, or other carbon molecules, until it reaches eight electrons in its second ring.  The “bond” formed means that the two molecules share their electrons – a single carbon molecule bound to four hydrogen molecules (each has one electron) thus effectively finds itself in possession of eight electrons in the outer ring, while each hydrogen molecule has two electrons in its ring.  This serves to make all five molecules much happier (more stable) than they were alone.


Carbon’s ability to bond four times is a pretty big deal – most other molecules can only form one to three bonds, such as oxygen (forms only two bonds) and nitrogen (forms three).  The ability to form four bonds gives carbon a remarkable amount of flexibility.  It can happily exist in a simple state, such as in carbon dioxide (CO2); in long chains, such as in our cell membranes; in ring structures, such as sugar; and myriad other complex structures.  Carbon is what diamonds, the hardest substance on earth, are made of.  It’s the backbone of all DNA – plant and animal – and as a result it’s found pretty much everywhere you look.  Its flexibility means it has myriad applications since it happily exists in so many forms – simple molecules, long chains, flat fence-like structures, rings, soccer ball shapes, etc.

Carbon allotropes wikipedia

A few examples of the different shapes carbon can form. Each dot at the intersection of lines represents a carbon molecule; connecting lines are bonds between molecules. Photo credit: Wikipedia.

While other molecules are necessary to sustain life, it’s safe to say that we couldn’t exist without carbon.  And even though silicon can also bond four times, it lacks the rest of the properties that make carbon so nifty: it isn’t stable in rings or spheres, and usually doesn’t exist in chains of more than six (compare to fatty acid chains of carbon that can be up to twenty-eight carbons long!).  For now, we’ll have to consign the dream of finding silicon-based life forms to the realm of science fiction.

Interested in a more thorough explanation of why silicon is less likely as a basis for life?  There’s a great article from the NASA Astrobiology Institute here.


1 thought on “Carbon: Four Times Cooler Than Most of the Other Molecules You Know

  1. Mack

    Hm, neat. Thanks, really interesting set of insights there. Convenient as well, coming across this in the middle of some research for a short story 🙂

    Obviously it’s a personal thing, but for this really was perfect. Detailed enough to consider and explore, simple enough to understand relations and observations.

    Very neat.


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