The BE&HM Series ~ Part V: Covalent Bonding & Molecules

In the last installment I discussed how atoms in Groups IA and IIA have only 1 and 2 e's in their valence shells respectively. These atoms satisfy the Octet Rule by losing those electrons to become +1 and +2 cations.  On the other side of the periodic table, Groups VIA and VIIA have almost filled valence shells containing 6 and 7 e's respectively.  One way these atoms can satisfy the Octet Rule is to take on 2 or 1 electron to become -2 or -1 anions.  Ionic compounds are formed between at least one cation and one anion so that the substance formed is electrically neutral, held together by the electrostatic force of attraction between oppositely charged ions.  Ionic compounds must contain at least one metal (cation) and one non-metal (anion).

Covalent compounds are bound by means of electron-sharing.  Covalent compounds are formed from non-metals only.

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Let's look at the carbon atom, C.  Atomic number 6, there are 2 e's in the first shell and 4 e's in the valence shell.  We don't even need to look at that, C is in Group IVA so has 4 valence e's.  In order to have a filled outer shell, C would have to either lose all 4 e's (it would then have 2e's) and become a +4 ion, or take on 4 e's and become a -4 ion.  

The formation of ions is a balancing act between two energetically favored states.  Neutrality is favored over charged, but filled valence shell is favored over partially filled.   Speaking of the Group A elements in biologically relevant contexts, the less favorable charged +/- 1 or 2 states are outweighed by the favorability of achieving a filled outer shell.  This is not the case for atoms like carbon.  Carbon has only +6 charge in its nucleus and sustaining either a + or -4 charge would be highly energetically unfavorable.  As favorable as a filled shell is in this case, it is not enough to outweigh the preference for neutrality.  So instead of donating or receiving electrons, carbon will share it's electrons with other atoms, which in turn share theirs.  

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Electron sharing and covalent bonding is best introduced using the hydrogen atom, H.  This simplest of atoms has one proton in the nucleus orbited by a single electron in a spherical s-orbital.  Thus H is an exception to the Octet Rule in that to achieve a filled valence shell it needs 2e's.  As with the carbon atom, with only the +1 in the nucleus, taking on an electron to become -1 is energetically unfavorable and filling the shell is not favorable enough to outweigh this.  Now hydrogen does ionize to the well known H+ ion (a proton), but only under certain conditions.  More often, however, hydrogen enters into electron sharing as depicted above.  Shown are two hydrogen atoms each sharing one electron.  The s-orbitals overlap as shown on the right and one can envision where each individual hydrogen "feels complete".  These two electrons don't spend all of their time fixed in the orientation shown, but they do spend the majority of their time between the two nuclei.  This pair of electrons, called a bonding pair,  forms a single covalent bond and is replaced by a single line in pictures of molecular structures, e.g.  H-H.  I like to think of the word covalent as co-valent where co = cooperating and valent = between the valence shells.  The bonding pair is not  "owned" outright by either of the constituent H atoms, indeed neither H "owns" it's original electron.   However since the H's are the same atoms, they share the two e's equally and each ultimately still "owns"  ½(2e) = 1e.   We'll address this concept of "ownership" in the next installment of this series.

The valence shell overlap model of covalent bonding can be more difficult to envision for the other atoms as p-orbitals become involved.  We won't concern ourselves with the geometry of these for the Group VIIA elements, a simplistic overlap is sufficient.  For example, below left, see how the molecule of hydrochloric acid, HCl is formed that completes both the H and Cl valence shells.  Below right, we see the electron configuration of the chlorine gas, Cl2, molecule.

              

Note how the electrons not involved in the overlapping are shown paired up.  This is to signify electrons occupying orbitals in pairs, with no more vacancy they are not available for bonding.  These are called, what else,  non-bonding pairs.  Although all electrons migrate around the molecule once a bond is formed, it is convenient to think of these non-bonding pairs as electrons belonging to, or "owned by" each constituent atom in a molecule.   All VIIA elements form diatomic molecules similar to the Cl2 molecule shown. 

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Now the two dimensional model of valence shell overlap starts to become less useful when we are talking about molecules with the more common central atoms C, N and O.  Take carbon, C, for example, and the simplest covalent compound it forms, methane, CH4.   The electron sharing can be shown with the basic 2D model at right.  The red circles represent the 4 valence e's contributed to the bonding by C, and the blue x's are one electron contributed by each H.  Looking at the black circles symbolizing the valence shells of each atom, C feels complete, satisfying the Octet Rule with 8 e's, and each H feels complete having its shell filled with 2 e's.   It is important to understand that the distinction between the source of the electrons is for illustration only.  Once the bond is formed -- a sharing agreement, if you will, has been entered into by the atoms -- neither atom "owns" any of the bonding electron pairs.

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But the two dimensional depiction can be misleading because molecules are three dimensional structures.  For the major central elements of relevance to biological molecules (C, N, O), the 4 valence electron pairs orient in tetrahedral-based structures.  At right, carbon is shown with its 4 hybridized valence orbitals (each containing 1e) and the red lines outline the 4-triangular faced tetrahedron.  This structure is determined by what is called VSEPR theory (Valence Shell Electron Pair Repulsion).  Pairs of electrons are "forced" to pair up, but each of these negatively charged pairs is repulsed by the others.  

Imagine you have two negatively charged spheres sitting on a table connected by a string, the like charges repel one another and they will try to get as far apart from one another as possible.  Constrained by the length of the string, this orientation will be when the string is straight as shown in (1).  If you grasp the string in the middle and rapidly drag the spheres downward, the string bends at the center, but this brings the opposite charges closer to one another (2).  They repel one another and with time (3) return to the equilibrium orientation (1) .  Imagine dragging these spheres around the table or displacing the center up and down, etc.  The opposite charges repel and "work" to re-establish their equilibrium orientation (1) that minimizes the repulsive forces between them. 

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This is what is going on in three dimensions with 4 negatively charged electrons (C atom) or 4 pairs (all VIIIA elements).  A 3D depiction of the formation of methane and molecular shape is shown at right.  We could rotate this molecule any which way in 3D and it will be the same.  All H-C-H angles are 109.5.  There is no reason to go into the various shapes, as they are infinite and this isn't an organic chemistry course.  However different atoms will distort the tetrahedron as will central atoms with non-bonding pairs occupying a  "legs" of the tetrahedron.  We have not discussed multiple bonds just yet, but these too influence molecular shape.  I'll come back and address some aspects of molecular shape, or more specifically atomic orientations in a molecule that are relevant to the concept of polarity which is very important, especially in terms of physico-chemical properties of covalent compounds.  The next installment of this series will include a discussion on polarity.

Multiple Bonds

The sharing of one pair of electrons, one from each atom, forms a single covalent bond, or just a single bond.  Adjacent atoms are able to engage in orbital overlap of more than one unfilled valence orbital.   When 2 pairs of e's are shared, we form a double bond, and when three pairs are shared, we form a triple bond.  There is no such thing as a quadruple bond. 

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The O2 molecule is the classic example of a double bond.  As a VIA element, oxygen atoms have 6 valence e's, 4 of which are paired as non-bonding pairs, and 2 unpaired electrons.  Each atom contributes these two unpaired electrons forming two bonding pairs.  The double bond is indicated in structures with the double line, thus O=O.   

The N2 molecule is the classic example of a triple bond.  As a VA element, nitrogen atoms have 5 valence e's, 2 of which are paired as non-bonding pairs, leaving 3 unpaired electrons.  Each atom contributes these three unpaired electrons forming three bonding pairs.  The triple bond is indicated in structures with the triple line, thus N≡N.   

Multiple bonds can have huge impact on molecular shape and properties.  Remember Tinker Toys?  Consider the diagrams below  The structure on the left was obtained from here.   We can envision any of the connections as single bonds, whether they be end-to-end (1,2), center-to-center (3), or center-to-end (4).  This "molecule" is pictured in one orientation.  But you can rotate the sticks in each of the holes!  So for example rotating bond 2 can have a huge impact on the shape.  And if we had other wheels (atoms) attached to the front, top or right-most wheels, the rotations about these bonds would make for virtually infinite configurations at any snapshot in time.

I created the diagram above right to illustrate multiple bonds using a Tinker Toy model.  On top we have a bond as in 1&2, full rotation of the wheels (atoms) at either end.  Underneath is a double bond.  If I tried to rotate the wheels, I could probably be able to have a nominal rotation from flexing/twisting the sticks holding them together, but I cannot rotate fully without breaking the bond.  The same would hold for a triple bond (not pictured).  If connection 2 in our contraption above left were instead to include 2 pegs, there would no longer be rotation around this bond thus limiting the number of unique orientations possible, and introducing rigidity into the structure.  When these.  To separate the atoms completely I would have to break/dislodge the number of sticks.  This relates in general to bond strength and energy:  triple > double > single bond.    All of these issues will be addressed if and when discussions require.  But before moving on, multiple bonds are quite important in biochemistry.  The terms saturated, monounsaturated and polyunsaturated when discussing fatty acids refer to the presence and number of double bonds which dictates both structure and stability.

How Many Bonds & What Type Can Different Atoms Form?

This section is more for the chemistry student because we're not going to discuss the various molecules that are possible, but I'm going to go through this anyway, as it can help drive the point home of various possibilities and how we can have so many different molecules made from C&H, or C&H&O, or C&H&N etc.   The number of covalent bonds a non-metal atom can and must form = 8 - Group #  (exception H that can only form one bond).  Note the word *must* in there, because these atoms must form this many bonds to satisfy the Octet Rule which is the driving force of the bonding in the first place.  Of relevance:

• Group VIIA (and H) form 1 single bond.  Therefore these elements are always terminal atoms in any molecule.
• Group VIA form 8-6 = 2 bonds.  These can be either 2 single bonds or 1 double bond.  When they form 2 single bonds they are either central atoms or part of the chain in a molecule, when they form a double bond they are terminal atoms.
• Group VA form 8-5 = 3 bonds.  These can be either 3 single bonds, a double and a single bond, or a triple bond.  The first two scenarios would be central or chain atoms in a molecule, the latter would be a terminal atom.  And lastly,
• Group IVA form 8-4 = 4 bonds.  Any of the following combinations are possible:  4 single, 2 double, 1 double & 2 single, 1 triple & 1 single.

Let's look at 2C's and variable H's we have the possibilities below left.  Add in just one O (still variable H's) and we have the possibilities below right.   Each of the CH compounds have the same atomic constituents but different chemical and physical properties.  Same for the CHO compounds.   The left two CHO compounds have the same molecular composition - 2C, 1O & 6H - but different properties.  The left one is ethanol, that is generally written C2H5OH to denote the structure, as C2OH6 tells us nothing and could be the middle compound.

Ionic or Covalent?

Again, this is more of an academic exercise, but supplements sometimes provide elements in different forms/compounds.  There is a third class of substances, metals (held together by, you guessed it, metallic bonding) that are composed of metals only or a metallic matrix with very little dispersed non-metal (e.g. carbon is often used to strengthen steel).  We don't usually refer to these as compounds per se and will discuss these no further.  So how to tell if a compound is ionic or covalent?  Simple!  If it contains a metal and a non-metal (one from the left, one from the right), it is ionic.  If it only contains non-metals from the right corner of the table (and H), it is covalent.  Yes, in chemistry there are always exceptions and there are substances that are hybrid molecules.  An example would be sodium sulfate, Na2SO4.  The bonding between the S & O's of the polyatomic sulfate ion is covalent, the bonding between the sodium and sulfate ions is ionic.  We'll discuss this in the next installment or two.

Another thing you may or may not have picked up on is that certain non-metals can form either ionic or covalent compounds depending on the nature of the atoms they combine with.  I used chlorine, Cl, as an example because it is often misrepresented in pseudoscientific literature.  The dissolved chloride ion is not only non-toxic, it is essential for life, chlorine atoms as part of large covalent compounds (like some pharmaceuticals) can be biologically active and/or innocuous in terms of toxicity, but chlorine gas is a toxic compound.  Context is HUGE in chemistry!

I plan to devote a few posts (not necessarily consecutive) to covalent compounds.  Why?  Because pretty much every compound involved in biochemistry is covalent.  Ionic compounds mostly exist as dissolved ions, including a large number of polyatomic (covalently bonded) anions.  When the time comes we'll discuss the important behavior of such.  But fatty acids, carbohydrates and proteins are all covalent molecules.  As proteins, then, hormones, enzymes and various other peptides are covalent molecules as well.  So too, almost every drug -- OTC/prescription/elicit -- and vitamin supplement and things like sterols and flavanoids  The electrochemical behavior of these molecules (recall, that's the major focus of where I'm going here) depends heavily on the nature of the bonding, atomic arrangement in these molecules, and various "functional groups". 

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Next up, polarity.