The BE&HM Series ~ Part III: The Main Group Elements, The Octet Rule and Ions

In the previous part of this series, Atoms & The Periodic Table, I went into how the structure of atoms is related to the construction of the Periodic Table of elements.  Some take home points:

• The mass of an atom is highly concentrated in the central nucleus made up of protons and neutrons
• The volume of an atom is determined by the space "swept out" by the essentially massless electrons orbiting the nucleus.
• Electrons travel in ever-widening paths that can be approximated as concentric spheres, and electrons will fill the available paths, aka shells & orbitals, closest to the nucleus before being forced to occupy the next level.  
• Each shell contains various sub-levels or orbitals.  Depending on the shell these include one s, three p, five d and 7 f orbitals.  Each orbital can contain 2 e's.
• Interactions between atoms are between the outermost electrons occupying what is called the valence shell.  
• The chemical behavior of an atom can be predicted based on this valence shell, particularly in the Main Group or Group A elements
• The horizontal rows of the periodic table correspond to the Principal Quantum Number or outermost occupied shell in the alphabetic designations K, L, M ...
• The vertical columns of the periodic table include atoms with similar valence shell electron configurations.

With regard to the last three bullet points, it is the Group Number of each column that is most predictive of how atoms will (or won't) interact with each other.  This post will focus on the "A" elements as these are the ones most commonly involved in biochemistry.  

The Octet Rule

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Most of the Group A elements have a valence shell with an s-orbital and three p-orbitals that can hold a combined 8 electrons.  There are exceptions to this, but the only relevant one we'll discuss is hydrogen, H, with its valence shell of a single s-orbital that can hold 2 e's.  All discussions until further notice, exclude hydrogen as we'll discuss why it is considered a non-metal despite sharing certain characteristics with the rest of the Group IA atoms it is similar to later.  

To understand the Octet Rule, let's consider the sodium atom, Na, depicted at right.  Sodium is atomic number 11 (and has a mass of 23 so its nucleus contains 12 neutrons) so there are 11 electrons orbiting the neutral sodium atom.  I've already discussed how the first shell can hold 2 e's, and the second 8 e's as pictured.  This leaves the 11th electron that must orbit in the next larger M shell, and it is all alone in that space.  Recall, the electrons occupy space by "sweeping it out" rather than filling it.  You have one electron trying to establish this atom's larger volume, trying to do the work of eight!  The 8 electrons in the (middle) L shell are establishing and defending this space against all approaching atoms that might collide with it well, but that single electron is out there on its own.  It is vulnerable to attack.

The valence shells of the Main Group elements have 8 spots to be filled by 8 electrons, and this is the most stable state of the atom, hence, the Octet Rule which merely states that these atoms react and combine to achieve the stable state of a full valence shell with 8 e's.  The remainder of this post will discuss how single atoms go about accomplishing this.

Ions:  Cations and Anions

Let's start with some definitions:

• Ion:  The charged particle formed when an atom loses or gains electrons so that the number of electrons associated with it differs from its atomic number (e.g. number of protons).  Recall protons and electrons have equal and opposite charges, a neutral atom #e = #P and charge = 0.
• Cation:  A positively charged ion formed when an atom loses electrons.  When electrons are lost #P > #e, so #(+) > #(-) and the atom has a net positive charge.  
• Anion:  A negatively charged ion formed when an atom gains electrons.  When electrons are gained #e > #P, so #(-) > #(+) and the atom has a net negative charge.  
• Which is which?  I remember this by ca+ion and aNion, the t looks like a + and the N is for negative.  

Now let's revisit the sodium, Na, atom with its vulnerable 11th electron.  While the neutral state is desirable, the 8-electron containing valence shell is moreso, and the easiest way for Na to attain that is to cut the 11th electron loose.  This is what happens, although it's probably more correct to say that the electron gets knocked off easily.  And now let's revisit the Periodic Table.

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The print is small, but looking for #11 in the table, you'll find Na in the third row on the left, and below the Na you'll see 3s1. Scan above and below in that column and notice that the number changes, but the s1 does not. All Group IA atoms share the same valence shell configuration: a single electron in the s-orbital of the outer shell. They all behave like Na, and are more stable losing that "extra" electron. In doing so they become cations with a +1 charge. For "A" groups, Group I = +1 . 

Looking one column to the right, we have the Group IIA elements. That's magnesium, Mg, to the right of Na. I'm betting most can predict what Mg does ... it loses it's 2 e's (still trying to do the work of 8) and becomes a +2 cation. For "A" groups, Group II = +2. Scan up and down that column, and each of these atoms does the same.

Before we jump all the way to the other side of the Periodic Table, let's state this rule:  For the Main (A) Group elements, the number of electrons in the valence shell is equal to the group number.  Quickie Quiz time ...

Q: How many valence electrons in sulfur, S?

A: I find S in the table and see it is a in Group VIA, thus sulfur has 6 electrons in its valence shell. 

Q: Which element is Iodine, I most similar too? F, C or Li? 

A: F, both I and F are Group VIIA elements and have 7 valence electrons.

So now let's skip all the way to the right side of the table, to Group VIIIA.  These atoms have the complete complement of 8 valence electrons.  They are "happy" all the way around -- neutral and with filled valence shells capable of "defending" their space.  These are commonly called the Inert or Noble Gases -- they tend to be gases (as opposed to solids or liquids) at ambient temperatures and they are unreactive.  The octet rule tells us why, because reactions are electron interactions/transactions.  These elements have no need to engage in such to achieve an energetically favorable state.  All matter is lazy -- it seeks out the lowest energy state.  These atoms are already there and don't need to give any other atoms the time of day!!

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And now we'll go one to the left to Group VIIA.  How many electrons in the valence shells of these atoms?  Seven, and they want eight, so what do you think happens here?  Why they "suck" in an electron and become -1 anions.  Let's use chlorine, Cl, atomic no. 17 depicted at right.  Compared to Na, those 7 e's in the valence shell do a pretty good job of defending the space of the M shell ... but 8 could do better.  When a chlorine atom gains that additional electron it has one more -1 (e) than +1 (P) charge and is thus negatively charged.  

One more column to the left and we have the last cases of Main Group atoms in biological context that form ions.  The group VIA, having 6 valence electrons will gain 2 e's to form -2 anions.  

Not all Atoms form Ions

What of the remaining Groups III, IV and V?  Group IIIA elements do not play a huge role in biochemical compounds so we'll ignore this entirely.  Group IV contains the big guy -- Carbon -- also the only biochemically relevant atom in this group.  Carbon with its four electrons, has a half filled valence shell.  It would have to either lose or gain four electron to become an ion to satisfy the Octet Rule.  This is not energetically favorable, so carbon meets the rule through other means.  So too the biochemically relevant  Group VA elements of nitrogen, N, and phosphorus, P -- they do not form ions to meet the Octet Rule.

Symbols and Terminology

Cations:

• Usually referred to as "parent atom ion":  e.g. the sodium ion
• Superscripts of +n indicate the number of electrons lost, no number indicates +1, sometimes multiple + are used.  Examples:    Na⁺ , Ca⁺² or Ca⁺⁺

Anions:

• Replace the suffix (ygen, ur, ine) with "ide":  e.g. the oxygen atom becomes an oxide ion, a bromine atom becomes a bromide ion.
• Superscripts of -n indicate the number of electrons gained, no number indicates -1.  Examples:   F^- or S^-2

Ion Size vs. Atom

I'll mention this here because it may become relevant later.  Atomic size increases more considerably down a column than across a period.  This is because as you go down a column, the atoms have electrons occupying the next larger valence shell whereas across a period you are just adding electrons to an already established volume.  The differences in ion size vs. their parent atom is most pronounced with those cations formed from IA and IIA atoms.  Refer to the picture of the sodium atom, and you can see that losing that one little teeny tiny essentially massless electron makes for a much smaller ion.  The size of ions is often important in transport across biological membranes and such, so I thought this was worthwhile mentioning.  On the other side, however, with the VIIA ions, adding an electron to an almost already filled shell has a negligible effect on the ion size vs. the parent atom.  

Ion Formation, Oxidation and Reduction

The formation of an ion from its parent atom is the most basic process of oxidation or reduction.  Oxidation involves losing electrons, therefore cation formation is oxidation.  Reduction involves gaining electron, therefore anion formation is reduction.

Every atom "comes with" a certain number of electrons in this universe.  All Na atoms come with 11 e's.  When one is knocked off, the remaining 11 protons are only associated with 10 electrons.  We say that Na+ is oxidized.  Conversely when a flourine atom, F, comes with 9 electrons for its 9 protons.  When it takes on an extra electron to become a flouride F- ion, we call the anion a reduced form of flourine. 

Oxidation and reduction to form monatomic ions of  IA, IIA and VIIA elements is generally irreversible in biochemical contexts.  We'll revisit this in coming days why you're in no danger from consuming chlorine atoms as chloride ions from table salt and having them convert to toxic chlorine gas in your body.

Where do the electrons come from?  Where do they go?

This post was getting long, so I decided to end it here with this question.  The electrons are either transferred from one or more atoms to another (or more) atoms in the creation of ions, or they are shared between two or more atoms.  In both cases, substances comprised of two or more different atoms -- called compounds -- are formed.  The means by which these substances are held together is called bonding.  When electrons are transferred, in the formation of a compound the individual atoms become ions that form, you guessed it, ionic bonds and ionic compounds.  When electrons are shared, covalent molecules are formed.   I will devote a separate post to each of these types of bonding and compounds.   I'll leave you with this referring specifically to the endogeneous biological context (dietary sources and supplements can be another matter):

• Groups IA and IIA elements are always in ionic form
• Groups IVA, VA and VIA elements do not form monatomic (single atom) ions
• Group VIIA elements are generally, but not always, in ionic form