Lewis Dot Structures

G. N. Lewis and His Valence Electron Dot Model

This post covers the symbolic method of recording valence electron distribution across a molecule, aka the Lewis Model.

Lewis Dot Symbols of Elements

Hello,

It is useful to consider Lewis Elemental Symbols when attempting to ascertain the overall valence electron distribution in a molecule.  A correctly written lewis dot symbol will show the maximum number of single electrons (i.e. dots) around the element symbol.  See the examples below.

The vacant positions are places where electrons (from other atoms) will fill in.  The "vacant position" squares are not normally part of a Lewis Symbol.  The blue squares are meant to visually separate the Lewis Symbols from each other.  Lewis symbols are quite valuable in determining the correct Lewis Dot Structure of a molecule.

Lewis Dot Structures

A correct Lewis Dot Structure will indicate the actual 3-dimensional structure (sometimes not very well) of a molecule.  However, a chemist knows what geometric shape goes with a certain number of surrounding atoms.  For example, a chemist knows that methane (CH4) has a tetrahedral geometry because carbon is surrounded by 4 atoms (and there are no non-bonding electrons).  Molecular Geometry will be covered in my next post.

Guidelines

There are certain "rules" to follow, which will simplify the Lewis Structure writing process.  I'll present these rules in the form of questions.

Number 1
Which elements are in the center of the molecule and which elements are on the outside?
answer Be, S, P, N, and C are center-tending elements.  H and all halogens (group VIIA) are outer-tending.  Oxygen is outer-tending but is often found between a pair of atoms (eg Dimethyl Ether; H3C:O:CH3).

Number 2
How many valence electrons should be surrounding a particular element within a molecule?
answer H wants 2. B wants 6. Be wants 4 (Be is interesting because, being a metal, can also form ionic bonds).  C, N, O, and all halogens want 8 (the proverbial stable-octet!).  P wants 8, but sometimes 10.  S wants 8, but sometimes 12.

Number 3
How many pairs of electrons are bonding for a particular element?
answer H and all halogens - 1.  Be - 2.  O - 2.  S - 2 but sometimes 6.  P - 3 but sometimes 5.  N - 3 always.  C - 4 always.

Number 4
How do I know when to use double and triple bonds (2 and 3 electron pairs)?
answer  Generally, if you wind up using more electrons than you are allowed, then you may need a double or triple bond.  There are a few standard situations for certain elements:  C + 2 atoms: one single bond and one triple bond or two double bonds.  C + 3 atoms: One double bond and two single bonds.  O + 1 atom: one double bond, always.  N + 2 atoms: One double bond and one single bond.

Lewis Dot Structure Examples

The following image shows two common examples: We have methane on top and hydrogen peroxide below.  Note that hydrogen is always on the outside and hydrocarbon compounds always have carbon in the middle.  Electrons are color-coded to help you keep track of them and arrows indicate where electrons "fill-in".  Also note; because we are using properly constructed Lewis Symbols, we really don't have to deal with an electron total count.  Obviously, the number of electrons are relatively small here, but even with more complex structures, the electron count is more of a final check than a "working part" of writing Lewis Dot-Structures.

The next example is for Nitrogen Dioxide (aka Nitrous Oxide).  Note the colored electron dots showing their before and after positions.  You also see "plus" and "minus" signs.  A non-bonding electron needs to be "borrowed" from Nitrogen to complete the single bond to Oxygen.  This is a transfer of a -1 charge to O resulting in a -1 charge on the Oxygen and a +1 charge at the non-bonding position where the electron was.  This separation of charge across the molecule can make it unstable, but the existence of resonance allows Nitrogen Dioxide to exist.  Additionally, the atoms could not be put together at the outset because single non-bonding electrons, generally, don't exist (with one exception for NO2 ; Can you find it?), therefore some electrons needed to "migrate" before bonding could occur!

The last example is for Propyne.  Propyne is a hydrocarbon and this type of compound consists of a Carbon chain surrounded by Hydrogen atoms.  It's best to complete the part of the molecule with single bonds only and then do the multiple bond part.  Note the additional comments in the figure.

That's all for this post.  I hope it was helpful!