Once the kits came out, the room got quieter. Building from a red ball with two holes drilled at an angle is easy. Working out why the holes are at an angle, and why the carbon in needed two sticks to each oxygen, is the part the kit cannot do for you. A Lewis structure is the piece of paper that decides what the model should look like before you build it.
It is also the shortest route to the answers you measured in Testing for Bond Type. If you can draw the structure, you can say whether the substance is ionic or molecular, whether the molecule is polar, and roughly how it will behave.
The dots are the valence electrons, and nothing else
A Lewis structure shows only the outer-shell electrons, because only those take part in bonding. A sulfur atom has sixteen electrons; its Lewis symbol shows six. The ten inner ones are real, and they are irrelevant to what sulfur bonds to.
For a main-group element the number of valence electrons is readable straight off the periodic table: group 1 has one, group 2 has two, group 13 has three, and so on up to group 18 with eight. That is not a coincidence — it is the periodic law from The Periodic Table as an Argument showing up as a counting rule.
The two kinds of bonding are drawn differently, and the difference is the point rather than a convention:
- Ionic. Draw the electron leaving. Sodium’s single dot moves to chlorine, and you write the result as two separate species in square brackets with their charges: beside carrying eight dots. Nothing is shared, so nothing is drawn between them.
- Covalent. Draw the shared pairs between the atoms, usually as a line for each pair. Pairs that belong to one atom alone are lone pairs and stay as dots on that atom.
A procedure that works
Follow this in order. Skipping to the answer works for water and fails for anything with a double bond.
- Count the total valence electrons for the whole species. Add one for each unit of negative charge, subtract one for each unit of positive charge.
- Choose a central atom — the least electronegative one that is not hydrogen. Hydrogen can form only one bond, so it is never central.
- Join every outer atom to the centre with a single bond. Subtract two electrons for each bond you drew.
- Spend the remaining electrons completing octets on the outer atoms first, as lone pairs.
- Anything still left goes on the central atom as lone pairs.
- If the central atom is now short of eight, move a lone pair from an outer atom into the space between them, making a double or triple bond.
- Count everything you have drawn. Bond pairs plus lone pairs must come to the total from step one, exactly.
Carbon dioxide runs like this. Total valence: . Carbon is the least electronegative, so it is central. Two single bonds use four electrons, leaving twelve, which complete the octets on both oxygens. Carbon now has only four electrons around it, so pull one lone pair from each oxygen into the bonding region — and you have , with two lone pairs left on each oxygen. Sixteen electrons, all accounted for.
Ethene works the same way: valence electrons, and the only arrangement that gives both carbons eight is a carbon–carbon double bond with two hydrogens on each. Every electron is in a bond and there are no lone pairs at all.
The octet rule and its real exceptions
The octet rule is a good rule with a bad reputation for being stated as if it were universal. It is not, and the exceptions are common enough that you will meet them this year.
- Hydrogen wants two, not eight. Its outer shell is the first shell and the first shell holds two. Helium is full at two as well.
- Beryllium is stable with four. In the beryllium has two bonding pairs and nothing else.
- Boron is stable with six. has three bonds around boron and no lone pair. This is why boron compounds are so reactive towards anything with a lone pair to offer — the shell has room.
- Period 3 and beyond can exceed eight. Phosphorus in has ten electrons around it, and sulfur in has twelve. These atoms are larger, so more neighbours physically fit, and they have energy levels available that period 2 atoms do not. You will often see this explained as “d orbitals”; that is a simplification, and the honest version is that the reason is more subtle than Grade 11 needs. What matters here is that nitrogen can never do it and phosphorus can, and that is a size argument.
- Some molecules have an odd number of electrons, so somebody has to be short. Nitrogen monoxide, , has eleven valence electrons. No arrangement gives everyone eight, and the molecule exists anyway.
When a structure will not close, check the count first. Most “exceptions” students meet are arithmetic.
What a model shows, and what it lies about
Build and from the same kit and the difference is immediately physical: one is bent, one is straight. That difference is the whole reason water dissolves salt and carbon dioxide does not dissolve much of anything, which is where Water and Solutions picks up.
The model is telling the truth about shape and connectivity. It is lying about three things, and it is worth knowing which:
- Scale. The sticks are far too long. Real atoms in a bond overlap; they do not sit at opposite ends of a rod.
- Stillness. Molecules stretch, bend, and rotate constantly. The model is a time-averaged snapshot of something that never stops moving.
- Particles that are not there. Build sodium chloride and you must build a lattice — a repeating block of alternating ions with each ion surrounded by six of the other kind. A student who builds one stuck to one has built something that does not exist. This is the single most useful thing a model kit can teach about ionic compounds, and it only works if you refuse to stop at two balls.
Between the drawing and the model sits the structural formula, which is the Lewis structure with the lone pairs left off and the bonds drawn as lines. It is what you will actually write most of the time, and it is enough for everything in Naming and Formulas and Predicting Products.
Curriculum connection
B2.4
draw Lewis structures to represent the bonds in ionic and molecular compounds [PR, C]
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B2.6
build molecular models, and write structural formulae, for molecular compounds containing single and multiple bonds (e.g., , , ), and for ionic crystalline structures (e.g., NaCl) [PR, AI, C]
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