CN− Lewis Structure
Cyanide ion · A potent poison, and an industrial reagent for extracting gold and making polymers.
The CN- Lewis structure has 10 valence electrons - the 1− charge adds one, drawn as one C≡N bond, with two lone pairs in total, all inside square brackets. The central C is sp hybridized and the ion is linear. The C carries a 1− formal charge, and that is the ion's 1− charge. The charge belongs to the ion as a whole, which is what the brackets around the drawing mean.
- 10 valence e−
- linear
- sp
- 1− charge
Calculated properties
| Total valence electrons | 10 |
|---|---|
| Bonding electrons | 6 (3 shared pairs) |
| Nonbonding electrons | 4 (2 lone pairs) |
| Electron domains (steric number) | 2 |
| VSEPR notation | AX1E1 |
| Electron geometry | linear |
| Molecular geometry | linear |
| Bond angle | no angle (only two atoms) |
| Hybridization | sp |
| Charge | 1− |
| Charge distribution | localised |
| Formal charges | C: -1 |
| Resonance structures | none |
| Molar mass | 26.018 g/mol |
C Carbon 4 valence e− N Nitrogen 5 valence e−
How to draw the CN− Lewis structure
Step through it. The bar shows how much of the 10-electron budget is spent at each stage - it never grows, which is the whole constraint.
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Skeleton. Join every connected pair with one bond. 1 bond spends 2 electrons.
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Outer octets. Spend electrons as lone pairs on the outer atoms. That uses the entire budget, and the central atom has nothing left over.
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Multiple bonds. The central atom is short of an octet, so two lone pairs on the outer atoms move in to be shared. The electron count does not change - the same electrons are just counted by both atoms now.
The reasoning at each step
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Count the valence electrons
Add up the outer-shell electrons every atom brings, then account for the charge. A 1- ion has gained one electron, so it is added to the budget. Skip this and every count that follows will be wrong.
4 (C) + 5 (N) + 1 (1- charge) = 10 valence electrons -
Work out what is bonded to what
Two atoms, one connection. The only question is the bond order.
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Join everything with single bonds first
Every connection starts as one shared pair. Draw them all before worrying about double bonds - the arithmetic in the next step is what tells you where the double bonds have to go.
1 bond x 2 = 2 electrons used, 8 left -
Work out how far short you are
Add up what every atom still needs to fill its shell - eight electrons for most atoms, two for hydrogen - minus what its single bonds already give it. Compare that with the electrons you have left. The difference decides everything that follows.
needs 12, has 8, short by 4 -
Turn the shortfall into multiple bonds
Being short means atoms have to share more. Every pair that moves from a lone pair into a bond counts twice - once for each atom - so a shortfall of 4 is covered by 2 extra shared pairs. That is where the triple bond comes from.
C#N -
Fill in the lone pairs
Every electron not in a bond sits as a lone pair on an atom. Each atom takes exactly what it needs to finish its shell - there is no choice left at this point.
C: 1, N: 1 (4 electrons) -
Check the formal charges
Formal charge is valence electrons, minus lone-pair electrons, minus the number of bonds. They are not real charges, and they are not the ion's charge either - but they must add up to it. Adding these up gives 1-, which is the check that the structure is complete.
C: 4 - 2 - 3 = -1 -
Work out the shape
Count the electron domains on C: 1 bonded group plus 1 lone pair. A double or triple bond still counts as one domain, because it points in one direction. That gives linear electron geometry; ignore the lone pairs and the atoms themselves sit in a linear arrangement. Two atoms can only lie in a straight line, so there is no bond angle to measure.
steric number 2 -> sp -> linear -
Decide whether it is polar
The atoms around the centre are not all the same, so the bond dipoles are different sizes and cannot cancel.
dipoles do not cancel -> polar
Where the 1− charge sits
The C carries a 1− formal charge, and that is the ion's 1− charge. The charge belongs to the ion as a whole, which is what the brackets around the drawing mean.
People also ask whether CN− is polar. Strictly it is neither: polar and nonpolar sort neutral molecules by whether their bond dipoles cancel. A 1− ion is pulled towards any nearby charge whatever its shape, which is why its salts dissolve in water regardless of geometry. The table below is still worth reading - it shows how unevenly each bond holds its electrons, which is what decides whether the charge ends up spread or concentrated.
| Bond | Electronegativity difference | Character |
|---|---|---|
| C–N | 0.49 | polar covalent |
Common questions
How many valence electrons does CN- have?
CN- has 10 valence electrons. The atoms bring 9 between them, and the 1− charge adds one. 6 of them are in bonds and 4 sit in lone pairs.
Why does CN- have a 1− charge?
The C carries a 1− formal charge, and that is the ion's 1− charge. The charge belongs to the ion as a whole, which is what the brackets around the drawing mean. The brackets around the drawing are part of the answer: they say the charge belongs to the whole ion rather than to whichever atom happens to be nearest.
What is the molecular geometry of CN-?
CN- is a linear ion. The central C has 1 bonded group and one lone pair, a steric number of 2, which gives linear electron geometry and a linear ion.
What is the hybridization of CN-?
The central C in CN- is sp hybridized. Steric number 2 means 2 orbitals have to be mixed, which is exactly what sp gives you.
Is CN- polar or nonpolar?
Neither, strictly. Polar and nonpolar describe neutral molecules, by asking whether their bond dipoles cancel out. CN- carries a net 1− charge, so it is pulled towards any nearby charge whatever its shape - which is why ionic compounds dissolve in water regardless of geometry. The question the shape does answer is whether the charge is spread evenly: it is not, because the atoms around the centre are not all the same, so the bond dipoles are different sizes and cannot cancel.