If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

# Worked example: Lewis diagram of the cyanide ion (CN⁻)

We can draw Lewis structures for polyatomic ions (ions containing multiple atoms) using the same stepwise procedure as for neutral molecules. In this video, we'll see how to construct the Lewis diagram of the cyanide ion (CN⁻). Created by Sal Khan.

## Want to join the conversation?

• Why is an extra electron on Carbon side, isn't Nitrogen more electronegative, so shouldn't an electron hover over instead?
• Carbon and nitrogen both need to have an octet, 8 valence electrons around them. This structure satisfies that requirement.

This has nothing to do with electronegativity which is the how strongly an element pulls shared pairs of electrons towards itself.
• I don’t understand how carbon has the last bond because carbon only has 4 valence electrons right, so how does it have 5 electrons on its side of the share.
• Carbon begins with four valence electrons as a single atom, but as a second period element wants to follow the octet rule and reach eight valence electrons. The three bonds of the triple bond are a total of six electrons and the lone pair is another two electrons giving carbon a total of eight electrons which is what it would like to have.

Same thing with nitrogen, it starts with five electrons as a single atom but also wants to attain eight valence electrons following the octet rule. And so it arrives at eight electrons in the same way the carbon does.

Hope that helps.
• Why sometimes I see charges on atoms instead of a bracket? Which one is appropriate?
Is bracket for ions?
• If the entire compound or atom has a charge, you'll see that represented as a '+' or '-' on the upper right of the chemical formula. For example, a fluoride ion is F^(-), a calcium ion is Ca^(2+), and a phosphate ion is PO4^(3-). If we have an ion as a Lewis structure, the ion should have a bracket around the entire structure to show the same thing as the chemical formula, the overall charge of the ion. So my listed examples would have their valance electrons drawn in with a bracket around the ion and the charge in the upper right. Sometimes in your Lewis structure you'll want to know exactly which atom in an polyatomic has the charge and so we call this formal charge. Formal charge tells you the charge of individual atoms in an ion (neutral molecules too). So in Sal's cyanide example the carbon would have a -1 formal charge and so we write that as a negative sign in a small circle next to the carbon. In my phosphate example, three of the four oxygens would have a -1 formal charge so we would write the same negative sign in a small circle next to those oxygens. Sometimes formal charge will be drawn in lieu of the overall charge in brackets in Lewis structures, and sometimes they will be drawn together. Hope this helps.
• We always learned that, carbon makes 4 bond. But here it makes 3, can you explain me
• there is a Coordinate Covalent bond here, where carbon actually offers a full set of 2 electrons in one bond, so it's essentially carbon supplying all the electrons and nitrogen supplying the empty orbital. This is a common question to ask before you've learnt the types of covalent bonds.
• Why does the bonding stop at a triple bond? Why not a quadruple bond, even a quintuple bond!
• There are higher order bonds beyond just triple bonds. Quadruple, quintuple, and even sextuple bonds are possible. You don't really view these higher bond orders for second and third period elements like carbon because the resulting molecules are often unstable. They accomplish greater stability by using smaller bond orders so there's no incentive for them to use higher ones.

Where these higher bond orders are more prevalent are the transition metal complexes, or molecules involving transition metals. Because atoms of these elements make use of more orbitals than second or third period elements they can share more electrons in bonds. Often this happens between two transition metal atoms within the same molecule.

As a side note, taking into consideration resonance, it also possible to get fractional bond orders instead of just whole number ones. It's possible to have things like 1.3 and 1.5 bonds as well.

Hope that helps.
• If I started with the carbon atom and not the nitrogen atom when distributing valence electrons, would it make any difference?
(1 vote)
• In this case no, because either way you need to form the triple bond between C and N
• Are atoms just allowed to use two electrons from their end to make a bond? Like nitrogen in this example grants carbon access to electrons by using two of its own valence electrons, instead of sharing one in exchange for an electron of carbon. Is that a rule that applies in all chemical settings?
• So there's a difference between the electrons contributed by each atom for a bond and the electrons shared in the final bond. For cyanide the carbon contributes four valence electrons and the nitrogen contributes five (plus the additional electron for the negative charge). In the final ion though, the carbon and nitrogen are sharing collectively six electrons in a triple bond which ultimately came from both atoms.

Hope that helps.
• Is helium also an exception to the octet rule?
(1 vote)
• Correct. Hydrogen, helium, and lithium follow the duet rule where they aim to have only two electrons in their valence shell.

Hope that helps.