Lewis Structures and Formal Charge

Transcript

0:00Introduction

Hello everyone! Okay, a small puzzle before we start. Take the sulphate ion. I can draw it with four single bonds, and then every atom has a perfect octet. But the structure usually drawn has two double bonds, and that puts twelve electrons around sulphur! So why would chemists break the octet rule on purpose? Strange, isn't it? Today, a simple tool called formal charge settles this. And in JEE, questions on Lewis structures and formal charges come up often. So, let's begin.

0:36Drawing a Lewis structure: four steps

So, how do we draw a Lewis structure? Four steps, okay? Step one: add up the valence electrons of all the atoms. But what if it is an ion? Take a negative ion. Do you add electrons, or remove them? Add them, yes, exactly! Negative charge means extra electrons. So add one for each negative charge, and subtract one for each positive charge. Careful here. Right? Step two: put the least electronegative atom in the centre, and join the others to it by single bonds. But see the brackets: not H, meaning never hydrogen. Hydrogen can share only one pair, so it always sits outside. Clear? Step three: complete the octets of the outer atoms first. Whatever electrons are left over go on the central atom. Okay? And step four: if the central atom is still short of an octet, borrow from outside. Turn a lone pair of an outer atom into a shared pair, making a double bond, or even a triple bond. We will watch this happen in example one. With me?

1:41Formal charge: three letters

Now, the second big idea: formal charge. Formal charge equals V, minus L, minus one half of B, with the one half written in front of B, okay? Let's meet the three letters. V is the number of valence electrons of the free atom, the atom on its own. Next, L, for the lone pairs. Suppose an atom has two lone pairs. Is L two, or four? Four, correct! L counts lone-pair electrons, not lone pairs. Two pairs means four electrons. This is a very common slip. Got it? And B is the number of bonding electrons, the shared ones. Half of B gives one for each shared pair: one for a single bond, two for a double, three for a triple. So half of B is just the number of shared pairs.

2:29Try it on ozone

Clear? Let's try it on ozone. On screen: oxygen, a double bond, oxygen, a single bond, and oxygen again. So V is six for every one of them. Then do they all carry the same formal charge? No! The bonding is different, so the charges are different. Take the double-bonded end. Two lone pairs, so L is four, and one double bond, so half of B is two. Six minus four minus two: zero. Okay? Now the centre. One lone pair, so L is two, and three shared pairs: two in the double bond plus one in the single bond. Six minus two minus three: plus one. Now you do the last one. The single-bonded end has three lone pairs. What is its formal charge? Six minus six minus one: minus one. Very good! Same element, three different formal charges. Right? Now add them up: zero, plus one, minus one. Zero. And ozone is a neutral molecule. That is no accident: the formal charges always add up to the charge of the species. Use this as a quick check on every answer. With me?

3:39Which Lewis structure is best?

Now, why do we need formal charge at all? Often one species can be drawn in more than one way, and we need a judge. Rule one: the best structure is the one with the smallest formal charges. Okay? Rule two: if there is a negative charge, it should sit on the more electronegative atom, the atom that pulls electrons harder. Example two needs rule one. Got it?

4:01Example 1: formal charges in carbon monoxide

Alright, example one. Draw the Lewis structure of carbon monoxide, and give the formal charges on both atoms. Pause the video and try it first. Done? Okay, let's check. Step one, count. Carbon brings four valence electrons, and oxygen brings six. Four plus six: ten electrons. Okay? Step two: with only two atoms, just join carbon and oxygen by one single bond. That uses two, so eight electrons are left. Step three: three lone pairs complete oxygen's octet. The last two go on carbon as a lone pair. So carbon has only four electrons around it. Not enough, right? Step four: borrow from oxygen. One lone pair of oxygen becomes a shared pair. That makes a double bond, and carbon has six. Still short, so one more: a triple bond, and carbon has eight. On screen, the three lines are the triple bond, and the two dots on each side are one lone pair on carbon and one on oxygen. Both octets complete. Clear? Now the formal charges. Oxygen is the more electronegative atom, so the negative charge must be on oxygen, right? Oh no, careful! Do not guess; use the formula. Carbon: four, minus two for its lone pair, minus three for the three shared pairs of the triple bond. Minus one! Oxygen: six minus two minus three: plus one. And the sum is zero, just right for a neutral molecule. Got it? Notice, the octets came first: step four built the triple bond, and the charges simply follow. And these charges explain the tiny dipole of carbon monoxide, and why it bonds through carbon. Interesting, isn't it?

5:47Example 2: best Lewis structure of the sulphate ion

Now example two, and this is our puzzle from the start. Compare two Lewis structures of the sulphate ion. On screen, S O four, with a small four below for four oxygens, and two minus on top for the charge. Structure a: four S O single bonds. Structure b: two S O double bonds and two single bonds. Which one does formal charge prefer? Pause and try. Hint: start with sulphur. Done? Okay, let's check. Structure a. Sulphur has six valence electrons, no lone pair, and four bonds. Six minus zero minus four: plus two. Each oxygen has three lone pairs and one bond: six minus six minus one, minus one. Okay? Quick check: plus two, plus four times minus one, gives minus two. That is exactly the charge of the ion. Good. Now structure b. Sulphur has two double bonds and two single bonds: six shared pairs, and still no lone pair. Six minus zero minus six: zero! Each doubly bonded oxygen has two lone pairs: six minus four minus two, zero. Each singly bonded oxygen is minus one, just like before. With me? Check: zero for sulphur, zero and zero for the double-bonded oxygens, minus one and minus one. Minus two again. Now, structure a gives every atom a perfect octet, but structure b gives sulphur twelve electrons. So structure a must be better, right? No, no, wait! Apply rule one: the smallest formal charges win. In structure a, sulphur carries plus two. In structure b, sulphur carries zero. So structure b, with the smaller charges, is preferred, and that is the one usually drawn. Clear? And twelve electrons around sulphur? That is allowed. Sulphur is in period three, so it can expand its octet, as we saw in the last chapter. So, puzzle solved! Formal charge picks structure b. See how neat that was?

7:50Quick revision: Lewis steps, V-L-B/2 and the best structure

So, quick revision. One: for a Lewis structure, count the electrons, put the right atom in the centre, fill the outer octets, then make multiple bonds if needed. Two: formal charge is V minus L minus half of B, and the charges add up to the charge of the species. Three: the best structure has the smallest formal charges, with any negative charge on the more electronegative atom. Now go and try the practice problems on this page. All the best!