Sharing, not stealing: what a covalent bond really is
In the last guide you met carbon and asked why it sits at the centre of life. The short answer was: it bonds. Now we slow that word down. There are two big ways atoms can join. In an *ionic* bond, one atom yanks an electron clean off another — table salt is sodium that handed an electron to chlorine, leaving Na+ and Cl- stuck together by raw plus-minus attraction. Carbon almost never does this. Instead carbon forms the covalent bond: two atoms each contribute one electron and *share* the resulting pair, like two people holding the ends of a single rope. Neither lets go, and the shared pair sits in the space between the two nuclei, gluing them together.
Why would sharing make atoms stick? Each shared pair sits where it feels the pull of *both* positive nuclei at once, and that double attraction is a lower-energy, more comfortable place for the electrons to be than floating around either atom alone. Lower energy means more stable, and nature rolls downhill toward stability. A covalent bond is simply nature's way of letting two atoms both be happier by pooling a pair of electrons. The whole of organic chemistry is the story of these shared pairs — where they sit, when they shift, and what happens when one atom tugs harder than the other.
Valence electrons and the octet rule
Only the *outermost* electrons of an atom take part in bonding — the ones in the highest shell, called the valence electrons. The inner electrons sit close to the nucleus, fully spoken for, and never join in. So the first thing to learn for any atom is: how many valence electrons does it bring to the table? A neat trick for the main-group elements organic chemists care about: the count equals the group number across the top of the periodic table. Hydrogen brings 1, carbon brings 4, nitrogen 5, oxygen 6, the halogens (F, Cl, Br, I) 7. Those five numbers will get you through almost all of organic chemistry.
Now the famous rule of thumb that drives the whole game. Atoms behave as if they want their valence shell to look like the nearest noble gas — and for almost everything in organic chemistry that means eight electrons in the outer shell. This is the octet rule: carbon, nitrogen, oxygen and the halogens are at their most stable surrounded by eight valence electrons, counting both shared and unshared. That is exactly why carbon, with its four lone electrons, makes *four* bonds — four shared pairs add four more electrons to its own four, reaching eight. Nitrogen makes three bonds, oxygen two, hydrogen one. The number of bonds an atom 'wants' is simply how many electrons it needs to borrow to hit a full shell.
Lone pairs, bonding pairs, and how to draw them
A Lewis structure (an electron-dot structure) is the chemist's way of drawing every valence electron in a molecule explicitly. There are only two places an electron pair can live. A *bonding pair* is shared between two atoms — we draw it as a line connecting them, and two such lines side by side mean a double bond, three a triple bond. A *lone pair* belongs to a single atom alone, unshared — we draw it as two dots sitting on that atom. Water, H2O, is the perfect starter: two O-H lines (the bonding pairs) and two pairs of dots on the oxygen (its lone pairs). Count around that oxygen — two bonds give four shared electrons, two lone pairs give four more — and you find its octet.
Drawing a Lewis structure is a recipe you can follow mechanically, and it is worth burning into memory because you will do it thousands of times. The goal is always the same: place every valence electron somewhere, and give every atom (except hydrogen and the known exceptions) a full octet. Here is the procedure, walked through the way you would do it on paper.
- Count the total valence electrons — add up every atom's contribution. For an ion, add one electron per negative charge and subtract one per positive charge.
- Sketch the skeleton: pick the central atom (usually the least electronegative one that isn't hydrogen — often carbon) and arrange the others around it. Connect each with a single bond.
- Subtract the electrons you just used (two per bond) and spread the rest as lone pairs on the outer atoms first, filling their octets.
- Put any leftover electrons on the central atom. If the central atom is still short of an octet, slide a neighbouring lone pair into a shared position — turning a single bond into a double or triple bond — until everyone is satisfied.
- Check your work: total the electrons in the drawing (lines count as 2, dots in pairs count as 2) and confirm it equals the count from step 1, with every atom at its target shell.
CO2 (carbon dioxide) total valence e- = 4 + 6 + 6 = 16
.. ..
O == C == O two C=O double bonds
'' ''
each O: 2 lone pairs + 1 double bond -> octet (8)
carbon: 2 double bonds, no lone pairs -> octet (8)
electrons used: 2 bonds x4 + 4 lone pairs x4 = 16 checkFormal charge: bookkeeping for electrons
Sometimes more than one octet-satisfying structure is possible, and you need a tie-breaker to decide which is best — or to label where a charge sits on an ion. That tool is formal charge: a careful piece of bookkeeping that asks, for each atom, "how many electrons does this atom *own* in my drawing, versus how many it owns when neutral and free?" The ownership rule is simple and fair: an atom owns *all* of its lone-pair electrons, but only *half* of each bonding pair (since it shares those with a partner). The formula is: formal charge = (valence electrons of the free atom) - (lone-pair electrons) - (half the bonding electrons), which is the same as valence minus lone-pair electrons minus number of bonds.
Try it on the hydronium ion, H3O+. Oxygen normally brings 6 valence electrons. In H3O+ it has three O-H bonds and one lone pair, so it owns the 2 lone-pair electrons plus half of the 6 bonding electrons (3), giving 5. Six it should own, five it actually owns: formal charge = 6 - 5 = +1. The positive charge sits on the oxygen, exactly as the +1 on the ion told us. Formal charge does not measure the *real* charge on an atom — that depends on electronegativity and is messier — but it pins the bookkeeping charge to the right atom, and that is what you draw and reason with.
Formal charge earns its keep in two big ways down the road. First, the best Lewis structure is usually the one with the *fewest* and *smallest* formal charges, with any negative charge resting on the most electronegative atom — a quick way to choose between rival drawings. Second, formal charge is how you keep track of the charged species at the heart of mechanisms: the electron-poor carbocation (a carbon with only three bonds and a +1 formal charge, missing its octet) and the electron-rich carbanion (a carbon with a lone pair and a -1 formal charge). When you later push electrons around with curved arrows, formal charge is the running tally that tells you the arrows added up.
One molecule, several drawings: a first look at resonance
Once in a while a single Lewis structure cannot honestly describe a molecule, because the electrons are spread out more than any one drawing of lines and dots can show. Take the nitrate ion, NO3-. You can draw the double bond going to any one of the three oxygens, giving three equally good structures — but the real ion has all three N-O bonds identical, each one and a third of a bond. The molecule is not flickering between the pictures; it is a single fixed thing that *all three drawings together* approximate. This is resonance, and the separate drawings are called contributors or resonance structures.
Why bother with this now, in the very first bonding guide? Because spreading electrons over several atoms is *stabilizing* — a charge or a multiple bond that is shared out is more comfortable than one crammed onto a single atom. That single idea quietly explains a huge amount of what is coming: why some acids are strong, why some carbocations survive long enough to react, why benzene is so weirdly stable. You do not need to master it yet. You only need to recognize, when you see two or three Lewis structures linked by a double-headed arrow, that they are humble attempts to draw one richer reality.
Why this is the first skill of all
It is tempting to treat Lewis structures as a tedious chore to get past on the way to the 'real' chemistry. That is backwards. Every reaction you will ever learn is, at bottom, a story about where electron pairs are and where they move to. A reaction starts from a correct picture of the bonds and lone pairs in the reactants and ends at a correct picture of the bonds and lone pairs in the products. If your starting drawing is wrong — a missing lone pair here, a miscounted bond there — every conclusion you draw from it will be wrong too, no matter how clever the rest of your reasoning is.
So the lone pairs you so carefully drew are not decoration. A lone pair on oxygen or nitrogen is precisely the electron-rich elbow that will reach out and attack an electron-poor carbon later on; a formal charge flags an atom that is desperate to give or take electrons. Reading those features off a structure is how chemists *predict* reactivity instead of memorizing it. Get into the habit, starting now, of never writing a molecule without its lone pairs and charges shown — your future self, three rungs up this ladder, will be grateful.
One last warning to carry forward: a flat Lewis structure tells you *what is bonded to what* and *where the electrons are*, but it deliberately says nothing about *shape*. The dots and lines are a connectivity map, not a photograph — water is drawn with a square corner but is really bent, methane looks flat on paper but is a three-dimensional tetrahedron. Bonding is the skeleton; shape is the flesh, and shape is the very next thing this rung will give you. Master the electron bookkeeping here, and you are ready to ask how those shared pairs arrange themselves in space.