Why counting is not enough
Rung zero treated electrons as tokens: carbon has four, oxygen has six, share until everyone balances. That gets you molecules. It does not get you reactions, because it cannot answer the only question that matters — which electrons move first, and where do they go?
For that you need to stop counting electrons and start locating them. Five models do it, and they are the last purely descriptive things in this ladder. Everything above this rung is about change.
The Orbital Box
An electron does not orbit like a planet. It occupies an orbital — a region of space where it is likely to be found. Three facts about orbitals carry the entire book.
First, an orbital holds at most two electrons. Second, orbitals have shape: s orbitals are spheres, p orbitals are dumbbells pointing along an axis. Third — and this is the one that does the work — an orbital is either filled or empty, and that difference is what makes one molecule attack another.
When two atoms approach, their orbitals combine. Two atomic orbitals always give two molecular orbitals: one lower in energy than either parent, one higher.
- Filled — electrons available to give
- Empty — room to accept
- Starting atomic orbitals
Hybrid Shape
Carbon's four valence electrons do not sit in four identical orbitals. They sit in one s and three p. Yet methane's four bonds are measurably identical, spaced evenly in three dimensions.
The resolution is hybridisation: the orbitals mix before bonding, averaging into a set of equivalent ones. How many p orbitals join the mix decides the shape, and the shape follows one rule — the results spread as far apart as they can.
- σ bond — the framework
- π bond — the exposed pair
Read the table right to left and it says something useful: every time carbon gains a π bond, it loses a direction. Four separate partners, then three, then two. Shape is not a fact to memorise about each molecule. It is a consequence you can derive from what carbon is bonded to.
The σ/π Split
This is the most load-bearing distinction in organic chemistry. Two orbitals can overlap in two different geometries, and the two results behave nothing alike.
A σ bond overlaps head-on, along the line joining the nuclei. The electrons sit between the atoms, shielded, held tightly. A π bond overlaps sideways, from two dumbbells lying parallel. Its electrons sit above and below the line of the bond, out in the open.
- π electrons — exposed and reactive
- σ framework — structural, unreactive
Three consequences follow immediately, and you will use all three constantly. π bonds are weaker, so they break first. π electrons are exposed, so they are what another molecule reaches. And a π bond cannot twist without the sideways overlap being destroyed, so double bonds are rigid — which is where molecular shape starts to have biological consequences.
The Polarity Gradient
A shared pair is rarely shared evenly. Atoms differ in how hard they pull on electrons — their electronegativity — and the pull rises as you go right and up the periodic table. Fluorine and oxygen pull hardest; carbon and hydrogen pull about equally, which is why plain hydrocarbon chains are so unreactive.
When a bond joins two atoms that pull unequally, the pair sits closer to the greedier one. That end becomes slightly negative, written δ−; the other end becomes slightly positive, δ+. Nothing has an actual charge. The electrons are just off-centre.
| Bond | Where the pair sits | What it means later |
|---|---|---|
| C–H | Almost exactly central | Nothing happens here. This is scaffolding. |
| C–O | Pulled hard toward oxygen | Carbon is electron-thin, so things attack it |
| C–Cl | Pulled toward chlorine | Carbon is thin, and chlorine can leave entirely |
| C–Li | Pulled back toward carbon | Carbon is electron-thick — now it attacks |
The Delocalisation Smear
So far every electron pair has belonged to one bond or one atom. Sometimes it belongs to several at once.
When p orbitals on adjacent atoms all lie parallel, their sideways overlap does not stop at one pair of atoms. The electrons spread across the whole aligned run. This is delocalisation, and it obeys one rule you can apply everywhere: spreading out is stabilising. Electrons repel each other, so giving them more room lowers the energy.
- Localised pair — one home
- Delocalised pair — many homes
- Atom framework
Push this to its limit and you get aromaticity. In benzene, six p orbitals form a closed ring with nowhere for the smear to end. The stabilisation is so large that benzene refuses reactions that would break the ring — it behaves like nothing else on this rung.
Carry this one upward carefully. Delocalisation is the reason some negative charges are comfortable and others are not, and that difference becomes the single most useful number in the subject two rungs from here.
Grok check
Prediction, not recall. Each of these is answerable from this rung alone.
- A carbon has one double bond and two single bonds. Predict its hybridisation and its bond angles — without looking up the molecule.
- cis and trans forms of a double bond do not interconvert at room temperature, but rotating a single bond is free. Explain both facts with one model.
- Which is more likely to be attacked by something electron-rich: the carbon of C–O, or the carbon of C–Li? Say why in terms of where the pair sits.
- Two negative charges, one pinned on a single oxygen and one spread over three atoms. Which molecule gives up its proton more readily?
Question 4 is the bridge. If your answer was "the spread one, because spreading is stabilising", you have already understood the mechanism behind acidity before meeting it. Next rung: the curly-arrow alphabet, where these five models start doing work. Or step back to the map.