Atoms join in a small number of ways, and the way they join decides why a drop of oil beads on water, why it disappears into a carrier, and why some of it leaves the room in minutes while the rest is still there at bedtime.
~11 min read
Everything in this course so far has taken things apart: matter into elements, elements into atoms, the set laid out on the periodic table. From here it puts things back together. Everything before this lesson is parts; everything after it is molecules.
A lone atom is a rare thing in nature, and the reason sits in its outermost shell. A full outer shell is settled and low in energy; the noble gases are born with one and are known for doing nothing at all. Every other atom sits short, and a bond is how the shortfall gets resolved.
There are three answers. An atom with a spare electron or two can give them away, exposing the full shell beneath. An atom one or two short can take them. Or two atoms can share, each putting an electron into a pair both of them count. Giving and taking are one bond seen from two sides — ionic. Sharing is covalent.
Carbon is the case that matters. Four of its six electrons sit in the outer shell, exactly half way, and stripping four off or forcing four on both cost more than the arrangement is worth. So carbon shares — four times over, including with other carbons. Chains, branches and rings follow, and so does every compound in every bottle you own.
| Atom | Outer electrons | Bonds it makes | Where you meet it |
|---|---|---|---|
| Hydrogen | 1 | 1 | Capping the ends. |
| Carbon | 4 | 4 | The backbone. 70–80 per cent of an oil by mass. |
| Nitrogen | 5 | 3 | Rare in oils, everywhere in proteins. |
| Oxygen | 6 | 2 | The atom that changes everything. |
That third column is a rule to use rather than a fact to memorise: when a drawn molecule looks wrong, counting bonds is how you find out why.
Sodium has one lone electron in its outer shell and chlorine is one short of a full one, so between them there is a trade. Nothing is shared: what holds Na⁺ and Cl⁻ together is only that opposite charges attract, and since a positive charge pulls on every negative charge nearby, an ionic compound forms a lattice rather than molecules.
Diagram to come
Three panels reading left to right. Panel one: a sodium atom with a single dot alone in its outer ring and a chlorine atom with seven dots in its outer ring, an arrow carrying the one electron across. Panel two: the same two particles, now drawn smaller and larger respectively, both with complete outer rings, marked with their charges and pulled together by a plain attraction arrow — no shared pair is drawn anywhere, and that absence is the point of the panel. Panel three: a cube of alternating small positive and large negative spheres in a regular grid, with one corner being taken apart by bent water molecules whose negative ends face the positive ion and whose positive ends face the negative one. A reader should see at a glance that nothing is shared here, and that there is no single molecule to point at.
Where a healer meets this is the supplement shelf. Every mineral in a capsule arrives as an ion, and an ion cannot be sold alone, so it comes paired: oxide, citrate, bisglycinate, carbonate. That partner decides how much of the capsule is actually magnesium by weight, and how readily the pair comes apart. Reading a supplement label has the arithmetic.
The last thing to say about ionic bonds is what they are absent from. A lattice is not volatile, so distillation leaves the minerals in the plant. Everything that reaches the bottle got there by being light and covalent.
Where neither atom can afford to give or to take, they share. Each puts one electron into a pair between them and counts both as its own, so both shells read as complete and both atoms are more stable than either was alone. Carbon does it four times over, which is why an oil built from three elements holds hundreds of compounds: a drop of lavender contains well over 150.
Single bond — one shared pair
The most common bond in an oil by a wide margin. The two atoms rotate freely about the axis, so a chain of them coils and uncoils constantly.
Double bond — two shared pairs
Shorter and stronger, and rotation stops dead: the atoms at each end lock into one flat plane, fixing the shape. The second pair sits further out from the axis, which makes it the most reachable thing on the molecule.
Triple bond — three shared pairs
Very strong, perfectly straight, rare in oils. Nitrogen gas, N≡N.
Diagram to come
Three stacked rows, one per bond type, each showing two carbon atoms with the shared pairs drawn as actual dots between them — one pair, then two, then three — beside the conventional line notation. On the single-bond row, curved arrows around the axis showing free rotation. On the double-bond row, a shaded flat plane through the four attached atoms, a crossed-out rotation arrow, the second pair drawn sitting above and below the axis rather than on it, and an oxygen molecule approaching that exposed region. On the triple-bond row, a straight rigid rod. Bond length must visibly shorten down the three rows. A reader should come away able to say why a double bond is both stiffer and more vulnerable than a single one.
The bonding pattern decides the shape, and shape is what gets recognised. When you inhale an oil its molecules reach receptors in the roof of your nose, and which receptor one fits is a question about its three-dimensional form. A double bond, by refusing to rotate, holds a shape definite enough to be recognised at all.
The other half of what a double bond does is less welcome. That exposed pair is reachable, and what reaches it is oxygen from the air, which adds across the bond and makes compounds that were not in the sealed bottle. Light and warmth speed it up. Age therefore matters for more than aroma: oxidised material is more likely to sensitise skin, and more double bonds means a shorter comfortable life.
Real constituents make it concrete. Limonene, the terpene that dominates citrus, is ten carbons in a ring with two double bonds — 60–75 per cent of lemon, 80–97 per cent of wild orange. Linalool is an open ten-carbon chain with two double bonds and one oxygen, 20–47 per cent of lavender. Beta-caryophyllene, 45–65 per cent of copaiba, is a nine-membered ring fused to a four-membered one.
A chain of carbons carrying as many hydrogens as it can — no rings, no double bonds — holds exactly 2n + 2 of them. Ten carbons fully loaded is C₁₀H₂₂, and every two hydrogens missing means one ring or one double bond. Limonene, C₁₀H₁₆, is six short: one ring plus two double bonds. Linalool, C₁₀H₁₈O, is four short — two double bonds, no ring, the oxygen ignored because it costs no hydrogens. Beta-caryophyllene, C₁₅H₂₄, is eight short of thirty-two: two rings plus two double bonds.
Then run it on menthol, C₁₀H₂₀O, 30–50 per cent of peppermint. Two short: one ring, and not a single double bond. Which tells you, before anyone says a word about shelf life, that peppermint holds up in a way a citrus oil does not.
So far there have been two settings: electrons handed over, or electrons shared. Almost everything real sits between them, and that in-between is the most useful idea in this course.
A shared pair is hardly ever shared evenly. Both atoms pull on it and some pull harder, measured on a scale of electronegativity: hydrogen 2.20, carbon 2.55, oxygen 3.44. Carbon to hydrogen differs by 0.35, slight enough to treat as none. Carbon to oxygen differs by 0.89, oxygen to hydrogen by 1.24.
Where the pull is uneven the pair sits nearer the greedier atom, leaving it a small negative charge and its partner a small positive one, written δ− and δ+. Not the full charge of an ion — the pair is still shared, simply shared badly and permanently.
Which leads to the conclusion the rest of the course rests on. A molecule built only from carbon and hydrogen has no meaningful partial charges: non-polar end to end, nothing for a charge to grip. Add one oxygen and a patch becomes uneven — and that patch is a functional group. Linalool is limonene’s ten carbons with an –OH bolted on, and it is less willing to evaporate, far more at home near water, entirely different to smell.
Geometry gets a vote too. Carbon dioxide, O=C=O, has two strongly polar bonds and is non-polar overall, because they point in opposite directions and cancel.
The question to bring to every molecule
Is it polar, or is it not? Non-polar keeps company with non-polar: it dissolves in oil and fat, sits on water, slips through a fatty membrane. Polar keeps company with polar: it dissolves in water and stays put. Nearly everything else in oil chemistry, the chemistry wheel included, hangs on those two sentences.
Diagram to come
One large bent water molecule in the centre: a big oxygen sphere with two smaller hydrogens below it at a clearly non-straight angle, the angle dimensioned with an arc. The two lone pairs drawn as paired dots on the far side of the oxygen, visibly pressing the hydrogens together — a reader must be able to see that the lone pairs cause the bend. Partial charges marked on the atoms, and one bold arrow for the net dipole running from between the hydrogens up through the oxygen. To one side, three more water molecules faded back, joined to the main one by dashed lines from a positive hydrogen to a negative oxygen, so the network is visible. Draw the angle to scale rather than schematically; the bend is the whole point and must not look like an accident of layout.
Water is two hydrogens bonded to one oxygen. If that were the whole story it would be a straight line, its two polar bonds would cancel exactly as carbon dioxide’s do, and water would be a non-polar gas. It is not a straight line. Oxygen brings six electrons to its outer shell: two go into the bonds with hydrogen, and four stay put as two lone pairs, bonded to nothing but still taking up room. Four regions of electron density push each other as far apart as they can, which is the corners of a tetrahedron, about 109.5° — and a lone pair, held by one nucleus rather than stretched between two, spreads wider and presses the bonds closer together. The measured angle in water is 104.5°.
Bent means the two bond dipoles do not cancel. They add, into one dipole running from between the hydrogens toward the oxygen: the oxygen end δ−, each hydrogen δ+. Everything below follows from that.
"Like dissolves like" is the phrase people repeat. A bond angle of 104.5° is the mechanism underneath it.
Bonds hold a molecule together. They are not what holds one molecule to the next — and for almost everything you notice about an oil, the second decides the outcome.
When an oil leaves a diffuser or the back of your hand, nothing breaks — the molecules lifting off are the ones that were in the bottle. Evaporating, spreading, feeling greasy or dry: every one is attraction between whole molecules rather than within them.
Two kinds do that work. Van der Waals forces are the weak, universal pull between any two molecules lying against each other, scaling with surface area. Hydrogen bonding is a much stronger special case, needing a hydrogen attached to oxygen or nitrogen and a lone pair to reach toward.
Size alone explains a great deal. Limonene, ten carbons, boils at roughly 176 °C; beta-caryophyllene, fifteen, at around 265 °C. Same forces, nothing different but bulk — ninety degrees between them.
Hydrogen bonding reorders things again. Alcohols and phenols carry an –OH, so they donate and accept. Esters, aldehydes, ketones and oxides have an oxygen but no –OH, so they accept only. Hydrocarbons do neither. That is the boiling-point ladder: limonene leaves at about 176 °C, linalool at about 198 °C — twenty-two degrees for one oxygen atom.
Evaporating and degrading are two different events
A diffuser supplies only the small energy needed to separate molecules, which is why diffusing does not damage an oil. Heat can still ruin one — but it is not breaking bonds, it is speeding oxygen’s attack on double bonds. One is physics and reverses; the other is chemistry and does not.
This is where volatility finally has a mechanism under it. A top note is small and mostly hydrocarbon, weakly held: lemon, wild orange, peppermint, off a strip within the hour. A middle note like lavender runs on esters and monoterpenols and lasts hours. A base note is sesquiterpene territory — copaiba, cedarwood, vetiver — with a hydrogen bond on top of the bulk. Days on a strip.
Put a drop of lemon on one hand and vetiver on the other and time them. Nothing in the result is about strength — it is surface area and hydrogen bonds. It is also why a base note fixes a blend.
| Interaction | Rough strength | What it decides |
|---|---|---|
| Carbon-carbon single bond | about 350 kJ per mole | Whether the molecule exists at all |
| Carbon-carbon double bond | about 610 kJ per mole | Shape, rigidity, where oxygen attacks |
| Hydrogen bond | about 20 kJ per mole | Boiling point, water solubility, how long it lingers |
| Van der Waals contact | about 1–5 kJ per mole each | Everything that scales with molecular size |
Weak forces win by numbers rather than by strength: one contact is nothing, but a sesquiterpene lying against its neighbours makes a great many at once, and the sum is what you are working against. The van der Waals pull exists because electrons are never still — at any instant a molecule’s electron cloud is slightly lopsided, which makes it a fleeting dipole and nudges its neighbour into a matching one.
Five sections of groundwork. Here is what it buys you.
Why oils and water separate
Water will not open its network for a molecule with nothing to trade, so a drop floats and spreads into a film. Excluded, not diluted.
Why a carrier works and water does not
Fractionated coconut oil is non-polar too, so like holds like: spread, slower evaporation, lower concentration.
Why dilution changes behaviour and not only strength
A neat drop is a small area at full concentration briefly; in a teaspoon of carrier it is a large area at low concentration for far longer.
Why an old citrus oil has changed
Oxygen has added across its double bonds: the aroma flattens, and the new compounds are likelier to sensitise skin. Monoterpene-heavy oils have the shortest life.
Why heat matters twice
Warmth drives molecules apart — evaporation — and speeds oxygen’s attack on double bonds — degradation. Distillation does the first, avoids the second.
Why an expressed citrus oil behaves differently in sunlight
Cold pressing takes everything in the peel, including furocoumarins — too heavy for steam to lift into a condenser, and the ones that react with UV on skin. After applying lemon, lime, grapefruit, bergamot or wild orange topically, keep the area out of direct sunlight and UV for up to 12 hours.
What are you about to do with the oil?
The same chemistry gives a different answer depending on where the drop is going.
One qualification. Bonding explains a molecule, and an essential oil is never one molecule. The research behind the oil effect found that isolating a constituent did not reproduce what the whole oil did. Chemistry tells you what each part can do, not what the mixture does. Safety and how to use oils hold the practical guidance; availability differs by country.
Keep oils away from eyes and inner ears, and out of reach of children. Dilute before applying to skin, and dilute more heavily for children, for sensitive skin and for the warm oils named above. Only take an oil internally if the label on your bottle states that it is for internal use. If you are pregnant, nursing, under medical care or taking medication, speak to your healthcare practitioner before use. Nothing here is intended to diagnose, treat, cure or prevent any disease.
You now have atoms joined into molecules, and two properties — polarity and size — that predict a great deal of their behaviour. Next comes the first real family of molecules, built from one repeating five-carbon unit.
Lesson 5 of 18 · Science