Foundations · Section 1 of 64

Atomic structure

Practice this — interactive lesson

Organic chemistry is, in the end, a book about where electrons go. Every mechanism you will draw is a story about electrons leaving one place and arriving at another, and the reason they move at all comes down to the structure of the atoms holding them. So the course starts here, one level below anything that looks like organic chemistry, because almost every rule that comes later is this chapter wearing a different hat.

The three particles

An atom is built from three particles. Protons carry a positive charge, sit in the nucleus, and have a mass of about 1 atomic mass unit (amu). Neutrons carry no charge, also sit in the nucleus, and also weigh about 1 amu. Electrons carry a negative charge, occupy the space around the nucleus, and weigh about 1/1800 amu — so little that for most purposes an atom's mass is just its protons and neutrons.

The number of protons is the element's atomic number (Z), and it is what defines the element. Change the proton count and you have a different element entirely; there is no such thing as a carbon atom with seven protons. Everything else about an atom — its charge, its mass, its reactivity — can vary while it stays the same element. This is worth stating plainly because organic chemistry spends a great deal of time moving electrons on and off carbon, and it is easy to lose track of the fact that none of that ever changes what the atom is.

Isotopes and mass number

An element's neutron count can vary. Atoms of the same element with different neutron counts are isotopes. The mass number (A) is protons + neutrons, written as a superscript: carbon-12 (¹²C) has 6 protons and 6 neutrons; carbon-13 (¹³C) has 6 protons and 7 neutrons. Both are carbon, both behave chemically almost identically, and both are present in every organic sample you will ever handle — about 98.9% ¹²C and 1.1% ¹³C.

Notice what changing the neutron count does and does not do. It changes the mass. It does not change the proton count, so the element is the same; and it does not change the electron count, so ¹²C and ¹³C form the same bonds, in the same places, at the same rates. Two atoms can differ in weight and be chemically indistinguishable.

CC¹²C — 98.9% of all carbon¹³C — 1.1%6 protons + 6 neutronsmass 126 protons + 7 neutronsmass 13
The two carbons in every sample you will handle. The extra neutron in ¹³C changes the mass and nothing else — the nucleus is not what bonds, so both atoms carry the same four valence electrons (teal) and form the same bonds at the same rates. That extra neutron is not useless to you: it is the entire basis of ¹³C NMR in Module 14.Same 4 valence electrons → chemically identical

That is the first piece of evidence for the claim this chapter opened with. If chemistry were about the whole atom, an 8% difference in mass would show up somewhere. It does not, because bonding is the business of electrons alone — and the nucleus matters only for how many of them an atom has.

Shells and valence electrons

Electrons occupy shells around the nucleus. The innermost (K) shell holds up to 2 electrons; the next (L) shell holds up to 8. A neutral atom has exactly as many electrons as protons. The electrons in the outermost occupied shell are the valence electrons, and they are the only ones that participate in bonding. Core electrons — everything in the filled inner shells — are held too tightly and are too well shielded by the outer electrons to take part in chemistry.

For the main-group elements that make up essentially all of organic chemistry, there is a shortcut worth memorizing now and never looking up again: an element's group number gives its valence electron count. Carbon is in group 14, so 4 valence electrons. Nitrogen, group 15, has 5. Oxygen, group 16, has 6. The halogens, group 17, have 7. Hydrogen has 1.

Worked example — building nitrogen (Z = 7)
NNeutral nitrogen, Z = 77 protons, 7 electrons2 core (grey) + 5 valence (teal)NHHHlone pairso nitrogen builds NH₃3 bonds + 1 lone pair= a full octet of 8
Filling nitrogen from the inside out. Two electrons finish the inner K shell; the remaining five go to the L shell and are the valence electrons. The L shell holds eight, so nitrogen is three short — which is why it makes exactly three bonds and keeps the leftover pair to itself.

Nitrogen has 7 protons, so a neutral nitrogen atom has 7 electrons. Fill from the inside out: 2 electrons complete the K shell, leaving 5 for the L shell. Those 5 L-shell electrons are nitrogen's valence electrons — matching its group number, 15.

The L shell holds 8, so nitrogen is 3 electrons short of full. That single number predicts nitrogen's entire structural behavior in this course: it forms 3 bonds and keeps 1 lone pair — a pair of electrons that stays on the nitrogen itself rather than being shared with a neighbour. Ammonia (NH₃), amines (R–NH₂), and the nitrogen of an amide all follow from it.

Why valence electrons matter: the octet rule

Atoms are most stable when their outer shell is completely full — 8 electrons for almost everything you will meet, and 2 for hydrogen, whose only shell is the K shell. An atom with a nearly-empty or nearly-full outer shell is comparatively unstable, and will gain, lose, or share electrons to reach a full octet. Sharing electrons between two atoms so that both can count them toward their own octet is precisely what a covalent bond is.

Counting to eight is the single most useful habit in the first half of this course. It tells you how many bonds an atom "wants": carbon needs 4 more electrons and so forms 4 bonds; nitrogen needs 3; oxygen needs 2; a halogen needs 1. When you later meet a structure that violates one of these counts — a carbon with only three bonds, an oxygen with three — that violation is not a typo. It is the interesting part of the molecule, and it will carry a charge.

Common slip: the octet rule counts electrons, not bonds. A carbon in ethene is bonded to only three things (two hydrogens and one carbon) but still has a full octet, because one of those bonds is a double bond contributing four shared electrons rather than two. Count electrons and the rule never lets you down; count neighbors and it will mislead you the first time a double bond shows up.

Ions: what "cation" and "carbocation" actually mean

An atom that loses an electron is left with more protons than electrons and is positively charged overall — a cation. An atom that gains an electron is negatively charged overall — an anion. The mnemonic that survives: a cation is "pawsitive."

This vocabulary reappears constantly, applied to carbon rather than to metals. A carbocation (Module 6, SN1 and E1) is a carbon that has lost a bonding pair and carries a positive charge with only six valence electrons — an incomplete octet, which is exactly why it is so reactive and so eager to accept a pair from somewhere. A carbanion is a carbon-centered anion: three bonds, one lone pair, a full octet, a negative charge, and a powerful appetite for anything electron-poor. Sodium giving up an electron and a carbon giving up a group it was bonded to are the same idea two hundred pages apart.

CarbanionC−RRR3 bonds + 1 lone pair8 electrons — charge −1Neutral carbonCHHHH4 bonds8 electrons — neutralCarbocationRRRC+empty p3 bonds, empty p orbital6 electrons — charge +1
The three states a carbon can be in, counted the way the octet rule asks — in electrons, not bonds. Take a bonding pair away from neutral carbon and you are left with six electrons and an empty p orbital (dashed): that hole is the whole reason a carbocation is so desperate for a pair. Leave the pair behind instead and carbon holds eight, a lone pair and a negative charge. You meet the cation in Module 6 and the anion in Module 11.
Sodium (Z = 11) has the configuration 2, 8, 1 — a single valence electron, marooned far from a full shell. Losing that one electron leaves it with a complete 2, 8 shell structure, which is why sodium gives it up so readily and why Na⁺ is what you actually encounter. One structural fact predicts most of the element's chemistry. This is the pattern the whole course runs on: get the electron picture right and the reactivity follows.

What carries forward

Three things from this section are load-bearing for everything after it. First, valence electron counts from group number — you will use these to draw every Lewis structure in Module 1 and to compute every formal charge. Second, the octet rule as an electron count, which is the test that tells you whether a structure you have drawn is plausible. Third, the cation/anion vocabulary, because the charged intermediates of Modules 6 through 11 are all described in exactly these terms. The next section, Orbitals, replaces the tidy fiction of "shells" with the real picture of where those valence electrons actually live.