Fluorine (configuration 2, 7) will tear an electron off almost anything it touches. Add one electron and you get the configuration of neon (2, 8), which does essentially nothing. The two are neighbors, one electron apart. That single electron, and the shell it sits in, is the whole story. This page is about reading an atom from its structure: building its shells, writing its configuration, and counting the few electrons that actually do chemistry.
An atom is almost entirely empty space around a tiny, dense nucleus of protons (and neutrons). The proton count, known as the atomic number (Z), is the atom's fingerprint: change it and you have a different element. The electrons live in shells at set distances out from the nucleus, filling from the inside out.
Here is the key idea: inner electrons are locked away, doing nothing chemically. Only the electrons in the outermost shell, called valence electrons, get close enough to another atom to be lost, gained, or shared. When you predict how an element behaves, you are almost always reasoning about that outer shell alone.
An atom's chemistry is almost entirely about its outermost electrons. Everything on this page (building the shells, writing the configuration, and counting valence) is in service of finding and reading that outer shell.
Neutrons add mass without charge, and they do not alter electron arrangements, so they do not change how an element bonds or reacts. Atoms of one element that differ only in neutron count are isotopes. They matter for mass and nuclear behavior rather than chemical reactivity, which is why we reason primarily from protons and electrons.
Electrons fill from the inside out: the first shell holds 2, the next 8. Step up one proton at a time and watch each new electron find its place. The valence electrons (the outer shell) stay in the accent color. The shell configuration is just the count in each shell, written out.
The details of transition metal electron configurations (including the energy crossover from potassium onward and the chromium/copper exceptions) are beyond the assessment boundary for general chemistry. You will not be tested on predicting these anomalies, but they are explained here to clarify how the rest of the periodic table fills.
The first three shells fill in the tidy way you'd expect: 2, then 8, then 8, taking you from hydrogen to argon. But the third shell isn't actually full at 8; it can hold up to 18. So why does potassium start a fourth shell before the third is finished?
Because at that point the fourth shell sits at slightly lower energy than the rest of the third, so the next electrons go there first. Potassium and calcium open shell four (…8, 1 and …8, 2). Only then does the third shell go back and fill from 8 up to 18 (the transition metal block) before the fourth shell picks up again at gallium.
"Slightly lower energy" deserves one more sentence, because it sounds like hand-waving and isn't. Inside shell 3 there's a region called the 3d subshell, which holds electrons 9 through 18 of that shell. After shell 3 reaches 8, the next open spot in shell 3 is 3d, and 3d sits above the first spot in shell 4 (called 4s). Electrons don't care about your shell numbering. They fill the lowest-energy vacancy, period. So electron 19 takes 4s. Then, once 4s holds two electrons, their repulsion nudges the energies again, 3d dips back underneath, and the next ten electrons (scandium through zinc, the transition metals) go back and finish shell 3 before shell 4 resumes at gallium.
Two transition metals break even this pattern: chromium (2, 8, 13, 1) and copper (2, 8, 18, 1) each pull one electron down from the fourth shell into the third, because a half-full or completely full third shell is a touch more stable. You don't need to predict these; just don't be thrown when the builder shows a lone electron in the outer shell where you expected two.
For the elements you'll reason about most (the main-group ones), valence is still a clean count of the outer shell. The wrinkle is exactly why the transition metals get set aside: their outer-shell count doesn't track their chemistry the simple way.
Atoms are most stable with a full outer shell: 8 valence electrons for most main-group elements (just 2 for the first row, hydrogen and helium). Everything an atom "wants" is the shortest route to that full shell, and the valence count tells you the route.
1–3 valence electrons → the atom loses them. Shedding a few is the short path to the full shell underneath. These are the metals; the fewer it has to lose, the more reactive (Group 1 are the eager ones).
5–7 valence electrons → the atom gains the rest. Pulling in one or two completes the octet. These are the reactive nonmetals; needing just one (Group 17) makes them the grabbiest.
8 valence electrons → the atom does nothing. The shell is already full, so there's no reason to lose or gain. These are the noble gases, and that's exactly why they're inert.
Whether an atom loses or gains, and how easily, is the engine behind every reactivity trend in 1-1b. The trends there (radius, ionization energy, electronegativity) are just this idea, measured and mapped across the table.
No diagram this time: write the shell configuration and the valence count yourself, then check. Separate the shells however you like (2, 8, 7 or 2 8 7 both work). Reading a diagram is recognition; producing the configuration is the skill that actually gets tested.
Stuck on one? Tap Reveal. The point is to pull it from your head, not recognize it on a page.
The number of protons in the nucleus is the , and it sets the element's identity. Electrons fill shells from the inside out: the first shell holds and the second holds . The electrons in the outermost shell are called the electrons, and they are the ones that react. An atom with 1–3 of them tends to them, while an atom with 5–7 tends to electrons, both routes toward a full outer shell of .
• core electrons · • valence electron
same outer count → same chemistry · the inner shells change, the pattern doesn't
Neon signs and fireworks rely on electron configuration. When atoms absorb energy, their electrons jump to higher shells. As they drop back down to their stable ground-state configuration, they release the exact energy difference as a photon of colored light, painting the sky with their atomic configurations.
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