The electronegativity trend
Electronegativity increases from left to right across a period, and decreases going down a group. Fluorine is the most electronegative element at 3.98 on the Pauling scale; caesium and francium are the least, both at 0.79.
Electronegativity is how strongly an atom pulls on the electrons it is sharing in a bond. Switch the map above to electronegativity and the gradient is obvious: bright at the top right, dark at the bottom left. Here are the same four walks as numbers.
| Walk | Electronegativity (Pauling) | What it does |
|---|---|---|
| Across period 2 | Li 0.98 → Be 1.57 → B 2.04 → C 2.55 → N 3.04 → O 3.44 → F 3.98 | increases at every step |
| Across period 3 | Na 0.93 → Mg 1.31 → Al 1.61 → Si 1.9 → P 2.19 → S 2.58 → Cl 3.16 | increases at every step |
| Down group 1, lithium to francium | Li 0.98 → Na 0.93 → K 0.82 → Rb 0.82 → Cs 0.79 → Fr 0.79 | decreases (with a tie) |
| Down group 17, the halogens | F 3.98 → Cl 3.16 → Br 2.96 → I 2.66 → At 2.2 | decreases at every step |
Both directions have one cause. Going right along a period, each element gains a proton but its outer electrons stay in the same shell, at about the same distance, so the pull gets stronger. Going down a group, each step adds a whole shell, pushing the outer electrons further out and screening them behind the filled shells underneath, so the pull weakens.
Notice that going down group 1 is not a perfectly clean slide: potassium and rubidium are both quoted at 0.82. And hydrogen, which sits at the top of group 1, is left out of that walk on purpose. It has one outer electron like the alkali metals but it is a nonmetal, and at 2.2 it is more electronegative than any element below it.
Electronegativity differences are what make a bond polar, which is where this trend earns its keep.
All four trends at a glance
| Property | Across a period (left to right) | Down a group | Why |
|---|---|---|---|
| Atomic radius (no data on the map) | decreases | increases | More protons pull the same shell inward. Each new period adds a whole shell. |
| First ionization energy | increases, with dips | decreases | A smaller atom holds its outer electron closer to more protons, so it costs more to remove. |
| Electron affinity | rises to a peak at the halogen, then collapses at the noble gas | decreases, roughly | An almost-full shell wants one more electron. A full one has nowhere to put it. |
| Electronegativity | increases | decreases | Same cause as ionization energy: more protons, smaller atom, tighter grip on shared electrons. |
Atomic radius is the one to hold on to, because it explains the other three. A smaller atom holds its outer electrons closer to more protons, so it takes more energy to remove one (ionization energy), it grips shared electrons harder (electronegativity), and it usually accepts a new one more readily (electron affinity). Everything else is a consequence.
Electronegativity or ionization energy?
These two get mixed up constantly, and the reason is that they move together: both rise across a period and fall down a group. But they answer different questions.
- Electronegativity is about electrons an atom is sharing. How hard does it pull on the pair in a covalent bond? It is a comparison between atoms, on a made-up scale with no units, and it only has meaning when the atom is bonded to something.
- First ionization energy is about an electron the atom already owns. How much energy does it take to pull that electron off completely? It is a measured quantity in kJ/mol, for a single atom on its own in the gas phase.
So an isolated fluorine atom has an ionization energy (1681 kJ/mol) but no meaningful electronegativity until you bond it to something. They track each other because both come down to the same two things, nuclear charge and atomic size, not because they measure the same thing.
First ionization energy
First ionization energy is the energy needed to remove the outermost electron from a neutral gas-phase atom, in kJ/mol. It increases across a period and decreases down a group, for the same reasons as electronegativity. Helium is the highest of any element at 2372 kJ/mol.
| Walk | First ionization energy (kJ/mol) | What it does |
|---|---|---|
| Across period 2 | Li 520 → Be 900 → B 801 → C 1087 → N 1402 → O 1314 → F 1681 → Ne 2081 | increases overall, dipping at B and O |
| Across period 3 | Na 496 → Mg 738 → Al 578 → Si 787 → P 1012 → S 1000 → Cl 1251 → Ar 1521 | increases overall, dipping at Al and S |
| Down group 1, lithium to francium | Li 520 → Na 496 → K 419 → Rb 403 → Cs 376 → Fr 380 | decreases overall, rising at Fr |
| Down group 17, the halogens | F 1681 → Cl 1251 → Br 1140 → I 1008 → At 899 | decreases at every step |
The dips are the interesting part, and they are not noise. Across period 2 the value falls at boron and again at oxygen; across period 3 it falls at aluminium and sulfur. Both are the same two effects repeating:
- Boron and aluminium start a new p subshell. That electron sits slightly further out and is shielded by the full s subshell beneath it, so it comes off more easily than the rising nuclear charge suggests.
- Oxygen and sulfur are where a p orbital first has to take a second electron. Two electrons crammed into one orbital repel each other, and that repulsion helps push one out.
Group 1 metals are lowest in their rows because losing one electron leaves them looking like a noble gas, which is exactly what makes them so reactive. Going down that group the values fall as expected until the very last step, where francium comes out slightly above caesium. That is not a typo in the data: in an atom that heavy the innermost electrons move fast enough for relativity to matter, which contracts the outer 7s orbital and pulls its electron in a little tighter.
Electron affinity
Electron affinity is the energy change when a gas-phase atom gains an electron (kJ/mol, where a larger positive number means more energy released, so the atom “wants” the electron).
The usual one-line version, “it increases across a period”, is worth being careful with. What actually happens is that it climbs to a peak at the halogen and then collapses at the noble gas. Across period 2 it runs from lithium at 60 up to fluorine at 328, then drops to −116 at neon, which has a full shell and nowhere to put another electron. It also dips at group 2 and group 15, mirroring the ionization-energy dips for the same subshell reasons.
There is one famous anomaly worth knowing: chlorine (349 kJ/mol) has a higher electron affinity than fluorine (328). Fluorine’s 2p shell is so small that an incoming electron meets real crowding, and the repulsion cancels part of the energy released. It is the standard exam example of a trend having a good reason to break.
Atomic radius
Atomic radius is the trend that explains the others, and it is the one the map cannot colour: our element dataset carries no radii, so rather than guess at 118 numbers the map leaves it out.
The behaviour is the mirror image of the rest. Radius decreases across a period, because the growing nuclear charge pulls the same shell inward, and increases down a group, because each period adds a shell. Two anchors worth memorising: across period 2, lithium is much larger than fluorine; down group 1, caesium is far larger than lithium.
What “periodic trends” means
Element properties repeat predictably as you move across a period (a row) or down a group (a column). That repetition is the whole point of the periodic table, and it is why the table is shaped the way it is rather than being a list.
The heat-map above colours every element by one measurable property at a time, shaded against the real minimum and maximum for that property. Elements with no measured value stay grey and are never guessed.
Using this with a class
Project the map, switch properties, and have students predict the next colour before you reveal it. A good follow-up once they have the trend: ask them to find the dips in ionization energy across period 2 and explain them, which forces the jump from “the trend goes up” to “here is what the electrons are doing”. Tie it to the Bohr model for the full atomic picture. It is free to embed on your own site or LMS using the snippet below.