Periodic Trends Lab
Same particles, opposite behaviour
Fluorine and potassium are built from the same three particles: protons, neutrons and electrons. Yet fluorine snatches electrons from almost anything, and potassium gives its outer electron away to almost anything.
The difference comes from one tug-of-war inside every atom. Once you see it, every trend on this page follows: atomic size, ion size, ionisation energy, electron affinity and electronegativity, and from those, how bonds and molecules behave.
The table stays pinned as you scroll and changes colour for each idea. Tap any element to see its value. Use the buttons above the table to switch properties yourself.
1 · The root cause
The tug-of-war: effective nuclear charge
The nucleus pulls on the outer electrons. The inner electrons sit in between and block part of that pull; this is called shielding. The pull that is left over is the effective nuclear charge, Zeff.
- Across a period, each element adds one proton, but the new electron joins the same shell. Electrons in the same shell shield each other poorly, so Zeff climbs: lithium 1.30, fluorine 5.20.
- Down a group, each row adds a whole new shell. Zeff stays about the same in groups 1 and 2 (sodium and potassium are both 2.20) and rises slowly further right. But the outer electron is now much farther from the nucleus.
These two rules carry the rest of the page. Tap elements on the table to compare their tug-of-war below.
A sketch, not to scale. Pull index = Zeff ÷ r², with fluorine set to 100. Ring size follows covalent radius.
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Slater's rules estimate shielding for an s or p electron: each other electron in the same shell shields 0.35, each electron one shell in shields 0.85, and each electron deeper shields 1.00.
They are a model, not a measurement. Other methods (Clementi–Raimondi) give larger values that rise more clearly down a group. The trends across a period agree.
Into the maths Degree
Chlorine: Z = 17, configuration 1s² 2s² 2p⁶ 3s² 3p⁵. For one 3p electron:
2 · Size
Atomic radius
Across a period, atoms get smaller: the same shell is pulled in by a stronger Zeff (lithium 128 pm, fluorine 57 pm).
Down a group, atoms get bigger: each row adds a shell, like adding a layer to an onion (fluorine 57 pm, chlorine 102 pm, bromine 120 pm).
1 pm (picometre) = 10⁻¹² m. Noble gases are hatched: they rarely form bonds, so their radii are estimates.
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An atom has no hard edge, so "radius" needs a definition. This table uses the covalent radius: half the distance between two identical atoms bonded together.
Gallium (122 pm) is barely bigger than aluminium (121 pm). The ten 3d electrons added in between (the thin strip in row 4) shield poorly, so gallium's outer electrons feel a stronger pull than expected. This is the d-block contraction.
Into the maths Degree
For a hydrogen-like orbital, the most likely distance of the electron grows with n² and shrinks with Zeff. The Orbital Lab sizes its clouds the same way.
3 · Size after a change
Ionic radius
Atoms that lose electrons shrink; atoms that gain electrons swell. On the table, the dashed ring is the atom and the filled circle is its common ion.
- Sodium (166 pm) becomes Na⁺ (102 pm): it loses its whole outer shell.
- Chlorine (102 pm) becomes Cl⁻ (181 pm): one extra electron, the same 17 protons, more repulsion.
Ions with the same number of electrons line up by proton count. All five below have 10 electrons. More protons pull the same cloud tighter.
Circles drawn to scale. Ionic radii for six-fold coordination (Shannon).
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Small, highly charged cations distort (polarise) large anions and pull their electron cloud towards themselves. This gives an ionic bond some covalent character (Fajans' rules).
It is why aluminium chloride behaves partly like a covalent compound while sodium chloride is ionic.
Into the maths Degree
Polarising power tracks charge density, roughly charge ÷ radius.
Al³⁺ is about 5.7 times stronger.
4 · Holding on
Ionisation energy
Ionisation energy is the energy needed to remove one electron from a gaseous atom. High means the atom holds on tightly.
- Across a period it rises (lithium 520 kJ/mol, neon 2081 kJ/mol).
- Down a group it falls (lithium 520, sodium 496, potassium 419).
Two dips break the pattern (outlined on the table), and they are evidence for subshells:
- Beryllium 900 → boron 801. Boron's electron sits in a 2p orbital, slightly higher in energy and shielded by the 2s pair.
- Nitrogen 1402 → oxygen 1314. Oxygen's fourth p electron has to share an orbital, and the pair repel each other.
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Successive ionisation energies reveal the shells. The huge jump shows where an electron has to come from an inner shell. Exam questions use this to deduce an element's group.
Log scale: each gridline is ten times the one below. Gold bar = the first electron taken from an inner shell.
Into the maths Degree
For a hydrogen-like atom the energy of an electron is set by Zeff and n. Hydrogen itself: 13.6 eV, which is 1312 kJ/mol (1 eV = 96.485 kJ/mol).
With Slater's Zeff this predicts the direction of the trends but not the exact values.
5 · Welcoming one more
Electron affinity
Electron affinity is the energy change when a gaseous atom gains one electron. Negative means energy is released: the atom welcomes the electron. Chlorine releases the most (−349 kJ/mol).
It generally gets more negative across a period and less negative down a group, with three exceptions:
- No stable negative ion (hatched): noble gases, beryllium, magnesium. The electron would have to start a new shell or subshell.
- Nitrogen ≈ 0: its half-filled p subshell resists.
- Fluorine (−328) releases less than chlorine (−349): it is so small that the new electron is crowded.
Darker cells release more energy. Some textbooks drop the minus sign (+349); same fact.
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The second electron affinity is always positive: adding an electron to an ion that is already negative costs energy (O⁻ + e⁻ → O²⁻, about +800 kJ/mol).
So why does O²⁻ exist in compounds? Because the lattice energy released when the solid forms pays that cost. Section 9 does the same accounting for salt.
6 · Pull inside a bond
Electronegativity
Electronegativity is how strongly an atom pulls shared electrons towards itself in a bond. The Pauling scale runs from about 0.8 to 4.0: fluorine is highest (3.98), potassium lowest here (0.82).
It rises across a period and falls down a group, for the same two reasons.
Helium, neon and argon have no value: they form no ordinary bonds.
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Mulliken's definition makes the link visible: electronegativity is roughly the average of how hard an atom holds its own electrons (ionisation energy) and how much it welcomes a new one (electron affinity).
Pauling built his scale from bond energies instead. The two scales agree closely.
Into the maths Degree
Pauling compared the real A–B bond energy with what a perfectly non-polar bond would have. The extra energy comes from the electronegativity difference. One common form, with energies in kJ/mol:
7 · Bonds
From sharing to giving: the bond spectrum
When two atoms bond, they share electrons. How fairly they share depends on the electronegativity difference, ΔEN.
- Equal pull (H–H, Cl–Cl, ΔEN = 0): equal sharing, a non-polar covalent bond.
- Unequal pull (H–Cl, ΔEN 0.96): the shared electrons lean towards chlorine. Chlorine becomes slightly negative (δ−), hydrogen slightly positive (δ+). A polar covalent bond.
- Very unequal pull (Na–Cl, ΔEN 2.23): the electron is effectively handed over. Na⁺ and Cl⁻: an ionic bond.
One continuous scale. The 0.4 and 1.7 boundaries are rules of thumb, so the zones blend.
Hydrogen fluoride (ΔEN 1.78) sits past the 1.7 line, yet it is still a covalent molecule. Treat the cut-offs as a guide.
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Pauling estimated a bond's percentage ionic character from ΔEN: H–Cl about 21%, H–F about 55%, Na–Cl about 71%. Even "ionic" NaCl keeps some sharing.
Into the maths Degree
H–Cl: ΔEN = 0.96 → 100 × (1 − e−0.230) ≈ 21%.
8 · Molecules
Polar bonds, polar molecules?
Picture each polar bond as an arrow pointing towards the more electronegative atom. A molecule is polar only if its arrows do not cancel out. Shape decides.
Non-polar means the arrows cancel. 2D sketches of 3D shapes. Grey arrows are bond dipoles, their length set by ΔEN; the gold arrow is their sum. D = debye, the unit of dipole moment (measured values).
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Lone pairs are what bend water and lift ammonia into a pyramid (VSEPR). Without them, both would be flat and non-polar. Try "Flatten it" on those two cards.
Into the maths Degree
Dipoles add as vectors. For water, with bond angle θ = 104.5°:
9 · Reactions, the bridge
Why atoms react
Atoms on the left hold their outer electrons loosely (low ionisation energy). Atoms on the right, except the noble gases, pull electrons strongly (high electronegativity). Put one of each together and an electron moves. Sodium and chlorine give salt.
The same logic explains reactivity down a group:
- Group 1 gets more reactive going down (lithium < sodium < potassium): the outer electron is farther away and easier to lose (520 → 496 → 419 kJ/mol).
- Group 17 gets less reactive going down (fluorine > chlorine > bromine): the incoming electron is pulled less strongly (3.98 → 3.16 → 2.96). That is why chlorine water turns potassium bromide solution orange: chlorine takes bromide's electron.
- Noble gases barely react: full shells, very high ionisation energy, no stable negative ion.
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Moving an electron from a gas-phase sodium atom to a gas-phase chlorine atom actually costs energy:
Salt still forms because the ions pack into a crystal and release lattice energy (about −787 kJ/mol). The full accounting is the Born–Haber cycle.
10 · Your turn
Try it yourself
Pick two elements, or tap them on the table while this section is on screen.
All trends on one page
| Property | Across → | Down ↓ | Because |
|---|---|---|---|
| Zeff | Rises | Similar or rises slowly | Same-shell electrons shield poorly |
| Atomic size | Shrinks | Grows | Stronger pull; then extra shells |
| Ion size | Shrinks | Grows | Cations lose a shell; anions gain repulsion |
| Ionisation energy | Rises | Falls | Tighter hold; then a farther electron |
| Electron affinity | More negative | Less negative | Same, with many exceptions |
| Electronegativity | Rises | Falls | Same two rules, inside a bond |
Quick check
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