# Periodic trends

Canonical: https://duckyhelper.com/learn/chemistry/periodic-trends/
Updated: 2026-10-01

Periodic trends are patterns in how atoms behave across the periodic table. Going left to right across a period, atoms get smaller, and ionization energy and electronegativity go up, because more protons pull on electrons in the same shell. Going down a group, atoms get bigger, and ionization energy and electronegativity go down, because the outer electrons sit in shells farther from the nucleus.

## Key ideas

Two ideas explain every trend on this page:

- **Pull from the nucleus.** Across a period, each element has one more proton but its outer electrons stay in the same shell. Inner electrons block, or **shield**, only part of that extra pull. So the pull on the outer electrons, called the **effective nuclear charge**, grows from left to right.
- **Distance.** Down a group, each element adds a whole new shell. The outer electrons are farther away and shielded by more inner electrons, so they are held more loosely.

**The main trends and why they happen**

| Property | What it means | Left to right across a period | Top to bottom down a group |
| --- | --- | --- | --- |
| Atomic radius | Size of the atom | Decreases (stronger pull) | Increases (more shells) |
| First ionization energy | Energy to remove the outermost electron, in kJ/mol | Increases | Decreases |
| Electronegativity | How strongly an atom pulls shared electrons in a bond | Increases | Decreases |
| Metallic character | How easily an atom loses electrons | Decreases | Increases |

So the smallest atoms with the highest ionization energy and electronegativity sit at the top right (fluorine, ignoring noble gases for electronegativity). The largest, most metallic atoms sit at the bottom left (cesium and francium).

### Ion size

Sizes on this page are in picometers (1 pm = \(10^{-12}\) m). Atom sizes are covalent radii and ion sizes are ionic radii.

- **Positive ions (cations) are smaller than their atoms.** \(\mathrm{Na^+}\) has lost its whole outer shell. Radius: Na 166 pm, \(\mathrm{Na^+}\) 102 pm.
- **Negative ions (anions) are larger than their atoms.** Extra electrons repel each other and spread out. Radius: Cl 102 pm, \(\mathrm{Cl^-}\) 181 pm.

> **Tip: Two exceptions AP Chemistry loves**
>
> Ionization energy does not rise smoothly. Mg (738 kJ/mol) is higher than Al (578), because aluminum's outer electron is in a 3p orbital that is a little higher in energy. P (1012) is higher than S (1000), because sulfur's fourth 3p electron has to pair up in an orbital, and the paired electrons repel each other.

## Worked examples

**Example 1: rank by atomic radius**

Problem: Rank Na, Mg, S and Cl from largest to smallest atomic radius.

1. All four are in period 3, so they have the same number of shells.
2. Moving right, each adds a proton, so the pull on the outer electrons grows and the atom shrinks.
3. Order by position, left to right: Na, Mg, S, Cl. Check with data (covalent radii): Na 166 pm, Mg 141 pm, S 105 pm, Cl 102 pm.

Answer: Na > Mg > S > Cl

**Example 2: rank by ionization energy**

Problem: Rank F, Cl, Br and I from highest to lowest first ionization energy.

1. All four are in group 17, so this is a down-the-group trend.
2. Each step down adds a shell, so the outer electron is farther away and easier to remove.
3. Highest at the top. Data: F 1681, Cl 1251, Br 1140, I 1008 kJ/mol.

Answer: F > Cl > Br > I

**Example 3: atom or ion?**

Problem: Which is larger in each pair: Mg or \(\mathrm{Mg^{2+}}\)? O or \(\mathrm{O^{2-}}\)?

1. \(\mathrm{Mg^{2+}}\) lost two electrons and its whole third shell, so it is smaller than Mg (141 pm vs 72 pm).
2. \(\mathrm{O^{2-}}\) gained two electrons. The same 8 protons now hold 10 electrons, which spread out, so it is larger than O (140 pm vs 66 pm).

Answer: Mg is larger than \(\mathrm{Mg^{2+}}\), and \(\mathrm{O^{2-}}\) is larger than O

## Common mistakes and how to fix them

- **Thinking more protons means a bigger atom.** Across a period, more protons pull harder, so atoms shrink. Fix: size grows only when you add a shell, which happens going down.
- **Explaining with "it wants a full octet".** Teachers want the real reason. Fix: use effective nuclear charge across a period and distance plus shielding down a group.
- **Mixing up ionization energy and electronegativity.** Ionization energy is about removing an electron from a lone atom. Electronegativity is about pulling shared electrons inside a bond.
- **Giving noble gases an electronegativity.** They rarely form bonds, so most tables leave them blank. Their ionization energies are the highest in each period, though.

**Practice questions**

1. Which atom has the largest atomic radius?
   A. Li
   B. Na
   C. K
   D. Mg

   Answer: K. K is lowest in group 1 among these, so it has the most shells (4). Li and Na have fewer shells, and Mg is to the right of Na in the same period.

2. Which element has the highest electronegativity?
   A. N
   B. O
   C. F
   D. Cl

   Answer: F. Electronegativity rises to the right and up. Fluorine is the top right of the elements that form bonds, with 3.98 on the Pauling scale.

3. Which element has the lowest first ionization energy?
   A. Cs
   B. Na
   C. Mg
   D. Cl

   Answer: Cs. Cesium is far down group 1. Its single outer electron is very far from the nucleus and well shielded, so it takes only 376 kJ/mol to remove.

4. Which has the higher first ionization energy, phosphorus or sulfur?
   A. P
   B. S

   Answer: P. This is an exception to the trend. Sulfur's fourth 3p electron must pair up, and paired electrons repel, so it is a little easier to remove. P is 1012 kJ/mol, S is 1000 kJ/mol.

5. Which is larger, a chlorine atom or a chloride ion, \(\mathrm{Cl^-}\)?
   A. Cl atom
   B. Cl- ion
   C. They are the same size

   Answer: Cl- ion. The ion has 18 electrons held by the same 17 protons. The extra electron adds repulsion, so the ion spreads out: 181 pm vs 102 pm.

## Frequently asked questions

### What is effective nuclear charge?

It is the net pull an outer electron feels from the nucleus after inner electrons block part of it. A simple estimate is protons minus inner (core) electrons. For sodium it is \(11 - 10 = 1\), and for chlorine \(17 - 10 = 7\), which is why chlorine holds its outer electrons much more tightly.

### Why do noble gases have the highest ionization energies?

A noble gas has the most protons in its period while its outer electrons are still in the same shell, so they feel the strongest pull. Removing one would also break a very stable filled shell. Helium has the highest first ionization energy of any element, 2372 kJ/mol.

### What is the difference between ionization energy and electron affinity?

Ionization energy is the energy needed to remove an electron from a neutral gas atom. Electron affinity is the energy change when a neutral atom gains an electron. Electron affinity follows the same rough trend but has more exceptions, so many high school courses focus on ionization energy.

## Sources

- [OpenStax Chemistry 2e, 6.5 Periodic Variations in Element Properties](https://openstax.org/books/chemistry-2e/pages/6-5-periodic-variations-in-element-properties), accessed 2026-10-01
- [NIST Atomic Spectra Database, ionization energies](https://physics.nist.gov/PhysRefData/ASD/ionEnergy.html), accessed 2026-10-01
- [PubChem Periodic Table (Pauling electronegativity)](https://pubchem.ncbi.nlm.nih.gov/periodic-table/), accessed 2026-10-01
- [Cordero et al., Covalent radii revisited, Dalton Transactions 2008](https://doi.org/10.1039/B801115J), accessed 2026-10-01
- [Shannon, Revised effective ionic radii, Acta Crystallographica A 1976](https://doi.org/10.1107/S0567739476001551), accessed 2026-10-01

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## Try asking Ducky

- "Check my ranking on question 3. Did I put the ions in the right order?"
- "Explain why aluminum has a lower ionization energy than magnesium, using my orbital diagram."
- "Give me five quick ranking questions and tell me which trend each one tests."

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