Chemistry Clarity · chemistryclarity.com

Periodic Trends

Why do atomic radius, ionization energy and electronegativity change across the periodic table?

IntermediateAtomic Structure & PeriodicityLast reviewed 3 October 2026
Download lesson (PDF)

What is it?

Periodic trends are the regular, predictable changes in the properties of elements as you move across a period (left to right) or down a group (top to bottom) of the periodic table. The four most important are:

  • Atomic radius: the size of an atom.
  • First ionization energy: the energy needed to remove one electron from each atom in a mole of gaseous atoms.
  • Electronegativity: how strongly an atom attracts the shared electrons in a chemical bond.
  • Ionic radius: the size of an ion.

Key idea

Two factors explain almost every trend: the number of shells (more shells means outer electrons are farther from the nucleus) and the effective nuclear charge (the net positive charge an outer electron actually feels).

Why does it matter?

  • It predicts reactivity. Group 1 metals get more reactive down the group because their outer electron is easier to remove.
  • It predicts bonding. The difference in electronegativity between two atoms tells you whether a bond is nonpolar, polar or ionic.
  • It organises facts. Instead of memorising the properties of 118 elements, you learn a few patterns and the reasons behind them.

How does it work?

1. Effective nuclear charge

Inner electrons shield outer electrons from the full nuclear charge. A simple estimate:

Zeff≈Z−SZ_\text{eff} \approx Z - S

where ZZ is the number of protons and SS is the number of inner (core) electrons. For sodium, Zeff≈11−10=1Z_\text{eff} \approx 11 - 10 = 1; for chlorine, in the same period, Zeff≈17−10=7Z_\text{eff} \approx 17 - 10 = 7.

  • Across a period: protons are added but the core stays the same, so ZeffZ_\text{eff} increases.
  • Down a group: a new shell is added each period, so outer electrons are farther away.

2. Atomic radius

  • Decreases across a period: a higher ZeffZ_\text{eff} pulls the same shell closer.
  • Increases down a group: each period adds a shell.

3. First ionization energy

  • Increases across a period: outer electrons are held more tightly.
  • Decreases down a group: the outer electron is farther away and more shielded. Group 1: Li 520.2, Na 495.8, K 418.8, Rb 403.0, Cs 375.7 kJ/mol.

4. Electronegativity

  • Increases across a period and decreases down a group, for the same reasons.
  • Fluorine is the most electronegative element (3.98 on the Pauling scale); caesium and francium are the least.

5. Ionic radius

  • Cations are smaller than their atoms: often a whole shell is lost, and the remaining electrons feel more nuclear pull.
  • Anions are larger than their atoms: the extra electrons repel each other, and the same nuclear charge holds more electrons.
  • In an isoelectronic series (same number of electrons), more protons means a smaller ion: OX2−>FX−>NaX+>MgX2+\ce{O^2-} > \ce{F-} > \ce{Na+} > \ce{Mg^2+} (all have 10 electrons).

Think of it like this

Think of the nucleus as a magnet and the outer electrons as paper clips. Across a period the magnet gets stronger while the clips stay at the same distance, so they are held more tightly. Down a group the clips sit farther away, behind more layers of other clips, so they are held more loosely.

More precisely

Ionization energy does not rise perfectly smoothly across a period. In period 2 there are two dips. Boron (800.6 kJ/mol) is lower than beryllium (899.5) because its outer electron is in a 2p orbital, slightly higher in energy than 2s. Oxygen (1313.9) is lower than nitrogen (1402.3) because oxygen’s fourth 2p electron is paired, and repulsion between the paired electrons makes one of them easier to remove.

Visualise it

Two outline periodic tables. The first shows atomic radius increasing towards the left and down a group, so the largest atoms are at the bottom left. The second shows ionization energy and electronegativity increasing towards the right and up a group, so the highest values are at the top right. Main-group trends; noble gases have no usual electronegativity.
Atomic radius grows towards the bottom left; ionization energy and electronegativity grow towards the top right.
Bar chart of first ionization energies in kJ/mol for period 2: Li 520.2, Be 899.5, B 800.6, C 1086.5, N 1402.3, O 1313.9, F 1681.0, Ne 2080.7. The overall trend increases from left to right, with dips at boron and oxygen highlighted.
First ionization energies across period 2. The general rise has two dips, at B and O.

Worked example

Worked example: Ranking atomic radius

Question: Arrange Na, Mg, K and Cl in order of increasing atomic radius.

  1. Na, Mg and Cl are in period 3, where radius decreases from left to right: Cl < Mg < Na.
  2. K is directly below Na in group 1, with one more shell: K > Na.
  3. Order: Cl < Mg < Na < K

Worked example: Explaining an ionization-energy trend

Question: Why is the first ionization energy of sodium (495.8 kJ/mol) lower than that of lithium (520.2 kJ/mol)?

  1. Na is [Ne] 3s¹ and Li is [He] 2s¹. Both have one outer s electron with Zeff≈1Z_\text{eff} \approx 1.
  2. Sodium’s outer electron is in the third shell, farther from the nucleus and shielded by more inner electrons.
  3. It is held less strongly, so less energy is needed to remove it.

Common mistake

Common mistake: Thinking more protons make a bigger atom across a period

Across a period atoms get smaller, even though they have more protons and electrons. The added electrons go into the same shell, while the extra protons pull that shell in more tightly.

Common mistake: Confusing ionization energy with electronegativity

Ionization energy is about removing an electron from an isolated gaseous atom. Electronegativity is about attracting shared electrons in a bond. They follow similar trends but measure different things.

Common mistake: Assuming an ion is the same size as its atom

NaX+\ce{Na+} is much smaller than Na (it has lost its whole third shell), and ClX−\ce{Cl-} is larger than Cl.

Notation note

  • Ionization energies are usually given in kJ/mol (sometimes in eV per atom).
  • Electronegativity has no units; the Pauling scale is the most common.
  • Zeff≈Z−SZ_\text{eff} \approx Z - S is a simple estimate. More detailed methods (such as Slater’s rules) give values that are not whole numbers.

Remember this

Remember this

  • Across a period: ZeffZ_\text{eff} increases, radius decreases, ionization energy and electronegativity increase.
  • Down a group: more shells, radius increases, ionization energy and electronegativity decrease.
  • Cations are smaller than their atoms; anions are larger.
  • Period 2 ionization-energy dips: B (2p electron) and O (paired 2p electron).

Test yourself

Check your understanding before moving on.

Flashcards

10 cards

  1. Question
    How does atomic radius change across a period and down a group?
    Answer

    Decreases across a period (left to right); increases down a group.

  2. Question
    What is effective nuclear charge?
    Answer

    The net positive charge felt by an outer electron, roughly Zeff≈Z−SZ_\text{eff} \approx Z - S (protons minus inner electrons).

  3. Question
    What is first ionization energy?
    Answer

    The energy needed to remove one electron from each atom in a mole of gaseous atoms.

  4. Question
    How does first ionization energy change across a period and down a group?
    Answer

    Increases across a period; decreases down a group.

  5. Question
    Which element is the most electronegative?
    Answer

    Fluorine (3.98 on the Pauling scale).

  6. Question
    Is a cation larger or smaller than its atom?
    Answer

    Smaller. It often loses a whole shell, and the remaining electrons are pulled in more strongly.

  7. Question
    Is an anion larger or smaller than its atom?
    Answer

    Larger. The added electrons repel each other, and the same nuclear charge holds more electrons.

  8. Question
    Order by size: NaX+\ce{Na+}, FX−\ce{F-}, MgX2+\ce{Mg^2+}, OX2−\ce{O^2-}
    Answer

    OX2−>FX−>NaX+>MgX2+\ce{O^2-} > \ce{F-} > \ce{Na+} > \ce{Mg^2+} (all have 10 electrons; more protons means a smaller ion)

  9. Question
    Why is the first ionization energy of B lower than that of Be?
    Answer

    Boron's outer electron is in a 2p orbital, which is higher in energy than 2s, so it is easier to remove.

  10. Question
    Why is the first ionization energy of O lower than that of N?
    Answer

    Oxygen's fourth 2p electron is paired; repulsion between the paired electrons makes one of them easier to remove.

Quiz

7 questions

  1. Question 1EasyWhich atom has the largest atomic radius?
    Show answer

    Answer: K

    Radius increases down a group. K is below Li and Na in group 1, and Cl is far to the right in period 3.

  2. Question 2EasyWhich element has the highest first ionization energy?
    Show answer

    Answer: F

    Ionization energy increases across a period and decreases down a group. F is furthest right in period 2.

  3. Question 3MediumWhy do atoms get smaller across a period?
    Show answer

    Answer: Effective nuclear charge increases while the outer shell stays the same

    Protons are added but the number of inner electrons stays the same, so the outer shell is pulled in more tightly.

  4. Question 4MediumWhich ion is the smallest?
    Show answer

    Answer: MgX2+\ce{Mg^2+}

    All four have 10 electrons. Mg²⁺ has the most protons (12), so it pulls its electrons in most tightly.

  5. Question 5MediumWhich is the correct order of increasing electronegativity?
    Show answer

    Answer: K < Na < Li < F

    Electronegativity decreases down group 1 (K 0.82 < Na 0.93 < Li 0.98), and fluorine is the highest of all (3.98).

  6. Question 6HardThe first ionization energy of oxygen is lower than that of nitrogen. Why?
    Show answer

    Answer: Oxygen has a paired 2p electron, and the repulsion makes it easier to remove

    N (2p³) has three unpaired electrons; in O (2p⁴) one orbital holds a pair, and the extra repulsion lowers the ionization energy.

  7. Question 7MediumUsing Zeff≈Z−SZ_\text{eff} \approx Z - S, estimate the effective nuclear charge felt by a valence electron of chlorine (Z = 17).
    Show answer

    Answer: 7

    Chlorine has 10 inner electrons (a [Ne] core), so Zeff≈17−10=7Z_\text{eff} \approx 17 - 10 = 7.

Notes and downloads

  • Worksheet

    Periodic Trends Worksheet

    9 questions on atomic radius, ionization energy, electronegativity and ionic radius. Answer key included.

    IntermediateFree

References

  1. Brown, T. L.; LeMay, H. E., Jr.; Bursten, B. E.; Murphy, C. J.; Woodward, P. M.; Stoltzfus, M. W. Chemistry: The Central Science, 15th ed.; Pearson, 2022.

Practise this topic with flashcards and a quiz at chemistryclarity.com/chemistry/periodic-trends/

Spotted a mistake? Let us know and we'll fix it.