Periodic Trends

Learning Objective 1.7.A Explain the relationship between trends in atomic properties of elements and electronic structure and periodicity.

Quick Notes

Periodicity: Recurring patterns in element properties due to repeating electron configurations in the periodic table.
Key trends: Ionization energy, atomic radius, ionic radius, electron affinity, electronegativity.
Coulomb’s Law: Attraction between opposite charges increases with greater nuclear charge (more protons) and decreases with larger distance (more shells).
Effective nuclear charge (Z_eff): Net positive charge felt by valence electrons = nuclear charge − shielding. Higher Z_eff → stronger attraction → smaller atom, higher ionization energy.
Ionization energy:
- Increases across a period – more protons, same shielding, stronger pull on electrons.
- Decreases down a group – more shells, greater distance and shielding.
Atomic radius:
- Decreases across a period – increasing Z_eff pulls electrons closer.
- Increases down a group – more shells increase size despite higher nuclear charge.
Ionic radius:
- Cations smaller than atoms (less electron repulsion).
- Anions larger than atoms (more electron repulsion).
- Across Period 3: Na⁺ → Al³⁺ decrease; P³⁻ → Cl⁻ decrease; big jump between cations and anions.
Electron affinity:
- More negative (favorable) across a period – greater nuclear attraction for extra electron.
- Less negative down a group – electrons added further from nucleus.
- Exceptions: filled sublevels (e.g., Mg, Ar) have low affinity.
Electronegativity:
- Increases across a period – smaller size, higher Z_eff.
- Decreases down a group – more shells, weaker attraction.
- Highest = F (4.0), lowest = Fr (~0.7–1.0).
Trends explained by electron configuration, nuclear charge, shielding, and distance of valence electrons from nucleus.

Full Notes

The periodic table is more than just a list of elements — it is structured to show repeating patterns in chemical and physical properties, known as periodicity. These patterns are explained by electron configurations and the arrangement of electrons in energy levels and sublevels.

Why Periodicity Occurs

AP Chemistry periodic table labelled with periods and groups to illustrate recurring trends in properties

Elements in the same group (vertical column) have similar outer electron configurations, giving them similar chemical properties.

Across a period (horizontal row), the number of protons and electrons increases, affecting atomic properties in predictable ways.

These trends are largely due to:

The number of electron shells (increasing down a group)
The effective nuclear charge (Z_eff) experienced by the outermost electrons
Shielding by inner electrons
The strength of the attraction between the nucleus and valence electrons

Coulomb’s Law and Periodic Trends

Coulomb’s Law:
AP Chemistry Coulombs law showing how force of attraction is proportional to charge of particles and distance between them F ∝ (q1 × q2) / r2 F ∝ (q₁ × q₂) / r²
Where F is the electrostatic force, q₁ and q₂ are the charges, and r is the distance between them.

As nuclear charge (number of protons) increases, the attraction to electrons increases, unless offset by shielding.

Ionization Energy – the energy required to remove an electron from an atom

Increases as a trend* across a period (left to right)

AP Chemistry graph showing general increase in first ionization energy across Period 3 from Na to Ar with small deviations

Decreases down a group (top to bottom)

AP Chemistry diagram showing first ionization energy decreasing down a group due to increased distance and shielding

Explanation:
Across a period, more protons = stronger attraction = harder to remove an electron.
Down a group, more shells = more distance and more shielding = easier to remove an electron.

Atomic Radius – distance from the nucleus to the outermost electron

AP Chemistry comparison of atomic radii across Period 3 elements Na to Ar, showing a general decrease

AP Chemistry schematic showing atomic radius decreasing across a period due to increasing nuclear charge

Decreases across a period

Increases down a group

AP Chemistry diagram showing atomic radius increasing down a group as additional energy levels are added

Explanation:
Across a period, increasing nuclear charge pulls electrons closer.

AP Chemistry illustration comparing sodium and chlorine to show stronger nuclear attraction leading to smaller atomic radius

Down a group, new energy levels are added, making the atom larger.

Ionic Radius – size of an ion

Cations (positive ions) are smaller than their atoms (less electron repulsion).
Anions (negative ions) are larger than their atoms (more electron repulsion).

ExampleIonic Radius Trend Across a period

AP Chemistry diagram showing ionic radii across Period 3: Na⁺, Mg²⁺, Al³⁺ decreasing, then a jump to P³⁻, S²⁻, Cl⁻ with decreasing radii

Metals (Na⁺, Mg²⁺, Al³⁺) lose electrons and form smaller ions than their atoms. The ionic radius decreases from Na⁺ to Al³⁺ as the nuclear charge increases whilst the number of electrons stays the same (the ions all have the same electron configuration), meaning greater attraction from nucleus pulls outer-shell of ion more tightly, decreasing ionic radius.

Non-metals (P³⁻, S²⁻, Cl⁻) gain electrons and form larger ions than their atoms.

There's a noticeable jump in ionic radius between Al³⁺ and P³⁻, due to the switch from cations to anions.

Note that Silicon (Si) and Argon (Ar) don’t readily form ions.

Electron Affinity – energy change that occurs when a gaseous atom gains an electron

Becomes more negative across a period (more favorable).
Becomes less negative (or positive) down a group.

AP Chemistry graph showing general trend of more negative first electron affinity across Period 3 with noted exceptions

Note: There are exceptions to the trend. Both Mg and Ar have electron affinity values close to 0 kJ mol⁻¹, indicating that they do not easily gain electrons. This is because they already have filled sublevels, so incoming electrons experience little attraction to the nucleus.

Electronegativity

Electronegativity is the power of an atom to attract electrons to itself.
It is a relative value and does not have units.

The most commonly used scale is the Pauling scale, where:

Fluorine (F) = 4.0 (most electronegative)
Group 1 metals = ~0.7 to 1.0 (least electronegative)

As a general trend

AP Chemistry periodic table arrows showing electronegativity increasing up and to the right, highest near fluorine

Across a period (left to right):
Electronegativity increases.
Atomic radius decreases and nuclear charge increases.
Shielding stays roughly the same.

Down a group (top to bottom):
Electronegativity decreases.
Atomic radius increases and shielding increases.
Outer electrons are further from the nucleus and less strongly attracted.

Highest: Fluorine
Lowest: Francium

Effective Nuclear Charge (Z_eff)

The effective nuclear charge is the net positive charge felt by an electron after accounting for shielding by inner electrons – it helps explain periodic trends outlined on this page, like atomic size and ionization energy.

AP Chemistry banner stating Z eff equals actual nuclear charge minus shielding from inner electrons

Z_eff = actual nuclear charge – shielding effect from inner electrons

As Z_eff increases, electrons are more tightly held.
Higher Z_eff means a smaller atomic radius and higher ionization energy.
Lower Z_eff means larger radius atomic radius and lower ionization energy.

Worked Example

Which has a higher first ionization energy: sodium (Na) or magnesium (Mg)?

Recall Definitions
First ionization energy is the energy required to remove the outermost electron from a neutral atom in the gas phase.
This depends on effective nuclear charge (Z_eff) and distance of the electron from the nucleus.
Compare Atomic Structures
Sodium (Na): 1s² 2s² 2p⁶ 3s¹
Magnesium (Mg): 1s² 2s² 2p⁶ 3s²
Both atoms have electrons in the same outer shell (n = 3).
Analyze Z_eff
Mg has one more proton than Na (12 vs. 11), but both have the same number of inner (core) electrons.
This means Mg has a greater Z_eff — its nucleus pulls on the valence electrons more strongly.
Apply Reasoning
Because Mg holds its outer electrons more tightly, it takes more energy to remove one.

Answer: Magnesium has a higher first ionization energy than sodium.

Summary

The periodic table is structured around predictable trends in atomic properties. These trends arise from differences in electron configurations, nuclear charge, and shielding.
Understanding periodicity allows chemists to predict chemical behavior and compare elements, even when experimental data is not available.
Key trends to remember:
- Ionization energy: ↑ across a period, ↓ down a group
- Atomic radius: ↓ across a period, ↑ down a group
- Electron affinity: more negative across a period
- Electronegativity: ↑ across a period, ↓ down a group
Use Coulomb’s Law, shell model, and shielding to explain these trends.