Periodic trends

Periodicity and Electron Configuration

The periodic table is organized by recurring patterns of properties that repeat at regular intervals because ground-state electron configurations themselves repeat as atomic number increases.

Elements in the same group (vertical column) share the same valence-shell electron configuration, so they display similar chemical behavior. Across a period (horizontal row), electrons fill the same principal energy level while nuclear charge rises one proton at a time. This repetition of valence configurations is what chemists call periodicity — the orderly recurrence of properties that the table is built to reveal.

The pattern is driven by completely or partially filled shells and subshells. A completely filled valence shell, as in the noble gases, is unusually stable, which makes those elements chemically unreactive. A partially filled valence subshell, by contrast, gives an atom characteristic reactivity because it can gain, lose, or share electrons to approach a more stable configuration.

Out of ScopeWriting electron configurations for elements that are exceptions to the aufbau principle. Do not memorize this for the exam.

Effective Nuclear Charge and Shielding

The net pull felt by a valence electron, called effective nuclear charge ($Z_{eff}$), underlies every periodic trend and is interpreted through Coulomb’s law and the shell model.

In the shell model, inner (core) electrons occupy energy levels between the nucleus and the valence electrons. These core electrons repel the valence electrons and partly cancel the nuclear charge, an effect called shielding [screening of outer electrons from the full nuclear pull]. As a result, a valence electron does not feel the full charge of all protons — it feels a smaller, effective charge.

Coulomb’s law explains why this matters. The attraction between the nucleus and an electron strengthens as charge increases and as distance decreases:

$${E}_{coulombic}∝\frac{{q}_{1}{q}_{2}}{r}$$

Across a period, protons are added while the number of core electrons stays the same, so $Z_{eff}$ increases and electrons are held more tightly. Down a group, valence electrons occupy shells of higher principal quantum number, lying farther from the nucleus with more shielding, so distance dominates even though $Z_{eff}$ rises slightly.

A simple approximation estimates $Z_{eff}$ for a valence electron by subtracting the shielding core electrons from the atomic number.

Equation used

$${Z}_{eff}=Z-S$$

Where: $Z_{eff}$ = effective nuclear charge (no units), Z = atomic number, number of protons (no units), S = number of shielding core electrons (no units).

A worked example estimates the effective nuclear charge experienced by a valence electron in a chlorine atom (Z = 17), whose core electrons fill the 1s, 2s, and 2p subshells.

Given

$$Z=17$$

$$S=10$$

Working

$${Z}_{eff}=Z-S$$

$${Z}_{eff}=17-10$$

Answer

$${Z}_{eff}=+7$$

MisconceptionStudents confuse atomic number Z with effective nuclear charge $Z_{eff}$. Z counts all protons; $Z_{eff}$ is the reduced pull a valence electron actually feels after shielding. Exam tip: Always subtract core electrons before reasoning about a trend.

Practice Problem: Estimate the effective nuclear charge felt by the valence electron of a sodium atom (Z = 11), which has 10 core electrons.

Answer

Equation used

$${Z}_{eff}=Z-S$$

Where:

$Z_{eff}$ = effective nuclear charge (no units),

Z = atomic number (no units),

S = number of shielding core electrons (no units).

Given

$$Z=11$$

$$S=10$$

Working

$${Z}_{eff}=11-10$$

Answer

$${Z}_{eff}=+1$$

Atomic and Ionic Radii

Atomic size follows a clear pattern: atomic radius decreases across a period and increases down a group, set by the competition between effective nuclear charge and principal energy level.

Across a period, $Z_{eff}$ increases, so the nucleus pulls the valence electrons inward and the radius shrinks. Down a group, each new period adds a shell of higher principal quantum number, placing valence electrons farther from the nucleus, so the radius grows despite the larger nuclear charge.

Ion formation changes the radius predictably:

In an isoelectronic series [ions with identical electron counts], the radius decreases as nuclear charge increases, because more protons pull the same number of electrons more tightly.

Illustrative ExampleNa⁺, Mg²⁺, and Al³⁺ are isoelectronic with neon; their radii decrease in that order because nuclear charge rises while electron count stays at 10.

Ionization Energy

Electron removal becomes harder as binding strengthens: ionization energy increases across a period and decreases down a group, mirroring changes in effective nuclear charge and atomic radius.

Ionization energy is the energy required to remove an electron from a gaseous atom. Across a period, rising $Z_{eff}$ and shrinking radius bind valence electrons more tightly, so more energy is needed to remove one. Down a group, valence electrons lie farther out and are more heavily shielded, so they are removed more easily and ionization energy falls.

Small deviations from the smooth increase appear within a period. They arise because a filled or half-filled subshell is slightly more stable, so removing an electron from the element just past that configuration requires less energy than the overall trend predicts.

Examiner InsightFRQs frequently ask students to explain a dip in first ionization energy. Tie the answer to subshell stability or $Z_{eff}$ — never to “the atom wanting a full shell.” Exam tip: Always anchor explanations in Coulombic attraction and electron configuration.

Graphing Activity: First Ionization Energy Across Period 3 Data:

Task: Construct a scatter or line graph from the data above. Include: (1) labeled axes with units, (2) appropriate scaling, (3) a title. Do not describe the finished graph — construct it yourself, then mark the two points that dip below the overall upward trend.

Electron Affinity and Electronegativity

An atom’s pull on electrons strengthens together for two properties: electron affinity and electronegativity both generally increase across a period and decrease down a group, because rising effective nuclear charge tightens an atom’s grip on electrons.

Electron affinity is the energy change when a gaseous atom gains an electron; a more negative value means electron addition is more favorable. Electronegativity is the ability of an atom in a bond to attract the shared bonding electrons. Both depend on the same Coulombic reasoning.

Across a period, higher $Z_{eff}$ and a smaller radius let the atom attract an extra or shared electron more strongly. Down a group, the added or shared electron sits farther from the nucleus with more shielding, weakening the attraction. Fluorine, with high $Z_{eff}$ and a small radius, has the greatest electronegativity.

MisconceptionStudents assume every atom releases energy on gaining an electron. Noble gases and full subshells resist electron addition, so their electron affinities are unfavorable. Exam tip: Treat noble gases as exceptions when comparing electron affinities.

Predicting Properties Using Periodicity

Because trends repeat dependably, periodicity allows estimation of unknown properties from neighboring elements even when no measured data is available.

By locating an element on the table, the relative size of its atomic radius, ionization energy, electron affinity, and electronegativity can be predicted from its row and column. Values can also be interpolated between known neighbors. For example, an unknown element’s atomic radius can be bracketed by the radii of the elements directly above and below it, or to its left and right, because the property changes smoothly along each direction.

This predictive power is grounded in the same Coulombic logic that produces the trends: knowing how $Z_{eff}$ and distance change across the table lets a chemist estimate a missing value with confidence and justify it mechanistically rather than by memorized fact.

Diagram: Periodic Trend Arrows Show a blank periodic table outline with directional arrows for atomic radius, ionization energy, electronegativity, and electron affinity. Each arrow must be labeled with the property and the direction of increase, and the diagram must mark fluorine as the most electronegative element. Exam tip: Memorize one master arrow (increase toward the upper right for IE, EN, EA; radius is the reverse).