Periodicity and atomic structure overview

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How does first ionisation energy change across a period?

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First ionisation energy generally increases across a period because stronger nuclear attraction makes electron removal harder.

Key concepts

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Definition of periodicity

Periodicity describes the repeating pattern of chemical and physical properties of elements when they are ordered by increasing proton number. Similar properties recur at regular intervals because of repeating electron configurations across periods. Limiting factors include sub-level filling and transition-metal behaviours that produce deviations from simple patterns.

Periods and groups

Periods are horizontal rows where principal energy levels increase by one across the row. Groups are vertical columns where outer-electron configurations are similar, causing similar chemical properties. Elements in the same group show similar valence electron counts; therefore, they show comparable bonding and reactivity under the same conditions.

Atomic structure controls periodic properties

Nuclear charge, electron shielding and the principal quantum number of the outer electrons control atomic size and reactivity. Higher nuclear charge increases attraction on electrons and reduces atomic radius; increased shielding and higher principal quantum number place outer electrons further from the nucleus and increase atomic radius. Cause leads to predictable effects on measurable properties.

Atomic radius trend

Across a period, atomic radius decreases because proton number increases while shielding remains approximately constant, causing stronger attraction of electrons and smaller atoms. Down a group, atomic radius increases because electrons occupy higher principal energy levels and shielding increases, reducing effective nuclear attraction on outer electrons.

Ionisation energy trend

First ionisation energy is the energy required to remove one electron from a gaseous atom. Across a period, first ionisation energy generally increases because increased nuclear charge holds electrons more tightly. Down a group, first ionisation energy generally decreases because outer electrons are further from the nucleus and experience greater shielding, so removal requires less energy. Exceptions arise from subshell arrangements and paired electrons.

Electronegativity and electron affinity

Electronegativity measures an atom's tendency to attract electrons in a bonded context and generally increases across a period and decreases down a group. Electron affinity describes the energy change when an atom gains an electron; trends follow similar causes but show more variation due to electronic subshell stability and pairing effects.

Metallic character

Metallic character reflects ease of losing electrons. Metallic character decreases across a period as atoms hold electrons more tightly and increases down a group as atoms lose electrons more readily due to larger radii and increased shielding. Transition metals and d-block effects modify simple group/period trends.

Key notes

Important points to keep in mind

Periodicity results from repeating electron configurations when elements are ordered by proton number.

Atomic radius decreases left to right across a period because nuclear charge increases with similar shielding.

Atomic radius increases down a group because electrons occupy higher principal energy levels and shielding rises.

First ionisation energy generally rises across a period and falls down a group due to nuclear charge and shielding effects.

Electronegativity generally increases across a period and decreases down a group.

Subshell filling, electron pairing and d-block electrons cause exceptions to simple periodic trends.

Effective nuclear charge (nuclear charge minus shielding) predicts many periodic trends.

Metallic character links to ease of losing electrons: decreases across a period, increases down a group.