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Periodic Table Trends
Electronegativity, Atomic Radius, and Ionization Energy
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The periodic table is arranged so that elements with similar properties appear in the same column (group). Moving across a period (row) or down a group, measurable properties - atomic radius, ionization energy, electronegativity, and electron affinity - change in predictable patterns. These trends emerge from the interplay between the number of protons (nuclear charge) and the shielding effect of inner electron shells.
Understanding these trends lets you predict chemical behavior without memorizing individual elements. A more electronegative element will attract bonding electrons more strongly. A larger atomic radius usually means lower ionization energy. The trends have exceptions - noble gases and certain transition metals - but the patterns hold for the vast majority of main-group elements.
Key Facts
- Atomic radius increases down a group (more electron shells) and decreases across a period (more protons pull electrons closer).
- Ionization energy: energy to remove the outermost electron - increases across a period, decreases down a group.
- Electronegativity: ability to attract bonding electrons - increases across a period and up a group. Fluorine is the most electronegative element.
- Electron affinity: energy change when gaining an electron - generally increases across a period (more exothermic).
- Key exception: noble gases have very high ionization energy and essentially zero electron affinity.
- Key exception: nitrogen has higher IE than oxygen due to its half-filled stable 2p subshell.
Vocabulary
- Atomic radius
- Half the distance between the nuclei of two identical atoms in a bond; a measure of atomic size.
- Ionization energy (IE)
- The energy required to remove the outermost electron from a neutral atom in the gas phase.
- Electronegativity
- A measure of an atom's ability to attract shared electrons in a covalent bond. Measured on the Pauling scale.
- Electron affinity
- The energy change when a neutral atom gains one electron. Negative values indicate exothermic processes.
- Effective nuclear charge
- The net positive charge experienced by an electron after accounting for shielding by inner electrons.
Common Mistakes to Avoid
- Thinking atomic radius always decreases going right. Across a period, electrons enter the same shell while proton count increases - they're pulled in closer, so radius decreases. But across periods where new shells begin (going down), radius increases.
- Assuming ionization energy always increases right. The nitrogen exception (higher IE than oxygen) trips up many students - always note the N and Cu/Zn exceptions.
- Confusing electronegativity with electron affinity. Electronegativity applies in bonding situations; electron affinity is measured for isolated atoms gaining electrons.
- Ignoring shielding when explaining trends. Inner electrons shield outer electrons from the nuclear charge, which is why atoms get larger going down a group even though nuclear charge increases.
Practice Questions
- 1 Place these in order from smallest to largest atomic radius: Na, K, Mg, Al.
- 2 Why does fluorine have a higher electronegativity than oxygen, even though oxygen has a greater electron affinity?
- 3 Explain why the second ionization energy is always larger than the first ionization energy for any element.