Periodicity

Trends in the Periodic Table

Atomic Radius• One half the distance between nuclei of identical

atoms that are bonded together• Decrease across a period due to increasing positive

nuclear charge• Increase down a group due to increasing number of

energy levels (outer electrons are farther from the nucleus)

Ionization Energy

Ion: An atom or group of atoms that has a positive (cation) or negative (anion) charge

Ionization: The formation of an ionIonization Energy: The energy

required to remove one electron from a neutral atom of an element, measured in kilojoules/mole (kj/mol)

Ionization Energy• IE increases for each successive

electron• Each electron removed experiences a

stronger effective nuclear charge• Greatest increase in IE comes when

trying to remove an electron from a stable, noble gas configuration.

Ionization Energy1. IE tends to increase across each period.

• Atoms are getting smaller, electrons are closer to the nucleus

2. IE tends to decrease down a group• Atoms are getting larger, electrons are farther from the

nucleus• Outer electrons becomes increasingly more shielded

from the nucleus by inner electrons

3. Metals have characteristically low IE4. Nonmentals have high IE5. Noble gases have a very high IE

Ionic Radius• Cations have a smaller ionic radius

than corresponding atom• Protons outnumber electrons• Less shielding of electrons

• Anions have a larger ionic radius than corresponding atom

• Electrons outnumber protons• Greater electron-electron repulsion

• Ion size decreases down a group

Electron AffinityEnergy change that occurs when an atom accepts an electron1. Halogens have the highest electron affinities2. Metals have low electron affinities3. Affinity tends to increase across a period and decrease

down a group

ElectronegativityThe ability of an atom to attract electrons when bonded1. Nonmetals have high electronegativity2. Metals have low electronegativity3. Electronegativity increases across a period and

decreases down a group

Electronegativity

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