Explaining Periodic Trends

Textbook Pages: 31-40

Periodic TrendsO Atomic RadiusO Ionization EnergyO Electron AffinityO Electronegativity

O When examining Periodic Trends ALWAYS look at:O Number energy levelsO Number of protons

O As we go down a group: O Energy levels increase

O As we go across a period (L to R): O Number of protons increases

Atomic RadiusO Definition: The distance from the center of an atom to

the boundary within which the electrons spend 90% of their time.

O Ex: Which atom would be larger: Be or Mg? Explain.O Be has 2 energy levelsO Mg has 3 energy levels

O Therefore, Mg is larger

O Ex: Which atom would have the smallest radius: Mg or Si? Explain.O Mg and Si both have 3 energy levelsO Si has more protons to attract the electrons

O Therefore, Si is smaller (smallest radius)

Atomic RadiusO In General:

O Down a Group:O Atomic Radius increases (more energy

levels)

O Across a Period (L to R):O Atomic Radius decreases (same energy

levels, more protons)

O The atomic radius of Bromine is larger than that of Nitrogen. Why do you think this is so?

Ionization EnergyO Definition: The amount of energy required to

remove an electron from the outermost energy level of an atom or ion (in the gaseous state).

O Ex: Which atom has the larger ionization energy: F or O? Why?O Both F and O have the same energy levelsO F has more protons to attract the same number

of energy levelsO F will hold the electrons more tightlyO Therefore, F has a larger ionization energy

Ionization EnergyO Ex. Would more energy be required to remove an electron

from a Ne atom or from a F ion?O Both F and Ne have the same number of energy

levelsO Ne has more protons to attract the electrons,

making them more difficult to removeO Therefore, Ne would require more energy to

remove an electron (higher ionization energy)O Ex: Which atom has the smallest ionization energy: Li or

Rb? Explain.O Li has 2 energy levels

O Holds its electrons more tightly because they are closer to the nucleus

O Rb has 5 energy levelsO Takes less energy to remove the outermost

electron because it is farther from the nucleusO Therefore, Rb has the smallest ionization

energy

Ionization EnergyO In General:

O Down a Group:O Ionization Energy decreases (more

energy levels)

O Across a Period (L to R):O Ionization Energy increases (same

energy levels, more protons)

Electron AffinityO Definition: Ability of an atom to attract

electrons.

O In General:O Down a Group:

O Electron Affinity decreases (more energy levels)

O Across a Period (L to R):O Electron Affinity increases (same energy

levels, more protons)

ElectronegativityO Definition: Ability of an atom to attract

electrons in a bond.

O In General:O Down a Group:

O Electronegativity decreases (more energy levels)

O Across a Period (L to R):O Electronegativity increases (same energy

levels, more protons)

Ionization EquationsO Atoms will lose or gain electrons so that they

are isoelectronic (have the same number of electrons) with the nearest noble gas

O This makes the atoms stable because their orbits are full (stable octet)

O Examples:O F + e− F−

O Mg 2e− + Mg2+

O Isoelectronic with Ne (noble gas)

Summarizing Trends in the Periodic Table

OTake a look at Page 38 in your textbook!!!