Periodicity

Periodicity - the trends in properties of elements across a period and down a group.

• Atmoic radius

• Ionic radius

• Ionization energy

• Electronegativity

• Electron afinitniy

Effective Nuclear Charge - the attraction of the positively charged nucleus acting on the valence electrons, taking into account electron shielding from core electrons

Atomic Radius

Used to describe the size of an atom.

Across the period: decreases

• increase of # of protons and # of electrons (in the same valence shell)

  • nuclear charge increases

  • no change in electron shielding

Down the group: increases

• increasing number of main electron shells where valence electrons are added to higher energy levels (further from nucleus)

• increasing number of inner shells

  • increases electron shielding effect (reduces attraction to valence electron from nucleus

Ionic Radius

• Radius of cation < Radius of parent atom

  • same # of protons, but less electrons —> smaller radius

  • less electron-electron repulsion

• Radius of anion > Radius of parent atom

  • greater electron-electron repulsion

• Isoelectronic ions: Ion with more protons will have a smaller radius

Ionization Energy

The energy required to remove one electron from each atom in one mole of gaseous atom

Across the period: increases

• atomic radius decreases, nuclear charge increases

  • valence electrons are closer to the nucleus

Down the group: generally decreases

• atomic radius increases down a group, so the valence electron is further away (more shielding)

Electronegativity

The relative attraction that an atom has for the shared electron in a covalent bond

Across the period: increases

• # of protons increase → nuclear charge increases

• atomic radius decreases, so nuclear attraction to electrons in the bond becomes stronger

Down the group: decreases

• atomic radius increases, so valence electrons are in a higher energy level

• nuclear attraction to electrons in the bond becomes weaker

Relationship between electronegativity difference and bond character

• Higher difference in electronegativity = higher ionic character

• Lower difference in electronegativity = higher covalent character

Electron Affinity

The enthalpy change when an electron is added to an isolated atom in a gaseous state

The energy released when one electron is added to each atom in one mole of gaseous atoms

• Stronger affinity for a free electron → more energy will be released

• Always a negative value

  • energy is released (exothermic reaction)

  • the more negative, the more exothermic

Across the period: stronger (more energy)

• # of protons increases

• stronger attraction between valence electrons and the nucleus

• Halogens (Group 17) are the strongest (most negative)

  • Gaining one more electron → noble gas (stable)

• Exception: when an electron is added to half-filled orbitals, it will be less exothermic, since the orbital already contains an electron

Down the group: weaker

• number of energy level increases → atomic radius increases

• weaker bond between added electrons and the nucleus

• Exception: in very small atoms, the electron affinity may be less exothermic, because of the electron-electron repulsions

  • F has less electron affinity than Cl, since F is so small that when an electron is added, there are electron-electron repulsions

Summary of Trends