Periodic Trends

Atomic Radiuse

• decreases across a row

  • REASON: nuclear charge is increasing pulling those electrons in closer to the nucleus

  • the distance of electrons from the nucleus gets smaller because it has more electrons

• increases down a column

  • the shell and the principle quantum number of the shell is getting larger

  • the electrons are farther from the nucleus

• Measure of radius

  • Diffraction technique or from bnod length of a diotomic molecule

    • half of the bond length is the atomic radius

• Importance of atomic radius: selective ion channels in any cells where a very rapid signaling is important

• MORE CHARGE=MORE FORCE

Ionization Energy

To form a positive ion, an electron must be removed from a neutral atom. This requires energy. The energy is needed to overcome the attraction between the positive charge of the nucleus and the negative charge of the electron.

*usually refers to the first ionization energy

Ionization Energy = final energy of the products - the energy of the reactants; energy required to remove an electron from a gaseous atom

first ionization energy = energy to remove the electron from the highest occupied atomic orbitation

second ionization energy= energy to remove the 2s electron for example in boron plus

third ionization energy= remove the next strongly bound electron

• Left to right = ionization energy increases

  • the nuclear charge is increasing

  • when the nuclear charge increases the Coulomb interaction is greater bringing electrons closer to the nucleus (more strongly bound)

• putting electrons into the same shell

  • on the average, the distance of those electrons from the nuclelus is about the same

*2p is higher than 2s in a multi electron atom

• the nuclear charge isn’t high enough to overcome the extra energy to access the 2p state

As you go down

• the ionization energy decreases

  • the principal quantum number is increasing because the shell is larger

  • the distance from the electron from the nucleus is larger

Electronegativity

• essentially the average of the ionization energy and the electron affinity; indicates the relative ability of its atoms to attract electrons in a chemical bond

• high electronegativity= good electron acceptor

  • high electron affinity

  • the energy released in adding an electron is large

  • located in upper right hand corner on the right hand side of the periodic table

  • increase as you go from left to right

• Low electronegativity=electron donor

  • electron affinity is low value

  • located in the lower left hand corner of the periodic table

  • decreases as you move down

• high x low electronegativity = ionic bonds

  • want to attract the electron from the lower corner

Ionic Radius

isoelectronic= having the same electron configuration

Relative size of the radii of ions vs their neutral partners

- 40% decrease in the radius

• if you add an electron to iodine to form iodine radius

  • increase in about 55% in the atomic radius