Periodic Properties of the Elements

1
Periodic Properties of the Elements

• Early versions of the Periodic table were
  constructed by Mendeleev and Meyer.
• We now know that the periodic properties are due
  to the electronic structure of atoms.
• Electronic structure explains the observed trends
  in
• Atomic size
• Ionization Energy
• Electron Affinity

2
Electron Shells

• Electrons in successive shells (n 1,2,3...)
  produce maxima in the electron density at
  increasing distances from the nucleus.
• The inner, completed electron shells are called
  core shells.
• The outermost electron shell is called the
  valence shell. It may include subshells of inner
  shells.
• Example The valence shell of 33As includes 4s,
  3d, and 4p subshells.

3
Atomic Size

• The radius of an atom is found from
  the distance between nuclei in
  a molecule.

4
Trends in Atomic Size

• Size increases going down a column of the
  periodic table.
• Size decreases from left to right in a row.

5
Trends in Atomic Size

• In a column, size increases with the addition of
  successive electron shells.
• Example The alkali metals

6
Effective Nuclear Charge

• The size of an orbital decreases with increasing
  Zeff for the valence electrons.

7
Effective Nuclear Charge

• The size of an orbital decreases with increasing
  Zeff for the valence electrons.

Be
4
8
Effective Nuclear Charge

• The size of an orbital decreases with increasing
  Zeff for the valence electrons.

B
5
9
Effective Nuclear Charge

• The size of an orbital decreases with increasing
  Zeff for the valence electrons.

C
6
10
Effective Nuclear Charge

• The size of an orbital decreases with increasing
  Zeff for the valence electrons.

O
8
11
Effective Nuclear Charge

• The size of an orbital decreases with increasing
  Zeff for the valence electrons.

F
9
12
Ionization Energy

• The energy required to remove an electron from an
  atom in its ground level.
• Example Hydrogen
• H(g) ? H(g) e?
• I ?E

1312 kJ/mol
13
Ionization Energy

• Energies can be defined for successive
  ionizations. For Mg
• Mg(g) ? Mg(g) e? I1 738 kJ
• Mg(g) ? Mg2(g) e? I2 1450 kJ
• Mg2(g) ? Mg3(g) e? I3 7730 kJ

For Al
14
Trends in Ionization Energy

• Ionization energy increases with Zeff for the
  valence electrons.
• I1 decreases going down a column of the periodic
  table.
• I1 increases from left to right in a row.

Alkali metals have the lowest ionization energies
in a period. Rare gases have the highest.
15
Trends in Ionization Energy

• Arrange the following with increasing ionization
  energy
• C, K, Mg, Na, Ne, Si
• K lt Na lt Mg lt Si lt C lt Ne
• Why is the ionization energy of N greater
  than O?
• The 2p subshell of N has a slightly higher
  energy due to electron repulsion.

16
Electron Affinity

• The change in energy when an electron is added to
  a gaseous atom.
• Cl(g) e? ? Cl?(g) ?E -349 kJ/mol
• A large, negative ?E indicates strong attraction
  between the atom and the added electron.
• A positive ?E indicates the addition of an
  electron is unfavorable.
• Ne(g) e? ? Ne?(g) ?E 40 kJ/mol

17
Metals and Nonmetals

• Metals are characterized by low ionization
  energy Nometals by high electron affinity.

18
Group 1A Alkali Metals

• Alkali metals react to lose electrons.
• These metals are very reactive with reactivity
  increasing down the group.
• Reaction with H2O
• Na(s) H2O(l) ??NaOH(aq) H2(g)

19
Groups 2A and 7A

• The alkaline earth metals have higher ion-ization
  energies than alkali metals and are less
  reactive.
• Ca H2O ? Ca(OH)2 H2

The halogens react to gain electrons. Reactivity
decreases going down the group.
Cl2 2 Br?? 2 Cl? Br2