1. Periodic Properties of the Elements
Chemistry 2 Honors Jeff Venables Northwestern High School

2. Development of the Periodic Table
As of 2008, there were 117 elements known.
How do we organize 117 different elements in a meaningful way that will allow us to make predictions about undiscovered elements?

3. Certain elements were missing from this scheme.
Arrange elements to reflect the trends in chemical and physical properties.
First attempt (Mendeleev and Meyer) arranged the elements in order of increasing atomic mass.
Certain elements were missing from this scheme.
Example: Mendeleev and Germanium.
Modern periodic table: arrange elements in order of increasing atomic number.
Groups – what do elements have in common?

4. Group – a column (vertical) on the periodic table.
Period - a row (horizontal) on the periodic table. There are 7 periods.
Group – a column (vertical) on the periodic table.
Group1 = alkali metals
Group 2 = alkaline earth metals
Groups 3-12 = transition metals
Group 13 = Boron group
Group 14 = Carbon group
Group 15 = Nitrogen group
Group 16 = Chalcogens
Group 17 = Halogens
Group 18 = Noble gases

5. Periodic Trends in Atomic Radii
As a consequence of the ordering in the periodic table, many properties of elements vary periodically.
As we move down a group, the atoms become larger.
As we move across a period, atoms become smaller.
There are two factors at work:
principal quantum number, n, and
the increasing numbers of protons.

7. Examples – Place each group of elements in order of increasing atomic radius:
S, Al, Cl, Mg, Ar, Na
K, Li, Cs, Na, H
Ca, As, F, Rb, O, K, S, Ga

8. Examples – Place each group of elements in order of increasing atomic radius:
S, Al, Cl, Mg, Ar, Na
Ar < Cl < S < Al < Mg < Na
K, Li, Cs, Na, H
H < Li < Na < K < Cs
Ca, F, As, Rb, O, K, S, Ga
F < O < S < As < Ga < Ca < K < Rb

9. Electron Configurations of Ions
Cations: electrons removed from orbital with highest principle quantum number, n, first:
Li (1s2 2s1)  Li+ (1s2)
Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5)
Anions: electrons added to the orbital with highest n:
F (1s2 2s2 2p5)  F- (1s2 2s2 2p6)

10. Write electron configurations for the following ions:
Al3+
S2-
Li+
Br-
Fe2+
Fe3+

11. Write electron configurations for the following ions:
Al3+ 1s22s22p6
S2- [Ne]3s23p6
Li+ 1s2
Br- [Ar]4s23d104p6
Fe2+ [Ar]3d6
Fe3+ [Ar]3d5

12. Trends in the Sizes of Ions
Ion size is the distance between ions in an ionic compound.
Ion size also depends on nuclear charge, number of electrons, and orbitals that contain the valence electrons.
Cations vacate the largest orbital and are smaller than the atoms from which they are formed.
Anions add electrons to the largest orbital and are larger than the parent atom.

14. O2- > F- > Na+ > Mg2+ > Al3+
For ions of the same charge, ion size increases down a group.
All the members of an isoelectronic series have the same number of electrons.
As nuclear charge increases in an isoelectronic series the ions become smaller:
O2- > F- > Na+ > Mg2+ > Al3+

15. Examples – Choose the larger species in each case:
Na or Na+
Br or Br-
N or N3-
O- or O2-
Mg2+ or Sr2+
Mg2+ or O2-
Fe2+ or Fe3+

16. Examples – Choose the larger species in each case:
Na or Na+
Br or Br-
N or N3-
O- or O2-
Mg2+ or Sr2+
Mg2+ or O2-
Fe2+ or Fe3+

17. Ionization Energy
Ionization energy, is the amount of energy required to remove an electron from a gaseous atom:
Na(g)  Na+(g) + e-.
The larger ionization energy, the more difficult it is to remove the electron.

18. Periodic Trends in Ionization Energies
Ionization energy decreases down a group.
This means that the outermost electron is more readily removed as we go down a group.
As the atom gets bigger, it becomes easier to remove an electron from the most spatially extended orbital.
Ionization energy generally increases across a period.
As we move across a period, the number of protons increases. Therefore, it becomes more difficult to remove an electron.

21. Examples – Put each set in order of increasing first ionization energy:
P, Cl, Al, Na, S, Mg
Ca, Be, Ba, Mg, Sr
Ca, F, As, Rb, O, K, S, Ga

22. Examples – Put each set in order of increasing first ionization energy:
P, Cl, Al, Na, S, Mg
Ca, Be, Ba, Mg, Sr
Ca, F, As, Rb, O, K, S, Ga
1. Na < Al < Mg < S < P < Cl
2. Ba < Sr < Ca < Mg < Be
3. Rb < K < Ga < Ca < As < S < O < F

23. Electron Affinities
Electron affinity (love of electrons) is the energy change when a gaseous atom gains an electron to form a gaseous ion:
Cl(g) + e-  Cl-(g)
Increases across a period.
Decreases down a group.

24. Electronegativity increases
Electronegativity: The ability of an atom in a molecule to attract electrons to itself.
Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Values are calculated from ionization energies and electron affinities.
Electronegativity increases
across a period and
up a group.

25. Electronegativity

26. Examples – put each set in order by increasing electronegativity:
Na, Li, Rb, K, Fr
Cl, Ca, F, P, Mg, S, K

27. Examples – put each set in order by increasing electronegativity:
Na, Li, Rb, K, Fr
Cl, Ca, F, P, Mg, S, K
Fr < Rb < K < Na < Li
K < Ca < Mg < P < S < Cl < F

28. Metals, Nonmetals, and Metalloids

29. Metals
Metallic character refers to the properties of metals (shiny or lustrous, malleable and ductile, oxides form basic ionic solids, and tend to form cations in aqueous solution).
Metallic character increases down a group.
Metallic character decreases across a period.
Metals have low ionization energies.
Metals form positive ions.

31. Nonmetals
Gain electrons to form negative ions.
Do not conduct electricity.
Metalloids
Metalloids have properties that are intermediate between metals and nonmetals.
Example: Si has a metallic luster but it is brittle.
Metalloids have found application in the semiconductor industry.