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Periodic Properties of the
Elements
Chemistry 2 Honors
Jeff Venables
Northwestern High School
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?
• 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?
• 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
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.
Examples – Place each group of elements in order of
increasing atomic radius:
1. S, Al, Cl, Mg, Ar, Na
2. K, Li, Cs, Na, H
3. Ca, As, F, Rb, O, K, S, Ga
Examples – Place each group of elements in order of
increasing atomic radius:
1. S, Al, Cl, Mg, Ar, Na
Ar < Cl < S < Al < Mg < Na
2. K, Li, Cs, Na, H
H < Li < Na < K < Cs
3. Ca, F, As, Rb, O, K, S, Ga
F < O < S < As < Ga < Ca < K < Rb
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)
Write electron configurations for the following ions:
1. Al3+
2. S23. Li+
4. Br5. Fe2+
6. Fe3+
Write electron configurations for the following ions:
1. Al3+
1s22s22p6
2. S2[Ne]3s23p6
3. Li+
1s2
4. Br[Ar]4s23d104p6
5. Fe2+
[Ar]3d6
6. Fe3+
[Ar]3d5
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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.
• 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+
Examples – Choose the larger species in each case:
1. Na or Na+
2. Br or Br3. N or N34. O- or O25. Mg2+ or Sr2+
6. Mg2+ or O27. Fe2+ or Fe3+
Examples – Choose the larger species in each case:
1. Na or Na+
2. Br or Br3. N or N34. O- or O25. Mg2+ or Sr2+
6. Mg2+ or O27. Fe2+ or Fe3+
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.
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.
Examples – Put each set in order of increasing first ionization energy:
1. P, Cl, Al, Na, S, Mg
2. Ca, Be, Ba, Mg, Sr
3. Ca, F, As, Rb, O, K, S, Ga
Examples – Put each set in order of increasing first ionization energy:
1. P, Cl, Al, Na, S, Mg
2. Ca, Be, Ba, Mg, Sr
3. 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
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.
• 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.
Electronegativity
Examples – put each set in order by increasing electronegativity:
1. Na, Li, Rb, K, Fr
2. Cl, Ca, F, P, Mg, S, K
Examples – put each set in order by increasing electronegativity:
1. Na, Li, Rb, K, Fr
2. Cl, Ca, F, P, Mg, S, K
1. Fr < Rb < K < Na < Li
2. K < Ca < Mg < P < S < Cl < F
Metals, Nonmetals, and Metalloids
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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.
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.