Chemistry Chapter 5&6
The Periodic Law Notes 5
Mendeleev’s Periodic Table
Dmitri Mendeleev
Modern Russian Table
Chinese Periodic Table
Stowe Periodic Table
A Spiral Periodic Table
Triangular Periodic Table
“Mayan”
Periodic
Table
Orbital filling table
Periodic Table with Group Names
The Properties of a Group:
the Alkali Metals
Easily lose valence electron
(Reducing agents)
React violently with water
Large hydration energy
React with halogens to form salts
Sublevel Blocks of the Periodic Table
Figure 5-5 p. 129
s-block
• Group 1
• Alkali Metals
–
–
–
–
ns1 is highest level
Silvery appearance
Soft—cut with knife
Highly reactive—never
found free in nature
– Low melting points
<100oC
• Group 2
• Alkaline-earth metals
– ns2 is highest level
– Harder & denser, w/
higher melting points
than Group 1
– Highly reactive—never
found free in nature
Special exceptions to s-block
• Hydrogen
– Has ns1
– Totally different
properties from alkali
metals
• Helium
– Has ns2
– Highest level is
completely full
– Stable like noble gases
d-block
• d sublevel for preceding energy level is filling
• d sublevel filling has some deviations—Group 11:
Cu, Ag, Au
– Outer s & d sublevels still have same # e-
• Transition elements: d-block metals w/ typical
metallic properties
–
–
–
–
Less reactive than Group 1 & 2
Exist free in nature
Good conductors of electricity
High luster (shiny)
•
•
•
•
p-block
All elements of Groups 13-18 except Helium
Properties vary greatly
Nonmetals (right hand end)
All six metalloids
– Brittle solids
– Some properties of metals and nonmetals
• Eight metals (left hand side and bottom of the block)
– Harder and denser then s-block alkaline-earth metals
– Softer and less dense than d-block metals
– Stable in the presence of air
• Group 17 Halogens
– Most reative of the nonmetal
– 7 electrons in outer shell
f-block
• Lanthanides & Actinides
• 14 elements—seven 4f orbitals are filling
• Lanthanides
– Similar reactivity to Group 2
– Shiny metals
• Actinides
– Only 1st four found in nature
– All are radioactive
Determination of Atomic Radius:
Half of the distance between nucli in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels
Table of
Atomic
Radii
Ionization Energy - the energy required to remove an
electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove
Tends to decrease down a group
Outer electrons are farther from the
nucleus
Ionization of Magnesium
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Table of 1st Ionization Energies
Another Way to Look at Ionization
Energy
Electron Affinity - the energy change associated
with the addition of an electron
Affinity tends to increase across a period
Affinity tends to decrease as you go down
in a Group or family
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
Table of Electron Affinities
Ionic Radii
Cations
Anions
Positively charged ions
Smaller than the corresponding
atom
Negatively charged ions
Larger than the corresponding
atom
Summation of Periodic Trends
Table of Ion Sizes
Electronegativity
A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase across
a period
Electronegativities tend to decrease down a
group or remain the same
Periodic Table of Electronegativities