The Periodic Law
Modern Chemistry © 2009
Holt, Rinehart, & Winston
Chapter 5, pp 133 - 165
Periodic Law
• Chemical & physical properties regularly
repeat when elements are listed according to
their atomic numbers
 Atomic number is equal to # of protons
 # of protons = # of electrons
 # of valence electrons = Group # for s block
= Group # - 10 for p block
o “Main-group elements” are those in the s & p blocks
o Properties are determined mostly by valence electrons
Contributors
• Stanislao Cannizzaro
– Reliable method of measuring atomic mass
• Dmitri Mendeleev
– Group elements of similar properties together
– Arrange groups according to their atomic masses
• Henry Moseley
– Rearrange elements according to atomic numbers
– Maintain groupings according to properties
Periodic Properties’ Trends
• Atomic & ionic radii: as go left & down
– Across the period, nuclear charge increases.
• As the positive charge increases, electrons are pulled in
more tightly, thereby the radius.
– Down the group, the # of energy levels increase.
• Cations have smaller ionic radii than the atom.
– When the valence electron shell is lost, the ion is
smaller
• Anions have larger ionic radii than the atom.
– Extra electrons do not compact as readily, mostly
because of their electrostatic repulsion for each other.
Periodic Properties’ Trends
• Ionization energy = energy required to remove
electron from neutral atom

as you go right across the period
o As nuclear charge increases, electrons are held more
tightly.

as you go down the group
o As energy levels are added, electrons are held more
loosely.
o Additional layers of electrons shielding of protons and
further repel valence electrons.
Periodic Properties’ Trends
• Electron affinity = energy change
accompanying acquisition of electrons by
neutral atoms.
 Most atoms release energy when get electron.
o Released energy is noted as a negative value.
 A + e-  A- + energy
F has the highest value of -339.9.
o Positive or less negative values indicate atom was
“forced” to get electron, i.e., energy was absorbed.
o Most atoms so forced will spontaneously lose electron.
 A + e- + energy  A-
Values are often listed as “(0)”.
Periodic Properties’ Trends
• Electronegativity = ability to attract electron in
a compound
• F has highest value.
 Fluorine becomes most like a noble gas (filled
valence shell = greater stability) when is F-.
• Fr/Cs has lowest.
 Alkali metals are energetically most stable when
they lose their valence electron, not gain one.
Additions to Periodic Table
• Noble gases
– 1894: Ar; 1895: He; 1898: Kr & Xe; 1900 Rn
• Lanthanides
– Early 1900s
– Very similar in chemical & physical properties
– Soft and shiny; reactivity like alkaline-earth metals
• Actinides
– All are radioactive; 1st 4 are natural, rest synthetic
– Those after uranium are “transuranium elements”.
The s-Block Elements
• Group 1 = alkali metals (but H is a nonmetal)
– Electron configuration for all end in ns1
– The most reactive of all elements (with air or H2O)
• Never found free in nature; stored in kerosene
– Are silvery and soft
• Group 2 = alkaline earth metals
– Electron configuration for all end in ns2
– Too reactive to be found free in nature
• Used in fireworks (Mg – white; Sr – red)
– Harder, denser, stronger than alkali metals
d-Block, or Transition, Elements
• Less reactive than elements of the s-block
• May have different # valence electrons/ group
– Sum of outer s & inner d electrons = group #
• Have typical metal characteristics
– Good conductors of electricity
– High luster
• Some are don’t react, staying free in nature
– Pt, Pd, & Au are least reactive metals
The p-Block Elements
• All have filled s-orbitals
• Includes metals, all metalloids, & nonmetals
– Metals: Al, Ga, In, Tl, Sn, Pb, Bi, Po, & Uuq & Uuh
– Metalloids: B, Si, Ge, As, Sb, Te, & (At)
• Properties of both metals and nonmetals
• Brittle solids; have luster
• Semi-conductors
– Nonmetals: noble gases, halogens, O, S, N, P, C, Se
• Halogens are the most reactive of all nonmetals
• Halogens have all 3 states of matter in their group
Trends for p- & d-blocks
• Atomic radii: going right across period
• Ionization energy (IE): going right across
period & going down the group for the 1st IE
– Outer s electrons are less shielded by d electrons
• Ion formation & radii: all lose ns2 electrons 1st
– Ions of a 2+ charge
in size across period
• Electronegativity: {F = 4.0; others are relative}
– d block: all are 1.1 – 2.54
– f block: all are 1.1 – 1.5
– Inversely proportional to atomic radii for both.