Periodic Properties of the
Elements
Chapter 8
Periodic Properties
-- a property that is predictable based on the
element’s position within the periodic table
• Atomic Radius (size)
• Ionization Energy
(ionization potential)
• Electron Affinity
• Electronegativity
• Metallic Character
The Periodic Table
• The modern periodic table was developed by
Dmitri Mendeleev
– Based on periodic law
• Which states that when elements are arranged in order
of increasing mass, their properties recur periodically
• Mass increased from left to right… elements
with similar properties fell into the same
columns
• Mendeleev's arrangement allowed him to predict the existence
of elements not yet discovered and even predicted (accurately)
their properties… Why?? How ???
What determines properties?
• Electron Configurations – show the particular
orbitals that are occupied for an atom
• These are (by default) in the ground state
– But can be used to show electrons “jumping” to
an excited state
• The e- configuration tells us hydrogen’s one
electron is in the 1s orbital (the lowest orbital
first – ground state)
– The highest energy level in the entire
configurations = valence electrons
• So how many valence electrons does Hydrogen have??
Orbital Diagrams
• The electron configuration can be represented
slightly differently by using orbital diagrams
• The electron spin is represented by the
direction of the arrow (up or down)
Electron Spin
• Two fundamental aspects of electron spin:
1. Spin, like a (-) charge, is a basic property of all
atoms. One electron does not spin more or less
than another (all have the same amount)
2. Only two possibilities:
+1/2 (up) and -1/2 (down)
4th and final QN  Spin quantum number (ms)
Pauli Exclusion Principle
• The Pauli Exclusion
Principle states no two
electrons in an atom
can have the same four
quantum numbers
– Even if they have the
same n, l, ml … then
they must have
different ms
(+1/2 and -1/2)
Degenerate Orbitals
• Orbitals that are present but do not have any
electrons in them are said to be degenerate
– Like the 3s, 3p, and 3d orbitals in hydrogen
– H1: 1s1
– There are no electrons past the 1s orbital
Energy Associated
• In general, the lower the l quantum number,
the lower the associated energy
E (s orbital) < E (p orbital) < E (d orbital) < E (f orbital)
Shielding
• Electrons feel the repulsive effect from other
electrons AS WELL AS the attractive effect
from the positively charged nucleus
• The electrons between cause a screen effect
or “shielding” that decreases the attractive
force of the (+) nucleus to outer electrons
Effective Nuclear
Charge (Zeff)
• The third electron of
lithium experiences Zeff
• The effective nuclear charge for this third
electron is approximately 1+
• 3+ from the nucleus and 2- from two
shielding electrons, leaving 1+ (the charge)
Penetration
• As the electron “penetrates” the cloud, it feels the
effective of the nucleus more fully because it is less
shielded
• If the electron could somehow get closer to the
nucleus it would feel the full 3+ attraction
• The s orbitals penetrate deeper than the p orbitals
and thus are lower in energy (closer to the nucleus,
smaller r (radius) – s < p < d < f
Things to notice…
• In the 4th and 5th principal levels, the effects of
penetration become so important that the 4s
orbital lies lower in energy than the 3 d
orbitals and the 5s lies lower in energy than
the 4 d orbitals
Things to notice…
• The energy separations between one set of
orbitals and the next become smaller
– The relative energy ordering of these orbitals can
actually vary among elements
– These variations result in abnormalities in econfigurations of the transition metals and their
ions
“Conceptual Connection”
• Lets take a look at page 325
Electron Configurations
• A systematic way of illustrating the increasing
energy of electrons
– They fill the lowest energy first and work their way
up (if ground state, of course)
– Two electrons per orbital  opposite spins
• ms QN of +1/2 and -1/2
Principles / Rules
• This patter of orbital filling is known as the
Aufbau Principle (German for “building up”)
• Hund’s Rule states when filling degenerate
orbitals, electrons fill them singly first, with
parallel spins, then double up with opposite
spins
Hund’s Rule = Pizza Party
e- configs… a summary
• Electrons occupy orbitals so as to minimize the
energy of the atom.. Lower energy orbitals fill
before moving on to the higher energy orbitals
• Orbitals hold no more than two electrons with
opposite spins (Pauli exclusion principle) – no 2
electrons can have the same 4 QN’s)
• Electrons fill orbitals singly with parallel spins
before doubling up (Hund’s Rule)
Review
• The s sublevel has only one orbital and can
only hold 2 electrons
• The p sublevel has three orbitals and can hold
6 electrons
• The d sublevel has five orbitals and can hold
10 electrons
• The f sublevel has seven orbitals and can hold
14 electrons
Practice
• Write out the electron configurations for each of
the following elements:
(a) Mg
(b) P
(c) Br
(d) Al
• What element has the following configuration:
1s22s22p63s23p54s23d4
More practice
• Write out the orbital diagram for sulfur and
determine the number of unpaired electrons
More practice
• Write out the electron configuration for Ge
and identify the valence electrons and the
core electrons
• Try the same for phosphorus
Using the “f” orbitals
Valence Electrons
• The chemical properties are largely due to the
number of valence electrons
• This explains the reactivity of hydrogen and
the inertness of helium
The Octet Rule
• Noble gases are least reactive
• Elements with configurations similar to noble
gases are most reactive
• The formation of cation and anions to obtain
their octet…
Atomic Size
• Most of the atom’s size is from the electrons
• The orbitals are just “high probabilities”
• How can we define atomic size?
Radius
• Two Ways:
– Van der Waals radius – one-half the distance
between adjacent nuclei in an atomic solid
(meaning they are not bonded but just ‘frozen’ up
against one another)
– Bonding atomic radius or covalent radius
• one-half the distance between two bonded atoms
• As is diatomics or other BONDED molecules
Atomic Radius
**TREND #1**
• A more general term, the atomic radius, refers
to the average bonding radii determined from
many bonded and non-bonded measurements
Atomic Radius
• As you move down a column (or family) in the
periodic table, atomic radius increases
– Why??
• As you move to the right across a period (or
row) in the periodic table the atomic radius
does what???
– It decreases!! Think about Zeff and shielding…
A General Trend:
Atomic Radius
Think about it…
• Why does it take 1312 kj/mol of energy to
remove the 1s electron from a hydrogen atom
but 5251 kJ/mol to remove it from He+
H
He+
1s1
1s1
• The core electrons efficiently shield electrons
in the outermost principal energy level from
nuclear charge, but outermost electrons doe
not efficiently shield one another from
nuclear charge
Transition Metals
• These do not change all that much
• As another proton is added, as is another
electron
– This goes into the highest energy level (outside)
• Because the 4s orbital fills before the 3d, this
ratio stays roughly constant, which makes the
radius stay roughly constant
Practice
• Choose the larger atom from each and explain
your choice… trends are not reasons…
(a) N or F
(b) C or Ge
(c) N or Al
**Page 338 for additional reasoning
(d) Al or Ge
Ions
• When making cations – subtract electrons
(Li) 1s22s1
(Li+) 1s22s0
• When making anions – add electrons
(F) 1s22s22P5
(F+) 1s22s22P6
Transition Metal Ions
• Transitions are a little more complicated…
– Remove electrons in the highest n-value orbitals
first, even if this does not correspond to the
reverse order of filling
(V)
[Ar] 4s23d3
(V2+)
[Ar] 4s03d3
Explanations
• Paramagnetic – contains unpaired electrons
and therefore is attracted by external
magnetic fields
Ag [Kr] 5s14d10
• Diamagnetic – contains only paired electrons
and therefore slightly repelled by external
magnetic fields (preferred)
Zn2+ [Ar] 4s03d10
Practice
• Write the e- configs and orbital diagrams for
each of the following ions. State whether they
are diamagnetic or paramagnetic
(a) Al3+
(b) S2(c) Fe3+
Diamagnetic
Diamagnetic
Paramagnetic
Hint: write out the atom’s config first and then
do what needs to be done to make the ion
Ionic Radii
**Trend #2**
• The radius of an ion is the ionic radius (same
concept as atomic radius)
• What happens to the radius as of an atom when
it becomes a cation? An anion?
Ionic Radii
• What happens to the radius as of an atom
when it becomes a cation? An anion?
Ionic Radii
• In general, cations are much smaller than their
corresponding atoms
• Anions are much larger than their
corresponding atoms
Practice
• Choose the larger atom or ion from each
(a) S or S2S2- is larger
(b) Ca or Ca2+
Ca is larger
(c) Br- of Kr
????????
The Br- is larger than the Kr atom because Br- has
one fewer proton than Kr which results in a lesser
pull on the electrons (Zeff) and thus a larger radius.
Isotopes
• Think about isotopes…
Would you expect C-12 and C-13 to have
different atomic radii??
The isotopes of atoms have the same atomic
radii because neutrons are relatively small and
are neutral in charge.
Ionization Energy (IE)
** Trend #3**
• The ionization energy (IE) of an atom or ion is
the energy required to remove an electron
from the atom or ion in the gaseous state
Ionization Energy (IE)
• The energy required to remove the first
electron is called the first ionization energy
• The second  the second ionization energy
Ionization Energy (IE)
Na(g)  Na+(g) + 1e-
IE1 = 496 kJ/mol
Na+(g)  Na2+(g) + 1e-
IE2 = 4,560 kJ/mol
Much harder because it does not want to lose
that second electron
Summary of IE
• IE usually decreases as you move down a
column because they are farther away from
the nucleus and are held less tightly
• IE generally increase as you move to the right
because electrons in the outermost principal
energy level experience a greater Zeff
Practice
• Choose the element with the higher first
ionization energy
(a) Al or S
S
(b) As or Sb (c) N or Si
As
N
(d) O or Cl
Cannot tell
Page 345-346 for explanations… remember trends are
not reasons 
Exceptions
• Elements would rather have one electron in
each orbital than one doubled up (and not all
doubled up)
Electron Affinity
**Trend #4**
•
• Electron Affinity (EA) of an atom or ion is the
energy change associated with the gaining of
an electron by an atom in the gaseous state
How much they ‘want’ an electron…
In a Nutshell…
• Most groups do not show a ‘trend’
• But, group IA electron affinity becomes more
positive as you move down the column (adding an
electron becomes less exothermic)
• EA generally becomes more negative (adding an ebecomes more exothermic) as you move to the right
across a period (or row) of the table
Metallic Character
**Trend #5
• Just as it sounds
– How ‘good’ it is at being a metal
•
•
•
•
•
Good conductors of heat and electricity
Malleability
Ductile
Shiny
Tend to lose electrons in a rxn
Metallic Character
• As you move to the right across the table, the
metallic character decreases
• As you move down a column (or family) in the
periodic table, metallic character increases
Practice
• Choose the more metallic element from each
(a) Sn or Te (b) P or Sb (c) Ge or In (d) S or Br
Sn
Sb
In
Page 350 – 351 for full explanations
Cannot tell
Conceptual Connection 8.5
• Try this problem on page 351 without looking
at the answer
• Validate it using your knowledge on the trends
(great test quality question)
Alkali Metals
• Their single valence electron (which prevents
them from the noble gas configuration) is
easily removed
Configurations: ns1 in the outer shell
(valence shell)
Alkali Metals
Halogens
• These need one more electron to obtain their
noble gas configuration, making them the
most reactive!!
Configurations: ns2np5 in the outer shell
(valence shell)
Halogens
Noble Gases
• The noble gases are the ‘happiest’ on the
table… they have already obtained the octet
and are stable.
– Least Reactive
– INERT!!
Configurations: ns2np6 in the outer shell
(valence shell)
Noble Gases
Noble Gases
• The high ionization energies and full outer quantum
levels make them very unreactive
• Before the 1960s, no noble gas compounds were
known
• Since then, two of the noble gases have been shown
to react with fluorine (the most reactive nonmetal on
the table), under extreme conditions
Noble Gases
Kr + F2  KrF2
Xe + F2  XeF2
Xe + 2 F2  XeF4
Xe + 3 F2  XeF6