Periodic Properties
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Ionic Electron Configurations
• You learned how to write the electron configurations
for ions.
• Just remember that the valence electrons get added
to or removed first!
• So you always add to or remove from the highest
principal energy level, or n value.
• Write the electron configuration for Cl-, Mn2+
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Ionic Electron Configurations
• Can you see why metals tend to lose electrons?
• Can you see why nonmetals tend to gain electrons?
• What kind of e- configuration are they trying to
achieve?
• How many valence s and p electrons is this?
• This is the basis for the Octet Rule: Elements in the
Main Group tend to gain, lose, or share electrons in
chemical rxns in order to have 8 valence electrons,
or filled s and p subshells like the Noble Gases.
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Ionic Radii
• You learned the Periodic Trend for atomic radii, that
is the radii for neutral atoms.
• What happens to the size of an atom when it loses
electrons to become a cation?
• Why does it become smaller? (think about Zeff and
the shell that it loses e- from)
• Likewise, what happens to an O atom when it gains
2 e- to become the oxide anion?
• Why does it get larger?
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Ionization Energy
• You know that an electron in an orbital may be
promoted to a higher energy level orbital if it
absorbs enough energy.
• But you also know that electrons may be removed
from an atom entirely.
• Although it is typically the valence electrons which
are removed, we can actually remove core electrons
as well.
• So if we put in enough energy, we can remove an
electron (or more than one) entirely.
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Ionization Energy
• The Ionization Energy is the energy required to
remove an electron from an atom (or ion) in the
gaseous state.
Li(g) → Li+(g) + e• The first Ionization Energy, IE1 or Ei1, is the energy
required to remove the outermost valence electron
from a neutral atom.
• The second IE, IE2, is the energy required to
remove the second valence electron from the +1
cation.
• What’s the IE3?
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Ionization Energy
• How does IE vary down a group or across a Period?
• Going down a group, you add a shell, so the valence electrons
are further and further from the nucleus. As the atom gets
larger, the valence electron is further away, and is loosely
held. So it is easier to remove, and the IE is smaller. So IE
decreases going down a group.
• Going across a Period, the atomic size is decreasing due to an
increasing Zeff. Both of these means that the valence electrons
are more tightly held, so the IE increases going across a
period from left to right.
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Ionization Energy
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Ionization Energy
• Look at the following graphs, do you see any exceptions, or
something you think is unusual?
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Ionization Energy
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Ionization Energy
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Ionization Energy
• There are exceptions to the general trends, but many of them
are explained easily.
• For example, the Alkaline Earth metals have higher IE1
values than expected, higher than Group 13 values.
• IE1 for Group 2 is removing an s electron, while IE1 for
Group 13 is removing a p electron in the same energy level.
• s electrons are slightly closer on average to the nucleus, are
more tightly held to the atom, and so are harder to remove
than a p electron on the same level.
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Ionization Energy
• Also, Group 16 tend to have lower than expected IE values,
and are often lower than the values for Group 15.
• This is due to electron-electron repulsion effects.
• Group 16 has a valence e- configuration of ns2np4, while
Group 15 has a valence e- configuration of ns2np3.
• Group 16 has the last valence electron doubled up in a p
orbital, so there is higher electron-electron repulsions in this
filled p orbital, so it is easier to remove an electron.
• Group 15, on the other hand, has no filled p orbitals, so the
electron-electron repulsions are lower, so it is harder to
remove an electron from these p orbitals.
• Group 15 has the p orbitals half-filled! More stable!
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Ionization Energy
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Ionization Energy
• Note that the IE values for the transition metals
don’t increase very much (d-block) going across a
Period.
• Remember that the ns2 electrons are the valence
electrons for the transition metals, so they are
removed first, NOT the d electrons.
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Ionization Energy
• What about higher IE values? How do you think
they compare to the IE1 values?
• For example, for Mg: how many valence electrons
does Mg have?
• Is there a trend for IE1, IE2, IE3 for Mg?
• Do you think that there will be a large increase in IE
for one of these?
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Ionization Energy
• Mg has 2 valence electrons, 3s2.
• When we remove the first 3s electron, we put in 738
kJ/mol and produce the Mg+ ion.
• Do you think that it will be harder or easier to remove a
second electron from a Mg+ electron?
• When we remove the second 3s electron, we must put in
1451 kJ/mol (or roughly double) and we produce the Mg2+
cation. So the more positive the charge on an ion, the
harder it is to remove an electron. This should make
sense as cations are smaller than the neutral atom.
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Ionization Energy
• What about the IE3 for Mg? Now what valence
electron are we trying to remove? We have removed
the 3s electrons, so we are now trying to remove a 2p
electron! This is a core electron.
• As we are now removing a core electron from a
lower energy level (closer to nucleus) as well as
removing an electron from a cation, the IE jumps
sharply (roughly 8 to 10 times greater). It now
requires 7733 kJ/mol to remove the third electron
from Mg.
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Ionization Energy
• We can relate this to the valence electrons of Group
2. They have 2 valence electrons, so they may be
removed fairly easily, but once the valence electrons
are removed, there is a huge increase in the next IE
as we are now removing core electrons.
• So you can always tell what group an atom is in
given the successive IE values, as they are giving you
the number of valence electrons!
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Ionization Energy
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Ionization Energy
• Example: An atom in Period 2 showed the following
IE values: IE1 = 750 kJ/mol; IE2 = 1500 kJ/mol; IE3
= 2800 kJ/mol; IE4 = 5000 kJ/mol; IE5 = 38,000
kJ/mol; IE6 = 47,000 kJ/mol. What group is this
atom a member of? What element is it?
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Ionization Energy
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Electron Affinity
• Electron Affinity, EA or Eea, is defined as the energy change
when an electron is added to an atom (or ion) in the gaseous
state:
F + e- → F• How has the electron configuration of F changed once an
electron is added? Can you see why nonmetals tend to gain
electrons?
• The electron affinity value gives us an idea of the attraction
that atom has for an electron.
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Electron Affinity
• For most atoms, energy is released when an electron is added
so the EA is negative.
• So the more negative the EA, the more likely the atom will
gain an electron, and the resulting anion is more stable.
• The Halogens have very negative EA values which should
make perfect sense as they only need 1 electron to achieve the
same electron configuration as a Noble Gas.
• However, for the Noble Gases and a few other atoms
(particularly Group 2 with their filled s subshell and to a
lesser extent Group 15 with their half-filled p subshell), the
EA is actually positive. This means that it is very difficult to
add electrons to these atoms.
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Electron Affinities
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Electron Affinity
• You can add more than 1 electron to an atom, but the second
electron is added to the anion:
O + e- → O- EA1 = -141 kJ/mol
O- + e- → O2- EA2 = 878 kJ/mol
• As you might expect (and as is similar to the successive IE
values), it is harder to add an electron to an anion as it has a
lower Zeff, it is larger, and electron-electron repulsions are
high.
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Electron Affinity
• However, oxygen and other nonmetals may gain more than 1
electron, but this is one of the reasons why it is very hard for
nonmetals to gain 3 or more electrons.
• For example, although N does form ionic compounds as
nitrides where the N has gained 3 electrons (N3-), it forms
many more covalent compounds.
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Electron Affinity
• How does the EA vary going down a Group and across a
Period from left to right?
• First of all, the EA is the Periodic Property which has the
most variation from the general trend.
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Electron Affinity
• That said, IN GENERAL, the EA becomes less negative as
you go down a group.
• However, the change is small and is not always followed.
• This is because the atomic radius is larger down the group
which would tend to make the EA less negative, but the
electron-electron repulsions are smaller in the larger atom,
which tends to make EA more negative!
• The EA generally becomes more negative across a Period
from left to right (or going towards the Halogens).
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Electron Affinities
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Metallicity
• The Periodic Table is also roughly arranged by metallic
character.
• Remember the general properties of metals (review them)?
• The more an element has these properties, the more metallic
it is.
• Metals tend to have lower IE values and less negative EA
values as they tend to lose electrons to become cations.
• Where is this on the Periodic Table?
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Metallicity
• Based on the above, you should be able to predict that as you
go down a Group, the metallicity increases.
• Why? The atomic radius increases, so the IE decreases,
and the EA becomes less negative (roughly). So it is
easier to remove an electron, so it is more metallic!
• What about going across a Period from left to right?
• Now the atomic radius is decreasing, the Zeff is increasing,
so the IE increases, and the EA becomes more negative
(roughly). So the metallicity decreases across a Period
from left to right.
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Metallicity
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Metallicity
• So what’s the most metallic element on the Periodic
Table? (Not that anyone works with it, it is too
radioactive and toxic!)
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