Periodic Trends
Atomic Size
Atomic Size
• The electron cloud doesn’t have a definite
edge.
• Scientists get around this by measuring
more than 1 atom at a time.
• Summary: it is the volume that an atom
takes up
Atomic Size
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Radius
• Atomic Radius = half the distance between two
nuclei of a diatomic molecule.
• A diatomic molecule is a molecule where an
element is bonded to itself. (Hydrogen, Oxygen,
Chlorine Bromine Fluorine Iodine Nitrogen)
Group trends – Atomic
Radius
• As we go down a
group the atoms
have more e-,
therefore more
energy levels and
the atoms get
bigger
H
Li
Na
K
Rb
Periodic Trends – Atomic Radius
The atomic radius decreases as you go
from left to right across a period.
Na
Mg
Al
Si
P
S Cl Ar
Explaining this trend
When moving across a period additional
protons (+) are in the nucleus and more
electrons (e-) are in the same energy level.
• The opposite charges in the nucleus and
the e- cloud cause the atom to be 'sucked'
together a little tighter.
Therefore the radius decreases, meaning
the atom becomes smaller.
Shielding
• Electrons on the outside energy level
(valence shell) have the inner energy
levels blocking the positive nucleus.
• As you go across the period:
– the nuclear charge (positive charge) gets
larger due to additional protons.
– valence e- are added to the outer shell.
– the blocking strength (shielding effect) of the
inner electrons remains the same.
• Further right in a period the valence e- will
have a greater attraction to the nucleus
because of the greater positive charge.
Shielding becomes less effective across the row;
2e- can shield +3 better than 2e- can shield +10.
As you go down a group:
- the valence e- are being added to a new energy
level further from the nucleus.
- there are additional levels of inner shielding
electrons, which more effectively shield valence efrom the positive nucleus.
Ex. Campfire
Electronegativity
Electronegativity
• The tendency for an atom to attract e- to
itself when it is chemically combined
(bonded) with another element.
• Large electronegativity means it has a
strong pull on an e- toward itself.
Group Trend
• Going down a group, the bigger the atom
so the further the valence electrons are
from the nucleus.
• This means they are better shielded from
the nuclear (+) charge and thus not as
attracted to the nucleus.
• Therefore, the electronegativity
decreases as you go down a group. (less
ability to attract an electron in a bond)
Period Trend
• Electronegativity increases from left to
right across a period
• When the nuclear charge increases (due
to increased protons), so will the
attraction that the atom has for its valence
electrons.
This means the electronegativity will
increase.
Electronegativity
Ionization Energy
Ionization Energy (IE)
• The amount of energy required to completely
remove an electron (e-) from a gaseous atom.
• Recall: removing one e- makes a +1 ion.
• The energy required to remove the first electron
is called the first ionization energy
X(g) + energy → X+ + e-
• Second and third ionization energies represents
losing a 2nd and then a 3rd e- from the same
atom.
• It can be shown as:
X+ (g) + energy X2+ (g) + eX2+ (g) + energy X3+ (g) + e• More energy required to remove 2nd e-, and still
more energy required to remove 3rd e-.
• The closer the e- is to the nucleus, the more
difficult it will be to remove.
Group Trends (I.E.)
• Ionization energy decreases down the group.
• Ex. Going from Be to Mg, IE decreases because:
– Mg outer e- is in the 3s sub-shell rather than the 2s.
– This is higher in energy and further from the nucleus.
– So the 3s e- is more easily removed, requiring less
energy.
• A similar decrease occurs in every group in the
periodic table.
Period Trends (IE)
IE generally increases from left to right.
Why?
The e- are attracted more strongly to the nucleus
(smaller radius).
It takes more energy to remove one e- from the
atom with stronger attraction, therefore IE
increases.
Ex. From Na to Ar (11 p+ to 18 p+), the attraction
of the protons to e- within the same energy level
increases.
Some Exceptions:
There is a decrease in IE from Mg
to Al….Why?
• Al is 1s2 2s2 2p6 3s2 3p1
It has one e- that is in a p sublevel.
• Mg is 1s2 2s2 2p6 3s2.
Mg - the ‘s’ sublevel is full – this gives it
a slight stability advantage and will require
more energy to let go of its e-.
Why is there a fall in IE from
phosphorus to sulfur?
• This can be explained in terms of e- pairing.
• Phosphorus - 1s2 2s2 2p6 3s2 3p3
• Sulfur - 1s2 2s2 2p6 3s2 3p4
•
As the p sublevel fills up, e- fill up the vacant
sub levels and are unpaired.
• Phosphorus’ configuration is more energetically
stable than sulfur’s because there are e- that are
unpaired.
• When e- are paired, there is some repulsion
which lessens their attraction to the nucleus.
• It becomes easier to remove!
• Having a half filled sublevel is more stable
than a partially filled sublevel.
• So… sulfur will break the expected trend and
want to lose an e- requiring less IE.
Why an exchange in e- ?
• Noble Gases have full energy levels.
• Atoms behave in ways to achieve noble
gas configuration.
2nd Ionization Energy
• The amount of energy required to remove
the 2nd e- from a gaseous atom.
• For elements that reach a filled or half
filled sublevel by removing 2 e- the 2nd IE
is lower than expected.
• Makes it easier to achieve a full outer shell
• True for s2 , the alkaline earth metals
which form +2 ions.
3rd IE
• Using the same logic s2p1 atoms have a
low 3rd IE.
• Atoms in the aluminum family form +3
ions.
• 2nd IE and 3rd IE are always higher than 1st
IE!!!
Ionization Energy
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Reactivity
Reactivity
• Reactivity refers to how likely or vigorously
an atom bonds (reacts) with other
substances.
• This is usually determined by how easily ecan be removed or gained.
Reactivity - for Metals:
Group - reactivity increases going down a
group. This is because the valence
e- is further from the (+) nucleus,
therefore is more easily removed.
(reacts more easily)
Period - reactivity decreases from left to
right. This is because the atom
becomes smaller, so the positive
nuclear charge holds more tightly
onto the valence e-. (less reactive)
Reactivity of metals
Reactivity -for Non-Metals
(not including the noble gases)
• Group - reactivity decreases going down the
group. There is the same number of
valence e- but the atom becomes bigger
due to having more energy levels. The
increased shielding means the positive
nucleus is less able to attract valence eof other atoms.
Reactivity -for Non-Metals
(not including the noble gases)
• Period - reactivity increases from left to
right. The atoms become smaller
(increased positive nucleus), and
are therefore able to attract
valence electrons more easily.
The further right in the period, the
less e- required for a full octet,
therefore group 17 are the most
reactive (only need to gain 1 e-).
Reactivity of Nonmetals
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Electron Affinity
Electron Affinity
• The ability of an atom to gain an electron.
• The stronger the attraction of the nucleus
and the valence electron gained, the more
energy is released.(the higher the electron
affinity)
Example
• Cl(g) + e- → Cl-(g) E.A. = -349 kJ/mole
Group Trends
• Going down a group, the bigger the atom
so the further the valence electrons are
from the nucleus.
• The greater the distance, the less the
attraction (due to the shielding of valence
electrons from the positive nucleus) and
so the less energy is released.
• Electron Affinity decreases as you go
down a group.
Period Trends
• From left to right within a period, the atoms
become smaller due to increased nuclear
charge (more protons) with no change in
shielding.
• Therefore, it is easier to attract valence
electrons and the electron affinity
increases.