Transcript Periodic Trends

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.
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
p+ are in the nucleus and more 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.
Ionization Energy
Ionization Energy (IE)
• The amount of energy required to completely
remove an e- from a gaseous atom.
• Recall: removing one e- makes a +1 ion.
• The energy required 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.
• Exceptions in 1st Ionization Energy
Exceptions in 1st Ionization Energy
Why is there a decrease in IE from Mg to Al?
• 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
orbitals 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!!!
Reactivity
Reactivity
• Reactivity refers to how likely or vigorously
an atom is to react with other substances.
• This is usually determined by how easily ecan be removed and how strongly atoms
want to take other atom's e- .
Reactivity - for Metals:
Period - reactivity decreases from left to
right
Group - reactivity increases going down a
group
Why?
- Elements located toward the left of the
periodic table (alkali metals) and near the
bottom easily lose their e-, resulting in
higher reactivity.
- Within the same group, the more e- an
atom has, the easier it will give it up.
(Ex. Li (3) and Fr (87))
Reactivity -for Non-Metals
• Period - reactivity increases from left to
right (not including the noble
gases)
• Group - reactivity decreases going down
the group. (not including the noble
gases)
• Why?
– Atoms are most stable when they have
noble gas electron configuration.
– Groups closest to the noble gases want
to gain an e- to become stable therefore
they have a more vigorous exchange of
e-.
– Elements within the same group vary
significantly in number of e- but contain
the same number of valence e-.
– The lower energy levels are found closer
to the nucleus, having a stronger desire to
complete their energy level and will react
more violently.
Shielding
• Electrons on the outside energy level
(valence e-) have the inner energy levels
blocking the positive force field (nucleus).
• These inside energy level e- shield (block)
the nuclear (pos) force field from the
valence e-
• As you go across the row the nuclear
charge (positive charge) gets larger
because protons are being added to the
nucleus.
• As you go across the row valence e- are
added to the valence shell but the valence
e- have the same shielding.
• The blocking strength (shielding effect) of
these inner e- is the same across the
period.
• 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 move down a group the valence e- are being
added to a new energy level further from the nucleus.
These new valence e have additional levels of inner
shielding e- and are more effectively shielded from the
positive charge.
Ex. Campfire
Electronegativity
Electronegativity
• The tendency for an atom to attract e- to
itself when it is chemically combined with
another element.
• How fair it shares.
• Large electronegativity means it has a
strong pull on an e- toward itself.
Group Trend
• The further down a group the farther the eis away from the nucleus and the more ean atom has.
• Going to the bottom of a group, the e-are
further away from the nucleus.
•
This means they are better shielded
from the nuclear (+) charge and thus not
as attracted to the nucleus.
•
For that reason the electronegativity
decreases as you go down the periodic
table.
Period Trend
• Electronegativity increases from left to
right across a period
• When the nuclear charge increases, so
will the attraction that the atom has for ein its outermost energy level.
This means the electronegativity will
increase
Electron Affinity
• The energy change associated with
adding an e- to a gaseous atom.
• Easiest to add to group 17 or 7A.
• Gets them to full energy level.
• Energy is often required (+) when adding
an e- to metals. Energy is given off (-)
when adding an e- to non-metals.
• EA decreases as we go down a group.
 Electron affinity decreases as we go down
a group because the atoms are getting
bigger and the valence electrons are not
attracted as strongly to the nucleus.
Ionic Size
• Cations form by losing e- (have a positive
charge).
• Cations are smaller than the atom they
come from.
• Metals form cations.
Ionic size
• Anions form by gaining electrons.
• Anions are bigger than the atom they
come from.
• Nonmetals form anions.
Configuration of Ions
• Ions always have noble gas configuration.
• Non-metals form ions by gaining electrons
to achieve noble gas configuration.
• They end up with the configuration of the
noble gas after them.
Periodic Trends
• Across the period nuclear charge increases so
the attractive force gets stronger and the atoms
get smaller when filling to the same energy level.
• Energy level changes between anions and
cations.
Li+1
B+3
Be+2
C+4
N-3
O-2
F-1
• http://www.mhhe.com/physsci/chemistry/a
nimations/chang_7e_esp/pem3s3_2.swf
Size of Isoelectronic ions
• Iso - same
• Isoelectronic ions have the same # of
electrons
• Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
• all have 10 electrons
• all have the configuration 1s22s22p6
Size of Isoelectronic ions
• Positive ions have more protons so they have a
smaller atomic radius.
• The greater the # of protons the stronger the
attraction to the same # of electrons.
• This will cause the atomic radius to be smaller.
-3
N
-2
O
F
Ne
+
Na
+3
Al
Mg+2
• http://www.mhhe.com/physsci/chemistry/a
nimations/chang_7e_esp/pem3s3_4.swf
• Organize the isoelectronic ions/atoms in
order from smallest to largest
• P-3, Ar, Cl-, K+, Ca2+, S2-, Sc3+,
Click here for a further explanation of
periodic trends.
Atomic size increases,
shielding constant
Ionic size increases
Ionization energy, electronegativity
Electron affinity INCREASE