Atomic and Ionic Size Trends. Mr. Shields Regents Chemistry U08 L04. Size Trends. Atomic Radii follows two trends: Radii increases going down a group Radii decreases going across a period. But how do we measure Atomic Radii?. Atomic Radii.
Mr. Shields Regents Chemistry U08 L04
But how do we measure Atomic Radii?
Atomic Radius is measured as ½ the distance between
Adjacent nuclei in a molecule.
Why not just measure ½ the diameter of the atom?
Hint: What’s the definition of an atomic orbital?
Atomic Radii Trends
How can we account for this trend?
The trend down a group may be easier to explain first
As we move down a group what happens to the principal
Energy level ?
As the principal energy level increases, electrons move
further and further from the nucleus.
Nucleus n=1 2 3 4
As we move down group 1 each successive S orbital electron
Is further from the nucleus and thus Atomic radii increases
Since the s orbital is further from the nucleus the radii of the
Atom increases. But …
The nuclear charge is also increasing since Atomic Number is increasing.
Increasing nuclear charge
Diminishes the rate of
Change of increasing
Atomic Radii Down a
Electrons are pulled
Toward the nucleus more
We’ve seen how atomic radii increases going down a group
But what happens when we go across a period?
We’ll, in fact atomic radiidecreases. But why?
We can begin to understand what is happening if we look at
Both the Atomic number and what principle energy level
electrons are being added to.
As we move across a period Atomic numbers increase
- Pos. Nuclear charge also increases so would expect
the electrons to be pulled closer to the nucleus.
So this could explain decreasing Atomic radius
- BUT … this same thing happens as we move down
a group. And for groups Atomic Radii increases as we
add more electrons? So why does radius increase in
groups but not across a period?
The difference is that when we go across a period electrons do
Not fill higher energy levels. They either occupy lower energy
Levels or the same energy level
Whereas when going down a group electrons occupy
successively Higher principle energy levels. For example …
1 2 3 13
n=3 2-8-1 2-8-2 2-8-3
n=4 2-8-8-1 2-8-8-2 2-8-9-22-8-18-3
n=5 2-8-18-8-1 - - -
But why doesn’t atomic radii remain about the same across
e-Atomic radii across periods
As atomic number increases across a row additional electrons
are added to the same (or lower) energy level. The effective
Nuclear charge may also be increasing and electrons are pulledin
More strongly towards the larger more positive nucleus.
Unlike when moving down a group, there are no new principal
energy levels being added to counteract the effect of increasing
nuclear charge and increasing effective nuclear charge.
Na: Effective nuclear charge = +1 S: Eff. Nuclear Charge = +6
Effective Nuc. Chg. (Grp I) = 1
Nuclear charge (Br) = 35
Effective Nuc. Chg. (Grp VII)= +7
We’ve now seen how atomic radii changes in Periods & Groups
But what happens when Atoms either gain or lose electrons to form ions?
How does Ionic Radii vary down Groups & across Rows?
The representative elements of groups 1, 2, 13 and 14 give up electrons to form +1, +2, +3 and +4 ions respectively
When atoms lose all their valence electrons they lose the
outermost quantum level (n).
Consider Aluminums electron configuration. What is it?
2-8-3 (principle energy levels 1, 2 and 3 are occupied)
What is the electron config after Al loses its 3 valence electrons?
2-8 (only principle energy levels 1 and 2 are occupied)
What is the charge on Aluminum?
Notice that even though the ionic electron config is the same
ionic radius gets progressively smaller moving across the period.
This happens because the positive eff. nuclear charge seen by the same number of electrons increases as we move across a the row
Variation in atomic and Ionic Radii
Let’s next look at the non-metals, for example Chlorine
Non-metals form ions by gaining electrons
Cl 2-8-7 2-8-8 Cl- (negative ion)
When we add electrons the effective nuclear charge per
electron decreases AND there is increases electron repulsion
So … you would expectthe ionic radius to increase and it does
Cl atomic radius = 99 nm
Cl- ionic radius = 181 nm
Moving down groups the principal energy level increases
- This is true for all atoms, Anions (-) & Cations (+)
Li 2-1 Li+ 2 F 2-7 F- 2-8
Na 2-8-1 Na+ 2-8 Cl 2-8-7 Cl- 2-8-8
K 2-8-8-1 K+ 2-8-8 Br 2-8-18-7 Br- 2-8-18-8
So atom & ionic size increases going down Groups
Going across periods Ionic size first decreases then jumps up
When Oxidation states change from positive to negative.
- after the jump up the downward trend in size continues
The ionic compound MgO
If we were to look at individual atoms Mg would
Actually be larger than Oxygen!