Transcript atomic radii

Mr. Shields
Regents Chemistry
U08 L04
1
Size Trends
Atomic Radii follows two trends:
1) Radii increases going down a group
1) Radii decreases going across a period
But how do we measure Atomic Radii?
2
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?
3
Atomic Radii Trends
Atomic Radii
increasing
Decreasing radii
How can we account for this trend?
4
Atomic Radii
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
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Atomic Radii in Groups
As we move down group 1 each successive S orbital electron
Is further from the nucleus and thus Atomic radii increases
n
Element
Radii (pm)
change
2
Li
155
-
3
Na
190
35
4
K
235
45
5
Rb
248
13
6
Cs
267
19
7
Fr
270
3
6
Atomic Radii in groups
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
Group.
Electrons are pulled
Toward the nucleus more
strongly
Nuclear
charge
Radii (pm)
change
Li (3)
155
-
Na (11)
190
35
K (19)
235
45
Rb (37)
248
13
Cs (55)
267
19
Fr (87)
270
3
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Atomic Radii across periods
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 radii decreases. 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.
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Atomic Radii across periods
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?
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Atomic radii across periods
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 …
n=3
n=4
n=5
1
2-8-1
2-8-8-1
2-8-18-8-1
group
2
2-8-2
2-8-8-2
-
3
2-8-9-2
-
13
2-8-3
2-8-18-3
-
But why doesn’t atomic radii remain about the same across
a group?
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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 pulled in
More strongly towards the larger more positive nucleus.
+11 2-8-1
e-
Na: Effective nuclear charge = +1
+16 2-8-6
e-
S: Eff. Nuclear Charge = +6
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.
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Nuclear charge (K) = 19
Effective Nuc. Chg. (Grp I) = 1
Nuclear charge (Br) = 35
Effective Nuc. Chg. (Grp VII)= +7
12
Ionic Radii
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?
Remember …
The representative elements of groups 1, 2, 13 and 14 give up
electrons to form +1, +2, +3 and +4 ions respectively
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Positive Ions
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?
14
Positive Ions
The loss of the outermost valence shell has two effects:
1) The atoms radius shrinks because it loses it’s outermost
principle quantum level
AND …
2) The Nucleus now has more positive charge than the total
negative charge from electrons. The larger effective nuclear
charge will now pull electrons in closer to the nucleus 15
Positive Ions
Notice that even though the ionic electron config is the same
ionic radius gets progressively smaller moving across the period.
Na
Mg
Al
Atomic No.
11
12
13
Ionic charge
+1
+2
+3
Atomic Radius (pm)
190
160
143
Ionic Radius (pm)
99
65
50
Ionic Elec. Config
2-8
2-8
2-8
This happens because the positive eff. nuclear charge seen
by the same number of electrons increases as we move
across a the row
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Variation in atomic and Ionic Radii
What’s going on
Here?
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Negative Ions
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 expect the ionic radius to increase and it does
Cl atomic radius =
Cl- ionic radius =
99 nm
181 nm
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Negative Ions
Moving down groups the principal energy level increases
- This is true for all atoms, Anions (-) & Cations (+)
Li 2-1
Na 2-8-1
K 2-8-8-1
Li+ 2
Na+ 2-8
K+ 2-8-8
F 2-7
Cl 2-8-7
Br 2-8-18-7
FClBr-
2-8
2-8-8
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
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20
The ionic compound MgO
Mg atom
O atom
Electron Config
2-8
2-8
If we were to look at individual atoms Mg would
Actually be larger than Oxygen!
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