Download 1) Configurations of ions 2) Trends in atom size (atomic

Chem 105
Friday 6 Nov 2009
1) Configurations of ions
2) Trends in atom size (atomic radius)
3) Trends in ion size
4) Ionization Potential
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Atomic Radius
- Measured in picometers (pm) 1 pm = 10-12 m
or Angstroms (Å)
1 Å = 100 pm = 10-8 cm
- Generally increase going down a group (down a column)
and decrease going across a period (L-to-R in a row)
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Group 1 Alkali metals
Group 8A Noble Gases
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One issue that arises with the term “Atomic Radius” is that the
numbers may differ depending on how they are obtained!
- “covalent radius” = half distance between bonded atoms
or
- “calculated radius” = distance out to arbitrary electron
density based on quantum mechanics calculation (Schrödinger
equation)
or
- “experimental” based on crystal of metal atoms = ½
interatomic distance
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Comparing the electron distribution in a H atom vs. H2 molecule
The 0.015 e-/Å3 contour
The 0.01 e-/Å3 contour
We define the “calculated atomic radius” =
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distance from nucleus out to electron density ~ 0.015 e-/Å3
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Comparing the electron distribution in a H atom vs. H2 molecule
H atom radius = 44 pm
H atom
H covalent radius = 37 pm
H2 molecule
H-H dist = 74 pm
Generally the covalent radius is smaller than the radius of free atom because the electrons in the
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molecule are attracted by two or more nuclei. This shrinks the whole electron cloud a bit.
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Crystal structure (experimental) of metallic sodium.
Na-Na distance = 365 pm; so, Na radius = 365/2 = 183 pm
365 pm
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Group 1 Alkali metals
Sodium: 184 pm
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Going from element-to-element DOWN a group,
you add a complete shell of electrons plus the
same number of protons in the nucleus.
For example, Group 2: Be, Mg, Ca…
e- in 1s orbital
e- in 2s orbital
4+
- 2e- -
Berylium atom
2e- in 3s orbital
-
12+
2e-
-
8e-
8e- in 2s,2p orbitals
Magnesium atom
Although the nuclear charge increases by 8+, adding a complete
inner shell of 8 electrons shields the outer shell electrons from the
increased positive charge.
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However, atoms get SMALLER going ACROSS a row Left-to-Right.
In this case, electrons are added to the same shell - on the periphery of the atom, and
the # of inner-shell electrons is constant.
The outer-shell electrons DO NOT shield each other from the increasing nuclear charge
because they are spread out with approximately same average distance from the
nucleus.
The nuclear charge increases by +1 for each electron added, and this added proton has
a much larger effect on all the electrons compared to the effect of the added electron.
12+
2e-
8e-
Magnesium atom
radius = 145 pm
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-
-
-
13+
2e8e-
-
Aluminum atom
radius = 118 pm
14+
2e8e-
-
Silicon atom
radius = 111 pm
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Transition metals decrease, then
increase slightly at end of the series
(Cu, Zn)
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21+
2e8e8e-
-
-
Scandium atom
radius = 144 pm
Decreases due to the increasing
nuclear charge which is not
shielded by outer electrons.
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-
-
-
26+
2e8e8e-
-
-
Iron atom
radius = 117 pm
-
-
-
-
30+
2e8e8e-
-
-
Zinc atom
radius = 125 pm
Increase (or constant depending on which
measurement). Now there are SO MANY
ELECTRONS in outer shell – they expand
due to mutual repulsion.
(covalent radii, Inorganic Chemistry, Miessler and Tarr, p 42)
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Sizes of Ions
Cations (remember ca + ion)
always SMALLER than corresponding atom (you’re removing electrons
– usually a whole shell - without changing the nuclear charge)
Anions
Always LARGER than corresponding atom (you’re adding electrons – to
complete a shell usually – without changing the nuclear charge.)
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12+
2e-
-
12+
2e8e-
-
8e-
Magnesium atom
radius = 145 pm
Magnesium 2+ ion
radius = 72 pm
Huge shrinkage - you’re stripping away the whole outer shell
-
17+
2e8e-
-
-
-
Chlorine atom
radius = 99 pm
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-
-
17+
2e8e-
-
-
Chloride ion (Cl-)
radius = 181 pm
Outer shell expands (a lot)
because you’re adding an
e- without adding a proton
in the nucleus.
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Atom and Common Anion Size Comparison
These 3
anions
have 10 e-
radii in
picometers pm
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Cations
radii in
picometers pm
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Place the following atoms in order of increasing
atomic radii: K, Mg, Ca, Rb...
K < Mg < Ca < Rb
K < Mg < Rb < Ca
Mg < Ca < K < Rb
K < Rb < Mg < Ca
Mg < K < Ca < Rb
N = 75
21%
9%
a
R
<
C
<
K
b
M
g
<
R
K
<
<
g
M
a
M
g
<
<
Ca
g
M
<
K
<
R
<
K
b
R
<
C
<
g
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<
M
Rb
Ca
K
K
C
a
C
<
R
<
a
Mg
b
1%
b
1%
b
1.
2.
3.
4.
5.
67%
Excellent work,
chem students! JK
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(OWL question on ion sizes)
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Ionization energy (= ionization potential) definition
A (g) ---> A+ (g) + 1eA+ (g) ---> A2+ (g) + 1e-
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∆E = 1st ionization energy (kJ/mol)
∆E = 2nd ionization energy (kJ/mol)
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First Ionization Energies
H(g) H+(g) + e-
ΔH = +1312.0 kJ
(For comparison, the thermite reaction gives off way less energy per mole of iron
oxide consumed.
Fe2O3(s) + 2 Al(s) Al2O3(s) + 2 Fe(l)
He(g) He+(g) + e-
ΔH = -851.5 kJ)
ΔH = +2372.3 kJ
The first ionization energy for helium is about twice the ionization energy for hydrogen
because each electron in helium feels the attractive force of two protons, instead of
one.
Far less energy is required to remove an electron from a lithium atom, which has
three protons in its nucleus.
Li(g) Li+(g) + e-
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ΔH = +572.3 kJ
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The 1st ionization energy (A --> A+ + e-)
decreases going down a group.
He
Ne
Group 8A
Ar
Kr
Xe
Li
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Na
K
Rb
Cs
Rn
Fr
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The 1st ionization energy (A --> A+ + e-)
decreases going down a group.
Li
Na
K
Rb
Cs
Fr
Group 1A
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Electron in 2s orbital
Li
Na
+
Electron in 3s orbital
Electron in 4s orbital
-
+
As you go down a group, the outermost
electron(s) are further from nucleus,
and are easier to remove.
-
This is the same order as the chemical
reactivity of these metals as reducing
agents. (“Donate electrons”) K > Na > Li
K
+
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Another way to put this is that, as you go
to larger atoms in the same group, the
effective nuclear charge decreases due
to shielding by inner electrons.
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The 1st ionization energy (A --> A+ + e-) generally
increases going (L-to-R) across a row.
He
Ne
Ar
Kr
You’re adding a proton in nucleus and
electron around the periphery of atom.
Electrons in the same shell do not shield
each other from the nuclear charge –
too spread out.
Xe
Li
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Na
K
Rb
Cs
Rn
Fr
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In Period 2, B,C, and N (and O, F, and Ne even more so) all
have slightly lower ionization potentials than expected.
He
Ne
l
ai
t
n
e
t
o
P
n
o
it
az
i
n
o
I
ts
1
F
N
H
Be
C
O
Mg
B
Li
0
1
2
3
Na
4
5
6
7
8
9 10 11 12 13
Atomic Number ->
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B and O have anomolously low 1st ionization potentials, which
means an outer electron is unexpectedly easy to remove.
Boron, carbon, and nitrogen outermost electrons go
into a p-orbital, which is less stable than an s-orbital
().
Oxygen, fluorine and neon outermost electrons go into
a two-electron p-orbital () . These are further
destabilized by electron-electron repulsion within the
orbital.
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The End
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