Periodic Trends: Size
Trends in Atomic Radii
Imagine transforming lithium to beryllium in two steps:
1. Add an electron to Li.
2. Then, add a proton (plus neutrons, but they don’t count) to
the nucleus.
As we go across the rows of the periodic table, the electron-electron
repulsion from adding an electron to a same-shell orbital (2s, 2p,
3s, 3p, etc.) is a decrease in effective charge that is always less than
the increase in nuclear charge.
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Ionic Radii
Adding or removing electrons from an atom will always change its
size:
These effects become quite large when adding an electron or removing it empties a valence shell (Na → Na+ ) or starts a new one
(Xe → Xe− .)
Isoelectronic Series
Once we add or remove electrons to form ions, we can compare
multiple species (anions, neutrals, cations) that have different nuclear charges but the same electronic configuration. Here are the
electronic configurations of oxygen, fluorine, neon, and sodium:
2
and here are O2− , F− , Ne, and Na+ :
3−
N
O2−
F−
Ne
Na+
Mg2+
Al3+
R (pm)
146
140
133
?
102
72
53.5
3
Z
7
8
9
10
11
12
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Ionization Energy
Neutral atoms are stable: they do not spontaneously lose electrons.
The process
X −→ X+ + e−
will always take energy, because the positively charged ion will
always be attracted to the negatively charged electron. The (positive) energy input required is the ionization energy, usually measured in eV or in kJ/mol, with
1 eV/atom = 96.49 kJ/mol
Elements that have multiple electrons will have multiple ionization
energies:
X(g) −→ X+ + e−
X+ −→ X2+ + e−
X16+ −→ X17+ + e−
Na
Mg
Al
Si
P
S
Cl
Ar
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[Ne]3s
[Ne]3s2
[Ne]3s2 3p1
[Ne]3s2 3p2
[Ne]3s2 3p3
[Ne]3s2 3p4
[Ne]3s2 3p5
[Ne]3s2 3p6
I1
3s
3s
3p
3p
3p
3p
3p
3p
I1 = 1st ionization energy
I2 = 2nd ionization energy
I17 = 17th ionization energy
I2
2p
3s
3s
3p
3p
3p
3p
3p
4
I3
I4
I5
I6
I7
2p
3s
3s
3p
3p
3p
3p
2p
3s
3s
3p
3p
3p
2p
3s
3s
3p
3p
2p
3s
3s
3p
2p
3s
3s
with energies in kJ/mol:
Na
Mg
Al
Si
P
S
Cl
Ar
I1
495.8
737.7
577.5
786.5
1011.8
999.6
1251.2
1520.6
I2
4562.4
1450.7
1816.7
1577.1
1907.5
2251.8
2297.7
2665.9
I3
I4
I5
I6
I7
7732.7
2744.8
3231.6
2914.1
3357
3822
3931
11577.5
4355.5
4963.6
4556.2
5158.6
5771
16090.6
6274.0
7004.3
6540
7238
21267.4
8495.8
9362
8781.0
27107.4
11018.2
11995.3
Chemistry—ionic or covalent—involves valence electrons
only.
Periodic trends in first ionization energies:
1. Ionization energy increases as a shell fills.
2. There is generally a small dip in ionization energy when an
electron is added after a filled subshell.
3. Adding a fourth electron to the p subshell produces a small
dip in ionization energy.
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4. Ionization energies tend to decrease moving down a column.
5. Putting it all together, ionization energies are lowest in the
lower left of the periodic table (Rb and Cs are most easily
ionized to Rb+ and Cs+ ) and highest in the upper right (He
to He+ ).
Transition Metals and Lanthanides
Electron Affinities
The electron affinity is to anions what the ionization energy is to
cations:
X(g) + e− −→ X−
energy change = EA1
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In general, electron affinities:
1. become more negative from left to right on the periodic table,
2. are positive for noble gases,
3. are positive for elements with s2 and d10 configurations,
4. become more negative moving from heavier elements within
a column to lighter—though the first-row elements (BCNOF)
are an exception to this.
As with ionization energy, you can also have a second, third, fourth,
etc. electron affinity; e.g.
X− + e− −→ X2−
energy change = EA2
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energy required to make cation
energy required/gained to make anion
energy gained by bringing ions together
Na −→ Na+ + e−
Cl + e− −→ Cl−
Na+ + Cl− −→ NaCl(s)
Electronegativity
There are multiple electronegativity definitions to use. The simplest to understand (based on what we’ve done so far) is the Mulliken definition:
IP + |EA|
ξ=
2
The electronegativity is, in many ways, a half-way point between
ionization energy and electron affinity, and gives a general measurement of whether an element is going to pick up or give away
an electron in any given compound. Electronegativity trends arise
from those of ionization energy and electron affinity:
1. Electronegativity tends to increase from left to right (Li to
F) and from bottom to to (Cs to Li, or I to F).
2. The non-metals in the upper right corner of the periodic table
are elements with high electronegativity (ξ ≥ 2.2), and tend
to gain electrons in chemical reactions—they are oxidants.
3. Metals (and in particular, metals in the lower left corner) are
elements with low electronegativity (ξ ≤ 1.8) and tend to lose
electrons in chemical reactions—they are reductants.
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always positive
sometimes negative, sometimes positive
always negative