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Electron affinities:
The electron affinity is defined (in the real world, not necessarily the world of the
“Zumster”) by the following relationship:
A + e- → A-
EA = -∆H
So, the vertical scale of the plot in Zumdahl should be reversed in sign to be consistent
with the conventions of the National Institute of Standard and Technology, for example.
Shown here is a plot to indicate how electron affinities vary across the first two rows of
the Periodic table.
Local maxima occur one atom before a
subshell is filled or half-filled. So, C
(2p2) is a local maximum, because C- is
spherically symmetric. F (2p5) is also a
local maximum, because the driving force
for forming the filled shell of neon is
particularly strong.
Li and Na are local maxima because of the
stability of the 2s2 and 3s2 configurations.
The plot below shows especially clearly
that the species with largest electron
affinities are the halogens, and F and Cl
have especially high values. On this
particular Periodic Table, the columns are
labeled 1 to 18, and it is the group 17 species with the largest affinities.
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Atomic Size
Another important periodic property is atomic radius. This parameter can be measured a
number of ways. Interatomic spacings in crystals (metals, rare gases) determined by
spectroscopy or by X-ray crystallography can give atomic separations. From
experimental data like these, a consistent set of atomic data can be determined.
General trends:
• Across a row, the principal quantum number is fixed. The atomic size
decreases because of the increase in nuclear charge.
•
Down a column, the outermost electrons are, on average, further from the
nucleus that the species immediately above it in the Periodic Table. Thus,
the atomic size increases down a column
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Isoelectronic Species
A very useful comparison with atoms (and molecules too) is the look at properties of
systems with the same electron configuration. Such species are isoelectronic.
Let’s look at Ne, with a 2p6 electron configuration:
Other species with that same configuration are: Na+, Mg2+. We get these species by
removing the 3s electron(s) from the first two atoms in the next row of the periodic table.
Also, F-, and O2- have the same rare gas configuration.
So, for the following isoelectronic series
O2-, F-, Ne, Na+, Mg2+
the nuclear charge increases left to right. So, we expect the size to decrease left to right.
Can you tell me why?
How about the energy to remove the last 2p electron? Should be highest for Mg2+ and
lowest for O2-, right?
Isoelectronic arguments are very important in chemistry
Chemical Bonding
We have already observed the phenomenon of chemical bonding in our studies of
thermochemistry. Remember when two atoms A(g) and B(g) form a bound species
AB(g), the potential energy of the bound system is lower than the energies of the
separated atoms.
A(g) + B(g) → AB(g)
∆H = -D(A-B)
We have described how real thermodynamic cycles can be used to determine accurate
bond energies. This is really how bond energies are determined!
What happens when a chemical bond is formed? The electrons on the individual atoms
are shared between the atoms. The extent of this sharing determines the type of
chemical bond that is formed. The next figure shows how in the hydrogen molecule, the
electrons are shared equally between the atoms. Clearly, in such a homonuclear system,
there is no way the electrons could not be shared equally on the average.
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Bonding Systematics in Molecular Hydrogen
Note the increase in electron density between the atoms, which is caused by the sharing
of the electron pair between the atoms. The electron pair bond is the basic bond type that
we will discuss.
There are essentially three different kinds of electron pair chemical bonds, which depend
of the extent to which the electrons are shared equally between the atoms. The H-H
system is the classic example of a system with perfect covalent binding. We will
consider these types of binding:
Types of chemical bonds:
• Covalent
• Ionic
• Metallic
Electronegativity
The physical parameter denoted electronegativity is what allows us to distinguish among
these types of bonding. Electronegativity is the tendency of an atom to attract the shared
electrons in a chemical bond.
The electronegativity is defined by Pauling is a function of ionization potential and
electron affinity of the atom. The difference in electronegativity between atoms that
bind together is a measure of the percent ionic character of the bond.
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Electronegativity is large (4.0 is the maximum value) for the fluorine, atom indicating its
ability to attract shared electrons. Metals tend to have small electronegativities,
consistent with their ability to donate their electrons to the other atom in a bond. The Cs
atom has the smallest electronegativity of any atom, with a value of 0.8.
∆(EN):
0
covalent
→
→
~3
ionic
Electronegativities are shown on the chart below: note that the most ionic compounds
results from species from the upper right of the Periodic Table bind with those from the
lower left.
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The next graph shows the percent ionic character associated with such systems:
Dipole Moments
The idea of percent ionic character leads naturally to a discussion of how the center of
positive and negative charge in a bond are displaced from one another.
If a bond has 100% ionic character, and the separation between the atoms is 10-10 m (1
Å), we have the following situation:
+
10-10 m
Suppose each charge is 1.6 × 10-19 Coulombs in magnitude
We define the dipole moment µ as QR = (1.6 × 10-19 Coulombs)(10-10 m) = 1.6 × 10-29
C-m. We define 1 Debye as 0.3336 × 10-29 C-m
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So, the dipole moment for this particular charge distribution is
1.6 × 10-29 /(0.3336 × 10-29 C-m /Debye) = 4.8 Debye
1 Debye is the dipole moment associated with equal and opposite charge of magnitue
0.208 e- separated by 1 Å.