Geos 306, Lecture 4
Electronic Origins of Bonding, continued
We assume that all the atoms in metallic bonds are alike, in that
they all have diffuse electron densities. They would be the cations in
Like the cation in ionic crystals, these atoms give up their valence
electrons, but instead of giving the electrons to some other specific
atom, they are redistributed to all, and are shared by all. The
model is called "electron gas".
Eg. Na metal. 1s22s22p63s1.
Each Na atom gives up its 3s1 electrons. We end up with an array
of positive ions behaving like a gas of negative charged space.
The electron gas behaves like the glue.
Each of the valence electrons has about the same energy and since
they are so diffuse to begin with, they are free to move around from atom
to atom, or even place to place, within the crystal with no significant energy
Electron density map of copper
Since we assume that all the valence electrons have been given to the electron
gas, then the atoms are closed shell and assumed to be
just as in the ionic model.
Electron density of copper. Compare this with the electron density of NaCl.
Note that every nearest neighbor atom has a bond in Cu, but not in NaCl.
Electron density map of NaCl
Metallic shiny luster
Easy tendency to form alloys
Metallic shiny luster
Eg Li metal, 1s22s1: Each atom has 8 nearest neighbor atoms, 6
2nd nearest neighbors, spherical symmetry on 2s electronic orbitals.
All these 2s orbitals interpenetrate to give de-localized or multi-centered
orbitals throughout the whole crystal. This produces a spectrum of energies,
or a "band" of energies, with only a small energy gap between the
ground state and the excited state. There are a large number of these low-lying
excited states, all with energy differences in the visible light range.
Therefore, after striking a metallic object, light is immediately
absorbed and an electron jumps up a quantum level into an excited state.
This is why metals are opaque, all photons are captured at the surface.
In addition, the electron quickly drops back to the ground state, thereby
emitting a photon, (l change between incident wavelength and emmitted wavelength is small because
(wavelength) of incoming photon matched the DE of the electron transfer
so well) and we observe reflection.
Because the metal is made of spherical atoms in a sea of electrons
the atoms tend to pack like marbles in a box, with non-directional bonding.
If you hit a metal, say with a hammer, then the metal tends to deform rather
than break like an ionic crystal. This is related to the way that the atoms
are packed, along with the non-directional bonding, you can glide or roll them easily
over each other. The motion is actually related to defects in the crystal structure that are known as dislocations.
If a metal is hammered enough, then it hardens (called cold working). This is because
the large number of defects stops the gliding process.
The electrons are not strongly bonded to any individual atom, so they
are free to travel.
Because the atoms are considered to be positive spheres in a sea of
electrons then any similar sized sphere can fit right in without too much
trouble. Even dissimilar sized (eg small H) can fit into the interstices.
Since the atoms are packed together like marbles, you can put
a lot of them into a given volume.
At high pressures most elements become metallic.
A type of bond that is intermediate between ionic and metallic is the covalent
bond where atoms in groups (or molecules) of 2 or more share electrons.
Eg H2, N2, O2.
H:H where the resulting electronic configuration mimics the closed shell
around each H atom simultaneously.
These groups of atoms are quite stable and do not react easily.
The most important covalent bond is formed with carbon. We could think
of some C giving up 4 e- and others accepting 4 e-,
or all the carbons give up 4 e- into a sea to form a metal,
but this just does not happen. Instead, the electrons are shared along individual bonds.
Organic chemistry is the study of the carbon bond.
The first thing that we notice is that if the atoms are sharing the e-
then the e- must be located in between the atoms. Therefore
the atoms do not have spherical shapes.
Lewis dot diagrams.
Eg Cl2, 1s22s22p63s2
3p5=Ne3s23p5 to form molecules (0 dimensional)
Eg S, Ne3s23p4 forming polymers (1 dimensional)
Eg P, Ne3s23p3 forming sheets (2 dimensional)
Eg Si, Ne3s23p2 forming tetrahedral framework (3 dimensional)
The angular relationship between bonds is largely a function of the number
of electron pairs.
Eg 2 pairs, linear, eg CO2
Eg 3 pairs, triangular, CO3
Eg 4 pairs, tetrahedral, SiO4
The electron density model of bonding
Recent work centered on the analysis of the electron density around bonded
pairs of atoms show that the models discussed above for ionic, metallic
and covalent are over simplified.
The electrons around an atom do not really deform all that much in order
to form a bond. In general, the atoms really are quite spherical.
Ionic bonds typically form between pairs of atoms where one atom has diffuse electron
density and the electron density of the other falls off rapidly.
Metallic bonds typically form between atoms where both have diffuse electron density.
Covalent bonds typically form between atoms where both have electron densities that
fall off rapidly.
Nature does not put names on the bonds, they simply form between pairs
of atoms. Every bond is unique, i.e. the MgO bond is different from the
SiO bond, which is different from the CH bond, and so on.
Links neutral molecules together if the molecules display some sort of
dipole behavior. Eg. The crystal structure of ice is stabilized by the polarized bonds formed between oriented water molecules.
Eg: Surface of NaCl, Na exposed, H2O bonds to it with the O-side of the molecule facing the Na atom.
Where the Cl is exposed, then the water molecule bonds to Cl with an orientation such that
the H-side of the molecule is towards the Cl.
If the bonding in a molecule leaves one part of the molecule with more
charge than another, then we have polar bonds. This is why the neutral molecule H2O
can dissolve rocksalt.
Another example is the CO2 quadrupole, forming dry ice.
We assume that H, with only 1 electron, should be able to bond to only
one other atom. But experimental evidence suggests that this is not so.
There is often a second weak polar bond on the proton side of H. We should consider that
the hydrogen bond is a special case of a polar bond. For example, brucite, Mg(OH)2.
The hydrogen bonds hold the layers together.
van der Waals
Bonding between dynamic dipoles.
Eg graphite: C = sp2 hybrid.
Eg. He is solid at low T, here is a cartoon of the He2 system in its crystal structure.
"The assumption of van der Waals forces is usually made in the absence of better information
and is an admission of ignorance about the detailed nature of the intermolecular interaction."
Zallen, Slade and Ward (1971)
Suggested Reading: Wenk and Bulakh, 12-22; Klein, 53-57; Nesse, 46-56