CHAPTER 4

Metals have quite a wide range of properties at normal temperatures, from liquid (like mercury) to extremely hard (like tungsten). Most are shiny, but not all are colorless - gold and copper have distinct colors. All metals conduct electricity, but not all equally. How are atoms arranged in a metal? Let us use aluminum as an example - most of us have something made of aluminum, such as a pan or aluminum foil. With modern microscopes, it is actually quite easy to “see” atoms nowadays, since we can use a variety of techniques to image where the Al atoms are in the solid structure.

3.1 Elements & Bonding
3.2 Interactions
3.3 Carbon
3.4 Metals

What emerges is a picture of aluminum nuclei and their core electrons, packed like spheres where one layer of spheres rests in the interstices of the underlying and overlying layers (but remember the uncertainty principle!)

In contrast to H-H or diamond, where the electrons involved in bonding are located (most probably) between the two nuclei, in aluminum and other metals, the valence electrons are not closely associated with each nucleus. Instead they are dispersed over the whole macroscopic piece of metal. Imagine that instead of two or three or four atomic orbitals combining to form molecular orbitals, a mole (6 x10²³) atomic orbitals are combined to produce a mole of molecular orbitals. As more and more MOs are formed - the energies between them will get smaller and smaller and for a macroscopic piece of metal (that is one you can see) for all intents and purposes the energy gap between the individual bonding MOs will be negligible. These orbitals produce what is an essentially a continuous band of (low energy) bonding molecular orbitals and a continuous band of (higher energy) anti-bonding molecular orbitals. The energy gap between the bonding and anti-bonding orbitals is called the band-gap, and in a metal this band-gap is quite small (recall that the gap between the bonding and antibonding MOs in diamond is very large). Moreover, in metals the bonding molecular orbitals (known as the valence band) are able to accommodate more electrons. This is because in metals there are typically fewer electrons than there are atomic orbitals. Consider aluminum: it has three valence electrons and in the ground (lowest energy) state has an electron configuration of 3s2 3p1, this implies that it has two unoccupied 3p orbitals (although what the meaning of an unoccupied orbital is not completely clear.) We can consider the bonding molecular orbitals in aluminum to be formed from all the available atomic orbitals - which means that there are many bonding MOs that are not occupied by electrons. The physical consequences of this are that the valence electrons can move relatively easily from MO to MO, since their energies are very close together. While nuclei and core electrons remain more or less locked in position, the valence electrons can “spread out” forming a kind of “electron sea” within the metal. When an electrical potential 80 is applied across the metal, electrons from an external source can easily enter the valence band and electrons can just as easily leave the metal. Electrical conductivity is essentially a measure of how easily electrons can flow through a substance. Metals typically have high conductivity due to the ease with which electrons can move from one MO to another, and the fact that each MO extends throughout the whole piece of metal. Since the numbers of electrons entering and leaving are the same, the piece of metal remains uncharged.

In this model, the atomic cores are packed together and are surrounded by a cloud of electrons that serve as the “glue” that binds them together. There are no discrete bonds in this type of structure. When a piece of metal is put under physical stress – for example it is stretched or deformed, the atoms can move relative to one another, but the electrons remain spread throughout the structure. Metals can often be slowly deformed into different shapes without losing their structural integrity or electrical conductivity – they are malleable!

They can be melted (increased atomic movement), become liquid, and then allowed to cool until they solidify; throughout this process, they retain their integrity and their metallic properties - they continue to conduct electricity. This is quite different from other substances (such as diamond or water). The hardness of a (solid) metal depends on how well its atoms packed together and how many electrons are contributed to the valence band of orbitals.

So why do some elements behave as metals and others do not? For example, graphite conducts electricity, but is not metallic in other ways - it is not malleable - and can’t be heated and molded into other shapes. The answer lines in the behavior of the molecular orbitals and the resulting bonds they can produce. Graphite has a rigid carbon backbone of carbon-carbon bonds that makes it strong and stable, but overlaying that is the set of delocalized molecular orbitals that spread out over the whole sheet. As a result graphite has some properties that are similar to diamond (stability and strength) and some that are like metals (electrical conductivity) and some that are a consequence of its unique sheet structure (slipperiness).

Why are metals shiny?

We see things because photons hit the back of our retinas, are absorbed by specialized molecules (proteins and associated pigment molecules). This leads to changes in protein structure and initiates a cascade of neuron-based cellular events that alters brain activity. So where do these photons come from. First and foremost they can be emitted from a source (the sun, a lightbulb, etc.) and such a source appears to shine and can be seen in the dark. Alternatively, photons can be reflected off a surface - most of the things we see do not emit light, but rather reflect it. A red T-shirt appears red because it absorbs other colors and reflects red light. Photons can also be refracted when they pass through a substance. A cut diamond sparkles because light is refracted as it passes though the material and exits from the many facets. Refraction is caused when photons bump into electrons, are absorbed and then (very shortly thereafter) are re-emitted, as they travel through a material. These processes take time, so the speed of light slows down. It can take a photon many thousands of years to move from the core to the surface of the sun because of all the collisions that it makes during the journey.

To explain why metals (and graphite) are shiny we invoke a combination of reflection, refraction, and the energy levels of molecular orbitals. When a photon of light is absorbed and reemitted, the electron moves from one orbital to another. Let us consider a piece of metal at room temperature. When a photon arrives at the metal’s surface it encounters the almost continuous band of molecular orbitals. Most photons, regardless of their wavelength, can be absorbed, promoting electrons up to a higher energy level. As the electrons drop back down to a lower energy level, the photons are re-emitted, resulting in the characteristic metallic luster – metals actually emit light. This does not mean metals glow in the dark (like a light bulb or the sun). Metals absorb and re-emit photons at room temperature.

The color of a particular metal (at room temperature) depends upon the range of wavelengths that are re-emitted. For most metals the photons re-emitted have a wide range of wavelengths, this makes the metallic surface colorless, or silvery. A few metals, such as copper and gold, absorb light in the blue region and re-emit light whose wavelengths are biased toward the red end region of the spectrum (400-700 nm) and therefore they appear yellowish. This is due to “relativistic effects” way beyond the scope of this book – but something to look forward to in your future physical chemistry studies!

Now we can also understand why metals emit light when they are heated. The kinetic energy of the atoms increases with temperature, which promotes electrons from low to higher energy orbitals, when these electrons lose that energy by returning to the ground state, it is emitted as light. The higher the temperature, the shorter the wavelength of the emitted light. As a filament heats up, it first glows red, and then increasing whiter as photons of more and more wavelengths are emitted.

3.1 Elements & Bonding
3.2 Interactions
3.3 Carbon
3.4 Metals

Question to answer:

• What properties indicate that a substance is metallic?
• Why are metals shiny?
• How can metallic properties be explained by the atomic-molecular structure of Al (for example)
• Why can we see through diamond but not aluminum - how about graphite?
• Why does aluminum (and for that matter all metals) conduct electricity? What must be happening at the atomic molecular scale for this to occur?
• What does the fact that diamond doesn’t conduct electricity tell you about the bonding in diamond?
• How do the bonding models for diamond and graphite explain the differences in properties between diamond, graphite, and a metal like aluminum?
• Why is is OK to use different models to describe bonding in different species?