The bonding of metals
When two atoms combine, different types of bonding can occur: covalent, ionic, and metallic. Silver, iron, platinum, gold, and copper all form metallic bonds. Unlike covalent bonding, metallic bonding is non-directional. The strong bond consists of positively charged metal atoms in fixed positions, surrounded by delocalized electrons. These delocalized electrons are often referred to as “a sea of electrons,” and can help explain why copper and gold are yellow and orange, while most other metals are silver.
BAND THEORY
The color of metals can be explained by band theory, which assumes that overlapping energy levels form bands.
The mobility of electrons exposed to an electric field depends on the width of the energy bands, and their proximity to other electrons. In metallic substances, empty bands can overlap with bands containing electrons. The electrons of a particular atom are able to move to what would normally be a higher-level state, with little or no additional energy. The outer electrons are said to be “free,” and ready to move in the presence of an electric field.
Some substances do not experience band overlap, no matter how many atoms are in close proximity. For these substances, a large gap remains between the highest band containing electrons (the valence band) and the next band, which is empty (the conduction band). As a result, valence electrons are bound to a particular atom and cannot become mobile without a significant amount of energy being made available. These substances are electrical insulators.Semiconductors are similar, except that the gap is smaller, falling between these two extremes.
When light strikes the surface of a metal, electrons in a lower energy level can be excited to a higher energy level. The distance between the levels represents the relative energy required to excite an electron. When four atoms combine, the outermost energy levels merge, providing four energy levels at a low energy and four energy levels at a higher energy. As the number of neighboring atoms increases, the spacing between the energy levels decreases. More overlap occurs and bands of low and high energy replace the distinct energy levels. As more atoms combine, the distance between the two bands decreases, the band gap decreases, and less energy is required for the electron to be excited from one band to the other. In metals, when very large numbers of atoms are brought close to each other, the low and high energy bands can overlap, forming a nearly continuous band of available energy levels, where electrons may move freely
Above the Fermi level, energy levels are empty (empty at absolute zero), and can accept excited electrons. The surface of a metal can absorb all wavelengths of incident light, and excited electrons jump to a higher unoccupied energy level. This creates current, which rapidly discharges to emit a photon of light of the same wavelength. So, most of the incident light is immediately re-emitted at the surface, creating the metallic luster we see in gold, silver, copper, and other metals. This is why most metals are white or silver, and a smooth surface will be highly reflective, since it does not allow light to penetrate deeply
Silver, gold and copper have similar electron configurations, but we perceive them as having quite distinct colors. Electrons absorb energy from incident light, and are excited from lower energy levels to higher, vacant energy levels. The excited electrons can then return to the lower energies and emit the difference of energy as a photon.
If an energy level (like the 3d band) holds many more electrons (than other energy levels) then the excitation of electrons from this highly occupied level to above the Fermi level will become quite important. Gold fulfills all the requirements for an intense absorption of light with energy of 2.3 eV (from the 3d band to above the Fermi level). The color we see is yellow, as the corresponding wavelengths are re-emitted. Copper has a strong absorption at a slightly lower energy, with orange being most strongly absorbed and re-emitted. In silver, the absorption peak lies in the ultraviolet region, at about 4 eV. As a result, silver maintains high reflectivity evenly across the visible spectrum, and we see it as a pure white. The lower energies (which in this case contain energies corresponding to the entire visible spectrum of color) are equally absorbed and re-emitted.
Silver and aluminum powders appear black because the white light that has been re-emitted is absorbed by nearby grains of powder and no light reaches the eye.
Metals are colored because the absorption and re-emission of light are dependent on wavelength. Gold and copper have low reflectivity at short wavelengths, and yellow and red are preferentially reflected, as the color here suggests. Silver has good reflectivity that does not vary with wavelength, and therefore appears very close to white.
Showing the variation of density of states and the excitation seen in gold, silver and copper (actually from the 3d band to above the Fermi level)
CRYSTAL FIELD 