Reading: Chapter 4, the Blue Planet
Mineralogy is the study of minerals. The reason there is a lecture in this class entitled, "Earth" is that all of the solid parts of the universe are composed of minerals and we cannot hope to understand it without some knowledge of mineralogy. Also, minerals have great economic value (e.g. coal and diamonds). Finally most gem stones are minerals are crystals and gems are cool.
The definition of minerals used by geologists is: a naturally occurring substance that is inorganic, solid and has a specific chemical composition and an ordered atomic arrangement. According to this definition, then, synthetic diamonds are NOT minerals, while diamonds pulled out of mines are.
Rocks are usually aggregates of minerals which can be physically separated
from one another. A mineral is a chemical element or compound whose composition
can be represented by a chemical formula, for example Quartz is written
in chemical parlance as SiO2, which means one atom of silicon
for every two atoms of oxygen. While defined in this way, there can be
a range of allowed compositions. For example olivine may vary in composition
4 to Fe2SiO4 which
means that Magnesium and iron can substitute for one another. Finally,
the last criterion is that a mineral has an ordered atomic arrangement.
This means that a mineral is a specific crystal. For example, diamond and
graphite (pencil lead) are both made exclusively from carbon; their chemical
compositions are identical, yet their crystal structures are completely
different. They are NOT the same mineral of course. They are polymorphs
In this strict sense, a glassy material is not a mineral, although the book treats glasses as "amorphous" (having no form) minerals.
The key to mineralogy is in chemical bonding. Why do elements bond? What's in it for them?
What defines an element is the number of protons. Elements can have different numbers of neutrons, making different isotopes of the the element, but they always have the same number of neutrons.
In order to achieve electrical neutrality, atoms must have the same number of electrons as protons. If not, then the particle is charged and is called an ion. If there are too few electrons, the ion is postively charged and is called a cation. Negatively charged ions (too many electrons) are called anions.
Electrons orbit the nucleus, but they don't just fly around anywhere. They must stick in particular orbits. This is the essence of "quantum physics" if you've ever heard of that. It means that the energy that electrons have can only be certain amounts - it is "quantized". This is something like the difference between "analog" recording and "digital" recording. In analog devices, things can have any value, whereas in digital devices, things have discrete values. So electrons can only fly around in particular orbits called shells and each shell can have only so many electrons. The first shell can hold two electrons, and the second shell contains eight.
The key to all this is that matter likes to be electronically neutral (same number of protons and electrons), BUT, electronic shells like to be "full".
The simplest form of bonding in crystals is ionic bonding. In Figure 4.2 in the book you will see a picture of a flourine (F) atom and a lithium (Li) atom. Flourine, if you look up in the periodic table, you will find has nine protons (Atomic number 9). So, it would normally have 9 electrons: 2 in the inner shell, and only 7 in the outer shell. Lithium has 3 protons, hence has two in the inner shell, leaving one alone in theouter shell. Thus if lithium "loaned" one electron to flourine, making Li+ and F- ions. These now have opposite charge. They therefore attract one another and boom we get the LiF molecule. This kind of bonding is called "ionic bonding"). Sodium chloride (table salt) is another form of crystal that uses ionic bonding. The salt crystal is a very regular, cubic one . The cubic arrangement of atoms can be seen if you look very closely at your salt during dinner (amaze your friends in the cafeteria!) because salt grains are usually little cubes.
Another way of bonding is by covalent bonding. These are formed when two atoms "share" electrons, instead of "loaning" them. A good example of this is the silicate anion (see Figure 4.4). Because silicon has four electrons in its outer shell, it likes to share electrons with the oxygen atoms (each of which has two "extra" in the outer shell).
One beautiful example of an covalent bonding is the carbon molecule of diamond . Here the four electrons get shared between four other carbon atoms and the result is a tight tetrahetral crystal that is the hardest naturally occuring substance.
Finally, electrons of metals are very loosely bound and tend to "float" among the various atoms. This causes the high electrical conductivity of metals.
The important factors controlling crystal structure are the number of neighboring atoms or ions and their sizes. Ionic size is controlled by the number of electronic shells and the charge. Charge is important because if the ion has lost electrons (is a cation), then the postive charge of the nucleus pulls harder on the remaining electrons. Therefore, cations tend to be smaller than anions.
When the ions pack together, the dominant structure is determined by the close packing of the anions and the cations just fit themselves into the spaces left. Think about trying to make a pile with a bunch of oranges and grapes. You would first build the pile with the oranges and just stick the grapes in the spaces.
Because the cations just slip into the holes, it sometimes is possible to slip in different cations as long as they have the same charge. This is why magnesium and iron often are interchangable in a particular crystal structure (say olivine or mica).
As discussed last time, the Earth's crust is made out of mostly oxygen and silicon, with aluminum, iron and calcium contributing a few percent as well (see Table 2.1 in the book). Therefore most minerals are made of these things. There are several major groups of minerals based on the dominant element:
1) those composed of silicon plus other elements (the silicates). A common silicate that you have probably all heard of is Quartz (the major component in BEACH SAND!) This is SiO2.
Silicates do the heavy lifting as far as building the earth goes. The fundamental structure of all silicates is a tetrahedron formed by oxygen atoms with a tiny silicon atom at its heart. This structure is the silicate anion (SiO44-). These tetrahedra can be isolated from one another (as in olivine), or linked by sharing oxygen atoms to form rings, single chains , double chains, sheets or framework silicates (see Figure 4.6 in book).
2) those composed of oxygen plus other elements (the oxides). One common oxide is the often blood-red hematite (Fe2O3), also known as a semi-precious stone specularite.
3) those composed of the carbonate ion (the carbonates), such as calcium carbonate (the major mineral in marble!). This has a chemical formula of CaCO3.
4)Those forming minerals with sulpher (the sulfides and sulfates). Fools gold, or pyrite (FeS2) is one of these.
5)Those formed of native elements, such as metallic copper.
6) Salts (halides), such as NaCl.
How can we tell the various minerals apart? Since we can't actually see the atoms, how can we know which ones are which? Geologists have developed a set of criteria for distinguishing the most common (about 30) minerals using easily observed features. These include: hardness, cleavage, fracture, luster, color (especially of a scratched surface or streak), density, crystal habit and some simple chemical properties (does it dissolve in acid?).
Note: you should read the chapter carefully several times and the material will start to sink in. The next lectures will depend on you understanding a little bit about mineralogy and some common minerals so study up!