Describe the basic chemistry involved in mineral formation and structures.
In this section, you will learn the basic chemistry associated with the different types and classes of minerals.
What You’ll Learn to Do
- Identify the building blocks of matter.
- Differentiate between different kinds of atomic bonds.
The Building Blocks of Matter
Why study the chemistry of minerals? Minerals are made of atoms, which have an impact on the behavior and characteristics of the mineral. Thus, to understand, explain, and predict the behavior of minerals, and rocks—which are made of minerals—we must understand some basic facts about atoms and how they behave. This requires a basic understanding of some chemistry. We will begin by constructing atoms in our thinking in terms of the three sub-atomic particles of which atoms are made.
Atoms consist of protons, neutrons, and electrons. Protons have a positive (+) electrical charge. Electrons have a negative (−) charge that is exactly equal and opposite to the electrical charge of a proton. Neutrons are electrically neutral.
Most of the mass of an atom is packed into its tiny nucleus. An atomic nucleus is made of protons and neutrons, which have approximately the same mass (about 1.67 × 10−24 grams). Electrons, on the other hand, are arranged in specific orbitals around the nucleus of an atom; they are also much smaller in mass than protons and neutrons, weighing only 9.11 × 10−28 grams, or about 1/1800 the weight of protons and neutrons.
Even though the mass of an electron is a tiny mass compared to the mass of a proton or a neutron, the electrons occupy most of the volume of an atom (see Figure 1).
A neutral atom has the same number of electrons as it does protons. An atom that has lost or gained any electrons is no longer an electrically neutral atom. That type of atom, which is not electrically neutral and has an electrical charge associated with it, is called an ion. Atoms that have gained electrons are negatively (−) charged ions, or anions. Atoms that have lost electrons are positively (+) charged ions, or cations.
It is also possible to have ions that are actually small groups of atoms bonded together. These are known as polyatomic ions. One example of a polyatomic ion is the carbonate ion, (CO3)2−, which has two extra electrons, giving it the net electrical charge of 2−.
We’ve just seen that a carbonate ion can also be called (CO3)2−. But what exactly does this mean?
First let’s look at the letters: CO. Atoms have chemical symbols; each element has been assigned one or two letters to represent it. Thus, C stands in for carbon, and O stands in for oxygen (all of these chemical symbols can be seen in the periodic table in Figure 2 below).
As we read above, the 2− means that the ion has two extra electrons. But what about the 3? This means there are three oxygen atoms in the ion. The number of atoms in a particular formula is always notated in subscript. Charge is always written in superscript at the end of the formula (a superscript at the beginning of the formula means something else—we’ll get to this when we discuss isotopes below). The parentheses around CO3 indicate that the charge belongs to the whole polyatomic unit rather than just the O3.
Thus the carbonate ion is one carbon atom (C), three oxygen atoms (O3), and two extra electrons (2−), which charge the whole polyatomic ion.
The Periodic Table
Naturally occurring atoms found in the earth range from hydrogen, with just one proton in its nucleus, to uranium, with 92 protons in its nucleus. These are the naturally occurring chemical elements, which includes such commonly known elements as carbon, oxygen, iron, and so on. The periodic table lists all the chemical elements in a way that tells us how many protons each of them has, how its electrons are arranged, and what the general chemical behavior of each element is.
As shown in Figure 2, the periodic table consists of eighteen groups and seven periods. Two additional rows of elements, known as the lanthanides and actinides, are placed beneath the main table. These elements are placed separately to make the table more compact. All the elements in a group have a similar chemical behavior. This is because all the elements in a group have a similar arrangement of electrons in their atoms, and it is the electron arrangement that determines the chemical behavior of an element.
For each element, the name, atomic symbol, atomic number, and atomic mass are provided. The atomic number is a whole number that represents the number of protons: each chemical element is distinguished by the number of protons in its nucleus. For example, every atom of the element oxygen has eight protons in its nucleus. That is why the atomic number of oxygen is 8. If an atom has greater or fewer than eight protons in its nucleus, it is not oxygen, it is some other chemical element. In the periodic table, the atomic number of each element is listed above the chemical symbol of the element.
The atomic mass, which is the average mass of different isotopes, is estimated to two decimal places. For example, hydrogen has the atomic symbol H, the atomic number 1, and an atomic mass of 1.01. The atomic mass is always larger that the atomic number. For most small elements, the atomic mass is approximately double the atomic number, as the number of protons and neutrons is about equal.
The elements are divided into three categories: metals, nonmetals and metalloids. These form a diagonal line from period two, group thirteen to period seven, group sixteen. All elements to the left of the metalloids are metals, and all elements to the right are nonmetals.
The periodic table was created to help chemists better understand elements and how they function. It is a map to elemental behavior.
Every atom of a specific element must have the same number of protons in its nucleus. This number is its atomic number. However, there is a range of possible numbers of neutrons it can have in its nucleus. The fact that atoms of a chemical element may have different numbers of neutrons results in each chemical element having several isotopes. Isotopes are atoms of a given chemical element that have different numbers of neutrons in their nuclei.
For example, while all atoms of the element oxygen have eight protons in their nuclei, those oxygen atoms may have eight, nine, or ten neutrons. The different numbers of neutrons in the nucleus distinguishes the three isotopes of oxygen. Oxygen-16 is the isotope of oxygen with 8 neutrons in its nucleus. The number 16 is called the atomic mass number. The atomic mass number is the total number of protons and neutrons in the nucleus of an isotope. From this definition, and knowing that all oxygen atoms have 8 protons in the nucleus, you can deduce that oxygen-17 is the oxygen isotope with 9 neutrons and oxygen-18 is the oxygen isotope with 10 neutrons. Abbreviated into symbols, the three isotopes of oxygen are written as 16O, 17O and 18O.
Isotopes are not very important for understanding minerals, but are important in understanding how to apply chemistry and nuclear physics to geology, such as how to use measurements of radioactive isotopes to measure the ages of rocks and minerals and how to use oxygen isotopes from layers of glacial ice to determine what the temperature of the earth was during an ice age.
Minerals form as a result of chemical reactions. Chemical reactions are driven mainly by the arrangement and rearrangement of electrons in atoms. In a mineral, the atoms are held together by chemical bonds, which derive from the electrons.
Electrons can be thought of as occupying energy levels, or shells, in an atom. The lowest-energy shell is closest to the nucleus. Each shell can accommodate only a limited number of electrons. The first shell can hold two electrons, the second and third shells can hold eight electrons, and the next two shells can hold eighteen electrons. Unless energy is added to an atom to excite it from its low-energy state, the electrons in the atom will occupy the lowest-energy shells available to them.
The Bohr model (see Figure 3) was developed by Niels Bohrs in 1913. In this model, electrons exist within principal shells. An electron normally exists in the lowest energy shell available, which is the one closest to the nucleus. Energy from a photon of light can bump it up to a higher energy shell, but this situation is unstable, and the electron quickly decays back to the ground state.
If atoms interact with other atoms, they can gain or lose electrons to the other atoms, or share electrons with other atoms. In an individual atom, the most stable arrangement is a full outer shell of electrons. Therefore, chemical reactions will occur, and chemical bonds will form that hold atoms together to each other, when atoms encounter other atoms and change their electron configurations toward more stable, lower-energy arrangements, which generally involves achieving full outer electron shells in the atoms.
This stable configuration—a full outer shell of electrons—is exemplified by the inert gases. In the periodic table the inert gases are the elements of group 18 or VIIIA, the last column on the right. Inert gases do not have to undergo any chemical reactions or form any chemical bonds with other atoms in order to have a full outer shell of electrons. The inert gases already have full outer shells of electrons. That is why they are chemically inert. Their electrons are as stable as can be arranged. For this reason, inert gases are extremely unlikely to undergo any chemical reactions and it is almost impossible for them to bond with any other atoms. Because they do not bond with any other atoms to form a liquid, a solid, a molecule, or a mineral, the inert gases consist of atoms that stay separate from each other, in the gas state.
Individual atoms of all the other chemical elements, when they are neutral atoms, do not have full outer shells of electrons like the inert gases do. Therefore, they do not have the most stable arrangement of electrons that they possibly can. That is why most chemical elements have a strong tendency to either gain or lose electrons, or to enter into other arrangements of their valence electrons, the electrons in their outer shell. Chemical reactions and chemical bonds are generally a result of electrons being rearranged within and among atoms to give the atoms full outer electron shells.
For an atom to lose or gain one electron takes less energy than to lose or gain two, which in turn takes less energy than to lose or gain a third electron. For an individual atom to gain or lose four electrons will only occur in extremely high-energy environments such as in a star. In common chemical reactions on earth, and in the formation of chemical bonds, no element will completely gain or lose four electrons. This limits the charges of atomic cations to +1, +2 or +3 and the charges of atomic anions to –1, –2, or –3.
Reading this far, you have learned about one group of elements in the periodic table, group 18, the inert gases. Another group of chemical elements in the periodic table is the alkali elements. The alkali elements compose group 1 or IA, the left hand column, including the elements sodium (Na) and potassium (K).
Hydrogen is not usually considered as an alkali element because, even though it is shown in group 1 in the periodic table. Hydrogen is so light and small, with just a single proton in its nucleus, that it has some unique behaviors and is considered in a class by itself.
The alkali elements have a single electron in their outer electron shell. If an alkali element loses a single electron, it becomes an ion with a +1 charge and a full outer shell. If an opportunity arises, alkali elements will easily turned into +1 cations.
Now look at group 17 or VIIA in the periodic table, which includes the chemical elements fluorine (F), chlorine (Cl) and so on. These are the halogen elements. If a halogen element gains a single electron, it becomes an ion with a–1 charge and a full outer electron shell. If an opportunity arises, halogen elements have a strong tendency to take in an extra electron and become -1 anions because by doing so they achieve a full outer shell of electrons, which is the most stable arrangement of electrons possible.
If sodium and chlorine atoms get together in the right conditions, such as in an evaporating solution of salt water, each sodium atoms will give up an electron to a chlorine atom. This turns the sodium atoms into sodium ions, Na+, and the chlorine atoms into chloride ions, Cl–. Opposite electrical charges attract, so the sodium ions and chloride ions will tend to stick together with each other, joined by what are called ionic bonds.
Not only will the sodium and chloride ions have a very strong tendency to join together with each other via ionic bonds, in most situations they will naturally arrange into a configuration where there is no wasted space and no wasted energy. This leads them to form the crystal lattice of the mineral halite. Halite is a mineral with the chemical formula NaCl, sodium chloride, in which the bonds between the atoms are all ionic bonds.
Look at the diagram of halite showing the sodium and chloride ions arranged into the crystal lattice. All the ionic bonds are at the same angle and the same distance, so they are all of equal strength. This is the lowest-energy arrangement of the ions, the most stable arrangement. If any of the ions were spaced located at different angles or at different distances, there would be extra energy available. This extra energy would drive the ions toward equal angles and distances from each other, until the extra energy is used up and the ions are arranged into their lowest energy state. That is why minerals form, as a natural way for atoms to arrange themselves into the lowest energy state currently available to them.
Some elements, such as carbon (C) and silicon (Si) have a half-full valence shell. (The valence shell is another name for the outer shell, where the most reactive electrons are.) If an element such as carbon were to gain 4 electrons or lose 4 electrons, it would have a full valence shell. However, it is very difficult for an atom to gain or lose four electrons—the energy barrier becomes too strong. Therefore, carbon and silicon, along with a few other elements, tend to form a different type of bond in which they share their outer electrons with other atoms, which in turn share their outer electrons with the carbon (or silicon) atom. The atoms all end up with a full outer shell of electrons, even though some or all of those electrons are being shared with neighboring atoms. This electron sharing keeps the atoms bonded together. This type of chemical bond is called a covalent bond.
It is not uncommon for covalent bonds to be relatively strong. An extreme example can be in diamond. Diamond is a mineral consisting of nothing but carbon atoms, so its chemical formula is simply C. Each carbon atom in the diamond crystal lattice is covalently bonded to—sharing its valence electrons with—four neighboring carbon atoms. A diamond crystal is held together by nothing but extremely strong covalent bonds in all directions, which makes diamond a very hard mineral, the hardest known.
Gold forms a naturally occurring mineral of more or less pure gold, Au, held together by another type of bond, the metallic bond. Metallic elements such as gold and copper, when they bond with other metallic elements, are sharing some of their electrons not just with adjacent atoms, but throughout the whole substance. That is why metallic substances such as copper, gold, and aluminum make such good electrical conductors, because it is so easy to get the “loose” electrons to respond through the whole extent of the metal.
Another type of chemical bond that occurs in some minerals is the hydrogen bond. Hydrogen bonds are caused by the positive and negative ends of polar molecules attracting each other strongly enough to hold each other in fixed positions. For example, water molecules can join together through hydrogen bonds to form the mineral known as ice. In a water molecule, H2O, each of the hydrogen atoms forms a covalent bond with the oxygen atom.
Check Your Understanding
Answer the question(s) below to see how well you understand the topics covered in the previous section. This short quiz does not count toward your grade in the class, and you can retake it an unlimited number of times.
Use this quiz to check your understanding and decide whether to (1) study the previous section further or (2) move on to the next section.