Lab 2: Minerals I -- Bonding


OK, so now you understand the basics of chemistry regarding protons and neutrons. Now, lets tackle electrons and how their arrangement determines whether or not two atoms will form a chemical bond to form a compound. Minerals are just compounds, so we will be learning about how minerals are built from their constituent atoms.


Carbon Atom


First, recall the illustration of carbon from the last section, repeated above. The nucleus has protons and neutrons, and each proton has a positive charge. For each proton, there is an electron somewhere within the structured cloud of electrons surrounding the nucleus. Protons and electrons cancel out each other's charge, such that, in a typical atom with equal numbers of protons and electrons, the atom itself has no charge, overall. Let's state that point explicitly:


In an atom with equal numbers of protons and electrons, which is the condition for atoms in their unbonded state (their ground state), the atom does not have a charge.

And, as you might guess, there is a companion statement to that one:


If an atom has unequal numbers of protons and electrons, the atom has a charge, either positive (# protons > # electrons) or negative (# electrons > # protons), and is called an ion.

This brings us to the simplest type of chemical bonding, ionic bonding, wherein atoms of opposite charge are attracted to one another and "stick together."


Ionic Bonding


How can some atoms have unequal numbers of protons and electrons? By loss or gain of one or more electrons.


Why would an atom loose or gain an electron? Because the structure of that structured cloud of electrons has a few special properties regarding charge and "stability."



Let's look at the periodic table again, this time scanning the elements for those which bond by ionic bonding, and for those which don't form bonds at all. It is always a good idea to start with something familiar when you can, in this case, good ol' table salt, NaCl (sodium chloride, going by the names salt, rock salt, and halite, for the proper mineral name). Sodium (Na) forms an ionic bond with chlorine (Cl). Why? When we examine the structured clouds of electrons around the nuclei of sodium and chlorine atoms, we see the answer. First, sodium:




Sodium has 11 protons, and therefore 11 electrons. The structure of the electron configuration works like this:


  1. Just around the nucleus, in a tighter fit than for electrons lying farther out, there can be only as many as two, count 'em, two (2) electrons. If we think of the structure of the electron cloud as a set of nested regions, or shells, we can say that two electrons can lie in the inner shell. There are nine more electrons in a sodium atom.
  2. The next region, or second shell, in the structure of the electron cloud can hold eight, count 'em, eight (8) electrons. There is one more electron in a sodium atom.
  3. The remaining electron sits out by its lonesome (1) in the next region, or third shell.




Chlorine has 17 protons, and therefore 17 electrons. Chlorine's electron structure works like this:


  1. The inner shell can only hold as many as two (2) electrons. Remember? So, there are 15 more electrons in a chlorine atom.
  2. The second shell can hold as many as eight (8) electrons. Remember? So, there are 7 more electrons.
  3. The remaining seven (7) electrons can fit in the third shell, which can hold as many as 18 electrons.


Look up at the two sections we just went over for sodium and chlorine, and scan the text for the numbers in parentheses. We have the electron structure for sodium as (2) (8) (1) and for chlorine we have it as (2) (8) (7). There is something telltale about this pattern of numbers that explains a whole lot about why certain elements form bonds with others. Recognizing the pattern leads to something called the octet rule, which states that atoms will tend to "seek" a count of eight (8) electrons in their outermost electron shell (this is called the valence shell). In the case of ionic bonding, they do this by either losing or gaining electrons. Notice that sodium (2) (8) (1) just has to loose that lonesome electron in the third shell to become sodium (2) (8). And chlorine (2) (8) (7) just needs to gain one electron to become chlorine (2) (8) (8). But when they each do this (loose or gain an electron), they take on an overall charge, as follows:


Sodium (2) (8) still has 11 protons, but with the loss of that lonesome electron, it only has 10 electrons, so now the atom has a +1 charge.


Chlorine (2) (8) (8) still has 17 protons, but with the gain of the electron, it now has 18 electrons, so now the atom has a -1 charge.


And that is why sodium and chlorine form an ionic bond. Together, when they come into close contact, the drive to satisfy the octet rule leads to the simultaneous loss for sodium and gain for chlorine of one electron, such that they each take on a charge, and these charges are opposite, and so "click," they become bonded together. Opposites do, in fact, attract (well, in this case, they do).


If you want an illustration of the NaCl ionic bond, even though illustrations of atoms can be deceiving, here 'tis:


Sodium-Chlorine ionic bond


Other pairs of elements will undergo ionic bonding in a similar fashion: LiF (lithium fluoride), KCl (potassium chloride), etc. We can see why, if we look at the number of electrons in the outermost electron shell (the valence shell) of atoms of the first twenty elements. Confirm the number of valence electrons for sodium and chlorine:



The table above helps to visualize the octet rule, which we now understand for the case of sodium and chlorine, in the growth of the mineral halite. We can learn more from this table. Scan the rows and check the number of electrons in the outer electron shell. Are there 8 electrons? If so, the octet rule is satisfied straightaway, and the element doesn't even need to form bonds with other elements - an element with eight electrons in its outermost (valence) shell, is inert, or non-reactive (doesn't form bonds). We call these the noble gases.


Noble Gases


Let's repeat the table with the noble gases marked (and show one exception to the octet rule, that should make sense to you):



What is it about helium, neon, and argon that make them noble gases (non-reactive)? Take a look at the number of electrons in the valence shell. Helium has 2, neon has 8, and argon has 8. For neon and argon, the octet rule has been satisfied, and there is no need for atoms of these elements to bond with atoms of other elements. What about helium? Well, recall that there are only up to two (2) electrons possible in the first shell, and this shell, being filled in helium, takes on a "satisfied" state when full. It turns out that, in a similar way for heavier elements, eight is a "happy number" for the number of electrons in the outermost shell, regarding whether or not an atom will bond or not, even if the outermost shell will hold more than eight electrons. You just have to remember the special case of 2 electrons in the outer shell of helium also fitting the bill. So, you understand the case of two for helium, what is this business about eight for heavier elements?


Ionic Bonding in Common Salt


Let's return to common salt, NaCl, the mineral halite. We learned that sodium, with its lone valence electron, gives it up in an ionic bond with chlorine, and afterwards has eight (8) electrons in its outer shell. And, chlorine, with its seven (7) valence electrons, can take on the electron, and afterwards has eight (8) electrons in its valence shell. When this ionic bond forms, both sodium and chlorine have eight valence electrons. Observe:



Look at sodium and chlorine in the table and appreciate the statement that after forming an ionic bond these elements have "taken on a noble gas configuration." (Compare to He, Ne, Ar).


So, where does this leave us, with regard to understanding why and how atoms bond with one another? We now understand why ionic bonds happen between Na and Cl, or between Li and F, or between K and Cl, but what about the other elements in that list and other combinations (and that's just the first 20 elements)?


There is another important type of chemical bond, called the covalent bond.


Covalent Bonding


A covalent bond is formed when two atoms share electrons. This type of bond is also important in minerals, and is considerably stronger than the ionic bond. Hard minerals such as quartz contain atoms covalently bonded. The name includes a reference to the fact that the shared electrons are in the valence shell of the atoms, covalent, and the prefix co- indicates the sharing relationship.


Several elements which occur in nature as gases, such as oxygen, hydrogen, and chlorine, exist as diatomic molecules, in which the two atoms are covalently bonded. Observe atom illustrations for these pairings and you'll readily see the sharing relationship for the electrons involved:


Hydrogen Gas Molecule


Oxygen Gas Molecule


Chlorine Gas Molecule


In all three, you see the shared electrons in the valence shell. In the case of hydrogen gas, H 2 , and chlorine gas, Cl 2 , there is a single valence electron (lonesome electron) in each atom before the bond. When the bond forms, those lonesome valence electrons are shared. Hydrogen gas attains the noble gas configuration of helium, because there are now two valence electrons (Recall that for the inner shell, if there are two electrons, then a parallel rule to the octet rule is satisfied). Chlorine gas attains the noble gas configuration of argon, as it now has eight valence electrons (octet rule is satisfied). For oxygen, there are two shared pairs of electrons, but the result satisfies the octet rule, as each oxygen atom attains a noble gas configuration.


Those are gases. We need to look at covalent bonds involved in minerals. The most important are the covalent bonds that form between silicon (atomic number 14) and oxygen (atomic number 8) in the silicate minerals, the most common minerals on Earth. First, do the simple analysis of the electron shell filling from the inside out for each element to appreciate the valence electrons:


Silicon , with its 14 electrons, has the shell configuration (2) (8) (4). There are four valence electrons. Silicon needs to loose these four or to gain four, or to share four.


Oxygen , with 8 electrons, has the shell configuration (2) (6). There are six valence electrons. With two electrons shared with another atom, the octet rule would be satisfied for oxygen.


Notice that if you consider a bond between a silicon atom (4 valence electrons) and an oxygen atom (6 valence electrons), you can't find a combination of shared pairs of electrons that would satisfy the octet rule, where both atoms after the bond has formed have eight valence electrons. And, if you try various combinations of one or more silicon atoms with one or more oxygen atoms, you won't find a solution. So, it is best for us to learn at this point what does, in fact, happen with these two elements. Four oxygen atoms bond to one silicon atom, producing an SiO 4 molecule. Covalent bonding like this can occur, without the octet rule being satisfied. The consequence is that the resulting molecule has either a surplus of electrons or a deficit of electrons, and thus has an overall charge. For SiO 4 , the charge is -4, because of the 4 valence electrons. The -4 charge is shown when we write the chemical formula, which is SiO 4 4- .


From that last paragraph, you hopefully can see that atoms can bond with one another without satisfying the octet rule, yet form stable molecules that act as "units" of chemistry, just like individual atoms do. They also have a size, just like individual atoms do. And, they have a charge (in this case -4) that results from either a surplus or deficit of electrons, per the octet rule. A molecule like this will "behave as a unit," just like an atom does, and will form bonds with atoms or other molecules by the same tendency to fulfill the octet rule.


The SiO 4 molecule is extremely important, in this "behaving as a unit" way, because it forms the basic building block for so many minerals. We haven't yet spoken about the three dimensional structure of molecules, but it is worth mentioning here, a bit ahead of the game, that the Si atom sits between the four oxygen atoms to form a pyramidal structure called the silicon-oxygen tetrahedron, tetra for four, and hedron for the four sides of the pyramidal structure. There are other important "building block" molecules in non-silicate mineral groups.


Other Types of Bonds


Hybrid Bonds


Atoms of several elements my bond together by a combination of ionic and covalent bonds. This is especially important in the silicate minerals, where silicon-oxygen tetrahedra, with their internal covalent bonds, are in turn bonded to other elements by ionic bonds.


Metallic Bonds


In metals such as copper, the valence electrons are, in a sense, free to move around within the metal without any attachment to a particular atom. This explains why metals such as copper are such conductors of electricity: electricity is the flow of electrons, which is happening as a normal matter of course in metals. If we find copper on the periodic table, we see that it has an atomic number of 29. The 29 electrons in copper are arranged as (2) (8) (18) (1). That lonesome electron is the one that freely moves between atoms of copper in a mass of the metal. These metallic bonds are not strong, which you see when you easily bend a copper wire.