Chemical Bonding

Let’s start thinking about chemical bonding by looking at a simple model – Na (atomic # 11) and Cl (atomic # 17) form a fairly simple bond to make salt (NaCl) – the mineral halite. We begin with the nucleus. The atomic number of an element is the number of protons in an atom. This number is always the same for any atom of the element. The other particles in the atom can vary in number (in different ways) but the number of protons does not. For sodium (Na) we therefore put 11 protons in the nucleus. For Cl the number is 17. The number of neutrons in the nucleus is variable among atoms of the element. That is, different atoms may have slightly different numbers of them. This does not affect the chemical properties of the atom – it bonds in the same ways with other atoms regardless of the number of neutrons. The only difference is that its total mass is different from another atom of the same element. Atoms of the same element with different masses (more or fewer neutrons) are called isotopes. Consequently we put ~11 neutrons in the nucleus of Na and ~17 in Cl. Incidentally, the atomic mass of an element given on the periodic table does not refer to the mass of any specific atom. It is the average mass of many atoms, some of which will have slightly different masses from the typical value. No atom weighs what that average mass is. Remember that they all have a whole number mass. The masses of the two atoms above will be ~22 (Na) and ~34 (Cl). Sodium can have mass ranging from 18-37 (7-26 neutrons) but the only stable one has mass=23. The others are radioactive and decay within seconds to a few years. Cl can range in mass from 28-51, but, again, most of these are unstable isotopes and decay very quickly. Masses 35 and 37 are stable.

11P ~11N

17P ~17N

Na (11) Cl (17)

11P ~11N

17P ~17N

That’s the simple story of the nucleus. The atomic number of the element is the number of (or approximate number of) each particle. The biggest error I see on tests is taking the atomic number and dividing by two – half to protons and half to neutrons, or even by 3 – adding electrons into the mix. Remember that the atomic mass of the atom is roughly twice the atomic number and this shouldn’t be a problem.

Na (11) Cl (17)

11P ~11N

17P ~17N

Moving on to the electrons, they go into “shells” around the nucleus in a Bohr model. Each shell has a set maximum number of electrons that can fit into it. (In heavier elements there may be “subshells” as well, but we don’t have to understand those to understand the chemistry of minerals.) The reason we assign a fixed number per shell will be obvious when we begin looking at examples of bonding. The first shell gets two electrons for every element except hydrogen, which only has one electron.

Na (11) Cl (17)

Beyond the first shell we assign eight electrons to each shell. Remember that in an electrically balanced atom the number of electrons is the same as the number of protons – that’s why it’s electrically neutral. But we’ll see when we begin seeing how bonds form that they will involve a certain amount of mobility of the electrons among atoms. At any one time (or even for substantiannl ong times) each atom in a bond may not have all “its” electrons or may have some of “its neighbor’s” electrons. The second shell in both Na and Cl atoms can take a full complement of 8 electrons, bringing the total to 10. Na now needs only one additional electron to bring its e- count into balance with its P+ count. Cl needs an additional 7.

11P ~11N

17P ~17N

Na (11) Cl (17)

11P ~11N

17P ~17N

One more shell, partly filled in both cases, brings the e- count to the right number – 11 for Na and 17 for Cl. Count the electrons and make sure they’re all there.

Na (11) Cl (17)

Counting electrons gets tiresome, particularly when you’re grading dozens of tests and you have to count them over and over. We will therefore use a shorthand version to symbolize the counts, like this: Remember the fill sequence: 2-8-8- … Now we are set to see how these two elements bond to make halite.

11P ~11N

17P ~17N 8e- 2e- 2e- 8e- 1e- 7e-

Na (11) Cl (17)

The e- in the inner shells do not become involved in bonding, only the ones in the outer shell. This is the valence shell and the e- in that shell (that can contribute to a bond) are valence electrons. The charge on the atom that results from the addition of loss of valence e- is the valence state or oxidation state (or simply the “charge”.) When these atoms bond the single e- in the valence shell of Na is transferred to the Cl valence shell. (This is sometimes said to be a “donation” but, in fact, the e- may sometimes return to the Na shell, of one that was originally around the Cl may hop over for a short visit.)

11P ~11N

17P ~17N 8e- 2e- 2e- 8e- 1e- 7e-

Na (11) Cl (17)

There are two important consequences of this transfer. The first is that both outer shells are now “filled”. We will see what that means shortly, but for now take it for granted that this in some important way makes each atom “stable” – that is, the atom is no longer able to enter into any other chemical bond. The other effect is that neither atom has the “correct” number of electrons. (This is why e- number can vary and why the atomic number is not defined as the number of electrons.) Sodium is one e- short, giving it a +1 charge (it has 11P+ and 10e-). Chlorine has one extra, giving it a -1 charge (it has 18e- and 17P+).

11P ~11N

17P ~17N 8e- 2e- 2e- 8e- 8e-

Na (11)

Cl (17)

The opposite charges on the two atoms are what hold them together in this bond. Electrical opposites attract. This is an ionic bond. Ionic bonds are relatively easy to break. You do it whenever you dissolve salt, in water or in saliva. The single e- in the valence shell of the Na is exactly the right number to move to the Cl, filling the outer shells of both (the second shell of the Na is now the outer one), and bonding them together. The Ratio of Na to Cl in the compound is therefore 1:1 – the number of the two atoms is the same. The formula reflects this: NaCl. (Subscript 1’s are simply dropped in chemical formulas.)

11P ~11N

17P ~17N

Na (11)

Cl (17)

8e- 2e- 2e- 8e- 8e-

We’ll symbolize an ionic bond in a couple of ways. If we are thinking of a bond between single atoms the symbolism will be a circle to represent each atom with one or more dots between to represent the number of valence e-. In NaCl there is, of course, 1 valence e-.

We’ll come back and examine the other way to represent ionic bonds, between structural parts of a compound, once we understand what those structural parts are.

Na Cl

1P 8P

~8N

H (1)

O (8)

1e- 2e- 6e-

Let’s look at another familiar example to see a slightly more complex bond: hydrogen (H) and oxygen (O). Hydrogen has an atomic number of 1, so it gets one proton. Remember that a typical hydrogen atom has no neutron, just the proton. It’s single e- shell has only the one e- in it to balance the P+. O has an atomic number of 8, so it gets 8 P and roughly the same number of N. Its first shell gets a pair of e- and the second gets 6 to bring the total to 8, balancing the protons.

1P 8P

~8N

H (1)

O (8)

2e- 7e-

Notice that this compound is different in a couple of ways from NaCl. In the first place, “donating” its e- to the O leaves the H without any electrons – without any shell, at all. This is not good, so we will see how it is managed soon. The other difference is that the e- from the one H in this diagram does not bring the outer shell of O up to 8. It still needs one e- to “fill”.

The solution to the second problem is to combine the O with two H atoms, not just one. This does bring the second shell of O up to 8 e-, making it stable. The not-very-surprising formula for this compound is (drum roll) …… H2O.

This does not solve the difficulty of the H atoms lacking e- shells.

1P

8P ~8N

H (1)

O (8)

2e-

8e-

1P

H (1)

1P

8P ~8N

H (1)

O (8)

2e-

That problem is solved by not transferring the e-’s long-term, but by sharing them equally among all three atoms. Instead of each atom having a dedicated valence shell, a compound valen(t) shell is formed, encircling all three atoms. This type of bond is called a covalent bond. Covalent bonds are much stronger than ionic bonds, and much harder to break. (This is why hydrolysis of pure water takes so long.)

1P

H (1)

8e-

(-) (+)

1P

8P ~8N

H (1)

O (8)

2e-

1P

H (1)

8e-

107°

As an aside, when these atoms bond the H atoms are not directly across from each other. Lines through the centers of the H atoms from the center of the O atom lie at about a 107° angle from each other. Because the O is the anion (-) and the H the cation (+) in this bond, each end of the molecule is slightly charged, in the opposite sense. We call this a polar molecule. This gives water some very interesting properties that we will return to later in the course.

x x

Covalent bonds are symbolized by showing the atoms in the same way as for ionic bonds. The overlap of the circles (representing the valence shells) indicates that the bond is covalent. The number of valence electrons involved in the bond is indicated by the number of dots in the overlap of the valence shells.

If the polar nature of the molecule is important to note, the basic structure can be illustrated like this. Also, some folks use “x” instead of a dot to represent the valence e-.

H H O

H H

O

6P ~6N

C (6)

2e- 4e-

Let’s try another covalently bonded substance (the familiar CO2) for practice and to begin to make a couple of new, related points. Check the atom models and make sure all the particles are placed correctly and that the correct number is present. The oxygen atoms, as always, will need 2 e- to fill their outer shell, so their valence state in the bond will be -2. That leaves the C to be the cation, but notice that 4 is a unique number in this context. It half-fills a shell. Whether the shell is half empty or half full is an interesting thing to ponder. We’ll get back to it very soon. In this atom the C will empty that shell, leaving it with a valence state of +4. Dumping 4 e- is difficult. Dumping one is not too troublesome, 2 is harder, true “donation” of 3 to make an ionic bond is very rare (or probably never happens -- I’m not sure.) With 4 they are never dumped. These always form strong covalent bonds.

8P ~8N

O (8)

2e-

6e-

8P ~8N

O (8)

2e- 6e-

6P ~6N 2e-

8P ~8N 2e- 8P

~8N 2e-

6P ~6N

C (6)

2e- 4e-

The preceding picture suggests that all these ions enter into a single covalent bond, but that is not exactly correct. CO2 differs from water (where there is a single bond) in this respect. If you count the total number of e- in the covalent shell you’ll find the problem: 4 from the single Si atom and 6 each from the two O atoms gives us 4+6+6 = 16 electrons in the shell. Because shells “fill” at the 8 e- level this is obviously not what happens. The diagram shows what happens instead. each O covalently bonds to the C. This is called a double bond. Because H only has one e- it can only form single bonds. O and C, with 6 and 4 valence electrons respectively, can and do form double bonds. We won’t be concerned with double bonds, but I wanted to explain the idea in case someone noticed the odd e- count here.

8P ~8N

O (8)

2e-

6e-

8P ~8N

O (8)

2e- 6e-

6P ~6N 2e-

8P ~8N 2e- 8P

~8N 2e-

8e- 8e-

O O C

Symbolically a double bond is easy to illustrate. It is covalent so the atom’s circles overlap. The number of valence e- is two per bond and there are two bonds shown among the three atoms. Bear in mind that individual e- may not “stick to” individual nuclei. Any one e- is free to move between the covalent shells, and even int the next molecule’s shell if the compound crystallizes. (Crystalline CO2 is “dry ice”.) In covalent bonds there is some loyalty of e- to a single molecule, but also some potential mobility. Soon we’ll see a bond type where that is not the case.

C C C

Carbon can bond to itself in the same way – double covalent bonds among triplets of C atoms, like this:

C C C

C

C

Carbon can also bond covalently to itself with single bonds. As you can probably imagine, this configuration makes a different sort of mineral from the one above.

Carbon bonds to many other atoms with both single and double covalent bonds. It therefore serves as a back-bone molecule for most biologically active molecules (carbohydrates, e.g.), and also most fuel molecules (hydrocarbons, e.g.). We’ll only be interested for now in how it occurs in minerals.

C C C C C C C C C etc. etc.

In the mineral graphite the C atoms are double bonded into sheets. (Remember always that what you see in a drawing is a 2D representation of a 3D structure. The C atoms go back into the picture and out toward the viewer in the same pattern) In the mineral diamond the other bond type is the one present: single bonds among all the atoms in the mineral. This makes diamond very strong in every direction, whereas graphite is only strong “across” the sheet. We’ll see the rest of its structure in the next slide. In addition to having strong bonds among all the atoms in the structure, the C atoms fit together more closely in the diamond structure, making it denser.

Because of the relatively large number of mobile e- in the C sheets, one side of the sheet can, at some random moment, have the majority of its e-, and the other will have fewer. While crystallizing these vaguely charged surfaces attract each other because of their opposite charges. Once this starts the charge differences on the opposite sides of each sheet are perpetuated. The weak electrical attractions are called Van der Waals bonds.

Metallic bonds are found, not surprisingly, in metals. The valence e- are almost entirely unattached to individual nuclei, but are free to roam randomly throughout the structure.It is almost as if there is a single valence shell for all the atoms.

There are two other types of chemical bonds that are of some little interest to geologists because of the properties they lend to certain minerals. These are described below.

Van der Waals Bonds Metallic Bonds

Covalently bonded sheets of carbon

C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C

weak

charge

differences

between

layers

C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C

C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C

C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C

C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C

C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C

C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C C=C=C=C=C=C=C=C=C=C=C=C=C

When you push a pencil point across paper the weak bonds between the sheets break easily, leaving a streak of stacks of C sheets behind as a line. This tendency to break in specific directions is called “cleavage”. We’ll come back to it later.

Metallic bonds are interesting for a couple of reasons. First, remember that the electrons are very mobile and move randomly most of the time, …

but if an electrical current is introduced to them the e- all race toward the + end of that field. Metallic bonds, in other words, give metals their great conductivity.

Secondly, the rather loose nature of the valence e- means that even though the atoms are tightly bonded to each other, as in a covalent bond, their exact locations is not as important as in a covalent bond. They can easily be stretched, bent, or hammered into new shapes. Copper is “drawn” (stretched) into the appropriate shape and diameter for electrical wire, and bending it to fit into a switch or junction box is fairly easy to do. Gold has these same properties. There is not really very much gold in the Georgia Capitol roof in Atlanta. It was hammered out into very thin “gold leaf” and then applied to an underlying structural roof.

Of these bond types we will go on to talk more about two: covalent and ionic. These are the most common bonds by far in minerals and lend the minerals many of their characteristics. We’ll not see much more about Van der Waals and metallic bonds. You should be generally familiar with how they work, but we really only bring them up because of the interesting and useful properties they impart to the minerals (and other compounds) that contain them. They also are blatantly obvious examples of how bonds help control mineral characteristics, which we’ll explore in lab.