Chemistry Essentials for Dummies
Chapter 6. Covalent Bonding
Writing Covalent Compound Formulas
You can predict the formula of an ionic compound, based on the loss and gain of electrons to reach a noble gas configuration, as I show you in Chapter 5. (For example, if you react Ca with Cl, you can predict the formula of the resulting salt: CaCl2.) However, you really can’t make that type of prediction for covalent compounds, because they can combine in many ways, and many different possible covalent compounds may result.
Most of the time, you have to know the formula of the molecule you’re studying. But you may have several different types of formulas, and each gives a slightly different amount of information. Oh, joy!
Empirical formulas
REMEMBER. An empirical formula indicates the different types of elements in a molecule and the lowest whole-number ratio of each kind of atom in the molecule. For example, suppose you have a compound with the empirical formula C2H6O. Three different kinds of atoms (C, H, and O) are in the compound, and they’re in the lowest whole-number ratio of two carbons to six hydrogens to one oxygen.
Molecular or true formulas
The molecular formula, or true formula, tells you the kinds of atoms in the compound and the actual number of each atom.
You may determine, for example, that the empirical formula C2H6O is actually the molecular formula, too, meaning that there are actually two carbon atoms, six hydrogen atoms, and one oxygen atom in the compound. However, you may instead find that the molecular formula is C4H12O2, C6H18O3, C8H24O4, or another multiple of 2:6:1.
Structural formulas: Dots and dashes
For ionic compounds, the molecular formula is enough to fully identify a compound, but it’s not enough to identify covalent compounds. Look at Figure 6-4. Both compounds have the molecular formula of C2H6O. That is, both compounds have two carbon atoms, six hydrogen atoms, and one oxygen atom.
However, these are two entirely different compounds with two entirely different sets of properties. The difference is in the way the atoms are bonded, or what’s bonded to what. The compound on the left, dimethyl ether, is used in some refrigeration units and is highly flammable. The one on the right, ethyl alcohol, is the drinking variety of alcohol. Simply knowing the molecular formula isn’t enough to distinguish between the two compounds. Can you imagine going into a restaurant, ordering a shot of C2H6O, and getting dimethyl ether instead of tequila?
Figure 6-4: Two possible compounds of C2H6O.
REMEMBER. Compounds that have the same molecular formula but different structures are called isomers of each other. To identify the exact covalent compound, you need its structural formula.
The structural formula shows the elements in the compound, the exact number of each atom in the compound, and the bonding pattern for the compound. The electron-dot formula and Lewis formula, which I cover in this section, are common structural formulas.
Basic bonds: Writing the electron-dot and Lewis formulas
The following steps explain how to write the electron-dot formula for a simple molecule — water — and provide some general guidelines to follow:
TIP. 1. Write a skeletal structure showing a reasonable bonding pattern using just the element symbols.
Often, most atoms are bonded to a single atom. This atom is called the central atom. Hydrogen and the halogens are very rarely, if ever, central atoms. Carbon, silicon, nitrogen, phosphorus, oxygen, and sulfur are always good candidates, because they form more than one covalent bond in filling their valence energy level. In the case of water, H2O, oxygen is the central element and the hydrogen atoms are both bonded to it. The bonding pattern looks like this:
It doesn’t matter where you put the hydrogen atoms around the oxygen. I put the two hydrogen atoms at a 90° angle to each other.
2. Take all the valence electrons from all the atoms and throw them into an electron pot.
Each hydrogen atom has one electron, and the oxygen atom has six valence electrons (VIA family), so you have eight electrons in your electron pot. Those are the electrons you use when making your bonds and completing each atom’s octet.
3. Use the N - A = S equation to figure the number of covalent bonds in this molecule.
In this equation,
• N equals the sum of the number of valence electrons needed by each atom. N has only two possible values — 2 or 8. If the atom is hydrogen, it’s 2; if it’s anything else, it’s 8.
• A equals the sum of the number of valence electrons available for each atom. A is the number of valence electrons in your electron pot. (If you’re doing the structure of an ion, you add one electron for every unit of negative charge or subtract one electron for every unit of positive charge.)
• S equals the number of electrons shared in the molecule. And if you divide S by 2, you have the number of covalent bonds in the molecule.
So in the case of water,
• N = 8 + 2(2) = 12 (eight valence electrons for the oxygen atom, plus two each for the two hydrogen atoms)
• A = 6 + 2(1) = 8 (six valence electrons for the oxygen atom, plus one for each of the two hydrogen atoms)
• S = 12 - 8 = 4 (four electrons shared in water), and S ÷ 2 = 4 ÷ 2 = 2 bonds
You now know that there are two bonds (two shared pairs of electrons) in water.
4. Distribute the electrons from your electron pot to account for the bonds.
You use four electrons from the eight in the pot, which leaves you with four electrons to distribute later. There has to be at least one bond from your central atom to the atoms surrounding it.
5. Distribute the rest of the electrons (normally in pairs) so that each atom achieves its full octet of electrons.
Remember that hydrogen needs only two electrons to fill its valence energy level. In this case, each hydrogen atom has two electrons, but the oxygen atom has only four electrons, so you place the remaining four electrons around the oxygen. This empties your electron pot. Figure 6-5 shows the completed electron-dot formula for water.
Figure 6-5: Electron-dot formula of H2O.
Notice that this structural formula shows two types of electrons: bonding electrons, the electrons that are shared between two atoms, and nonbonding electrons, the electrons that aren’t being shared. The last four electrons (two electron pairs) that you put around oxygen aren’t being shared, so they’re nonbonding electrons.
TIP. If you want the Lewis formula, all you have to do is substitute a dash for every bonding pair of electrons in the electron-dot formula. Figure 6-6 shows the Lewis formula for water.
Figure 6-6: The Lewis formula for H2O.
Double bonds: Writing structural formulas for C2H4O
Drawing the structural formula for a molecule that contains a double or triple bond can be a bit tricky (see the earlier section “Dealing with multiple bonds”). In those cases, your equations may tell you that you have more covalent bonds that you know what to do with.
For example, here’s an example of a structural formula that’s a little more complicated — C2H4O. The compound has the following framework:
Notice that it has not one but two central atoms — the two carbon atoms. You can put 18 valence electrons into the electron pot: four for each carbon atom, one for each hydrogen atom, and six for the oxygen atom.
Now apply the N - A = S equation:
✓ N = 2(8) + 4(2) + 8 = 32 (two carbon atoms with eight valence electrons each, plus four hydrogen atoms with two valence electrons each, plus an oxygen atom with eight valence electrons)
✓ A = 2(4) + 4(1) + 6 = 18 (four electrons for each of the two carbon atoms, plus one electron for each of the four hydrogen atoms, plus six electrons for the oxygen atom)
✓ S = 32 - 18 = 14, and S ÷ 2 = 14 ÷ 2 = 7 covalent bonds
Put single bonds between the carbon atoms and the hydrogen atoms, between the two carbon atoms, and between the carbon atom and oxygen atom. That’s six of your seven bonds.
There’s only one place that the seventh bond can go, and that’s between the carbon atom and the oxygen atom. It can’t be between a carbon atom and a hydrogen atom, because that would overfill hydrogen’s valence energy level. And it can’t be between the two carbon atoms, because that would give the carbon on the left ten electrons instead of eight. So there must be a double bond between the carbon atom and the oxygen atom. The four remaining electrons in the pot must be distributed around the oxygen atom, because all the other atoms have reached their octet. Figure 6-7 shows the electron- dot formula.
Figure 6-7: Electron-dot formula of C2H4O.
If you convert the bonding pairs to dashes, you have the Lewis formula of C2H4O, as in Figure 6-8.
Figure 6-8: The Lewis formula for C2H4O.
Grouping atoms with the condensed structural formula
I like the Lewis formula because it enables you to show a lot of information without having to write all those little dots. But it, too, is rather bulky. Sometimes chemists (who are, in general, a lazy lot) use condensed structural formulas to show bonding patterns. They may condense the Lewis formula by omitting the nonbonding electrons (dots) and grouping atoms together and/or by omitting certain dashes (covalent bonds). For instance, condensed formulas often group all the hydrogens bonded to a particular carbon atom.
Figure 6-9 shows a couple of condensed formulas for C2H4O.
Figure 6-9: Condensed structural formulas for C2H4O.