Chemistry for Dummies

Part II. Blessed Be the Bonds That Tie

Chapter 7. Covalent Bonds: Let's Share Nicely

In This Chapter

· Seeing how one hydrogen atom bonds to another hydrogen atom

· Defining covalent bonding

· Finding out about the different types of chemical formulas

· Taking a look at polar covalent bonding and electronegativity

· Accepting the unusual properties of water

Sometimes when I’m cooking, I have one of my chemistry nerd moments and start reading the ingredients on food labels. I usually find lots of salts, such as sodium chloride, and lots of other compounds, such as potassium nitrate, that are all ionically bonded (see Chapter 6). But I also find many compounds, such as sugar, that aren’t ionically bonded.

If no ions are holding a compound together, what does hold it together? What holds together sugar, vinegar, and even DNA? In this chapter, I discuss the other major type of bonding: covalent bonding. I explain the basics with an extremely simple covalent compound, hydrogen, and I tell you some cool stuff about one of the most unusual covalent compounds I know — water.

Content Bond Basics

An ionic bond is a chemical bond that comes from the transfer of electrons from a metal to a nonmetal, resulting in the formation of oppositely charged ions — cations (positive charge) and anions (negative charge) — and the attraction between those oppositely charged ions. The driving force in this whole process is achieving a filled valence energy level, completing the atom’s octet. (For a more complete explanation of this concept, see Chapter 6.)

But many other compounds exist in which electron transfer hasn’t occurred. The driving force is still the same: achieving a filled valence energy level. But instead of achieving it by gaining or losing electrons, the atoms in these compounds share electrons. That’s the basis of a covalent bond.

A hydrogen example

Hydrogen is #1 on the periodic table — upper left corner. The hydrogen found in nature is often not comprised of an individual atom. It’s primarily found as H2, a diatomic (two atom) compound. (Taken one step further, because a molecule is a combination of two or more atoms, H2 is called a diatomic molecule,')

Hydrogen has one valence electron. It’d love to gain another electron to fill its Is energy level, which would make it isoelectronic with helium (because the two would have the same electronic configuration), the nearest noble gas. Energy level I can only hold two electrons in the Is orbital, so gaining another electron fills it. That’s the driving force of hydrogen — filling the valence energy level and achieving the same electron arrangement as the nearest noble gas.

Imagine one hydrogen atom transferring its single electron to another hydrogen atom. The hydrogen atom receiving the electron fills its valence shell and reaches stability while becoming an anion (H). However, the other hydrogen atom now has no electrons (H+) and moves further away from stability. This process of electron loss and gain simply won’t happen, because the driving force of both atoms is to fill their valence energy level. So the H2 compound can’t result from the loss or gain of electrons. What can happen is that the two atoms can share their electrons. At the atomic level, this sharing is represented by the electron orbitals (sometimes called electron clouds) overlapping. The two electrons (one from each hydrogen atom) “belong” to both atoms. Each hydrogen atom feels the effect of the two electrons; each has, in a way, filled its valence energy level. A covalent bond is formed — a chemical bond that comes from the sharing of one or more electron pairs between two atoms. The overlapping of the electron orbitals and the sharing of an electron pair is represented in Figure 7-1 (a).

Figure 7-1: The formation of a covalent bond in hydrogen.

Another way to represent this process is through the use of an electron-dot formula. In this type of formula, valence electrons are represented as dots surrounding the atomic symbol, and the shared electrons are shown between the two atoms involved in the covalent bond. The electron-dot formula representations of H2 are shown in Figure 7-1 (b).

Most of the time, I use a slight modification of the electron-dot formula called the Lewis structural formula; it’s basically the same as the electron-dot formula, but the shared pair of electrons (the covalent bond) is represented by a dash. The Lewis structural formula is shown in Figure 7-1 (c). (Check out the section, “Structural formula: Add the bonding pattern,” for more about writing structural formulas of covalent compounds.)

REMEMBER. In addition to hydrogen, six other elements are found in nature in the diatomic form: oxygen (02), nitrogen (N2), fluorine (F2, chlorine (Cl2), bromine (Br2), and iodine (I2). So when I talk about oxygen gas or liquid bromine, I’m talking about the diatomic compound (diatomic molecule).

Here’s one more example of using the electron-dot formula to represent the shared electron pair of a diatomic compound: This time, look at bromine (Br2), which is a member of the halogen family (see Figure 7-2). The two halogen atoms, each with seven valence electrons, share an electron pair and fill their octet.

Figure 7-2: The covalent bond formation of Br2.

Comparing covalent bonds with other bonds

Ionic bonding occurs between a metal and a nonmetal. Covalent bonding, on the other hand, occurs between two nonmetals. The properties of these two types of compounds are different. Ionic compounds are usually solids at room temperature, while covalently bonded compounds can be solids, liquids, or gases. There’s more. Ionic compounds (salts) usually have a much higher melting point than covalent compounds. In addition, ionic compounds tend to be electrolytes, and covalent compounds tend to be nonelectrolytes. (Chapter 6 explains all about ionic bonds, electrolytes, and nonelectrolytes.) I know just what you’re thinking: “If metals react with nonmetals to form ionic bonds, and nonmetals react with other nonmetals to form covalent bonds, do metals react with other metals?” The answer is yes and no.

Metals don’t really react with other metals to form compounds. Instead, the metals combine to form alloys, solutions of one metal in another. But there is such a situation as metallic bonding, and it’s present in both alloys and pure metals. In metallic bonding, the valence electrons of each metal atom are donated to an electron pool, commonly called a sea of electrons, and are shared by all the atoms in the metal. These valence electrons are free to move throughout the sample instead of being tightly bound to an individual metal nucleus. The ability of the valence electrons to flow throughout the entire metal sample is why metals tend to be conductors of electricity and heat.

Understanding multiple bonds

I define covalent bonding as the sharing of one or more electron pairs. In hydrogen and the other diatomic molecules, only one electron pair is shared. But in many covalent bonding situations, more than one electron pair is shared. This section shows you an example of a molecule in which more than one electron pair is shared.

Nitrogen (N2) is a diatomic molecule in the VA family on the periodic table, meaning that it has five valence electrons (see Chapter 4 for a discussion of families on the periodic table). So nitrogen needs three more valence electrons to complete its octet.

A nitrogen atom can fill its octet by sharing three electrons with another nitrogen atom, forming three covalent bonds, a so-called triple bond. The triple bond formation of nitrogen is shown in Figure 7-3.

A triple bond isn’t quite three times as strong as a single bond, but it’s a very strong bond. In fact, the triple bond in nitrogen is one of the strongest bonds known. This strong bond is what makes nitrogen very stable and resistant to reaction with other chemicals. It’s also why many explosive compounds (such as TNT and ammonium nitrate) contain nitrogen. When these compounds break apart in a chemical reaction, nitrogen gas (N2) is formed, and a large amount of energy is released.

Figure 7-3: Triple bond formation in N2.

There are no salt molecules!

A molecule is a compound that is covalently bonded. It's technically incorrect to refer to sodium chloride, which has ionic bonds, as a molecule, but lots of chemists do it anyway. It's kind of like using the wrong fork at a formal dinner. Some people may notice, but most don't notice or don't care. But just so you know, the correct term for ionic compounds is formula unit.

Carbon dioxide (CO2) is another example of a compound containing a multiple bond. Carbon can react with oxygen to form carbon dioxide. Carbon has four valence electrons, and oxygen has six. Carbon can share two of its valence electrons with each of the two oxygen atoms, forming two double bonds. These double bonds are shown in Figure 7-4.

Figure 7-4: Formation of carbon dioxide.

Naming Binary Covalent Compounds

Binary compounds are compounds made up of only two elements, such as carbon dioxide (CO2). Prefixes are used in the names of binary compounds to indicate the number of atoms of each nonmetal present. Table 7-1 lists the most common prefixes for binary covalent compounds.

Table 7-1. Common Prefixes for Binary Covalent Compounds

Number of Atoms






















In general, the prefix mono- is rarely used. Carbon monoxide is one of the few compounds that uses it.

Take a look at the following examples to see how to use the prefixes when naming binary covalent compounds (I’ve bolded the prefixes for you):

CO2 carbon dioxide

P4O10 tetraphosphorus decoxide (Chemists try to avoid putting an a and an o together with the oxide name, as in decaoxide, so they normally drop the a off the prefix.)

SO3 sulfur trioxide

N2O4 dinitrogen tetroxide

This naming system is used only with binary, nonmetal compounds, with one exception — MnO2 is commonly called manganese dioxide.

So Many Formulas, So Little Time

In Chapter 6,1 show you how to predict the formula of an ionic compound, based on the loss and gain of electrons, to reach a noble gas configuration. (For example, if you react Ca with Cl, you can predict the formula of the resulting salt — CaCl2.) You really can’t make that type of prediction with 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 formula: Just the elements

The 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 that you have a compound with the empirical formula C2H6O. Three different kinds of atoms are in the compound, C, H, and O, and they’re in the lowest whole-number ratio of 2 C to 6 H to 1 O. So the actual formula (called the molecular formula or true formula) may be C2H6O, C4H12O2, C6H18O3, C8H24O4, or another multiple of 2:6:1.

Molecular or true formula: Inside the numbers

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.

For ionic compounds, this formula is enough to fully identify the compound, but it’s not enough to identify covalent compounds. Look at the Lewis formulas presented in Figure 7-5. Both compounds have the molecular formula of C2H6O.

Figure 7-5: Two possible compounds of C2H6O.

Both compounds in Figure 7-5 have two carbon atoms, six hydrogen atoms, and one oxygen atom. The difference is in the way the atoms are bonded, or what’s bonded to what. These are two entirely different compounds with two entirely different sets of properties. The one on the left is called dimethyl ether. This compound is used in some refrigeration units and is highly flammable. The one on the right is ethyl alcohol, 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 and ordering a shot of C2H6O and getting dimethyl ether instead of tequila?

It's always important to KISS

A lot of molecules obey the octet rule: Each atom in the compound ends up with a full octet of eight electrons filling its valence energy level. However, tike most rules, the octet rule does have exceptions. Some stable molecules have atoms with just 6 electrons, and some have 10 or 12.1 point out some examples of compounds that don't obey the octet rule in the section, "What Does Water Really Look Like? The VSEPR Theory,” later in this chapter, but for the most part in this book, I concentrate on situations in which the octet rule is obeyed.

I pretty much stick to the KISS principle—Keep It Simple, Silly. Electron-dot formulas are used quite a bit by organic chemists in explaining Why certain compounds react the way they do and are the first step in determining the molecular geometry of a compound.

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.

Structural formula: Add the bonding pattern

To write a formula that stands for the exact compound you have in mind, you often must write the structural formula instead of the molecular 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 are examples of structural formulas.

Writing the electron-dot formula for water

The following steps explain how to write the electron-dot formula for a simple molecule — water — and provide some general guidelines and rules to follow:

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, centred 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. In the section, “What Does Water Really Look Like? The VSEPR Theory,” later in this chapter, you see why I put the hydrogen atoms at a 90- degree angle to each other, but it really doesn’t matter when writing electron-dot (or Lewis) formulas.

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 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. If you’re doing the structure of an ion, you add one electron for every unit of negative charge if it’s an anion or subtract one electron for every unit of positive charge if it’s a cation. A is the number of valence electrons in your electron pot.

S equals the number of electrons shared in the molecule. And if you divide 5 by 2, you have the number of covalent bonds in the molecule.

So in the case of water,

N = 8 + 2(2) = 12 (8 valence electrons for the oxygen atom, plus 2 each for the two hydrogen atoms)

A = 6 + 2(1) = 8 (6 valence electrons for the oxygen atom, plus 1 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 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 the remaining four electrons are placed around the oxygen. This empties your electron pot. The completed electron-dot formula for water is shown in Figure 7-6.

Figure 7-6: Electron-dot formula of H2O.

Notice that there are actually two types of electrons shown in this structural formula: bonding electrons, the electrons that are shared between two atoms, and nonbonding electrons, the electrons that are not being shared. The last four electrons (two electron pairs) that you put around oxygen are not being shared, so they’re nonbonding electrons.

Writing the Lewis formula for water

If you want the Lewis formula for water, all you have to do is substitute a dash for every bonding pair of electrons. This structural formula is shown in Figure 7-7.

Figure 7-7: The Lewis formula for H2O.

Writing the Lewis formula for C2H4O.

Here’s an example of a Lewis 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 (2 carbon atoms with 8 valence electrons each, plus 4 hydrogen atoms with 2 valence electrons each, plus an oxygen atom with 8 electrons)

A = 2(4) + 4(1) + 6 = 18 (4 electrons for each of the two carbon atoms, plus 1 electron for each of the 4 hydrogen atoms, plus 6 valence electrons for the oxygen atom)

S = 32 - 18 = 14, and S/2 = 14/2 = 7 bonds

Add single bonds between the carbon atoms and the hydrogen atom, 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. The electron-dot formula is shown in Figure 7-8.

Figure 7-8: Electron-dot formula of C2H4O.

If you convert the bonding pairs to dashes, you have the Lewis formula of C2H4O, as shown in Figure 7-9.

Figure 7-9: The Lewis formula for C2H4O.

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 and grouping atoms together and/or by omitting certain dashes (covalent bonds). A couple of condensed formulas for C2H4O are shown in Figure 7-10.

Figure 7-10: Condensed structural formulas for C2H4O.

Some Atoms Are More Attractive Than Others

When a chlorine atom covalently bonds to another chlorine atom, the shared electron pair is shared equally. The electron density that comprises the covalent bond is located halfway between the two atoms. Each atom attracts the two bonding electrons equally. But what happens when the two atoms involved in a bond aren’t the same? The two positively charged nuclei have different attractive forces; they “pull” on the electron pair to different degrees. The end result is that the electron pair is shifted toward one atom. But the question is, “Which atom does the electron pair shift toward?” Electronegativities provide the answer.

Attracting electrons: Electronegativities

Electronegativity is the strength an atom has to attract a bonding pair of electrons to itself. The larger the value of the electronegativity, the greater the atom’s strength to attract a bonding pair of electrons. Figure 7-11 shows the. electronegativity values of the various elements below each element symbol on the periodic table. Notice that, with a few exceptions, the electronegativities increase, from left to right, in a period, and decrease, from top to bottom, in a family.

Electronegativities are useful because they give information about what will happen to the bonding pair of electrons when two atoms bond. For example, look at the Cl2 molecule. Chlorine has an electronegativity value of 3.0, as shown in Figure 7-11. Each chlorine atom attracts the bonding electrons with a force of 3.0. Because there’s an equal attraction, the bonding electron pair is shared equally between the two chlorine atoms and is located halfway between the two atoms. A bond in which the electron pair is equally shared is called a nonpolar covalent bond. You have a nonpolar covalent bond anytime the two atoms involved in the bond are the same or anytime the difference in the electronegativities of the atoms involved in the bond is very small.

Now consider hydrogen chloride (HCl). Hydrogen has an electronegativity of 2.1, and chlorine has an electronegativity of 3.0. The electron pair that is bonding HCl together shifts toward the chlorine atom because it has a larger electronegativity value. A bond in which the electron pair is shifted toward one atom is called a polar covalent bond. The atom that more strongly attracts the bonding electron pair is slightly more negative, while the other atom is slightly more positive. The larger the difference in the electronegativities, the more negative and positive the atoms become.

Now look at a case in which the two atoms have extremely different electronegativities — sodium chloride (NaCl). Sodium chloride is ionically bonded (see Chapter 6 for information on ionic bonds). An electron has transferred from sodium to chlorine. Sodium has an electronegativity of 1.0, and chlorine has an electronegativity of 3.0. That’s an electronegativity difference of 2.0 (3.0 - 1.0), making the bond between the two atoms very, very polar. In fact, the electronegativity difference provides another way of predicting the kind of bond that will form between two elements.

Figure 7-11: Electronegativities of the elements.


Electronegativity Difference

Type of Bond Formed

0.0 to 0.2

nonpolar covalent

0.3 to 1.4

polar covalent

> 1.5


The presence of a polar covalent bond in a molecule can have some pretty dramatic effects on the properties of a molecule.

Polar covalent bonding

If the two atoms involved in the covalent bond are not the same, the bonding pair of electrons are pulled toward one atom, with that atom taking on a slight (partial) negative charge and the other atom taking on a partial positive charge. In most cases, the molecule has a positive end and a negative end, called a dipole (think of a magnet). Figure 7-12 shows a couple of examples of molecules in which dipoles have formed. (The little Greek symbol by the charges refers to a partial charge.)

Figure 7-12: Polar covalent bonding in HF and NH3.

In hydrogen fluoride (HF), the bonding electron pair is pulled much closer to the fluorine atom than to the hydrogen atom, so the fluorine end becomes partially negatively charged and the hydrogen end becomes partially positively charged. The same thing takes place in ammonia (NH3); the nitrogen has a greater electronegativity than hydrogen, so the bonding pairs of electrons are more attracted to it than to the hydrogen atoms. The nitrogen atom takes on a partial negative charge, and the hydrogen atoms take on a partial positive charge.

The presence of a polar covalent bond explains why some substances act the way they do in a chemical reaction: Because this type of molecule has a positive end and a negative end, it can attract the part of another molecule with the opposite charge.

In addition, this type of molecule can act as a weak electrolyte because a polar covalent bond allows the substance to act as a conductor. So if a chemist wants a material to act as a good insulator (a device used to separate conductors), the chemist would look for a material with as weak a polar covalent bond as possible.

Water: A realty strange molecule

Water (H2O) has some very strange chemical and physical properties. It can exist in all three states of matter at the same time. Imagine that you’re sitting in your hot tub (filled with liquid water) watching the steam (gas) rise from the surface as you enjoy a cold drink from a glass filled with ice (solid) cubes. Very few other chemical substances can exist in all these physical states in this close of a temperature range.

And those ice cubes are floating! In the solid state, the particles of matter are usually much closer together than they are in the liquid state. So if you put a solid into its corresponding liquid, it sinks. But this is not true of water. Its solid state is less dense than its liquid state, so it floats. Imagine what would happen if ice sank. In the winter, the lakes would freeze, and the ice would sink to the bottom, exposing more water. The extra exposed water would then freeze and sink, and so on, until the entire lake was frozen solid. This would destroy the aquatic life in the lake in no time. So instead, the ice floats and insulates the water underneath it, protecting aquatic life. And water’s boiling point is unusually high. Other compounds similar in weight to water have a much lower boiling point.

Another unique property of water is its ability to dissolve a large variety of chemical substances. It dissolves salts and other ionic compounds, as well as polar covalent compounds such as alcohols and organic acids. In fact, water is sometimes called the universal solvent because it can dissolve so many things. It can also absorb a large amount of heat, which allows large bodies of water to help moderate the temperature on earth.

Water has many unusual properties because of its polar covalent bonds. Oxygen has a larger electronegativity than hydrogen, so the electron pairs are pulled in closer to the oxygen atom, giving it a partial negative charge. Subsequently, both of the hydrogen atoms take on a partial positive charge. The partial charges on the atoms created by the polar covalent bonds in water are shown in Figure 7-13.

Figure 7-13: Polar covalent bonding in water.

Water is a dipole and acts like a magnet, with the oxygen end having a negative charge and the hydrogen end having a positive charge. These charged ends can attract other water molecules. The partially negatively charged oxygen atom of one water molecule can attract the partially positively charged hydrogen atom of another water molecule. This attraction between the molecules occurs frequently and is a type of intermolecular force (force between different molecules).

Intermolecular forces can be of three different types. The first type is called a London force or dispersion force. This very weak type of attraction generally occurs between nonpolar covalent molecules, such as nitrogen (N2), hydrogen (H2), or methane (CH4). It results from the ebb and flow of the electron orbitals, giving a very weak and very brief charge separation around the bond.

The second type of intermolecular force is called a dipole-dipole interaction. This intermolecular force occurs when the positive end of one dipole molecule is attracted to the negative end of another dipole molecule. It’s much stronger than a London force, but it’s still pretty weak.

The third type of interaction is really just an extremely strong dipole-dipole interaction that occurs when a hydrogen atom is bonded to one of three extremely electronegative elements — 0, N, or F. These three elements have a very strong attraction for the bonding pair of electrons, so the atoms involved in the bond take on a large amount of partial charge. This bond turns out to be highly polar — and the higher the polarity, the more effective the bond. When the O, N, or F on one molecule attracts the hydrogen of another molecule, the dipole-dipole interaction is very strong. This strong interaction (only about 5 percent of the strength of an ordinary covalent bond but still very strong for an intermolecular force) is called a hydrogen bond. The hydrogen bond is the type of interaction that’s present in water (see Figure 7-14).

Figure 7-14: Hydrogen bonding in water.

Water molecules are stabilized by these hydrogen bonds, so breaking up (separating) the molecules is very hard. The hydrogen bonds account for water’s high boiling point and ability to absorb heat. When water freezes, the hydrogen bonds lock water into an open framework that includes a lot of empty space. In liquid water, the molecules can get a little closer to each other, but when the solid forms, the hydrogen bonds result in a structure that contains large holes. The holes increase the volume and decrease the density. This process explains why the density of ice is less than that of liquid water (the reason ice floats). The structure of ice is shown in Figure 7-15, with the hydrogen bond indicated by dotted lines.

Figure 7-15: The structure of ice.

What Does Water Really Look Like? The VSEPR Theory

The molecular geometry of a molecule, how the atoms are arranged in threedimensional space, is important for chemists to know because it often explains why certain reactions will or won’t occur. In the area of medicine, for example, the molecular geometry of a drug may lead to side reactions. Molecular geometry also explains why water is a dipole (a molecule with a positive end and a negative end, like a magnet) and carbon dioxide is not.

The VSEPR (Valence Shell Electron-Pair Repulsion) theory allows chemists to predict the molecular geometry of molecules. The VSEPR theory assumes that the electron pairs around an atom, whether they’re bonding (shared between two atoms) or nonbonding (not shared), will try to get as far apart from each other in space to minimize the repulsion between themselves. It’s like going to a fancy party and seeing someone else wearing the exact same outfit. You’re gonna try to stay as far away from that person as possible!

Electron-pair geometry is the arrangement of the electron pairs, bonding and nonbonding, around a central atom. After you determine the electron-pair geometry, you can imagine the nonbonding electrons being invisible and see what’s left. What’s left is what I call the molecular geometry, or shape, the arrangement of the other atoms around a central atom.

To determine the molecular geometry, or shape, of a molecule using the VSEPR theory, follow these steps:

1. Determine the Lewis formula (see “Covalent Bond Basics,” earlier in the chapter) of the molecule.

2. Determine the total number of electron pairs around the central atom.

3. Using Table 7-2, determine the electron-pair geometry. (Table 7-2 relates the number of bonding and nonbonding electron pairs to the electron-pair geometry and molecular shape.)

4. Imagine that the nonbonding electron pairs are Invisible, and use Table 7-2 to determine the molecular shape.

Table 7-2. Predicting Molecular Shape with the VSEPR Theory

Total Number of Electron Pairs

Number of Bonding Pairs

Electron-pair Geometry

Molecular Geometry







trigonal planar

trigonal planar



trigonal planar

bent, V-shaped



trigonal planar









trigonal pyramidal




bent, V-shaped



trigonal bipyramidal

trigonal bipyramidal



trigonal bipyramidal




trigonal bipyramidal




trigonal bipyramidal









square pyramidal




square planar

Even though you normally don’t have to worry about more than four electron pairs around the central atom (octet rule), I put some of the less common exceptions to the octet rule in Table 7-2. Figure 7-16 shows some of the more common shapes mentioned in the table.

To determine the shapes of water (H2O) and ammonia (NH3), the first thing you have to do is determine the Lewis formula for each compound. Follow the rules outlined in the section, “Structural formula: Add the bonding pattern” (the N - A = S rules), and write the Lewis formulas as shown in Figure 7-17.

Figure 7-16: Common molecular shapes.

Figure 7-17: Lewis formulas for H2O and NH3.

For water, there are four electron pairs around the oxygen atom, so the electron-pair geometry is tetrahedral. Only two of these four electron pairs are involved in bonding, so the molecular shape is bent or V-shaped. Because the molecular shape for water is V-shaped, I always show water with the hydrogen atoms at about a 90-degree angle to each other — it’s a good approximation of the actual shape.

Ammonia also has four electron pairs around the nitrogen central atom, so its electron-pair geometry is tetrahedral, as well. Only one of the four electron pairs is nonbonding, however, so its molecular shape is trigonal pyramidal. This shape is like a three-legged milk stool, with the nitrogen being the seat — the lone pair of nonbonding electrons would then stick straight up from the seat. You’d get a surprise if you sat on an ammonia stool!