In This Chapter

► Giving and receiving electrons in ionic bonding

► Sharing electrons in covalent bonding

► Understanding molecular orbitals

► Shaping up molecules with VSEPR theory and hybridization

► Tugging at the idea of polarity

Any atoms are prone to public displays of affection, pressing themselves against WWW Other atoms in an intimate electronic embrace called Bonding. Atoms bond with one another by playing various games with their valence electrons. In this chapter, we describe the basic rules of those games.

Because valence electrons are so important to bonding, problems involving bonding sometimes make use of Electron dot structures, Symbols that represent valence electrons as dots surrounding an atom’s chemical symbol. You should be able to draw and interpret electron dot structures for atoms as shown in Figure 5-1. This figure shows the electron dot structures for elements in the periodic table’s first two rows; notice that the valence shells progressively fill moving from left to right. To determine the electron dot structure of any element, count the number of electrons in that element’s valence shell. Then draw that number of dots around the chemical symbol for the element. Chapter 4 describes some of the factors that determine whether atoms gain or lose electrons to form ions. You should make sure to understand those patterns before attacking this chapter.

Figure 5-1:

Electron dot structures

For elements in the first

Two rows of

The periodic table.

IA IIA IIIA IVA VA VIA VIIA VIIIA H – He:

■ • • •

Li • »Be» «B« »C» SIMS JO! IF! INeJ

Pairing Charges with Ionic Bonds

Atoms of some elements, like metals, can easily lose valence electrons to form Cations (atoms with positive charge) that have stable electron configurations. Atoms of other elements, like the halogens, can easily gain valence electrons to form Anions (atoms with negative charge) with stable electron configurations. Cations and anions experience Electrostatic attraction To one another because opposite charges attract. So, a cation will snuggle up to an anion, given the chance. This event is called Ionic bonding, And it happens because the energy of the ioni-cally bonded ions is lower than the energy of the ions when they are separated.

You can think of an ionic bond as resulting from the transfer of an electron from one atom to another, as shown in Figure 5-2 for sodium and chlorine. Metals (like sodium) tend to give up their electrons to nonmetals (like chlorine) because nonmetals are much more Electronegative (they more strongly attract electrons within a bond to themselves). The greater the difference in electronegativity between the two ions, the more Ionic (or completely uneven in sharing of electrons) is the bond that forms between them.

Figure 5-2:

The transfer of an electron from sodium to

Chlorine to Na – ^~\*Cl:-Na+ IClf

Form an •• ••

Ionic bond between the Na+ cation and the Cl-anion.

Although ions are often individual, charged atoms, there are also many examples of Polyatomic ions (charged particles made up of more than one atom). Examples of common polyatomic ions are ammonium, NH4+, and sulfate, SO42-. We cover polyatomic ions in detail in Chapter 6.

When cations and anions associate in ionic bonds, they form Ionic compounds. At room temperature, most ionic compounds are solid because of the strong electrostatic forces that hold together the ions within them. The ions in ionic solids tend to pack together in a Lattice, A highly organized, regular arrangement that allows for the maximum electrostatic interaction between anions and cations. The geometric details of the packing can differ among different ionic compounds, but a simple lattice structure is shown in Figure 5-3. Flip to Chapter 6 for full details on ionic compounds.

The strong electrostatic forces that hold together ionic lattices result in the high melting and boiling points that are common among ionic compounds (see Chapter 10 for general

Information on melting and boiling points). Although it may take a great deal of thermal energy to disrupt ionic bonds, ionic compounds are usually easily dissolved in water or in other Polar solvents (fluids made up of molecules that have unevenly distributed charge). When the solvent molecules are polar, they can engage in favorable interactions with the ions that help to compensate for disrupting the ionic bonds. For example, polar water molecules can interact well with both sodium cations (Na+) and chlorine anions (Cl-). Water molecules are polar because they have distinct and separate bits of positive and negative charge. Water molecules can orient their positive bits toward Cl – and their negative bits toward Na+. Positive charges attract negative charges and vice versa, so these kinds of interactions are favorable — they require less energy. So, water dissolves solid NaCl quite well because the water-ion interactions can compete with the (Na+)-(Cl-) interactions.

Figure 5-3:

The lattice structure of an ionic solid, sodium chloride.

Cl ‘

Q.

A.

When ionic compounds are melted or dissolved, so the individual ions can move about, the resulting liquid is a very good conductor of electricity. Ionic solids, however, are often poor conductors of electricity.

Salts Are a common variety of ionic compound. A salt is formed from the reaction between a base and an acid. For example, hydrochloric acid reacts with sodium hydroxide to form the salt sodium chloride and water:

HCl(a<7) + NaOH(ag) — NaCl(aq) + H2O(l)

Note that Aq Indicates that the substance is dissolved in water, in an Aqueous Solution.

Why do metals tend to form ionic compounds with nonmetals?

Metals are much less electronegative than nonmetals, meaning that they give up valence electrons much more easily. Nonmetals (especially group VIIA and

VIA nonmetals) very easily gain new valence electrons. So, metals and non-metals tend to form bonds in which the metal atoms entirely surrender valence electrons to the nonmetals. Bonds with extremely unequal electron-sharing are called ionic bonds.

1. What is the electron dot structure of potassium fluoride?

Solve It

2. The ionic compound lithium sulfide forms between the elements lithium and sulfur. In which direction are electrons transferred to form ionic bonds, and how many electrons are transferred?

Solve It

3. Magnesium chloride is dissolved into a beaker of water and a beaker of rubbing alcohol until no more compound will dissolve. Electrical circuits are set up for each beaker in which wires lead from a battery into the solution, and a separate set of wires leads from the solution to a light bulb. The bulb connected to the aqueous solution circuit glows more brightly than the bulb connected to the alcohol solution circuit. Why?

Solve It

Sharing Charge with Covalent Bonds

.«S)MS> Sometimes the way for atoms to reach their most stable, lowest-energy states is to share valence electrons. When atoms share valence electrons, we say that they are engaged in Covalent bonding. The very word Covalent Means "together in valence." Compared to ionic bonding, covalent bonding tends to occur between atoms of similar electronegativity, most especially between nonmetals.

Just as ionic bonds tend to form in such a way that both atoms end up with completely filled valence shells, the atoms involved in covalent bonds tend to share electrons in such a way that each ends up with a completely filled valence shell. The shared electrons are attracted to the nuclei of both atoms, forming the bond. The simplest and best studied covalent bond is the one formed between two hydrogen atoms, shown in Figure 5-4. Separately, each atom has only one electron with which to fill its 1 S Orbital. By forming a covalent bond, each atom lays claim to two electrons within the molecule of dihydrogen. The figure shows various ways in which a covalent bond can be represented, explicitly depicting the valence shells (a), by using electron dot structures (b), or by signifying a shared pair of electrons with a single line (c). The latter two ways to show bonding are referred to as Lewis structures.

Figure 5-4:

Three repre-

Sentations of the formation of a,, , .. .. …..

, . (b) H* + *H ->■ H! H

Covalent

Bond in dihydrogen.

^^^^^^ (c) H* + *H -»- H-H

Atoms can share more than a single pair of electrons. When atoms share two pairs of electrons, they are said to form a Double bond, And when they share three pairs of electrons they are said to form a Triple bond. Examples of double and triple bonds are shown with electron dot and line structures in Figure 5-5.

Figure 5-5:

The formation of double bonds in carbon dioxide and triple bonds in dinitrogen.

C – + 2 -o:

-c-o:

:mi::: n: (:n – n:)

A few guidelines can help you figure out the correct Lewis structure for a molecule if you know the molecule’s formula. As an example, we work out the Lewis structure of formaldehyde, CH2O (Figure 5-6 can help you follow along):

1. Add up all the valence electrons for all the atoms in the molecule.

These are the electrons you can use to build the structure. Account for any extra or missing electrons in the case of ions. For example, if you know your molecule has +2 charge, remember to subtract two from the total number of valence electrons. In the case of formaldehyde, C has four valence electrons, each H has one valence electron, and O has six valence electrons. The total number of valence electrons is 12.

2. Pick a "central" atom to serve as the anchor of your Lewis structure.

The central atom is usually one that can form the most bonds, which is often the atom with the most empty valence orbital slots to fill. In larger molecules, some trial-and-error may be involved in this step, but in smaller molecules, some choices are obviously better than others. For example, carbon is a better choice than hydrogen to be the central atom because carbon tends to form four bonds, whereas hydrogen tends to form only one bond. In the case of formaldehyde, carbon is the obvious first choice because it can form four bonds, while oxygen can form only two, and each hydrogen can form only one.

+

3. Connect the other, "outer" atoms to your central atom using single bonds only.

Each single bond counts for two electrons. In the case of formaldehyde, attach the single oxygen and each of the two hydrogen atoms to the central carbon atom.

4. Fill the valence shells of your outer atoms. Then put any remaining electrons on the central atom.

In our example, carbon and oxygen should each have eight electrons in their valence shells; each hydrogen atom should have two. However, by the time we fill the valence shells of our outer atoms (oxygen and the two hydrogens), we have used up our allotment of 12 electrons.

5. Check whether the central atom now has a full valence shell.

If the central atom has a full valence shell, then your Lewis structure is drawn properly — it’s formally correct even though it may not correspond to a real structure. If the central atom still has an incompletely filled valence shell, then use electron dots (nonbonding electrons) from outer atoms to create double and/or triple bonds to the central atom until the central atom’s valence shell is filled. Remember, each added bond requires two electrons. In the case of our formaldehyde molecule, we must create a double bond between carbon and one of the outer atoms. Oxygen is the only choice for a double-bond partner, because each hydrogen can accommodate only two electrons in its shell. So, we use two of the electrons assigned to oxygen to create a second bond with carbon.

Sometimes a covalent bond is formed in which one atom donates both electrons to the bond, with the other atom contributing no electrons. This kind of bond is called a Coordinate cova-lent bond. Atoms with lone pairs are capable of donating both electrons to a coordinate cova-lent bond. A Lone pair Consists of two electrons paired within the same orbital that aren’t used in bonding. Even though covalent bonding usually occurs between nonmetals, metals can engage in coordinate covalent bonding. Usually, the metal receives electrons from an electron donor called a Ligand.

1. C(4 e) + H(1 e) + H(1 e) + 0(6 e) = 12 e

2. Carbon is central atom; it can form more bonds (4) than 0, H.

: 0 :

*0*

Figure 5-6: 3. H C H

Putting

Together

A Lewis " C

Structure. / \ 4. H H

Sometimes a given set of atoms can covalently bond with each other in multiple ways to form a compound. This situation leads to something called Resonance. Each of the possible bonded structures is called a Resonance structure. The actual structure of the compound is a Resonance hybrid, A sort of average of all the resonance structures. For example, if two atoms are connected by a single bond in one resonance structure, and the same two atoms are connected by a double bond in a second resonance structure, then those atoms are connected by a bond in the resonance hybrid that is worth 1>2 single bonds. A common example of resonance is found in ozone, O3, shown in Figure 5-7.

Figure 5-7: :0*j0:j0j

Resonance structures of ozone, shown in two representations.

- +

^: 0"—0=0:

10=0—0: j

-+

+

Or

+

Q.

Draw a Lewis structure for propene, C3H6.

A. First, add up the total valence electrons. Each carbon contributes 4 electrons, and each hydrogen contributes 1, for a total of 18 valence electrons. Next, pick a central atom. The best choice is a carbon atom because carbon can form four bonds, more than any hydrogen. Connect the remaining atoms to the central carbon with single bonds. To connect all the atoms into one molecule, the central carbon must be connected to each of the two other carbon atoms. These connections use up 16 of the 18 valence

Electrons, leaving 2 electrons that you can place onto one of the carbon atoms. One carbon atom in the structure still requires two additional electrons to fill its valence shell. The only way to fill this shell is to create a carbon-carbon double bond. Only one arrangement of hydrogen atoms to the three carbons allows you to fill all the carbon valence shells, as you can see in the following figure:

HH

H — C — C = C

H

/

C

\

4. Bertholite is the common name for dichlo-rine, a toxic gas that has been used as a chemical weapon. Why is bertholite most certainly a covalently bonded compound? What is the most likely electron dot structure of this compound?

Solve It

5. When aluminum chloride salt is dissolved in water, aluminum (III) cations become surrounded by clusters of six water molecules to form a "hexahydrated" aluminum cation, Al(H2O)63+. Being a group IIIA metal, aluminum easily gives up its valence electrons. The oxygen atom in water possesses two lone pairs. What kind of bonding most likely occurs between the aluminum and the hydrating water molecules?

H

H

Solve It

6. Benzene, C6H6, is a common industrial solvent. The benzene molecule is based on a ring of cova-lently bonded carbon atoms. Draw two acceptable Lewis structures for benzene. Based on the structures, describe a likely resonance hybrid structure for benzene.

Solve It

Occupying and Overlapping Molecular Orbitals

Chapter 4 describes how electrons occupy distinct orbitals within atoms. When atoms cova-lently bond to form molecules, the shared electrons are no longer constrained to those atomic orbitals, but occupy Molecular orbitals, Larger regions that form from the overlap of atomic orbitals. Just as different atomic orbitals are associated with different levels of energy, so are molecular orbitals. A stable covalent bond forms between two atoms because the energy of the molecular orbital associated with the bond is lower than the combined energies associated with the atomic orbitals of the separated atoms.

Because electrons have wave-like properties, atomic orbitals can overlap in different ways depending on the relationship between the waves of the shared electrons.

In one mode, the electron waves interact Favorably (with low energy) and together occupy a Bonding orbital.

In another mode, the electron waves interact unfavorably within a higher-energy Anti-bonding orbital.

The energy relationships between unbound atoms and different types of molecular orbitals are summarized within Molecular orbital diagrams, Such as the one shown for dihydrogen in Figure 5-8. In this figure, two hydrogen atoms each contribute a single electron from a 1 S Orbital to a sigma (o) Bonding orbital. The low-energy bonding orbital is favored over the higher-energy sigma antibonding (o*) Orbital. This illustrates a general principle: Given a choice between high – and low-energy states, molecules prefer the low-energy states. This preference for lower energy is what is meant by Favorable (low energy) versus Unfavorable (high energy).

Figure 5-8:

A molecular orbital diagram for the formation of dihydrogen.

Energy

H! H

H

H

In addition to differences in the interaction of electron waves, covalent bonds can differ based on the shape of the molecular orbitals.

When atomic orbitals overlap in such a way that the resulting molecular orbital is symmetrical with respect to the Bond axis (the line connecting the two bonded atoms), we say that a O Bond (sigma bond) Is formed.

When atomic orbitals overlap in such a way that the resulting molecular orbital is symmetric with the bond axis in only one plane, we say that a N Bond (pi bond) Is formed.

Sigma bonds are stronger than pi bonds because the electrons within sigma bonds lie directly between the two atomic nuclei. The negatively charged electrons in sigma bonds therefore experience favorable (as in, low-energy) attraction to the positively charged nuclei. Electrons in pi bonds are farther away from the nuclei, so they experience weaker attraction.

Sigma bonds form when S Or P Orbitals overlap in a head-on manner. Single bonds are usually sigma bonds. Pi bonds form when adjacent P Orbitals overlap above and below the bond axis. These situations are depicted in Figure 5-9.

Figure 5-9:

Formation of a sigma bond(o) from two S Orbitals, and formation of a pi bond (n) from two adjacent P Orbitals.

Ss