Chapter 5

Ionic Bonding

In This Chapter

Finding out why and how ions are formed

Understanding how ions create chemical bonds

Deciphering the formulas of ionic compounds

Naming ionic compounds

Connecting conductivity and ionic bonds

In this chapter, I introduce you to ionic bonding, the type of bonding that holds salts together. I discuss simple ions and polyatomic ions: how they form and how they combine. I also show you how to predict the formulas of ionic compounds and how chemists detect ionic bonds.

Forming Ions: Making Satisfying Electron Trades

In nature, achieving a filled (complete) valence energy level is a driving force of chemical reactions, because when that energy level is full, elements become stable, or “satisfied” — stable elements don’t lose, gain, or share electrons.

The noble gases — the VIIIA elements on the periodic table — are extremely nonreactive because their valence energy level (outermost energy level) is filled. However, the other elements in the A families on the periodic table do gain, lose, or share valence electrons to fill their valence energy level and become satisfied.

Because filling the valence energy level usually involves filling the outermost s and p orbitals, it’s sometimes called the octet rule — elements gain, lose, or share electrons to reach a full octet (eight valence electrons: two in the s orbital and six in the p orbital).

In this section, I explain how atoms gain or lose electrons to form ions and achieve stability. I also explain how ions can consist of single atoms or a group of atoms. (For info on achieving stability by sharing electrons, flip to Chapter 6.)

Gaining and losing electrons

When an atom gains or loses an electron, it develops a charge and becomes an ion. In general, the loss or gain of one, two, or sometimes even three electrons can occur, but an element doesn’t lose or gain more than three electrons.

Losing an electron to become a cation: Sodium

Ions that have a positive charge due to the loss of electrons are called cations. In general, a cation is smaller than its corresponding atom. Why? The filled energy level determines the size of an atom or ion, and a cation gives up enough electrons to lose an entire energy level.

Consider sodium, an alkali metal and a member of the IA family on the periodic table. Sodium has 1 valence electron and 11 total electrons, because its atomic number is 11. It has an electron configuration of 1s22s22p63s1. (See Chapter 2 for a review of electron configurations.)

By the octet rule, sodium becomes stable when it has eight valence electrons. Two possibilities exist for sodium to become stable: It can gain seven more electrons to fill energy level 3, or it can lose the one 3s electron so that energy level 2 (which is already filled at eight electrons) becomes the valence energy level.

So to gain stability, sodium loses its 3s electron. At this point, it has 11 protons (11 positive charges) and 10 electrons (10 negative charges). The once-neutral sodium atom now has a single positive charge [11 (+) plus 10 (–) equals 1+]. It’s now an ion, an atom that has a charge due to the loss or gain of electrons. You can write an electron configuration for the sodium cation:

Na+: 1s22s22p6

Note that if an ion simply has 1 unit of charge, positive or negative, you normally don’t write the 1; you just use the plus or minus symbol, with the 1 being understood.

Atoms that have matching electron configurations are isoelectronic with each other. The positively charged sodium ion (cation) has the same electron configuration as neon, so it’s isoelectronic with neon. So does sodium become neon by losing an electron? No. Sodium still has 11 protons, and the number of protons determines the identity of the element.

There’s a difference between the neutral sodium atom and the sodium cation: one electron. As a result, their chemical reactivities are different and their sizes are different. Because sodium loses an entire energy level to change from a neutral atom to a cation, the cation is smaller.

Gaining an electron to become an anion: Chlorine

Ions with a negative charge due to the gain of electrons are called anions. In general, an anion is slightly larger than its corresponding atom because the protons have to attract one or more extra electrons. The attractive force is slightly reduced, so the electrons are free to move outward a little.

Chlorine, a member of the halogen family — the VIIA family on the periodic table — often forms anions. It has seven valence electrons and a total of 17 electrons, and its electron configuration is 1s22s22p63s23p5. So to obtain its full octet, chlorine must lose the seven electrons in energy level 3 or gain one at that level.

Because elements don’t gain or lose more than three electrons, chlorine must gain a single electron to fill energy level 3. At this point, chlorine has 17 protons (17 positive charges) and 18 electrons (18 negative charges). So chlorine becomes an ion with a single negative charge (Cl–). The neutral chlorine atom becomes the chloride ion. The electronic configuration for the chloride anion is

Cl–: 1s22s22p63s23p6

The chloride anion is isoelectronic with argon. The chloride anion is also slightly larger than the neutral chlorine atom. To complete the octet, the one electron gained went into energy level 3. But now there are 17 protons attracting 18 electrons, so the electrons can move outward a bit.

Looking at charges on single-atom ions

In the periodic table, the roman numerals at the top of the A families show the number of valence electrons in each element. Because atoms form ions to achieve full valence energy levels, that means you can often use an element’s position in the periodic table to figure out what kind of charge an ion normally has. Here’s how to match up the A families with the ions they form:

IA family (alkali metals): Each element has one valence electron, so it loses a single electron to form a cation with a 1+ charge.

IIA family (alkaline earth metals): Each element has two valence electrons, so it loses two electrons to form a 2+ cation.

IIIA family: Each element has three valence electrons, so it loses three electrons to form a 3+ cation.

VA family: Each element has five valence electrons, so it gains three electrons to form an anion with a 3– charge.

VIA family: Each element has six valence electrons, so it gains two electrons to form an anion with a 2– charge.

VIIA family (halogens): Each element has seven valence electrons, so it gains a single electron to form an anion with a 1– charge.

Determining the number of electrons that members of the transition metals (the B families) lose is more difficult. In fact, many of these elements lose a varying number of electrons so that they form two or more cations with different charges.

Seeing some common one-atom ions

Table 5-1 shows the family, element, ion name, and ion symbol for some common monoatomic (one-atom) cations.

Table 5-1 Common Monoatomic Cations
Family	Element	Ion Name	Ion Symbol
IA	Lithium	Lithium cation	Li+
	Sodium	Sodium cation	Na+
	Potassium	Potassium cation	K+
IIA	Beryllium	Beryllium cation	Be2+
	Magnesium	Magnesium cation	Mg2+
	Calcium	Calcium cation	Ca2+
	Strontium	Strontium cation	Sr2+
	Barium	Barium cation	Ba2+
IB	Silver	Silver cation	Ag+
IIB	Zinc	Zinc cation	Zn2+
IIIA	Aluminum	Aluminum cation	Al3+

Table 5-2 gives the same information for some common monoatomic anions.

Table 5-2 Common Monoatomic Anions
Family	Element	Ion Name	Ion Symbol
VA	Nitrogen	Nitride anion	N3–
	Phosphorus	Phosphide anion	P3–
VIA	Oxygen	Oxide anion	O2–
	Sulfur	Sulfide anion	S2–
VIIA	Fluorine	Fluoride anion	F–
	Chlorine	Chloride anion	Cl–
	Bromine	Bromide anion	Br–
	Iodine	Iodide anion	I–

Possible charges: Naming ions with multiple oxidation states

The electrical charge that an atom achieves is sometimes called its oxidation state. Many of the transition metal ions (the B families) have varying oxidation states because these elements can vary in how many electrons they lose. Table 5-3 shows some common transition metals that have more than one oxidation state.

Table 5-3 Common Metals with More than One Oxidation State
Family	Element	Ion Name	Ion Symbol
VIB	Chromium	Chromium (II) or chromous	Cr2+
		Chromium (III) or chromic	Cr3+
VIIB	Manganese	Manganese (II) or manganous	Mn2+
		Manganese (III) or manganic	Mn3+
VIIIB	Iron	Iron (II) or ferrous	Fe2+
		Iron (III) or ferric	Fe3+
	Cobalt	Cobalt (II) or cobaltous	Co2+
		Cobalt (III) or cobaltic	Co3+
IB	Copper	Copper (I) or cuprous	Cu+
		Copper (II) or cupric	Cu2+
IIB	Mercury	Mercury (I) or mercurous	Hg22+
		Mercury (II) or mercuric	Hg2+
IVA	Tin	Tin (II) or stannous	Sn2+
		Tin (IV) or stannic	Sn4+
	Lead	Lead (II) or plumbous	Pb2+
		Lead (IV) or plumbic	Pb4+

Notice that these cations can have more than one name. Here are two ways to name cations of elements that have more than one oxidation state:

Current method: Use the metal name, such as chromium, followed by the ionic charge written as a roman numeral in parentheses, such as (II). For example, Cr2+ is chromium (II) and Cr3+ is chromium (III).

Traditional method: An older way of naming ions uses -ous and -ic endings. When an element has more than one ion, do the following:

• Give the ion with the lower oxidation state (lower numerical charge, ignoring the + or –) an -ous ending.

• Give the ion with the higher oxidation state (higher numerical charge) an -ic ending.

So for chromium, the Cr2+ ion is named chromous and the Cr3+ ion is named chromic.

Grouping atoms to form polyatomic ions

Ions can be polyatomic, composed of a group of atoms. For example, take a look at Table 5-3 in the preceding section. Notice anything about the mercury (I) ion? Its ion symbol, Hg22+, shows that two mercury atoms are bonded together. This group has a 2+ charge, with each mercury cation having a 1+ charge. The mercurous ion is classified as a polyatomic ion.

Similarly, the symbol for the sulfate ion, SO42–, indicates that one sulfur atom and four oxygen atoms are bonded together and that the whole polyatomic ion has two extra electrons: a 2– charge.

Polyatomic ions are treated the same as monoatomic ions (see “Naming ionic compounds,” later in this chapter). Table 5-4 lists some important polyatomic ions.

Table 5-4 Some Important Polyatomic Ions
Ion Name	Ion Symbol
Sulfite	SO32–
Sulfate	SO42–
Thiosulfate	S2O32–
Bisulfate (or hydrogen sulfate)	HSO4–
Nitrite	NO2–
Nitrate	NO3–
Hypochlorite	ClO–
Chlorite	ClO2–
Chlorate	ClO3–
Perchlorate	ClO4–
Chromate	CrO42–
Dichromate	Cr2O72–
Arsenite	AsO33–
Arsenate	AsO43–
Phosphate	PO43–
Hydrogen phosphate	HPO42–
Dihydrogen phosphate	H2PO4–
Carbonate	CO32–
Bicarbonate (or hydrogen carbonate)	HCO3–
Cyanide	CN–
Cyanate	OCN–
Thiocyanate	SCN–
Peroxide	O22–
Hydroxide	OH–
Acetate	C2H3O2–
Oxalate	C2O42–
Permanganate	MnO4–
Ammonium	NH4+
Mercury (I)	Hg22+

Creating Ionic Compounds

Ionic bonding, the bonding that holds the cations and anions together, is one of the two major types of bonding in chemistry. (I describe the other type, covalent bonding, in Chapter 6.)

An ionic bond occurs between a metal and a nonmetal. The metal loses electrons (to becomes a positively charged cation), and a nonmetal gains those electrons (to become a negatively charged anion). The ions have opposite charges, so they’re attracted to each other. This attraction draws them together into a compound.

In this section, you look at how ionic bonding works, and you see how to write formulas for and name ionic compounds.

Making the bond: Sodium metal + chlorine gas = sodium chloride

The transfer of an electron creates ions — cations (positive charge) and anions (negative charge). Opposite charges attract each other, so the cations and anions may come together through an ionic bond. An ionic bond is a chemical bond (a strong attractive force that keeps two chemical elements together) that comes from the electrostatic attraction (attraction of opposite charges) between cations and anions. Together, the ions form a compound.

For instance, sodium, a metal, can fill its octet and achieve stability by losing an electron. Chlorine, a nonmetal, can fill its octet by gaining an electron. (See the earlier section “Gaining and losing electrons” for details on the octet rule.) If the two are in the same container, then the electron that sodium loses can be the same electron that chlorine gains. The Na+ cation attracts the Cl– anion and forms the compound NaCl, sodium chloride.

Compounds that have ionic bonds are commonly called salts. In sodium chloride — table salt — a crystal is formed in which each sodium cation is surrounded by six different chloride anions and each chloride anion is surrounded by six different sodium cations.

Different types of salts have different crystal structures. Cations and anions can have more than one unit of positive or negative charge if they lose or gain more than one electron. In this fashion, many different kinds of salts are possible.

Figuring out the formulas of ionic compounds

When an ionic compound is formed, the cation and anion attract each other, resulting in a salt. This section shows you how to write the formula of that salt.

Balancing charges: Magnesium and bromine

Suppose you want to know the formula, or composition, of a compound that results from reacting a metal and a nonmetal. You start by putting the two atoms side by side, with the metal on the left. Then you add their charges.

Figure 5-1 shows this process for magnesium and bromine. (Forget about the crisscrossing lines for now. I explain them in the upcoming section “Using the crisscross rule.”)

Figure 5-1: Figuring the formula of magnesium bromide.

The electron configurations for magnesium and bromine are

Magnesium (Mg): 1s22s22p63s2

Bromine (Br): 1s22s22p63s23p64s23d104p5

Magnesium, an alkaline earth metal, has two valence electrons that it loses to form a cation with a 2+ charge. The electron configuration for the magnesium cation is

Mg2+: 1s22s22p6

Bromine, a halogen, has seven valence electrons, so it gains one electron to complete its octet (eight valence electrons) and form the bromide anion with a 1– charge. The electron configuration for the bromide anion is

Br1–: 1s22s22p63s23p64s23d104p6

When writing the formula of a compound, the compound must be neutral. That is, it needs to have equal numbers of positive and negative charges. So after writing the atoms, you need to balance their charges.

The magnesium ion has a 2+, so it requires two bromide anions, each with a single negative charge, to balance the two positive charges of magnesium. So the formula of the compound that results from reacting magnesium with bromine is MgBr2.

Using the crisscross rule

A quick way to determine the formula of an ionic compound is to use the crisscross rule: Take the numerical value of the metal ion’s superscript (forget about the charge symbol) and move it to the bottom right-hand side of the nonmetal’s symbol — as a subscript. Then take the numerical value of the nonmetal’s superscript and make it the subscript of the metal. (Note that if the numerical value is 1, it’s just understood and not shown.)

To see how to use this rule, look back at Figure 5-1. For magnesium and bromine, you make magnesium’s 2 a subscript of bromine and make bromine’s 1 a subscript of magnesium (but because it’s 1, you don’t show it). You get the formula MgBr2.

So what happens if you react aluminum and oxygen? Figure 5-2 shows the crisscross rule for this reaction. You get Al2O3.

Figure 5-2: Figuring out the formula of aluminum oxide.

Compounds involving polyatomic ions work in exactly the same way. For example, here’s the compound made from the ammonium cation (NH4+) and the sulfide anion (S2–):

(NH4)2S

Notice that because you need two ammonium ions (two positive charges) to neutralize the two negative charges of the sulfide ion, you enclose the ammonium ion in parentheses and add a subscript 2.

After you use the crisscross rule, reduce all the subscripts by a common factor, if possible, to get the right formula.

For example, suppose that you want to write the compound formed when calcium reacts with oxygen. Calcium, an alkaline earth metal, forms a 2+ cation, and oxygen forms a 2– anion. So you may predict that the formula is

Mg2O2

But you need to divide each subscript by 2 to get the correct formula:

MgO

Naming ionic compounds

When you name inorganic compounds, you write the name of the metal first and then the nonmetal, adding an -ide ending to the nonmetal (for compounds involving monatomic ions).

Suppose, for example, that you want to name Li2S, the compound that results from the reaction of lithium and sulfur. You first write the name of the metal, lithium, and then write the name of the nonmetal, adding an -ide ending so that sulfur becomes sulfide:

Li2S: Lithium sulfide

Ionic compounds involving polyatomic ions follow the same basic rule: Write the name of the metal first, and then simply add the name of the nonmetal. However, with polyatomic anions, it’s not necessary to add the -ide ending:

(NH4)2CO3: Ammonium carbonate

K3PO4: Potassium phosphate

Dealing with multiple oxidation states

When the metal involved is a transition metal with more than one oxidation state (see Table 5-3, earlier in the chapter), there can be more than one way to correctly name the compound, based on how you name the metal.

For example, suppose that you want to name the compound formed between the Fe3+ cation and the cyanide ion, CN–. The preferred method is to use the metal name followed in parentheses by the ionic charge written as a roman numeral: iron (III). But an older naming method, which is still sometimes used (so it’s a good idea to know it), is to use -ous and -ic endings.

The ion with the lower oxidation state (lower numerical charge, ignoring the + or –) gets an -ous ending, and the ion with the higher oxidation state (higher numerical charge) gets an -ic ending. So because Fe3+ has a higher oxidation state than Fe2+, it’s called a ferric ion.

After you write the name of the metal, name the nonmetal. So the compound Fe(CN)3 can be named

Fe(CN)3: iron(III) cyanide, or ferric cyanide

Getting names from formulas and formulas from names

Sometimes figuring out the charge on an ion can be a little challenging (and fun), so take a look at how to name FeNH4(SO4)2. I show you earlier in Table 5-4 that the sulfate ion has a 2– charge, and from the formula you can see that there are two of these ions. Therefore, you have a total of four negative charges. Table 5-4 also indicates that the ammonium ion has a 1+ charge, so you can figure out the charge on the iron cation:

Ion	Charge
Fe	?
NH4	1+
(SO4)2	(2–) × 2

Because you have a 4– charge for the sulfates and a 1+ for the ammonium, the iron must be a 3+ to make the compound neutral. So the iron is in the iron (III), or ferric, oxidation state. You can name the compound:

FeNH4(SO4)2: Iron (III) ammonium sulfate, or ferric ammonium sulfate

And finally, if you have the name, you can derive the formula and the charge on the ions. For example, suppose that you’re given the name cuprous oxide. You know that the cuprous ion is Cu+ and the oxide ion is O2–. Applying the crisscross rule (from the earlier section “Using the crisscross rule”), you get the following formula:

Cuprous oxide: Cu2O

Bonding Clues: Electrolytes and Nonelectrolytes

Scientists can get some good clues about the type of bonding in a compound by discovering whether a substance is an electrolyte or a nonelectrolyte. Ionically bonded substances act as electrolytes, but covalently bonded compounds, in which no ions are present, are commonly nonelectrolytes.

Electrolytes are substances that conduct electricity in the molten state or when dissolved in water. For instance, sodium chloride is an electrolyte because it conducts an electrical current when dissolved in water. If you were to melt pure NaCl (which requires a lot of heat!) and then check the conductivity of the molten salt, you’d find that the molten table salt also conducts electricity. In the molten state, the NaCl ions are free to move and carry electrons, just as they are in the saltwater solution.

Substances that don’t conduct electricity when in these states are called nonelectrolytes. Table sugar, or sucrose, is a good example of a nonelectrolyte. You can dissolve sugar in water or melt it, but it won’t have conductivity. No ions are present to transfer the electrons.