# Introductory Chemistry Online/Chemical Bonding and Nomenclature
## 3.1 Compounds, Lewis Diagrams and Ionic Bonds

If we take two or more atoms and bond them together chemically so that they now behave as a single substance, we have made a chemical compound. We will see that the process of bonding actually involves either the sharing, or the net transfer, of electrons from one atom to another. The two types of bonding are covalent, for the sharing of electrons between atoms, and ionic, for the net transfer of electrons between atoms. Covalent or ionic bonding will determine the type of compound that will be formed.
In Chapter 1, we used atomic theory to describe the structure of the fluorine atom. We said that neutral fluorine has nine protons in its nucleus (an atomic number of 9), nine electrons surrounding the nucleus (to make it neutral), and the most common isotope has ten neutrons in its nucleus, for a mass number of 19. Further, we said that the nine electrons exist in two energy levels; the first energy level contains two electrons and is written 1s$^2$. The second energy level contains seven electrons, distributed as 2s$^2$ 2p$^5$. The outermost electron level in any atom is referred to as the valence shell. For the representative elements (remember, this includes all of the elements except for the transition metals), the number of electrons in the valence shell corresponds to the Group number of the element in the periodic table. Group 1A elements will have one valence electron, Group 6A elements will have six valence electrons, and so on. Fluorine is a Group 7A element and has seven valence electrons. We can show the electron configuration for fluorine using a Lewis diagram (or electron-dot structure), named after the American chemist G. N. Lewis, who proposed the concepts of electron shells and valence electrons. In a Lewis diagram, the electrons in the valence shell are shown as small “dots” surrounding the atomic symbol for the element.Equation 1
When more than four electrons are present in the valence shell, they are shown as pairs when writing the Lewis diagram (but never more than pairs). Lewis diagrams for the atoms in the second period are shown in Figure 3.1. As you look at the dot-structures in Figure 3.1, please understand that it makes no difference where you place the electrons, or the electron pairs, around the symbol, as long as pairs are shown whenever there are four or more valence electrons.
If you examine the Lewis diagram for neon (Ne) in Figure 3.1, you will see that the valence shell is filled; that is, there are eight electrons in the valence shell. Elements in Group 8A of the periodic table are called noble gasses; they are very stable and do not routinely combine with other elements to form compounds (although today, many compounds containing noble gasses are known). Modern bonding theory tells us that this stability arises because the valence shell in the noble gasses is completely filled. When the valence shell is not full, theory suggests that atoms will transfer or share electrons with other atoms in order to achieve a filled valence shell… that is, the electron configuration of the noble gasses. Chemical bonding can then be viewed as a quest by atoms to acquire (or lose) enough electrons so that their valence shells are filled, that is, to achieve a “noble gas configuration”. This is often referred to as the “octet rule”; the desire for elements to obtain eight electrons in the valence shell (except of course for helium where the noble gas configuration is two valence electrons).
Atoms can achieve a noble gas configuration by two methods; the transfer of electrons from their valence shells to another atom, or by sharing electrons with another atom. If you examine the Lewis diagram for lithium (Li), you will see that it has only one valence electron. If lithium was to transfer this electron to another atom, it would be left with two electrons in the 1s-orbital (denoted as 1s$^2$). This is the same electron configuration as helium (He), and so by losing this electron, lithium has achieved a noble gas configuration. Because electrons carry a negative charge, the loss of this electron leaves lithium with a single positive charge. This is the lithium cation and it is shown as Li$^+$.Equation 2
Returning to fluorine (F), in order to achieve the 2s$^2$ 2p$^6$ configuration of neon (Ne), fluorine needs to gain one valence electron. Because fluorine has gained one electron, it now has one negative charge. This is the fluoride anion and it is shown as F$^-$. The transfer of electrons in order to achieve a noble gas configuration is the process known as ionic bonding, and this will be covered in more detail later in this chapter.Equation 3
Example 3.1 Drawing Lewis Diagrams
Exercise 3.1 Drawing Lewis Diagrams

## 3.2 Covalent Bonding

A second method by which atoms can achieve a filled valence shell is by sharing valence electrons with another atom. Thus fluorine, with one unpaired valence electron, can share that electron with an unshared electron on another fluorine to form the compound, F$_2$ in which the two shared electrons form a chemical bond holding the two fluorine atoms together (Figure 3.2). When you do this, each fluorine now has the equivalent of eight electrons in its valence shell; three unshared pairs and one pair that is shared between the two atoms. Note that when you are counting electrons, the electrons that are shared in the covalent bond are counted for each atom, individually. A chemical bond formed by sharing electrons between atoms is called a covalent bond. When two or more atoms are bonded together utilizing covalent bonds, the compound is referred to as a molecule.
There is a simple method, given below, that we can use to construct Lewis diagrams for diatomic and for polyatomic molecules:
Begin by adding up all of the valence electrons in the molecule. For F$_2$, each fluorine has seven, giving a total of 14 valence electrons.
Next, draw your central atom. For a diatomic molecule like F$_2$, both atoms are the same, but if several different atoms are present, the central atom will be to the left (or lower) in the periodic table.
Next, draw the other atoms around the central atom, placing two electrons between the atoms to form a covalent bond.
Distribute the remaining valence electrons, as pairs, around each of the outer atoms, so that they all are surrounded by eight electrons.
Place any remaining electrons on the central atom.
If the central atom is not surrounded by an octet of electrons, construct multiple bonds with the outer atoms until all atoms have a complete octet.
If there are an odd number of valence electrons in the molecule, leave the remaining single electron on the central atom.
Let’s apply these rules for the Lewis diagram for chlorine gas, Cl$_2$. There are 14 valence electrons in the molecule. Both atoms are the same, so we draw them next to each other and place two electrons between them to form the covalent bond. Of the twelve remain electrons, we now place six around one chlorine (to give an octet) and then place the other six around the other chlorine (our central atom). Checking, we see that each atom is surrounded by an octet of valence electrons, and so our structure is complete Figure 3.3
All of the Group 7A elements (the halogens), have valence shells with seven electrons and all of the common halogens exist in nature as diatomic molecules; fluorine, F$_2$; chlorine, Cl$_2$; bromine, Br$_2$ and iodine I$_2$ (astatine, the halogen in the sixth period, is a short-lived radioactive element and its chemical properties are poorly understood). Nitrogen and oxygen, Group 5A and 6A elements, respectively, also exists in nature as diatomic molecules (N$_2$ and O$_2$). Let’s consider oxygen; oxygen has six valence electrons (a Group 6A element). Following the logic that we used for chlorine, we draw the two atoms and place one pair of electrons between them, leaving 10 valence electrons. We place three pairs on one oxygen atom, and the remaining two pairs on the second (our central atom). Because we only have six valence electrons surrounding the second oxygen atom, we must move one pair from the other oxygen and form a second covalent bond (a double bond) between the two atoms. Doing this, each atom now has an octet of valence electrons. Figure 3.4
Nitrogen has five valence electrons. Sharing one on each atom gives the first intermediate where each nitrogen is surrounded by six electrons (not enough!). Sharing another pair, each nitrogen is surrounded by seven electrons, and finally, sharing the third, we get a structure where each nitrogen is surrounded by eight electrons; a noble gas configuration (or the “octet rule”). Nitrogen is a very stable molecule and relatively unreactive, being held together by a strong triple covalent bond. Figure 3.5
As we have constructed Lewis diagrams, thus far, we have strived to achieve an octet of electrons around every element. In nature, however, there are many exceptions to the “octet rule”. Elements in the first row of the periodic table (hydrogen and helium) can only accommodate two valence electrons. Elements below the second row in the periodic table can accommodate, 10, 12 or even 14 valence electrons (we will see an example of this in the next section). Finally, in many cases molecules exist with single unpaired electrons. A classic example of this is oxygen gas (O$_2$). We have previously drawn the Lewis diagram for oxygen with an oxygen-oxygen double bond. Physical measurements on oxygen, however, suggest that this picture of bonding is not quite accurate. The magnetic properties of oxygen, O$_2$, are most consistent with a structure having two unpaired electrons in the configuration shown in Figure 3.6.
In this Lewis diagram, each oxygen atom is surrounded by seven electrons (not eight). This electronic configuration may explain why oxygen is such a reactive molecule (reacting with iron, for example, to form rust); the unpaired electrons on the oxygen molecule are readily available to interact with electrons on other elements to form new chemical compounds.
Another notable exception to the “octet rule” is the molecule NO (nitrogen monoxide). Combining one nitrogen (Group 5A) with one oxygen (Group 6A) gives a molecule with eleven valence electrons. There is no way to arrange eleven electrons without leaving one electron unpaired. Nitric oxide is an extremely reactive molecule (by virtue of its unshared electron) and has been found to play a central role is biochemistry as a reactive, short-lived molecule involved in cellular communication.Figure 3.7
As useful as Lewis diagrams can be, chemists tire of drawing little dots and, for a shorthand representation of a covalent bond, a short line (a line-bond) is often drawn between the two elements. Whenever you see atoms connected by a line-bond, you are expected to understand that this represents two shared electrons in a covalent bond. Further, the unshared pairs of electrons on the bonded atoms are sometimes shown, and sometimes they are omitted (see Figure 3.8). If unshared pairs ore omitted, the chemist reading the structure is assumed to understand that they are present.Figure 3.8

## 3.3 Lewis Representation of Ionic Compounds

As mentioned in Section 3.1, elements can also transfer electrons to another element in order to achieve the noble gas configuration. Consider sodium. Sodium (Na) is in Group 3A. Its first energy level is completely filled (1s$^2$ ), the second energy level is also filled (2s$^2$ 2p$^6$) and there is a single electron in the third energy level (3s$^1$). Energetically, the easiest way for sodium to achieve an octet of valence electrons is by transferring its valence electron to an acceptor atom. This will leave the sodium atom with the same electron configuration as neon (1s$^2$2s$^2$ 2p$^6$), and will satisfy the “octet rule”. Because sodium loses one electron (with its negative charge) the sodium atom must now have a positive charge. Atoms, or covalently bound groups of atoms with a positive charge are called cations, and the sodium cation is written as Na$^+$. Figure 3.9
If the acceptor atom in the example above was chlorine, the third valence shell would now be filled, matching the electron configuration of argon. Because the chlorine atom has accepted an electron, with its negative charge, the chlorine atom must now have a negative charge. Atoms, or covalently bound groups of atoms with a negative charge are called anions, and the chlorine anion is written as Cl$^-$. Although electron transfer has occurred between the two atoms in this example, there is no direct bond holding the sodium cation and the chlorine anion together, other than the simple electrostatic attraction between the two charged atoms. This is the real difference between ionic and covalently bound atoms; covalent molecules are held together in a specific geometry which is dictated by the electrons that they share. Ionic compounds are held together by simple electrostatic attraction and, unless these atoms are present in organized crystals, there is no defined geometric order to this attraction. Ionic compounds are often referred to as salts. Some simple examples of ionization to form ionic salts are shown in Figure 3.10. Line-bond shorthand is used in Table 3.1 to show a series of common hydrogen-containing compounds and ions formed from second-row elements.
Example 3.2 Drawing Lewis Diagrams for Simple Covalent Molecules
Exercise 3.2 Drawing Lewis Diagrams for Simple Covalent Molecules

## 3.4 Identifying Molecular and Ionic Compounds

The tendency for two or more elements to combine and form a molecule that is stabilized by covalent bonds (a molecular compound) can be predicted simply by the location of the various elements on the periodic table. In Chapter 1, we divided the elements in the periodic table into (seemingly) arbitrary groupings; the metals, the non-metals, the semi-metals, and so on (See Figure 3.11). These groupings are not arbitrary, but are largely based on physical properties and on the tendency of the various elements to bond with other elements by forming either an ionic or a covalent bond. As a general rule of thumb, compounds that involve a metal binding with either a non-metal or a semi-metal will display ionic bonding. Compounds that are composed of only non-metals or semi-metals with non-metals will display covalent bonding and will be classified as molecular compounds. Thus, the compound formed from sodium and chlorine will be ionic (a metal and a non-metal). Nitrogen monoxide (NO) will be a covalently bound molecule (two non-metals), silicon dioxide (SiO$_2$) will be a covalently bound molecule (a semi-metal and a non-metal) and MgCl$_2$ will be ionic (a metal and a non-metal). Later in this chapter we will see that many covalent compounds have bonds that are highly polarized with greater electron density around one atom than the other. These compounds are often described as having “ionic character” and these types of covalent bonds can often be readily broken to form sets of ions.
Example 3.3 Identifying Molecular Compounds
Exercise 3.3 Identifying Molecular Compounds

## 3.5 Polyatomic Ions

The compound NaOH has wide industrial use and is the active ingredient in drain cleaners. Based on the discussion in the previous section, we would expect NaOH to be an ionic compound because it contains sodium, a Group 1A metal. Hydrogen and oxygen, however, are nonmetals, and we would expect these to bond together covalently. This compound, called sodium hydroxide, is an example of an ionic compound formed between a metal ion (sodium) and a polyatomic ion (HO$^–$). Charged groups of atoms, like HO$^–$, that are bonded together covalently are called polyatomic ions. Within an ionic compound a polyatomic behaves as a single unit forming salts with other cations or anions.
Using the rules described in Section 3.2, we can draw a Lewis diagram for HO$^–$. Oxygen has six valence electrons and hydrogen has one, for a total of seven. The central atom in our structure will be hydrogen (it is to the left of oxygen in the periodic table). Next, because this is a polyatomic ion with a single negative charge, we add the extra electron to the central atom, pair the electrons and then draw the two atoms bonded together. Next, the six remaining electrons are distributed around the oxygen to form an octet. Finally, the polyatomic ion is enclosed in brackets with the charge as a superscript to show that the ion behaves as a single unit.
Polyatomic ions are very common in chemistry and formulas and charges for common polyatomic ions are collected in Table 3.2. It is essential that you memorize these and be able to correlate the name, the composition and the charge for each of them, as they will be discussed freely throughout the remainder of the course and you will be expected to know these in General Chemistry.
Example 3.4 Lewis Diagrams for Polyatomic Ions
Construct a Lewis diagram for the polyatomic ion CO$_3$$^{2–}$.
Solution: Oxygen has six valence electrons and carbon has four; therefore in CO$_3$$^{2–}$ there will be a total of 22 valence electrons, plus two additional electrons from the 2- charge. The central atom in our structure will be carbon (it is to the left of oxygen in the periodic table). Next, we draw the carbon (our central atom) with its’ four electrons and add the additional two electrons from the charge. The three oxygens are placed around the carbon and the electrons are arranged to form the three covalent bonds. Next, the 18 remaining electrons are distributed around the oxygens so that they all have a full octet. The carbon, however, is only surrounded by six electrons. To remedy this, we move one electron pair in to form a double bond to one of the oxygen atoms. Finally, the polyatomic ion is enclosed in brackets with the charge as a superscript.
The structure for this ion can also be drawn using line-bond shorthand, as shown below:
We need to understand that the process of placing electrons into a particular bond in a compound is an artificial aspect of building Lewis diagrams. In fact, the electrons are added to the polyatomic ion, but it is impossible to know exactly where they went.
Exercise 3.4 Drawing Lewis Diagrams for Polyatomic Ions.
Draw Lewis diagrams for NO$_2$$^-$ and NH$_4$$^+$.

## 3.6 Resonance

In Example 3.3, we constructed a Lewis diagram for the carbonate anion. Our final structure showed two carbon-oxygen single bonds and one carbon-oxygen double bond. The structure that we drew is shown below, along with two other possible representations for the carbonate anion. These structures differ only in the position of the carbon-oxygen double bond.
So which of these is correct? Actually, they all are! These are all “proper” Lewis diagrams for a covalent structure having constant geometry, and the diagrams differ only in the manner that we have arbitrarily arranged the electrons. These Lewis diagrams are called resonance forms. For the carbonate anion, there are three equivalent resonance forms that can be drawn. It is important to note that the electrons are not “hopping” between the atoms, but that the electrons are spread evenly between the carbon and all three oxygens and that each carbon-oxygen bond has a bond-order of 1.33 (one and one-third covalent bonds). A structure such as this is called the resonance hybrid (Figure 3.14), and although it most clearly represents the actual bonding in the compound, it is often difficult to understand the nature of the bonding when structures are represented as resonance hybrids. A full discussion of resonance is beyond the scope of an introductory text and, for structures such as the carbonate anion, we will accept any of the proper resonance forms shown above.

## 3.7 Electronegativity and the Polar Covalent Bond

If we were to construct a Lewis diagram for molecular hydrogen (H$_2$), we would pair the single valence electrons on each atom to make a single covalent bond. Each hydrogen would now have two electrons in its valence shell, identical to helium. The mathematical equations chemists use to describe covalent bonding can be solved to predict the regions of space surrounding the molecule in which these electrons are likely to be found. A particularly useful application of these calculations generates a molecular surface that is color coded to show electron density surrounding the molecule. This type of molecular surface is called an electrostatic potential map, and the calculated map for the hydrogen molecule is shown in Figure 3.15.
When this type of calculation is done for molecules consisting of two (or more) different atoms, the results can be strikingly different. Consider the molecule HF. Hydrogen, with one valance electron, can share that electron with fluorine (with seven valence electrons) to form a single covalent bond. A ball-and-stick model of HF is shown in Figure 3.16, along with the electrostatic potential map for the HF molecule.
In this electrostatic potential map, blue is used to indicate low electron density (a relative positive charge) and red indicates high electron density (a relative negative charge); the colors light blue, green, yellow and orange indicate the increasing charge gradient. The molecule HF is clearly very polar, meaning that a significant difference in electron density exists across the length of the molecule. The electrostatic potential map for HF contrasts significantly with that for H$_2$, where the charge was quite symmetrical (a uniform green color). Hydrogen fluoride (HF) can be described as a very polar molecule, while hydrogen (H$_2$) is nonpolar.
The origin of the polarization of the HF covalent bond has to do with electronegativity, an inherent property of all atoms. Within the periodic table, there is a trend for atoms to attract electrons towards themselves when they are bonded to another atom (as in HF). Atoms that tend to strongly attract electrons have a high electronegativity, relative to atoms that have a relatively low tendency to attract electrons towards themselves. The modern electronegativity scale was devised by Linus Pauling in 1932 and, in the Pauling scale, atoms in the periodic table vary in electronegativity from a low of 0.8 for cesium to a maximum of 4.0 for fluorine. A periodic table, modified to show electronegativity trends is shown in Figure 3.17.
Figure 3.17 A periodic table, modified to show the trend in electronegativity, increasing from francium (0.7) to fluorine (4.0).
In the molecule HF, the electronegativity of the hydrogen is 2.2 and fluorine is 4.0. This difference leads to the profound polarization of the HF covalent bond which is apparent in the electrostatic potential map.
Polarizations of covalent bonds also occur in more complex molecules. In water, the oxygen has an electronegativity of 3.5; hydrogen is 2.2. Because of this, each of the H-O bonds is polarized with greater electron density towards the oxygen. Within the molecule, H$_2$O, the effect of this polarization becomes apparent in the electrostatic potential map, as shown in Figure 3.18. The end of the molecule with the oxygen has a high electron density and the hydrogen ends are electron deficient. We will see in later chapters that the polarization of water, caused by the difference in electronegativities, gives water the special properties that allows it to dissolve ionic compounds, and basically support life as we know it. Within organic chemistry (the study of carbon-containing molecular compounds), you will appreciate that the relative reactivity of organic molecules with each other is largely dependent on the polarization of covalent bonds in these molecules.

## 3.8 Exceptions to the Octet Rule

Returning briefly to classical Lewis diagrams. Consider the diagram shown below for the molecule SF$_4$. In constructing this diagram, six valence electrons are placed around the sulfur and seven valence electrons are placed around each fluorine. As we attempt to pair these to form covalent bonds, we note that there are “too many” electrons on the sulfur! We clearly cannot form a covalent bond using three electrons, so we split one pair, move the single electrons into bonding position and form bonds with the remaining two fluorines. The “extra pair” of electrons just sits there on the sulfur and does not participate directly in the bonding.
This is an example of valence expansion. In general, elements below the second period in the periodic table (S, Se, Te, etc.) will commonly have 10 – 12 electrons in their valence shells. As in SF$_4$, these electrons are not directly involved in the formation of covalent bonds, but they affect the overall reactivity of the particular molecule.
Examples of compounds that exhibit valence expansion from Groups 4A – 8 are shown in Table 3.3. In general, all molecular compounds containing elements that appear below the second row in the periodic table are capable of valence expansion and you need to be very careful when you are drawing Lewis diagrams for these compounds. As we saw in Section 3.1 for the molecule nitrogen oxide (NO), stable molecules also exist in which atoms are not surrounded by an octet of electrons. Another example of this is the molecule BF$_3$, which is shown in the following example.
Example 3.5 Lewis Diagrams for BF$_3$
Construct a Lewis diagram for the molecule BF$_3$.
Solution: Boron has three valence electrons and fluorine has seven. The central atom in our structure will be boron (it is to the left of fluorine in the periodic table). Next, we draw the boron (our central atom) with its’ three electrons and place the three fluorines around the boron with the electrons arranged to form the three covalent bonds. Each of the fluorines have a full octet. The boron, however, is only surrounded by six electrons. Because of this, the boron in BF$_3$ is a powerful electron acceptor and forms strong complexes with electrons from other compounds. In Chapter 8 we will see that this property is called Lewis acidity and BF$_3$ is a very powerful Lewis acid.

## 3.9 Common Valence States and Ionic Compounds

In an ionic compound the total number of charges on the cations must equal the total number of charges on the anions; that is, the compound must be neutral. In Section 3.2 we described how a sodium atom can donate an electron to another atom in order to form an ion with a full octet of electrons in its outermost electron shell (the same electron configuration as Ne). The charge on the sodium atom, when this happens, is now 1+, because it has eleven protons in its nucleus, but is only surrounded by 10 electrons. Lithium, likewise, can lose one electron to form Li$^{1+}$ and be left with the same electron configuration as He. In fact, all of the Group 1A metals can lose a single electron to form 1+ ions. Elements in Group 2A can each lose two electrons to form 2+ ions and achieve a noble gas configuration. In fact, the group that a main-group element is associated with in the periodic table will dictate the valence (or charge) of its corresponding ion. Metals in Groups 1A, 2A and 3A will form ions with 1+, 2+ and 3+ charges, respectively.
Main-group nonmetals can easily achieve an octet of valence electrons by accepting electrons from other elements. Thus Group 5A elements can accept three electrons to form 3- ions, Group 6A elements accept two electrons to form 2- ions and Group 7A elements (the halogens) accept one electron to form 1- ions. For example, oxygen (Group 6A) needs to accept two electrons to achieve the electron configuration of neon. This gives oxygen a total of 10 electrons, but it only has eight protons in its nucleus (its atomic number is 8), therefore, the oxygen ion has a net charge of 2- (O$^{2-}$).
To write a formula for an ionic compound composed of main group elements (or containing polyatomic ions) you need to adjust the ratio of anions and cations so that the resulting molecule is electrically neutral. For example, consider an ionic compound containing sodium and chlorine. Lithium is a Group 1A element and will form a 1+ ion; fluorine is a Group 7A element and will form a 1- ion. Neutrality is achieved when one lithium is paired with one fluorine, or LiF. For a compound composed of calcium and chlorine, the Group 2A calcium will form a 2+ ion while chlorine forms a 1- ion. To achieve neutrality, there must be two chlorines for every calcium, and the formula must be as CaCl$_2$. Aluminum (Group 3A) will form a 3+ ion. If this was paired with oxygen (Group 6A) which forms a 2- ion neutrality would only be achieved if two Al$^{3+}$ ions (for a total of six positive charges) were paired with three O$^{2-}$ ions (a total of six negative charges).
Consider a compound consisting of sodium and the polyatomic ion sulfate (SO$_4$$^{2-}$). Sodium (Group 1A) yields a 1+ cation and so there must be two sodiums in the compound for every sulfate (which has a 2- charge), or Na$_2$SO$_4$. For a compound containing calcium (Group 2A) and nitrate (NO$_3$$^-$), two nitrate anions must be present for every calcium 2+ cation. In a compound containing multiple copies of a polyatomic ion, the entire ion is enclosed in parenthesis with a subscript to indicate the number of units. Thus the compound from calcium and nitrate would be written as Ca(NO$_3$)$_2$.

## 3.10 Nomenclature of Ionic Compounds

The simplest ionic compounds consist of a single type of cation associated with a single type of anion. Nomenclature for these compounds is trivial; the cation is named first, followed by the anion. If the anion is a single element, the suffix ide is added to the root name of the element. A set of common anions from Figure 3.19 is shown in Table 3.4.
When you are constructing names for ionic compounds, you do not use “multipliers” to indicate how many cations or anions are present in the compound. For example NaI is named sodium iodide; Na$_2$S is named sodium sulfide; CaCl$_2$ is named calcium chloride. The chemist reading the name is assumed to have sufficient knowledge to pair the elements properly based on their common valence states. There are exceptions to this simple nomenclature, however. Many transition metals exist as more than one type of cation. Thus, iron exists as Fe$^{2+}$ and Fe$^{3+}$ cations (they are referred to as “oxidation states”, and will be covered in detail in Chapter 5). When you are naming an ionic compound containing iron, it is necessary to indicate which oxidation state the metal has. For metals, the oxidation state is the same as the charge. Thus Fe$^{2+}$ in a compound with chloride would have a formula FeCl$_2$ and would be named iron (II) chloride, with the oxidation state (the charge on the iron) appearing as a Roman numeral in parenthesis after the cation. The cation Fe$^{3+}$ paired with oxygen would have the formula Fe$_2$O$_3$ and would have the name iron (III) oxide. Some common transition metals with variable oxidation states are shown in Table 3.5.
The procedure for naming ionic compounds contain polyatomic ions is identical to that described above for simple ions. Thus, CaCO$_3$ is named calcium carbonate; Na$_2$SO$_4$ is named sodium sulfate; (NH$_4$)$_2$HPO$_4$ (a compound with two polyatomic ions) is named ammonium hydrogen phosphate; and Pb$^{2+}$ paired with SO$_4$$^{2-}$, PbSO$_4$ is named lead (II) sulfate.
Example 3.6 Nomenclature of Ionic Compounds
Write a correct chemical formula for each of the following ionic compounds:
a. Calcium bromide b. Aluminum oxide
c. Copper (II) chloride d. Iron (III) oxide
Write a proper chemical name for each of the following ionic compounds:
e. Li$_2$S f. CaO
g. NiCl$_2$ h. FeO
Solutions:
a. Calcium is 2+, bromide is 1-; CaBr$_2$.
b. Aluminum is 3+, oxide is 2-; Al$_2$O$_3$.
c. From the oxidation state that is given, copper is 2+, chloride is 1-; CuCl$_2$.
d. From the oxidation state, iron is 3+, oxide is 2-; Fe$_2$O$_3$.
e. We don’t use multipliers, so this is simply lithium sulfide.
f. This is simply calcium oxide.
g. We don’t have to specify an oxidation state for nickel, so this is nickel chloride.
h. We must specify that iron is 2+ in this compound; iron (II) oxide.
Exercise 3.6 Nomenclature of Ionic Compounds
Write a correct chemical formula for each of the following ionic compounds:
a. Sodium phosphide
b. Iron (II) nitrite
c. Calcium hydrogen phosphate
d. Chromium (III) oxide
Write a proper chemical name for each of the following ionic compounds:
e. NaBr
f. CuCl$_2$
g. Fe(NO$_3$)$_3$
h. (NH$_4$)$_3$PO$_4$

## 3.11 Nomenclature of Molecular Compounds

The nomenclature of simple binary molecular compounds (covalently bonded compounds consisting of only two elements) is slightly more complicated than the nomenclature of ionic compounds because multipliers must be used to indicate the ratio of the elements in the molecule; the multiplier mono is only used for the second element in a compound. The multipliers that are used are shown in Table 3.6.
Further, when you are naming a molecular compound, you must also decide which element should be listed first. In general, elements appearing to the left or lower in the periodic table are listed first in the name. Once you have decided on the order, the second element is named using the element root and ide, just like in ionic compounds. Thus, for CCl$_4$, carbon is to the left of chlorine (Group 4A vs. Group 5A), so it is listed first. There are four chlorines, so the multiplier tetra is used, and the name is carbon tetrachloride. Compounds containing hydrogen are generally an exception, and the hydrogen is listed as the first element in the name. Thus, H$_2$S would be named using the multiplier di to indicate that there are two hydrogens and mono to indicate that there is only one sulfur, or, dihydrogen monosulfide.
For the molecule SO$_2$; they are both Group 6A elements, but sulfur is lower in the periodic table (Row 3 vs. Row 2) so it is first in the name. There are two oxygens, so the multiplier is di and the name is sulfur dioxide.
For the molecule NO; nitrogen is to the left of oxygen (Group 6A vs. Group 5A) so it is first in the name. There is one oxygen, so the multiplier is mono and, following the rules, the name would be “nitrogen monooxide”. In this case, however, the second “o” in the name is dropped (to allow for easier pronunciation ) and the name is shortened to nitrogen monoxide. Distinguish this from another oxide of nitrogen, N$_2$O$_4$. Again nitrogen is first and needs the multiplier di. There are four oxygens, so the multiplier is tetra, but once again the multiplier is shortened (again, the “a” is dropped) and the name is dinitrogen tetroxide.
Example 3.7 Nomenclature of Molecular Compounds
Write a correct chemical formula for each of the following molecular compounds:
a. Chlorine monofluoride b. Dihydrogen monosulfile
c. Carbon tetrabromide d. Bromine
Example 3.7 Continued
Write a proper chemical name for each of the following molecular compounds:
e. IF f. PCl$_3$
g. I$_2$ h. N$_2$F$_2$
Solutions:
a. ClF b. H$_2$S
c. CBr$_4$ d. Br$_2$
e. Iodine monofluoride f. Phosphorus trichloride
g. Iodine h. Dinitrogen difluoride

## Study Points

In covalent bonding, electrons are shared between atoms. In ionic bonding, electrons are transferred from one atom to another. Compounds that are formed using only covalent bonds are termed molecular compounds.
The outermost electron level in any atom is referred to as the valence shell. The electron configuration of the valence shell of an atom can be shown graphically using a Lewis diagram (or electron-dot structure). The arrangement of “dots” around the chemical symbol for the element are shown singly up through four electrons, and then paired until eight electrons are present.
To form ions from individual elements, electrons are added or subtracted from the valence shell in order to completely fill the shell with eight electrons (the octet rule). The charge on the ion reflects the electrons added or removed. Representative elements through Group IIIA will lose electrons to form cations, while those in groups IVA – VIIA will gain electrons and form anions.
A covalent bond is constructed in a Lewis diagram by pairing a set of unpaired electrons from two different atoms. For the purposes of the “octet rule”, a pair of shared electrons is counted as two electrons for each atom. Multiple covalent bonds (double bonds and triple bonds) are used, if necessary, to give each bonded atom a full octet (except, of course, for helium and hydrogen). When two or more atoms are bonded together utilizing covalent bonds, the compound is referred to as a molecule.
As a rule of thumb, ionic bonds will be formed whenever the compound contains a metal. Covalent bonding will be observed in compounds containing only semimetals or nonmetals.
Groups of covalently bonded semimetals or nonmetals which are charged are called polyatomic ions. Common examples include sulfate dianion, nitrate anion, phosphate trianion, etc. These polyatomic ions are commonly paired with metals forming ionic compounds.
Many (but not all) polyatomic ions can be drawn in two or more equivalent Lewis representations. These are called resonance forms of the ion. The actual electronic structure of the ion is a combination of these Lewis structures and is called the resonance hybrid.
The electronegativity of an element is a measure of the tendency of that element to attract electrons towards itself. Electronegativities range from 0.6 to 4.0, with fluorine as the most electronegative element (a value of 4.0). The general trend in the periodic table is for electronegativity to increase from the lower left-hand corner (Fr) to the upper right-hand corner (F).
Covalent bonds formed between atoms with different electronegativity will be polarized with the greatest electron density localized around the most electronegative atom. The effect of electronegativity on electron distribution within a molecule can be shown using a computer-calculated electrostatic potential map where colors are used to represent electron density.
Elements in periods 3 – 7 can accommodate more that eight electrons in there valence shells. This phenomena is called valence shell expansion and molecules involving these elements may have 10 – 14 valence electrons in properly drawn Lewis diagrams. Exceptions to the “octet rule” also exist where the valence shell contains less than eight electrons, or contains unpaired electrons.
When naming simple, binary ionic compounds, the cation is named first using the name of the element, followed by the anion, where the suffix ide is added to the root name of the element. Multipliers are not used. For transition metals in which the metal can assume a variety of oxidation states (different positive charges), the charge of the metal ion is shown in the name using Roman numerals, in parenthesis, following the name of the element (i.e., iron (III) chloride).
When naming simple binary molecular compounds (compounds containing only covalent bonds) the least electronegative element is (generally) named first, followed by the second element, where the suffix ide is again added to the root name of the element. In molecular compounds multipliers are used to indicate the number of each atom present (mono-, di-, tri-, tetra-, etc.) with the exception that mono is not used for the first element in the compound.