4.3: Ionic Compounds

 

 

In this section we will build on the the ideas of the previous two sections to show how metals and nonmetals combine to make compounds, as well as to demonstrate the utility of the Periodic Table in predicting the products of such reactions. To recap, the following two ideas are critical:

1. Elements in the same column of the Periodic Table tend to behave similarly, forming products with other elements in the same ratio. In other words, if lithium forms a compound with element E in a ratio of 2:1, Li₂E, the other alkali metals probably will too, making Na₂E, K₂E, etc.;  
2. A complementary relationship exists between metals and nonmetals: the former lose electrons easily while the latter often take additional electrons. 

Point 2 above has an immediate consequence: metals and non-metals readily react to form ionic compounds, that is, compounds that consist of anions and cations. Specifically, nonmetals take electrons from metals, making the former anions and the latter cations. As solids the ions are usually held together tightly by electrostatic attractions called ionic bonds. One of the most familiar chemicals to everyone, one that is essential to life and has been traded since antiquity for use as a spice and preservative, is the product of such a complementary union between metals and nonmetals: common table salt, or sodium chloride. Sodium sits amongst the other alkali metals at the far left of the Periodic Table. You will also find it listed on nutrition labels, just beneath cholesterol (Figure 4-10). But there is no metal in potato chips, at least ideally. The sodium referred to on labels is derived from the same element as the metal listed on the Periodic Table, but it is not the pure element, but an ionic compound of it. Sodium chloride is formed in a violent reaction between metallic sodium and nonmetallic chlorine, a highly toxic yellowish green gas that was used as a poison gas in World War I. This illustrates a critical concept in chemistry: the properties of a given compound can radically differ from those of its component elements. 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Figure 4-10. A nutrition label for potato chips; notice that the metallic elements sodium and potassium are listed explicitly. Is there metal in your chips? (Image from a post on an unrelated health topic at Chemistry StackExchange: https://chemistry.stackexchange.com/...fat-as-healthy)

 

Sodium provides us with a good example of why metallic elements are actually quite difficult to find in nature. Pure sodium is a typical metal in some ways, being highly reflective when pure, an excellent conductor, etc., (although it would be a lousy metal to make things out of, taking malleability to an absurd level by being so soft that it can be cut with a plastic knife!), it is extreme in others, reacting vigorously (i.e., explosively!) with water (Figure 4-11). As you might expect for such a reactive element, it is not naturally found on Earth’s surface in any form other than the sodium cation, Na⁺. Why? It is so reactive, that any metallic sodium that may have been present when the Earth was formed has long since reacted with water or oxygen. All that is left are the ionic products. Sodium is usually found either in close proximity of anions in the form of solid ionic compounds like NaCl, or as dissolved ionic compounds in water. Ocean water, for example, is 3.5% NaCl by weight.

 

Figure 4-11: (top) Elemental sodium is a metal but is extremely soft; in the above image it is being cut with a common kitchen knife (image credit: "Sodium metal" by Thorius is licensed under CC BY-NC-SA 2.0.)        (right) Sodium metal is also exceptional reactive. The image at right captures the explosive nature of sodium's reaction with water. The light trails are due to individual pieces of burning sodium metal that are ejected from the reaction container. (image credit: "Sodium metal and water" by Thorius is licensed under CC BY-NC-SA 2.0.) Note: a good video that shows both of these aspects of sodium can be viewed here.

 

While most metallic elements are not as reactive as sodium (fortunately!), they are usually not found in nature as pure metals either but, like sodium, as cations that are associated with anions of some type; the tendency of pure metals to lose electrons in our terrestrial environment, usually to oxygen, is so pronounced that most metals in the Earth’s crust reacted long ago to form ionic compounds (Figure 4-12). The only notable exceptions are silver and gold, both of which can be found as pure metals, as numerous gold-rushes all over the world attest. They are sometimes called “inert metals” owing to their relative lack of reactivity compared to most other metallic elements. 

 

Figure 4-12. These rusting steel objects had the typical characteristics of metals when they were forged: they were malleable, ductile, and were probably shiny. Reaction with atmospheric oxygen (O₂), converts the iron atoms to Fe³⁺ ions in rust, changing their properties. Rust is not malleable but flaky, and has no metallic reflectivity. (Image source: "Rusting" by lamdogjunkie is licensed under CC BY 2.0.)

 

 

 

 

The formation of sodium chloride via direct reaction between the pure elements can be represented by the following balanced equation:

\[ \ce {2 Na (s) +  Cl2 (g)  -> 2 NaCl (s)} \]

As mentioned above, this reaction is exceptionally vigorous, accompanied by a sort of fire that does not require oxygen to support it (chlorine plays the role that oxygen normally plays in normal combustion). The intense heat and light evolved by this reaction is direct evidence that a substantial loss of potential energy results from the combination of the elements. Recall that energy can neither be created nor destroyed, so the energy released upon reaction of these two elements must have been contained in the elements themselves. As we saw previously in the formation of H₂, the potential energy of reactants is rooted in how their respective electrons are arranged. In that case, formation of the covalent H₂ bond decreased the potential energy of the isolated atoms in part because the negatively charged electrons moved closer to positive nuclei. In other words, by sharing their respective electrons, the two hydrogen atoms decrease their potential energy by allowing the negative charges to share two positive nuclei. The decrease in potential energy that occurs upon formation of NaCl is also due to the movement of electrons, albeit of a very different sort. Specifically, an electron is transferred from the metal (sodium) to the nonmetal (chlorine), generating a pair of oppositely charged ions that are very difficult to pull apart. Why does this happen? Because an electron that is relatively far away from to the sodium nucleus – it resides in sodium’s “sea of electrons” – moves to a location that is closer to chlorine’s nucleus, thereby decreasing its potential energy. In the reaction, the electrons from sodium disrupt the chlorine-chlorine bond, converting neutral Cl₂ molecules to chloride anions. The result is the formation of an ionic, crystalline product. As the balanced equation above indicates, two sodium atoms, each providing one electron, are necessary to convert the two chlorine atoms of Cl₂ to a pair of chloride ions.

 

At A Glance: Types of Chemical Bonds

Covalent: bonds consisting of electrons shared between atoms in the same molecule

Ionic: Electrostatic attraction between anions and cations

 

At the atomic level, the structure of sodium chloride is simple: it is a three-dimensional network of alternating positive and negative ions (see Figure 4-13). As such, individual molecules of NaCl do not normally exist  – they do not come as a molecular unit consisting of one sodium atom and one chlorine atom in the same way that molecular compounds such as H₂O exists or any of the examples of organic compounds we discussed in Chapter 1. In water, the three component atoms in a given H₂O molecule are tightly connected and are closer to each other than atoms in other water molecules. Even in ice, where water molecules are rigidly packed together, individual molecules are discernible owing to their component atoms' close proximity. In contrast, each chloride anion in NaCl is surrounded by six sodium cations, all an equal distance away; think of the chloride ion as being at the origin in an x, y, z, coordinate system and the sodium ions are placed at identical distances along the positive and negative directions of the three axes. Each cation is likewise surrounded by six anions. This picture should clarify why a molecular view of NaCl is incorrect: any given chloride ion interacts directly not with one, but six, sodium ions, and each sodium ion interacts with six chloride ions. Because ionic compounds are composed of lattices of ions, as opposed to discrete molecules, their formulas are more properly referred to as empirical formulas, which provide the simplest whole number ratio of the component ions in the compound: there are equal numbers of sodium cations and chloride ions, hence the empirical formula is simply NaCl. Empirical formulas are distinguished from molecular formulas in that the latter give the actual number of atoms for each element in a molecule of a given compound [14]; since molecules of NaCl do not exist, no there is no corresponding molecular formula.

 

Figure 4-13. (left) Crystals of NaCl have a distinct geometric pattern that also exists at the atomic level (right). The structure of sodium chloride consists of a three-dimensional array of alternating anions and cations, where each anion is equidistant to six cations and vice versa; in other words, there are no discernible molecules of NaCl, where a particular cation is very close to a particular anion, as would be expected in molecular compounds.

Figure 4-13 also provides some insights into the physical properties of ionic compounds.  We've mentioned several times that these compounds are often crystalline solids; this is a consequence of the geometric arrangement of their component ions. The exact crystal structure possessed by an ionic compound emerges from oppositely charged ions arranging themselves in a manner that minimizes their overall potential energy. How? There exists a tension in solids such as this, where the attractive forces of anions and cations is opposed by repulsive forces of like charged ions (anions repel other anions and cations repel other cations). In the spatial arrangement shown above, the distance between anions and cations is minimized, stabilizing the structure, while ions having the same charge are kept as far apart as possible, minimizing the destabilizing effect of the repulsive forces. Structures like this can be quite brittle, however, as mechanical stresses can shift the position of the ions, changing the nearest neighbor interactions from attractive to repulsive in nature, fracturing the crystal. In addition, because oppositely charged ions are difficult to physically separate, many ionic compounds have very high melting points. These characteristics of sodium chloride, namely its high melting points and crystalline structure, emerge from its ionic nature and not because it consists specifically of the element sodium and chlorine. As a consequence, these traits are frequently shared by other ionic compounds because such compounds also consist of large collections anions and cations geometrically arranged to minimize their potential energy, although the specific crystal structures adopted by ionic compounds depends on a variety of factors, including the ratio of anions to cations and their respective sizes.

Given the numbers of metallic and nonmetallic elements, you can probably imagine that there are a large number of possible cation/anion permutations. Indeed there are and, what's more, you can predict the formula for most of them quite easily. In the preceding discussion relating to the stoichiometry of the reaction between sodium and chlorine we laid out the logic of why two sodium atoms are required to react with one Cl₂ molecule: to maintain charge balance, equal numbers of Cl^- anions and Na⁺ cations are necessary. The need to maintain charge balance is the key conceptual insight in predicting the ratio of cations to anions in any ionic compound.  In other words, the formulas of other ionic compounds are a function of the respective charges on the component ions. For example, if the charge on the cation was +2, as is the case with the magnesium ion, Mg²⁺, then two chloride ions would be needed to balance the charge, making the formula MgCl₂. The charges on sodium and magnesium ions are +1 and +2, respectively, in virtually all of their compounds. This consistency is critical in being able to predict how many such ions will be in a given compound. But how would you know these particular charges? Is it a bit of chemical trivia that needs to be memorized? No. Ionic charges are easily determined by the Periodic Table. It is, in fact, the consistency of these ionic charges that gives rise to the repeating patterns that shapes to the Periodic Table. It is what Mendeleev was making use of in its development. You will note that sodium is in the first column, labeled 1A in Figure 4-2 (and shared again, below). The other metals in the same column also form cations with a +1 charge (called monocations). [15] The elements in column 2A  all form ions with +2 charges almost exclusively (they are called dications). Likewise, the metals in Column 3A tend to form ions with a +3 charge, although this tendency is not as absolute as we saw for Columns 1A and 2A.

To summarize: metals in Columns nA, where n is the column number, will form cations having a charge of +n, although exceptions become more commonplace if n ≥ 3. 

 

In the above paragraph, we referred to the “A” label on certain columns of the Periodic Table. Elements that are positioned in these columns are known as main group elements and include metals on the left-hand side of the Periodic Table and a combination of metals and nonmetals on the right. Also note in the Periodic Table shown here, that the columns of the main group elements are interrupted, between Column 2A and 3A, by ten columns bearing “B” designations; the elements in these columns are all metals and are collectively referred to as the transition metals; they include some of the more familiar metallic elements such as iron, copper, nickel and gold. Predicting the cationic charges of these elements is not as straightforward as with the main group elements. For reasons we will describe later, most cations of transition metals tend to have +2 or +3 charges. For example, chromium, in Column 6B, is often found as Cr²⁺ or Cr³⁺, as is iron and several others. The main exception to this is seen with the coinage metals: copper, silver and gold. They all reside in the same column (a coincidence?), Column 1B and, as you may suspect from the number of the column, easily form the monocations Cu⁺, Ag⁺ and Au⁺, although higher charges are not uncommon; copper is most often found as Cu²⁺ and gold can be Au³⁺. Thus while the Periodic Table summarizes much information concerning the elements, the patterns that emerge are not inviolate and have many subtleties. You need not memorize the charges of ions of the transition metals (or any elements!) but, for the time being, be aware that they seem to follow different "rules" than the main group elements (and this simply means that the rules we have articulated are not actually rules, but patterns that depend on multiple factors and, therefore, can be messier than we might like). [16]

What about nonmetals? What patterns exist regarding the charges of their ions? Good question. Let’s take a closer look at chlorine, sodium’s reaction-partner in the formation of table salt. Chlorine, a halogen, which sits in Column 7A, gains one electron in the formation of sodium chloride, forming the monoanion, Cl^-. If the patterns exhibited by the metals hold across the Periodic Table, then we would anticipate that the other elements in the same column as chlorine also form ions with a -1 charge. This would be correct, as these elements form the anions F-, Cl-, Br-, and I-. It is for this reason that alkali metals and halogens form ionic compounds in a 1:1 ratio: this is the required ratio of ions that maintains overall neutrality of the compound. Charge balance, in other words.

The example of chlorine shows that we can't use the column number to predict anionic charges as directly as we can with cations (it doesn't form an ion with a -7 charge!) but, nevertheless, it's not difficult as we'll see momentarily. We'll use oxygen to illustrate a logical approach to the question and then see what pattern emerges. Oxygen reacts with sodium according to the following equation:

\[ \ce {4 Na (s) + O2 (g)  ->  2 Na2O (s)} \]

Given what we’ve already covered about sodium, that is, it always forms the Na⁺ ion when reacting with nonmetals, we can see by inspection that the charge on the anion formed by oxygen must be -2 to balance the charge in the compound Na₂O. This is one example of a much larger trend: when oxygen forms an ion, it is usually the O^2-, or oxide, ion. As would be expected, other non-metals in the same column also form ions with a -2 charge. For example, the compound between sodium and sulfur is Na₂S.

Thus nonmetals in column 7A form anions with a charge of -1, those in column 6A form ions with a charge of -2. The general pattern is seems to be: when a non-metal reacts with a metal to form a binary ionic compound, it is converted to an ion with a charge of (n-8), or if you prefer, -(8-n), where n is the number of the column on the Periodic Table in which it resides. To put it another way, each atom of the non-metal will accept 8-n electrons from whatever metal it reacts with. Thus oxygen, in Group 6A, takes on two electrons, while fluorine, in Group 7A, takes on only one. This is borne out repeatedly, among different nonmetallic elements and across many of their compounds.

 

At a Glance: Ionic Charges

Ionic charges of main group metals: is usually equal to the column number (especially for Group 1A and 2A)

Ionic charges of nonmetals: equals n-8, where n is the nA column number.

 

 

Exercise

Problem 4.1. Elemental nitrogen, which exists as the diatomic species, N₂, is an exceptionally unreactive non-metal. It does, however, reluctantly react with metallic lithium to make a binary ionic compound. Given their respective positions on the Periodic Table, predict the empirical formula for the compound formed between these two elements, give its name, and write the balanced equation for its formation.

Solution. Lithium, in column 1A, makes the Li⁺ ion, while nitrogen, in Column 5A, makes N^3-, as discussed immediately above. To achieve charge balance we therefore need three lithium cations per nitride anion, hence the empirical formula of lithium nitride is Li₃N. The balanced equation for its formation is:

\[ \ce{3 Li (s) + ½ N2 (g)   →   Li3N (s)} \nonumber \]

Exercise

Problem 4.2. State whether or not binary ionic compounds could form if the following pairs of elements were to react and, if so, give the expected empirical formula:
    a) lithium and sulfur
    b) iron and nickle
    c) phosphorus and magnesium
    d) nitrogen and chlorine
    e) aluminum and fluorine
    f) barium and bromine

 

Thus far we have limited our discussion to monatomic ions, that is, those consisting of a single atoms that have a different number of electrons than protons. There are, however, many important ions that are composed of multiple atoms. These are best viewed as molecules that bear a non-zero charge because the sum of their electrons is different than the sum of their protons. The carbonate ion, for example, has the formula CO₃^2-. It has a total of 30 protons (24 from 3 oxygen atoms and 6 from the carbon) but 32 electrons, giving the entire molecule a -2 charge. Such charged molecules are called polyatomic ions, and play a huge role in biochemical and environmental systems. The acidity of your blood, to cite an important example, is remarkably stable over the course of your day (and life!) owing to the presence and reactions of polyatomic ions such as carbonate. Many minerals contain polyatomic ions such as carbonates and phosphates. If you pursue any work in a chemically-related field, you will encounter polyatomic ions so we list below a few particularly important examples. 

Table 4-3. Selected Polyatomic Ions
	Formula	Name
Common Polyatomic Cations	NH4	ammonium
	H3O^+	hydronium
Common Polyatomic Anions	OH^-	hydroxide
	CO3^2^-	carbonate
	HCO3^-	bicarbonate (or hydrogen carbonate)
	SO4^2^-	sulfate
	HSO4^-	bisulfate (or hydrogen sulfate)
	NO3^-	nitrate
	NO2^-	nitrite
	PO4^3^-	phosphate
	HPO4^2^-	hydrogen phosphate
	CH3CO2^-	acetate

 

You may wonder where these "extra" electrons come from. Unlike the direct reaction between elements of the type we saw in the formation of NaCl, polyatomic ions are usually formed via acid/base reactions. We discuss those in more detail in Chapter 5, so that question will be addressed in detail at that point. We include polyatomic ions at the moment because they have properties similar to those described above for ionic compounds consisting of monatomic ions. Calcium carbonate, the main component of blackboard chalk as well as seashells, is a solid with a very high melting point and is crystalline when pure.

We conclude this section with a brief explanation of naming ionic compounds; as you will see, it is considerably more straightforward than naming organic compounds.

 

 

Chemical Nomenclature: Ionic Compounds

As is true with organic compounds, there exists naming systems for inorganic compounds as well. We introduce below the naming system for ionic compounds, but note that a separate system exists for covalently bonded molecules that are inorganic in nature.

Ionic compounds are named by specifying the cation first and then the anion. Note that because charge balance must be followed, there is no need to indicate how many ions of each type are present in a compound because the ion names provide the charges, and the ratios of the ions is such that the compound is neutral overall. To name cations and anions, the following rules apply:

• For cations of Group 1A and Group 2A metals, the name of the cation is exactly the same as the name of the element. Thus the "sodium" in "sodium chloride" is understood to be the Na⁺ cation, not elemental sodium. Most other monatomic cations are from elements that can assume more than one ionic charge. Iron, for example, is usually either Fe²⁺ or Fe³⁺. In such cases, the charge on the ion is written explicitly in parentheses in roman numerals immediately after the element name. So, FeCl₂, which has Fe²⁺ ions (how do you know?) is named iron (II) chloride, while FeCl₃ is named iron (III) chloride. Polyatomic cations are referred to by their systematic names, which often have an -ium suffix appended to a base name.
• To name monatomic anions, the root of the element name is appended with an -ide suffix. That's why the chlorine ion in sodium chloride is called "chloride". As we've seen multiple times already, the O^2- ion is called oxide. Other examples include the commonly encountered anions sulfide (S^2-), nitride (N^3-), fluoride (F^-), bromide (Br^-) and iodide (I^-). Polyatomic anions are indicated by their systematic names, some of which are listed in Table 4-3. 

Problem 4-3. Name the following ionic compounds.
    a) Na₃P
    b) CaCl₂
    c) NH₄Cl
    d) LiCO₃
    e) CoBr₂
    f) Pb(NO₃)₂

Problem 4-4. Provide the formulas for the following ionic compounds.
    a) calcium fluoride
    b) rubidium oxide
    c) potassium hydroxide
    d) ammonium acetate
    e) magnesium sulfate
    f) tin(II) sulfide

 

 

 

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Notes and References.

[14] We already encountered empirical formulas in a very different context. recall, when determining the composition of compounds via combustion analysis the ratio of atoms you calculate is the empirical formula, not the molecular formula. For example, C₅H₁₀ and C₆H₁₂ both have the same empirical formula, CH₂. You also need the molar mass of compounds to determine their molecular formula if you know the empirical formula.

[15] Hydrogen is the only nonmetal in this column but, as we’ve already seen, it too forms a monocation, H⁺, readily. Interestingly, under extremely high pressure, hydrogen becomes metallic and shares many physical properties with the other alkali metals; the core of Jupiter is thought to consist primarily of metallic hydrogen.

[16] It is worth noting that even the main group metals do not follow simple “rules”. As we mentioned earlier, cationic charges of main group elements in Columns 3A-6A frequently deviate from that predicted by their column number, but these exceptions are interesting in that they too exhibit a regular pattern. Specifically, if the ion does not have a charge of +n, it usually has a charge of +(n-2) instead. Thus thallium, a highly toxic metal in Column 3A, made (in)famous in the movie The Young Poisoner’s Handbook, easily forms Tl⁺ ions, and lead, in Column 4A, is often either Pb⁴⁺ or Pb²⁺. Aluminum, on the other hand, forms the Al³⁺ ion exclusively.