Ionic Solids and Lattice Energy

In this page, we'll discuss the attractive forces between ions in an ionic solid, as well as talk about something we've left off for a while, how to name ionic substances.

We briefly touched on a concept called lattice energy back in the "States of Matter" chapter when discussing crystal structures. To refresh your memory, the lattice energy is the amount of energy released when an ionic crystal forms, in kilojoules per mole, with a negative value signifying a release of energy and a product-favored reaction. It would seem that the more then anion (negative ion) wants to gain electrons, as indicated by a very negative electronegativity or electron affinity, and the more the cation (negative ion) wants to lose its electrons, as evidenced by a low ionization energy, the more negative the lattice energy would be. However, the cations with the lowest ionization energy (rubidium, cesium, and francium, for example) are also very large atoms. They take up a lot of room in the crystal structure, meaning that the component ions cannot get as close together, so the ionic bond is actually weaker. For this reason, ionic solids like LiF and NaF have higher lattice energies than ones such as KF or RbF. Larger, less electronegative cations (such as I or At) also make the lattice energy more positive.

Using Hess' Law, a lattice energy can be estimated for an ionic solid provided you know five factors: the enthalpies of formation of gases for both reactants (representing the energy needed to convert the elemental reactants, in whatever form, to single-atom gases for reaction), the ionization energy of the cation (to create a positive ion), the electron affinity of the anion (to create a negative ion), and the enthaly of formation of the solid ionic solid from its gaseous state (representing the energy released when the gaseous product is reconverted to a solid). We'll use that ever-popular example, NaCl:

1: Na(s) Na(g)    107 kJ/mol
2: 1/2 Cl₂(g) Cl(g)    122 kJ/mol
3: Na(g) Na⁺ + e^-    496 kJ/mol
4: Cl(g) + e^- Cl^-(g)    -349 kJ/mol
5: NaCl(g) NaCl(s)    -786 kJ/mol

Sum: Na(s) + 1/2 Cl₂(g) NaCl(s)     -410 kJ/mol

The "e^- in the 4th and 5th steps represents the electron transferred between the sodium and chlorine atoms. You'll see a lot more of the electron symbol in the "Redox Reactions" chapter.

Knowing what to call all these ionic solids is probably another useful skill, but first we need to discuss the transition metals and how they form ions. While all transition metals lose electrons to become cations, some form more than one ion, which seems to disagree with the octet rule. The discrepancy is explained by the presence of d orbitals, the filling of which can impact the ions formed. Iron, for example, forms both 2+ and 3+ ions; lead forms 2+ and 4+ ions. Let's look at lead's electron configuration:

[Xe]6s²4f¹⁴5d¹⁰6p²

In lead, the 6s, 4f, and 5d subshells are completely filled, while the 6p shell is only partly full. The removal of two electrons from the 6p subshell explains the +2 ion, and the +4 is produced by taking two more from the 6s subshell. The 4f subshell has so many electrons that it is very stable, making ionization difficult.

Predicting the ions formed by transition metals is difficult, and it is easier to simply memorize the more common ones: iron form 2+ and 3+ ions; and titanium forms 2+ and 4+ ions, as does lead; copper forms 1+ and 2+ ions; and tin also forms 2+ and 4+ ions. These are the most important polyionic transition metals, but several more are occasionally encountered (including gold, cobalt, bismuth, and manganese). Note that mercury can form Hg⁺, Hg²⁺, and the unusual Hg₂²⁺ ion.

For ionic solids formed by metals in which cation does not form multiple ions and simple nonmetals (not polyatomic ions, which we will discuss in a moment), the name is simple. The name of the cation is simply its element, but the anion takes the name of the ion (usually an "ide" replaces the "ium" or "ine" of the element). No mention of the number of anions or cations is used. For example, NaCl is sodium chloride, MgF₂ is magnesium fluoride, H₂S is hydrogen sulfide, K₂O is potassium oxide, and CaI₂ is calcium iodide.

Ionic solids in which the cation forms multiple ions are written in one of two ways. The first (the preferred method) gives the postive charge in Roman numberals in parentheses after the cation. For instance, FeO is iron(II) oxide, while Fe₂O₃ is iron(III) oxide; CuCl₂ is copper(II) chloride, but CuBr is copper(I) bromide.

Another less-common method uses special names for each different ion. The specialized names are shown below:

Ion:	Cu+	Cu2+	Fe2+	Fe3+	Hg+	Hg22+	Hg2+	Pb2+	Pb4+	Sn2+	Sn4+
Name:	Cuprous	Cupric	Ferrous	Ferric	Mercurous	Mercurous	Mercuric	Plumbous	Plumbic	Stannous	Stannic

The other metals which form multiple charges have no alternate names, so the Roman-numeral system must be used. For example, SnBr₄ may be written as tin(IV) bromide or as stannic chloride; HgO may be mercury(II) oxide or mercuric oxide, while HgCl is mercury(I) chloride or mercurous chloride.

In molecules like CaCO₃, calcium carbonate, we know the cation is calcium, but what do we call the polyatomic anion? Fortunately, there is a system for naming polyatomic ions with oxygen in them.

We start with the "most common" oxygen-containing ion; usually, it has four or three oxygens, with the formula XO₄ or XO₃. The ion is given the name of the "X" element, replacing the "ium" or "ine" with "ate." This ion often has a 2- charge, but some have 3- or 1- charges as well. For example, SO₄^2- is the sulfate ion, NO₃^- is the nitrate ion, ClO₃^- is the chlorate ion, CrO₄^2- is the chromate ion, and PO₄^3- is the phosphorate ion.

The number of oxygens in the ion may vary; however, the charge on the ion does not. One more oxygen than normal adds a "per" to the front of the ion. For example, ClO₃^- is the chlorate ion, while ClO₄^- is the perchlorate ion; MnO₃^- is the manganate ion, while MnO₄^- makes this the permanganate ion.

Subtracting one oxygen changes the "ate" suffix to "ite." PO₄^3- is the phosphorate ion, but with one less oxygen (PO₃^3-) it becomes the phosphorite ion. Removing an oxygen from the nitrate ion, NO₃^-, makes it the nitrite ion, NO₂^-.

Finally, removing one more oxygen from an ion retains the "ite" suffix but adds a "hypo" to the front of the name. For example, ClO₃^- is the chlorate ion, ClO₂^- becomes the chlorite ion, and ClO^- is the hypochlorite ion. PO₄^3- is the phosphorate ion, but with two less oxygen atoms, it is PO₂^3-, the hypophosphorite ion.

Those rules above take care of most polyatomic ions, but a few others should be mentioned:
NH₄⁺, the ammonium ion
HCO₃^-, the bicarbonate ion
CN^-, the cyanide ion
OCN^-, the cyanate ion
Cr₂O₇^2-, the dichromate ion
OH^-, the hydroxide ion
S₂O₃^2-, the thiosulfate ion

For example, using the definitions and rules above, Fe₃(PO₃)₂ is either iron(II) phosphorite or ferrous phosphorite. NH₄ClO₄ is ammonium perchlorate. Cu(NO₂)₂ is either copper(I) nitrite or cuprous nitrite. Al(CN)₄ is aluminum(IV) cyanide (don't eat that one!).

This wraps up our discussion of ionic solids. Next, we will revisit covalent bonds and how to predict what bonds are formed in covalent substances.

| |

home site map/index background information chemical reactions states of matter acid/base reactions

thermodynamics redox reactions atoms and molecules organic chemistry nuclear chemistry

help features reference thinkquest