Ions and Bonding
Chemistry Fundamentals
Course 1 · Chapter 4 · Ions and Ionic Bonding: Atoms That Carry a Charge
Table salt is sodium chloride, and it is made of two elements that on their own are very different: Chapter 3 described sodium as a soft, highly reactive alkali metal and chlorine as a halogen gas. Wikipedia's Ionic bonding article defines the link between them: "Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions." This chapter explains what an ion is, how atoms become ions by gaining or losing electrons, why the metals and non-metals of Chapter 3 tend to go different ways, how to write a formula by balancing charges, and what ionic compounds are like. It builds on Chapter 2, where an atom's protons and electrons were counted, and it is the first chapter to use electrons to explain behaviour, a promise made at the end of Chapter 3.
What an Ion Is
According to Wikipedia's Ion article, an ion is "an atom or molecule with a net electrical charge". The charge appears because the counts of protons and electrons no longer match: "The net charge of an ion is not zero because its total number of electrons is unequal to its total number of protons." In Chapter 2 a neutral atom had equal numbers of the two, so an ion is what you get when that balance is broken by gaining or losing electrons. The protons stay put; only electrons move.
Cation
The article defines a cation as "a positively charged ion with fewer electrons than protons". Losing electrons makes an atom positive.
Anion
An anion is "a negatively charged ion with more electrons than protons". Gaining electrons makes an atom negative.
Writing the charge
The article says singly charged ions are marked with just a + or a −, such as Na+ and F−. A charge of two or three is written with the number, as in 3+ and 2− later in this chapter.
Valence Electrons: Which Electrons Move
Not all of an atom's electrons take part. Wikipedia's Valence electron article defines them: "valence electrons are electrons in the outermost shell of an atom", and says they can take part in bonding "if the outermost shell is not closed". It also gives a rule for reading the count from the periodic table: "In groups 1-12, the group number matches the number of valence electrons", and for groups 13 to 18 "the units digit of the group number matches the number of valence electrons", with helium as the exception. So by that rule sodium (group 1) has one valence electron and chlorine (group 17) has seven. That is the electron explanation of group behaviour that Chapter 3 put off: atoms in one column have the same number of outer electrons, so they behave alike.
Why do atoms gain or lose them? The Octet rule article describes "the strong tendency for main-group elements to bond in such a way that each atom has eight electrons" in their outer shell. The Valence electron article connects this to the noble gases of Chapter 3: an atom with a closed outer shell tends to be chemically inert, which fits the low reactivity noted for group 18. Sodium "easily" loses its single valence electron "to form a positive ion", while chlorine "requires only one additional valence electron to form a closed shell". Shells beyond this point (what a shell holds, subshells, orbitals) are left for a later look; here the rule of thumb is enough.
Sodium Chloride: The Worked Example
The Ionic bonding article walks through this case: "the sodium atoms each lose an electron, forming cations (Na+)", and chlorine atoms gain those electrons and become anions, Cl−. The Octet rule article says it directly: "Sodium is observed to transfer one electron to chlorine during the formation of sodium chloride", and "Adding an electron to chlorine produces the chloride ion (Cl−)". The Valence electron article agrees that in solid sodium chloride the sodium valence electron "is transferred to chlorine to form an ionic bond". The two ions attract in a 1:1 ratio, giving NaCl.
| Step | Sodium | Chlorine |
|---|---|---|
| Group (Chapter 3) | 1, an alkali metal | 17, a halogen |
| Valence electrons (group rule) | 1 | 7 |
| What happens | loses 1 electron | gains 1 electron |
| Result | Na+ (cation) | Cl− (anion) |
Counting protons and electrons (Chapter 2) shows the charge. Sodium's atomic number is 11 and chlorine's is 17 (check, these are from general knowledge, not a page I read in this chapter). A neutral sodium atom has 11 protons and 11 electrons; after losing one, 11 protons and 10 electrons remain, and 11 - 10 = +1. A chlorine atom has 17 and 17, and after gaining one it has 17 protons and 18 electrons, so the net charge is 17 - 18 = -1. (The subtractions 11 - 1 = 10 and 17 + 1 = 18 were run in Python.) These counts match the Na+ and Cl− charges given by the pages. The pages do not themselves give the atomic numbers, so they are shown as unverified.
Why Metals Form Cations and Non-metals Form Anions
The Ion article says it plainly: "metals will lose electrons to form positively charged ions", while non-metals gain electrons to form negatively charged ones. The Ionic bonding article adds that "in the simplest case, the cation is a metal atom and the anion is a nonmetal atom". This links to Chapter 3: metals lie to the left and centre of the table, non-metals toward the right, and the halogens just before the noble gases. Chlorine, one electron short of a full outer shell, takes one. Sodium, one electron beyond a closed shell, gives one away. The pages do not explain in energy terms why sodium finds this easier than chlorine does; that is a deeper question left alone here.
Ionic Bonding as Attraction
The bond itself is the pull between opposite charges. The Ion article says "cations and anions attract each other and readily form ionic compounds", and the Ionic bonding article calls the bond "the electrostatic attraction between oppositely charged ions". The Ionic compound article says the ions are held together "by electrostatic forces called ionic bonds". Electrons are not shared here; they are handed over. Chapter 5 turns to the other main case, where atoms share electrons.
Writing a Formula from Charges
The Ionic bonding article gives the rule: "to maintain charge neutrality, strict ratios between anions and cations are observed." So a formula is found by choosing the smallest numbers of each ion that make the total charge zero. For sodium chloride, +1 and -1 cancel at 1:1, giving NaCl, as the Sodium chloride article's formula and "1:1 ratio" show.
A case with unequal charges comes from the Ionic compound article, which says that in Fe2(SO4)3 each iron ion has a 3+ charge, balancing sulfate ions at 2− each. Check the balance: two iron ions give 2 x 3 = +6, three sulfate ions give 3 x 2 = -6, and the total is 6 - 6 = 0 (all run in Python). Notice how the subscripts come from the charges: the number of each ion needed is found by making the totals equal, and the sulfate ion is written in brackets because it is a group of atoms carrying the charge together. The article describes such ions as polyatomic.
| Compound | Ions | Charge check |
|---|---|---|
| Sodium chloride, NaCl | Na+ and Cl− | (+1) + (-1) = 0 |
| Iron sulfate, Fe2(SO4)3 | Fe 3+ and SO4 2− | 2 x (+3) + 3 x (-2) = 0 |
Properties of Ionic Compounds
Several properties follow from the picture of charged ions held by strong attraction. The pages state the following.
| Property | What the page says | Source |
|---|---|---|
| Melting point | "Ionic compounds generally have a high melting point", rising with the ion charges. Sodium chloride melts at 801 °C (the infobox lists 800.7 °C). | Ionic bonding; Sodium chloride |
| Boiling point | Salts of small ions "typically have high melting and boiling points". Sodium chloride is given as 1465 °C in the text but 1,413 °C in the infobox, so the page disagrees with itself by 52 °C (Python). | Ionic compound; Sodium chloride |
| Electrical conduction | "Ionic compounds conduct electricity when molten or in solution, typically not when solid." The Ionic compound page says solids are "almost always electrically insulating" and become "highly conductive" when melted or dissolved, because the ions become mobile. | Ionic bonding; Ionic compound |
| Brittleness | "Most salts are very brittle." | Ionic compound |
| Dissolving | "When simple salts dissolve, they dissociate into individual ions." Solubility is highest in polar solvents such as water. | Ionic compound |
The conduction rule is a good way to remember what is going on. Electric current through a liquid is carried by charged particles that are free to move. In the solid the ions are locked in place, so there is no current; melting or dissolving lets them move. Spelling out that reason is my own wording (check), though the pages give the same cause for the Ionic compound page's conductivity statement. The pages do not say why brittleness occurs, so this chapter does not explain it.
A Lattice, Not a Molecule
A water molecule is a small, countable unit. Ionic compounds are different. The Ionic bonding article says ionic solids form extended lattices, so "it is usually not possible to distinguish discrete molecular units". The Ionic compound article calls the ions "part of a continuous three-dimensional network". The Sodium chloride article gives the real arrangement: it "can be represented as a face-centered cubic (fcc) lattice with a two-atom basis", and "each ion is surrounded by six ions of the opposite charge". So a grain of salt is one huge array in which every Na+ sits among six Cl− and every Cl− among six Na+, not a heap of separate NaCl molecules.
That is why NaCl is called a formula, not a molecular formula: it gives the simplest ratio of ions. The Ionic compound article uses the phrase "formula unit" for one such group, as in "the two iron ions in each formula unit" of Fe2(SO4)3. I did not read how that article defines the term, so treat "formula unit" as the smallest repeating ratio (check).
Where to Read More
- Encyclopaedias. The Wikipedia articles Ion, Ionic bonding, Ionic compound, Valence electron, Octet rule and Sodium chloride are the sources used here. See the note above for how much of each was read.
- A school textbook. Look for a table of common ions and their charges, which this chapter does not reproduce. Ion also has a table of common anions that I did not quote.
- Practical. Dissolve a little salt in a glass of water and look up how a conductivity test would differ from pure water, then compare with the page's statement on solutions.
Exercises
Predict an ion from position. Using the group rule from the Valence electron article, give the number of valence electrons of potassium (group 1) and fluorine (group 17). Say whether each is more likely to lose or gain electrons, how many, and write the ion with its charge. State which is the cation and which the anion, and how the number of electrons compares with the number of protons in each ion (use a general atomic number of your own and mark it as such).
Balance a formula from charges. (a) Using the Fe 3+ and sulfate 2− charges given in this chapter, show that Fe2(SO4)3 has zero total charge. (b) For this exercise only, suppose an ion X has charge 2+ and an ion Y has charge 1−. Find the formula of the compound by balancing the charges, and show the check. (c) Suppose instead X has charge 3+ and Y has charge 2−: find the formula and check it.
Decide from properties. Substance P is a hard, brittle solid that melts at a high temperature. It does not conduct electricity as a solid, but it does when melted and when dissolved in water. Substance Q is a soft solid that conducts electricity as a solid. Which, if either, matches the pages' description of an ionic compound, and which statements support the decision? Then say why a lattice picture, not a molecule picture, suits your answer, and what the pages do not tell you about substance P.
Quick Reference
- Ion: an atom or molecule with a net electrical charge; protons and electrons no longer match
- Cation / anion: positive (fewer electrons than protons) / negative (more electrons than protons)
- Valence electrons: outer-shell electrons; for groups 1 to 12 the group number, for groups 13 to 18 the units digit (helium excepted)
- Octet rule: main-group atoms tend to bond so each has eight outer electrons
- Sodium chloride: Na loses one electron to become Na+, Cl gains one to become Cl−, 1:1 ratio, NaCl
- Who forms what: metals lose electrons to form cations; non-metals gain electrons to form anions
- Bond: electrostatic attraction between oppositely charged ions
- Formula: choose ion numbers so the total charge is zero, as in Fe2(SO4)3: 2 x 3 = 6 and 3 x 2 = 6
- Properties: high melting points, brittle, conduct when molten or dissolved but typically not as solids
- Structure: a three-dimensional lattice; in sodium chloride each ion has six neighbours of opposite charge
- Check: the atomic numbers of sodium and chlorine; the reading of "formula unit"; my own memory aids