Stable Atoms and Ionic Bonding – Formation of Ionic and Metallic Bonds (Senior 4 Chemistry) | Ganzaa

Formation of Ionic and Metallic Bonds

Lessons

Senior 4 · Chemistry

Formation of Ionic and Metallic Bonds

3 lessons

Lesson 1 of 3

Stable Atoms and Ionic Bonding

25 min

A soft metal, a poison gas and table salt

Sodium is a soft, silvery metal that fizzes on water and has to be stored under oil. Chlorine is a greenish-yellow gas that attacks the lungs. Lower a piece of warm sodium into a jar of chlorine and it burns with a bright yellow flame, leaving a white solid: sodium chloride, the salt in your food. Its properties are nothing like either element. To understand why, you need to know what happens to the electrons when the two react, and what holds the product together. That is the subject of this unit, which uses the electron configurations from Unit 2.

Stable configurations

The noble gases, helium to radon, exist as single atoms and hardly react at all. Their outer shells are full: $1 s^{2}$ for helium (a duet) and $ns^2\,np^6$ for the others (an octet). Other atoms can reach the same kind of configuration by losing, gaining or sharing electrons. Metals on the left of the periodic table have only one to three outer electrons and low ionisation energies, so they tend to lose them. Non-metals on the right are one to three electrons short of an octet and tend to gain or share.

Be careful with the idea of "stability". Removing an electron from a sodium atom costs 496 kJ per mole; a lone $\mathrm{Na^{+}}$ ion is not lower in energy than a sodium atom. What makes ionic compounds stable is the energy released when the ions attract each other, as Worked example 1 shows.

Definition

Octet rule: atoms tend to lose, gain or share electrons until they have eight electrons in their outer shell (two for hydrogen, lithium and beryllium ions: the helium duet).

Ionic (electrovalent) bond: the electrostatic attraction between oppositely charged ions, formed when one or more electrons are transferred from a metal atom to a non-metal atom. A positive ion is a cation; a negative ion is an anion.

First electron affinity: the energy change when one mole of gaseous atoms each gains one electron to form one mole of gaseous 1− ions. $\mathrm{Cl(g)} + e^- \rightarrow \mathrm{Cl^{-}(g)}$ , $\Delta H = -349\,\mathrm{kJ\,mol^{-1}}$ .

Dot-and-cross diagrams

A dot-and-cross diagram shows where each electron came from: crosses for the metal's electrons and dots for the non-metal's. For ions, the outer shell is usually enough. Each ion is drawn in square brackets with its charge at the top right.

In sodium chloride, sodium's single 3s electron moves to chlorine's 3p subshell. $\mathrm{Na^{+}}$ ([Ne]) and $\mathrm{Cl^{-}}$ ([Ar]) both have noble-gas configurations. Magnesium has two outer electrons and each chlorine atom needs one, so two chloride ions form for every magnesium ion: $\mathrm{MgCl_2}$ .

Worked example 1: does making ions release energy?

Use the data to find the energy change for $\mathrm{Na(g)} + \mathrm{Cl(g)} \rightarrow \mathrm{Na^{+}(g)} + \mathrm{Cl^{-}(g)}$ . Then explain why sodium chloride forms. First ionisation energy of Na = $+496\,\mathrm{kJ\,mol^{-1}}$ ; first electron affinity of Cl = $-349\,\mathrm{kJ\,mol^{-1}}$ ; energy released when gaseous $\mathrm{Na^{+}}$ and $\mathrm{Cl^{-}}$ ions form one mole of solid NaCl = $787\,\mathrm{kJ\,mol^{-1}}$ .

Solution.

\Delta H = (+496) + (-349) = +147\,\mathrm{kJ\,mol^{-1}}

Making the separate gaseous ions absorbs 147 kJ per mole. But when the ions come together into a solid lattice, 787 kJ per mole is released, far more than 147.

Answer: $+147\,\mathrm{kJ\,mol^{-1}}$ . Sodium chloride forms because of the strong attraction between the ions in the lattice, not because $\mathrm{Na^{+}}$ and $\mathrm{Cl^{-}}$ are stable on their own.

When does an ionic bond form?

Table 1.1: Conditions that favour ionic bonding

Condition	Why it helps	Typical elements
The metal has low ionisation energies	little energy is needed to form the cation	Groups 1 and 2
The non-metal has an exothermic electron affinity	energy is released when the anion forms	Groups 16 and 17
A large difference in electronegativity	one atom holds electrons much more strongly than the other, so the electrons transfer rather than being shared	about 1.8 or more on the Pauling scale (Unit 4)
A large lattice energy	small, highly charged ions attract strongly and release a lot of energy	$\mathrm{Mg^{2+}}$ with $\mathrm{O^{2-}}$

The last condition explains a surprise. Adding a second electron to $\mathrm{O^{-}}$ absorbs energy ( $+798\,\mathrm{kJ\,mol^{-1}}$ ), because the electron is repelled by a negative ion. Oxide ions still form in compounds such as MgO, because the lattice of small $\mathrm{Mg^{2+}}$ and $\mathrm{O^{2-}}$ ions releases an exceptionally large amount of energy.

Worked example 2: aluminium oxide

Aluminium (Z = 13) reacts with oxygen (Z = 8). Write the configurations of the two ions formed and deduce the formula of the compound.

Solution. Al is $1s^2\,2s^2\,2p^6\,3s^2\,3p^1$ : it loses its 3 outer electrons to form $\mathrm{Al^{3+}}$ , $1s^2\,2s^2\,2p^6$ . O is $1s^2\,2s^2\,2p^4$ : it gains 2 electrons to form $\mathrm{O^{2-}}$ , $1s^2\,2s^2\,2p^6$ . Both are isoelectronic with neon.

The charges must balance. The lowest common multiple of 3 and 2 is 6: two $\mathrm{Al^{3+}}$ give $+ 6$ and three $\mathrm{O^{2-}}$ give $- 6$ .

Answer: $\mathrm{Al_2O_3}$ . Six electrons are transferred: two aluminium atoms lose three each, and three oxygen atoms gain two each.

Warning

Common mistakes. (1) Drawing shared pairs between the ions: in an ionic diagram, no electrons are shared. (2) Leaving out the square brackets or the charges. (3) Writing "sodium wants to lose an electron because it is stable": forming $\mathrm{Na^{+}}$ needs energy; the lattice energy makes the overall change favourable. (4) Forgetting that one dot-and-cross pair shows only one formula unit: solid NaCl is a giant lattice, not a molecule.

Safety. Sodium burning in chlorine is a teacher demonstration only, done in a fume cupboard behind a safety screen: chlorine is toxic, and sodium reacts violently with water and can ignite.

Tip

To get the formula of an ionic compound, cross over the sizes of the charges and simplify: $\mathrm{Al^{3+}}$ and $\mathrm{O^{2-}}$ give $\mathrm{Al_2O_3}$ ; $\mathrm{Ca^{2+}}$ and $\mathrm{O^{2-}}$ give $\mathrm{Ca_2O_2}$ , which simplifies to CaO.

Allery Chemistry: Ionic bonding (11 min). The video covers what ions are, how they form, ionic equations and the physical properties of ionic compounds. Watch the first part for how electrons are transferred, and keep the rest for the next lesson.

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Practice

Explain, in terms of electron configurations, why neon is unreactive.
Draw dot-and-cross diagrams (outer shells only) for: (a) potassium fluoride, (b) calcium oxide, (c) sodium sulfide.
Write the formulae of the ionic compounds formed between: lithium and nitrogen; calcium and fluorine; aluminium and chlorine; magnesium and nitrogen.
Give the full electron configurations of the ions in calcium chloride. Which noble gas are they isoelectronic with?
The first ionisation energy of potassium is $+419\,\mathrm{kJ\,mol^{-1}}$ and the first electron affinity of bromine is $-325\,\mathrm{kJ\,mol^{-1}}$ . Calculate the energy change for $\mathrm{K(g)} + \mathrm{Br(g)} \rightarrow \mathrm{K^{+}(g)} + \mathrm{Br^{-}(g)}$ . Comment on the sign.
Suggest why magnesium forms $\mathrm{Mg^{2+}}$ in its compounds and not $\mathrm{Mg^{3+}}$ .

Show answersHide answers

Neon is $1s^2\,2s^2\,2p^6$ : its outer shell is full, its ionisation energy is very high and an extra electron would have to go into a new shell, so it has no tendency to lose, gain or share electrons.
(a) $[\mathrm{K}]^{+}$ with 8 crosses (3s, 3p) and $[\mathrm{F}]^{-}$ with 7 dots and 1 cross. (b) $[\mathrm{Ca}]^{2+}$ with 8 crosses and $[\mathrm{O}]^{2-}$ with 6 dots and 2 crosses. (c) Two $[\mathrm{Na}]^{+}$ ions, each with 8 crosses, and one $[\mathrm{S}]^{2-}$ with 6 dots and 2 crosses.
$\mathrm{Li_3N}$ ; $\mathrm{CaF_2}$ ; $\mathrm{AlCl_3}$ (a simple ionic picture; Unit 4 shows it is largely covalent); $\mathrm{Mg_3N_2}$ .
$\mathrm{Ca^{2+}}$ and $\mathrm{Cl^{-}}$ are both $1s^2\,2s^2\,2p^6\,3s^2\,3p^6$ , isoelectronic with argon.
$+419 - 325 = +94\,\mathrm{kJ\,mol^{-1}}$ . It is positive: forming the separate gaseous ions needs energy; the compound forms because the lattice energy released is much larger.
The third electron would have to be removed from the full 2p subshell, much closer to the nucleus and far less shielded, so the third ionisation energy is very large (7733 kJ mol⁻¹). The extra lattice energy of an $\mathrm{Mg^{3+}}$ compound could not repay it.

Quick check

Which pair of elements is most likely to form an ionic compound?

Key points

Noble gases have full outer shells; other atoms lose, gain or share electrons to reach similar configurations.
An ionic bond is the electrostatic attraction between cations and anions formed by electron transfer from a metal to a non-metal.
Dot-and-cross diagrams: ions in square brackets, charges outside, no shared electrons.
Forming gaseous ions usually absorbs energy; ionic compounds are stable because of the large lattice energy released when the ions come together.

Sources and credits

Image: Halite crystals (four).jpg by James St. John, via Wikimedia Commons, licensed under CC BY 2.0.
Ionisation energies, electron affinities and the lattice energy of NaCl are standard reference values, as listed in the CRC Handbook of Chemistry and Physics.
Video: Ionic bonding, Allery Chemistry, embedded from YouTube.
The text, dot-and-cross diagrams, table, worked examples and Practice items were written and drawn by Ganzaa.

1A soft metal, a poison gas and table salt

2Stable configurations

3Dot-and-cross diagrams

4When does an ionic bond form?

A soft metal, a poison gas and table salt

Stable configurations

Dot-and-cross diagrams

When does an ionic bond form?