Chemical Bonding and Molecular Structure · Lesson 3 of 12
Ionic Bonds and Crystal Stability
“Follow electron transfer into a three-dimensional crystal and explain why lattice formation stabilises ionic solids.”
• Explain how ionisation and electron gain produce cations and anions. • Identify energetic conditions that favour ionic-bond formation. • Describe an ionic solid as a three-dimensional lattice rather than isolated molecule pairs. • Define lattice enthalpy and distinguish dissociation from lattice formation.
Electron transfer is only the beginning
The sodium–chlorine example is useful, but it leaves an important question. Removing an electron from sodium requires energy. Why should the overall process lead to a stable compound? To answer, follow the energetic changes and include the many attractions present in the finished crystal.
Forming an ionic compound involves both the ease of producing ions and the way those ions arrange in the solid. The stable product is an extended structure. Completing an octet on isolated gaseous ions is therefore not enough to explain the stability of the whole solid.
The energy needed to make ions
A gaseous metal atom becomes a cation when one or more electrons are removed. Removing an electron from an isolated gaseous atom always requires an input of energy because the electron is attracted to the positive nucleus. This energy change is called ionisation enthalpy.
A gaseous atom becomes an anion when it gains an electron. Its electron-gain enthalpy is the enthalpy change for adding that electron to the isolated gaseous atom in its ground state. Electron gain can release energy or require energy; it is not automatically favourable for every atom or every successive electron.
Low ionisation enthalpy makes cation formation easier. A large negative electron-gain enthalpy makes the first electron-gain step more favourable. These tendencies help explain why many ionic compounds combine metallic cations with nonmetallic anions. They are tendencies within a complete energy balance, not a guarantee that any chosen pair of atoms will form a stable salt.
Problem
Suppose forming a gaseous cation requires 500 kJ per mole and electron gain releases 350 kJ per mole. Is the combined ion-production step favourable by itself?
- 1.Write energy required as positive and energy released as negative: and .
- 2.Add the two changes: . This pair of steps requires a net input.
- 3.This does not settle the stability of a solid ionic compound. Energy released as the ions assemble into a lattice must still be considered.
What the crystal adds
Positive and negative ions attract through electrostatic forces. In a crystal, each ion interacts with many surrounding ions: opposite charges attract and like charges repel. A stable lattice arrangement produces a favourable overall balance of these interactions.
The rock-salt structure of sodium chloride extends in three dimensions. Each sodium ion has six nearest-neighbour chloride ions and each chloride ion has six nearest-neighbour sodium ions. The formula NaCl gives the simplest 1:1 ratio of ions, rather than identifying an isolated NaCl molecule inside the crystal.
The actual crystal structure depends on ionic sizes, packing, and the geometry of the entire arrangement. One cannot calculate the crystal energy accurately by adding only the attraction between one neighbouring pair. Repulsions and interactions with more distant ions also contribute.
Problem
Use the chapter’s gaseous-ion data to explain why ion formation alone is incomplete.
- 1.Removing an electron from gaseous sodium requires about . Adding that electron to gaseous chlorine releases about .
- 2.The net of these two steps is . Producing separated gaseous ions from the gaseous atoms is therefore energetically uphill.
- 3.Bringing the gaseous ions into the NaCl lattice releases about . Combining this with the two ion-production steps gives .
- 4.This comparison starts with gaseous atoms. It is not the complete enthalpy of forming solid NaCl directly from sodium metal and chlorine gas, because preparing the gaseous atoms also involves energy changes.
The ionisation-plus-electron-gain calculation starts with gaseous atoms. A reaction beginning with an elemental solid and a diatomic gas has additional steps. State the starting and finishing states whenever interpreting an energy value.
Lattice enthalpy and its sign
Two opposite processes are useful: separating a crystal into gaseous ions, and forming a crystal from those gaseous ions. They involve the same pair of states, so their enthalpy changes have equal magnitudes and opposite signs. This chapter defines lattice enthalpy using the separation process.
The energy required to separate one mole of an ionic solid completely into its constituent gaseous ions.
Reverse this process and the enthalpy change becomes approximately . The negative value describes lattice formation. A large positive dissociation enthalpy means a large energy input is needed to pull the crystal apart; it corresponds to a large release when the same lattice is built from gaseous ions.
Problem
A solid has lattice dissociation enthalpy . What is the enthalpy change when 0.50 mol forms from its gaseous ions?
- 1.For formation of one mole, reverse the separation process, giving .
- 2.For 0.50 mol, multiply: .
- 3.The system releases 450 kJ. The negative sign describes the direction of energy transfer, not a negative amount of substance.
Many simple cations are metallic, but the ammonium ion is a useful exception. It contains only nonmetals and still acts as a cation in ionic compounds. The N–H bonds within ammonium are covalent, while attraction between ammonium and a counter-ion contributes to an ionic crystal. A single compound can therefore contain more than one type of bonding interaction.
Check your understanding
Keep ion production, lattice separation, and lattice formation distinct. The energy sign becomes easier to interpret when the process is written explicitly.
Quiz
Which condition generally makes cation formation easier?
Why is the lattice important in understanding NaCl stability?
Using the chapter’s definition, lattice enthalpy refers to which process?
If lattice dissociation has enthalpy +800 kJ per mole, the reverse process has enthalpy:
Which statement about ionic compounds is accurate?
Practice Problems
- Write gaseous-ion formation equations for sodium and chlorine and explain the enthalpy sign in each.
- Explain why the 1:1 formula NaCl does not mean that each sodium ion interacts with only one chloride ion.
- Calculate the sum of a +420 kJ per mole ionisation step, a −300 kJ per mole electron-gain step, and a −700 kJ per mole lattice-formation step. State the starting and finishing particles for your sum.
- A salt has lattice dissociation enthalpy +760 kJ per mole. Calculate the energy change when 2.0 mol forms from gaseous ions.
- Explain why octet formation by itself is not an adequate energy argument for an ionic crystal.
- Describe how covalent and ionic interactions can both occur in ammonium chloride.
Key Takeaways
• Ion formation and lattice formation must both be considered when explaining ionic-solid stability. • Ionisation requires energy; electron gain may release or absorb it. • Ionic crystals are extended three-dimensional arrangements of cations and anions. • Lattice enthalpy here describes separation into gaseous ions and is positive. • Lattice formation is the reverse process and releases the corresponding energy.