Anand Classes provides complete notes on Lattice Enthalpy of Ionic Crystals and Factors Affecting Lattice Enthalpy for Class 11 Chemistry, JEE, and NEET. Lattice enthalpy is defined as the energy released when one mole of an ionic solid is formed from its gaseous ions, and it plays a vital role in determining the stability of ionic compounds. These notes explain the concept in detail along with the influencing factors such as ionic size, ionic charge, and crystal packing, making it easier for students to prepare for board exams as well as competitive exams like JEE and NEET. Click the print button to download study material and notes.
What is Lattice Enthalpy of Ionic Crystals ?
The stability of ionic crystals is determined in terms of their lattice enthalpy.
Definition
Lattice enthalpy is the amount of energy released when one mole of an ionic crystal is formed from its constituent ions in the gaseous state.
Formation of one mole of an ionic solid is represented as:
$$
\mathrm{M^{+}(g) + X^{-}(g) \rightarrow MX(s)} \quad \Delta H = \text{Lattice enthalpy} = -U
$$
- $U$ represents the lattice enthalpy.
- The negative sign indicates that energy is released during crystal formation.
- Ionic bond formation lowers the potential energy of the system due to strong electrostatic attractions between oppositely charged ions.
Reverse Process
If one mole of a solid ionic compound is broken into its gaseous ions, energy is absorbed, which is numerically equal to the lattice enthalpy:
$$
\mathrm{MX(s) \rightarrow M^{+}(g) + X^{-}(g)} \quad \Delta H = +U
$$
- Here, $U$ is positive because energy is required to overcome the interionic attractions.
Hence, lattice enthalpy may also be defined as:
The energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
Importance of Lattice Enthalpy
- A higher lattice enthalpy indicates greater stability of the ionic compound because greater amount of energy is required to break one mole of a solid ionic compound into its gaseous ions.
- The magnitude of lattice enthalpy gives an idea of the strength of interionic forces.
Factors Affecting Lattice Enthalpy
1. Size of the Ions
- Smaller the size of the ions, lesser is the internuclear distance. Consequently, the interionic attractions will be high and the lattice enthalpy will also be large.
Smaller ionic size ⇒ shorter internuclear distance ⇒ stronger interionic attraction ⇒ higher lattice enthalpy. - For example:
Ionic radius of K$^+$ = 133 pm (larger)
Ionic radius of Na$^+$ = 95 pm (smaller) - Therefore:
$$
\mathrm{Lattice\ enthalpy\ of\ NaCl = 758.7\ kJ\ mol^{-1} > KCl = 681.4\ kJ\ mol^{-1}}
$$
2. Charge on the Ions
- Larger the magnitude of charge on the ions, greater will be the attractive forces between the ions. Consequently, the lattice enthalpy will be high.
Greater ionic charge ⇒ stronger electrostatic attraction ⇒ higher lattice enthalpy. - Coulombic attraction varies directly with the product of the charges ($q_1, q_2$):
$$
F \propto \frac{q_1 q_2}{r}
$$ - Compounds with higher ionic charges (e.g., MgO, Al$_2$O$_3$) have much larger lattice enthalpies than those with singly charged ions (e.g., NaCl).
General Trends
- Uni-univalent compounds (e.g., NaCl, KCl) → Lower lattice enthalpy
- Uni-bivalent compounds (e.g., CaF$_2$) → Higher lattice enthalpy
- Bi-bivalent compounds (e.g., MgO, CaO) → Highest lattice enthalpy
Note : Compounds with small, highly charged ions exhibit the strongest ionic bonding.
Example of Exceptional Cation
Most ionic compounds have metallic cations and non-metallic anions.
However, the ammonium ion ($\mathrm{NH_4^{+}}$)—made entirely of non-metallic elements—acts as a cation in many ionic compounds, such as NH$_4$Cl.
Lattice Enthalpies of Some Ionic Solids
The following table shows the lattice enthalpies of different types of ionic solids. Values of lattice enthalpies are given in kJ/mol.
| Uni–univalent solids | Lattice enthalpy | Bi–univalent solids | Lattice enthalpy | Bi–bivalent solids | Lattice enthalpy |
|---|---|---|---|---|---|
| LiF | –1033 | CaF$_2$ | –2581 | BeO | –3125 |
| CsF | –748 | CaCl$_2$ | –2254 | MgO | –3932 |
| NaCl | –758 | MgF$_2$ | –2882 | MgS | –3254 |
| NaBr | –752 | MnCl$_2$ | –2525 | ZnO | –4032 |
| LiI | –140 | ||||
| CsI | –601 | ||||
| AgCl | –895 | ||||
| AgI | –795 |
Observations
- Uni–univalent solids (e.g., NaCl, LiF) have lower lattice enthalpies than bi–univalent and bi–bivalent solids.
- Bi–bivalent solids (e.g., MgO, ZnO) show the highest lattice enthalpies because of:
- Smaller ionic sizes
- Higher ionic charges, leading to stronger electrostatic attraction.
Key Points to Remember
- Higher lattice enthalpy ⇒ higher stability of the ionic crystal.
- Lattice enthalpy depends on:
- Charge on ions (direct relationship)
- Size of ions (inverse relationship)
Short Answer Conceptual Types Questions (SAT) on Lattice Enthalpy
Q1. What is lattice enthalpy?
Lattice enthalpy is the energy released when one mole of an ionic crystal is formed from its gaseous ions, or the energy required to separate one mole of a solid ionic compound into gaseous ions.
Q2. Why do compounds with small, highly charged ions have higher lattice enthalpy?
Smaller ions are closer together, and higher charges increase electrostatic attraction, resulting in stronger interionic forces and higher lattice enthalpy.
Q3. How is lattice enthalpy related to stability of ionic compounds?
Higher lattice enthalpy ⇒ higher stability because more energy is required to break the crystal into ions.
Q4. Give an example of an exception to metallic cation rule.
The ammonium ion ($\mathrm{NH_4^+}$) is a cation made entirely of non-metallic elements.
Multiple Choice Questions (MCQs) With Answers and Explanation on Lattice Enthalpy
1. Which factor increases lattice enthalpy?
A) Larger ionic size
B) Higher ionic charge
C) Lower ionic charge
D) None of the above
Answer: B (Higher ionic charge increases electrostatic attraction)
2. Lattice enthalpy is:
A) Always endothermic
B) Always exothermic
C) Exothermic for formation, endothermic for dissociation
D) Independent of ionic charge
Answer: C (Energy is released on formation, absorbed on breaking)
3. Among LiF, NaCl, and CsF, which has the highest lattice enthalpy?
A) LiF
B) NaCl
C) CsF
D) All same
Answer: A (LiF has the smallest cation, so strongest interionic attraction)
Assertion Reason Type Questions With Answers and Explanation on Lattice Enthalpy
Assertion (A): MgO has higher lattice enthalpy than NaCl.
Reason (R): Mg$^{2+}$ and O$^{2-}$ ions have higher charges than Na$^+$ and Cl$^-$.
A) Both A and R are true, R is correct explanation of A
B) Both A and R are true, R is not correct explanation
C) A is true, R is false
D) A is false, R is true
Answer: A
Assertion (A): CsI has lower lattice enthalpy than LiF.
Reason (R): Cs$^+$ and I$^-$ ions are larger in size, increasing internuclear distance.
A) Both A and R are true, R is correct explanation of A
B) Both A and R are true, R is not correct explanation
C) A is true, R is false
D) A is false, R is true
Answer: A
Case Study based on Lattice Enthalpy
Passage:
Consider the formation of NaCl and MgO. NaCl has lattice enthalpy –758 kJ mol$^{-1}$, while MgO has –3932 kJ mol$^{-1}$. The higher lattice enthalpy of MgO is due to smaller ions and higher ionic charges, making the ionic bond stronger and the crystal more stable.
Questions:
- Which compound has stronger ionic bonds?
- Why is lattice enthalpy of MgO much higher than NaCl?
- How does lattice enthalpy relate to stability?
Answers:
- MgO
- Because Mg$^{2+}$ and O$^{2-}$ have higher charges and smaller sizes, increasing electrostatic attraction.
- Higher lattice enthalpy ⇒ more stable ionic crystal.