Animated Solution for Chemistry - s and p-Block Elements: The correct set from the following in which both pairs are in correct order of melting point is

When we talk about the melting point of an ionic compound, we are essentially discussing how much thermal energy is required to break the rigid, three-dimensional crystal lattice into a liquid state. The stronger the electrostatic forces holding the ions together, the higher the melting point. This strength is quantified by a concept known as Lattice Energy (L.E.).

Lattice energy is governed by Coulomb's Law, which states that the force of attraction between two oppositely charged particles is directly proportional to the product of their charges and inversely proportional to the square of the distance between them. For lattice energy, the relationship simplifies to:

Here, $z_{+}$ and $z_{-}$ are the charges of the cation and anion, respectively, and $r_{+}$ and $r_{-}$ are their ionic radii. This formula gives us a clear hierarchy: Charge is the dominant factor, and when charges are equal, size becomes the tiebreaker.

LiF vs LiCl

Let's apply this to our first pair: Lithium Fluoride ($LiF$) and Lithium Chloride ($LiCl$).

First, we check the charges. In both compounds, the lithium ion has a $+1$ charge ($L{i}^{+}$), and the halide ions have a $-1$ charge ($F^{-}$ and $C{l}^{-}$). Because the product of the charges is identical ($|(+1)(-1)|=1$), the charge factor cannot help us differentiate their lattice energies.

We must move to the tiebreaker: ionic size. The lithium cation is common to both, so the difference lies entirely in the anions. Fluorine is in the second period of the periodic table, while chlorine is in the third. Therefore, the fluoride ion ($F^{-}$) is significantly smaller than the chloride ion ($C{l}^{-}$).

A smaller anion means that the distance between the centers of the positive and negative ions ($r_{+}+r_{-}$) is shorter in $LiF$ than in $LiCl$. According to our formula, a smaller denominator leads to a larger lattice energy. Consequently, the electrostatic attraction in $LiF$ is stronger, making its melting point higher than that of $LiCl$.

Result: $LiF>LiCl$

MgO vs NaCl

Now, let's examine the second pair: Magnesium Oxide ($MgO$) and Sodium Chloride ($NaCl$).

We start again by checking the charges. In $MgO$, the magnesium ion has a $+2$ charge ($M{g}^{2+}$) and the oxide ion has a $-2$ charge ($O^{2-}$). In $NaCl$, the sodium ion has a $+1$ charge ($N{a}^{+}$) and the chloride ion has a $-1$ charge ($C{l}^{-}$).

Let's calculate the charge products: For $MgO$: $|(+2)(-2)|=4$ For $NaCl$: $|(+1)(-1)|=1$

The charge product for $MgO$ is four times greater than that of $NaCl$. Because charge is the dominant factor in determining lattice energy, this massive difference in electrostatic attraction completely overshadows any minor differences in ionic radii. The $+2$ and $-2$ ions pull on each other with immense force, creating a highly stable crystal lattice.

Therefore, the lattice energy of $MgO$ is vastly superior to that of $NaCl$, resulting in a much higher melting point.

Result: $MgO>NaCl$

By systematically applying the principles of lattice energy, we have determined that the correct order for both pairs is $LiF>LiCl$ and $MgO>NaCl$. Looking at our options, this perfectly matches Option (a).

Whenever you face questions comparing the thermal stability or melting points of ionic compounds, always remember the golden rule: Check the charges first, and if they tie, check the sizes!