Problem 137 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following statement(s) is/are true? (a) Stability of alkali metal peroxide increases with increase in atomic number (b) Hydration energy of \(\mathrm{AgF}\) is higher than its lattice energy (c) Anhydrous \(\mathrm{MgCl}_{2}\) cannot be prepared by direct heating of \(\mathrm{MgCl}_{2} .6 \mathrm{H}_{2} \mathrm{O}\) (d) The milk of magnesia used as antacid is chemically \(\mathrm{MgO}+\mathrm{MgCl}_{2}\)

Step-by-Step Solution

Answer

Statements (a) and (c) are true.

Step 1: Analyze Statement (a)

The statement "Stability of alkali metal peroxide increases with increase in atomic number" refers to the trend of the stability of peroxides of alkali metals like lithium, sodium, potassium, etc. As we move down the group in the periodic table, the size of the alkali metals increases, which leads to greater ionic character in the peroxides and greater stability. This statement is true.

Step 2: Analyze Statement (b)

The statement "Hydration energy of \(\mathrm{AgF}\) is higher than its lattice energy" refers to the comparison between the energy released when forming hydrated ions and the energy required to break the ionic lattice. The lattice energy of silver fluoride (\(\mathrm{AgF}\)) is generally higher than its hydration energy, meaning this statement is false.

Step 3: Analyze Statement (c)

The statement "Anhydrous \(\mathrm{MgCl}_{2}\) cannot be prepared by direct heating of \(\mathrm{MgCl}_{2} .6 \mathrm{H}_{2} \mathrm{O}\)" concerns the preparation of magnesium chloride without water. During heating, \(\mathrm{MgCl}_{2} .6 \mathrm{H}_{2} \mathrm{O}\) decomposes to magnesium oxide and not \(\mathrm{MgCl}_{2}\), thus the statement is true.

Step 4: Analyze Statement (d)

The statement "The milk of magnesia used as antacid is chemically \(\mathrm{MgO}+\mathrm{MgCl}_{2}\)" suggests a specific composition for milk of magnesia. Milk of magnesia is actually a suspension of magnesium hydroxide \(\mathrm{Mg(OH)}_{2}\), making this statement false.

Key Concepts

Ionic Lattice

In the world of chemistry, an ionic lattice is a crystalline structure formed by the regular arrangement of ions. These structures are extremely stable due to the strong electrostatic forces between the oppositely charged ions. The simplest analogy is to think of it like a huge 3D grid where each positive ion is surrounded by negative ions, and vice versa.

The strength of an ionic lattice is determined by two key factors: the charge on the ions and the size of the ions. Higher charges increase the force of attraction, making the lattice stronger. Meanwhile, smaller ions pack more closely together, leading to stronger interactions and a more robust structure.

• Charge on the Ions: More charge means stronger attraction and a more stable lattice.
• Size of the Ions: Smaller ions create tighter and stronger lattice structures due to closer packing.

These properties help explain why certain ionic compounds are incredibly hard and have high melting points. Understanding ionic lattices is crucial, not just in theory, but in real-world applications like material science and engineering.

Hydration Energy

Hydration energy refers to the energy change accompanying the dissolution of ions in water. When an ionic compound dissolves, the ions separate and each ion becomes surrounded by water molecules. This process releases energy known as hydration energy.

Hydration energy varies depending on the nature of the ions involved. It is typically higher for smaller, highly charged ions because they exert a stronger attractive force on the water molecules. This increased energy release can often stabilize the dissolved ions in solution.

• Ion Size: Smaller ions have higher hydration energy due to stronger attraction to water molecules.
• Ion Charge: Ions with higher charges have greater hydration energy, as they interact more strongly with water.

Compared to lattice energy (the energy required to break the ionic lattice), if hydration energy is higher, the compound readily dissolves in water. For example, the case of \( \mathrm{AgF} \) having hydration energy less than its lattice energy, means this compound is less likely to dissolve, as more energy is required to dissociate the lattice than is gained from hydration.

Alkali Metals

Alkali metals include elements like lithium, sodium, and potassium found in the first column of the periodic table. Known for being highly reactive, especially with water, they have several distinctive properties and trends.

Upon moving down the group, alkali metals decrease in ionization energy, making it easier for them to lose an electron and form positive ions. This correlates with increased reactivity. Their ionic radius increases, yet they maintain a consistent charge of +1. As products like peroxides form (e.g., sodium peroxide \( \text{Na}_2\text{O}_2 \)), the larger size of the metal allows for greater stability due to the spreading out of the charge and reduced electron electron repulsions.

• Reactivity: Increases down the group as ionization energy decreases.
• Size Trend: Atomic and ionic size increase, causing compounds like peroxides to become more stable.

These trends highlight why alkali metals are easily recognizable with their unique set of chemical properties.

Periodic Table Trends

Periodic table trends describe how certain chemical properties change in a predictable way across the periodic table. These trends include variations in atomic size, ionization energy, electronegativity, and more as you move across a period or down a group.

One significant trend in groups like the alkali metals (Group 1) is the increase in size and decrease in ionization energy as you move down the group. This means that elements become easier to ionize, influencing their reactivity. Conversely, as you move from left to right across a period, atomic size decreases while ionization energy and electronegativity increase, resulting in a variety of chemical behaviors.

• Atomic Size: Increases down a group, decreases across a period.
• Ionization Energy: Decreases down a group, increases across a period.
• Electronegativity: Generally decreases down a group and increases across a period.

Understanding these trends is essential for predicting how different elements will interact in chemical reactions, allowing for more precise scientific predictions and applications.