Question

Account for the following observations: (a) The density of \(\alpha\) -graphite is less than that of diamond. (b) In group \(13,\) the melting point of \beta-rhombohedral B is much higher \((2453 \mathrm{K})\) than that of \(\mathrm{Al}(933 \mathrm{K})\) (c) The group 1 metals tend to exhibit lower values of \(\Delta_{\mathrm{a}} H^{\circ}(298 \mathrm{K})\) than metals in the \(d\) -block.

Short answer

Graphite's layered structure leads to lower density than diamond's 3D structure. Boron's covalent network solid causes its higher melting point compared to aluminum's metallic bonds. Group 1 metals have weaker bonds, lowering their enthalpy of atomization compared to d-block metals.

Step-by-step solution

Understand Atomic Structure

Graphite is composed of layers of carbon atoms arranged in hexagonal patterns, allowing for space between layers, making it less dense. Diamond has a denser, 3-dimensional tetrahedral structure.

Compare Bonding and Structure in Group 13

Boron forms strong covalent bonds in a network solid, resulting in a high melting point. Aluminum, in contrast, has metallic bonding, which is generally weaker, leading to a lower melting point.

Explain Enthalpy of Atomization in Group 1 vs d-block

Group 1 metals have larger atomic radii and weaker metallic bonds compared to d-block metals, resulting in lower enthalpy of atomization.

Key concepts

Atomic Structure

Atomic structure is fundamental to understanding why different forms of carbon, like graphite and diamond, have varying densities. In graphite, carbon atoms bond in flat layers that form hexagonal shapes. These layers are held together by weak forces, allowing them to slide over each other. This results in a structure that is less dense.

In contrast, diamond has carbon atoms bonded in a 3-dimensional tetrahedral arrangement. This structure is highly compact and rigid, making it denser than graphite. The unique configurations of atoms and bonds in these structures results in diverse physical properties.

Understanding atomic structure helps explain why substances exhibit different characteristics, such as density and hardness. Atom arrangements and bond strengths define the properties of materials.

Bonding and Structure

In chemistry, bonding and structure are crucial for explaining the properties of elements, such as those found in group 13, including boron and aluminum. Boron forms a robust network of covalent bonds, creating a giant structure known as \(\beta\)-rhombohedral B. This structure makes boron extremely resistant to heat.

Aluminum, on the other hand, forms metallic bonds. Metallic bonds occur when electrons freely move among a lattice of metal atoms, which is often weaker than covalent bonds. This explains why aluminum has a lower melting point compared to boron.

Comprehending the type of bonding and resulting structures helps predict and explain the physical properties of elements and compounds, such as melting points and mechanical strength.

Enthalpy of Atomization

Enthalpy of atomization refers to the energy required to break all bonds in one mole of atoms in the gaseous state. Group 1 metals, such as lithium and sodium, often have lower enthalpy of atomization compared to d-block metals.

This is primarily because group 1 metals have larger atomic radii, leading to weaker metallic bonds. Fewer electrons are available to participate in bonding, which results in less energy required to atomize the metal.

D-block metals, however, have more electrons involved in bonding, forming stronger bond structures and requiring more energy to break these bonds. Understanding enthalpy of atomization is essential for studying the energy changes in chemical reactions and the stability of metals.