Question

Which of the following types of forces bind together the carbon atoms in diamond? (a) Dipolar (b) Van der Waal's (c) Ionic (d) Covalent

Short answer

Diamond carbon atoms are bound by covalent forces.

Step-by-step solution

Understand carbon bonding

In order to answer the question, we need to determine how carbon atoms bond in diamond. Carbon has four electrons in its outer shell and forms strong bonds by sharing electrons with other atoms.

Identify possible bond types

Next, evaluate the types of bonds: Dipolar and Van der Waal's forces are typically weaker intermolecular forces and are not suitable for binding atoms within a molecule or crystal structure, like diamond. Ionic bonds, involve electron transfer, which usually happens between metals and non-metals.

Consider covalent bonding

Consider the characteristics of covalent bonds. Diamond is known for its hardness and high melting point, characteristics typical of covalent bonds, where atoms share electrons to achieve a full outer shell.

Match diamond properties

Match the diamond's known properties, such as its strength and high melting point, with the characteristics associated with covalent bonds. This step confirms the type of bond present in diamond.

Key concepts

Covalent Bonds

Covalent bonds are a fundamental type of chemical bond where atoms share pairs of electrons. This sharing allows each atom to achieve the electron configuration of a noble gas, resulting in a stable configuration.
In the case of carbon, it has four valence electrons and needs four more to fill its outer shell. By sharing these electrons with other carbon atoms, carbon can form strong covalent bonds. This results in each carbon atom having a complete set of eight electrons in its outer shell.

• Covalent bonds are typically very strong, resulting in high melting and boiling points for covalently bonded substances.
• The shared electron pairs create a strong attraction between the atoms involved, making covalent structures generally insoluble in water.
• Covalent compounds exist in various forms, with diamond being a prominent example of a network solid where atoms are continuously bonded throughout the structure.

So, when we talk about diamond, it's critical to understand that each of the carbon atoms is covalently bonded to four other carbon atoms, creating an extremely robust three-dimensional network.

Crystal Structure of Diamond

The crystal structure of diamond is what gives it its exceptional properties. Diamond forms a lattice structure known as a cubic crystal system.
In this structure, each carbon atom is tetrahedrally coordinated, meaning each carbon atom shares electrons with four other carbon atoms, forming a very regular and repeating pattern.

• This tetrahedral structure leads to an incredibly strong and rigid lattice network.
• The crystal lattice in diamonds is characterized by having every bond a part of the structure, contributing to its unparalleled hardness.
• This tightly packed structure ensures that there are very few topologically weak points within the diamond, resulting in its notable resistance to scratching and wear.
• The uniformity of the structure also contributes to its ability to refract and reflect light, giving diamonds their famous sparkle.

Such a perfectly structured array of atoms, bound tightly by strong covalent bonds, explains why diamonds are one of the hardest known materials.

Properties of Diamond

Diamond is renowned not only for its beauty but for several physical properties that make it unique among materials. First and foremost, due to its crystal structure and the strength of its covalent bonds, diamond is extremely hard.
In fact, it is the hardest known natural material, making it an excellent material for cutting tools and other applications requiring durability.

• Besides hardness, diamond also has a very high thermal conductivity. This means that diamond can effectively dissipate heat, making it valuable in applications like heat sinks and high-performance electronics.
• Diamonds are also sparingly low in electrical conductivity because there are no free electrons or ions to carry a charge, unlike in metals.
• Optical properties are another highlight, as diamonds boast an extraordinary ability to refract light, contributing to their brilliance and sparkle.

These properties, combined with rarity, make diamonds a highly sought-after material both for industrial applications and as a gemstone.