Question

Identify each solid as molecular, ionic, or atomic. (a) \(\mathrm{SF}_{6}(s)\) (b) \(C(s)\) (c) \(\mathrm{MgCl}_{2}(s)\) (d) \(\mathrm{Ti}(\mathrm{s})\)

Short answer

\(\mathrm{SF}_{6}(s)\) is molecular, \(C(s)\) is atomic, \(\mathrm{MgCl}_{2}(s)\) is ionic, and \(\mathrm{Ti}(\mathrm{s})\) is atomic.

Step-by-step solution

Understanding Molecular Solids

A molecular solid consists of atoms or molecules held together by intermolecular forces like hydrogen bonding, dipole-dipole interactions, or London dispersion forces. These are generally nonmetals bonded with nonmetals.

Classify \(\mathrm{SF}_{6}(s)\)

Since \(\mathrm{SF}_{6}\) is composed of the nonmetal sulfur (\(\mathrm{S}\)) bonded to six fluorine atoms (\(\mathrm{F}\)), this compound is a molecular solid.

Understanding Atomic Solids

Atomic solids are made up of atoms connected by covalent bonds, which form a giant covalent structure. Examples include pure elements like carbon in its diamond or graphite form.

Classify \(C(s)\)

Carbon (\(C\)) as a solid (such as diamond or graphite) consists of atoms bonded together in a network lattice. Hence, carbon is an atomic solid.

Understanding Ionic Solids

Ionic solids consist of positive and negative ions held together by ionic bonds. These compounds are generally formed between metals and nonmetals.

Classify \(\mathrm{MgCl}_{2}(s)\)

Magnesium chloride (\(\mathrm{MgCl}_{2}\)) is made from the metal magnesium (\(\mathrm{Mg}\)) and the nonmetal chlorine (\(\mathrm{Cl}\)), held together by ionic bonds. Therefore, \(\mathrm{MgCl}_{2}\) is classified as an ionic solid.

Classify \(\mathrm{Ti}(\mathrm{s})\)

Titanium (\(\mathrm{Ti}\)) is a metal that exists in solid form as a lattice of titanium atoms. It doesn't consist of molecules or ions, but rather a metal lattice with delocalized electrons, which makes it an atomic solid.

Key concepts

Molecular Solids

Molecular solids are materials in which atoms or molecules are held together by comparatively weaker intermolecular forces, rather than by ionic or covalent bonds. In molecular solids, the components are usually nonmetal atoms or molecules, like water (H_2O) in its solid state, ice. For example, in sulfur hexafluoride (SF_6(s)), the molecules are held together by London dispersion forces, which are the weakest intermolecular forces. These solids typically have lower melting and boiling points, and they are often soft or flexible. They are also poor conductors of heat and electricity because of the lack of free electrons or ions.

It's essential for students to understand that the physical properties of molecular solids—like their melting point and hardness—are directly related to the strength of the intermolecular forces holding the molecules together. The type of intermolecular force present depends on the specific molecules involved and can include dipole-dipole interactions, hydrogen bonds, and London dispersion forces.

• Dipole-dipole interactions occur when the positive part of one polar molecule is attracted to the negative part of another.
• Hydrogen bonds are a specific type of dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine.
• London dispersion forces are temporary fluctuating dipoles that occur even in non-polar molecules.

Atomic Solids

Atomic solids, on the other hand, are composed of atoms linked together by covalent bonds. These solids are characterized by their network lattices or structures in which each atom is covalently bonded to its neighbors forming a continuous network throughout the material. One of the most famous examples of an atomic solid is carbon in the form of diamond. In diamond, each carbon atom is bonded to four other carbon atoms creating a very strong and rigid structure. This accounts for its impressive properties such as high hardness and high melting points.

These bonds are much stronger than the intermolecular forces found in molecular solids, which is why atomic solids are typically much harder and have higher melting points. Graphite, another allotrope of carbon, also falls under this category, but it has a different structure, consisting of layers of carbon atoms that can slide over each other, which makes it a good lubricant and is the reason it feels slippery to the touch. Atomic solids are usually good conductors of heat and electricity, especially when the atoms involved have free electrons, like in the case of metallic solids.

• Covalent network solids like diamond, silicon carbide (SiC), and quartz (SiO_2) have high melting points and are very hard.
• Metallic solids, such as titanium (Ti(s)), have delocalized electrons that allow for electrical conductivity and malleability.

Ionic Solids

Ionic solids are formed when metal and nonmetal atoms come together to form a crystal lattice structure based on the strong electrostatic forces of attraction between positively and negatively charged ions. These ionic bonds occur as a result of the complete transfer of one or more electrons from a metal atom to a nonmetal atom, generating ions that attract each other. A common example here is magnesium chloride (MgCl_2(s)), where the magnesium (Mg^{2+}) ions and chloride (Cl^-) ions are held together in a repeating pattern that extends in three dimensions.

The ionic bonds are much stronger than intermolecular forces, which generally means ionic solids have high melting and boiling points. Additionally, while they typically do not conduct electricity in their solid state, they can conduct when melted or dissolved in water as the ions are free to move. Students should recognize characteristics such as the brittleness, hardness, and the ability to conduct electricity in solution, as distinctive features of ionic solids.

• They often form crystals with characteristic shapes due to their regular lattice structures.
• When dissolved in water, they can conduct electricity due to the dissociation into ions.

Intermolecular Forces

Intermolecular forces are the attractive forces that act between molecules. These forces are not as strong as covalent or ionic bonds but they are significant in determining the physical properties of molecular solids and liquids. We've earlier touched upon the different types of intermolecular forces such as hydrogen bonds, dipole-dipole interactions, and London dispersion forces. Understanding these forces is key to comprehending why certain substances have particular melting and boiling points, why some materials are gases at room temperature while others are liquids or solids, and how substances dissolve in solvents.

Recognizing the strength of these forces helps explain, for instance, why the boiling point of water is unusually high for such a small molecule - it's because of the strong hydrogen bonds between water molecules. Similarly, the reason why oil and water don't mix is based on the differences in intermolecular attractions: water molecules are strongly attracted to each other via hydrogen bonds, and non-polar oil molecules are held together by London dispersion forces, leading to poor solubility of oil in water.

• Intermolecular forces are crucial for understanding phase changes, solubility, and surface tension.

Covalent Bonds

Covalent bonds are the strong bonds that occur when nonmetal atoms share pairs of electrons. The shared electrons allow each atom to achieve a noble gas configuration, satisfying the octet rule. These bonds form the basic structural unit for many compounds, whether they are simple molecules like water or oxygen (O_2) or complex network solids like quartz (SiO_2). A single covalent bond represents a pair of shared electrons, and molecules can have single, double, or triple covalent bonds.

Covalent bonding is fundamental to organic chemistry and is what allows for the vast diversity of organic compounds, including biological macromolecules like proteins and DNA. It is the sharing of electron pairs between atoms that gives these substances their unique properties, and variations in covalent bond strength and arrangement can dramatically alter the physical characteristics of the compound.

• Covalent bonds allow atoms to achieve stable electron configurations.
• They are directional, which means they influence the shape and structure of molecules.

Ionic Bonds

Ionic bonds, distinct from covalent bonds, form when there is a complete transfer of one or more electrons from a metal to a nonmetal atom, resulting in the formation of ions. The positively charged metal ion (cation) and the negatively charged nonmetal ion (anion) are then held together by the strong electrostatic forces of attraction between opposite charges. Ionic compounds, such as sodium chloride (NaCl), typically form crystalline solids that have high melting and boiling points due to the strong attraction between the ions.

It's useful to understand that even though they result in seemingly rigid structures, ionic compounds can conduct electricity when melted or dissolved in water, as these conditions allow ions to move freely. Moreover, when discussing ionic bonding with students, highlighting the role of electron transfer in the formation of ions helps clarify the distinction between ionic and covalent bonds, two fundamentally different types of chemical bonds with different physical properties and behavior.

• These bonds are responsible for the crystalline structures and brittleness of ionic solids.
• They also govern the solubility and electrical conductivity in the molten state or aqueous solutions.