Chapter 12: Problem 81
Identify each solid as molecular, ionic, or atomic. (a) \(\mathrm{H}_{2} \mathrm{~S}(s)\) (b) \(\mathrm{KCl}(s)\) (c) \(\mathrm{N}_{2}(s)\) (d) \(\mathrm{NI}_{3}(s)\)
Short Answer
Expert verified
(a) Molecular solid, (b) Ionic solid, (c) Molecular solid, (d) Molecular solid.
Step by step solution
01
Identify Characteristics of Molecular Solids
Molecular solids are composed of atoms or molecules held together by intermolecular forces. These typically have relatively low melting points and are poor conductors of electricity in the solid state. Molecules usually have non-metal atoms held together by covalent bonds.
02
Identify Characteristics of Ionic Solids
Ionic solids are composed of cations and anions. They are held together by electrostatic attractions, known as ionic bonds. Typically, they have high melting points and conduct electricity when molten or dissolved in water.
03
Identify Characteristics of Atomic Solids
Atomic solids contain single atoms of an element. These atoms can be held together by metallic bonds, covalent bonds (network solids), or weak London dispersion forces (noble gases). They have varying melting points depending on the type of atomic bonding.
04
Classify \(\mathrm{H}_{2} \mathrm{~S}(s)\)
\(\mathrm{H}_{2} \mathrm{S}(s)\) consists of molecules made up of non-metal atoms hydrogen and sulfur, held together by covalent bonds. Hence, it is a molecular solid.
05
Classify \(\mathrm{KCl}(s)\)
\(\mathrm{KCl}(s)\) is a compound formed from a metal (potassium) and a non-metal (chlorine) which typically form ionic bonds. Thus, \(\mathrm{KCl}(s)\) is an ionic solid.
06
Classify \(\mathrm{N}_{2}(s)\)
\(\mathrm{N}_{2}(s)\) is a diatomic molecule consisting of nitrogen atoms held together by a covalent bond, with the molecules held together by London dispersion forces. It is a molecular solid.
07
Classify \(\mathrm{NI}_{3}(s)\)
\(\mathrm{NI}_{3}(s)\) consists of molecules made up of non-metal atoms (nitrogen and iodine), held together by covalent bonds, so it is categorized as a molecular solid.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Molecular Solids
Molecular solids are a class of crystalline materials where the individual particles are molecules, held together by relatively weak intermolecular forces, such as van der Waals forces, dipole-dipole interactions, or hydrogen bonds. Unlike ionic or metallic solids, the constituents of molecular solids are covalently bonded to each other within the molecule.
These solids are characterized by having low melting points compared to ionic or atomic solids due to the weaker forces holding the molecules in place. Additionally, molecular solids are typically poor conductors of electricity because they do not have free-moving charged particles. Examples of molecular solids include ice (frozen water, H_2O), dry ice (CO_2), and solid hydrogen sulfide (H_2S), which was mentioned in the exercise.
These solids are characterized by having low melting points compared to ionic or atomic solids due to the weaker forces holding the molecules in place. Additionally, molecular solids are typically poor conductors of electricity because they do not have free-moving charged particles. Examples of molecular solids include ice (frozen water, H_2O), dry ice (CO_2), and solid hydrogen sulfide (H_2S), which was mentioned in the exercise.
Ionic Solids
Ionic solids consist of positively charged cations and negatively charged anions held together by strong ionic bonds, which are electrostatic attractions between oppositely charged ions.
These types of solids typically exhibit high melting points due to the considerable energy required to break the numerous ionic bonds in a lattice structure. They are also known for conducting electricity when dissolved in water or melted, as this provides the ions the ability to move freely. Ionic solids include common table salt (NaCl), potassium chloride (KCl), and many other salts.
These types of solids typically exhibit high melting points due to the considerable energy required to break the numerous ionic bonds in a lattice structure. They are also known for conducting electricity when dissolved in water or melted, as this provides the ions the ability to move freely. Ionic solids include common table salt (NaCl), potassium chloride (KCl), and many other salts.
Atomic Solids
Atomic solids are materials where individual atoms are the fundamental building blocks of the solid structure. These atoms may be held together by different types of bonds or forces, including metallic bonds, network covalent bonds, or van der Waals forces.
Examples of atomic solids include metals like iron (Fe), which are bound by metallic bonds, diamond (a form of carbon), which is held together by network covalent bonds, and noble gases like solid argon (Ar), where atoms are held by London dispersion forces. The melting points of atomic solids can vary widely depending on the strength of the bonds that hold the atoms together.
Examples of atomic solids include metals like iron (Fe), which are bound by metallic bonds, diamond (a form of carbon), which is held together by network covalent bonds, and noble gases like solid argon (Ar), where atoms are held by London dispersion forces. The melting points of atomic solids can vary widely depending on the strength of the bonds that hold the atoms together.
Intermolecular Forces
Intermolecular forces are the attractions between molecules, which influence many physical properties of a substance, such as boiling points, melting points, and solubilities. These forces are generally much weaker than the intramolecular forces, like covalent and ionic bonds, that hold atoms together within a molecule or compound.
There are several types of intermolecular forces, including London dispersion forces, dipole-dipole interactions, and hydrogen bonding. London dispersion forces are the weakest and universal, present between all atoms and molecules. Dipole-dipole interactions occur between polar molecules, and hydrogen bonds are a special case of dipole-dipole interactions that occur when hydrogen is bound to a highly electronegative atom like oxygen, nitrogen, or fluorine.
There are several types of intermolecular forces, including London dispersion forces, dipole-dipole interactions, and hydrogen bonding. London dispersion forces are the weakest and universal, present between all atoms and molecules. Dipole-dipole interactions occur between polar molecules, and hydrogen bonds are a special case of dipole-dipole interactions that occur when hydrogen is bound to a highly electronegative atom like oxygen, nitrogen, or fluorine.
Ionic Bonds
Ionic bonds are strong attractions between oppositely charged ions, which typically form when a metal atom donates one or more electrons to a non-metal atom, resulting in the formation of cations and anions. The resulting ionic compounds are generally solid at room temperature and form crystal lattice structures.
Ionic bonds are responsible for some of the critical properties of ionic solids, such as their high melting and boiling points. Because ionic bonds involve the transfer of electrons, ionic compounds tend to have high polarity and can conduct electricity when molten or in solution, as the ions are free to move and carry a charge.
Ionic bonds are responsible for some of the critical properties of ionic solids, such as their high melting and boiling points. Because ionic bonds involve the transfer of electrons, ionic compounds tend to have high polarity and can conduct electricity when molten or in solution, as the ions are free to move and carry a charge.
Covalent Bonds
Covalent bonds are a type of chemical bond where two atoms share one or more pairs of electrons. This kind of bond typically occurs between non-metal atoms, which have similar electronegativities and thus a similar tendency to attract electrons.
Covalent compounds can be gases, liquids, or solids at room temperature, depending on the nature of the intermolecular forces present. A notable feature of many substances with covalent bonds is that they often have lower melting and boiling points than ionic compounds, owing to the weaker intermolecular forces between the molecules. In the case of molecular solids, like H_{2}S as seen in the exercise, the molecules are held together in a solid form by these covalent bonds and weaker intermolecular forces.
Covalent compounds can be gases, liquids, or solids at room temperature, depending on the nature of the intermolecular forces present. A notable feature of many substances with covalent bonds is that they often have lower melting and boiling points than ionic compounds, owing to the weaker intermolecular forces between the molecules. In the case of molecular solids, like H_{2}S as seen in the exercise, the molecules are held together in a solid form by these covalent bonds and weaker intermolecular forces.
Melting Points
The melting point of a solid is the temperature at which it changes state from solid to liquid. It is a fundamental physical property that can be used to distinguish between different types of crystalline solids.
Molecular solids, with their weaker intermolecular forces, typically have lower melting points compared to ionic or atomic solids, which have stronger bonds. Ionic solids often have high melting points due to the strong ionic bonds that require a lot of energy to break. Atomic solids, depending on the type of bonding, can have a wide range of melting points, with network covalent solids like diamond having extremely high melting points, while solids with van der Waals forces like noble gases having very low melting points.
Molecular solids, with their weaker intermolecular forces, typically have lower melting points compared to ionic or atomic solids, which have stronger bonds. Ionic solids often have high melting points due to the strong ionic bonds that require a lot of energy to break. Atomic solids, depending on the type of bonding, can have a wide range of melting points, with network covalent solids like diamond having extremely high melting points, while solids with van der Waals forces like noble gases having very low melting points.
Electrical Conductivity
Electrical conductivity is the measure of a material's ability to conduct an electric current. In solids, this is largely dependent on the presence and mobility of charged particles.
Molecular solids are generally poor conductors because they do not have free electrons or ions. Ionic solids, in contrast, conduct electricity when molten or dissolved in water but not in the solid state because the ions are fixed in place within the crystal lattice. Atomic solids, like metals, can be excellent conductors of electricity due to the sea of delocalized electrons that are free to move throughout the solid, thus allowing the flow of electric current.
Molecular solids are generally poor conductors because they do not have free electrons or ions. Ionic solids, in contrast, conduct electricity when molten or dissolved in water but not in the solid state because the ions are fixed in place within the crystal lattice. Atomic solids, like metals, can be excellent conductors of electricity due to the sea of delocalized electrons that are free to move throughout the solid, thus allowing the flow of electric current.
