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Molecular solid

Molecular solid is a chemistry topic covered in the lgStudy science library. This page brings together a partial reference excerpt, illustrations, worked examples, real-world applications and a short study plan, so you can understand Molecular solid rather than just read about it. In short: A molecular solid is a solid consisting of discrete molecules. The cohesive forces that bind the molecules together are van der Waals forces, dipole–dipole interactions, quadrupole interactions, π–π interactions, hydrogen bonding, halogen bonding, London dispersion forces, and in some molecular solids, coulombic interactions.

Molecular solid — main illustration
Molecular solid — illustration

Key takeaways

  • Molecular solid belongs to chemistry; place it in that map before memorising details.
  • Learn the definition first, then one example that makes the definition concrete.
  • Connect Molecular solid to a quantity you can measure, compute or draw — that is where exam questions come from.
  • Reproduce the core statement of Molecular solid from memory before moving on to harder problems.

Reference excerpt

A molecular solid is a solid consisting of discrete molecules. The cohesive forces that bind the molecules together are van der Waals forces, dipole–dipole interactions, quadrupole interactions, π–π interactions, hydrogen bonding, halogen bonding, London dispersion forces, and in some molecular solids, coulombic interactions. Van der Waals, dipole interactions, quadrupole interactions, π–π interactions, hydrogen bonding, and halogen bonding (2–127 kJ mol−1) are typically much weaker than the forces holding together other solids: metallic (metallic bonding, 400–500 kJ mol−1), ionic (Coulomb’s forces, 700–900 kJ mol−1), and network solids (covalent bonds, 150–900 kJ mol−1). Intermolecular interactions typically do not involve delocalized electrons, unlike metallic and certain covalent bonds. Exceptions are charge-transfer complexes such as the tetrathiafulvane-tetracyanoquinodimethane (TTF-TCNQ), a radical ion salt. These differences in the strength of force (i.e. covalent vs. van der Waals) and electronic characteristics (i.e. delocalized electrons) from other types of solids give rise to the unique mechanical, electronic, and thermal properties of molecular solids. Molecular solids are poor electrical conductors, although some, such as TTF-TCNQ are semiconductors (ρ = 5 × 102 Ω−1 cm−1). They are still substantially less than the conductivity of copper (ρ = 6 × 105 Ω−1 cm−1). Molecular solids tend to have lower fracture toughness (sucrose, KIc = 0.08 MPa m1/2) than metal (iron, KIc = 50 MPa m1/2), ionic (sodium chloride, KIc = 0.5 MPa m1/2), and covalent solids (diamond, KIc = 5 MPa m1/2). Molecular solids have low melting (Tm) and boiling (Tb) points compared to metal (iron), ionic (sodium chloride), and covalent solids (diamond). Examples of molecular solids with low melting and boiling temperatures include argon, water, naphthalene, nicotine, and caffeine (see table below). The constituents of molecular solids range in size from condensed monatomic gases to small molecules (i.e. naphthalene and water) to large molecules with tens of atoms (i.e. fullerene with 60 carbon atoms).

Composition and structure Molecular solids may consist of single atoms, diatomic, and/or polyatomic molecules. The intermolecular interactions between the constituents dictate how the crystal lattice of the material is structured. All atoms and molecules can partake in van der Waals and London dispersion forces (sterics). It is the lack or presence of other intermolecular interactions based on the atom or molecule that affords materials unique properties.

Van der Waals forces

Argon, is a noble gas that has a full octet, no charge, and is nonpolar. These characteristics make it unfavorable for argon to partake in metallic, covalent, and ionic bonds as well as most intermolecular interactions. It can though partake in van der Waals and London dispersion forces. These weak self-interactions are isotropic and result in the long-range ordering of the atoms into face centered cubic packing when cooled below -189.3. Similarly iodine, a linear diatomic molecule has a net dipole of zero and can only partake in van der Waals interactions that are fairly isotropic. This results in the bipyramidal symmetry.

Dipole-dipole and quadrupole interactions

For acetone dipole-dipole interactions are a major driving force behind the structure of its crystal lattice. The negative dipole is caused by oxygen. Oxygen is more electronegative than carbon and hydrogen, causing a partial negative (δ-) and positive charge (δ+) on the oxygen and remainder of the molecule, respectively. The δ- orient towards the δ+ causing the acetone molecules to prefer to align in a few configurations in a δ- to δ+ orientation (pictured left). The dipole-dipole and other intermolecular interactions align to minimize energy in the solid state and determine the crystal lattice structure.

A quadrupole, like a dipole, is a permanent pole but the electric field of the molecule is not linear as in acetone, but in two dimensions. Examples of molecular solids with quadrupoles are octafluoronaphthalene and naphthalene. Naphthalene consists of two joined conjugated rings. The electronegativity of the atoms of this ring system and conjugation cause an aromatic ring current resulting in a quadrupole. For naphthalene, this quadrupole manifests in a δ- and δ+ accumulating within and outside the ring system, respectively. Naphthalene assembles through the coordination of δ- of one molecules to the δ+ of another molecule. This results in 1D columns of naphthalene in a herringbone configuration. These columns then stack into 2D layers and then 3D bulk materials. Octafluoronaphthalene follows this path of organization to build bulk material except the δ- and δ+ are on the exterior and interior of the ring system, respectively.

Hydrogen and halogen bonding

A hydrogen bond is a specific dipole where a hydrogen atom has a partial positive charge (δ+) to due a neighboring electronegative atom or functional group. Hydrogen bonds are amongst the strong intermolecular interactions know other than ion-dipole interactions. For intermolecular hydrogen bonds the δ+ hydrogen interacts with a δ- on an adjacent molecule. Examples of molecular solids that hydrogen bond are water, amino acids, and acetic acid. For acetic acid, the hydrogen (δ+) on the alcohol moiety of the carboxylic acid hydrogen bonds with other the carbonyl moiety (δ-) of the carboxylic on the adjacent molecule. This hydrogen bond leads a string of acetic acid molecules hydrogen bonding to minimize free energy. These strings of acetic acid molecules then stack together to build solids.

A halogen bond is when an electronegative halide participates in a noncovalent interaction with a less electronegative atom on an adjacent molecule. Examples of molecular solids that halogen bond are hexachlorobenzene and a cocrystal of bromine 1,4-dioxane. For the second example, the δ- bromine atom in the diatomic bromine molecule is aligning with the less electronegative oxygen in the 1,4-dioxane. The oxygen in this case is viewed as δ+ compared to the bromine atom. This coordination results in a chain-like organization that stack into 2D and then 3D.

Coulombic interactions

… excerpt ends here. Continue reading the full article.

Illustrations

Molecular solid: Models of the packing of molecules in two molecular solids, carbon dioxide or Dry ice (a),[1] and caffeine (c).[2] The gray, red, and purple balls represent carbon, oxygen, and nitrogen, respectively. Images of carbon dioxide (b) and caffeine (d) in the solid state at room temperature and atmosphere. The gaseous phase of the dry ice in image (b) is visible because the molecular solid is subliming.
Models of the packing of molecules in two molecular solids, carbon dioxide or Dry ice (a),[1] and caffeine (c).[2] The gray, red, and purple balls represent carbon, oxygen, and nitrogen, respectively. Images of carbon dioxide (b) and caffeine (d) in the solid state at room temperature and atmosphere. The gaseous phase of the dry ice in image (b) is visible because the molecular solid is subliming.
Molecular solid: Van der Waals and London dispersion forces guide iodine to condense into a solid at room temperature.[22] (a) A lewis dot structure of iodine and an analogous structure as a spacefill model. Purple balls represent iodine atoms. (b) Demonstration of how van der Waals and London dispersion forces guide the organization of the crystal lattice from 1D to 3D (bulk material).
Van der Waals and London dispersion forces guide iodine to condense into a solid at room temperature.[22] (a) A lewis dot structure of iodine and an analogous structure as a spacefill model. Purple balls represent iodine atoms. (b) Demonstration of how van der Waals and London dispersion forces guide the organization of the crystal lattice from 1D to 3D (bulk material).
Molecular solid: The dipole-dipole interactions between the acetone molecules partially guide the organization of the crystal lattice structure.[23] (a) A dipole-dipole interaction between acetone molecules stacked on top of one another. (b) A dipole-dipole interaction between acetone molecules in front and bock of each other in the same plane. (c) A dipole-dipole interaction between acetone molecules flipped in direction, but adjacent to each other in the same plane. (d) Demonstration of how quadrupole-quadrupole interactions are involved in the crystal lattice structure.
The dipole-dipole interactions between the acetone molecules partially guide the organization of the crystal lattice structure.[23] (a) A dipole-dipole interaction between acetone molecules stacked on top of one another. (b) A dipole-dipole interaction between acetone molecules in front and bock of each other in the same plane. (c) A dipole-dipole interaction between acetone molecules flipped in direction, but adjacent to each other in the same plane. (d) Demonstration of how quadrupole-quadrupole interactions are involved in the crystal lattice structure.
Molecular solid: The quadrupole-quadrupole interactions between the naphthalene molecules partially guide the organization of the crystal lattice structure.[24] (a) A lewis dot structure artificially colored to provide a qualitative map of where the partial charges exist for the quadrupole. A 3D representation of naphthalene molecules and quadrupole. (b) A 3D representation of the quadrupole from two naphthalene molecules interacting. (c) A dipole-dipole interaction between acetone molecules flipped in direction, but adjacent to each other in the same plane. (c) Demonstration of how quadrupole-quadrupole interactions are involved in the crystal lattice structure.
The quadrupole-quadrupole interactions between the naphthalene molecules partially guide the organization of the crystal lattice structure.[24] (a) A lewis dot structure artificially colored to provide a qualitative map of where the partial charges exist for the quadrupole. A 3D representation of naphthalene molecules and quadrupole. (b) A 3D representation of the quadrupole from two naphthalene molecules interacting. (c) A dipole-dipole interaction between acetone molecules flipped in direction, but adjacent to each other in the same plane. (c) Demonstration of how quadrupole-quadrupole interactions are involved in the crystal lattice structure.
Molecular solid: The hydrogen bonding between the acetic acid molecules partially guides the organization of the crystal lattice structure.[26] (a) A lewis dot structure with the partial charges and hydrogen bond denoted with blue dashed line. A ball and stick model of acetic acid with hydrogen bond denoted with blue dashed line. (b) Four acetic acid molecules in zig-zag hydrogen bonding in 1D. (c) Demonstration of how hydrogen bonding are involved in the crystal lattice structure.
The hydrogen bonding between the acetic acid molecules partially guides the organization of the crystal lattice structure.[26] (a) A lewis dot structure with the partial charges and hydrogen bond denoted with blue dashed line. A ball and stick model of acetic acid with hydrogen bond denoted with blue dashed line. (b) Four acetic acid molecules in zig-zag hydrogen bonding in 1D. (c) Demonstration of how hydrogen bonding are involved in the crystal lattice structure.

Worked examples

Example 1 — a first encounter with Molecular solid

Start with the simplest possible case. Write down what Molecular solid claims or describes in one sentence, then invent the smallest concrete situation in which that sentence is true. In chemistry, the smallest case is usually a single object, a single equation or a single measurement. Check that every symbol or term in your sentence has a meaning in that case.

Example 2 — changing one variable

Take the situation from Example 1 and change exactly one quantity: double it, halve it, or set it to zero. Predict what should happen to Molecular solid before you calculate. Comparing your prediction with the result is the fastest way to find out whether you understand the idea or only the words.

Example 3 — an exam-style question

Typical questions about Molecular solid ask you to (a) state it precisely, (b) apply it to given data, and (c) explain a limitation. Practise writing all three answers in under five minutes; the third part is what separates a full-mark answer from an average one.

Applications of Molecular solid

In research
Molecular solid appears in chemistry research whenever the underlying quantities have to be modelled precisely. Papers usually cite it as a starting assumption and then explore where it breaks down.
In technology and industry
Engineering practice reuses Molecular solid in design rules, simulations and safety margins. Knowing the idea lets you read a specification sheet and understand why the numbers look the way they do.
In the classroom
Molecular solid is common in secondary-school and first-year university syllabi. It links to neighbouring topics Chemical compounds, so understanding it makes those chapters shorter.
In everyday life
Look for Molecular solid outside the textbook — in sport, cooking, traffic, electronics or the sky above you. An example you found yourself is remembered far longer than one you were given.

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How to study Molecular solid in 20 minutes

  1. Read the reference excerpt below once, without taking notes.
  2. Close the page and write down what Molecular solid means in your own words.
  3. Compare your version with the excerpt and mark what you missed.
  4. Work through the three examples above with pen and paper.
  5. Explain Molecular solid out loud to somebody else — or to Teacher Smith in the lgStudy chat.

Frequently asked questions

What is Molecular solid in simple terms?

A molecular solid is a solid consisting of discrete molecules. The cohesive forces that bind the molecules together are van der Waals forces, dipole–dipole interactions, quadrupole interactions, π–π interactions, hydrogen bonding, halogen bonding, London dispersion forces, and in some molecular sol…

Why does Molecular solid matter?

Because it connects several chemistry ideas at once: it gives you a definition you can apply, a quantity you can calculate, and a way to check whether a result is plausible.

How should I study Molecular solid?

Read the excerpt, restate it from memory, then work through the examples and applications listed on this page. The five-step study plan above takes about twenty minutes.

What does this page cover?

It gives you a compact reference excerpt plus original lgStudy explanations, examples, applications and study material on Molecular solid.

Tags

  • Chemical compounds

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