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Molecular solid

"Molecular crystal" redirects here. For a solid network of atoms covalently bound together, see Network covalent bonding.
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.

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.[3][4][5][6][7][8][9][10] Van der Waals, dipole interactions, quadrupole interactions, π–π interactions, hydrogen bonding, and halogen bonding (2–127 kJ mol−1)[10] are typically much weaker than the forces holding together other solids: metallic (metallic bonding, 400–500 kJ mol−1),[4] ionic (Coulomb’s forces, 700–900 kJ mol−1),[4] and network solids (covalent bonds, 150–900 kJ mol−1).[4][10] 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.[5] 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.[3][4][5][8]

Molecular solids are poor electrical conductors,[4][5] although some, such as TTF-TCNQ are semiconductors (ρ = 5 x 102 Ω−1 cm−1).[5] They are still substantially less than the conductivity of copper (ρ = 6 x 105 Ω−1 cm−1).[8] Molecular solids tend to have lower fracture toughness (sucrose, KIc = 0.08 MPa m1/2)[11] than metal (iron, KIc = 50 MPa m1/2),[11] ionic (sodium chloride, KIc = 0.5 MPa m1/2),[11] and covalent solids (diamond, KIc = 5 MPa m1/2).[12] Molecular solids have low melting (Tm) and boiling (Tb) points compared to metal (iron), ionic (sodium chloride), and covalent solids (diamond).[4][5][8][13] Examples of molecular solids with low melting and boiling temperatures include argon, water, naphthalene, nicotine, and caffeine (see table below).[13][14] The constituents of molecular solids range in size from condensed monatomic gases[15] to small molecules (i.e. naphthalene and water)[16][17] to large molecules with tens of atoms (i.e. fullerene with 60 carbon atoms).[18]

Melting and boiling points of metallic, ionic, covalent, and molecular solids
Type of Solid Material Tm (°C) Tb (°C)
Metallic Iron 1,538[13] 2,861[13]
Ionic Sodium chloride 801[13] 1,465[13]
Covalent Diamond 4,440[13] -
Molecular Argon -189.3[13] -185.9[13]
Molecular Water 0[13] 100[13]
Molecular Naphthalene 80.1[13] 217.9[13]
Molecular Nicotine -79[13] 491[13]
Molecular Caffeine 235.6[13] 519.9[14]
  1. ^ Simon, A.; Peters, K. (1980). "Single-Crystal Refinement of the Structure of Carbon Dioxide". Acta Crystallogr. B. 36 (11): 2750–2751. doi:10.1107/s0567740880009879.
  2. ^ Lehmann, C. W.; Stowasser, Frank (2007). "The Crystal Structure of Anhydrous Beta-Caffeine as Determined from X-ray Powder-Diffraction Data". Chemistry: A European Journal. 13 (10): 2908–2911. doi:10.1002/chem.200600973. PMID 17200930.
  3. ^ a b Hall, George (1965). Molecular Solid State Physics. Berlin, Germany: Springer-Verlag.
  4. ^ a b c d e f g Fahlman, B. D. (2011). Materials Chemistry. Berlin, Germany: Springer.
  5. ^ a b c d e f Schwoerer, M.; Wolf, H. C. (2007). Organic Molecular Solids. Weinheim, Germany: Wiley-VCH.
  6. ^ Omar, M. A. (2002). Elementary Solid State Physics. London, England: Pearson.
  7. ^ Patterson, J.; Bailey, B. (2010). Solid-State Physics. Berlin, Germany: Springer.
  8. ^ a b c d Turton, R. (2010). The Physics of Solids. New York, New York: Oxford University Press Inc.
  9. ^ Keer, H. V. (1993). Principles of Solid State. Hoboken, New Jersey: Wiley Eastern Limited.
  10. ^ a b c Israelachvili, J. N. (2011). Intermolecular and Surface Forces. Cambridge, Massachusetts: Academic Press.
  11. ^ a b c Varughese, S.; Kiran, M. S. R. N.; Ramamurty, U.; Desiraju, G. R. (2013). "Nanoindentation in Crystal Engineering: Quantifying Mechanical Properties of Molecular Crystals". Angewandte Chemie International Edition. 52 (10): 2701–2712. doi:10.1002/anie.201205002. PMID 23315913.
  12. ^ Field, J. E., ed. (1979). The Properties of Diamonds. New York, New York: Academic Press.
  13. ^ a b c d e f g h i j k l m n o p Haynes, W. M.; Lise, D. R.; Bruno, T. J., eds. (2016). CRC Handbook of Chemistry and Physics. Boca Raton, Florida: CRC Press.
  14. ^ a b O'Neil, M. J., ed. (2013). The Merck Index - An Encyclopedia of Chemicals, Drugs, and Biologicals. Cambridge, United Kingdom: Royal Society of Chemistry.
  15. ^ Barret, C. S.; Meyer, L. (1965). Daunt, J. G. (ed.). Low Temperature Physics: The Crystal Structures of Argon and Its Alloys. New York, New York: Springer.
  16. ^ Eisenberg, D.; Kauzmann, W. (2005). The Structures and Properties of Water. Oxford, UK: Oxford University Press.
  17. ^ Harvey, G. R. (1991). Polycyclic Aromatic Hydrocarbons: Chemistry and Carcinogenicity. Cambridge, UK: Cambridge University Press.
  18. ^ Jones, W., ed. (1997). Organic Molecular Solids: Properties and Applications. Boca Raton: CRC Press.

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