Intermolecular Forces and Types of Solids

• Focus on the types of solids and intermolecular forces
• Applies knowledge of forces to structural properties of matter

Intermolecular Forces in Solids

• Solids are held together tightly, which restricts movement of constituent molecules
• Intermolecular forces are significant in solids
• Classifying solids is based on the type of intermolecular forces holding them together

Types of Solids

• Solids are classified into four types:
• Molecular
• Metallic
• Ionic
• Covalent network

Molecular Solids

• Molecular solids consist of atoms or molecules held together by intermolecular forces
• Examples of molecular solids include:
  • Solid neon (melting point of -248°C)
  • Solid CO₂ (dry ice)
  • Solid H₂O (ice)
  • Iodine
  • Sugar

Metallic Solids

• Metals typically have high (but variable) melting points
• Mercury melts at -39°C while tungsten melts at 3410°C
• Intermolecular forces in metals are strong
• Valence electrons of neighboring atoms in metals form an 'electron soup'
• These electrons are called "delocalized electrons”
• This 'delocalized electron soup' can be thought of as a 'glue' that holds the positive nuclei of the metal atoms together

Ionic Solids

• Consider sodium chloride (AKA table salt) as an example of ionic solids
• An ionic solid consists of cations and anions held together by the electrical attraction of opposite charges (ionic bonds)
• Ionic bonds are technically considered an intramolecular force (a force within a molecule)

Ionic Bonding

• Ionic compounds have very high melting points and are brittle
• Ionic compounds are non-conductive in a solid state, since e- are held tightly in place
• They are conductive in a liquid state, if molten they can conduct (e- in motion!)
• Forces in ionic compounds are derived from the creation of a crystal lattice, composed of alternating negative and positive ions
• Ionic bonds are stronger than dipole-dipole interactions because of the full charge on ions

Covalent Network Solid

• Consider carbon dioxide (CO₂) and silicon dioxide (SiO₂)
• CO₂ b.p. = -78.5°C
• If London forces are the only intermolecular force, a prediction is the b.p. of Sio₂ is slightly more than CO₂ (because it is heavier)
• SiO₂ has a m.p. of 1650°C, and b.p. of 2230°C
• SiO₂ is a covalent network solid, unlike the molecular solid CO₂
• Covalent network solids consist of atoms held together in large networks/chains by covalent bonds
• Each atom is covalently bonded, forming a 3-dimensional network
• Diamond, graphite, silicon, and asbestos are covalent network solids
• This relates to a special type of compound called an allotrope

Allotropes

• Allotropes are compounds that consist of the same element but have different physical properties
• Allotropes of carbon differ in the pattern of covalent bonds between carbon atoms
• Graphite and diamond are carbon allotropes

Graphite

• Graphite consists of 'sheets' of covalently bonded carbons
• These sheets can slide over one another easily, making graphite useful for writing (i.e. it is the 'lead' in pencils!)
• Graphite is an electrical conductor

Diamond

• Diamond, like graphite, is a covalent network solid
• Diamond forms a 3D lattice of carbon atoms instead of sheets
• This 3D lattice form gives diamond its characteristic hardness

Intermolecular Forces

• Attractions between molecules are not nearly as strong as intramolecular attractions that hold compounds together
• Intermolecular forces control physical properties such as boiling/melting points, vapor pressures, and viscosities
• Intermolecular forces as a group are van der Waals forces
• There 3 types of van der Waals forces:
• Dipole-dipole interactions
• London dispersion forces
• Hydrogen bonding (special type of dipole-dipole force)

Dipole-Dipole interactions

• Molecules with permanent dipoles are attracted to each other
• The positive end of one is attracted to the negative end of the other and vice-versa
• These forces are only important when the molecules are close to each other
• The more polar the molecule, the higher is its boiling point

London Dispersion Forces

• Electrons in the 1s orbital of helium repel each other (and, therefore, tend to stay far away from each other), but occasionally wind up on the same side of the atom.
• At that instant, then, the helium atom is polar, with an excess of electrons on one side and a shortage on the other side
• Another helium nearby would have a dipole induced in it, as the electrons on the one helium repel the electrons in the cloud on the other helium
• London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole
• These forces are present in all molecules, whether they are polar or nonpolar
• The tendency of an electron cloud to distort in this way is called polarizability.

Induced Dipoles

• Dipole-induced dipole forces are forces of attraction between a polar molecule and a non-polar molecule that has an induced dipole due to the nearby polar molecule
• Ion-induced dipole forces are forces of attraction between an ion and a non-polar molecule that has an induced dipole due to the nearby ion

Factors Affecting London Forces

• The shape of the molecule affects the strength of dispersion forces: long, skinny molecules tend to have stronger dispersion forces than short, fat ones
• This is due to the increased surface area in n-pentane
• The strength of dispersion forces tends to increase with increased molecular weight
• Larger atoms have larger electron clouds, which are easier to polarize

Dipole-Dipole vs Dispersion Forces

• If two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force
• If one molecule is much larger than another, dispersion forces will likely determine its physical properties

Explaining Trends

• The nonpolar series (SnH₄ to CH₄) follow the expected trend
• The polar series follows the trend from H₂Te through H₂S, but water is quite an anomaly

Hydrogen Bonding

• The dipole-dipole interactions when H is bonded to N, O, or F are unusually strong
• We call these interactions hydrogen bonds
• Hydrogen bonding arises due to the high electronegativity of nitrogen, oxygen, and fluorine

Ion-Dipole Interactions

• Ion-dipole interactions are forces of attraction between partial charges on polar molecules and ions
• The strength of these forces make it possible for ionic substances to dissolve in polar solvents

Intermolecular Force Examples

• HBr: dipole-dipole and dispersion forces between HBr molecules, is a polar molecule
• CH4: dispersion forces, is nonpolar
• SO2: dipole-dipole and dispersion forces between SO₂ molecules, is a polar molecule

Summarizing Types of Crystals

• Ionic Crystals
• Formed from lattice points occupied by cations and anions
• Held together by electrostatic attraction
• Hard, brittle, and have high melting points
• Poor conductors of heat and electricity
• Covalent Crystals
• Formed from lattice points occupied by atoms
• Held together by covalent bonds
• Hard, and have high melting points
• Poor conductors of heat and electricity
• Molecular Crystals
• Formed from Lattice points occupied by molecules
• Held together by intermolecular forces
• Soft, and have low melting points
• Poor conductors of heat and electricity
• Metallic Crystals
• Formed from lattice points occupied by metal atoms
• Held together by metallic bonds
• Soft to hard, and have low to high melting points
• Good conductors of heat and electricity

Explores solids, intermolecular forces, and their impact on matter's structural properties. Covers molecular, metallic, ionic, and covalent network solids, with examples.