Intermolecular Forces and Molecular Properties Explained

Intermolecular Forces (IMFs)

The forces of attraction between molecules.

Dipole-Dipole Forces

Occurs in polar molecules; the positive end of one molecule is attracted to the negative end of another.

Example of Dipole-Dipole Forces

HCl, SO₂

London Dispersion Forces

Also known as Van der Waals Forces; occurs in all molecules, especially significant in nonpolar molecules due to temporary dipoles.

Strength of London Dispersion Forces

Weakest of all IMFs, but increases with molecule size and surface area.

Example of London Dispersion Forces

Noble gases (like Ar), hydrocarbons (like CH₄)

Hydrogen Bonding

A special strong case of dipole-dipole forces that occurs when H is bonded to N, O, or F.

Example of Hydrogen Bonding

Water (H₂O), ammonia (NH₃), HF

Ion-Dipole Forces

Occurs in mixtures of ionic compounds and polar molecules.

Example of Ion-Dipole Forces

Na⁺ dissolved in water (the attraction between Na⁺ and the negative pole of H₂O).

Comparison of Strength of IMFs

Hydrogen Bonding > Dipole-Dipole > London Dispersion (in small molecules).

Polar Molecules

Molecules that have a positive side and a negative side due to uneven sharing of electrons.

Dipole

A separation of charge within a molecule, resulting in a slightly positive side (δ⁺) and a slightly negative side (δ⁻).

Example of Dipole-Dipole Attraction

In HCl, Cl (δ⁻) is attracted to H (δ⁺) of another HCl molecule.

Determining if a Molecule has Dipole-Dipole Forces

Check if the molecule is polar; if yes, it has dipole-dipole forces.

Step 1 to Determine Polarity

Check bond polarity by looking at the difference in electronegativity between atoms.

Step 2 to Determine Polarity

Check the shape of the molecule; if polar bonds cancel each other out, the molecule is nonpolar.

Example of Polar Molecule

HCl is polar because Cl is more electronegative than H and there is no canceling out.

Example of Nonpolar Molecule

CO₂ is nonpolar because the polar bonds cancel each other out.

TL;DR for Dipole-Dipole Forces

If the molecule is polar, it has dipole-dipole forces; if not, it only has London dispersion forces.

Molecular Geometry

The 3D shape of a molecule that affects its overall polarity.

Polar Bonds

Bonds between atoms with different electronegativities, causing a dipole moment.

Dipole Cancellation

Occurs when polar bonds in a symmetrical molecule pull equally and oppositely.

Symmetrical Molecule

A molecule with a balanced arrangement of polar bonds, leading to nonpolar characteristics.

Asymmetrical Molecule

A molecule with an uneven distribution of polar bonds, leading to polar characteristics.

CO₂

A nonpolar molecule with a linear shape where dipoles cancel.

H₂O

A polar molecule with a bent shape where dipoles do not cancel.

CF₄

A nonpolar molecule with a tetrahedral shape where dipoles cancel.

CHCl₃

A polar molecule with a tetrahedral shape but asymmetrical distribution of dipoles.

VSEPR Theory

A theory used to predict the shape of molecules based on electron pair repulsion.

CH₃OH

Methanol, a polar molecule with polar bonds and a bent shape.

Molecule Type IMFs

Classification of intermolecular forces based on molecular polarity and bonding.

CH₄

A nonpolar molecule that exhibits only London dispersion forces.

HCl

A polar molecule that exhibits London and dipole-dipole forces.

CH₃OH IMFs

Methanol exhibits London dispersion, dipole-dipole, and hydrogen bonding.

Temporary Dipoles

Momentary uneven distribution of charge in molecules due to electron movement.

Electronegativity

The tendency of an atom to attract electrons in a bond.

Net Dipole

The overall dipole moment of a molecule resulting from the vector sum of all dipoles.

Tug-of-War Analogy

A metaphor used to explain how dipoles can cancel each other out in symmetrical molecules.

Molecular Shape

The geometric arrangement of atoms in a molecule, influencing its polarity.

Bonds in CH₃OH

C-H (slightly polar), C-O (polar), O-H (polar).

Strength of Hydrogen Bonding

5-40 kJ/mol, strongest intermolecular force.

Strength of Dipole-Dipole Forces

3-10 kJ/mol, weaker than hydrogen bonds but stronger than dispersion forces.

Boiling Point and IMF Strength

Boiling point is directly related to IMF strength; stronger forces result in higher boiling points.

CH₄ (methane)

Nonpolar molecule with only London dispersion forces, resulting in the lowest boiling point.

CH₃Cl (chloromethane)

Polar molecule with dipole-dipole and London dispersion forces, resulting in medium boiling point.

CH₃OH (methanol)

Polar molecule with hydrogen bonding, dipole-dipole, and London dispersion forces, resulting in the highest boiling point among the three.

C₈H₁₀ (xylene)

Nonpolar hydrocarbon with London dispersion forces only, medium-low boiling point due to larger size.

SiO₂ (silicon dioxide)

Not held together by IMFs; a network covalent solid with an extremely high boiling point.

Br₂ (bromine)

Nonpolar molecule with London dispersion forces, higher boiling point than C₈H₁₀ due to larger size.

Final Ranking of Boiling Points

C₈H₁₀ < Br₂ < CH₃OH < SiO₂.

IMFs in Hydrogen Bonding

If a molecule has hydrogen bonding, it automatically has dipole-dipole and London dispersion forces.

Pro Tip for London Forces

In very large molecules, London forces can be stronger than dipole-dipole forces.

Quick Decision Flow for IMFs

Determine if N, O, or F is bonded to H for hydrogen bonding; check for polarity for dipole-dipole.

Polar molecule

A molecule that has a net dipole moment due to the presence of polar bonds.

Solid state

A state of matter where molecules are closely packed and vibrate in fixed positions, leading to a definite shape and volume.

Liquid state

A state of matter where molecules are close together but can move around each other, resulting in a definite volume but no definite shape.

Viscosity

A measure of a fluid's resistance to flow, influenced by intermolecular forces.

Boiling point

The temperature at which a liquid's vapor pressure equals the external pressure surrounding the liquid.

Melting point

The temperature at which a solid becomes a liquid.

Stronger IMFs

Intermolecular forces such as hydrogen bonding, ionic bonds, and network covalent bonds that lead to higher melting and boiling points.

Weaker IMFs

Intermolecular forces such as London dispersion that lead to lower melting and boiling points.

Solubility

The ability of a substance to dissolve in a solvent, often described by the phrase 'like dissolves like'.

H₂O (water)

A polar molecule with hydrogen bonding that has an even higher boiling point than methanol.

Solubility in polar solvents

Polar molecules tend to dissolve in polar solvents, especially those that can form hydrogen bonds.

Metallic solids

Solids held together by metallic bonds, characterized by a sea of delocalized electrons.

Covalent solids

Solids where atoms are held together by covalent bonds, forming a network structure.

Ionic solids

Solids formed from ionic bonds between positively and negatively charged ions.

Molecular solids

Solids composed of molecules held together by intermolecular forces, such as van der Waals forces.

Metallic bonds

Bonds that occur between metal atoms, where electrons are delocalized and free to move across the entire solid.

Sea of electrons

The delocalized electrons in metallic bonds that hold the metal atoms together.

Good conductors of heat and electricity

A property of metallic solids due to the free movement of electrons.

Malleable and ductile

Properties of metals that allow them to be hammered into shapes or drawn into wires.

High melting and boiling points

A property of metallic solids due to the strong attraction between metal atoms and free electrons.

Shiny appearance

A property of metallic solids due to the movement of electrons.

Examples of metallic solids

Na (sodium), Cu (copper), Fe (iron).

Ionic bonds

Bonds that form between positively charged ions (cations) and negatively charged ions (anions) through electrostatic attraction.

High melting and boiling points (ionic)

A property of ionic solids due to strong electrostatic forces.

Brittle

A property of ionic solids where applying pressure can cause repulsion between ions.

Conduct electricity when dissolved in water or melted

A property of ionic solids where ions can move freely in liquid or aqueous state.

Hard and rigid

A property of ionic solids due to the strong attraction between ions.

Examples of ionic solids

NaCl (sodium chloride), MgO (magnesium oxide), CaF₂ (calcium fluoride).

Covalent bonds

Bonds that form when atoms share electrons to achieve a full outer electron shell.

Covalent solids (Network Solids)

Solids where atoms are connected by an extended network of covalent bonds.

Extremely high melting and boiling points (covalent)

A property of covalent solids due to very strong covalent bonds throughout the entire structure.

Very hard

A property of covalent solids due to the strong bonds.

Poor conductors of electricity

A property of covalent solids because there are no free electrons or ions.

Insoluble in water

A property of covalent solids due to the strength of the covalent bonds.

Examples of covalent solids

Diamond (pure carbon), Graphite (also pure carbon, but different structure), SiO₂ (silicon dioxide) (quartz).

Intermolecular forces

Forces that hold molecular solids together, including London dispersion forces, dipole-dipole interactions, or hydrogen bonds.

Low melting and boiling points (molecular)

A property of molecular solids because the forces between molecules are much weaker than covalent or ionic bonds.

Soft and brittle (molecular)

A property of molecular solids where molecules can easily slip past each other.

Nonconductive (molecular)

A property of molecular solids because there are no free electrons or ions.

Soluble in nonpolar solvents

A property of molecular solids depending on the type of intermolecular force.

Examples of molecular solids

H₂O (water) (hydrogen bonds), CO₂ (carbon dioxide) (London dispersion), I₂ (iodine) (London dispersion).

Example of metallic solid

Copper (Cu), Iron (Fe), Sodium (Na)

Example of ionic solid

Sodium chloride (NaCl), Magnesium oxide (MgO), Calcium fluoride (CaF₂)

Covalent bonds (network solids)

If the substance is made of atoms connected by covalent bonds that form a large, continuous network, it's a network solid. These solids are very hard, brittle, and have very high melting points.

Example of covalent solid

Diamond (C), Silicon dioxide (SiO₂), Quartz (SiO₂)

Molecular forces (molecular solids)

If the substance is made of molecules held together by intermolecular forces (like London dispersion, dipole-dipole, or hydrogen bonding), it's a molecular solid.