Liquids and Solids | General Chemistry 2

Liquids and solids are studied in this chapter: intermolecular forces, phase transitions, heating curves, phase diagrams, properties of liquids, vapor pressure, relative humidity, crystal structures and forces, colloids.

Intermolecular Forces

Intramolecular forces: forces within a molecule (i.e. chemical bonds)
Intermolecular forces: forces between molecules


Different types of intermolecular forces:

- Ion-dipole (40-600 kJ.mol-1)
- Hydrogen bond (10-40 kJ.mol-1)
- Dipole-dipole (5-25 kJ.mol-1)
- Ion-induced dipole (3-15 kJ.mol-1)
- Dipole-induced dipole (2-10 kJ.mol-1)
- London dispersion (0.05-40 kJ.mol-1)

 

Hydrogen bond: electrostatic interaction that occurs between a hydrogen atom bonded to N, O or F and an electronegative atom that has a lone pair of electrons
 

 

Dipoles: molecules with a molecular dipole
Induced dipoles: nonpolar molecules in which a temporary dipole moment has been induced by a charged or polar molecule
 

Isolated O2 molecule is nonpolar
When a polar water molecule approaches, the electron distribution in O2 is rearranged
⇒ O2 is now polar ⇒ this is an induced dipole

 

London dispersion forces: electrostatic interaction that occurs between two nonpolar molecules
Electrons within a molecule are constantly moving around
⇒ at any given moment: asymmetrically distribution of these electrons
⇒ creation of an instantaneous dipole moment
⇒ induced dipole on a neighboring nonpolar molecule

Phase Transitions

Endothermic phase transitions:

Solid to Liquid: fusion/melting (ΔH0fus)
Liquid to Gas: vaporization (ΔH0vap)
Solid to Gas: sublimation (ΔH0sub = ΔHfus + ΔHvap ⇒ Hess’s law)

 

Exothermic phase transitions:

Liquid to Solid: freezing (-ΔH0fus)
Gas to Liquid: condensation (-ΔH0vap)
Gas to Solid: deposition (-ΔH0sub)

Heating Curves

 

Segment A:

substance is a solid
As heat is added to the system, T increases according to: 

q = n CΔT

n = number of moles (in mol)
CP = molar heat capacity of the solid at constant pressure (in J.K-1.mol-1)

 

Segment B:

phase transition from solid to liquid
As heat is added to the system, T stays the same:

q = n ΔH0fus

n = number of moles (in mol)
ΔH0fus = enthalpy change for fusion (in J.mol-1)

 

Segment C:

substance is a liquid, as heat is added to the system, T increases according to:

q = n CP' ΔT

n = number of moles (in mol)
CP' = molar heat capacity of the liquid (in J.K-1.mol-1) ≠ CP from segment A

 

Segment D:

phase transition from liquid to gas
As heat is added to the system, T stays the same:

q = n ΔH0vap

n = number of moles (in mol)
ΔH0vap = enthalpy change for vaporisation (in J.mol-1)

 

Segment E:

substance is a liquid
As heat is added to the system, T increases according to:

q = nCP''ΔT

n = number of moles (in mol)
CP'' = molar heat capacity of the gas (in J.K-1.mol-1

 

 

Phase Diagrams

Phase diagram: diagram which displays the regions of all phases of a pure substance

 


Triple point: all 3 phases coexist in equilibrium
Critical point: point at which 2 phases become indistinguishable from one another

Properties of Liquids

Viscosity η:

Measure of the resistance of a liquid to flow
Stronger intermolecular forces = more viscous
Higher temperature = less viscous
Long and flexible molecules have higher viscosities than shorter-chain or spherical molecules

 

Surface Tension (in J.m2):

Energy required to increase the surface area of a liquid
Stronger intermolecular forces = higher surface tension
Surfactants reduce surface tension by decreasing the intermolecular forces between adjacent molecule at the surface

 

Capillary Action:

Tendency of a liquid to rise against gravity up a small-diameter tube (a capillary)
⇒ adhesive forces

Vapor Pressure

Pressure exerted by the vapor on the liquid
The pressure increases until equilibrium is reached (rate of evaporation = rate of condensation)
At equilibrium, the vapor pressure is constant (= equilibrium vapor pressure)
A liquid has a unique equilibrium vapor pressure at each temperature


Normal boiling point: temperature at which vapor pressure = atmospheric pressure (1 atm)

Volatile substance: substance with a high vapor pressure at normal temperature

Relative Humidity

Relative humidity φ expressed as a percentage:
 

φ = PH2OP0H2O x 100

PH2O = partial pressure of the water vapor in the air
P0H2O = equilibrium vapor pressure of water at the same temperature

 

P0H2O decreases when T decreases ⇒ φ increases when T decreases

Dew point: temperature at which the relative humidity φ = 100%

Crystal Structures

Crystal: solid material whose constituents (atoms, molecules, ions) are arranged in a highly ordered microscopic structure ⇒ crystal lattice

Unit cell: smallest subunit of a crystal lattice that can be used to generate the entire lattice

- simple cubic: 1 atom per unit cell (at the corner)
length of an edge = 2 x radii of the atom
- body-centered cubic: 2 atoms per unit cell (at the corner and at the center of the unit cell)
length of the main diagonal = 4 x crystallographic radii
- face-centered cubic: 4 atoms per unit (at the corner and at the center of each face)
Length of the diagonal of a face = 4 x crystallographic radii

Crystal Forces

Crystals can be classified according to the forces between the constituent particles:

- ionic crystal:
coulombic charge-charge attractions between cations and anions (ex: NaCl (s))
⇒ hard and brittle, high melting point, poor electrical conductor

- molecular crystal:
van der Waals interactions between molecules (ex: I2 (s))
⇒ soft, low melting point, poor electrical conductor

- network crystal:
extended network of covalent bonds between atoms (ex: C(s), diamond)
⇒ very hard, high melting point

- amorphous:
various attractive forces between groups of molecules (ex: plastic)
⇒ no sharp melting point

- metallic crystals: cations at lattice points and delocalized electrons (ex: Ag(s), Cu(s))
⇒ good electrical conductor

Colloids

Colloid: mixture of a solvent and suspended particles
⇒ heterogeneous mixture but particles are too small to be seen. Particle size in colloids: 1 to 100 nm ⇒ larger than simple molecules, but small enough to remain in suspension and not settle out


Homogeneous mixture: solution, particle size <1 nm
Heterogeneous mixture: suspension, particle size > 100 nm