Liquids and Solids by Owen Borville October 24, 2025
Intermolecular forces and the Physical Properties of Liquids and Solids:
The condensed phases: Intermolecular forces are attractive forces that hold particles together in the condensed phases. The magnitude and type of intermolecular forces is what determines whether the particles that make up a substance are gas, liquid, or solid.
Surface tension is the amount of energy required to stretch or increase the surface of any liquid by a unit area. The stronger the intermolecular forces, the higher the surface tension.
Capillary action is the movement of a liquid up a narrow tube. Two types of forces bring about capillary action: Cohesion is the attraction between like molecules. Adhesion is the attraction between unlike molecules.
Viscosity is a measure of a fluid's resistance to flow. The higher the viscosity the more slowly a liquid flows. Liquids that have higher intermolecular forces have higher viscosities.
Vapor pressure of liquids is also dependent on intermolecular forces. If a molecule at the surface of a liquid has enough kinetic energy, it can escape to the gas phase in a process called vaporization.
The vapor pressure increases until the rate of evaporation equals the rate of condensation:
H2O(l) ⇌ H2O(g)
Evaporation: H2O (l) => H2O(g)
Condensation: H2O(l) <= H2O(g)
When the forward process and the reverse process are occurring at the same rate, the system is in dynamic equilibrium.
The vapor pressure increases until the rate of evaporation equals the rate of condensation.
H2O(l) ⇌ H2O(g)
The vapor pressure increases with temperature.
The Clausius-Clapeyron equation relates the natural log of vapor pressure and the reciprocal of absolute temperature:
InP= - ΔHvap/RT + C
InP = natural log of vapor pressure
ΔHvap = the molar heat of vaporization
R = the gas constant (8.314 J/K*mol)
T = the Kelvin temperature
C is an experimentally determines constant
The Clausius-Clapeyron equation: InP = ΔHvap/RT + C
Plotting In P versus 1/T is a straight line moving downward with a slope of -ΔH/R
ΔH is assumed to be independent of temperature.
The Clausius-Clapeyron equation can be rearranged into a two-point form:
In(P1/P2) = ΔHvap/R(1/T2-1/T1)
Boiling Point is the temperature at which the vapor pressure equals the external atmospheric pressure. Boiling point varies with external pressure. Boiling point varies with the magnitude of intermolecular forces.
Properties of Solids: Melting Point is the temperature at which the energies of the individual particles enable them to break free of their fixed positions.
Vapor pressures of solids are usually very low at room temperature, with the exception of certain materials like moth balls, naphthalene, dry ice (solid carbon dioxide), and ice.
Amorphous Solids and Crystalline Solids: Non-crystalline (amorphous) quartz glass (no ordered internal structure) versus Crystalline glass (ordered internal atomic structure).
Amorphous solids lack a regular three-dimensional arrangement of atoms. Glass is an amorphous solid and a fusion product.
Crystalline solids possess rigid and long-range order. Its atoms, molecules, or ions occupy specific positions. Unit cells are the basic repeating structural unit of a crystalline solid. There are seven types of unit cells:
Simple cubic
Tetragonal
Orthorhombic
Rhombohedral
Monoclinic
Triclinic
Hexagonal
Coordination number is defined as the number of atoms surrounding an atom in a crystal lattice. The value of the coordination number indicates how tightly the atoms are packed together. The basic repeating unit in the array of atoms is called a simple cubic cell.
There are three types of cubic cells: Primitive cubic, body-centered cubic, and face-centered cubic.
In body-centered cubic cells (bcc), the spheres in each layer rest in the depressions between spheres in the previous layer. The coordination number is 8.
In a face-centered cubic cell (fcc), the coordination number is 12.
Most of a cell's atoms are shared by neighboring cells. A corner atom is shared by eight unit cells. An edge atom is shared by four unit cells. A face-centered atom is shared by two unit cells.
A simple cubic cell has the equivalent of only one complete atom contained within the cell and 8 atoms at corners.
A body-centered cubic cell has two equivalent atoms: 8 atoms at corners = one equivalent atom, plus one atom at the center = 2 equivalent atoms.
A face-centered cubic cell contains four complete atoms: 8 atoms at corners = one equivalent atom + 6 atoms on faces = 3 equivalent atoms =>4 equivalent atoms total.
Hexagonal close-packed (hcp) structure: Close packing starts with a layer of atoms. Atoms in the second layer fit into the depressions of the first layer. Site directly over an atom in first layer.
Cubic close-packed (ccp) structure: site NOT directly over an atom in first layer.
Closest-packing = hexagonal close-packing (hcp) => to cubic close-packing (ccp) corresponds to a face-centered cubic cell.
In crystalline solids, edge length (a) and radius (r) are related in simple cubic, body-centered cubic, and face-centered cubic structure.
Ionic crystals are composed of charged ions that are held together by Coulombic attraction. The unit cell of an ionic compound can be defined by either the positions of the anions or the positions of the cations. Crystal structures of three ionic compounds are simple cubic lattice (CsCl), Zincblende structure (ZnS), and fluorite structure (CaF2).
Covalent crystals feature atoms that are held together in an extensive three-dimensional network entirely by covalent bonds.
Molecular crystals feature lattice points that are occupied by molecules. The attractive forces between them are van der Waals forces and/or hydrogen bonding.
Metallic crystals feature every lattice point occupied by an atom of the same metal. Electrons are delocalized over the entire crystal. Delocalized electrons make metals good conductors. Large cohesive forces resulting from delocalization makes metals strong.
Types of Crystals and Their General Properties:
Ionic Crystals: Cohesive forces are Coulombic attraction and dispersion forces. Hard, brittle, high melting point, poor conductor of heat and electricity. Ex. NaCl, LiF, MgO, CaCO3
Covalent Crystals: Cohesive forces are covalent Bonds. Hard, brittle, high melting point, poor conductors of heat and electricity. Ex. C (diamond), SiO2(quartz)
Molecular Crystals: Cohesive forces are dispersion and dipole-dipole forces, hydrogen bonds. Soft, low melting point, poor conductors of heat and electricity. Ex. Ar, CO2, I2, H2O, C12H22O11
Metallic Crystals: Cohesive forces are metallic bonds. Variable hardness and melting point. Good conductors of heat and electricity. Ex. All metallic elements including Na, Mg, Fe, Cu.
Phase Changes: A phase is a homogenous part of a system that is separated from the rest of the system by a well defined boundary. When a substance goes from one phase to another phase, it has undergone a phase change.
Examples of Phase Changes:
H2O(l) => H2O(s) Freezing of water
H2O(l) => H2O(g) Evaporation (or vaporization) of water
H2O(s) => H2O(l) Melting or fusion of ice
H2O(g) => H2O(l) Condensation of water vapor
CO2(s) => CO2(g) Sublimation of dry ice
Six Possible Phase Changes Are:
Melting (fusion)= solid to liquid, temperature rising
Vaporization= liquid to gas, temperature rising
Sublimation= solid to gas, temperature rising
Deposition= gas to solid, temperature decreasing
Condensation= gas to liquid, temperature decreasing
Freezing= liquid to solid, temperature decreasing
Boiling point of a substance is the temperature at which its vapor pressure equals the external atmospheric pressure.
Molar heat of vaporization (ΔHvap) is the amount of heat required to vaporize a mole of substance at its boiling point.
Critical temperature of a substance is the temperature above which its gas cannot be liquified.
Critical pressure is the minimum pressure that must be applied to liquefy a substance at its critical temperature.
The transformation of a liquid to a solid is called freezing. The reverse process is called melting or fusion.
The melting point (freezing point) of a solid (or liquid) is the temperature at which the solid and liquid phases coexist in equilibrium.
ice ⇌ water
H2O(s) ⇌ H2O(l)
In dynamic equilibrium, the forward and reverse processes are occurring at the same time. These relationships can be shown graphically using heating curves plotted with temperature and time.
Molar Heat of Fusion (ΔHfus) is the energy required to melt 1 mol of a solid.
Sublimation is the process by which molecules go directly from the solid phase to the vapor phase.
Deposition is the reverse process of sublimation. Deposition is the transition from vapor to solid phase.
Molar enthalpy of sublimation (ΔHsub) of a substance is the energy required to sublime 1 mol of a solid.
ΔHsub = ΔHfus + ΔHvap
Phase diagrams are graphical charts that summarize the conditions at which a substance exists as a solid, liquid or gas. Pressure and temperature are plotted on the axes and the regions of solid, liquid, and gas are shown on the chart. Triple point is the only combination of pressure and temperature where three phases of a substance exist in equilibrium on the phase diagram.
Intermolecular forces and the Physical Properties of Liquids and Solids:
The condensed phases: Intermolecular forces are attractive forces that hold particles together in the condensed phases. The magnitude and type of intermolecular forces is what determines whether the particles that make up a substance are gas, liquid, or solid.
Surface tension is the amount of energy required to stretch or increase the surface of any liquid by a unit area. The stronger the intermolecular forces, the higher the surface tension.
Capillary action is the movement of a liquid up a narrow tube. Two types of forces bring about capillary action: Cohesion is the attraction between like molecules. Adhesion is the attraction between unlike molecules.
Viscosity is a measure of a fluid's resistance to flow. The higher the viscosity the more slowly a liquid flows. Liquids that have higher intermolecular forces have higher viscosities.
Vapor pressure of liquids is also dependent on intermolecular forces. If a molecule at the surface of a liquid has enough kinetic energy, it can escape to the gas phase in a process called vaporization.
The vapor pressure increases until the rate of evaporation equals the rate of condensation:
H2O(l) ⇌ H2O(g)
Evaporation: H2O (l) => H2O(g)
Condensation: H2O(l) <= H2O(g)
When the forward process and the reverse process are occurring at the same rate, the system is in dynamic equilibrium.
The vapor pressure increases until the rate of evaporation equals the rate of condensation.
H2O(l) ⇌ H2O(g)
The vapor pressure increases with temperature.
The Clausius-Clapeyron equation relates the natural log of vapor pressure and the reciprocal of absolute temperature:
InP= - ΔHvap/RT + C
InP = natural log of vapor pressure
ΔHvap = the molar heat of vaporization
R = the gas constant (8.314 J/K*mol)
T = the Kelvin temperature
C is an experimentally determines constant
The Clausius-Clapeyron equation: InP = ΔHvap/RT + C
Plotting In P versus 1/T is a straight line moving downward with a slope of -ΔH/R
ΔH is assumed to be independent of temperature.
The Clausius-Clapeyron equation can be rearranged into a two-point form:
In(P1/P2) = ΔHvap/R(1/T2-1/T1)
Boiling Point is the temperature at which the vapor pressure equals the external atmospheric pressure. Boiling point varies with external pressure. Boiling point varies with the magnitude of intermolecular forces.
Properties of Solids: Melting Point is the temperature at which the energies of the individual particles enable them to break free of their fixed positions.
Vapor pressures of solids are usually very low at room temperature, with the exception of certain materials like moth balls, naphthalene, dry ice (solid carbon dioxide), and ice.
Amorphous Solids and Crystalline Solids: Non-crystalline (amorphous) quartz glass (no ordered internal structure) versus Crystalline glass (ordered internal atomic structure).
Amorphous solids lack a regular three-dimensional arrangement of atoms. Glass is an amorphous solid and a fusion product.
Crystalline solids possess rigid and long-range order. Its atoms, molecules, or ions occupy specific positions. Unit cells are the basic repeating structural unit of a crystalline solid. There are seven types of unit cells:
Simple cubic
Tetragonal
Orthorhombic
Rhombohedral
Monoclinic
Triclinic
Hexagonal
Coordination number is defined as the number of atoms surrounding an atom in a crystal lattice. The value of the coordination number indicates how tightly the atoms are packed together. The basic repeating unit in the array of atoms is called a simple cubic cell.
There are three types of cubic cells: Primitive cubic, body-centered cubic, and face-centered cubic.
In body-centered cubic cells (bcc), the spheres in each layer rest in the depressions between spheres in the previous layer. The coordination number is 8.
In a face-centered cubic cell (fcc), the coordination number is 12.
Most of a cell's atoms are shared by neighboring cells. A corner atom is shared by eight unit cells. An edge atom is shared by four unit cells. A face-centered atom is shared by two unit cells.
A simple cubic cell has the equivalent of only one complete atom contained within the cell and 8 atoms at corners.
A body-centered cubic cell has two equivalent atoms: 8 atoms at corners = one equivalent atom, plus one atom at the center = 2 equivalent atoms.
A face-centered cubic cell contains four complete atoms: 8 atoms at corners = one equivalent atom + 6 atoms on faces = 3 equivalent atoms =>4 equivalent atoms total.
Hexagonal close-packed (hcp) structure: Close packing starts with a layer of atoms. Atoms in the second layer fit into the depressions of the first layer. Site directly over an atom in first layer.
Cubic close-packed (ccp) structure: site NOT directly over an atom in first layer.
Closest-packing = hexagonal close-packing (hcp) => to cubic close-packing (ccp) corresponds to a face-centered cubic cell.
In crystalline solids, edge length (a) and radius (r) are related in simple cubic, body-centered cubic, and face-centered cubic structure.
Ionic crystals are composed of charged ions that are held together by Coulombic attraction. The unit cell of an ionic compound can be defined by either the positions of the anions or the positions of the cations. Crystal structures of three ionic compounds are simple cubic lattice (CsCl), Zincblende structure (ZnS), and fluorite structure (CaF2).
Covalent crystals feature atoms that are held together in an extensive three-dimensional network entirely by covalent bonds.
Molecular crystals feature lattice points that are occupied by molecules. The attractive forces between them are van der Waals forces and/or hydrogen bonding.
Metallic crystals feature every lattice point occupied by an atom of the same metal. Electrons are delocalized over the entire crystal. Delocalized electrons make metals good conductors. Large cohesive forces resulting from delocalization makes metals strong.
Types of Crystals and Their General Properties:
Ionic Crystals: Cohesive forces are Coulombic attraction and dispersion forces. Hard, brittle, high melting point, poor conductor of heat and electricity. Ex. NaCl, LiF, MgO, CaCO3
Covalent Crystals: Cohesive forces are covalent Bonds. Hard, brittle, high melting point, poor conductors of heat and electricity. Ex. C (diamond), SiO2(quartz)
Molecular Crystals: Cohesive forces are dispersion and dipole-dipole forces, hydrogen bonds. Soft, low melting point, poor conductors of heat and electricity. Ex. Ar, CO2, I2, H2O, C12H22O11
Metallic Crystals: Cohesive forces are metallic bonds. Variable hardness and melting point. Good conductors of heat and electricity. Ex. All metallic elements including Na, Mg, Fe, Cu.
Phase Changes: A phase is a homogenous part of a system that is separated from the rest of the system by a well defined boundary. When a substance goes from one phase to another phase, it has undergone a phase change.
Examples of Phase Changes:
H2O(l) => H2O(s) Freezing of water
H2O(l) => H2O(g) Evaporation (or vaporization) of water
H2O(s) => H2O(l) Melting or fusion of ice
H2O(g) => H2O(l) Condensation of water vapor
CO2(s) => CO2(g) Sublimation of dry ice
Six Possible Phase Changes Are:
Melting (fusion)= solid to liquid, temperature rising
Vaporization= liquid to gas, temperature rising
Sublimation= solid to gas, temperature rising
Deposition= gas to solid, temperature decreasing
Condensation= gas to liquid, temperature decreasing
Freezing= liquid to solid, temperature decreasing
Boiling point of a substance is the temperature at which its vapor pressure equals the external atmospheric pressure.
Molar heat of vaporization (ΔHvap) is the amount of heat required to vaporize a mole of substance at its boiling point.
Critical temperature of a substance is the temperature above which its gas cannot be liquified.
Critical pressure is the minimum pressure that must be applied to liquefy a substance at its critical temperature.
The transformation of a liquid to a solid is called freezing. The reverse process is called melting or fusion.
The melting point (freezing point) of a solid (or liquid) is the temperature at which the solid and liquid phases coexist in equilibrium.
ice ⇌ water
H2O(s) ⇌ H2O(l)
In dynamic equilibrium, the forward and reverse processes are occurring at the same time. These relationships can be shown graphically using heating curves plotted with temperature and time.
Molar Heat of Fusion (ΔHfus) is the energy required to melt 1 mol of a solid.
Sublimation is the process by which molecules go directly from the solid phase to the vapor phase.
Deposition is the reverse process of sublimation. Deposition is the transition from vapor to solid phase.
Molar enthalpy of sublimation (ΔHsub) of a substance is the energy required to sublime 1 mol of a solid.
ΔHsub = ΔHfus + ΔHvap
Phase diagrams are graphical charts that summarize the conditions at which a substance exists as a solid, liquid or gas. Pressure and temperature are plotted on the axes and the regions of solid, liquid, and gas are shown on the chart. Triple point is the only combination of pressure and temperature where three phases of a substance exist in equilibrium on the phase diagram.
