Chapter 11: Intermolecular Forces, Liquids, and Solids

Chapter 11: Intermolecular Forces, Liquids, and Solids

intermolecular forces – forces that exist between molecules

11.1 A Molecular Comparison of Liquids and Solids

gases

average kinetic energy of the molecules is larger than average energy of attractions between molecules
lack of strong attractive forces allows gases to expand

liquids

denser than gases
have a definite volume
attractive forces not strong enough to keep molecules from moving allowing liquids to hold shape of container

solids

intermolecular forces hold molecules together and keep them from moving
not very compressible
crystalline – solids with highly ordered structures

Gas	assumes both the volume and shape of container is compressible diffusion within a gas occurs rapidly flows readily
Liquid	Assumes the shape of the portion of the container it occupies Does not expand to fill container Is virtually incompressible Diffusion within a liquid occurs slowly Flows readily
Solid	Retains its own shape and volume Is virtually incompressible Diffusion within a solid occurs extremely slowly Does not flow

state of substance depends on balance between the kinetic energies of the particles and interparticle energies of attraction

kinetic energies depends on temperature and tend to keep particles apart and moving

interparticle attractions draw particles together

condensed phases – liquids and solids because particles are close together compared to gases

increase temperature forces molecules to be closer together ® increase in strength of intermolecular forces

11.2 Intermolecular Forces

intermolecular forces weaker than ionic or covalent bonds
many properties of liquids reflect strengths of intermolecular forces
three types of intermolecular forces: dipole-dipole forces, London dispersion forces, and hydrogen-bonding forces

also called van der Waals forces
less than 15% as strong as covalent or ionic bonds
electrostatic in nature, involves attractions between positive and negative species

11.2.1 Ion-Dipole Forces

Ion-Dipole Force – exists between an ion and partial charge at one end of a polar molecule
Polar molecules are dipoles
magnitude of attraction increases as either the charge of ion or magnitude of dipole moment increases

11.2.2 Dipole-Dipole Forces

dipole-dipole force – exists between neutral polar molecules

effective only when polar molecules are very close together

weaker than ion-dipole forces

for molecules of approximately equal mass and size, the strengths of intermolecular attractions increase with increasing polarity

increase dipole moment ® increase boiling point

11.2.3 London Dispersion Forces

interparticle forces that exist between nonpolar atoms or molecules

motion of electrons can create an instantaneous dipole moment

molecules have to be very close together

polarizability – ease in which the charge distribution in a molecule can be distorted

greater polarizability ® more easily electron cloud can be distorted to give momentary dipole

larger molecules have greater polarizability

London dispersion forces increase with increasing molecular size

Dispersion forces increase in strength with increasing molecular weight

Molecular shape affects intermolecular attractions

greater surface contact ® greater boiling point and London dispersion forces

dispersion forces operate between all molecules

comparing relative strengths of intermolecular attractions:

1) comparable molecular weights and shapes = equal dispersion forces

differences in magnitudes of attractive forces due to differences in strengths of dipole-dipole attractions
most polar molecule has strongest attractions

2) differing molecular weights = dispersion forces tend to be the decisive ones

differences in magnitudes of attractive forces associated with differences in molecular weights
most massive molecular has strongest attractions

11.2.4 Hydrogen Bonding

hydrogen bonding – special type of intermolecular attraction that exists between the hydrogen atom in a polar bond and an unshared electron pair on a nearby electronegative ion or atom

hydrogen bond with F, N, and O is polar

density of ice is lower than that of liquid water

when water freezes the molecules assume the ordered open arrangement ® makes ice less dense than water

a given mass of ice has a greater volume than the same mass of water

structure of ice allows the maximum number of hydrogen bonding interactions to exist

11.2.5 Comparing Intermolecular Forces

dispersion forces found in all substances
strengths of forces increase with increases molecular weight and also depend on shape
dipole-dipole forces add to effect of dispersion forces and found in polar molecules
hydrogen bonds tend to be strongest intermolecular force

11.3 Some Properties of Liquids

two properties of liquids: viscosity and surface tension

11.3.1 Viscosity

viscosity – resistance of a liquid to flow
the greater the viscosity the more slowly the liquid flows
measured by timing how long it takes a certain amount of liquid to flow through a thin tube under gravitational forces
can also be measured by how long it takes steel spheres to fall through the liquid
viscosity related to ease with which individual molecules of liquid can move with respect to one another
depends on attractive forces between molecules, and whether structural features exist to cause molecules to be entangled
viscosity decreases with increasing temperature

11.3.2 Surface Tension

surface tension – energy required to increase the surface area of a liquid by a unit amount

surface tension of water at 20° C is 7.29 x 10^-2 J/m²

7.29 x 10^-2 J/m² must be supplied to increase surface area of a given amount of water by 1 m²

cohesive forces – intermolecular forces that bind similar molecules

adhesive forces – intermolecular forces that bind a substance to a surface

capillary action – rise of liquids up very narrow tubes

11.4 Phase Changes

11.4.1 Energy Changes Accompanying Phase Changes

phase changes require energy

phase changes to less ordered state requires energy

melting process of solid called fusion

heat of fusion – enthalpy change of melting a solid

D H_fus water = 6.01 kJ/mol

heat of vaporization – heat needed for vaporization of liquid

D H_vap water = 40.67 kJ/mol

melting, vaporization, and sublimation are endothermic

freezing, condensation, and deposition are exothermic

11.4.2 Heating Curves

heating curve – graph of temperature of system versus the amount of heat added

used to calculate enthalpy changes

supercooled water – when water if cooled to a temperature below 0° C

11.4.3 Critical Temperature and Pressure

critical temperature – highest temperature at which a substance can exist as a liquid

critical pressure – pressure required to bring about liquefaction at critical temperature

the greater the intermolecular attractive forces, the more readily gases liquefy ® higher critical temperature

cannot liquefy a gas by applying pressure if gas is above critical temperature

11.5 Vapor Pressure

vapor pressure – measures tendency of a liquid to evaporate

11.5.1 Explaining Vapor Pressure on the Molecular Level

dynamic equilibrium – condition when two opposing processes are occurring simultaneously at equal rates
vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor states are in dynamic equilibrium

11.5.2 Volatility, Vapor Pressure, and Temperature

volatile – liquids that evaporate readily
vapor pressure increases with increasing temperature

11.5.3 Vapor Pressure and Boiling Point

liquids boil when its vapor pressure equals the external pressure acting on the surface of the liquid
temperature of boiling increase with increasing external pressure
normal boiling point – boiling point of a liquid at 1 atm
higher pressures cause water to boil at higher temperatures

11.6 Phase Diagrams

phase diagrams – graphical way to summarize conditions under which equilibria exist between the different states of matter
three important curves:

1) vapor pressure curve of liquid
shows equilibrium of liquid and gas phases
normal boiling point = point on curve where pressure at 1 atm
2) variation in vapor pressure of solid at it sublimes at different temperatures
3) change in melting point of solid with increasing pressure
slopes right as pressure increases
higher temperatures needed to melt solids at higher pressures
melting point of solid identical to freezing point

differ only in temperature direction from which phase change is approached
melting point at 1 atm is the normal melting point

triple point – point at which all three phases are at equilibrium
gas phase stable at low pressures and high temperatures
solid phase stable at low temperatures and high pressures
liquid phase – stable between gas and solids

11.6.1 the Phase diagrams of H₂O and CO₂

melting point of CO₂ increases with increasing pressure

melting point of H₂O decreases with increasing pressure

triple point of H₂O (0.0098° C and 4.58 torr) at lower pressure than CO₂ (-56.4° C and 5.11 atm)

solid CO₂ does not melt but sublimes

CO₂ does not have a normal melting point but a normal sublimation point

CO₂ absorbs energy at ordinary temperatures

11.7 Structures of Solids

crystalline solid – solid whose atoms, ion, or molecules are ordered in well-defined arrangements

flat surfaces or faces that make definite angles
regular shapes

amorphous solid – solid whose particles have no orderly structure

lack well-defined faces and shapes
mixtures of molecules that do not stack together well
large, complicated molecules
intermolecular forces vary in strength
does not melt at a specific temperature but soften over a temperature range

11.7.1 Unit Cell

unit cell – repeating unit of a solid
crystal lattice – three-dimensional array of points, each representing an identical environment within the crystal
three types of cubic unit cell: primitive cubic, body-centered cubic, and face-centered cubic
primitive cubic – lattice points at corners only
body-centered cubic – lattice points at corners and center
face-centered cubic – lattice points at center of each face and at each corner

11.7.2 The Crystal structure of Sodium Chloride

total cation-to-anion ratio of a unit cell must be the same as that for entire crystal

11.7.3 Close Packing of Spheres

structures of crystalline solids are those that bring particles in closest contact to maximize the attractive forces
most particles that make up solids are spherical
two forms of close packing: cubic close packing and hexagonal close packing
hexagonal close packing – spheres of the third layer that are placed in line with those of the first layer
coordination number – number of particles immediately surrounding a particle in the crystal structure
both forms of close packing have coordination number of 12

11.8 Bonding in Solids

Type of Solid	Forms of Unit Particles	Forces Between Particles	Properties	Examples
Molecular	Atoms of molecules	London dispersion, dipole-dipole forces, hydrogen bonds	Fairly soft, low to moderately high melting point, poor thermal and electrical conduction	Argon, methane, sucrose, dry ice
Covalent-network	Atoms connected in a network of covalent bonds	Covalent bonds	Very hard, very high melting point, often poor thermal and electrical conduction	Diamond, quartz
Ionic	Positive and negative ions	Electrostatic attractions	Hard and brittle, high melting point, poor thermal and electrical conduction	Typical salts
Metallic	atoms	Metallic bonds	Soft to very hard, low to very high melting point, excellent thermal and electrical conduction, malleable and ductile	All metallic elements

11.8.1 Molecular Solids

molecular solids – atoms or molecules held together by intermolecular forces
soft, low melting points
gases or liquids at room temperature from molecular solids at low temperature
properties depends on strengths of forces and ability of molecules to pack efficiently in three dimensions
intermolecular forces that depend on close contact are not as effective

11.8.2 Covalent-Network Solids

covalent-network solids – atoms held together in large networks or chains by covalent bonds

hard, high melting points

11.8.3 Ionic Solids

ionic solids – ions held together by ionic bonds
strength depends on charges of ions
structure of ionic solids depends on charges and relative sizes of ions

11.8.4 Metallic Solids

metallic solids – metal atoms
usually have hexagonal close-packed, cubic close-packed, or body-centered-cubic structures
each atom has 8 or 12 adjacent atoms
bonding due to valence electrons that are delocalized throughout entire solid
strength of bonding increases as number of electrons available for bonding increases
mobility of electrons make metallic solids good conductors of heat and electricity

11.2.5 (continued)