Chemistry A Molecular Approach AP Edition Chapter 11 Liquids, Solids, and Intermolecular Forces 

11.1 Climbing Geckos and Intermolecular Forces 

• There are three states, or phases, of matter: solid, liquid, and gas
• Solids and liquids are condensed states
• Solids and liquids are more similar to each other than they are to gases 
• Particles are closer together and have stronger attractive forces 
• Intermolecular forces are the attractive forces between all atoms and molecules 
• Intermolecular forces hold solids and liquids together 
• Intermolecular forces allow geckos to stick to walls 
• When thermal energy is high, matter is gaseous 
• When thermal energy is low, matter is a liquid or a solid 

11.2 Solids, Liquids, and Gases: A Molecular Comparison 

• The densities of solids and liquids are greater than the densities of gases 
• Solids and liquids have similar density and molar volumes 
• The molecules in solids and liquids are very close together 
• The molecules in gases are very far apart for their size
• In liquids, unlike solids, molecules have the freedom to move around each other
• Liquids assume the shape of their containers because the molecules are free to flow 
• Liquids and solids are not easily compressed because the molecules cannot be pushed any closer together
• Gases have large spaces in between molecules, and if the pressure is great enough they can be easily compressed 
• Crystalline solids are solids thats molecules are arranged in well ordered 3D arrays 
• Amorphous solids are solids that have to long range order to their molecules

Changes between States 

• States of matter can be changed by using temperature or pressure 
• Increasing pressure causes more density, which can change a gas into a liquid 
• Gases in their liquid form will occupy less space 

11.3 Intermolecular Forces: The Forces That Hold Condensed States Together 

• Intermolecular forces determine if something is a solid, liquid, or gas at a given temperature 
• Moderate to strong intermolecular forces are often liquids and solids at room temperature 
• Weak intermolecular forces tend to be gases at room temperature 
• Molecules with temporary charges are attracted to each other because the potential energy decreases as distance decreases 
• Intermolecular forces are weaker than bonding forces 
• Intermolecular forces are between molecules 

Dispersion Forces 

• Dispersion forces are also called London Dispersion forces
• Dispersion forces are present in all molecules and atoms 
• All atoms have electrons, so they have dispersion forces 
• Dispersion forces are a temporary change that comes from the electrons being unevenly shared 
• Charge separation is called an instantaneous or a temporary dipole 
• The positive end of an atom is attracted to the negative end of another atom. This attraction is the dispersion force 
• Polarize is the formation of a dipole moment 
• The magnitude of the force depends on how easy electrons can move or polarize 
• The movement and polarization of electrons depend on how large the electron cloud is 
• In larger electron clouds, electrons are not held as tightly and can polarize easier 
• Dispersion forces increase as molar mass increases 
• Shape also influences dispersion forces  

Dipole Dipole Force

• Dipole dipole forces can be found in all polar molecules 
• Polar molecules have partial negative and partial positive regions 
• Permanent dipoles do not switch from being positive or negative. They are always one or the other 
• The positive end of a permanent dipole being attracted to the negative end of another permanent dipole is the dipole dipole force 
• Polar molecules have higher melting and boiling points that nonpolar molecules (of the same mass) 
• Miscibility is the ability to mix without separating into two states, for example water and pentane 
• Polarity determines miscibility 
• Polar molecules are miscible with polar molecules. They are not miscible with nonpolar molecules 
• Nonpolar molecules are miscible with nonpolar molecules 

Hydrogen Bonding 

• Polar molecules that have hydrogen atoms can have hydrogen bonding 
• The hydrogen atoms will typically be connected to fluorine, oxygen, or nitrogen atoms 
• Hydrogen atoms will have partial positive charges while the other atoms will have partial negative charges 
• The strong attractions between these atoms is a hydrogen bond 
• Hydrogen bonds are NOT chemical bonds 
• Hydrogen bonds are the stronger than dipole dipole and dispersion forces 

Ion Dipole Forces 

• Ion dipoles occur when an ionic compound is mixed with a polar compound 
• Ion dipoles are important in aqueous solutions 
• Ion dipole forces are the strongest type of intermolecular force
  

11.4 Intermolecular Forces in Action: Surface Tension, Viscosity, and Capillary Action 

Surface Tension 

• Liquids tend to minimize their surface tension 
• Surface tension is the energy required to increase the surface area by a unit amount 
• Surface tension decreases as intermolecular forces decrease 
• Surface tension allows water droplets to be sphere shaped 

Viscosity 

• Viscosity is the resistance of a liquid to flow 
• Viscosity is measure in posie (P)
• Viscosity is 1g /cm x s
• Viscosity is greater in substances with stronger intermolecular forces 
• Viscosity depends on molecular shape
• Viscosity increases with increasing molar mass, magnitude of dispersion forces, and increasing length 
• Viscosity depends on temperature 

Capillary Action

• Capillary action is the ability of a liquid to flow against gravity and up a tube 
• Capillary action results from the combination of cohesive and adhesive forces
• Cohesive forces are the attractions between molecules in a liquid 
• Adhesive forces are the attractions between molecules and the surface of the tube 
• Adhesive forces cause the liquid to spread across the surface of the tube
• Cohesive forces cause the liquid to stick together 
• If adhesive forces are greater than cohesive forces, the liquid will climb up the tube 
• The thinner the tube, the higher the liquid will rise 

11.5 Vaporization and Vapor Pressure 

The Process of Vaporization 

• The higher the temperature, the greater the average energy of the collection of molecules 
• Some molecules have more thermal energy than others
• Some molecules have less thermal energy than others 
• The transition from the liquid state to the gas state is vaporization 
• The opposite of vaporization, the transition from gas to liquid, is condensation 
• Vaporization occurs quicker at higher temperatures 
• Vaporization occurs quicker on larger surfaces 
• The rate of vaporization increases as intermolecular force strength decreases 
• Liquids that vaporize easily are volatile 
• Liquids that do not vaporize easily are nonvolatile 

The Energetics of Vaporization

• Vaporization is an endothermic process because it energy is absorbed to change from a liquid to a gas 
• Energy is needed to break the intermolecular forces 
• Condensation is exothermic
• Heat is released when a gas condenses to a liquid 
• The heat (or enthalpy) of vaporization is the amount of heat required to vaporize one mole of a liquid to gas 
• The heat of vaporization is always positive because the process is endothermic 
• When a substance condenses, the value will be negative 

Vapor Pressure and Dynamic Equilibrium 

• Once water molecules enter the gas state, some will start condensing back into a liquid 
• As the number of gaseous water molecules increases, so does the rate of condensation 
• Once the rate of vaporization and condensation are equal, they have reached dynamic equilibrium 
• Vapor pressure is the pressure of a fas in dynamic equilibrium with its liquid 
• Vapor pressure depends of intermolecular forces and temperature 
• Weak intermolecular forces have higher vapor pressure 
• Strong intermolecular forces have low vapor pressure
• When a balanced system is disturbed, it will work its way back to equilibrium and try to minimize the disturbance
• When temperature increases, vapor pressure rises 
• Boiling point is when the liquids vapor pressure is equal to the external pressure 
• At the boiling point, thermal energy is enough that molecules break free of their neighbors and enter the gas state 
• Normal boiling point is when vapor pressure equals 1 atm
• Lower pressures result in lower boiling points 
• Once the boiling point is reached, heating the temperature will not cause the process to occur faster 
• As long as the liquid is present, temperature cannot rise above its boiling point 
• Vapor pressure and temperature have a exponential relationship 

The Critical Point: The Transition to an Unusual State of Matter 

• As temperature and pressure increase, the density of a gas increases 
• As temperature increases, the density of a liquid decreases 
• A supercritical fluid is a substance that forms above the critical point 
• The temperature where supercritical fluids occur is the critical temperature 
• The pressure where supercritical fluids occur is the critical pressure 
• Supercritical fluids have properties of both liquids and gases 

11.6 Sublimation and Fusion 

Sublimation 

• Sublimation is the transition form solid to gas 
• The opposite of sublimation, transitioning from gas to solid, is deposition 
• The vapor pressure of the solid is the pressure of a gas in dynamic equilibrium with its solid 
• Sublimation will occur at a greater rate 
• Colder temperatures will lower the vapor pressure 

Fusion 

• An increase in thermal energy causes molecules to vibrate faster 
• At the melting point, molecules have enough energy to turn from solids into liquids 
• Melting is the transition from solid to liquid 
• Melting is also called fusion 
• The opposite of melting, the transition from liquid to solid, is freezing 
• Once the melting point is reached, going to a higher temperature will not speed up the process

Energetics of Melting and Freezing

• Melting is endothermic 
• The heat of fusion is the amount of heat required to melt one mole of a solid 
• The heat of fusion is a positive value 
• Freezing is exothermic 
• Different substances will have different heat of fusions 
• Heat of fusion is generally less than the heat of vaporization because liquids and solids are similar 

11.7 Heating Curve for Water

• Heating curves show at what temperatures each phase occurs
  
• Heating curves have 5 segments: Solid warming, melting, liquid warming, vaporization, and gas warming 
• Melting and vaporization are constant because these are the transition stages 
• Melting and vaporization are in equilibrium and the temperature will remain constant 
• In the warming stages, temperature increases linearly 
• The specific heat capacity is different for different substances 

11.8 Phase Diagrams 

• A phase diagram shows a map of a substance based on pressure and temperature 

The Major Features of a Phase Diagram 

• The y-axis displays pressure in the unit torr
• The x-axis displays the temperature in Celsius 
• The three main regions of the phase diagram represent solid, liquid, and gas
• The phrase diagram represents conditions where that phase is occurring
• Low temperatures and higher pressures tend to be solids 
• High temperatures and low pressures tend to be gases
• Middle temperatures and middle pressures tend to be liquids 
• Each of the lines, or curves, represent the temperatures and pressures where equilibrium between states is occurring 
• The line separating liquid and gas is the vaporization curve 
• The sublimation curve separates solids and gases 
• The fusion curve separates solids and liquids 
• The triple point is where all three phases occur in equilibrium 
• The critical point is the temperature and pressure above which a supercritical fluid exists 

Navigation within a Phase Diagram 

• As temperature rises, move right along the x-axis 
• As pressure rises, move up along the y-axis 

The Phase Diagrams of Other Substances 

• The fusion curve for carbon dioxide and iodine have positive slopes 
• The fusion curve for water has a negative slope 

11.9 Water: An Extraordinary Substance 

• Life is impossible without water 
• Water has a low molar mass, but is a liquid at room temperature 
• Water has a high boiling point 
• Water molecules are bent in shape 
• Water molecules are highly polar 
• Water's high polarity allows it to dissolve other polar and ionic compounds 
• Water is the main solvent in living organisms 
• Water has a high specific heat capacity 
• Water expands when it freezes, most substances contract 

11.10 Crystalline Solids: Determining Their Structure by X-Ray Crystallography 

• X-ray distraction allows scientists to determine the arrangement of atoms and measure the distance between them 
• Waves interact through interference 
• Constructive interference occurs when two waves interact with crests and troughs aligning with each other 
• Destructive interference occurs when two waves interact with crests and trough aligning with the other 
• Interference patterns consists of alternating light and dark lines 
• Atoms of crystal structures act similarly 
• The pattern of a diffraction reveals the spacing between atoms 

11.11 Crystalline Solids: Unit Cells and Basic Structures 

• The arrangement of atoms in a crystalline solid is called the crystalline lattice 
• Crystalline lattices are represented with small collections of atoms, ions, or molecules
• This representation is called a unit cell 
• When unit cells are repeated, an entire lattice is produced 
• The lattice point is a space occupied by an atom, ion, or molecule 
• Unit cells are classified by their symmetry 
• Cubic unit cells have equal edge lengths and 90 degree angles
• A unit cell may seem like it contains 8 atoms, but in reality there is only 1
• The coordination number is the number of atoms which each atom is in direct contact with 
• Packing efficiency is the percentage of the volume of the unit cell occupied by the spheres 
• The higher the coordination number, the higher the packing efficiency 
• Body centered unit cells contain two atoms per cell 
• Face centered unit cells contain four atoms per cell

Closest Packed Structures 

• Crystal structures have lots of empty space 
• Packing efficiency can be increased by not placing the next layer directly on top of the first 
• In hexagonal closest packing, the third layer is lined up directly above the first layer 
• In cubic closest packing, the third layer is lined up off the first layer 

11.12 Crystalline Solids: The Fundamental Types 

• There are three categories of crystalline solids: Molecular, ionic, and atomic 
• Atomics are classified into non bonded, metallic, and network covalent 

Molecular Solids 

• Molecular solids are solids whose units are molecules 
• Ice is an example of a molecular solid 
• They are held together by intermolecular forces 
• They tend to have low melting points 

Ionic Solids 

• Ionic solids are made up of ions 
• NaCl and CaF2 are examples of ionic solids 
• They are held together by coulombic interactions between cations and anions 
• Ionic solids have higher melting points than molecular solids 

Atomic Solids

• Atomic solids are made up of atoms 
• Nonbonding atoms are held together by weak dispersion forces 
• Nonbonding atoms form closest packing structure to maximize interactions 
• Nonbonding atoms have very low melting points 
• Nonbonding atomic solids are noble gases in their solid form 
• Metallic atomic solids are held together by metallic bonds 
• Metals form closest packed crystal structures 
• Some metals have high melting points while other metals have very low melting points 
• Network covalent atomic solids are held together by covalent bonds 
• Network covalent solids do not form closest packed structures  

11.13 Crystalline Solids: Band Theory 

• The band theory combines atomic orbitals of the atom within a solid crystal to from orbitals that are not localized on individual atoms 
• Electrons become mobile when they make a transition from the highest occupied molecular orbital into higher energy empty molecular orbitals 
• Occupied molecular orbitals are the valence band
• Unoccupied orbitals are conduction bands 
• When a metal is heated, electrons are excited and go to a higher molecular orbital 
• An energy gap, the band gap, exists between valence and conduction bands in semiconductors and insulators 

Doping: Controlling the Conductivity of Semiconductors

• Dope semiconductors contain holes in the valence band