Lecture 10

The Liquid and Solid States

Shaun Williams, PhD

Examples of Solids, Liquids, and Gases

Photographs of a solid, a liquid, and a gas.

Changes of State

Transitions between these states
Also called phase changes
6 main phase changes:
- Evaporation (also called vaporization)
- Condensation
- Freezing
- Melting (also called fusion)
- Sublimation
- Deposition

Liquid-Gas Phase Changes

Evaporation (vaporization)
- At the surface of a liquid, some molecules may have sufficient kinetic energy to escape into the gas state.
- Heat is required to maintain the temperature needed for evaporation.
  - Evaporation is an endothermic process.

An imagine depicting water molecules transitioning from the condensed liquid phase to the gas phase.

Liquid-Gas Phase Changes (cont.)

Images showing that in a closed container holding liquid water, water vapor builds up over time.

Condensation
- Transition from a gas to a liquid
- Occurs when gas particles cannot escape the container and thus, come into contact with a liquid
- An exothermic process (energy is released)

Equilibrium

A state in which opposing processes occur at equal rates
An equilibrium is designated by a double arrow, such as: \[ \text{liquid} \rightleftharpoons \text{gas} \]
In the above equilibrium, the rates of evaporation and condensation are equal.
The gas produced by evaporation exerts a pressure on the liquid below it.
- At equilibrium, this pressure is called vapor pressure.
  - Vapor pressure increases as temperature increases.

Boiling Point

The temperature at which boiling occurs
- Boiling occurs when the vapor pressure equals the external pressure of the atmosphere
Normal boiling point
- Occurs when the atmospheric pressure is exactly 1 atm

A plot showing three different compounds vapor pressures as a function of temperture.

Liquid-Solid Phase Changes

Freezing
- The average kinetic energy of a liquid molecules fall when it is cooled
- If the $ KE_{av} $ falls low enough, then the molecules become fixed in position into the solid
- Therefore, freezing is the conversion of a liquid into a solid

Liquid-Solid Phase Changes (cont.)

Freezing point
- The temperature at which the freezing of a liquid into a solid state occurs
  - As with boiling temperature, the two states are in equilibrium with one another. \[ \text{solid} \rightleftharpoons \text{liquid} \]
- Normal freezing point
  - The temperature when the freezing equilibrium is achieved under a pressure of 1 atm.

Freezing

A photograph of an ice cube melting and a image of molecular ice crystal and liquid water.

Melting

Melting
- Phase change from solid to liquid
- Reverse of freezing
- Opposite process of freezing
- Also called fusion
Melting point
- Same temperature as freezing point

Images showing solid water transitioning to them more disordered liquid water and crystalling sodium chloride transitioning to the more dissordered liquid salt.

Solid-Gas Phae Changes

Sublimation
- Evaporation of a solid
- Occurs when a solid has a high vapor pressure
- Solid can change directly from a solid to a gaseous state without going through the liquid state
Deposition
- Can go directly from gas to solid without passing through the liquid state
- Reverse of sublimation

A photograph of gaseous iodine depositing on the bottom of a cold evaporating dish.

Cooling Curve

Shows the phase changes of a substance in a graph plotting temperature versus heat removed at constant pressure

A plot of the temperature on a sample over time as heat is removed at a constant temperature. The temperature drops except when the phase changes are occuring (temperature is constant during phase changes).

Heating Curve

Shows the phase changes of a substance in a graph plotting temperature versus heat added at constant pressure

A plot of the temperature on a sample over time as heat is added at a constant temperature. The temperature increases except when the phase changes are occuring (temperature is constant during phase changes).

Energy Changes

Each change of state takes place at constant temperature
Any phase change is accompanied by an energy change
- The change in energy for each process is called the heat of that process ($q$) \[ q=m \times C \times \Delta T \]
- Molar heat of fusion
  - Energy required to melt 1 mole of a substance
- Molar heat of vaporization
  - Energy required to evaporate 1 mole of a substance

Intermolecular Forces

An attractive force that operates between molecules
There are many kinds of intermolecular forces:
- London dispersion force
- Dipole-dipole force
- Hydrogen-bonding force

A photo showing the yellowing chlorine gas, reddish-brown bromine liquid, and purple iodine crystals.

The Forces

Type of Force	Type of Interaction	Occurrence
London dispersion force	A temporary dipole in one molecule induces the formation of a temporary dipole in a nearby molecule and is attracted to it.	All atoms and molecules
Dipole-Dipole Force	Polar molecules (permanent dipoles) attract one another	Polar molecules
Hydrogen-Bonding Force	Two dipoles, one containing hydrogen to an electronegative element and the other containing an electronegative element, attract one another.	Polar molecules containing unpaired molecules and a hydrogen bonded to nitrogen, oxygen, or fluorine

London Dispersion Forces

Instantaneous dipole
- A temporary dipole formed when the electrons in an atom or nonpolar molecule happen to be more on one side in an instant in time, causing it to be more negative than normal and the opposite side positive
Induced dipole
- Positive end of the dipole exerts an attractive force on nearby electrons, causing an adjacent atom to develop into another temporary dipole

London Dispersion Forces (cont.)

London dispersion force
- The attraction between temporary dipoles
- Occurs between atoms and molecules
- Only intermolecular force in nonpolar substances
- Tend to be stronger the larger the atom or molecule
- Relatively weak forces

A series of images showing electron clouds fluctuating in time and forming a short lived bulge and therefore a small dipole moment.

Dipole-Dipole Forces

Attraction between polar molecules
Occurs when the partially positive end of one molecule attracts the partially negative end of another molecule
Generally stronger than London dispersion forces

An image showing the dipole moment on one SO2 molecule attracting the dipole moment on a neighboring SO2 molecule.

Boiling Points

A plot showing the boiling points of groups 17 diatomics, group 14 with hydrogen, and group 18 atoms. In each case the trend has the lighter species having a lower boiling point and the general trend is linear.

Hydrogen Bonding

Special type of dipole-dipole force
Only occurs in molecules that contain hydrogen bonded to a small, highly electronegative element
Stronger than a regular dipole-dipole force
Important force in living systems by stabilizing molecular shapes

A plot showing the boiling points of molecules involving groups 14, 15, 16, and 17 with hydrogen atoms. The molecules H2O, HF, and NH3 boiling at considerably higher temperatures than their family's linear trend predicts.

Hydrogen Bonding Examples

Images of hydrogen bonding in water, ammonia in water, ammonia by itself, and hydrogen fluoride.

Trends in Intermolecular Forces

Remember:
- London forces exist between all atoms and molecules.
- Dipole-dipole moments exist in only polar compounds.
- Hydrogen bonds only exist in polar compounds that contain hydrogen.
In terms of strength (magnitude), intermolecular forces compare as shown below:

All molecules have London Dispersion Forces, polar molecules have dipole-dipole forces, and polar molecules with HF, HO, or HN bonds have hydrogen bonding.

\[ \text{London Disp.} \lt \text{Dipole-Dipole} \lt \text{Hydrogen Bonds} \]

Properties of Liquids

Are related to the distance between particles and to intermolecular forces
Particles in a liquid are much closer together than the particles in a gas
Liquid particles are not fixed, as they are in a solid
Three common properties of liquids:
- Density
- Viscosity
- Surface tension

Density

Remember: \[ \text{density}=\frac{\text{mass}}{\text{volume}} \]
Densities of the states of matter are related to the distance between particles
Most substances are denser as solids than as liquids because their molecules or atoms are closer together
- Water is an exception in that ice is less dense than liquid water

Viscosity

Liquids and gases are fluids.
- A fluid is any substance that can flow
Viscosity is the resistance of a substance to flow.
- Generally, the viscosity of liquids is low
Viscosities generally vary by increasing with the magnitude of their intermolecular forces and with molecular size.

Surface Tension

The amount of work required to increase the surface area of a liquid by a unit amount
Causes a liquid surface to behave like a stretched membrane
The greater the intermolecular forces in a liquid, the greater the surface tension
Surface tension decreases as temperature increases

An image showing the hydrogen bonding of a water molecule in the center of a liquid versus a water molecule at the surface of the liquid (at the surface it has half as many hydrogen bonds.

Meniscus

Either a concave or convex curved surface of a liquid produced by intermolecular forces
A concave surface occurs when the intermolecular forces between the liquid and the glass are greater than the intermolecular forces among the liquid molecules
A convex surface occurs when the intermolecular forces between the liquid and the glass are less than the intermolecular forces among the liquid molecules

A picture of the concave surface of water in a test verus the convex surface of mercury in a test tube.

Properties of Solids

Amorphous and Crystalline Solids

Amorphous solid
- Occurs when the temperature of a liquid drops rapidly, resulting in the particles solidifying in a partially disordered state
- Particles are somewhat randomly arranged
- Lacks regular form
Crystalline solid
- Occurs when a liquid solidifies slowly, allowing the array of particles to become well ordered
- Most solids are of this type
- Results in the symmetrical arrangement of planar faces

Amorphous and Crystalline Solids Examples

$A picture of quartz crystal showing the appearance of a crystalling solid and another photograph of broken window glass showing the random nature of the fractures indicating the amorphous nature of window glass.$

Crystals and Crystal Lattices

Crystal
- An orderly, repeating, three-dimensional assembly of fundamental particles (atoms, molecules, or ions)
Crystal lattice
- Pattern forms by repeating arrays of crystal
- Also called crystal structure
- Atoms, molecules, or ions are arranged in close-packed structures
  - Orderly arrangement that is an efficient way to use space

Closest-Packed Arrangements

Photographs of the surface of a pineapple, the cells in a bee hive, the celling in a leaf, and the dimples on a golf ball. Each is an example of packing things into a limited space.

Types of Crystalling Solids

Solids show a wide range of properties such as:
- Melting point
- Hardness
- Malleability
The types of forces holding the fundamental particles together help explain the properties
4 Types of crystalline solids
- Metallic solids
- Ionic solids
- Molecular solids
- Network solids

Types of Crystals

Type of Solid	Fundamental Particles	Attractive Forces	Properties
Metallic	Atoms	Attractions between nuclei and delocalized electrons	Low melting point & soft; or high melting point & hard; good heat & electrical conductors; malleable & ductile
Ionic	Cations and Anions	Ionic bonds	High melting point; hard, brittle; nonconductors when solid; electrical conductors when melted
Molecular	Polar Molecules	Dipole-dipole forces	Low to moderate melting point; variable hardness; may be brittle; nonconductors
Molecular	Nonpolar Molecules	London dispersion forces	Low melting point; soft; poor heat conductors; electrical insulators
Netword	Atoms	Covalent bonds	Very high melting point; very hard; somewhat brittle; non- or semiconductors

Metallic Solids

Valence electrons move freely through all parts of a metal
Attractions between atoms of a metal are delocalized, and therefore it is easy to move atoms by applying force
Ductile and malleable

A photograph of copper, zinc, and brass metals.

Alloys

Forms when a metal is mixed with one or more additional metallic or nonmetallic elements
Have properties different than those of their parent elements

A image of the mixed atoms in a titanium alloyed with nickel.

Ionic Solids

Contain cations and anions arranged in crystalline solids
Electrostatic forces (ionic bonds) hold together ionic crystals
High melting points
Hard and brittle
Solids are not electrical conductors, but melted or dissolved, they become good conductors

A image of crystals of sodium chloride, zinc sulfide, calcium fluoride, and cesium chloride.

Superconductors

Offer no resistance to the conduction of electrical current
Repel magnetic fields
Although the nature of superconductors is not entirely understood, the crystal structure of the solid material does play a role
Often formed from silicon and/or other metalloids

An image of a generic superconductor made up of yttrium, barium, copper, and oxygen.

Molecular Solids

Intermolecular forces between molecules hold a molecular solid together
- Nonpolar solids are held together by London dispersion forces
  - Form soft crystals with low melting points
  - Electrical insulators
  - Poor conductors of heat
- Polar solids are held together by London dispersion forces and dipole-dipole forces
  - Typically harder than nonpolar solids
  - Low to moderate melting points
  - Electrical insulators (no ions)

Example of a Molecular Solid

An image of carbon dioxide molecules arranged in a crystal that is dry ice.

A photograph of a diamond, a network solid where all the carbon atoms are covalently bound to four neighboring carbon atoms.

Example of a Network Solids

A photograph of sand paper, the sand on which is silicon dioxide (a network solid). A photograph of graphite in pencil lead, which is a network solid of carbon where each carbon atom is covalently bonded to three neighbors on a plane.

Lecture 10 The Liquid and Solid States