Chapter 11: States of matter;

liquids and solids

Chemistry 1062: Principles of Chemistry II

Andy Aspaas, Instructor

Gases, liquids, and solids

• Gas: compressible fluid

– Mostly empty space– Particles in constant random motion– Vapor: gas form of substance normally solid or liquid

• Liquid: incompressible fluid

– Tightly packed particles– Still in constant random motion

• Solid: incompressible and rigid

– Tightly packed particles– Particles only vibrate about their fixed sites

Phase transitions

• Melting (fusion): transition from a solid to a liquid

• Freezing: transition from a liquid to a solid

• Vaporization: change of a solid or liquid to a vapor

– Sublimation: change of a solid directly to a vapor

• Condensation: change of a gas to a liquid or a solid

– Deposition: change of a gas directly to a solid

Vapor pressure

• Vapor pressure of a liquid: partial pressure of vapor over the liquid, measured at equilibrium– Vaporization and condensation are happening

continually– When the rate of condensation equals the rate of

vaporization, equilibrium has been reached• Vapor pressure depends on temperature

– More kinetic energy in the liquid makes vaporization occur more quickly

• Volatile liquids and solids have high vapor pressure

Boiling point and melting point

• Boiling point: temperature at which the vapor pressure of a liquid equals the atmospheric pressure– Bubbles form in the liquid when it’s heated

enough that its vapor pressure equals external pressure

• Once boiling begins, the temperature of the liquid stops rising

• Freezing point: temperature at which a pure liquid changes to a solid– Identical to melting point: temperature at which a

solid changes to a liquid

Heat of phase transition

• Heat of fusion: amount heat required to melt a solid

H2O(s) H2O(l); ∆Hfus = 6.01 kJ/mol

• Heat of vaporization: amount of heat required to vaporize a liquid

H2O(l) H2O(g); ∆Hvap = 40.7 kJ/mol

• Heating curve:no temperaturechange during a phase transition

Phase transitions

• Heat must be added for melting or vaporization

– Endothermic processes– Positive value of ∆H

• Heat must be removed for condensation or freezing

– Exothermic processes– Negative value of ∆H

Clausius-Clapeyron equation

• Vapor pressure of a liquid depends on temperature– Also depends on ∆Hvap for that liquid

• Clausius-Clapeyron equation:

– The two-point form removes the constant B from the equation

€

lnP =−ΔHvap

RT+ B

€

lnP2

P1

=ΔHvap

R

1

T1

−1

T2

⎛

⎝ ⎜

⎞

⎠ ⎟

Phase diagrams

• Phase diagram for water:

– Curves show exptl points at which phases are in equilibrium

• Solid-liquid line (melting-point curve)

– Leans left if liquid is more dense than solid (as in water)

– Otherwise leans right

Phase diagrams

• Curve AC gives the vapor pressure at any temperature

• Triple point, A: point at which all 3 phases exist in equilibrium

• Sublimation occurs when heating at pressures below triple point

• Critical point, C: point above which gas and liquid are no longer distinct

– Supercritical fluid: at temperature and pressure above critical point

Liquid state: surface tension

• Surface tension: energy required to increase the surface area of a liquid by a unit amount– Liquids will minimize their amount of surface area

(since intermolecular forces pull molecules at the surface inward)

• Dissolved substances reduce surface tension by interrupting intermolecular forces

• Capillary rise: when a liquid is attracted to a solid capillary (like water to glass) it will rise in the capillary to minimize surface tension at the inverted meniscus

Viscosity

• Viscosity: resistance to flow in a liquid or gas

– Strong intermolecular forces increase viscosity and cause a liquid to flow more slowly (like syrup)

• Motor oil viscosity in SAE units

– 10W/30 has SAE 10 in winter, 30 in summer (because viscosity of engine oil increases as temperature decreases)

Intermolecular forces

• Dipole-dipole force: attractive interaction between polar molecules

– Polar molecules have a dipole moment+ and – sides of the molecule resulting from

electronegativity differences– Negative side of one molecule will be attracted to the

positive side of another molecule• London dispersion forces: attractive forces resulting from

instantaneous induced dipoles caused by random electron groupings in an atom

– High molecular weight = higher London forces– Compact molecule = lower London forces

Predicting boiling points, vapor pressure, and viscosity

• Polar molecules have stronger intermolecular forces than nonpolar molecules of similar MW (dipole-dipole)

• Larger molecules have stronger intermolecular forces than smaller molecules of similar polarity (London)

• Strong intermolecular forces

– Low vapor pressure (less chance for molecule to break free)

– High boiling point (high temperature required to raise vapor pressure to atmospheric pressure)

– High viscosity (molecules cling strongly together)

Hydrogen bonds

• Hydrogen bonds: much stronger than dipole-dipole or London forces

• Interaction between H and Y in –X–H·····Y–– Where X and Y are F, O, or N (electronegative

atoms)– H has a strong + when covalently bonded to an

electronegative atom• Explains why CH3OH boils at 65 °C and CH3F boils

at –78 °C (both have similar MW and polarities)

Classification of solids

• Molecular solid: solid that consists of atoms or molecules held together by intermolecular forces– H2O(s), CO2(s)

• Metallic solid: any solid metal contains a regular arrangement of positive metal nuclei surrounded by a sea of delocalized electrons

• Ionic solid: consists of cations and anions held together by electrostatic attractions of opposite charges (ionic bonds)

• Covalent network solid: consists of atoms held together in large networks or chains by covalent bonds

Melting point of solids

• Molecular solid: low melting point (below 300 °C)– Only weak intermolecular forces to overcome

• Ionic and covalent network solids: high melting point (800 °C or above)– Ionic and covalent bonds require much more

energy to be broken– Larger charges make even higher mp’s in ionic

solids• Metals: variable (Hg: –39 °C; W: 3410 °C)

Hardness and conductivity of solids

• Molecular solids: generally soft and brittle• Ionic solids: generally hard and brittle• 3-dimensional covalent network solids are

tremendously hard• Metals are malleable since the atoms can move past

each other• Metals (and molten ionic solids) conduct electricity

because electrons can move freely

Crystalline solids

• Crystalline solids have an ordered 3-dimensional structure of particles

• Amorphous solids have their particles locked into random positions

• Crystal lattice: 3-dimensional geometric arrangement of particles in a crystal

– Unit cell: smallest boxlike unit from which a crystal lattice can be constructed

– Unit cells are stacked in 3 dimensions

Cubic unit cells

• Simple cubic unit cell: atoms only at the corners of a cubic unit cell

– Only 1/8 of a corner atom is contained in a unit cell

• Body-centered cubic: atom at the center of a unit cell as well as the corners

• Face-centered cubic: atoms in the center of each of the faces as well as the corners

– Only 1/2 of a face atom is contained in a unit cell

Cubic unit cells

Molecular solids; closest packing

• In simple molecular solids (such as individual noble gas atoms), nondirectional London forces govern the packing of the atoms together to form a solid

• Spheres pack together most closely in a hexagonal honey-comb type arrangement on one layer

– A second layer can be nested in half of the crevices left by the first layer

– A third layer can be nested in the crevices in the second layer, but there are two possible positions for this layer

Close packing arrangements

• Hexagonal close packing (hcp): when 3rd layer is identical to 1st (in x holes left by 2nd layer)

• Cubic close packing (ccp): when 3rd layer is different (in y holes left by 2nd layer)

Cubic close-packing

• Cubic close-packing has a face-centered cubic lattice structure (tilted on its side)

Coordination number

• Coordination number: number of nearest-neighbor adjacent atoms any one atom can have

• Close packing structures: CN = 12

– hcp, ccp/face-centered cubic

• Body-centered cubic: CN = 8

• Square cubic: CN = 4

Ionic solids

• 3 common structures: CsCl structure, NaCl structure, and zinc blende (cubic ZnS) structure

• CsCl structure

– 1 of eachper unit cell

– When ions are similarin size

– Expandedsimple cubic

NaCl structure

• Expanded face-centered cubic structure

• When cation and anion are moderately different in size

Zinc blende structure

• When cation and anion are significantly different in size

• Anion is expanded body-centered cubic

• Cation is in 4 opposite tetrahedral quadrants completely inside the unit cell

Unit cell calculations

• Length of side of a unit cell

– Volume of a unit cell• Mass of a unit cell

–Mass of individual atoms»Molar mass

• Forward or backwards!