Untitled

Molecular Comparison of Liquids and Solids

• Physical properties of liquids and solids depend on intermolecular forces, which are forces between molecules.
• Substances' physical properties are understood using kinetic-molecular theory.
• Gases are compressible and take the shape/volume of their container; gas molecules are far apart with minimal interaction.
• Liquids are nearly incompressible and assume the shape, but not the volume, of their container.
• Liquid molecules are closer than gas molecules but can still slide past each other.
• Solids are incompressible with definite shape/volume; solid molecules are closely packed and cannot easily slide.
• Solids and liquids are condensed phases.
• Converting a gas to liquid/solid requires molecules to get closer, achieved by cooling or compressing the gas.
• Converting a solid to liquid/gas requires molecules to move apart, achieved by heating or reducing gas pressure.

Intermolecular Forces

• The covalent bond holding a molecule together is an intramolecular force; the attraction between molecules is an intermolecular force.
• Intermolecular forces are weaker than intramolecular forces (e.g., 16 kJ/mol for intermolecular vs. 431 kJ/mol for HCl).
• When a substance melts or boils, intermolecular forces break; when a substance condenses, intermolecular forces form.
• Boiling/melting points reflect intermolecular force strength.
• High boiling points indicate strong attractive forces.
• Van der Waals forces are intermolecular forces between neutral molecules.
• An ion-dipole force is an interaction between an ion (e.g., Na+) and a polar molecule/dipole (e.g., water), it is the strongest intermolecular force.
• Ion-dipole forces are important for solutions of ionic substances in polar liquids, like NaCl(aq).

Dipole-Dipole Forces

• Dipole-dipole forces exist between neutral polar molecules, where polar molecules attract.
• The partially positive end of one molecule attracts the partially negative end of another.
• Polar molecules need to be close for strong dipole-dipole interactions.
• Dipole-dipole forces are weaker than ion-dipole forces.
• When two molecules have similar mass/size, dipole-dipole forces increase with increasing polarity.

London Dispersion Forces

• London dispersion forces are the weakest intermolecular forces.
• Adjacent nonpolar molecules can affect each other.
• Helium atoms have a symmetrical electron distribution; at any instant, both electrons might be on one side, forming an instantaneous dipole.
• This instantaneous dipole induces another in an adjacent molecule/atom.
• The attraction between temporary dipoles is the London or dispersion force, existing among all molecules.
• Strength Factors: Molecules must be very close for attraction; polarizability is the ease with which an electron cloud distorts.
• Larger molecules (more electrons) are more polarizable, and London dispersion forces increase with molecular weight. Shape is another factor.
• Greater surface area yields greater dispersion forces.
• Dispersion forces are smaller between spherical molecules than cylindrical ones
• n-pentane vs. neopentane is an example.

Hydrogen Bonding

• Boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high because their intermolecular forces are strong.
• Hydrogen bonding is a special dipole-dipole interaction.
• H-bonding requires an H bonded to an electronegative element (F, O, and N) and an unshared electron pair on a nearby electronegative ion/atom (F, O, or N).
• Electrons in the H-X bond (X = electronegative element) lie near X, so H presents an almost bare proton.
• H-bonds are strong and have biological significance by stabilizing protein/DNA structure.
• Solid molecules are usually more closely packed than liquids, thus solids are usually denser than liquids.
• Hydrogen bonding interactions are random in liquid water.
• Ice has an ordered, open structure for optimized H-bonding.
• In water the H-O bond is 1.0 angstrom; the O…H hydrogen bond is 1.8 angstrom.
• Water molecules arrange in an open hexagon, with each H+ pointing toward O's lone pair.
• Thus, ice is less dense and floats, forming an insulating layer on water bodies, allowing aquatic life to survive in winter, this is because water expands on freezing.

Comparing Intermolecular Forces

• Dispersion forces are in all substances, with strength based molecular shapes and weights.
• Dipole-dipole forces add to dispersion forces, exclusive to polar molecules.
• H-bonding, the strongest force involving neutral species, is a special dipole-dipole interaction.
• H-bonding is vital for H compounds of N, O, and F.
• If ions are involved, ion-dipole interactions (if a dipole is present) and ionic bonding are possible.
• Ion-dipole interactions are stronger than H-bonds.
• Ionic/covalent bonds are stronger than these interactions.

Some Properties of Liquids

• Viscosity resists a liquid's flow, where flow stems from molecules sliding over each other.
• Viscosity increases with stronger intermolecular forces and the tendency of molecules to entangle, but, viscosity decreases when temperature increases.
• Bulk molecules are equally attracted to their the surrounding molecules.
• Surface molecules attract inward toward the bulk molecules.
• Because of unequal attraction, surface molecules pack closely, causing the liquid to behave as the molecules have "skin."
• Surface tension is the energy to increase a liquid's surface area.
• Stronger intermolecular forces raise surface tension.
• Water possessing high surface tension due to H-bonding.
• Hg(/) has very strong metallic bonds, resulting in an even higher surface tension.
• Cohesive forces are intermolecular forces that bind molecules to one another.
• Adhesive forces bind molecules to a surface.
• Meniscus: The shape of a liquid in a tube is called the meniscus, this is because the container may have either greater adhesion or cohesion than molecules in the liquid do themselves.
• If force of adhesion is greater than force of cohesion, then the liquid is more attracted to the container, resulting in a "U" shape.
• Example: Water is a "U" shape because it sticks easily to glass.
• If force of cohesion is greater than force of adhesion, then the liquid pulls away and forms a curved surface.
• Example: Mercury forms a curved meniscus in glass.
• Capillary action: Liquids rise in narrow tubes until adhesive/cohesive forces balance gravity.

Phase Changes

• Phase changes are changes of state, where matter in one state converts to another.
• Sublimation: solid to gas
• Melting/fusion: solid to liquid
• Vaporization: liquid to gas
• Deposition: gas to solid
• Condensation: gas to liquid
• Freezing: liquid to solid
• Energy changes occur during phase changes.
• Sublimation and melting/fusion (∆Hfus > 0) are endothermic, involving heat of fusion.
• Vaporization: ∆Hvap > 0) is endothermic, involving heat of vaporization.
• Deposition, condensation, and freezing are exothermic.
• The heat of fusion (enthalpy of fusion) is less than the heat of vaporization.
• All phase changes are possible under the right conditions.
• Ex: Water sublimes when snow disappears without forming puddles.
• Heat solid --> heat liquid --> boil --> heat gas is endothermic.
• Cool gas --> condense --> cool liquid --> freeze --> cool solid is exothermic.
• Heating curve: A plot of temperature changes versus heat added is a heating curve.
• During a phase change, temperature is constant, so, added energy disrupts intermolecular interactions.
• The points on a such curves are used to calculate ∆Hfus and ∆Hvap.
• Supercooling: When a liquid cools to the freezing point, but still remains a liquid.
• You can plot the temperature rise against the heat absorbed if you take 1 kg of water from -10°C up to 150°C.

Critical Temperature/Pressure

• Gases can be liquefied by increasing pressure at a suitable temperature.
• Critical temperature: the highest temperature at which a substance can exist as a liquid.
• Critical pressure: the necessary pressure for liquefaction at a critical temperature.
• The greater the intermolecular forces, the easier it is to liquefy a substance and the higher the critical temperature.

Vapor Pressure

• Some molecules on a liquid's surface have enough energy to escape the bulk liquid's attraction, moving to gas phase.
• As gas phase molecules increase, some strike and return to the liquid.
• After some time the pressure of the gas will be constant.
• Dynamic Equilibrium: A condition in which two opposing processes occur simultaneously at equal rates.
• Equilibrium vapor pressure: The pressure exerted when liquid and vapor are in dynamic equilibrium.
• Volatility increases when equilibrium is never established, the vapor continues to form, eventually leading to the liquid drying out in the process.
• Liquids that evaporate easily are said to be volatile.
• The higher the temperature, the higher the average heat energy and the faster the liquid evaporates.
• Liquids boil when the external pressure at the surface of the liquid equals the vapor pressure.
• Normal boiling point is the boiling point at 760 mm Hg or 1 atm.
• As external pressure increases, the boiling point temperature increases as well.
• Two Ways to boil liquid: Increase temperature or decrease pressure.
• Water has a high boiling point at high pressure, so when food is cooked is is cooked at a higher temperature.

Phase Diagrams

• Phase diagram: a pressure vs. temperature plot summarizing equilibria between phases; they indicate which phase exists at a given temperature and pressure.
• Vapor-pressure curve is a feature of a phase diagram, it states that as temperature increases, vapor pressure increases as well.
• Critical point: Provides the critical temperature and pressure for gas.
• Normal melting point: melting point at 1 atm.
• Triple point: Temperature and pressure at which three phases are in equilibrium.
• In a phase diagram, any temperature/pressure combination not on a curve represents a single phase.

Phase Diagrams of H2O and CO2

• In general, increasing pressure favors the more compact phase which is usually a solid.
• Water is one of the few substances whose solid is less dense than liquid.
• The melting point curve for water slopes to the left.
• For water the triple point occurs at 0.0098°C and 4.58 mm Hg.
• Normal melting (freezing) point is 0°C.
• Normal boiling point is 100°C.
• The critical point is 374°C and 218 atm.
• Freeze-drying: frozen food in a low-pressure chamber (< 4.58 torr), and ice sublimes.
• For carbon dioxide the triple point occurs at -56.4°C and 5.11 atm.
• Normal sublimation point is –78.5°C, because at 1 atm, CO2 sublimes, but does not melt.
• Critical point occurs at 31.1°C and 73 atm.

Structures of Solids

• Crystalline solids have a well-ordered, definite arrangement of molecules, atoms, or ions.
• An example of a crystalline solid would be quartz, salt, and sugar.
• The intermolecular forces are similar in strength, and tend to melt at specific temperatures.
• Amorphous solids don't have an orderly arrangement of molecules, atoms, or ions.
• Examples of amorphous solids: rubber, glass
• Amorphous solids have intermolecular forces that vary in strength and tend to melt over a range of temperatures.
• Crystalline solids have an ordered, repeating structure, where the smallest repeating until in crystal is a unit cell, the 3D stacking of unit cells is the crystal lattice.
• There are three types of cubic unit cells: Primitive, Body-centered, and Face-centered.
• Atoms are positioned at the corners of a simple cube, with each atom being shared by eight unit cells in primitive cubic cells.
• Body-centered cubic cells-atoms are at the corners of a cube, plus one in the center of it's body.
• Atoms in the corner of a cube are shared by eight unit cells, and the corner atom is enclosed by one unit cell in face-centered cubic cells.
• There are atoms at the corners of a cube plus one atom in the center of each face of the cube in face-centered cubic cells.
• Eight unit cells share it's corner atoms, and two share the face atoms in face-centered cubic cells.
• Face-centered cubic lattice has two equivalent ways to decide a unit cell:
  • Cl- (larger) ions at the corners of the cell,
  • Na+ (smaller) ions at the corners of the cell.
• Cations to anion rations in a unit cell are the same for a crystal.

Close Packing Spheres

• Crystalline solids maximize attractive forces between particles.
• Particles modeled by spheres, where each atom or ion is a sphere.
• Molecular crystals form by close packing of molecules.
• Maximum intermolecular forces in crystals are met by close packing of spheres.
• Spheres are tightly packed, but there are still spaces between them called interstitial holes.
• A crystal is built layering of close packed spheres.
• Sphere-placement can only happen in only one place for the second layering.
• There are two choices to make for the third layer:
  • Third layer eclipses the first (ABAB arrangement), and is called hexagonal close packing (hcp);
• The third layer is in a different position to the first (ABCABC arrangement), and is called cubic close packing (ccp).
• The unit cell of a ccp crystal is face-centered cubic.
• In both close-packers structures, each sphere is surrounded by 12 spheres
• The coordination number is the number of spheres directly surrounding a central sphere.
• If unequally sized spheres are used, the smaller spheres go in the interstitial holes; an example of this is Li20.
• The larger O² Ions assume the cubic closed-packed structure with the maller Li*ions in the holes.
• When waves pass through narrow slits they are diffracted.
• When waves pass through a diffraction grating (many narrow slits), they interact to form a diffraction pattern (light/dark bands).
• Diffraction, works when light wavelengths are close to slits size.
• The typical spacing between crystal layers (2–20 angstrom), is the typical wavelength range that X-rays use.
• X-ray diffraction (X-ray crystallography):
• When X-rays get passed through a crystal, they are detected on a photographic plate, that have a bright spot at the center (incident beam), along with a diffraction pattern.
• The way a close-packing is arranged produces a unique diffraction pattern.
• By knowing diffraction patters we can calculate the positions of atoms to determine the crystal structure, by using the diffraction pattern.

Bonding in Solids

• There are four kinds of solids: molecular, covalent network, ionic, and metallic.
• Molecular solids consist of atoms/molecules held together by intermolecular forces.
• Weak intermolecular forces give rise to low melting points, which include dipole-dipole, London dispersion, and H-bonds.
• Molecular solids are generally soft, can pack efficiently because packing is important (since the solid is not regular spheres), exhibity poor thermal/electrical.
• Ar, CH4, CO2, and sucrose are examples of molecules that make up molecular solids.
• Covalent-network solids made of atoms held together in chains/networks with covalent bonds; they have very high melting points and are much harder than molecular solids.
• Strong solids result from covalent bonds that connect the atoms.
• An example of this would be diamond, graphite, quartz (SiO2), silicon carbide (SiC), and boron nitride (BN)
• Each C atom of diamond has coordination number of 4, is tetrahedral, has a 3D array structure.
• Diamond bonds very and has a high melting point (3550°C).
• Each atom of graphite is arranged in a planar hexagonal ring, that are placed on top of each other.
• The adjacent C atoms in each layer are close to the arrangement seen in benzene (1.42 angstrom vs. 1.395 A), which allows electrons to flow in delocalized pattern that's good for conductivity.
• The distance between layers of graphite is large, where a large space (3.41 angstrom) is connected by weak forces.
• Layers can slide easily as graphite is a good lubricant.
• Ionic solids are ions held together by ionic bonds and force of electrostatic attraction.
• F = k*Q1+Q2/d, where a higher charge and smaller length makes stronger the force The charges/lengths are determined by the atoms

Lattice Types

• The NaCI lattice has a cation-anion ration of 1:1, a Face-centered cubic (FCC) arrangement, as well as 6 coordination numbers.
• Other compounds in the same arrangement in NaCI include LiF, KC!, AgCl, and CaO.
• The CsCl structure has cation-anion ratios of 1:1, has 8 coordination numbers, and differs from NaCI since Cs is much bigger then NA
• Zinc blende features an arrangement of ions.
• Each has coordination of 4/7, and the S has a Face-centered cubic (FCC) arrangement, however, each ions are placed uniquely
• Other compounds with arrangements like this include are CuC
• Another arrangement includes fluoride (CaF2) with a FCC arrangement, each Calcium has unique placement, for example. Other compounds that features arrangements like Barium(II) Chloride, And Lead(II) Fluoride
• Metallic solids consist entirely of metal atoms.

Metallic Solid Properties

• Can be soft or hard, has high melting point, has good electro/thermal stability, is able to be cut.
• Metals like Al, Cu, and Fe can exist as bodies/Faces with unique arrangements.
• The coordination number for an atoms is either 8/12, but those don't explain strong boding, due the the amount of not bond atoms.
• It bonds too strongly to be bonded, metals nuclei is able to flow in delocalized arrangements.
• Metals conducts when it mobile and moves to other places.