Chem: IMFs, Phase Behavior, Solutions, and Colloids

for chem 2 first test

intermolecular forces

attractive forces between molecules that arise from interactions between temporary, partial, or full charges in atoms, ions, and molecules

London (dispersion) forces

weak attractive forces caused by instantaneous-induced dipole interactions, all atoms and molecules

dipole-dipole

attractive forces between permanent dipoles, all polar molecules

hydrogen bonding

strong attractive forces between molecules with hydrogen bonded to a very electronegative atom, all molecules with hydrogen bonded to N, O, or F

ion-dipole

strong attractive forces between ions and molecules with a permanent dipole moment, any solutions formed between soluble ionic compounds and polar molecules

gas to solid and solid to gas

gas to solid: deposition

solid to gas: sublimation

when phase transitions occur

when the temperature (or pressure) allows for partial or complete breaking (or forming) of intermolecular forces between molecules

London-dispersion strength increases when

increasing molar mass

increasing surface area

highest boiling points

London-dispersion (compare molar mass and surface area) → polar → hydrogen bonding

vapor

a gas formed from molecules that escape from the surface of a liquid

vapor pressure

the equilibrium partial pressure of a vapor at a given temperature (stronger IMFs mean lower vapor pressure) (increases with increasing temperature) (normal boiling point is the temperature at which a substance’s vapor pressure is at 1 atm)

Clausius-Clapeyron equation

relationship between temperature and vapor pressure

enthalpy of vaporization

the amount of energy required to vaporize a substance per mole

heating curve

heat of vaporization and fusion: at a phase transition temperature, all of the heat added or removed goes into breaking or forming IMFs

otherwise, the temperature increases linearly when a substance is heated based on its specific heat at a given phase

phase diagram

a plot of pressure vs. temperature that shows the range of phase behavior for a given substance

triple point

point at which 3 phases are in equilibrium

critical point

point at which there is no distinction between the liquid and gas phases (substance forms a supercritical fluid)

phase equilibrium curves

points at which 2 phases are at equilibrium

crystalline solids

composed of atoms, ions, or molecules that have a long range repeating order called a crystal lattice

unit cell

the smallest divisible unit that repeats in all directions (most common is a cubic cell with equal edge lengths and 90* angles)

simple cubic

one atom at each vertex of a cube, 1 atom per unit cell, edge length= 2r

body centered cubic

one atom at each vertex and the center of a cube, 2 atoms per unit cell, edge length= 4r/square root of 3

face centered cubic

one atom at each vertex and each face of a cube, 4 atoms per unit cell, edge length= 2xsquare root of 2xr

molecular solids

a crystalline solid, held together by IMFs, relatively low melting points

ionic solids

crystalline solids, held together by electrostatic forces, relatively high melting points

nonbonding solids

atomic solid, held together by London-dispersion forces, relatively low melting points

covalent network solids

atomic solid, held together by covalent bonds, extremely high melting points (ex: diamond, graphite, quartz)

metallic solids

atomic solid, cations in an “electron sea”, variable melting points

solutions

homogeneous (uniform) mixtures where solute particles are individual atoms and molecules

solute

pure substance that is being dissolved

solvent

the pure substance that is doing the dissolving

enthalpy of solution

enthalpy of lattice: energy absorbed to break the solid lattice into individual ions, endothermic, >0

H₂O molecules make room for ions and that is a tiny bit endothermic

enthalpy of hydration: energy released when water molecules surround ions in solution, exothermic, <0

miscibility

the ability for pure substances to mix without forming layers, depends on the enthalpy of solution and the natural tendency to maximize entropy (randomness)

“like dissolves like”

pure substances with similar intermolecular forces are likely miscible

solubility

the amount of solute that must be added to a given amount of solvent to form a saturated solution (solid in liquid increases with increasing temperature) (gas in liquid decreases in increasing temperature and increases with increasing pressure) (Henry’s Law)

molarity (M)

moles of solute / liters of solution

mole fraction (X_A)

moles of A / total moles in solution (total will equal 1)

molality (m)

moles of solute / kilograms of solvent (does not change with temperature)

percent by mass (parts per hundred)

(mass of solute / mass of solution) x 100%

colligative properties

solution properties that depend on concentration but not the identity of the solute (vapor pressure lowering, freezing point depression, boiling point elevation, osmotic pressure)

Raoult’s Law

relates the pressure of a solution to the mole fraction of the solvent

Dalton’s Law

for ideal solution of volatile substances, each substance contributes to the total vapor pressure

freezing point and boiling point

the change in the boiling point and freezing point temperatures depend on the molality (m) of the solution

osmosis

the flow of solvent from low to high concentration through a semi-permeable membrane

osmotic pressure

the amount of pressure needed to stop osmosis in a system

the van’t Hoff factor (i)

the ratio between the moles of ions and formula units dissolved in solution, estimated from the chemical formula of strong electrolytes, experimentally it is lower than expected due to ion pairing (ex: Ca(NO₃)₂ = 3)

colloids

are homogeneous mixtures where solute particles are 1nm-1um in diameter, colloid particles are large enough to scatter visible light, distinguished using the Tyndall effect

colloidal particles: solid or liquid

solvent: gas

aerosol (ex: smoke, fog)

colloidal particles: gas

solvent: solid or liquid

foam (ex: pumice, whipped cream)

colloidal particles: liquid

solvent: liquid

emulsion (ex: milk)

colloidal particles: liquid

solvent: solid

gel (ex: jelly)

colloidal particles: solid

solvent: liquid

sol (ex: paint)

soaps

acts as emulsifiers, ionic head that is hydrophilic, non-polar tail that is hydrophobic, form micelles that encase dirt and oils