MCAT Chemistry

Phases of Matter

MCAT Chemistry > Phases of Matter

Background

Phase Equilibria occurs in an isolated system and where phase changes are reversible

MCAT Phases of Matter - Phases Equilbria

Evaporation/Vaporization is the transition from liquid to gas, when atoms on top of a solution gain enough kinetic energy to eject and escape. Each time a high-energy particle escapes, the temperature of the solution decreases because evaporation is an endothermic process.

Given enough energy, a liquid will completely evaporate. In a covered/closed container, molecules get trapped and exert a countering pressure which forces some of the gas back into liquid phase through condensation. Vapor pressure increases as temperature increases since more molecules escape from the top

Condensation is when a force such as a lower temperature or higher pressure is exerted on a gas such that it is forced back into the liquid state

Boiling Point is the temperature at which vapor pressure equals the ambient (external) temperature and evaporation/vaporization occurs throughout the liquid

Fusion/Melting is the transition from solid to liquid and occurs as the vibrational motions of a solid increase from heat application, as atoms have greater freedom of movement, and energy ultimately disperses.

Melting Point of a solid is the temperature at which it changes to a liquid

Solidification or Crystallization or Freezing is the transition from liquid to solid

Freezing Point of a liquid is the temperature at which it changes into a solid

Sublimation is when a solid transforms directly to gas (dry ice/CO2)

Deposition is the reverse of sublimation when a gas deposits directly to a solid

MCAT Phases of Matter - Phases of Matter

Triple Point is the point at which solids, liquids and gases exist in equilibrium

Critical Point is the theoretical phase boundary that exists when two densities of molecules in different phases become equal and there is no distinction between the phases

Solids

Solids have rigidity, resistance to flow and are usually the densest state of a molecule. Water and Ice are an exception to this. Motion is limited to vibration (KPE) and volume doesn't change with pressure.

Amorphous Arrangement lacks ordered 3D arrangement (glass, plastic, wax) and tends to melt/solidify over a large range of temperatures

Crystalline Arrangement is an ordered structure lattice with repeating ion, atom, or molecular patterns (most solids) and has distinct melting points

Ionic Solids are infinite aggregates of positive and negatively (+/-) charged ions that repeat based on alternating cation/anions and do not contain discrete molecules. Empirical formulas are used to calculate the lowest common denominator

Physical properties include...

  • High Helting Point (MP)
  • High Boiling Point (BP)
  • Poor Electrical Conductivity in Solid State
  • High Electrical Conductivity in Molten/Aqueous States
  • Relatively immiscible

Most ionic compounds are structured in the following formations

  • Simple Cubic
  • Body-centered cubic
  • Face centered cubic

Metallic Solids are metals that are stacked as close together as possible. They have a high melting point and high boiling point because of strong covalent interactions. Pure metallic masses are spheres of similar radii stacked in a staggered formation

Liquids

Liquids are fluids that flow and conform to the shape of a container. They have a high degree of freedom through diffusion. Water can mix with stuff (like dissolves like) and provides a solvent for which reactions may occur in.

Miscibility is the degree to which water will mix with a substance

MCAT Phases of Matter - Miscible Solution

Miscible = Ethanol and Water

MCAT Phases of Matter - Immiscible Solution

Immiscible = Oil and Water (creates an emulsion if left to separate)

Colligative Properties are physical properties of solutions dependent upon the concentration of dissolved particles but not chemical identity. For example there is a measureable change in boiling point of a solution compared to that of the pure solvent

Vapor Pressure Depression (Raoult's Law) states when solute is dissolved in a solvent, the new solvent has a lower vapor pressure than that of the pure solvent

Compound A pureform = Xa = 1.0, BP = 100°C

Compound of B pureform = Xb = 1.0, BP = 80°C (more volatile)

ΔP = P°A – PA

ΔP = XBP°A

PA = XBP°B

Molality is the moles of solute per kilogram of solvent

M = (mol solute)/(kg solvent)

Boiling Point Elevation is when a solute is dissolved in a solvent and the boiling point of the solution is greater than that of the pure solvent. Adding solute results in a decrease in vapor pressure of the solvent. If vapor pressure of a solution is lower than that of a pure solvent, more energy (higher temperature) must be obtained before equilibrium is

ΔTb = ikBM

  • I = van't Hoff factor (moles of particles dissolved per mole of solute particles, NaCl = 2)

Freezing Point Depression is when a greater amount of energy must be removed from the solution (lower temp) in order for solution to solidify

ΔTf = = ikfm

Osmotic Pressure is the pressure, which needs to be applied to a solution in order to prevent the inward flow of water across a semipermeable membrane (such as the cell). It is also the pressure needed to fully nullify osmosis.

Π = iMRT

  • R = ideal gas constant (8.21 x 10-2 (Latm)/(molK) or 8.314 J/Kmol)

Gasses

Gasses are classified as fluids because they can flow and conform to the volume of a container. However, atoms of gases move rapidly and far away from each other.

Pressure (Pa) is measured through a mercury barometer

1 atm = 760 mm Hg = 760 torr = 101.325 kPa

Ideal Gasses represent hypothetical gases whose molecules have no intermolecular forces and occupy no volume. Deviation occurs at high pressures and low temperatures

Real Gases differ from ideal gases when exposed to high pressure, low temperature or close molecular proximity. At high pressure, volume is higher than an ideal gas predicts, and occurs until gas enters the liquid phase.

Kinetic Molecular Theory of Gases explains behaviors of gases. Gases show similar physical characteristics and behavior regardless of their particular chemical identity

At low temperatures, average kinetic energy decreases, and attractive intermolecular forces become more significant. As the temperature is reduced and the condensation/boiling point is approached, the intermolecular forces cause gas to have smaller volume than predicted, and gasses eventually condense to liquids. The closer to the boiling point, the less ideal the behavior

  • 1) Gases consist of particles whose volumes are negligible compared to container
  • 2) Gas atoms or molecules exhibit no intermolecular attractions or repulsions
  • 3) Gas particles are in continuous, random motion, and always undergo collisions with other gas particles and the container walls
  • 4) Collisions between any two gas particles are elastic, meaning there is a conservation of both momentum and kinetic energy
  • 5) The average kinetic energy of gas particles is proportional to the absolute temperature (Kelvin) and is the same for all gases at a given temperature, irrespective of chemical identity or atomic mass

Average Molecular Speeds is related to absolute temperature. To find velocity, figure out average kinetic energy and calculate corresponding speed

KE = ½mv2 = 3/2kT

  • k = Boltzmann Constant = 5.67 x 10-8 J/Sm2K4

μrms = √3RT/Mm

  • R = 8.314 J/kmol

Graham's Law of Diffusion and Effusion states that the movement of gas molecules through a mixture (air) is diffusion. Kinetic molecular theory predicts that heavier gases will diffuse more slowly than lighter ones because of differing average speeds All particles have same average kinetic energy at same temperature

Effusion is the flow of gas particles from one compartment to another through a small opening. At same temperature and pressure, rates of effusion are proportional to average speeds. A gas with a molar mass of 4x another gas will travel ½ as fast as the lighter gas

r1/r2 = √((Mm)2/(Mm)1)

  • r = rates of diffusion for two gases

Ideal Gas Behavior explains how all gases at a constant temperature and pressure occupy the same volume that is directly proportional to the number of moles of gas present

  • One mole of any gas = 22.4 L at STP

n/V= k

n1/V2 =n1/V2

Ideal Gas Law

PV = nRT

  • n = number of moles
  • R = 8.21 x 10-2 (L.atm)/(mol.K)

Density is the ratio of mass per volume and expressed in g/L

ρ = M/V = P(Mm)/RT

V2 = V1 (P1/P2)/(T2/T1)

Boyle's Law states for a gas at a constant temperature (isothermal) the volume of a gas is inversely proportional to its pressure

P1V1 = P2V2

Law of Charles/Gay-Lussac states at a constant pressure (isobaric), the volume of the gas is proportional to its absolute temperature (K)

V/T = k

V1/T1 =V2/T2

V/T = nR/P = constant

Dalton's Law of Partial Pressure states when 2+ gases are independently behaving in one container and no chemical interactions occurring, pressure exerted as if only one gas

PT = PA + PB + PC

PA = PTXA

XA = na/nt

Van Der Waals Equation of State

(P + (n2a)/V2)(V – nb) = RT

  • a = constant, corrects attractive forces (smaller for small/less polarizable atoms, ex: He)
  • b = constant, corrects for volume (larger for large/more polarizable atoms, ex: HCl, NH3)

MCAT Chemistry

Chemistry Topics