Phases and Phase Equilibria

Question 1

Convert between Kelvin and Celsius

Answer 1

K = °C + 273

Question 2

Convert between Farenheight and Celsius

Answer 2

F = °C x 1.8 + 32

Question 3

Pressure

Answer 3

Pressure is the force exerted over an area: P = F/A

Question 4

What happens to pressure at higher elevations?

Answer 4

Pressure decreases at higher elevations

Question 5

Simple mercury barometers were historically particularly useful because the high density of mercury allows a barometer to be only approximately 1 m tall to measure normal fluctuations in atmospheric pressure, which determined the height of the mercury. If a barometer reads a height of 1000 mmHg, what height would be read on a barometer using water instead? The density of mercury is 13.5 g/mL and that of water is 1 g/mL

Answer 5

A simple barometer setup has the height of the mercury controlled by the atmospheric pressure. We’re told that the reading of a barometer is 1000 mmHg, which means, using P = ρgh, that the pressure is P = (13.5 kg/L)(10 m/s2)(1 m) = 135 N/m2. That is the same pressure that a water barometer would experience. The issue is that water has a lower density. So we can set it up as 135 = (1)(10)h and therefore h = 13.5 m. This should make sense since mercury is 13.5 times denser so the water would need to be 13.5 higher to achieve the same pressure.

Question 6

Ideal Gases

Answer 6

ideal gases occupy 22.4 L per mol of molecules.

Question 7

Definition of Ideal Gas

Answer 7

• An ideal gas consists of pointy dots moving about randomly and colliding with one another and with the container wall. The ideal gas obeys the kinetic molecular theory of gases and has the following properties.
  
  • Random molecular motion.
  • No intermolecular forces.
  • No (negligible) molecular volume.
  • Perfectly elastic collisions (conservation of total kinetic energy).

Question 8

At what pressures and temperatures can you treat gases as ideal gases?

Answer 8

• You can treat gases as ideal gases at:
  
  • Low pressures
  • High temperatures

Question 9

Ideal Gas Law

Answer 9

PV=nRT, where P is pressure, V is volume, n is # mols of gas, R is the gas constant, and T is temperature

Question 10

Combined Gas Law

Question 11

Boyle’s Law

Answer 11

Boyle’s law: at constant temperature, P₁V₁ = P₂V₂

Question 12

Charles Law

Answer 12

Charles’ law: at constant pressure

Question 13

Kinetic molecular theory of gases - Assumptions

Answer 13

• Random molecular motion.
• No intermolecular forces.
• No (negligible) molecular volume.
• Perfectly elastic collisions (conservation of total kinetic energy).

Question 14

The kinetic theory holds the following concepts about pressure and temperature:

Answer 14

• Pressure is equally distributed over the walls of the container because molecular motion is random.
• Higher temperature means the molecules are traveling faster, lower temperatures means slower molecules.

Question 15

Diffusion

Answer 15

random molecular motion, causing a substance to move from an area of higher concentration to an area of lower concentration (diffusion down its concentration gradient).

Question 16

Effusion

Answer 16

random molecular motion, causing a substance to escape a container through a very small openning.

Question 17

Graham’s Law

Answer 17

(Rate1/Rate2) = √(M2/M1)

Question 18

Kinetic Theory in terms of temperature and mass

Answer 18

(v1/v2) = √(m2/m1)

Question 19

Van der Waals’ equation

Answer 19

b for bounce. The term with the constant b is the repulsion term. The greater b is, the more repulsion, which leads to greater pressure.
a for attraction. The term with the constant a is the attraction term. The greater a is, the more attraction, which leads to less pressure.

Question 20

Partial pressure

Answer 20

Partial pressure = a component of the total pressure exerted by a species in a gas mixture.

The total pressure of a mixture of gas = The sum of all the partial pressures.

Question 21

Mole Fraction

Answer 21

Mole fraction = a component (fraction) of the total # mols that belongs to a species in a gas mixture. Mole fraction for species A = # mols of A / # mols of the entire gas mixture. = # mols of A / Σ # mols of A, B, C …

Question 22

Dalton’s law relating partial pressure to composition

Answer 22

P_i = χ_i·P_total

P_total = ΣP_i = Σχ_i·P_total

Question 23

Hydrogen bonding

Answer 23

Hydrogen bonding is a weak interaction between a partially positive H and a partially negative atom.
Technically, hydrogen bonds are a special type of dipole-dipole interaction.
Hydrogen bonding increases the boiling point.
Partially positive H are also called hydrogen bond donors. They are hydrogens that are bonded to either F, O, or N.
Partially negative atoms are also called hydrogen bond acceptors. They are most commonly F, O, or N.

Question 24

Dipole-Dipole Interactions

Answer 24

All polar molecules exhibit dipole-dipole interactions. This is where the polar molecules align such that opposites attract.
Dipole-dipole interactions increase the boiling point, though not as significantly as hydrogen bonding.
Dipole interactions are stronger the more polar the molecule is.

Question 25

Ion-Dipole Interactions

Answer 25

Ion-dipole interactions are similar to dipole-dipole interactions, but it’s stronger because it is no longer an interaction involving just partial charges. Instead, it is an interaction between a full charge (ion) and a partial charge (dipole).

Ion-dipole interactions get stronger when you have larger charge magnitude of the ion, and large polarity of the dipole molecule.

Question 26

Van der Waals’ forces (London dispersion forces)

Answer 26

Dispersion forces exists for all molecules, but are only significant for non-polar molecules. For polar molecules, dipole forces are predominant.
Dispersion forces result from induced and instantaneous dipoles.

• Induced dipoles: when a polar molecule interacts with a non-polar molecule, then polar molecule induces a dipole in the non-polar molecule.
• Instantaneous dipoles: Non-polar molecules have randomly fluctuating dipoles that tend to align with one another from one instant to the next.

Question 27

Phase Diagram

Question 28

Molality

Answer 28

Molality = mols of solute / mass (in kg) of solvent.

Question 29

Colligative properties

Answer 29

Colligative properties = properties that depend on the # of solute particles, but not on the type.

Question 30

Van’t Hoff Factor (i)

Answer 30

all colligative properties take into consideration of the Van’t Hoff factor. Basically, it means convert concentration to reflect the total number of particles in solution. For example, glucose has i of 1 because it doesn’t break up in solution. NaCl has i of 2, because in solution, it breaks up into 2 particles Na+ and Cl-.

Question 31

vapor pressure lowering (Raoult’s law)

Answer 31

ΔP = χ_solute·P°_solvent

Question 32

boiling point elevation

Answer 32

ΔT_b = k_b·m·i

Question 33

freezing point depression

Answer 33

ΔT_f = -k_f·m·i

Question 34

osmotic pressure

Answer 34

π = MRT *i

π is the osmotic pressure; M is the molarity in mol/L; R is ideal gas constant; T is the temperature in K.

Question 35

Osmosis

Answer 35

Osmosis is the movement of solvent across a semi-permeable membrane from an area of low solute concentration (high solvent concentration) to an area of high solute concentration (low solvent concentration).

Question 36

Henry’s Law

Answer 36

P_solute = k [solute]