5. Phases and Solution

What determines the phases of matter?

Intermolecular forces (strong = solid, weak = gas)

Solids

Fixed volume and shape; atoms vibrate in place; crystalline or amorphous

Crystalline Solids

Regular arrangement of atoms (lattice); Ex. NaCl

How does H₂O solvate NaCl?

Dissociation by separating NaCl via ion-dipole interaction

Lattice Energy

Amount of energy required to separate a solid into component cations and anions (usually very large amount)

Amorphous Solids

No repeating arrangement of atoms (SiO₂; silica)

Liquids

Fixed volume, not shape (non-compressible); have viscosity

Viscosity

Resistance of a liquid to flow

Gases

No fixed shape or volume; compressible (no constant density)

How does decreasing volume affect gas density?

Density increases since IMFs increase (particles more close together)

How does increasing volume affect gas density?

Density decreases since IMFs decrease (particles far apart)

Ethane vs. Ethanol

Ethane is more likely to be a gas and ethanol is more likely to be liquid/solid (Ethane only has LDF, Ethanol can H-bond due to OH group)

Affect of high temperature on phases:

High temperatures confer more kinetic energy, making it easier to overcome IMFs → separate

Affect of low temperatures on phases:

Slow down molecules and lower the kinetic energy, making it more likely a molecule will stay together/form IMFs

Affect of high pressure on phases:

Force particles to interact more closely, making it more likely for solids or liquids to form

Phase Changes:

1. Melting/fusion

2. Freezing

3. Vaporization

4. Condensation

5. Sublimation

6. Deposition

Melting/fusion

Solid → Liquid

Freezing

Liquid → Solid

Vaporization

Liquid → Gas

Condensation

Gas → Liquid

Sublimation

Solid → Gas

Deposition

Gas → Solid

Which phase changes are endothermic?

Melting/Fusion, evaporation, sublimation (require energy to break apart, +ΔH)

Which phase changes are exothermic?

Freezing, condensation, deposition (release energy upon forming IMF, -ΔH)

Heat of Fusion

Amount of heat used to disrupt intermolecular interactions during melting/fusion at a constant temperature

Melting Point of Water =

0^o C (273 K)

Boiling Point of Water

100^o C (373 K)

Heat of Vaporization =

Heat required to convert a liquid to gas at a constant temperature

When is temperature constant?

During a phase change (because heat only goes toward breaking IMFs, not increasing temperature)

Units of Heat of Fusion/Evaporation

kJ/mol

What does temperature measure?

The kinetic energy of particles

Specific heat Capacity (c)

Amount of heat required to raise the temperature of 1 unit mass of substance by 1^o C (J/g*C)

Specific Heat Capacity of Water

4.184 J/g*C

Why does water have a large heat capacity?

Due to hydrogen bonding; water can take a lot of heat without its temperature changing

q_f/q_v (heats of fusion/vaporization) =

nΔH; n=moles

q =

mcΔT

When should you use q=mcΔT?

Between phase changes when temperature is changing

When should you use q=nΔH?

When a phase change is occurring.

q = mL

Variation of q=nΔH, where L = latent heat of fusion

Phase Diagrams

Shows the temperature and pressure conditions for phase changes of a substance; solid lines are phase changes

SLUG (Phase Diagram)

Solid → Liquid → Gas in clockwise order

Why is water phase diagram unique?

Ice is less dense than water (unlike other solid forms of substances), so an increase in pressure results in melting rather than freezing

Triple Point in Phase Diagram

Solid, Gas, and Liquid at equilibrium

Critical Point in Phase Diagram

End of liquid-gas interface (liquid is a supercritical fluid with properties of both gas and liquid)

Vapor Pressure

Pressure exerted by gas molecules; at equilibrium, # of molecules escaping to gas phase = # of molecules condensing to liquid

Vapor pressure increases with…

Increasing temperature since KE increases

When vapor pressure = atmospheric pressure,

Boiling point is reached (Ex. water boils at 100 degrees, 1 atm)

STP =

1 atm, 273 K (0 C)

1 atm =

760 mmHg

1 mol gas @ STP =

22.4 L

Ideal Gas Laws (Kinetic Molecular Theory):

1. Gases are made of atoms/molecules in continuous, random motion

2. Average KE is directly proportional to temp; all gases have the same KE at a given temp

3. Gas particles have no volume

4. Gas particles don’t exert attractive or repulsive forces on each other, only the container

5. All collisions are elastic, KE is conserved

When do gases deviate from ideal behavior?

High pressures and low temperatures (because likely to become liquid)

When do gases behave ideally?

High temperatures and low pressure

For gas law equations, what temperature scale should you use?

KELVIN always

Celsius → Fahrenheit:

9/5C + 32 (Approximate 2C^o +32)

Boyle’s Law (constant temp and moles):

P₁V₁ = P₂V₂ (PV = constant)

Charle’s Law (Constant Pressure and moles):

V₁/T₁ = V₂/T₂ (V/T = constant)

Gay-Lussac’s Law (Constant Volume and moles):

P₁/T₁ = P₂/T₂ (P/T = constant)

What does Boyle’s Law tell us?

If volume increases, pressure decrease; vice versa

What does Charle’s Law tell us?

When temperature increases, volume increases; vice versa

What does Gay-Lussac’s Law tell us?

When pressure increases, temperature increases; vice versa

Ideal Gas Law

PV = nRT (R = .08206 L*atm/K*mol)

Van der Waals Equation (For non-ideal gases):

nRT = (P + an²/V²)(V - nb)

a = attractive force

b = volume occupied by mol of gas

n = moles of gas

Van der Waals Accounts for:

1. Volume taken up by gas molecules (V - nb)

2. Attractive forces experienced by gas (P + an²/V²)

   Focus on knowing that non-ideal gases can be modeled, and that ideal gases have high T, low P

Mole fraction (X_gas)

Number of moles of a given gas divided by total moles of gas

X_gas =

n_gas/n_total

Derived Mole Fraction (X_gas) =

(n_gas/n_total) = (P_gasV/RT)/(P_totalV/RT) = P_gas/P_total

Partial Pressure

Pressure that a gas in a mixture exerts if it took up the same volume by itself

P_gas (Partial Pressure) =

X_gasP_total

Dalton’s Law

Total pressure of a mixture equals sum of partial pressures

P_total =

ΣP_partial = P_gas1 + P_gas2…

Solution

Homogenous mixture of evenly distributed particles (usually containing solids dissolved in liquids on MCAT)

Dissolution

Solute is dissolved in a solvent (based on solute-solvent interactions - like dissolves like)

Like dissolves like

Hydrophilic (charged or polar) dissolves hydrophilic; vice versa

Molarity (M)

Moles solute/Liters solution (mol/L)

Molality

Moles solute/kg solvent (mol/kg)

ppm (parts per million)

mg/L

ppt (parts per thousand)

g/L

Colligative Properties

1. Vapor pressure reduction

2. Boiling point elevation

3. Freezing point reduction

4. Osmotic pressure

High vapor pressure =

Low boiling point

Vapor Pressure Reduction

Results in higher boiling points

Raoult’s Law (Vapor Pressure Reduction):

P = X_AP_A^o

P = vapor pressure of solution (reduced)

X_A = mole fraction of solvent

P_A^o = vapor pressure of pure solvent

Boiling Point Elevation:

ΔT_b = iK_bm

ΔT_b = how much boiling point was elevated

i = van ‘t Hoff factor (ionization factor)

K_b = boiling point elevation constant

m = molal solute concentration (mol solute/kg solvent)

Ionization (van ‘t Hoff) Factor

Number of ions a mole of solute

What is i of NaCl?

2; Na⁺ and Cl^-

What is i of C₆H₁₂O₆ (glucose)?

0; no ions

Osmosis

Net flow of solvent through a semipermeable membrane from high solvent to low solvent (low solute → high solute); trying to lower concentration of solute

Osmotic Pressure (Π)

Pressure required to prevent osmosis (when solute and solvent are =)

Π (Osmotic Pressure) =

iMRT

i = van ‘t Hoff

M = Molarity (mol solute/L)

R = .08206

T = Kelvin

Solubility

Extent to which various substances will dissolve in a solvent

Solubility equilibrium

Solute dissolves at the same rate is precipitates out of solution

K_sp (Solubility Product Constant) =

[Ion 1]^x[Ion 2]^x (Solids and liquids not included)

When solution is not in equilibrium, use:

Q_sp (same equation, but can compare to K_sp)

High K_sp means…

The compound is soluble

Low K_sp means…

The compound is insoluble

Q_sp > K_sp

Supersaturated solution, precipitate will form

Q_sp<K_sp

Undersaturated solution, more solute can dissolve

Q_sp = K_sp

Solution is at equilibrium (dissolution = precipitation)

Common Ion Effect

Decrease in solubility of a species when a soluble salt is added into the mixture (related to Le Chatlier)

Electrolytes

In solution, charged cations and anions conduct electricity (strong acids, ionic compounds with high K_sp)