Chemistry 101: Gases, Liquids, and Solids

A phase is a physically distinct form of a substance that can be separated from another form. Generally phase is used interchangeably with “state of matter”.

Solid, liquid, and gas are the three phases of matter that most chemistry courses, as well as the AP Chemistry exam, will test.

Temperature (Heat, Enthalpy) can be added or removed. In general, an increase in temperature drives the phase change solid⇒liquid⇒gas.

Pressure (Volume) can be added or removed. In general, an increase in Pressure (or decrease in Volume) drives the phase change gas⇒liquid⇒solid.

1. solid⇒liquid is fusion (melting)

2. liquid⇒solid is freezing (solidification)

1. gas⇒liquid is condensation

2. liquid⇒gas is vaporization (boiling)

1. gas⇒solid is deposition

2. solid⇒gas is sublimation

• A represents the Solid phase region.

• B represents the Liquid phase region.

• C represents the Gas phase region.

• A is Fusion (or Melting)

• B is Freezing (or Solidification)

• A is Vaporization (or boiling)

• B is Condensation

• A is Sublimation

• B is Deposition

• A is the Triple Point.

• B is the Critical Point.
  Note: though the concept of Plasma exists, the AP Chem exam does not explicitly test this as a phase.

The critical temperature and critical pressure combine to be the critical point:

• the critical temperature is the temperature above which a distinct liquid-to-gas vaporization can no longer be accurately determined.

• the critical pressure is the pressure above which a distinct gas-to-liquid condensation can no longer be accurately determined.

The triple point is the point at which a substance can exist in equilibrium in all three states (solid, liquid, and gas).

This means that instantaneously, a substance at its triple point can interconvert between any phase.
Ex: water at its triple point would exist as ice, liquid water, and steam - all at one temperature and pressure.

Water has a slightly negative sloped Solid/Liquid boundary.
This is due to water’s Solid phase being less dense than its Liquid phase.

The difference is due to intermolecular forces.

The attraction of molecules to each other determines at which temperature mixtures will have phase changes, and subsequently the amount of heat necessary to make that transition.

From strongest (1) to weakest (4), they are:

1. Ionic Forces

2. Hydrogen Bonds

3. Dipole-Dipole

4. London Dispersion Forces

Dipole forces occur when a molecule with polar bonds has the geometry to become overall polar. This results in partially positive and negative charges, and these opposite charges attract other charged molecules in the mixture.

Hydrogen bonds are a product of H being covalently bonded to either O, N, or F. This is an extreme form of the dipole-dipole force.

The high electronegativity difference between these atoms and hydrogen creates strong dipoles, which consequently result in the strongest dipole-dipole interactions between molecules.

Dispersion forces (also called Van der Waals forces) are an instantaneous polarization of molecules that would otherwise be non-polar.

This process is also referred to as an induced dipole because there is an instantaneous reorganization of the electron cloud leading to partial polarization. The more valence electrons present in either a molecule or in a mixture, the more possible repulsion of electron clouds, and the more induced dipoles.

Ionic forces occur when two oppositely charged ions (cations and anions) attract each other.

In general, ions attract to a polar molecule first, then attract each other - hence this is often called Ion-Dipole force. (Two Ions attracting each other would otherwise prefer to Ionic Bond, which is an intramolecular force, not an intermolecular force.)

The stronger the intermolecular forces present, the higher the boiling point. This is due to the molecules being more tightly attracted to each other is the liquid phase.

In order of highest (1) to lowest (4) boiling points:
Ionic Forces
Hydrogen Bonds
Dipole-Dipole
Dispersion Forces

1. Ionic Forces

2. Hydrogen Bonds

3. Dipole-Dipole

4. London Dispersion Forces

Hv is the heat of vaporization. It represents the energy necessary to convert a liquid to a gas.

Hf stands for the heat of fusion. This is the amount of energy that is necessary to convert a solid to a liquid.

• At A, fusion is occurring.
  Hf is used to calculate the heat needed for plateau A.

• At B, vaporization is occurring.
  Hv is used to calculate the heat needed for plateau B.

Zero slope means no temperature change. This occurs at these times because any heat input is dedicated to breaking bonds between molecules, instead of increasing temperature.

Ex: during the vaporization of water, heat is needed to break the hydrogen bonds between liquid water molecules and convert it to gaseous steam.

q = mcΔT

This formula calculates heat required to raise a certain mass of a substance by a change in temperature ΔT.

where:
q = heat required in J
m = the sample's mass in g
c = the substance's specific heat in J/g*K
ΔT = temperature change in K

Water's specific heat is c = 1 cal/gºK.

In SI units, c ≈ 4.2 J/gºK.

q = mHL

This formula calculates heat required to change a certain mass of a substance from one phase to another. HL would be Hf for the first plateau and Hv for the second plateau.

where:
q = heat required in J
m = mass in g
HL = latent heat of phase change

100 cal.

At 20 ºC, water is a liquid and has specific heat c = 1 cal/gºC.
Since temperature is changing and phase is not changing within the temperature range 20-25, use the formula q = mcΔT.
hence: q = 20(1)(5) = 100cal.

STP is 0º C and 1 atm

Chemistry exams occasionally refer to Standard Ambient Temperature and Pressure (SATP), which you should know is 25ºC and 1 atm.

The kinetic molecular theory assumes an idealized version of gas, which makes calculating relationships easier. There are four assumptions:

1. A gas molecule has no volume (point molecule).

2. The collisions between gas molecules are completely elastic.

3. There are no dissipative forces due to collisions.

4. The average kinetic energy of gas molecules is directly proportional to the temperature of the gas.

Charles' Law states that the volume of a gas is directly proportional to temperature while at a constant pressure.

Equation:

The system's volume will double also.

Charles's Law indicates that at a constant pressure, the temperature and volume of a gas are directly proportional.

Boyle's Law states that the volume of a gas is inversely proportional to its pressure while at a constant temperature.

Equation:

The system's volume will double.

Boyle's Law indicates that at a constant temperature, the pressure and volume of a gas are inversely proportional.

Avogadro's Law states that the volume of a gas is directly proportional to the number of moles at a constant temperature and pressure.

V / n = k

where:
V = volume in L
n = number of moles
k = proportionality constant of the specific gas

The system's volume will decrease to 1/3 also.

Avogadro's Law indicates that at a constant pressure and temperature, the number of moles and volume of a gas are directly proportional.

The Ideal Gas Law combines Charles', Boyle's, and Avogadro's Laws into one:

PV = nRT

where:
P = Pressure in atm or Pa
V = Volume in L
n = Number of moles
R = Ideal gas constant .082 L(atm)/mol(K) or 8.31 J/mol(K)
T = Temperature in K

At STP, one mole of gas has a volume of 22.4 L

This is called the standard molar volume, and is true for a mole of any ideal gas at STP.

This value can be calculated using the Ideal Gas Law, but is worth memorizing.

PA = xAPtotal

where:
PA = pressure from gas A
xA = mole fraction of A (moles of A divided by total moles of gas)
Ptotal = total pressure of the sytem from all moles of gas

Ptotal = PA + PB + PC + …

Dalton's Law states that the total pressure of a system of ideal gases can be thought of as the sum of all of the partial pressures that each gas exerts.

U = KEavg = (3/2)nRT

where:
U = potential internal energy in J
KEavg = average kinetic energy in J
n = number of moles of gas
R = Ideal gas constant .082 L(atm)/mol(K) or 8.31 J/mol(K)
T = Temperature in K

Gases deviate from ideal at:

1. Low temperatures

2. High pressures ( or very low volume)

Gases deviate from ideal at low temperatures, because the molecules have very low kinetic energy (hence low velocity) and will start to attract based on the intermolecular forces.

Gases deviate from ideal at high pressures or very low volumes, because the molecules have very little actual space to move in and will start to exhibit characteristics more like a fluid than a gas.

(P + a(n/V)2)(V - nb) = nRT

where:
P = Pressure in atm or Pa
V = Volume in L
n = number of moles
R = Ideal gas constant .082 L(atm)/mol(K) or 8.31 J/mol(K)
T = Temperature in K
a = attraction between molecules due to intermolecular forces
b = actual volume taken up by gas molecules

Note: the van der Waals equation predicts the behavior of non-ideal gases, taking into account the intermolecular attractive forces of the gas molecules and the space taken up by the non-point molecules.

Polar gases will deviate more from ideal.

(P + a(n/V)2)(V - nb) = nRT

Remember: 'a' is the attractiveness between molecules. When 'a' is big (polar=attractive) the pressure term will be larger. Effectively, the gas will experience more pressure holding it together than the ideal gas law predicts.

Larger molecule gases will deviate more from ideal.

(P + a(n/V)2)(V - nb) = nRT

Remember: 'b' is the bigness of the actual molecules. When 'b' is big (larger size=big) the volume term will be smaller. Effectively, the gas will have less space to move in than the ideal gas law predicts.

The absolute temperature is any temperature that is given in Kelvin (K).

0 K is the temperature at which all molecules are assumed to cease moving completely. To convert Kelvin to Celcius:

K = C + 273.15