VI. Phases of matter: Solid, Liquid, and Gas
Key focus of this chapter: phase change for water
This chapter focuses on the phase change for water and gives concise summaries of the important things about the phase diagram of H2O and CO2, heating curve of H2O, colligative properties, Raoult’s law, osmotic pressure, and Hery’s law in more detail.
A. Solid, liquid, and gas
Classification | Features |
Solid | • Particles are rigid structures due to very strong attraction. • Regularly positioned particles – crystalline solid (ionic, metallic, covalent network, molecular solids). • Irregularly positioned particles – amorphous solid (any glass). |
Liquid | • Particles can flow and have strong attraction. • Cohesion: attraction force between liquid molecules - causing surface tension (resistance of increasing surface area) - causing viscosity (resistance of fluidity) • Adhesion: attraction force between liquid molecule and other substance |
Gas |  |
B. Phase change
: phase of solid, liquid, and gas is changed by temperature and pressure.
Fig. 1 Phase change
Phase changes | Features | Phase changes | Features | ||
Condensation | • Phase change from gas to liquid by bond formation | Decreasing enthalpy (ΔH) and entropy (ΔS) change | Vaporization | • Phase change from liquid to gas by breaking bond | Increasing enthalpy (ΔH) and entropy (ΔS) change |
Freezing | • Phase change from liquid to solid by bond formation | Fusion (melting) | • Phase change from solid to liquid by breaking bond | ||
Deposition | • Phase change from gas to solid by bond formation | Sublimation | • Phase change from solid to gas by breaking bond • Ex/ iodine, dry ice | ||
C. Phase diagram of H2O and CO2
Fig. 2 Phase diagram of H2O and CO2
 |  |
• Negative slope between solid and liquid boundary → decreasing melting point with higher pressure • Positive slope between liquid and gas boundary | • Positive slope between solid and liquid boundary → increasing melting point with higher pressure • Positive slope between liquid and gas boundary |
- Triple point
– coexisting point of solid, liquid, and gas in equilibrium
- Critical point
– indistinguishable line between liquid and gas
– impossible to be liquefied or vaporized by changing pressure or temperature
D. Heating curve of H2O
Fig. 3 Heating curve of H2O
 |  |  |  | |
Ice | 2.0 | 36.6 | ||
Ice → water | 6.0 | |||
Water | 4.2 | 75.4 | ||
Water → vapor | 40.7 |
Qs/ How much energy is required to convert 2mol of ice at -10oC to water at 50 oC?
Sol/Q1 = molar heat capacity of ice: 36.6 J/(mol•oC)
= 36.6 J/(mol•oC) × 2 mol × 10oC
= 732 J
Q2 = ΔHfusion (ice → water): 6.0 kJ/mol
= 6,000 J/mol × 2 mol
= 12,000 J
Q3 = molar heat capacity of water: 75.4 J/(mol•oC)
= 75.4 J/(mol•oC) × 2 mol × 50oC
= 7,540 J
Therefore, Q1 + Q2 + Q3 = 732 J + 12,000 J + 7,540 J
= 20,272 J
E. Colligative properties
: impure solution mixed as a solute in a pure solvent has a higher boiling and lower freezing point than the pure solvent.
1. Features of colligative properties
- Higher boiling point than the pure solvent → lower vapor pressure than the pure solvent
- Lower freezing point than the pure solvent
- Osmotic pressure is present
 |
**Freezing point of solution (mixing solute and pure solvent) is lower than freezing point of the pure solvent → melting point of impure solid is lower than the melting point of pure solid (see fig.4)
2. The changing of boiling and freezing point of colligative solution
Fig. 4 The graph of colligative properties for ΔTb, ΔTf, Psol
Classification | Features |
 |  |
 |  |
** van ’t Hoff factor (mole number of solute ions), i
- Ex/ CH3OH: non-electrolyte (CH3OH) → i = 1
LiCl: 2 moles of ions (Li+ + Cl–) → i = 2
CaCl2: 3 moles of ions (Ca+ + 2Cl–) → i = 3
Qs/ Calculate the freezing point of 21g of LiCl in 1000g of water. (Kf of water = 1.86oC•kg/mol)
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F. Raoult’s law
: Psoln (vapor pressure of a solution that is a nonvolatile solute mixed with a pure solvent) is equal to Psolv (vapor pressure of pure solvent) times Xsolv (mole fraction of solvent).
 |
** Induced Raoult’s law
 |
- The vapor pressure of the solution is lower than the vapor pressure of the pure solvent because the surface area of a vaporizable solvent mixed with a non-volatile solute is decreased.
Fig. 5 Vapor pressure of pure solvent and solution
Qs/ Calculate the vapor pressure of the solution when 4 moles of K2SO4 is dissolved in 12 moles of H2O at 45oC. (The vapor pressure of water at 45 oC = 70 mmHg.)
 |
G. Osmotic pressure (P)
: the pressure required to keep the same heights between the different heights of the solution and pure solvent.
Fig 6. Osmotic pressure between solvent and solution
** Osmosis: the movement of the solvent molecules from high solvent concentration area to low solvent concentration area through a porous membrane.
 |
Qs/ 116g of NaCl is dissolved in 4L of the solution at 27 oC. Calculate the osmotic pressure. M.W. of NaCl: 23+ 35 = 58g/mol
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H. Henry’s law
: the solubility of gas into liquid is proportional to the partial pressure on a liquid.
Fig. 7 Solubility of gas into liquid by the partial pressure
 |
Qs/ Calculate the solubility of an unknown gas in water when the partial pressure on the water is 3 atm at 23oC and the Henry’s-law constant of the unknown gas, k, is 5.0 × 10-2 mol/L•atm.
Sol/ Solubility of gas = k•P
= (5.0 × 10-2 mol/L•atm) • (3 atm)
= 0.15 mol/L