Properties of Solids

Characteristics

1. Definite shape and volume

   • Particles in fixed positions
   • Strong intermolecular forces hold structure

2. High density

   • Particles closely packed
   • Little empty space

3. Incompressible

   • Cannot be compressed significantly
   • Particles already at minimum distance

4. Very low compressibility

   • Particles vibrate in place
   • Cannot flow or move past each other

Types of Solids

1. Crystalline Solids

• Ordered, repeating 3D structure
• Sharp melting point
• Examples: NaCl, diamond, ice

Types of crystalline solids:

a) Ionic Solids

• Ions held by electrostatic forces
• Hard, brittle, high melting points
• Conduct when molten/dissolved
• Example: NaCl, CaCO₃

b) Molecular Solids

• Molecules held by IMFs (LDF, dipole-dipole, H-bonding)
• Soft, low melting points
• Do not conduct electricity
• Example: Ice (H₂O), dry ice (CO₂), I₂

c) Covalent Network Solids

• Atoms covalently bonded throughout
• Very hard, very high melting points
• Do not conduct (except graphite)
• Example: Diamond, SiO₂ (quartz), SiC

d) Metallic Solids

• Metal cations in sea of electrons
• Malleable, ductile, lustrous
• Conduct electricity and heat
• Variable melting points
• Example: Na, Fe, Cu, Au

2. Amorphous Solids

• No long-range order
• Broad melting range
• Examples: Glass, rubber, plastics

Phase Diagram Features (Solids)

Triple point: Temperature and pressure where all 3 phases coexist

Normal melting point: Temperature at 1 atm where solid ↔ liquid

Properties of Liquids

Characteristics

1. Definite volume, no definite shape

   • Takes shape of container
   • Particles mobile but close together

2. High density (similar to solids)

   • Particles still close
   • Slightly more space than solids

3. Incompressible

   • Little empty space between particles

4. Can flow

   • Particles slide past each other
   • Viscosity depends on IMF strength

Key Liquid Properties

1. Viscosity

Definition: Resistance to flow

Factors affecting viscosity:

• Stronger IMFs → Higher viscosity
• Larger molecules → Higher viscosity
• Higher temperature → Lower viscosity (more KE overcomes IMFs)

Examples:

• Water: Low viscosity (H-bonding, but small molecule)
• Glycerol: High viscosity (3 OH groups, extensive H-bonding)
• Motor oil: High viscosity (large molecules, strong LDF)

2. Surface Tension

Definition: Energy required to increase surface area of liquid

Cause: Molecules at surface experience net inward force

• Bulk molecules: attracted equally in all directions
• Surface molecules: only attracted downward and sideways
• Creates "skin" on liquid surface

Factors:

• Stronger IMFs → Higher surface tension
• Water has high surface tension (H-bonding)
• Allows some insects to walk on water

3. Capillary Action

Definition: Spontaneous rise of liquid in narrow tube

Two competing forces:

Adhesion: Attraction between liquid and tube surface

• Liquid "sticks" to walls

Cohesion: Attraction between liquid molecules (IMFs)

• Liquid sticks to itself

Result:

• If adhesion > cohesion: Liquid rises (water in glass)
• If cohesion > adhesion: Liquid depressed (mercury in glass)

Meniscus:

• Water: Concave (curves up at edges)
• Mercury: Convex (curves down at edges)

4. Vapor Pressure

Definition: Pressure exerted by vapor in equilibrium with liquid

Dynamic equilibrium:

$\text{Liquid} \rightleftharpoons \text{Vapor}$

Rate of evaporation = Rate of condensation

Factors affecting vapor pressure:

a) Temperature:

• Higher T → Higher vapor pressure
• More molecules have enough KE to escape

b) IMF Strength:

• Stronger IMFs → Lower vapor pressure
• Harder for molecules to escape liquid
• Example: Water (H-bonding) has lower VP than ethanol at same T

Volatile liquids: High vapor pressure (weak IMFs)

• Examples: Acetone, diethyl ether

Nonvolatile liquids: Low vapor pressure (strong IMFs)

• Examples: Glycerol, motor oil

Normal Boiling Point: Temperature where vapor pressure = 1 atm (760 mmHg)

Properties of Gases

Characteristics

1. No definite shape or volume

   • Fills entire container
   • Expands to occupy all available space

2. Very low density

   • Particles far apart
   • Mostly empty space (~99.9%)

3. Highly compressible

   • Can squeeze particles closer
   • Volume decreases with pressure

4. Complete mixing

   • Particles move randomly
   • Rapid diffusion

Kinetic Molecular Theory (KMT)

Five postulates:

1. Gas particles are in constant, random motion

   • Move in straight lines until collision

2. Negligible particle volume

   • Volume of particles ≈ 0 compared to container
   • Mostly empty space

3. No intermolecular forces

   • Particles don't attract or repel (ideal gas assumption)

4. Elastic collisions

   • No energy lost in collisions
   • Total KE conserved

5. Average kinetic energy ∝ Temperature

   • Higher T → Higher average KE
   • $\overline{KE} = \frac{3}{2}RT$

Gas Pressure

Cause: Collisions of gas particles with container walls

Factors affecting pressure:

1. Temperature (↑T → ↑P at constant V, n)

   • Higher KE → more forceful, frequent collisions

2. Volume (↓V → ↑P at constant T, n)

   • Smaller space → more collisions per area

3. Number of moles (↑n → ↑P at constant T, V)

   • More particles → more collisions

4. Molar mass (indirect effect)

   • At same T, all gases have same average KE
   • Lighter gases move faster, heavier gases move slower

Phase Changes

Types of Phase Changes

Endothermic (absorb energy):

1. Melting (fusion): Solid → Liquid
2. Vaporization: Liquid → Gas
3. Sublimation: Solid → Gas (direct)

Exothermic (release energy): 4. Freezing: Liquid → Solid 5. Condensation: Gas → Liquid 6. Deposition: Gas → Solid (direct)

Energy of Phase Changes

Enthalpy of fusion (ΔHfus): Energy to melt 1 mole

• Breaking IMFs in solid structure
• Example: H₂O: ΔHfus = 6.01 kJ/mol

Enthalpy of vaporization (ΔHvap): Energy to vaporize 1 mole

• Breaking all remaining IMFs
• Example: H₂O: ΔHvap = 40.7 kJ/mol

Key relationship:

$\Delta H_{vap} > \Delta H_{fus}$

Why: More IMFs to break going liquid → gas than solid → liquid

Heating Curve

Temperature vs. Heat Added graph:

Features:

1. Solid heating: Temperature increases

   • Slope depends on specific heat of solid

2. Melting plateau: Temperature constant at melting point

   • Energy goes to breaking IMFs, not increasing KE
   • Length depends on ΔHfus

3. Liquid heating: Temperature increases

   • Slope depends on specific heat of liquid

4. Boiling plateau: Temperature constant at boiling point

   • Energy goes to breaking IMFs
   • Length depends on ΔHvap (longer than melting)

5. Gas heating: Temperature increases

   • Slope depends on specific heat of gas

Calculation:

$q = n \cdot \Delta H_{phase}$ (during phase change)

$q = m \cdot c \cdot \Delta T$ (within one phase)

Clausius-Clapeyron Equation

Relates vapor pressure to temperature:

$\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{vap}}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)$

Where:

• $P_1$, $P_2$ = vapor pressures
• $T_1$, $T_2$ = temperatures (K)
• $\Delta H_{vap}$ = enthalpy of vaporization
• $R$ = 8.314 J/(mol·K)

Linear form:

$\ln P = -\frac{\Delta H_{vap}}{R} \cdot \frac{1}{T} + C$

Plot: ln P vs. 1/T gives straight line with slope = $-\frac{\Delta H_{vap}}{R}$

Water's Unique Properties

Due to hydrogen bonding:

1. High boiling point (100°C)

   • Much higher than other hydrides (H₂S: -60°C)

2. High specific heat

   • Moderates temperature changes
   • Important for climate, organisms

3. High heat of vaporization

   • Cooling effect of sweating

4. Ice less dense than water

   • Ice floats
   • H-bonding creates open crystal structure
   • Unique among common substances

5. High surface tension

   • Capillary action in plants
   • Water transport in trees

Summary Table

Applications

1. Distillation: Separates based on different boiling points
2. Refrigeration: Uses phase changes to transfer heat
3. Weather: Water vapor pressure affects humidity, precipitation
4. Materials science: Solid structure determines material properties
5. Biological systems: Cell membranes, protein folding affected by IMFs

This means: Ice has larger volume than same mass of water

Step 2: Typical behavior of substances

Most substances:

• Solid is MORE dense than liquid
• Particles pack closer in solid (ordered structure)
• Solid sinks in its own liquid
• Example: Solid paraffin sinks in liquid paraffin

Step 3: Explain water's unusual behavior

In liquid water:

• H₂O molecules in constant motion
• H-bonding exists but constantly breaking/reforming
• Average of ~3.4 H-bonds per molecule
• Molecules relatively close but not in fixed positions
• Density at 4°C: 1.000 g/cm³ (maximum density)

In ice (solid water):

• H₂O molecules in fixed crystalline structure
• Each molecule forms 4 H-bonds (maximum)
• Tetrahedral arrangement around each molecule
• Creates hexagonal crystal structure with lots of empty space
• Open, cage-like structure
• Density at 0°C: 0.917 g/cm³

Step 4: Why the open structure?

Hydrogen bonding geometry:

• Each O has 2 H atoms (H-bond donors)
• Each O has 2 lone pairs (H-bond acceptors)
• Total: Can form 4 H-bonds per molecule

In ice:

• All 4 H-bonds form simultaneously
• Tetrahedral geometry (109.5° angles)
• This specific geometry requires more space
• Molecules held at fixed distances
• Creates channels and cavities in structure

In liquid water:

• H-bonds constantly breaking/reforming
• Not all 4 H-bonds present at once
• Molecules can pack more closely
• Less open space

Step 5: Quantitative comparison

Ice: 0.917 g/cm³ Water: 1.000 g/cm³

Ratio: Ice is about 8-9% less dense than water

Volume increase: When water freezes, it expands by ~9%

• This is why pipes burst in winter
• Why ice cubes expand in freezer

Step 6: Why is this unusual?

For most substances:

• Freezing brings particles closer together
• Solid is 10-20% MORE dense than liquid
• Increased order = tighter packing

For water:

• Freezing creates more H-bonds
• More H-bonds = more rigid, open structure
• Increased order = LESS dense

Only a few substances behave this way:

• Water (most important)
• Silicon
• Bismuth
• Antimony

Answer:

Ice floats because it is less dense than liquid water (0.917 g/cm³ vs. 1.000 g/cm³).

Explanation:

In ice, water molecules form a hexagonal crystalline structure where each H₂O molecule forms 4 hydrogen bonds in a tetrahedral arrangement. This creates an open, cage-like structure with large empty spaces between molecules.

In liquid water, H-bonds are constantly breaking and reforming, so molecules are not locked into this rigid structure and can pack more closely together, resulting in higher density.

Why unusual:

For most substances, the solid phase is MORE dense than the liquid because particles pack closer in the ordered solid structure. Water is unusual because the specific geometry required by hydrogen bonding (tetrahedral, 4 H-bonds per molecule) forces molecules APART, creating an open structure that is less dense than the liquid.

Importance:

1. Aquatic life: Ice forms on surface of lakes/oceans, insulating water below and allowing life to survive winter
2. Climate: Ice caps reflect sunlight (albedo effect)
3. Weathering: Water expanding when freezing can crack rocks
4. Plumbing: Pipes can burst when water freezes and expands

Without this property, bodies of water would freeze from bottom up, potentially killing aquatic ecosystems.