Gases, Liquids, and Solids: States of Matter and Their Properties

Chapter 8: Gases, Liquids, and Solids

8.1 States of Matter and Their Changes

The three primary states of matter—solid, liquid, and gas—are distinguished by the arrangement and energy of their particles. Changes between these states are called phase changes and involve energy transfer as heat.

• Change of State: The transformation of a substance from one state (solid, liquid, gas) to another.

• Enthalpy (ΔH): The heat absorbed or released during a phase change.

• Entropy (ΔS): The change in molecular disorder or freedom during a process.

• Melting (Fusion): Solid to liquid; heat is absorbed (endothermic, ΔH > 0), and disorder increases (ΔS > 0).

• Freezing: Liquid to solid; heat is released (exothermic, ΔH < 0), and disorder decreases (ΔS < 0).

• Equilibrium Temperature: The temperature at which the enthalpy and entropy changes balance, resulting in equilibrium between phases (e.g., melting point, boiling point).

Example: Ice melting at 0°C absorbs heat and increases molecular disorder.

8.2 Intermolecular Forces

Intermolecular forces are attractions between molecules or atoms that influence physical properties such as boiling and melting points. They are negligible in gases but significant in liquids and solids.

• London Dispersion Forces: Present in all molecules due to temporary shifts in electron density. Strength increases with molecular weight and surface area.

• Dipole-Dipole Forces: Occur between polar molecules; positive and negative ends attract each other.

• Hydrogen Bonding: A strong dipole-dipole interaction between hydrogen bonded to O, N, or F and another O, N, or F atom.

Example: Water (H2O) forms four hydrogen bonds per molecule, resulting in high boiling point.

Type of Force	Occurs Between	Relative Strength
London Dispersion	All molecules	Weak
Dipole-Dipole	Polar molecules	Moderate
Hydrogen Bonding	H with O, N, or F	Strong

8.3 Gases and the Kinetic-Molecular Theory

The kinetic-molecular theory explains the behavior of gases based on particle motion and energy.

• Gas particles move randomly and independently, with negligible attractive forces.

• The volume of gas particles is much smaller than the space between them.

• Average kinetic energy is proportional to temperature (in kelvins).

• Collisions are elastic; total kinetic energy remains constant.

• An ideal gas perfectly follows these assumptions.

8.4 Pressure

Pressure (P) is the force per unit area exerted on a surface. It is commonly measured in atmospheres (atm), millimeters of mercury (mmHg or torr), pascals (Pa), or pounds per square inch (psi).

• Atmospheric Pressure: The pressure exerted by the weight of air at sea level (1 atm = 760 mmHg = 101,325 Pa = 14.7 psi).

• Barometer: Measures atmospheric pressure using a column of mercury.

• Manometer: Measures gas pressure in a container by comparing mercury levels.

Example: To convert 55 psi to atm: $55 \text{ psi} \times \frac{1 \text{ atm}}{14.7 \text{ psi}}$

8.5 Boyle’s Law: The Relation between Volume and Pressure

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

• Mathematical form: $P_1 V_1 = P_2 V_2$

• If pressure increases, volume decreases, and vice versa.

Example: Compressing a gas from 1.0 atm to 9.5 atm reduces its volume by a factor of 9.5.

8.6 Charles’s Law: The Relation between Volume and Temperature

Charles’s Law states that the volume of a gas is directly proportional to its temperature (in kelvins) at constant pressure and amount of gas.

• Mathematical form: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

• If temperature doubles, volume doubles.

Example: Air inhaled at room temperature expands when warmed to body temperature.

8.7 Gay-Lussac’s Law: The Relation between Pressure and Temperature

Gay-Lussac’s Law states that the pressure of a gas is directly proportional to its temperature (in kelvins) at constant volume and amount of gas.

• Mathematical form: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$

• If temperature doubles, pressure doubles.

Example: Heating an aerosol can increases the internal pressure.

8.8 The Combined Gas Law

The combined gas law merges Boyle’s, Charles’s, and Gay-Lussac’s laws for a fixed amount of gas.

• Mathematical form: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

• Allows calculation of changes in pressure, volume, or temperature.

Example: Determining the temperature needed to maintain a certain pressure and volume for a gas sample.

8.9 Avogadro’s Law: The Relation between Volume and Molar Amount

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

• Mathematical form: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$

• Standard Temperature and Pressure (STP): 0°C (273.15 K) and 1 atm.

• Standard molar volume at STP: 22.4 L per mole for an ideal gas.

Example: Calculating the number of moles of air in a room using its volume and the molar volume at STP.

8.10 The Ideal Gas Law

The ideal gas law combines the relationships among pressure, volume, temperature, and amount of gas into a single equation.

• Mathematical form: $PV = nRT$

• R (gas constant): $0.0821 \text{ L·atm·mol}^{-1}\text{·K}^{-1}$ or $8.314 \text{ J·mol}^{-1}\text{·K}^{-1}$

Example: Calculating the number of moles of air in the lungs or the pressure in a gas cylinder.

8.11 Partial Pressure and Dalton’s Law

Partial pressure is the pressure a single gas in a mixture would exert if it occupied the entire volume alone. Dalton’s Law states that the total pressure of a gas mixture is the sum of the partial pressures of each component.

• Mathematical form: $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

• Partial pressure can be calculated as: $P_{\text{gas}} = \text{(mole fraction)} \times P_{\text{total}}$

Example: Calculating the partial pressures of oxygen, nitrogen, water vapor, and argon in humid air.

8.12 Liquids

Liquids have molecules in constant motion, with intermolecular forces holding them together. Some molecules escape into the gas phase, creating vapor pressure.

• Vapor Pressure: The pressure exerted by vapor in equilibrium with its liquid.

• Vapor pressure increases with temperature and depends on intermolecular forces.

• Boiling Point: The temperature at which vapor pressure equals atmospheric pressure.

• Viscosity: Resistance to flow; increases with stronger intermolecular forces.

• Surface Tension: Caused by differences in intermolecular forces at the surface versus the interior of a liquid.

8.13 Solids

Solids are classified based on the arrangement and bonding of their particles.

• Crystalline Solids: Ordered, repeating structures.

  • Ionic Solids: Composed of ions (e.g., NaCl).

  • Molecular Solids: Composed of molecules held by intermolecular forces (e.g., ice).

  • Covalent Network Solids: Atoms linked by covalent bonds in a giant network (e.g., diamond).

  • Metallic Solids: Metal cations in a sea of electrons; conduct electricity and are malleable.

• Amorphous Solids: Disordered, non-crystalline (e.g., glass, tar).

Type	Constituent Particles	Forces	Examples
Ionic	Ions	Ionic bonds	NaCl
Molecular	Molecules	Intermolecular forces	Ice
Covalent Network	Atoms	Covalent bonds	Diamond
Metallic	Metal cations	Metallic bonds	Iron
Amorphous	Atoms/molecules	Random arrangement	Glass

8.14 Changes of State Calculations

Calculating the heat required for phase changes involves the heat of fusion (melting) and heat of vaporization (boiling).

• Heat of Fusion (ΔHfus): Heat required to melt 1 g (or 1 mol) of a substance at its melting point.

• Heat of Vaporization (ΔHvap): Heat required to vaporize 1 g (or 1 mol) of a substance at its boiling point.

• When a substance is at its phase change temperature, added heat changes the phase, not the temperature.

• Strong intermolecular forces result in higher heats of fusion and vaporization.

Example: To calculate the heat required to melt 0.300 mol of naphthalene, multiply the mass (in grams) by the heat of fusion.

Substance	Melting Point (°C)	Boiling Point (°C)	ΔHfus (kJ/mol)	ΔHvap (kJ/mol)
Water	0	100	6.01	40.7
Naphthalene	80	218	19.0	45.0
Additional info: ...	...	...	...	...

Summary Concept Map

This chapter connects the properties and behaviors of gases, liquids, and solids to the underlying molecular structure and intermolecular forces. Understanding these concepts is essential for predicting and explaining physical changes and properties in chemical systems.