Properties of States of Matter

Properties of States of Matter

• Entropy: The tendency towards disorder. Solids have low entropy, liquids intermediate, and gases high entropy.
• Solids: Particles are closely packed, have high density, are incompressible, and have particles in fixed positions. Diffusion is very low.
• Liquids: Particles are close together but can move past each other. Have a moderate density and are relatively incompressible. Diffusion is slower than in gases, but faster than in solids.
• Gases: Particles are far apart, have low density, and are easily compressible. Particles are in constant, random motion, resulting in high diffusion rates.

Types of Crystalline Solids

• Ionic crystals: Held together by ionic bonds, hard, high melting points, and do not conduct electricity.
• Covalent network crystals: Covalent bonds, exceptionally hard and high melting points, do not conduct electricity. Examples include diamond, silicon carbide, quartz.
• Metallic crystals: Metallic bonds, hard, high melting points, and conduct electricity.
• Covalent molecular crystals: Covalent bonds, soft, low melting points, and are poor conductors.

Liquids

• Volume/Shape: Takes the shape of the container it is in but does not completely fill it.
• Attractive Forces: Liquids have stronger attractive forces than gases but weaker than solids. Particles can move past each other, giving them fluidity.
• Viscosity: Resistance to flow, higher viscosity indicates greater resistance.
• Surface Tension: Uneven forces at the surface of a liquid
• Capillary Action: The movement of liquids against gravity up a thin tube due to intermolecular forces between the liquid and the tube material.

Gases

• Volume/Shape: Takes the shape and volume of its container.
• Forces of Attraction: Weak attractive forces allow gases to expand and move freely.
• Diffusion: Gases diffuse relatively quickly due to the constant, random motion of their particles.
• Viscosity: Gases have very low viscosity, flowing readily.

Changes of State

• Phase Equilibrium: The rate of one phase change equals the rate of the opposing phase change, with no change in the amount of either phase.
• Vapor Pressure: The pressure exerted by a vapor in equilibrium with its liquid. Vapor pressure increases with increasing temperature.

Phase Diagrams

• Triple Point: The temperature and pressure at which all three phases (solid, liquid, and gas) coexist in equilibrium.
• Critical Point: The combination of temperature and pressure above which a liquid-gas transition cannot occur.

Kinetic Molecular Theory (KMT)

• Assumptions: Gases are composed of tiny particles in constant motion, individual gas particles possess no volume, there is no force of attraction between the particles.

Gas Laws

• Boyle's Law: Pressure and volume are inversely related at a constant temperature.
• Charles' Law: Volume and temperature are directly related at a constant pressure.
• Gay-Lussac's Law: Pressure and temperature are directly related at a constant volume.
• Avogadro's Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
• Combined Gas Law: Combines Boyle's, Charles' and Gay-Lussac's Laws.
• Ideal Gas Law: Relates pressure, volume, moles, and temperature for gases, with the gas constant R.
• Dalton's Law of Partial Pressures: The total pressure of a mixture of gases is the sum of the partial pressures of the individual gases.
• Graham's Law of Effusion: The rate of effusion of a gas is inversely proportional to the square root of its molar mass.

Density of Gases

• Density can be calculated using the Ideal Gas Law.

Molar Mass

• Molar mass is the mass of one mole of a substance (g/mol).

Gas Stoichiometry

• Stoichiometric relationships are used for reactions involving gases based off of mole ratios.