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Honors Chemistry Review Guide: States of Matter & Gas Laws

1. States of Matter Characteristics

• Solid:

  • Definite shape and volume.

  • Particles are tightly packed and vibrate in place.

  • Strong intermolecular forces.

• Liquid:

  • Definite volume, but no definite shape (takes the shape of the container).

  • Particles are close together but can flow past each other.

  • Weaker intermolecular forces than solids.

• Gas:

  • No definite shape or volume.

  • Particles are far apart and move freely.

  • Very weak intermolecular forces.

• Plasma:

  • Ionized gas (atoms have lost electrons).

  • High energy and temperature.

  • Found in stars, lightning, and some laboratory conditions.

2. Gas Laws

• Boyle’s Law (Pressure-Volume Relationship):

  P1V1=P2V2P_1 V_1 = P_2 V_2

  • For a fixed amount of gas at constant temperature, pressure and volume are inversely related.

• Charles’s Law (Temperature-Volume Relationship):

  V1T1=V2T2\frac{V_1}{T_1} = \frac{V_2}{T_2}

  • For a fixed amount of gas at constant pressure, volume and temperature are directly proportional.

• Gay-Lussac’s Law (Pressure-Temperature Relationship):

  P1T1=P2T2\frac{P_1}{T_1} = \frac{P_2}{T_2}

  • For a fixed amount of gas at constant volume, pressure and temperature are directly proportional.

• Ideal Gas Law:

  PV=nRTPV = nRT

  • Describes the behavior of ideal gases.

  • PP = Pressure, VV = Volume, nn = Moles, RR = Ideal gas constant, TT = Temperature in Kelvin.

• Real Gases:

  • Deviate from ideal gas behavior at high pressure and low temperature due to intermolecular forces and the volume occupied by gas molecules.

  • Real gases are more likely to condense into liquids at low temperatures or high pressures.

• Finding Molar Mass Using the Ideal Gas Law:

  M=dRTPM = \frac{dRT}{P}

  Where:

  • MM = Molar mass, dd = Density, RR = Ideal gas constant, TT = Temperature, PP = Pressure.

3. Key Vocabulary

• Absolute Zero:

  • The lowest possible temperature, 0 K (−273.15°C), where all molecular motion stops.

• STP (Standard Temperature and Pressure):

  • Defined as 0°C (273.15 K) and 1 atm pressure.

  • At STP, 1 mole of an ideal gas occupies 22.4 L.

• Boiling Point:

  • The temperature at which the vapor pressure of a liquid equals the external pressure, causing the liquid to change to a gas.

  • Increases with increasing atmospheric pressure.

• Vapor Pressure:

  • The pressure exerted by a vapor in equilibrium with its liquid or solid phase.

  • Increases with temperature.

• Phase Diagram:

  • A graphical representation of the physical states of a substance under different conditions of temperature and pressure.

  • Shows regions for solid, liquid, and gas states, and the lines where phase transitions occur.

4. Atmospheric Pressure

• Sea Level:

  • Standard atmospheric pressure is 1 atm, which is equivalent to 101.3 kPa or 760 mm Hg.

• At High Altitudes:

  • Atmospheric pressure decreases due to lower amounts of air above.

• At Low Altitudes:

  • Atmospheric pressure increases.

5. Kinetic Theory of Gases

• Assumptions:

  • Gases are composed of tiny particles that are far apart relative to their size.

  • Gas particles are in constant, random motion.

  • Collisions between gas particles and the walls of the container are elastic (no energy loss).

  • There are no intermolecular forces in an ideal gas.

  • The average kinetic energy of gas particles is proportional to the temperature of the gas.

• Kinetic Energy Formula:

  KE=32kTKE = \frac{3}{2} k T

  Where:

  • kk = Boltzmann constant, TT = Temperature in Kelvin.

6. Specific Heat Formula

• Specific Heat (C):

  • The amount of heat energy required to raise the temperature of 1 gram of a substance by 1°C (or 1 K).

  q=mcΔTq = mc\Delta T Where:

  • qq = Heat absorbed or released (Joules),

  • mm = Mass of the substance (grams),

  • cc = Specific heat capacity (J/g°C or J/g K),

  • ΔT\Delta T = Change in temperature (°C or K).

7. Phase Changes

• Types of Phase Changes:

  • Melting: Solid → Liquid (Endothermic).

  • Freezing: Liquid → Solid (Exothermic).

  • Vaporization (Boiling/Evaporation): Liquid → Gas (Endothermic).

  • Condensation: Gas → Liquid (Exothermic).

  • Sublimation: Solid → Gas (Endothermic).

  • Deposition: Gas → Solid (Exothermic).

• Heating Curve:

  • A graph showing the temperature change of a substance as heat is added.

  • Flat regions represent phase changes (melting or boiling).

• Vapor Pressure Curve:

  • Shows the relationship between temperature and vapor pressure.

  • As temperature increases, vapor pressure increases.

8. Phase Diagrams

• Triple Point: The temperature and pressure at which solid, liquid, and gas coexist in equilibrium.

• Critical Point: The point where the liquid and gas phases become indistinguishable.