Unit 7: States and Properties of Matter

Unit 7 Study Guide Answer Key

Concept 1: States of Matter

Vocabulary

• Thermal energy: The sum of the kinetic and potential energy of particles in an object.

• Kinetic energy: Energy in motion.

• Potential energy: Stored energy.

• Elastic collision: When no kinetic energy (KE) is lost overall in the collision of particles.

• Temperature: Measure of the average kinetic energy of particles in an object.

• Diffusion: Spontaneous mixing of particles caused by random motion.

• Effusion: The process where gas particles pass through tiny openings.

• Surface tension: The force that pulls adjacent parts of a liquid’s surface together, minimizing the surface area.

Objectives

1. Explain the kinetic-molecular theory.

   • All matter is made of small particles.

   • Those particles are in constant random motion.

   • This motion causes the particles to collide with each other and the container they are in.

2. Explain the relationship between temperature, kinetic energy, and the velocity of particles in a substance.

   • Temperature = the average kinetic energy of particles in an object.

   • Kinetic energy formula: $KE = rac{1}{2}mv^2$.

   • If you increase temperature, you increase the kinetic energy of particles; thus if mass is constant, the velocity of the particles increases.

3. Differentiate between the three main states of matter:

   State	Volume	Shape	Particle Movement
   Solid	Definite	Definite	Vibrate
   Liquid	Definite	Indefinite	Slide/Flow
   Gas	Indefinite	Indefinite	Freely move and collide

4. Differentiate between an intermolecular force and a chemical bond.

   • Chemical bonds: Form within molecules between atoms.

   • Intermolecular forces: Attractions between different molecules that make up a substance.

5. Explain the relationship between intermolecular forces and each state of matter.

   • Ideal gases: Have no intermolecular forces at work, thus no attractive or repulsive forces, allowing particles to move freely.

   • Liquids: Have stronger intermolecular forces at work that hold the particles more closely together while still allowing mobility.

   • Solids: Have the strongest intermolecular forces that hold the particles in rigid and organized formations.

6. Distinguish between ideal and real gases.

   • Ideal gas: A hypothetical gas that perfectly follows all of the kinetic-molecular theory with no attractive forces between particles.

   • Real gas: A gas that doesn’t behave entirely according to the kinetic-molecular theory due to exhibiting attractive forces.

7. Describe the properties of ideal gases.

   • Expand to fill the container they are in.

   • Fluid: Particles flow past each other.

   • Low density: Particles are incredibly far apart.

   • Compressibility: Can force particles close together by decreasing volume.

   • Readily diffuse and effuse.

8. Describe the properties of liquids, highlighting those that are unique to them.

   • Fluid: Particles flow past each other.

   • Relatively high density.

   • Incompressibility: Can’t really compress them more than they already are.

   • Can diffuse, but more slowly.

   • Surface tension: Unique to liquids; a force that gives the surface of a liquid a minimal surface area.

   • Capillary action: Unique to liquids; attraction of the surface of a liquid to a solid allows it to move against the pull of gravity.

9. Explain how both gases and liquids can be described as fluids.

   • Fluids are substances that flow. Both gases and liquids have weak enough intermolecular forces that the particles are able to flow past each other.

10. Differentiate between crystalline and amorphous solids.

    • Crystalline solids: Made of crystals, with particles arranged in an orderly geometric pattern.

    • Amorphous solids: Particles are still organized but in a much more random arrangement.

11. Differentiate between liquids and amorphous solids.

    • Liquids have weaker intermolecular forces than amorphous solids and no definite shape. Amorphous solids can exhibit liquid properties (like fluidity) but have stronger intermolecular forces at work to hold their definite shape.

12. Be able to identify a state of matter based on a description of its properties.

Concept 2: Phase Changes

Vocabulary

• Heat of fusion: The amount of energy, as heat, needed to turn 1 mole of solid into a liquid at its melting point.

• Heat of vaporization: The amount of energy, as heat, needed to turn 1 mole of liquid into a gas at its boiling point.

• Melting point: Temperature at which a solid becomes a liquid due to the kinetic energy of particles overcoming the attractive intermolecular forces that hold their order.

• Freezing point: The temperature at which a liquid turns into a crystalline solid.

• Volatile liquid: Liquids that readily evaporate due to weak attractive forces between particles.

• Vapor pressure: The temperature at which the vapor pressure of the liquid is equal to atmospheric pressure.

• Boiling point: A way to describe quantities in chemical reactions.

• Triple point: Indicates the temperature and pressure conditions necessary for a solid, liquid, and gas of a substance to coexist at equilibrium.

Objectives

1. Explain whether or not a phase change is physical or chemical.

   • It is physical because the identity of the substance doesn’t change, just the energy of the particles within.

2. Relate phase changes to temperature and the kinetic-molecular theory.

   • As temperature increases, the kinetic energy of particles increases, eventually reaching a point in each substance where there is enough energy for the particles to overcome the attractive intermolecular forces that hold them together and thus they enter a new phase/state of matter.

3. Differentiate between melting and freezing.

   • In melting, enough energy in the form of heat is added to change from solid → liquid.

   • In freezing, enough energy in the form of heat is removed to change from liquid → solid.

4. Differentiate between vaporization and condensation.

   • In vaporization, enough energy in the form of heat is added to change from liquid → gas.

   • In condensation, enough energy in the form of heat is removed to change from gas → liquid.

5. Differentiate between sublimation and deposition and give an example of each.

   • In sublimation, a solid changes directly into a gas (e.g., dry ice to carbon dioxide gas).

   • In deposition, a gas changes directly into a solid (e.g., frost formation).

6. Explain what is happening to particles during evaporation vs. boiling.

   • Evaporating occurs only at the surface of the liquid when particles escape the liquid state and become gas due to a pressure change.

   • Boiling occurs throughout the liquid and is the result of a temperature change that causes liquid particles to turn into bubbles of vapor.

7. Draw a diagram to summarize the different phase changes.

8. Explain what it means for particles in a substance in a closed container to be in equilibrium.

   • In a liquid, particles will be both vaporizing and condensing at equal rates, moving between liquid and gas at the same rate due to the pressure.

9. Be able to interpret a heating curve, cooling curve, or phase diagram.

Concept 3: Gas Laws

Vocabulary

• Pressure: Force per unit area on a surface.

• Atmospheric pressure: Pressure exerted by the atmosphere (shell) that surrounds Earth.

• STP: The atmospheric pressure at sea level (1 atm) and the freezing point of water at sea level (273 K).

• Partial pressure: Pressure of each gas in a mixture.

• Molar volume: The volume that 1 mole of a gas occupies at STP.

Objectives

1. Summarize the relationship between variables in each of the following gas laws:

   • Dalton’s Law: The total pressure of a gas mixture is the sum of the partial pressures of gases that make it up.

   • Boyle’s Law: Volume is inversely proportional to pressure; as one goes up, the other goes down.

   • Charles’s Law: Volume is directly proportional to temperature; as one goes up, the other goes up.

   • Gay-Lussac’s Law: Pressure is directly proportional to temperature; as one goes up, the other goes up.

   • Combined Gas Law: Puts pressure, volume, and temperature in relationship to each other when the amount of gas is constant.

     • Relationship: $P_1V_1 = P_2V_2$.

   • Avogadro’s Law: Equal volumes of gases contain equal numbers of molecules; as gas volume goes up, the amount of gas goes up.

   • Ideal Gas Law: $PV = nRT$.

2. Explain why each of the following gas laws makes sense according to the kinetic-molecular theory:

   • Dalton’s Law: Pressure is additive because the collisions of gas particles in constant motion are additive.

   • Boyle’s Law: If temperature is unchanged, then particles do not have more or less kinetic energy when volume or pressure are changed. As volume decreases, pressure increases because the particles have less space in which to collide; thus, the frequency of collisions with the walls increases, and so does the pressure within those walls.

   • Charles’s Law: Increasing temperature increases the kinetic energy of the particles, resulting in more collisions with the container, which would increase pressure. Since pressure is constant, the volume must increase to compensate for the increase in kinetic energy.

   • Gay-Lussac’s Law: Increasing temperature increases the kinetic energy of the particles, resulting in more collisions with the container, which also increases the pressure.

3. Be able to interpret models and analyze data that represent the relationship between temperature, pressure, and volume of a gas.

4. Be able to make predictions about changes in variables if other variables are held constant and defend those predictions with the gas laws.

5. Be able to use the mathematical equations that represent each gas law to solve for unknown variables.

6. Use the standard molar volume of a gas to calculate mass, molecules, or volume.

Practice Problems

7. Find the volume of a gas at 2.00 atm and 215 K if its original volume is 30.0 L at 0.250 atm and 408 K.

   • Given:

   • $V_1 = 30.0 ext{ L}$

   • $P_1 = 0.250 ext{ atm}$

   • $T_1 = 408 ext{ K}$

   • $P_2 = 2.00 ext{ atm}$

   • $T_2 = 215 ext{ K}$

   • Using combined gas law:
     $rac{P_1V_1}{T_1} = rac{P_2V_2}{T_2}$
     Solving gives:
     $V_2 = rac{(P_1V_1)(T_2)}{(P_2T_1)}$
     $V_2 = rac{(0.250)(30.0)(215)}{(2.00)(408)}$
     $V_2 = 1.98 ext{ L}$

8. If you have 0.275 moles of gas in a container at -198 ℃ with a pressure of 1.75 atm, what is the volume of the gas?

   • Given:

   • $n = 0.275 ext{ mol}$

   • $T = -198 ext{ °C} = 75 ext{ K}$

   • $P = 1.75 ext{ atm}$

   • Using Ideal Gas Law:
     $PV = nRT$
     Where $R = 0.0821 rac{L imes atm}{mol imes K}$
     Solving gives:
     $V = rac{(nRT)}{P} = rac{(0.275)(0.0821)(75)}{1.75} \ <br>ewline = 0.966 ext{ L}$