Conceptual Chemistry Unit 6 – States of Matter

Objective 1 Describe, at the molecular level, the difference between a gas, liquid, and solid phase.

Solids Definite shape Definite volume Particles are vibrating and packed close together. The particles do not flow.

Crystalline Solids Particles are arranged in an organized pattern. Example: Diamond

Amorphous Solids Particles are not organized in an orderly fashion. Example: Glass

Liquids Indefinite shape Definite volume Liquids will take the shape of a container, but they maintain the same volume. Particles are touching and packed close together. The higher energy allows the particles to move around each other.

Viscosity A liquid’s resistance to flow.

Gases Indefinite shape Indefinite volume Gases take the shape of a container. They also occupy the volume of the container no matter how big or small it is. High energy motion

Plasma High energy matter A common example is the sun. Super high energy gas particles that lost electrons. Plasma is the most common form of matter in the Universe.

States of Matter Property Solid (s) Liquid (l) Gas (g) Particle Spacing Close Great Energy Low Medium High Motion Shape Definite Indefinite Volume

Objective 2 Describe states of matter using the kinetic molecular theory.

Kinetic Molecular Theory can explain the behavior of matter in its different states. Kinetic Molecular Theory: Explains the states of matter based on the concept that the particles in all forms of matter are in constant motion. Kinetic Energy: Energy an object has due to its motion.

Kinetic Energy and Kelvin Temperature Temperature: the average kinetic energy of the particles in a material As particles are heated, they absorb energy, thus increasing their average kinetic energy and their temperature. Motion stops at absolute zero (0 Kelvin). Kelvin temperature scale reflects the relationship between temperature and average kinetic energy. It is directly proportional.

Objective 3 Describe changes in states of matter with respect to kinetic energy and temperature.

Energy and Phase Changes During a phase change, all energy goes to motion until phase change is done. The temperature does not change until the phase change is done.

Melting Solid  Liquid Example 1 Example 2

Freezing Liquid  Solid Example 1

Evaporation/Boiling Liquid  Gas Example 1

Condensation Gas  Liquid Example

Sublimation Solid  Gas Example Opposite of Sublimation? Deposition

Objective 4 Describe the different variables that define a gas.

Kinetic Theory of Gases Gases are mostly empty space. The molecules in a gas are separate, very small, and very far apart.

Kinetic Theory of Gases Gas molecules are in constant, chaotic motion. Collisions between gas molecules are elastic (there is no energy gain or loss).

Kinetic Theory of Gases The average kinetic energy of gas molecules is directly proportional to the absolute temperature. Gas pressure is caused by collisions of molecules with the walls of the container.

Gases doing all of these things! Behavior of Gases Gases have weight. Gases take up space. Gases exert pressure. Gases fill their containers. Gases doing all of these things!

Variables that Describe a Gas Volume: measured in L, mL, cm3 (1 mL = 1 cm3) Amount: measured in moles (mol), grams (g) Temperature: measured in Kelvin (K) K = ºC + 273 Pressure: measured in mm Hg, torr, atm, etc.  P = F / A (force per unit area)

P = F /A Moderate Force (about 100 lbs) Small Area (0.0625 in2) Enormous Pressure (1600 psi)

Large Surface Area (lots of nails) Bed of Nails Moderate Force Small Pressure P = F / A Large Surface Area (lots of nails)

Units of Pressure 1 atm = 760 mm Hg 1 atm = 760 torr 1 atm = 1.013 x 105 Pa 1 atm = 101.3 kPa

Boyle’s Law As P, V and vice versa…. Inverse relationship For a given number of molecules of gas at a constant temperature, the volume of the gas varies inversely with the pressure. As P, V and vice versa…. Inverse relationship P1V1 = P2V2

Boyle’s Law and Kinetic Molecular Theory How does kinetic molecular theory explain Boyle’s Law? Gas molecules are in constant, random motion. Gas pressure is the result of molecules colliding with the walls of the container. As the volume of a container becomes smaller, the collisions over a particular area of container wall increase…the gas pressure increases!

Pressure-Volume Calculations Example: Consider the syringe. Initially, the gas occupies a volume of 8 mL and exerts a pressure of 1 atm. What would the pressure of the gas become if its volume were increased to 10 mL?

Equation for Boyle’s Law P1V1 = P2V2 where: P1 = initial pressure V1 = initial volume P2 = final pressure V2 = final volume

P1V1 = P2V2 Using the same syringe example, just “plug in” the values: (1 atm) (8 mL) = (P2) (10 mL)

P1V1 = P2V2 (1 atm) (8 mL) = P2 (10 mL) P2 = 0.8 atm

P1V1 = P2V2 (1.2 atm)(12 L) = (3.6 atm)V2 V2 = 4.0 L Example: A sample of gas occupies 12 L under a pressure of 1.2 atm. What would its volume be if the pressure were increased to 3.6 atm? (assume temp is constant) P1V1 = P2V2 (1.2 atm)(12 L) = (3.6 atm)V2 V2 = 4.0 L

P1V1 = P2V2 (200 kPa)(28 L) = (P2)(17 L) P2 = 329 kPa Example: A sample of gas occupies 28 L under a pressure of 200 kPa. If the volume is decreased to 17 L, what be the new pressure? (assume temp is constant) P1V1 = P2V2 (200 kPa)(28 L) = (P2)(17 L) P2 = 329 kPa

Temperature – Volume Relationships What happens to matter when it is heated? It EXPANDS. What happens to matter when it is cooled? It CONTRACTS. Gas samples expand and shrink to a much greater extent than either solids or liquids.

Charles’ Law As T , V  and vice versa…. Direct relationship The volume of a given number of molecules is directly proportional to the Kelvin temperature. As T , V  and vice versa…. Direct relationship Video Clip 1, Clip 2

Temperature – Volume Relationship Doubling the Kelvin temperature of a gas doubles its volume. Reducing the Kelvin temperature by one half causes the gas volume to decrease by one half… WHY KELVIN? The Kelvin scale never reaches “zero” or has negative values.

Converting Kelvin To convert from Celsius to Kelvin: add 273. Example: What is 110 ºC in Kelvin? 110 ºC + 273 = 383 K

Converting Kelvin To convert from Kelvin to Celsius: subtract 273. Example: 555 K in Celsius? 555 K - 273 = 282 ºC

T2 = 746 K V1 = 117 mL; T1 = 100 + 273 = 373 K V2 = 234 mL; T2 = ??? Example: A sample of nitrogen gas occupies 117 mL at 100.°C. At what temperature would it occupy 234 mL if the pressure does not change? V1 = 117 mL; T1 = 100 + 273 = 373 K V2 = 234 mL; T2 = ??? V1 / T1= V2 / T2 T2 = 746 K

Example: A sample of oxygen gas occupies 65 mL at 28. 8°C Example: A sample of oxygen gas occupies 65 mL at 28.8°C. If the temperature is raised to 72.2°C, what will the new volume of the gas? V1 = 65 mL; T1 = 28.8 + 273 = 301.8 K V2 = ??? mL; T2 = 72.2 + 273 = 345.2 K V1 / T1= V2 / T2 V2 = 74.3 mL

Temperature – Pressure Relationships Picture a closed, rigid container of gas (such as a scuba tank) – the volume is CONSTANT. What would happen to the kinetic energy of the gas molecules in the container if you were to heat it up? How would this affect pressure? States of Matter Interactive Egg in a Bottle:! Video Clip

Temperature – Pressure Relationships Raising the Kelvin temperature of the gas will cause an INCREASE in the gas pressure. WHY? With increasing temperature, the K.E. of the gas particles increases – they move faster! They collide more often and with more energy with the walls of the container.