Unit 6 Gases, Phase Changes and Introduction to Thermochemistry

Unit 6Gases, Phase Changes and Introduction to Thermochemistry Characteristics of Gases Pressure Kinetic-Molecular Theory The Gas Laws Partial Pressures Effusion and Diffusion Real Gases Part I: Gases

Properties of Gases • Three phases of matter • solid • liquid • gas Definite shape and volume Definite volume, shape of container Shape and volume of container

Properties of Gases • A gas is a collection of molecules that are very far apart on average. • In air, gas molecules occupy only 0.1% of the total volume. • In liquids, molecules occupy ~ 70% of the total space.

Properties of Gases • Gases are highly compressible. • Volume decreases when pressure is applied. • Gases form homogeneous mixtures with each other regardless of the identities or relative proportions of the different gases. • Water and gasoline = heterogeneous mixture. • Water vapor and gasoline vapor = homogeneous mixture.

Properties of Gases • Properties of gases vary depending on their composition. • Air: ~ 78% N2 and ~ 21% O2 • CO2: colorless, odorless • CO: colorless, odorless, highly toxic • NO2: toxic, red-brown, irritant • N2O: colorless, sweet odor (laughing gas)

Pressure • Four quantities are commonly needed to describe a gas: • amount of gas (n) • Temperature (T) • Volume (V) • Pressure (P)

Pressure • Gases exert pressure on the objects in their surroundings. • Pressure is caused by collisions between the gas molecules and objects with which they are in contact. • Pressure:the force exerted on a unit area P = F A

Pressure • Atmospheric pressure:the pressure exerted by gas molecules in the air on all objects exposed to the atmosphere • Atmospheric pressure varies with altitude.

Pressure Why does atmospheric pressure decrease with increasing altitude? • Gravity decreases • Density of gas decreases • Fewer gas molecules • Fewer collisions • Lower pressure

Pressure • Many different units used to report pressure. • millimeters of Hg (mm Hg) • inches of Hg (in. Hg) • pounds per square inch (psi) • atmosphere (atm) • torr (torr) • pascal (Pa) = SI base unit • kilopascal (kPa) Must know units and abbreviations!!

Pressure • Relationships between different pressure units: 1 atm = 760 mm Hg = 760 torr = 29.92 in. Hg = 14.7 psi = 1.01325 x 105 Pa Must be able to interconvert between units. Memorize the ones in red…I’ll give you the others. You must know that 1 kPa = 1000 Pa

Pressure Example: The measured pressure inside the eye of a hurricane was 669 torr. What was the pressure in atm?

Pressure Example: On a nice sunny day in Chicago the barometric pressure was 30.45 in. Hg. What was the pressure in Pa?

Pressure Example: On Titan, the largest moon of Saturn, the atmospheric pressure is 1.631 Pa. What is the pressure in atm?

Kinetic Molecular Theory • The behavior of gases can be described and explained usingkinetic molecular theory. • the “theory of moving molecules” • You must know the basic ideas that are part of kinetic molecular theory.

Kinetic Molecular Theory • Gases consist of large numbers of molecules that are in continuous, random motion. • The combined volume of all the molecules of the gas is negligible compared to the total volume in which the gas is contained. • i.e. the molecules are very far apart on average

Kinetic Molecular Theory • Attractive and repulsive forces between gas molecules are negligible. • Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change as long as the temperature remains constant. • Collisions are perfectly elastic.

Kinetic Molecular Theory • The average kinetic energy of the molecules is proportional to the absolute temperature. • At any given temperature all molecules of a gas have the same average kinetic energy. • As T (in K) increases, KE increases.

Gas Laws • Four variables are needed to define the physical condition or state of any gas: • Temperature (T) • Pressure (P) • Volume (V) • Amount of gas (moles: n) • Equations relating these variables are known as the gas laws.

P Gas Laws Consider a fixed amount of gas that is confined to a container with a certain volume. At a specific temperature, the gas sample will exert a certain pressure on the container.

P P Volume decreases Gas Laws What will happen to the pressure if the volume is decreased?

Gas Laws • As the volume of a fixed quantity of gas decreases, the pressure increases because: • gas molecules are more tightly packed together • i.e. denser • more collisions between gas molecules and the container • greater pressure

Gas Laws • Boyle’s Law: • The volume of a fixed quantity of gas maintained at constant temperature is inversely proportional to the pressure. • Mathematically, V = k x 1 or PV = k or P1V1 = P2V2 P at constant temperature and quantity of gas

Gas Laws • As liquid nitrogen (-196oC) is poured over a balloon, the volume of the balloon decreases.

Gas Laws • Charles’ Law: • The volume of afixed amount of gas maintained at constant pressureis directly proportional to itsabsolute temperature. V = k x T orV = k or V1 = V2 T T1 T2 At constant pressure and quantity of gas Remember:T must be in Kelvin

Gas Laws • On a molecular level, as the temperature of a gas maintained at constant pressure decreases, • KE decreases • fewer collisions between gas molecules and the environment (i.e. container) • volume decreases in order to maintain constant pressure

Gas Laws What happens when you “blow up” a balloon?

Gas Laws What happens when you “blow up” a balloon? • the number of moles of gas (n) increases • and • the volume of the gas (balloon) increases

Gas Laws • Avogadro’s Law: • The volume of a gas maintained at constant temperature and constant pressure is directly proportional to the number of moles of the gas. • Mathematically, V = constant x n At constant temperature and pressure

Gas Laws • At any given temperature and pressure, as the amount of gas increases, • the number of gas molecules increases • the number of collisions between gas molecules and the environment (container) increases • the volume must increase in order to maintain constant pressure

Gas Laws • In a chemical reaction, we use the coefficients to tell us how many moles or molecules are used or produced in a chemical reaction. N2 (g) + 3 H2 (g)  2 NH3 (g) • 1 mole of nitrogen reacts with 3 moles of hydrogen to produce 2 moles of ammonia

Gas Laws • Since the volume of a gas is directly proportional to the number of moles of gas at constant temperature and pressure, we can also use the coefficients to represent the volume of a gas involved in a reaction. (Avogadro’s Hypothesis) N2 (g) + 3 H2 (g)  2 NH3 (g) • 1 liter of nitrogen reacts with 3 liters of hydrogen to produce 2 liters of ammonia

Gas Laws • Boyle’s Law, Charles’ Law, and Avogadro’s Law can be combined to make a more general gas law: • Ideal Gas Law: PV = nRT where P = pressure V = volume n = moles T = temperature (K) R = gas constant

Gas Laws • The value of the gas constant (R) depends on the units of P, V, n, and T. • T must always be in Kelvin • n is usually in moles • If P (atm) and V (L), • then R = 0.08206 atm.L mol.K • If P (torr) and V (L), • then R = 62.36 L.torr mol.K I will give you these on the test.

Unit 6 Gases, Phase Changes and Introduction to Thermochemistry