Introduction to Gases & Atmospheric Chemistry

Gases – All around us…. - found all around us (literally) and very important to our daily lives - e.g. breathing, heating, medical uses, scuba diving, air bags - other examples?? - Ivory Soap – have you ever wondered?

Comparison of Solids, Liquids and Gases Solids – hold a definite shape Definite volume Particles are closely packed Can’t be compressed Very little movement of particles Strong attractive forces exist between particles Particles have low Ek (little motion, Ek is kinetic energy)

Liquids - indefinite shape (takes shape of container) Definite volume Particles easily move past each other Slightly compressible Fluid (liquids flow) Weak attractive forces exist between particles Particles have more Ek (significant motion)

Gases - indefinite shape (takes shape of container) Indefinite volume (particles spread out to fill container) Particles have large spaces between them Most of the volume of a gas is empty space Particles rarely touch each other Easily compressible Fluid (gases flow) Weakest attractive forces exist between particles Particles have high Ek (constant motion)

Kinetic Molecular Theory The properties and behavior of gases can be explained by the these five main points – 1. All matter is composed of tiny particles 2. The particles are in constant motion (  temperature,  E K ) 3. Forces of attraction and repulsion exist between the particles 4. Very large spaces exist between gas particles 5. Force of attraction between particles is very small 6. Particles have elastic collisions (energy is neither gained nor lost)

Gas Properties – Volume, V Volume – amount of space that a gas occupies measured in mL or L (or cm 3 ) 1000 mL = 1.0 L 1 mole of any gas occupies 22.4 L at STP

Gas Properties – Temperature, T The Celsius scale is determined based on water freezing at 0  C and boiling at 100  C. This is quite arbitrary. The Kelvin scale, however, is an absolute scale Comparing the two scales Kelvin = Degrees Celsius Degrees Celsius = Kelvin Convert the following: 150  C = K290K=  C 1250  C= K400K=  C

Absolute Zero This is believed to be the lowest temperature possible for a gas to achieve and is often called Absolute Zero. Hypothetically, all molecular motion stops at this temperature. Absolute Zero =  C

Gas Properties – Pressure, P Force - a push or a pull Pressure - force per unit area Pressure = example: Step on foot with running shoes Step on foot with high heels Force is the same, but the pressure is different

Pressure results from the collision of gas particles with the walls of the container in which the gas is contained SI units - Pascal, Pa (derived unit - a combination of other units) Other units – atm, mm Hg, psi, torr Gas pressure depends on: 1. Area: Increase Area, Decrease Pressure Decrease Area, Increase Pressure Area is Inversely related to Pressure

Measuring Pressure - Allow gas to exert a pressure on something that exerts a pressure back ie. Usually a column of Hg Barometer - any device used to measure Earth’s atmospheric pressure Force holds Hg in column or it would run out Earth’s atmosphere pushes on pool and supports column When two forces balance, level of Hg becomes constant  atm. pressure, Hg falls  atm. pressure, Hg rises

Pressure Conversions At sea level, 1 atm = kPa = 760 mmHg = 760 torr = in. Hg = 14.7 p.s.i

Convert the following: a) 3.25 x 105 Pa = kPa b) 97.8 kPa = psi c) 675 mmHg = atm d) 300 kPa = mmHg e) 798 mmHg = atm f) 13.2 psi = kPa g) 16.4 psi = mmHg h) 1.39 x 103 kPa = atm i) 670 Pa = kPa

Gas Properties – Number of Moles, n Number of moles = n One mole represents 6.02 x particles Why would one mole of any gas have the same volume at STP?

STP – Standard Temperature & Pressure As a reminder, STP is O  C and kPa What would STP be in K and atm?

Summary - Units and Conversions Volume: V - L or mL Temperature: T - °C or K Kelvin= Celsius amount of gas measured in moles: n Pressure: P - Pa or kPa, atm, mm Hg, torr STP kPa and 0 °C

Intro to the Gas Laws - file:///C:/Program%20Files/PhET/simulations/si ms8c14.html?sim=Gas_Properties

Boyle’s Law Discovered by Robert Boyle in 1662 pressure is inversely proportional to volume mathematical formula- P 1 V 1 =P 2 V 2 A real life application is a syringe. As volume increases, molecules must travel farther to impact into the walls of the container, as such the pressure is decreased.

Boyle’s Law Example 50 mL of oxygen is at atmospheric pressure. What is the pressure when there is a volume is compressed to 35 mL? V 1 = 50 mLV 1 P 1 = V 2 P 2 P 1 = 101 kPaP 2 = V 1 P 1 /V 2 V 2 = 35 mLP 2 = (50mL)(101kPa)/35 mL P 2 = ?P 2 = kPa

Boyle’s Law Demonstration Example: syringe As the water is drawn up into the syringe, the volume of the air inside is decreased. The air molecules have less space to travel, and therefore, collide into the walls of the container more often. The pressure increases because there are more molecule impacts per unit time. This pressure holds the water in the syringe.

Charles’ Law Discovered by Jacques Charles in 1787 volume is proportional to temperature mathematical formula: V 1 = V 2 T 1 T 2 An example would be a hot air balloon rising and falling. Hot air is pushed into the balloon. The air expands and fills the balloon. To come down, air from the balloon is released.

Charles’ Law Example A gas occupies 100 mL at 20 °C. At what temperature will it occupy 200 mL? V 1 = 100mLV 1 /T 1 = V 2 /T 2 T 1 = 20 °C = 293KT 2 = V 2 T 1 /V 1 V 2 = 200 mLT 2 = (200mL)(293K)/(100mL) T 2 = ?T 2 = 586 K

Charles’ Law Demonstration Example: Hot Water Balloon When the water bottle is dispersed into the hot water, the air molecules begin to expand, rise and move faster. So, the molecules are pushed up from the bottle and inflate the balloon.

Pressure- Temperature Law Also called the Gay-Lussac law Pressure is proportional to temperature mathematical formula: P 1 = P 2 T 1 T 2 An example would be how bike tires lose pressure during the winter. As air cools, it contracts. So a bike tire will be flat after the winter months because the air inside has decreased in temperature.

Pressure- Temperature Law Example Temperature is increased from 40 K to 200 K. If the final pressure was 80 kPa, what was the initial pressure? T 1 = 40 KP 1 /T 1 = P 2 /T 2 P 1 = ?P 1 = P 2 T 1 /T 2 T 2 = 200 KP 1 = (80kPa)(40K)/200K P 2 = 80 kPaP 1 = 16 kPa

Pressure- Temperature Law Demonstration Example: beads in a bottle The beads represent air molecules in a container. There are the same number of beads in each, however, the second bottle has a higher temperature. The bottles are shaken to represent the movement of the molecules. At higher temperatures, molecules move faster and therefore, impact the container walls more often. This demonstrates the increase in pressure.

Review

Common Mistakes Confusing the laws temperature problems can only be solved using the Kelvin scale difference between SATP and STP – SATP- 100 kPa and 25 °C – STP kPa and 0 °C

ANY QUESTIONS ?