Gas Laws Unit (part 1): Chapter 10• Gas Behavior Basics• Kinetic-Molecular Theory• Temperature and Pressure relationships• Gas Laws

– Boyle’s Law,– Charle’s Law , – Gay – Lussac’s Law – Combined Gas Law

Learning Objectives…

• To describe the characteristics and behaviors that all gases have in common

• To use Kinetic Molecular Theory (KMT) to describe how the behavior of gas particles explains the physical properties of gases

Gas Basics…• Gases are all around us but we can’t

always see them• A gas is usually not a single element or

compound but a mixture of them

Gas Basics…

• Air is a mixture of gases that acts like a single gas• 78% Nitrogen• 21% Oxygen• 1% Other gases

Gas Basics…

• Most gases are made of molecules– Can be diatomic or polyatomic– N2, O2, CO2, CH4, etc.

• Some gases exist as single atoms– He, Ne, Ar, Kr, Xe, Rn

• We use the term “gas particles” to refer to an individual atom OR a molecule of gas.

F F

How Do Gases Behave?• Regardless of their chemical identity,

gases have very similar behavior– 1 mole of any gas

at 0o C & 1 atm occupies a volume of 22.4 L

– STP = 0o C & 1 atm • Kinetic Molecular

Theory (KMT)explains the behavior of gases

Temperature time…

• Get up out of your seats and find a partner from some other area of the room and discuss the warm-up question for the day…

• The next slide will help your discussion and or calculations…

Temperature conversions…

• Kelvin temperature from Celsius: 0° C + 273 = kelvin temp

• Temperatures in Celsius and Fahrenheit[(° F) – 32] x (5/9) = ° C(oC x 9/5) + 32 = o F

Kinetic-Molecular Theory…

• 1. Gases consist of small particles that have mass

• Evidence of mass– a flat b-ball weighs less than one with air in it

Kinetic-Molecular Theory…

• 2. The particles in a gas are separated from each other by large distances– The volume of gas

particles themselves is considered to be zero b/c they are so small compared to the volume of their containers

Kinetic-Molecular Theory…• Evidence of low density

– Gases are easy to compress. Moderate squeezing (pressure) on a gas will decrease its volume.• Compressing a gas changes Its density…

Kinetic-Molecular Theory…• 3. The particles of a gas are in constant,

rapid, random motion– They travel in straight lines, changing

direction only when they collide with each other or the edges/walls of the container

– KE dependsonly on temp• Hotter = faster

Kinetic-Molecular Theory…

• Evidence of constant random motion– Gases fill their containers completely

• when a balloon fills with air, the air is distributed evenly throughout the balloon

Kinetic Molecular Theory…– Gases “diffuse” and “effuse”

• Diffusion = when gas molecules move through each other from areas of higher to lower concentration

– This is why you can smell food cooking in another area of a house

• Effusion = when gas molecules move through a hole so tiny that they have to pass through one particle at a time

– This is why balloons deflate slowly even when tied off

• The lighter the gas,(or smaller the atom) the faster it will diffuse/effuse (Graham’s Law)

Kinetic-Molecular Theory …

Which picture shows diffusion? Effusion?

(a)

(b)

Kinetic Molecular Theory…

Do these show diffusion? Effusion? Both?

(a) (b)

Kinetic-Molecular Theory …

• A balloon filled with which gas will deflate the quickest and why?– O2: 32 g/mol– N2: 28 g/mol– H2: 2 g/mol– He: 4 g/mol

Kinetic-Molecular Theory…

• 4. Gases exert pressurebecause they constantlycollide with each otherand with the walls of thecontainer in which they are held.

Pressure…

• Get up out of your seats and find a partner from some other area of the room and discuss any units of pressure (air pressure) that you may be aware of or familiar with…

Conversions for Pressure Units…

One (1) atmosphere of pressure equals…1 atmosphere = 101.325 kPa (kilopascals)

1 atm. = 760 mm Hg 1 atm. = 14.7 lbs./in2 (or psi)

1 atm. = 760 torr 1 torr = 1 mm Hg

101,325 Pascals = 101.325 kPa

Kinetic-Molecular Theory …

– Every time a gas particle collides with a container wall it exerts an outward push or force on the wall

– The outward force spread over the area of the container is called “pressure”

Kinetic-Molecular Theory…

–Evidence of pressurea. The pressure a gas exerts on the inner walls

of a balloon gives the balloon its shape.b. Pressure increases

when more air is addedto a basketball–More gas particles

more collisions with the walls greater force increased pressure

Kinetic-Molecular Theory…– Assuming a constant volume and temp., the total

pressure of a mixture of gases is the sum of the pressures of the component gases. • Dalton’s Law of Partial Pressures

Kinetic-Molecular Theory

– The pressure of a gas depends on its temp• Higher temp = higher pressure (& vice-versa)

• Higher temp = faster moving particles = collisions are more frequent and at higher speeds = more forceful = increased pressure.

Low Temp High TempMed Temp

Each of these balloons has

the same amount of gas

inside!

Kinetic-Molecular Theory…

• 5. When gas particles collide, they do not sloooow dooooown (slow down)– Collisions are “elastic.” In a perfectly elastic

collision, no kinetic energy is lost

Kinetic-Molecular Theory…

• 6. Gas particles exert no force on one another.– Gas particles don’t slow down and condense

into a liquid because the attractive forces between them are very weak

Real vs. Ideal Gases

• Most of the time, gases behave exactly as the KMT predicts they will

• When they do, we call them “Ideal Gases”• However, KMT makes some faulty

assumptions…– 1) gases have no attraction for each

other (no Intra-molecular forces)– 2) gases have no volume

Real vs. Ideal Gases

• So under certain conditions, real gases do not behave like KMT predicts they will– At very low temps, gas particles slow down

enough that the IMFs they supposedly don’t have between them become significant

– At very high pressures, the volume they supposedly don’t have is significant and compression becomes very difficult

Real vs Ideal Gases

• Even though ideal gases do not exist, “real gases” behave like ideal gases under most conditions so we can still use the gas laws to describe them