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Gases

Mar 20, 2016

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Gases. Exploring Gases. Make a table: Demo # Prediction Observation. Kinetic Theory (Observing Properties of Gases ). Gases are tiny particles, that have mass but a small volume (Volume assumed to = 0 ) Gases in constant random motion - PowerPoint PPT Presentation
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Page 2: Gases

Exploring GasesMake a table:Demo # Prediction Observation

Page 3: Gases

Kinetic Theory(Observing Properties of Gases)

•Gases are tiny particles, that have mass but a small volume ▫ (Volume assumed to = 0)

•Gases in constant random motion•Collisions are elastic with walls and each other (No attraction between molecules)

•Kinetic energy (movement) depends on temperature – (high temperature is more movement)

Page 4: Gases

Measuring Gases

•Amount (n)moles (number of)

•Volume (V) liters

•Temperature (T)Kelvin K= C + 273

•Pressure (P)mm Hg or atm

Page 5: Gases

Particles colliding with objects•Pressure=force (Pascal) areaGas pressure

gas particles colliding with objects

Atmospheric pressureair particles colliding with objects

Pressure

Page 6: Gases

Barometer

Measures atmospheric pressure

Page 7: Gases

Measuring pressure

Barometer760 mm Hg = 1 atmosphere = 101,300 Pascals

=14.7 lb/in2

STPStandard temperature 0oCStandard pressure 1 atm

Page 8: Gases

Pressure problemsIf 760 mm Hg = 1 atm

•Convert 793 mm Hg to atm•Convert 3.5 atm to mm Hg

Page 9: Gases

Manometer problemsMeasures pressure in a closed container

Page 10: Gases

Manometersgas

Pressure755 mmHg

Pressure755 mmHg

mercury

Page 11: Gases

Manometers855 mmHg

Pressure? mmHg

100 mm Hg

Page 12: Gases

Dalton’s Law of Partial PressuresEach of the components of a gas

mixture contributes some of the collisions

Each component contributes part of the total pressure.

Mathematically..PT = P1 + P2 + P3…..

Page 13: Gases

What is temperature?•Measure’s average kinetic energy of

particles.▫Higher temp means higher energy

More energy means faster particles▫Lower temp means lower energy

Less energy means slower particles

•When particles move faster, they collide more often and with more force.

Page 14: Gases

Temperature and Energy Revisited

•If there are a fixed number of gas particles in a container

•And it has a fixed pressure•What happens when it is heated up?

–The particles go faster–They collide more–The volume goes up

Page 15: Gases

Combined Gas lawCombines •Boyles, •Charles, •Gay-Lussac law

P1V1 = P2V2 T1 T2

Page 16: Gases

Boyles Law(P,V)

At a constant T,N, the volume varies indirectly with the pressure

P1V1 = P2V2

Page 17: Gases

Lab: Boyles Law

•Purpose:To observe changes in

pressure with the volume changes

Page 18: Gases

Charles Law(T,V)

•At a constant P, N the volume varies directly with the Kelvin temperature

V1 = V2T1 T2

Page 19: Gases

Gay-Lussacs Law(P,T)

At a constant V,N the pressure varies directly with the Kelvin temperature.

P1 = P2T1 T2

Page 20: Gases

Lab: Pressure/temperaturePurpose: To determine the absolute zero

using Gay-Lussac’s Law

Page 21: Gases

Absolute zero demonstrationTemperature at which all molecules stop moving

0 K -273 oC

Page 22: Gases

Avagadro’s Law•If you hold pressure and temperature

constant ▫Like at standard temperature and

pressure Which are?

•Volume and moles are related

V n

Page 23: Gases

Lab: combined gas lawPurpose: To determine the volume of 1 mole of a gas using the combined gas law

Reaction: Hydrochloric acid and Mg

Page 24: Gases

Ideal Gas LawAt STP 1 mole=22.4LIf not at STP use Ideal Gas Law

P V = n R TR=ideal gas constant

(0.0821 L atm ) moles K

Page 25: Gases

Ideal Gas Law•All the gas laws are related.•By the pressure, volume, temperature

and number of particles (moles or n)PV = constantV/T = constantn/V = constantP/T = constant

Page 26: Gases

Gas Proportions•There are four variables in the Gas Laws

▫Pressure▫Volume▫Temperature▫Moles

•We can intuit each gas law using KMT•For example:

▫If moles and temperature are constant▫How does the volume and pressure compare?

Page 27: Gases

Gas Stoichiometry•C8H18, octane, combusts in your car’s

engine. If the cylinder is 0.500 L and the oxygen intake is at 45oC and 1.05 atm, how many grams of octane are needed to completely react with the oxygen?

Page 28: Gases

Gas Collected Over Water

If the water level in the flask is equal to the surrounding water, than the inside pressure is equal to the outside pressure. Pin = PO2 + PH2O = P atmospheric

PH2O = 21 torr

Page 29: Gases

Pressure of Collected Gas•The vapor pressure of water @ 20.0 C is

17.54 mmHg•How many mmH2O is this?

▫What data do you need?▫Mercury d = 13.7 g./ml▫240. mmH2O

•If 100.0 ml of oxygen is collected over 20.0 C ▫If the atmospheric pressure is 739 mmHg,

what is the pressure of oxygen?▫How many moles of Oxygen gas?▫How many atoms of oxygen

Page 30: Gases

Pressure EqualizationWhich pressure is Higher?

Page 31: Gases

Pressure EqualizationWhich pressure is Higher?

Page 32: Gases

Motion of Gases•At the same temperature, two samples of

gas have the same average kinetic energy•What has more kinetic energy, a bus

moving at 5 mph or a baby on a tricycle moving at 10 mph?

• If mass is important let’s consider molecular motion.

Page 33: Gases

Diffusion•Gases at the same temp have the same

average KE•More massive gases must be moving

slower than less massive gases at the same temp

•If I let out a smelly gas of a flask, how does it get to your nose? What path does it take?

•When a gas spreads out or dissolves into the air, we call this diffusion

Page 34: Gases

Graham’s Law•The rate of diffusion is directly

proportional the speed of the molecule•The bigger the molecule, the ______ the

speed of the molecule (at the same temp)•The “bigger” really means molar mass.•We can compare rates or velocities

Page 35: Gases

Graham’s Law va = Mb

vb Ma

The ratio of the velocities of gas molecules is proportional to the Inverse square root of their molar masses

Page 36: Gases

Problem•The rate of diffusion of an unknown gas is

four times faster than the rate of oxygen gas. Calculate the molar mass of the unknown gas and identify it.

va = 1 = Mb

vb 4 32 g/molMb = 2 What gas has a molar mass of 2?

Page 37: Gases

Dalton’s law•The total pressure equals the sum of the

partial pressures of the gases in the container

PT = P1 + P2 + P3 + ……..

Page 38: Gases

Try thisAir contains oxygen, nitrogen, carbon

dioxide and other gases. What is the pressure due to oxygen in mm Hg if

PT= 1 atmPN=593.4 mm HgPCO2 = 46.78 mm Hg

Page 39: Gases

problems•The pressure on 2.50 L of anesthetic gas is changed from 765 mm Hg to 304 mm Hg. What is the new volume if the temperature is constant?

Page 40: Gases

problems

•A balloon inflated in air conditioning at 27oC has a volume of 4.0L. It is heated to 57oC. What is the new volume?

Page 41: Gases

Practice

A gas has a volume of 17.3 mL at 3.5 atm. What is the volume if the pressure is increased to 6.7 atm?

A can contains a gas at 50oC and has a volume of .5L. When released what is its new volume at 20oC?

Page 42: Gases

Try this

If 87.6 mL of hydrogen gas is collected at a room temperature of 23oc and room pressure of 742 mmHg, what will the volume be at STP?

Page 43: Gases

problems

•A gas has a volume of 6.8L at 327oC. What is its volume at 36oC?

Page 44: Gases

Molar Volume

1 Mole = 22.4 L

At STPStandard temp 0oC

Standard pressure 1 atm

Page 45: Gases

Pressure /Temperature•In a sealed container, with a fixed volume

and fixed number of particles•What happens to the pressure, if the

temperature of the system is increased? Why?

•The pressure and the temperature vary directly. Just like in Charles Law.

Page 46: Gases

•1. Why should the thistle tube be under the water level?

• 2. Why was the first bottle “let go”?• 3. Why were the bottles placed upside

down on the lab bench?• 4. What was this method called for

collecting gas • using a pneumatic trough and pushing

water out?• 5. Why did the splint go out inside the

bottle?• 6. What was the clear, colorless liquid

produced? • 7. Write a chemical reaction for its

production.

Page 47: Gases

Lab: Combined gas law

1.Take room temperature and pressure2. Get about 5 cm (or less) Mg and mass.

Tie onto copper wire3. Pour 15 ml of HCl into eudiometer and

fill to top with water4. Put Mg into top of eudiometer.

Stopper.5. Put finger over the hole, turn upside

down and place into big beaker of water.

6. When reaction is complete, put finger over hole and transfer to large graduated cylinder to measure volume of gas collected.

Page 48: Gases

Kinetic Theory Real vs IdealValid if not at Low temperatures

– attractive forces apply

high pressures – volume of particles - attractive forces apply

Page 49: Gases
Page 50: Gases

Questions1. What happens to the energy of the particles of gas

when you put the flask into cold water?2. Why do we use Kelvin when calculating gas law

problems? (Hint – is the Celsius temperature directly proportional to pressure below zero degrees?)

3. Predict the volume of the gas at 0 K from your data. You do this by getting an equation an plugging in the numbers.

4. Compare your answer with the real answer. Why are they different? What could affect your results?

5. A flask has a pressure of 1.0 atm at 25 C. What is the pressure at –40 C? (remember to convert to Kelvins)