Chapter 10. Gases

Chapter 10. Gases

Watch Bozeman Videos & other videos on my website for additional help: Big Idea 2: • Gases

10.1 Characteristics of Gases

Read p. 398-401.

Shape Volume Compressible Flow/Speed Type of mixtures formed

Density units

Solids

Liquids

Gas

Although relatively few substances exist as gases under typical conditions, they are

very important.

• Supports life, waste receptacle for exhaust gases, shields us from harmful

radiation, etc

Gases are made up of particles that have (relatively) large amounts of energy.

• They are highly compressible – when pressure is applied to a gas, its volume

decreases (Boyle’s Law).

• A gas has no definite shape or volume and will expand to fill as much space

as possible (entire container – homogeneously).

• As a result of the large amount of empty space in a volume of gas, gases are

easily compressed.

10.2 Pressure

Read p. 401-404.

All gases exert Pressure on their surroundings.

• Pressure – the force exerted by gas molecules as they strike the surfaces

around them.

The gases most familiar to us are the ones that make up our atmosphere (N2, O2, Ar, CO2, Ne, He, CH4, …)

• Together they exert atmospheric pressure on us, and on the earth.

Atmospheric pressure is measured with a barometer (Invented by Italian physicist Torriceli). • If a tube is completely filled with mercury and then inverted into a container

of mercury open to the atmosphere, the mercury will rise 760 mm up the tube.

2 factors that can change atmospheric pressure:

1. Altitude

• At higher elevations there is a smaller amount of “air” above you, so

the atmospheric pressure would be less (barometer reads < 760 mm

Hg).

• At lower elevations – greater atmospheric pressure.

2. Weather (humidity)

• In the tropics, there’s more H2O molecules in the air. So atmospheric

pressure is greater (Bad hair days)

• In dry regions and deserts, humidity is very low. Few H2O molecules

in the air leads to lower atmospheric pressure.

Important non SI units used to express gas pressure include:

atmospheres (atm)

torr (torriceli)

millimeters of mercury (mm Hg)

Pascal (Pa)

kilopascal (kPa)

pounds per square inch (psi)

1 atm = 760. mm Hg = 760. torr = 101, 325 Pa = 101.325 kPa = 14.7psi

Example #1: Perform the following conversions

1. 657 mmHg to torr

2. 830 torr to atmospheres

3. 7.8 kPa to atmospheres

4. 1200 Pa to mmHg

Barometer with Hg Pic (We will draw 3 different pics here)

Example #2:

The height of the mercury in the open ended arm is 136.4mm, and the height in the

arm in contact with the gas in the flask is 103.8mm. What is the pressure of the gas

in the flask (a) in atmospheres, (b) in torr. (Draw and Label the barometer)

10.3 The Gas Laws

Read p. 404-407.

The equations that express the relationships

among T (temperature), P (pressure),

V (volume), and n (number of moles of gas) are known

as the GAS LAWS.

Boyle’s Law: The Pressure-Volume Relationship

• What is constant: temperature and moles

• Relationship: inversely proportional

o This means that as the pressure increases, the volume decreases (and

vice versa).

o It’s helpful to think about the fact that pressure is caused by the

number of collisions of particles with the wall of the container.

o Balloon scenario – squeezing a balloon decreases its volume and

increases pressure.

▪ This will bring the gas molecules closer and the gas molecules

will collide MORE with each other (and walls of the balloon).

• Graph : plot of P vs.V is a hyperbola.

THINK about it: The working of the lungs illustrates Boyle’s Laws.

• As we breathe in, the diaphragm moves down, and the ribs expand;

therefore, the volume of the lungs increases.

• According to Boyle’s law, when the volume of the lungs increases, the

pressure decreases; therefore, the pressure inside the lungs is less than the

atmospheric pressure.

• Atmospheric pressure forces air into the lungs until the pressure once again

equals atmospheric pressure.

• As we breathe out, the diaphragm moves up and the ribs contract; therefore,

the volume of the lungs decreases, pressure increases, and air is forced out.

I’m Robert Boyle,

and I came up

with my law way

back in 1622.

P1 V1 → P2 V2

P1 and V1 are the original conditions. P2 and V2 are the new conditions.

• Units :

o The units for P1/P2 need to be the same. Any units can be used (atm,

mm Hg, etc).

o The units for V1/V2 need to be the same. Any units can be used (mL,

L, cm3, etc).

Example: Calculate the following

If a 1.23 L sample of a gas at 53.0 torr is put under pressure up to a value of 240.

torr at a constant temperature, what is the new volume?

Charles’s Law: The Temperature-Volume Relationship

• What is constant: pressure and moles

• Relationship: directly proportional

o This means that the volume of a gas increases with increasing

temperature (and vice versa).

o If the temperature is increased, the gas particles gain kinetic energy,

move around more and occupy more space.

o Balloon scenario – consider a hot air balloon. What happens when

heat (fire) is turned on? The gas molecules will move faster and

collide more causing the balloon to expand (volume increases).

Boyle was “Very Pretty”

I’m Jacques

Charles, and even

though my friend

Gay-Lussac

published the

official law, he was

so nice to name it

after me, since I

originally had the

idea in 1787!

• Graph : plot of V vs.T is straight (linear) line.

V1 T2 → V2 T1

T1 and V1 are the original conditions. T2 and V2 are the new conditions.

• Units :

o The units for Temperature must be in KELVIN! (no exceptions)

▪ Remember: K = oC + 273

▪ Absolute zero: 0 K = –273 oC.


L, cm3, etc).


A sample of gas at 15oC and 1 atm has a volume of 2.58L. What volume will this

gas occupy at 38oC?

Charles liked to watch”TV”

Gay-Lussac’s Law: Pressure and Temperature relationships

• What is constant: volume and moles


o This means that the pressure of a gas increases with increasing

temperature (and vice versa).

o If the temperature is increased, the gas particles gain kinetic energy

and move around more. The particles will have more energy and

collisions with the walls of the container (and with other gas

molecules) will be more forceful and the pressure will increase.

• Graph : plot of P vs.T is straight (linear) line.

P1 T2 → P2 T1

T1 and P1 are the original conditions. T2 and P2 are the new conditions.

• Units :





mm Hg, etc).


A gas at 25oC in a closed container has its pressured raised from 150. atm to 160.

atm. What is the final temperature of the gas?

Avogadro’s Law: The Moles-Volume Relationship

• What is constant: temperature and pressure


o This means that the volume of a gas increases with increasing number

of moles (and vice versa).

o You put more moles of gas in a balloon , the volume will expand.

o As more moles of a gas are placed into a container (or balloon) if

conditions of temperature and pressure are to remain the same, the gas

must occupy a larger volume.

o What’s interesting is that he had another component to his law –

equal volumes of gases at the same temperature and pressure

contain equal numbers of particles (moles), even if the gases are

different!!!

o What that means is that if you have 1 mole of one gas, such as CO2,

and also 1 mole of a different gas, such as N2, that both of those moles

will occupy the same volume, as long as their pressure and

temperature is the same.

• Graph : plot of V vs. n is straight (linear) line.

V1 n2 → V2 n1

V1 and n1 are the original conditions. V2 and n2 are the new conditions.

• Units :

o The units for “n” must be in MOLES! (no exceptions)

▪ So, if you are given grams, molecules, formula units, etc., then

you must convert to moles

▪ Remember:

molar mass (g) = 1mol = 6.02 x 1023 atoms/molecules/formula units


L, cm3, etc).

V1 n2 → V2 n1

V1 and n1 are the original condition. V2 and n2 are the new conditions.

Can also use Stoichiometry!!! (See 10.4 why we can)


A 13.1 L sample of 0.502 moles of O2 is held under conditions of 1.00 atm and 25.0 oC. If all of the O2 is then converted to Ozone (O3) what will be the volume of

ozone? Hint: Need a balanced equation first.

10.4 Combined Gas Law & The Ideal-Gas Equation

Read p. 408-412.

COMBINED Gas Law: all gas laws in one equation (Boyles, Charles, Gay-Lussac, & Avogadro’s Laws)

• You can combine the four gas laws into one equation: When you work your

math, it is best to start with the Combined Gas Law Equation. It includes

pressure, volume and temperature relationships and is really a combination

of all three-gas laws – even adding in Avogadro’s Law with “n”. All you do

is write down the combined equation and then cross off the variables that

remain constant (if any). If a variable is not mentioned in a problem

(frequently amount), you can assume that variable has remained constant.

• What is constant: any variable can be constant

1’s are the original conditions. 2’s are the new conditions.

22

22

1

1

1

1

Tn

VP

Tn

VP=

• Units :


mm Hg, etc).


L, cm3, etc).




o The units for “n” must be in MOLES! (no exceptions)

Example: Calculate the following Diborane gas (B2H6) has a pressure of 3.45torr at a temperature -15oC and a volume of 3.48L. If conditions change such that the temperature is 36oC and the pressure is 468torr, what will be the volume of the sample?

GAS LAWS OVERVIEW

LAW INDEP/DEP VARIABLES

CONTROL VARIABLES

MATH RELATIONSHIP

FORMULA

BOYLE’S

CHARLES’S

GAY LUSSAC’S

AVOGADRO’S

COMBINED

Remember:

An independent variable (x-axis) – a variable that stands alone and isn't changed by the other variables

you are trying to measure. For example, someone's age might be an independent variable. In fact, when

you are looking for some kind of relationship between variables you are trying to see if the independent

variable causes some kind of change in the other variables, or dependent variables.

A dependent variable (y-axis) - is something that depends on other factors. For example, a test score

could be a dependent variable because it could change depending on several factors such as how much

you studied, how much sleep you got the night before you took the test, or even how hungry you were

when you took it. Usually when you are looking for a relationship between two things you are trying to

find out what makes the dependent variable change the way it does.

Ideal gas equation: An ideal gas is a hypothetical gas whose P, V, and T behavior is completely described by the ideal-gas equation.

The way to identify that you have this type of problem (and not one of the previous types) is that CONDITIONS WILL NOT BE CHANGING!!! You will not have both initial and final values for the variables PV = nRT

• Units : *** Have to have certain units to plug into equation***(no

exceptions!)

o P = pressure (atm)

o V = volume (L)

o n = number of moles (mol)

o T = temperature (K)

o R = gas constant = 0.08206 𝐿 ° 𝑎𝑡𝑚

𝑚𝑜𝑙 ° 𝐾

▪ The units of “R” force us to have P, , N, and T in specific units

when doing the calculation. Let me show you why in the

examples below.

▪ Other numerical values of R in various units are given in Table

Ch 10.2. (Let’s write them down in case you want to use them)

P R value R units

atm 0.08206 𝐿 ⋅ 𝑎𝑡𝑚

𝑚𝑜𝑙 ⋅ 𝐾

Torr (or mmHg) 62.4 𝐿 ⋅ 𝑡𝑜𝑟𝑟

𝑚𝑜𝑙 ⋅ 𝐾

KPa 8.314 𝐿⋅𝐾𝑃𝑎

𝑚𝑜𝑙⋅𝐾 OR

𝐽

𝑚𝑜𝑙⋅𝐾

Let’s do some Algebra….Rearrange Ideal Gas Equation to solve for each variable: P = V = n = T =

Helpful Tip: PV = nRT problems don’t typically have a P1, V1, P2, T2, etc

situations. There’s usually just one situation. Initial values/final values type

of problems require using Boyles, Charles, Gay-Lussac, Avogadro’s, or

Combined Gas Law.

Examples: Calculate the following

1. Assuming ideal behavior, how many moles of Helium gas are in a sample

that has a volume of 8.12 L at a temperature of 0.00 oC and a pressure of

1.20 atm?

Gas Stoichiometry

• We can use Stoichiometry ☺ sometimes instead of using gas law equations,

but only under 1 condition!!!

• Suppose you have 1 mol of an ideal gas at 0oC (273K) and 1 atm pressure.

Then what volume does the gas occupy?

• Using PV=nRT….

• Thus is MOLAR VOLUME of an ideal gas at 0oC (273K) and 1 atm.

• The conditions of standard temperature (0°C) and standard pressure (1

atm) are called “STP conditions”.

• So… 1 mol of any gas behaving ideally at STP occupies a volume of

22.4L.

1 mol = 22.4L at STP This is Molar volume

• BEWARE: If conditions are not at STP, we cannot use this conversion.

We will learn what to do in these situations in another unit.

o If you have a liquid or solid, you cannot use this conversion!

Examples: Calculate the following

1. How many liters of oxygen gas at STP are required to react with 7.98 liters

of ethane gas at STP in a combustion reaction? (HINT: chemical

changes…need a balanced equation)

2. A process requires 42 L of ammonia. How many moles of nitrogen should

be used, assuming the reaction was performed under standard conditions?

(HINT: chemical changes…need a balanced equation)

3. 5.33 x 1023 molecules of hydrogen gas reacts with excess chlorine gas, how

many liters of hydrogen chloride gas could potentially be formed at STP? (HINT: chemical changes…need a balanced equation)

10.5 Further Applications of the Ideal-Gas Equation

Read p. 412-415.

Molar Mass and Gas Densities

• How can the ideal gas equation be used to solve for density and molar mass

of a gas?

• Remember: Density d = 𝒎

𝑽 (m = mass, V = volume)

• Gases use density units of grams/Liters.

• If you rearrange the ideal-gas equation with “M” as molar mass (g/mol) we

get….

(Beware not to confuse “M” molar mass with Molarity from Ch 4)

PV = nRT and V = 𝒎

𝒅

(from density equation… Substitute V density equation into ideal gas)

𝑷𝒎

𝒅= 𝒏𝑹𝑻

(rearrange variables to get the desired units of molar mass = g/mol)

𝑚

𝑛=

𝑑𝑅𝑇

𝑃

𝑴 = 𝒅𝑹𝑻

𝑷

Now rearrange the molar mass equation of a gas to solve for density (d):

𝒅 = 𝑴𝑷

𝑹𝑻

The “Meow-Meow” formula… cats put “DiRT” over their “Pee”!!!


Find the molar mass of a gas at 27 oC and 1.50 atm that has a density of 1.95 g/L.

10.6 Gas Mixtures and Partial Pressures

Read p. 415-418. Dalton observed:

• This law is used when you have a mixture (solution) of 2 or more gases

which do not react chemically.

• For a mixture of gases in a container, the total pressure exerted is the SUM

of the individual partial pressures. (as long as volume, temperature, and

moles are constant!!!)

o Partial pressure is the pressure exerted by a specific gas in gaseous

mixture.

• Dalton’s law of partial pressures: In a gas mixture the total pressure is

given by the sum of partial pressures of each component:

Pt = PA + PB + PC …for however many gases you have!

Pt = total pressure PA + PB + PC = different gases

o What is constant: Volume, temperature, and moles

o Units: any pressure units (as long as all pressures are the same units)

Since gas molecules are so far apart, we can assume that they behave ideally.

• And if each gas behaves ideally, then PV = nRT can also be used to find the

presuure of each gas. Once you solve for each of the gas pressures, then

ADD the pressures up to solve for the TOTAL pressure.

𝑃𝐀 = 𝑛𝐴 𝑅𝑇

𝑉 𝑃𝐵 =

𝑛𝐵 𝑅𝑇

𝑉 𝑃𝐶 =

𝑛𝐶 𝑅𝑇

𝑉 so Pt = PA + PB + PC…..

• OR if you know the moles of each gas in the mixture, you can add up the

total moles (nt)and solve for Pt using PV = nRT.

𝑃𝑡 = 𝑛𝑡 𝑅𝑇

𝑉

Mole Fractions: Using Dalton’s Partial Pressures

• Mole Fraction: the ratio of the # of moles of a given component in a gas

mixture to the total # of moles in the mixture.

𝑋 𝐴 = 𝑛 𝐴

𝑛 𝑡𝑜𝑡𝑎𝑙=

𝑛 𝐴

𝑛 𝐴 + 𝑛 𝐵 + 𝑛 𝐶 … .

Remember: A, B, C are different gases.

• If you wanted to solve for the pressure of only 1 gas in a mixture of other

gases….

PA = XAPt = (𝑛 𝐴

𝑛 𝑡𝑜𝑡𝑎𝑙 ) 𝑃𝑡

PA = pressure for a certain gas (any units of pressure can be used) XA = mole fraction (nA/nt) Pt = total pressure of all gases in container (any units of pressure can be used)

10.7 Kinetic-Molecular Theory – VERY IMPORTANT!!!

Read p. 418-421. The Kinetic-Molecular theory (KMT)was developed to explain gas behavior.

• A model that explains why real gases approach the behavior of the ideal gas

law (PV=nRT)

• KMT attempts to explain the properties of an ideal gas. (6 postulates below)

• Kinetic molecular theory gives us an understanding of pressure and

temperature on the molecular level.

o Ideal Gas Behavior: Higher Temperatures, Lower pressure

• It is a theory of moving molecules.

Summary: (KNOW THIS WELL)

1. Gases consist of molecules whose separation is much larger than the size of

the molecules themselves. Meaning – gases are mostly empty space! They

move in straight line paths and constant random directions.

2. The gas particles are so small compared with the distances between them

that the volume of the individual gas particles is assumed to be ZERO

(negligible). See ch 10.9 – This is UNTRUE for REAL gases.

3. Intermolecular forces (forces between gas molecules) are negligible. They

never attract, repel, or collide with each other. See ch 10.9 – This is

UNTRUE for REAL gases.

4. Particles in a gas collide frequently with the sides of the container and less

frequently with each other. All collisions are elastic, which means No

energy is gained or lost as a result of the collisions,

but it can be transferred from one molecule to

another. Energy can be transferred between

molecules during collisions, but the average

kinetic energy is constant at constant

temperature.

5. The collisions of gas particles with the walls of the

container cause the “pressure” exerted by the gas.

The magnitude (size) of the pressure is determined

by how often and how hard the molecules strike.

6. The average kinetic energy (KE) of the gas

molecules is directly proportional to the absolute temperature. KE increases

as T increases = greater motion of the gas particles = velocity or speed (u) of

gas molecules increase.

Read p.419 really well and Figure 10.13a. Draw this figure. Understand the red vs. blue line. Which has the higher average kinetic energy?

Application to the Gas-Laws

• So what would happen to gas molecules if there was an increase in volume

of the container (at constant temperature):

o As volume increases at constant temperature, the average kinetic

energy of the gas remains constant (KE directly related to T, not V).

o Therefore, u (velocity/speed) is constant.

o The gas molecules have to travel further to hit the walls of the

container.

o Leads to less collusions

o Therefore, pressure decreases. (Boyles law)

• So what would happen to gas molecules if there was an increase in

temperature (at constant volume):

o If temperature increases at constant volume, the average kinetic

energy of the gas molecules increases. (directly related)

o There are more collisions with the container walls.

o Therefore, u (velocity/speed) increases.

o The change in momentum in each collision increases (molecules

strike harder).

o Therefore, pressure increases. (Gay-Lussac Law)

Let’s Watch Michael Farabaugh Ch 10 video on Maxwell-Boltzman diagrams – provides graphical representation of energies/velocities of particles at a given temperature. (AP loves these graphs..hint hint)

The diagram at right can be found on the bottom of

p. 403. It shows the distribution of molecular speeds

for nitrogen gas at two different temperatures.

True or False: At a given temperature, each

molecule in a gas sample is moving at the

same speed.

True or False: As the temperature of a gas sample

is increased, the average speed of the gas molecules

increases.

True or False: The distribution curve tends to broaden

as the temperature is increased, indicating that the range

of molecular speeds increases at a higher temperature.

(c) The average kinetic energy per molecule of a gas is equal to ½ mv2

(where m = mass and v = velocity). This information helps us to compare the

average speeds of different gases at a given temperature. How should the following

gases be arranged in order of increasing average speed at 298 K: N2(g), Ne(g),

CO2(g)

(d) What does the diagram below tell us about the relationship between molecular

speed and molar mass?

10.8 Molecular Effusion and Diffusion

Read p. 421-426.

According to figure 10.14 (read p.422), which of these gases has the largest

molar mass? List them from smallest to largest. (KNOW THIS PICTURE)

According to Figure 10.16 p. 424, because pressure and temperature are

constant in this figure but volume changes, which other quantity in the ideal

gas equation must change? Compare two different balloons, one of Argon

gas and one of Helium gas. Explain this process of effusion within the

balloons.

M

RTu

3=

Root Mean Speed: rms (velocity) – essentially the same as the “average velocity of the gas particles” in a sample.

o Symbol is u or urms

o The speed of a gas is related to its mass

o Consider two gases at the same temperature: the lighter gas has a

higher rms (speed) than the heavier gas.

Mathematically:

o R = 8.314 J/K.mol

(new unit of R since we are discussing energy = Joules = J)

o M = molar mass in kg/mol (not g/mol)

o T = Kelvin

Example #1: Calculate the rms speed of the molecules in a sample of Helium and

Nitrogen gas at 25 oC.

a. Helium

b. Nitrogen

Example #2: Consider 3 gases all at 298K: HCl, H2, and O2. List the gases in

order of increasing average speed.

Example #3: How is the rms speed of N2 molecules in a gas sample changed by

a. Increase in temperature

b. Increase in volume

c. Mixing with a sample of Ar at the same temperature.

Effusion vs Diffusion: Graham’s Law **Two consequences of the dependence of molecular speeds on mass are:

1. Diffusion is the spread of one substance throughout a space or process by

which a gas spreads through a space occupied by another gas.

o Diffusion is faster for light gas molecules.

▪ Example: Mixing gases – opening a bottle of perfume or

cooking in kitchen. Eventually its odor (gas particles) can be

detected in another location of house.

o Diffusion is slowed by collisions of gas molecules with one another.

▪ Example: Consider someone opening a perfume bottle: It takes

awhile to detect the odor, but the average speed of the

molecules at 25oC is about 515 m/s (1150 mi/hr).

Effusion

1

2

2

1

M

M

r

r=

2. Effusion is the escape of gas molecules through a tiny hole into an

evacuated space. (balloon)

o Only those molecules which hit the small hole will escape through it.

o Therefore, the higher the rms speed the more likely it is that a gas

molecule will hit the hole.

o Although slightly different from diffusion, the MATH is the same

Lab Experiment: Diffusion process between HCl and NH3

o Glass type and add cotton balls to each side.

o One end of tube add 5 drops of acid while at the same time add 5 drops of

base.

o Put stoppers in glass ends so no gas escapes.

o Start timing reaction to find a “white ring on the glass where the acid and

base meet to form SOLID NH4Cl”.

o What is happening in the experiment?

o The acid and base particles evaporate off the cotton and travel inside

tube.

o Distances traveled are NOT the same due to MOLAR MASSES

(Graham’s Law equation).

Graham’s Law (Rate of Diffusion) – the rate of the mixing of the gases

o Consider two gases with molar masses, M1 and M2, and with diffusion rates,

r1 and r2, respectively.

M1 and M2 = molar mass units can be in g/mol.

Example #1: Let’s Calculate the Rate of Diffusion between HCl and NH3: What does this value mean? Example #2: Calculate the “rate of effusion” of N2 and O2 gases.

10.9 Real Gases: Deviations from Ideal Behavior – AP Free Response Question!!!!

Read p. 426-429. Real Gases:

o There is no such thing as an ideal gas, but we assume that real gases

“behave ideally”.

o According to the KMT (Ch 10.7), 2 of the postulates are UNTRUE for

real gases:

o #2 – Real gases DO have volume themselves

o #3 – Real gases DO have interparticle interactions, they do attract,

repel, and collide with each other.

o We make an assumption about real gases “behaving ideally”, from a

calculation standpoint. The math is not that far off.

o In other words, assuming that real gases behave ideally allows us to use the

ideal gas law for ALL gases (PV=nRT)

o KMT postulates #2 & #3 allow us to simplify the math involved.

OTHERWISE, the ideal gas law would have to be replaced with the “real

gas law” equation.

o See Ch 10.9: Van Der Walls Equation (NOT ON AP EXAM ☺)

nRT

PV=

See graph above: For 1 mol of an ideal gas, PV/RT = 1 for all pressures.

o The higher the pressure the more the deviation from ideal behavior.

For 1 mol of an ideal gas, PV/RT = 1 for all temperatures. • As temperature increases, the gases behave more ideally. Deviations from ideal behavior - THIS IS WHERE IDEAL GASES DIFFER FROM REAL GASES!!!!!

At high pressures and low temperatures gas particles come close enough to one

another to make the two postulates of the Kinetic Theory below, invalid.

KMT #2. The gas particles are so small compared with the distances between them

that the volume of the individual gas particles is assumed to be ZERO (negligible).

HOWEVER with real gases….

o When the gas is pressurized into a small space, the gas particles size

becomes more significant compared to the total volume.

o When gases are compressed to high pressures, the size of the gas particles is

no longer negligible compared to the total space occupied by the gas (its

total volume).

o Therefore, the observed total volume occupied by the gas under these real

conditions is actually large since the gas particles are now occupying a

significant amount of that total volume.

KMT #3. Intermolecular forces (forces between gas molecules) are negligible. They never attract, repel, or collide with each other. HOWEVER with real gases….

o Low temperature means less energy, so the particles are attracted to one

another more.

o The actual pressure of a gas is lower than one would expect when assuming

there were no attractive forces between the particles. Because, in a real gas,

the particles are attracted to one another, they collide with the walls with

less force, and the observed pressure is less than it an ideal gas.

Chapter 10. Gases

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Transcript of Chapter 10. Gases