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GASES 

• Gases are one of the most pervasive 

aspects of our environment on the 

Earth. We continually exist with 

constant exposure to gases of all 

forms. 

• The steam formed in the air during a 

hot shower is a gas. 

• The Helium used to fill a birthday 

balloon is a gas. 

• The oxygen in the air is an essential 

gas for life. www.elearnuganda.net

GASES 

A windy day or a still day is a result of the difference in pressure of gases 

in two different locations. A fresh breeze on a mountain peak is a study in 

basic gas laws. 


Important Characteristics of Gases 

1) Gases are highly compressible 

An external force compresses the gas sample and decreases its 

volume, removing the external force allows the gas volume to 

increase. 

2) Gases are thermally expandable 

When a gas sample is heated, its volume increases, and when it is 

cooled its volume decreases. 

3) Gases have high viscosity 

Gases flow much easier than liquids or solids. 

4) Most Gases have low densities 

Gas densities are on the order of grams per liter whereas liquids 

and solids are grams per cubic cm, 1000 times greater. 

5) Gases are infinitely miscible 

Gases mix in any proportion such as in air, a mixture of many gases. 


Substances That Are Gases under 

Normal Conditions 

Substance Formula MM(g/mol) 

• Helium He 4.0 

• Neon Ne 20.2 

• Argon Ar 39.9 

• Hydrogen H 2 2.0 

• Nitrogen N 2 28.0 

• Nitrogen Monoxide NO 30.0 

• Oxygen O 2 32.0 

• Hydrogen Chloride HCl 36.5 

• Ozone O 3 48.0 

• Ammonia NH 3 17.0 

www.elearnuganda.net 

• Methane CH 4 16.0

Kinetic Molecular Theory 

• To fully understand the world around us 

requires that we have a good understanding 

of the behavior of gases. The description of 

gases and their behavior can be approached 

from several perspectives. 

• The Gas Laws are a mathematical 

interpretation of the behavior of gases. 

• However, before understanding the 

mathematics of gases, a chemist must have 

an understanding of the conceptual 

description of gases. That is the purpose of 

the Kinetic Molecular Theory. 


Kinetic Molecular Theory 

• The Kinetic Molecular Theory is a single set of 

descriptive characteristics of a substance known as 

the Ideal Gas. 

• All real gases require their own unique sets of 

descriptive characteristics. Considering the large 

number of known gases in the World, the task of 

trying to describe each one of them individually 

would be an awesome task. 

• In order to simplify this task, the scientific 

community has decided to create an imaginary gas 

that approximates the behavior of all real gases. In 

other words, the Ideal Gas is a substance that does 

not exist. 

• The Kinetic Molecular Theory describes that gas. 

While the use of the Ideal Gas in describing all real 

gases means that the descriptions of all real gases 

will be wrong, the reality is that the descriptions of 

real gases will be close enough to correct that any 

errors can be overlooked. 


The Nature of Gases 

Three basic assumptions of the kinetic 

theory as it applies to gases: 

1. Gas is composed of particles- usually 

molecules or atoms 

–Small, hard spheres 

–Insignificant volume; relatively far 

apart from each other 

–No attraction or repulsion between 

particles 



2. Particles in a gas move rapidly in 

constant random motion 

–Move in straight paths, changing 

direction only when colliding with one 

another or other objects 

–Average speed of O 2 in air at 20 o C is 

an amazing 1660 km/h! 

(1.6km=1mile) 



3. Collisions are perfectly elasticmeaning 

kinetic energy is transferred 

without loss from one particle to 

another- the total kinetic energy remains 

constant 

Newtonian Cradle- 

Where the collisions between the balls elastic? 

Yes, because kinetic energy was transferred 

with each collision 


• Why did the balls eventually stop 

swinging? The collisions were not 

perfectly elastic, some kinetic energy 

was lost as heat during each collision. 

• At constant temperatures and low to 

moderate pressures, collisions between 

gas particles are perfectly elastic 


THE KINETIC THEORY OF GASES 

Remember the assumptions 

• Gas consists of large number of particles 

(atoms or molecules) 

• Particles make elastic collisions with each 

other and with walls of container 

• There exist no external forces (density 

constant) 

• Particles, on average, separated by distances 

large compared to their diameters 

• No forces between particles except when 

they collide 


What happens to a ball when it 

drops? 

The Which potential is is converted energy into 

of to the the potential potential ball energy energy of 

in the the ball………….. 

…..but in reality the ball 

loses height and 

eventually stops bouncing 

Which Is converted is converted to kinetic to 

kinetic energy energy in the ball in the 

ball 

Why does this 

happen? 


How does the bouncing ball lose 

energy? 

• Through friction with the air (air 

resistance) 

• Through sound when it hits the floor 

• Through deformation of the ball 

• Through heat energy in the bounce 



IDEAL GAS MODEL 

• The gas consists of objects with a defined mass 

and zero volume. 

• The gas particles travel randomly in straight-line 

motion where their movement can be described by 

the fundamental laws of mechanics. 

• All collisions involving gas particles are elastic; 

the kinetic energy of the system is conserved even 

though the kinetic energy among the particles is 

redistributed. 

• The gas particles do not interact with each other or 

the with the walls of any container. 

• The gas phase system will have an average kinetic 

energy that is proportional to temperature; the 

kinetic energy will be distributed among the 

particles according to a Boltzmann type of 

distribution. 


Boltzman Distribution. The behaviour of 

the gas molecules under the action of 

gravity. 


Maxwell Distribution. Experiment with 

Galton board demonstrates the 

statistical sense of Maxwell distribution. 


Ideal Gas Model 

Kinetic Molecular Theory (KMT) for an ideal 

gas states that all gas particles: 

• are in random, constant, straight-line motion. 

• are separated by great distances relative to 

their size; the volume of the gas particles is 

considered negligible. 

• have no attractive forces between them. 

• have collisions that may result in the transfer 

of energy between gas particles, but the total 

energy of the system remains constant. 


Brownian motion. Chaotic motion of 

minute particle suspended in a gas or 

liquid 


This animation illustrates the concept of 

free path length of molecules in a gas. 


Ideal vs. Non-Ideal Gases 

x 

x 

• Kinetic Theory Assumptions 

– Point Mass 

– No Forces Between Molecules 

– Molecules Exert Pressure Via Elastic 

Collisions With Walls 


(courtesy F. Remer)

Ideal vs. Non-Ideal Gases 

• Non-Ideal Gas 

– Violates Assumptions 

• Volume of molecules 

• Attractive forces of molecules 


(courtesy F. Remer)

Deviations from ideal behaviour 

• A real gas is most like an ideal gas when the 

real gas is at low pressure and high 

temperature. 

• At high pressures gas particles are close 

therefore the volume of the gas particles is 

considered. 

• At low temperatures gas particles have low 

kinetic energy therefore particles have some 

attractive force 

• Example 

• Dry ice, liquid oxygen and nitrogen 


Ideal Gases 

• Behave as described by the ideal gas 

equation; no real gas is actually ideal 

• Within a few %, ideal gas equation describes 

most real gases at room temperature and 

pressures of 1 atm or less 

• In real gases, particles attract each other 

reducing the pressure 

• Real gases behave more like ideal gases as 

pressure approaches www.elearnuganda.net zero.

Atmospheric Pressure 

• Weight of column of air above your head. 

• We can measure the density of the 

atmosphere by measuring the pressure it 

exerts. 


Effect of Atmospheric Pressure on 

Objects at the Earth’s Surface 



Atmospheric Pressure 

Pressure = Force per Unit Area 

Atmospheric Pressure is the weight of 

the column of air above a unit area. For 

example, the atmospheric pressure felt 

by a man is the weight of the column of 

air above his body divided by the area 

the air is resting on 

P = (Weight of column)/(Area of base) 

Standard Atmospheric Pressure: 

1 atmosphere (atm) 

14.7 lbs/in 2 (psi) 

760 Torr (mm Hg) 

1013.25 KiloPascals or Millibars (kPa = 

N/m 2 ) 


Pressure Measurement 

Torricelli's Barometer 

• Torricelli determined from this 

experiment that the pressure of the 

atmosphere is approximately 30 

inches or 76 centimeters (one 

centimeter of mercury is equal to 13.3 

millibars. He also noticed that height 

of the mercury varied with changes in 

outside weather conditions. 

For climatological and meteorological purposes, standard sea-level pressure 

is said to be 76.0 cm or 29.92 inches or 1013 millibars 



Atmospheric pressure results from 

the collisions of air molecules with 

objects 

–Decreases as you climb a mountain 

because the air layer thins out as 

elevation increases 

Barometer is the measuring 

instrument for atmospheric 

pressure; dependent 

www.elearnuganda.netupon 

weather

Common Units of Pressure 

Unit Atmospheric Pressure Scientific Field 

pascal (Pa); 1.01325 x 10 5 Pa SI unit; physics, 

kilopascal(kPa) 101.325 kPa chemistry 

atmosphere (atm) 1 atm* Chemistry 

millimeters of mercury 760 mmHg* Chemistry, medicine, 

( mm Hg ) biology 

torr 760 torr* Chemistry 

pounds per square inch 14.7 lb/in 2 Engineering 

( psi or lb/in 2 ) 

bar 1.01325 bar 


Meteorology, 

chemistry, physics

Converting Units of Pressure 

Problem: A chemist collects a sample of carbon dioxide from the 

decomposition of limestone (CaCO 3 ) in a closed end manometer, the 

height of the mercury is 341.6 mm Hg. Calculate the CO 2 pressure in 

torr, atmospheres, and kilopascals. 

Plan: The pressure is in mmHg, so we use the conversion factors from 

Table 5.2(p.178) to find the pressure in the other units. 

Solution: 

converting from mmHg to torr: 

P CO2 (torr) = 341.6 mm Hg x 

1 torr 

= 341.6 torr 

1 mm Hg 

converting from torr to atm: 

1 atm 

P CO2 ( atm) = 341.6 torr x = 0.4495 atm 

760 torr 

converting from atm to kPa: 

P CO2 (kPa) = 0.4495 atm x 

101.325 kPa 

1 atm 


= 45.54 kPa

Change in Pressure 

Change in average 

atmospheric pressure with 

altitude. 



Gas Pressure – defined as the 

force exerted by a gas per unit 

surface area of an object 

–Due to: a) force of collisions, and b) 

number of collisions 

–No particles present? Then there 

cannot be any collisions, and thus no 

pressure – called a vacuum 



Manometers 

Manometers measure a pressure difference by balancing the 

weight of a fluid column between the two pressures of interest 

Rules of thumb: 

• When evaluating, start from the known 

pressure end and work towards the 

unknown end 

• At equal elevations, pressure is 

constant in the SAME fluid 

• When moving down a manometer, 

pressure increases 

• When moving up a manometer, 

pressure decreases 

• Only include atmospheric pressure on 

open ends 



Manometers 


Manometers 

Example 2 

Find the pressure at 

point A in this open u- 

tube manometer with an 

atmospheric pressure P o 

P D = γ H2O x h E-D + P o 

P c = P D 

P B = P C - γ Hg x h C-B 

P A = P B 


P = γ x h + P O

The Gas Laws 

• What would Polly 

Parcel look like if she 

had no gas molecules 

inside? 

zero molecules = zero pressure inside 

zero pressure inside = zero force on the 

inside 


Gas Law Variables 

• In order to describe gases, mathematically, it 

is essential to be familiar with the variables 

that are used. There are four commonly 

accepted gas law variables 

• Temperature 

• Pressure 

• Volume 

• Moles 


Temperature 

• The temperature variable is always symbolized as T. 

• It is critical to remember that all temperature values 

used for describing gases must be in terms of 

absolute kinetic energy content for the system. 

• Consequently, T values must be converted to the 

Kelvin Scale. To do so when having temperatures 

given in the Celsius Scale remember the conversion 

factor 

• Kelvin = Celsius + 273 

• According to the Kinetic Molecular Theory, every 

particle in a gas phase system can have its own 

kinetic energy. Therefore, when measuring the 

temperature of the system, the average kinetic 

energy of all the particles in the system is used. 

• The temperature variable is representing the position 

of the average kinetic energy as expressed on the 

Boltzmann Distribution. 


Pressure 

• The pressure variable is represented by the 

symbol P. 

• The pressure variable refers to the pressure 

that the gas phase system produces on the 

walls of the container that it occupies. 

• If the gas is not in a container, then the 

pressure variable refers to the pressure it 

could produce on the walls of a container if it 

were in one. 

• The phenomenon of pressure is really a force 

applied over a surface area. It can best be 

expressed by the equation 


Pressure 

• Consider the Pressure equation and the impact of 

variables on it. 

• The force that is exerted is dependent upon the 

kinetic energy of the particles in the system. If the 

kinetic energy of the particles increases, for 

example, then the force of the collisions with a given 

surface area will increase. This would cause the 

pressure to increase. Since the kinetic energy of the 

particles is increased by raising the temperature, 

then an increase in temperature will cause an 

increase in pressure. 

• If the walls of the container were reduced in total 

surface area, there would be a change in the 

pressure of the system. By allowing a given quantity 

of gas to occupy a container with a smaller surface 

area, the pressure of the system would increase. 


Pressure 

• As this container of gas 

is heated, the 

temperature increases. 

As a result, the average 

kinetic energy of the 

particles in the system 

increases. 

• With the increase in 

kinetic energy, the force 

on the available amount 

of surface area increases. 

As a result, the pressure 

of the system increases. 

• Eventually,......................... 

.Ka-Boom 


Volume 

• The Volume variable is represented by the symbol V. 

It seems like this variable should either be very easy 

to work with or nonexistent. 

• Remember, according to the Kinetic Molecular 

Theory, the volume of the gas particles is set at zero. 

Therefore, the volume term V seems like it should be 

zero. 

• In this case, that is not true. The volume being 

referred to here is the volume of the container, not 

the volume of the gas particles. 

• The actual variable used to describe a gas should be 

the amount of volume available for the particles to 

move around in. In other words 


Volume 

• Since the Kinetic Molecular Theory 

states that the volume of the gas 

particles is zero, then the equation 

simplifies. 

• As a result, the amount of available 

space for the gas particles to move 

around in is approximately equal to the 

size of the container. 

• Thus, as stated before, the variable V is 

the volume of the container. 


Moles 

• The final gas law variable is the quantity of gas. This is always 

expressed in terms of moles. The symbol that represents the 

moles of gas is n. Notice that, unlike the other variables, it is in 

lower case. 

• Under most circumstances in chemistry, the quantity of a 

substance is usually expressed in grams or some other unit of 

mass. The mass units will not work in gas law mathematics. 

Experience has shown that the number of objects in a system 

is more descriptive than the mass of the objects. 

• Since each different gas will have its own unique mass for the 

gas particles, this would create major difficulties when working 

with gas law mathematics. 

• The whole concept of the Ideal Gas says that all gases can be 

approximated has being the same. Considering the large 

difference in mass of the many different gases available, using 

mass as a measurement of quantity would cause major errors 

in the Kinetic Molecular Theory. 

• Therefore, the mole will standardize the mathematics for all 

gases and minimize the chances for errors. 


Conclusions 

There are four variables used mathematically for describing a 

gas phase system. While the units used for the variables may 

differ from problem to problem, the conceptual aspects of the 

variables remain unchanged. 

1. T, or Temperature, is a measure of the average kinetic energy of 

the particles in the system and MUST be expressed in the 

Kelvin Scale. 

2. P, or Pressure, is the measure of the amount of force per unit of 

surface area. If the gas is not in a container, then P represents 

the pressure it could exert if it were in a container. 

3. V, or Volume, is a measure of the volume of the container that 

the gas could occupy. It represents the amount of space 

available for the gas particles to move around in. 

4. n, or Moles, is the measure of the quantity of gas. This 

expresses the number of objects in the system and does not 

directly indicate their masses. 


Gas Laws 

• (1) When temperature is held constant, the density of a 

gas is proportional to pressure, and volume is inversely 

proportional to pressure. Accordingly, an increase in 

pressure will cause an increase in density of the gas and 

a decrease in its volume. – Boyles’s Law 

• (2) If volume is kept constant, the pressure of a unit 

mass of gas is proportional to temperature. If 

temperature increase so will pressure, assuming no 

change in the volume of the gas. 

• (3) Holding pressure constant, causes the temperature of 

a gas to be proportional to volume, and inversely 

proportional to density. Thus, increasing temperature of 

a unit mass of gas causes its volume to expand and its 

density to decrease as long as there is no change in 

pressure. - Charles’s Law 



Boyle’s Law 

• Hyperbolic Relation Between Pressure and 

Volume 

T 1 T 2 T 3 T 3 >T 2 >T 1 

(courtesy F. Remer) 

p 

isotherms 

V 

p – V Diagram 



Charles’ Law 

• Linear Relation Between Temperature and 

Pressure 

V 1

Charles’ Law 

Real data must be 

obtained above 

liquefaction 

temperature. 

Experimental curves for 

different gasses, 

different masses, 

different pressures all 

extrapolate to a 

common zero. 


Another version of Charles Law 


Compression and expansion of 

adiabatically isolated gas is accompanied 

by its heating and cooling. 



The Gas Laws 

• What would Polly 

Parcel look like if she 

had a temperature of 

absolute zero inside? 

absolute zero = no molecular motion 

no molecular motion = zero force on 

the inside 


Ideal Gas Law 

The equality for the four variables involved 

in Boyle’s Law, Charles’ Law, Gay-Lussac’s 

Law and Avogadro’s law can be written 

PV = nRT 

R = ideal gas constant 


PV = nRT 

R is known as the universal gas constant 

Using STP conditions 

P V 

R = PV = (1.00 atm)(22.4 L) 

nT (1mol) (273K) 

n 

T 

= 0.0821 L-atm 

www.elearnuganda.net mol-K

Learning Check 

What is the value of R when the STP value 

for P is 760 mmHg? 


Solution 

What is the value of R when the STP value 

for P is 760 mmHg? 

R = PV = (760 mm Hg) (22.4 L) 

nT (1mol) (273K) 

= 62.4 L-mm Hg 

mol-K 



Dinitrogen monoxide (N 2 O), laughing gas, 

is used by dentists as an anesthetic. If 2.86 

mol of gas occupies a 20.0 L tank at 23°C, 

what is the pressure (mmHg) in the tank in 

the dentist office? 


Solution 

Set up data for 3 of the 4 gas variables 

Adjust to match the units of R 

V = 20.0 L 

T = 23°C + 273 

n = 2.86 mol 

20.0 L 

296 K 

2.86 mol 

P = ? ? 


Rearrange ideal gas law for unknown P 

P = nRT 

V 

Substitute values of n, R, T and V and 

solve for P 

P = (2.86 mol)(62.4L-mmHg)(296 K) 

(20.0 L) (K-mol) 

= 2.64 x 10 3 mm Hg 



A 5.0 L cylinder contains oxygen gas 

at 20.0°C and 735 mm Hg. How many 

grams of oxygen are in the cylinder? 


Solution 

Solve ideal gas equation for n (moles) 

n = PV 

RT 

= (735 mmHg)(5.0 L)(mol K) 

(62.4 mmHg L)(293 K) 

= 0. 20 mol O 2 x 32.0 g O 2 = 6.4 g O 2 


1 mol O 2

Molar Mass of a gas 

What is the molar mass of a gas if 0.250 g of 

the gas occupy 215 mL at 0.813 atm and 

30.0°C? 

n = PV = (0.813 atm) (0.215 L) = 0.00703 mol 

RT 

(0.0821 L-atm/molK) (303K) 

Molar mass = g = 0.250 g = 35.6 g/mol 

mol 

0.00703 mol 


Density of a Gas 

Calculate the density in g/L of O 2 gas at STP. 

From STP, we know the P and T. 

P = 1.00 atm 

T = 273 K 

Rearrange the ideal gas equation for moles/L 

PV = nRT PV = nRT P = n 

RTV RTV RT V 


Substitute 

(1.00 atm ) mol-K = 0.0446 mol O 2 /L 

(0.0821 L-atm) (273 K) 

Change moles/L to g/L 

0.0446 mol O 2 x 32.0 g O 2 = 1.43 g/L 

1 L 1 mol O 2 

Therefore the density of O 2 gas at STP is 

1.43 grams per liter 


Formulas of Gases 

A gas has a % composition by mass of 

85.7% carbon and 14.3% hydrogen. At 

STP the density of the gas is 2.50 g/L. 

What is the molecular formula of the 

gas? 


Formulas of Gases 

Calculate Empirical formula 

85.7 g C x 1 mol C = 7.14 mol C/7.14 = 1 C 

12.0 g C 

14.3 g H x 1 mol H = 14.3 mol H/ 7.14 = 2 H 

1.0 g H 

Empirical formula = CH 2 

EF mass = 12.0 + 2(1.0) = 14.0 g/EF 


Using STP and density ( 1 L = 2.50 g) 

2.50 g x 22.4 L = 56.0 g/mol 

1 L 1 mol 

n = EF/ mol = 56.0 g/mol = 4 

14.0 g/EF 

molecular formula 

CH 2 x 4 = C 4 H 8 


Gases in Chemical Equations 

On December 1, 1783, Charles used 1.00 x 10 3 

lb of iron filings to make the first ascent in a 

balloon filled with hydrogen 

Fe(s) + H 2 SO 4 (aq) FeSO 4 (aq) + H 2 (g) 

At STP, how many liters of hydrogen 

gas were generated? 


Solution 

lb Fe g Fe mol Fe mol H 2 L H 2 

1.00 x 10 3 lb x 453.6 g x 1 mol Fe x 1 mol H 2 

1 lb 55.9 g 1 mol Fe 

x 22.4 L H 2 = 1.82 x 10 5 L H 2 

1 mol H 2 

Charles generated 182,000 L of hydrogen to fill his 

air balloon. 



How many L of O 2 are need to react 28.0 g 

NH 3 at 24°C and 0.950 atm? 

4 NH 3 (g) + 5 O 2 (g) 4 NO(g) + 6 H 2 O(g) 


Solution 

Find mole of O 2 

28.0 g NH 3 x 1 mol NH 3 x 5 mol O 2 

17.0 g NH 3 4 mol NH 3 

= 2.06 mol O 2 

V = nRT = (2.06 mol)(0.0821)(297K) = 52.9 L 

P 

0.950 atm 


Mixture of gases 


Reacting mixture of gases 



A.If the atmospheric pressure today is 745 

mm Hg, what is the partial pressure (mm 

Hg) of O 2 in the air? 

1) 35.6 2) 156 3) 760 

B. At an atmospheric pressure of 714, what is 

the partial pressure (mm Hg) N 2 in the air? 

1) 557 2) 9.14 3) 0.109 


Solution 

A.If the atmospheric pressure today is 745 

mm Hg, what is the partial pressure (mm 

Hg) of O 2 in the air? 

2) 156 

B. At an atmospheric pressure of 714, what is 

the partial pressure (mm Hg) N 2 in the air? 

1) 557 


Partial Pressure 

Partial Pressure 

Pressure each gas in a mixture would exert 

if it were the only gas in the container 

Dalton's Law of Partial Pressures 

The total pressure exerted by a gas mixture 

is the sum of the partial pressures of the 

gases in that mixture. 

P T = P 1 + P 2 + P 3 + ..... 


Partial Pressures 

The total pressure of a gas mixture depends 

on the total number of gas particles, not on 

the types of particles. 

STP 

P = 1.00 atm 

P = 1.00 atm 

1.0 mol He 

0.50 mol O 2 

+ 0.20 mol He 

+ 0.30 mol N 2 


Health Note 

When a scuba diver is several hundred feet 

under water, the high pressures cause N 2 from 

the tank air to dissolve in the blood. If the 

diver rises too fast, the dissolved N 2 will form 

bubbles in the blood, a dangerous and painful 

condition called "the bends". Helium, which is 

inert, less dense, and does not dissolve in the 

blood, is mixed with O 2 in scuba tanks used for 

deep descents. www.elearnuganda.net


A 5.00 L scuba tank contains 1.05 mole of 

O 2 and 0.418 mole He at 25°C. What is the 

partial pressure of each gas, and what is 

the total pressure in the tank? 


Solution G20 

P = nRT P T = P O + P He 

V 

2 

P T = 1.47 mol x 0.0821 L-atm x 298 K 

= 7.19 atm 

5.00 L (K mol) 


Micro Effusion 


Macro Effusion

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