School Science Lessons
(UNPh20)
2024-07-29
Gas Laws
Contents
20.1.01 Kinetic theory of gases
20.1.02 Combined gas equation
20.10 Boyle's law
20.1.0a Boyle's law, Experiments
20.2.0 Charles's law
20.2.0a Charles's law, Experiments
20.3.0 Gas laws
20.4.0 Thermodynamics
20.1.0 Boyle's law
"Air & Water Pressure Kit" (Boyle's law), potato gun (Commercial).
PV = constant, if temperature is constant (from a pressure / volume graph)
Robert Boyle (1627-1691), For a given mass of gas at constant temperature, the volume is inversely proportional to the pressure, PV = constant.
The Boyle's law relationship would be true only for an ideal gas with particles that occupy no space, have no forces between them, and have perfectly elastic collisions between them and between the particles and the walls of the container.
The pressure of a gas is caused by the gas particles colliding with the walls of the container.
V1 / V2 = P2 / P1, P1V1 = P2 V2
The volume of a gas enclosed in a cylinder is halved when the piston
is pushed down half way.
This action doubles the number of molecules per cubic centimetre, so there are twice as many collisions with the walls of the cylinder that causes the pressure to double.
P1V1 = P2 V2, 1 × 1 = P2 × 0.5, So P2 = 2.
Constant temperature, Boyle's law, P1V1= P2V2, pV = constant
20.1.0a Boyle's law, Experiments
20.1.1 Air pump, (Experiment)
20.1.2 Lift weight by blowing, the work done by gas pressure, (Experiment)
20.1.3 Plastic syringes and air pressure, Boyle's law, (Experiment)
20.1.4 Potato gun pneumatic launcher, (Experiment)
20.1.5 Pressure effect on gas volume of syringe, (Experiment)
20.1.6 Scuba diving and Boyle's law
20.1.7 Volume and pressure of air, Boyle's law, (Experiment)
20.2.0 Charles's law
Constant gas pressure, Charles's law (Gay-Lussac's law), V1 / T1 = V2 / T2, V / T = constant
Charles's law (Gay-Lussac's law), V1 / T1 = V2 / T2, V / T = constant
See diagram 20.0.0: Contraction of gas obeying Charles's law.
(J. A. C. Charles 1746-1823) (J. L. Gay-Lussac 1778-1850)
For a given mass of gas at constant pressure, the volume, V, is proportional to the absolute or Kelvin temperature of the gas, T,
Volume = constant × T, V / T = constant provided pressure is constant.
Volume is proportional to temperature, V1 / T1 = V2 / T2, where T = absolute temperature, (Absolute zero = -273oC)
For a given mass of gas at constant pressure, the volume, V, increases by 1/ 273 rd of its volume at 0oC for every Celsius degree rise in temperature:
Vt = Vo[1 + (1 / 273)t]
The new internal pressure, P2, of a litre of gas at 20oC, enclosed at pressure 770 mmHg, then heated to 100oC = P1 / T1 = P2 / T2,
(V1 = V2), 770 / 293 = P2 / 373, So P2 = 980 mm Hg
Charles's law (J. A. C. Charles 1746-1823)
For a given mass of gas at constant pressure, the volume is proportional to the absolute thermodynamic temperature, K., V = KT.
So all gases have the same coefficient of expansion at constant pressure.
However, this is strictly true only at low pressures and high temperatures.
Gay-Lussac's law (J. L. Gay-Lussac 1778-1850):
When gases combine chemically, the volumes of the reactants and the volume of the gaseous product bear simple relationships to each other under the same conditions of temperature and pressure.
Experiment: The effect of temperature on the volume of a gas
See diagram 20.1.0: Charles's law.
V1 / T1 = V2 / T2 Constant Pressure, volume / temperature graph, V / T = constant, PVT relationship.
Experiment
Quantitative treatment of ideal gases.
Boltzman's Constant, approximations used for real gases
Put a drop of oil into the capillarity tube to seal a column of air.
Measure the length of the trapped column of air rather than its volume.
Use a spring band to fix a capillarity tube and a thermometer together,
put them into a beaker.
Use the scale intervals on a thermometer to measure length.
To read easily, the lowest level of a column of air trapped in a capillarity tube is better to meet at 0oC on the thermometer.
Then record the length of the column of air by using the scale nearest to top of the column of air.
Measure a set of values of length of a column of air and temperature between 0oC and 100oC.
Before the experiment, mix crushed ice and water in a bottle of mineral water, put them into a beaker.
Record your temperature readings in a suitable position of set up the table.
Then pour tap water into the beaker, heat it by an alcohol burner.
As the temperature reaches 40oC, 60oC, 80oC, record the length of the column of air trapped in a capillarity tube each.
Remove the alcohol burner before taking records.
As water boils, record the last reading.
Draw a graph of temperature t and length of the column of air L so your graph can show how of volume of air varies with temperature.
20.2.0a Charles's law, Experiments
20.2.1 Bicycle pump
20.2.2 Carbon dioxide gas cylinder
20.2.3 Charles's law, Use oil instead of mercury for school Charles's law experiments
20.2.4 Cold air is heavier than warm air, inverted paper bag balance
20.2.5 Expansion indicator
20.2.6 Gas-filled bulb and U-tube manometer
20.2.7 Heat air and cool air
20.2.8 Heat flask with hands
20.2.9 Heated air expands in a balloon
20.2.10 Hot air balloons
20.3.0 Gas laws
Gas laws relate the pressure, temperature and volume of gas
20.3.1 Boltzmann constant / Ideal gas constant
20.3.2 Dalton's law of partial pressures
Density of gases at STP (Table)
20.5.0 Gas expansion caused by heat
20.3.3 Gas pressure
20.3.4 Henry's law and decompression sickness, the bends
20.3.5 Pressure law, constant volume
20.3.6 Standard atmosphere
20.3.7 Standard temperature and pressure, STP, density of gases
20.3.8 Universal gas equation, Ideal gas constant
20.5.0 Gas expansion caused by heat
20.5.1 Expansion of air, flask or bottle
20.5.2 Expansion of air, inverted test-tube
20.5.3 Expansion of air in a balloon
20.4.0 Thermodynamics
20.4.1 Thermodynamics, isothermal change and adiabatic change
First law of thermodynamics
Second law of thermodynamics, (Experiment)
Third law of thermodynamics
20.4.4 Adiabatic change
See diagram 20.4.04 Adiabatic process (Wikipedia, edited).
20.4.5 Adiabatic compression
20.4.6 Adiabatic cooling, (Experiment)
20.4.7 Adiabatic processes, (Experiment)
20.4.8 Cincinnati flask, (Experiment)
20.4.9 Crookes' radiometer, (Experiment)
20.4.10 Expansion cloud chamber, (Experiment)
20.4.11 Heat cycles, Carnot cycle, Stirling engine, (Experiment)
22.4.12 Heat transfer and laws of thermodynamics, (Experiment)
22.4.13 Heat transfer coefficient, 1/thermal insulation
22.4.14 Heat transfer by Dewar flask (vacuum flask, "Thermos" flask)
22.4.15 Reduce heat loss with heat insulation materials
20.4.16 Isothermal change, Isothermal change and adiabatic change, Diesel engine
Layers of the atmosphere, lapse rate, auroras, (adiabatic lapse rates)
20.4.17 Maxwell's Demon, (Experiment)
20.1.01 Kinetic theory of gases
The temperature of a gas is proportional to the mean kinetic energy of translation of the molecules of the gas.
Two gases at the same temperature have the same mean kinetic energy.
The pressure of a gas is caused by the impact of its molecules on the wall of the containing vessel.
Pressure, p = 1 / 3 nmVrms2, where p = pressure, m = mass of a molecule, n = number of molecules,
Vrms = root mean square velocity. p = 1/3 ρ Vrms2, where ρ = density of the gas
The density of a gas is inversely proportional to the absolute temperature.
20.1.02 Combined gas equation
The equations for Charles's law, Boyle's law and the Pressure law can be combined as follows:
PV / T = constant (depending on the nature and the mass of that gas), or
P1V1 / T1 = P2V2 / T2, where Temperature, T, should be in absolute scale, Kelvin scale, K.
20.1.1 Air pump
Use of a syringe needle may be not allowed in some school systems.
If temperature is constant, when you compress gas and reduce its volume, its pressure will increase, and vice versa.
Insert a piston covered with some glycerine into a 100 mL pump with a valve.
Rotate the piston inside the pump several times to make the glycerine distributed evenly.
This can insulate the air inside the pump from outside completely.
Open the valve and suck up 60 mL air into the pump.
Measure the volume of the air with the scale on the pump.
Close the valve to insulate the air in the pump from outside.
Push the piston to compress the air volume to about 2 / 3 of the original, i.e. about 40 mL.
Release the piston that will come back to the original position.
Pull the piston out with effort to expand the air volume in the pump to about 80 mL.
Release the piston that will come back to the original position.
The reason of coming back of the piston is the pressure difference between two sides of the piston.
At constant temperature, the more the air volume inside the pump is compressed, the more pressure it has.
As the air volume expands, the pressure decreases.
When the piston is compressed, as the air pressure inside the pump is higher than that of outside, the air inside the pump will push the piston back to its original position.
When you pull out the piston, the air pressure inside the pump becomes less, the atmospheric pressure outside pushes the piston back.
You can do the experiment with a large glass syringe instead of a pump.
Close the hole with the fingers used as a valve.
20.1.2 Lift weight by blowing, the work done by gas pressure
See diagram 20.2.3: Lift weight by blowing.
20.1.3 Plastic syringes and air pressure, Boyle's law
"Vacuum Container & Pump", reduced pressure increases balloon size, (industrial product)
"Vacuum Stoppers", creates near vacuum in plastic syringes (industrial product)
See diagram 12.301: Syringes and air pressure.
[Some school systems do not allow the use of syringes in the classroom.]
1. With the tip sealed, use a syringe to compress air or to produce a partial vacuum.
Attach a small piece of plastic tubing to seal the tip with a pinch clamp or seal the syringe by pushing the tip into a wooden block drilled to the appropriate size.
With this base as a platform, use the syringe in a vertical position as a balance for measuring weight by air compression.
You can quantify all the following experiments, because syringes are already graduated.
2. Fill the syringe with a small amount of air and hang it inverted to serve as a "spring type" balance.
3. Compress moist air within a syringe to cause water condensation and make "artificial rain".
4. Attach a length piece of plastic tubing to make a simple syringe pump.
5. Put water in the tube to make an air thermometer or use 12 m of tubing to make a water barometer.
6. Couple two syringes with a piece of tubing to show pressure changes within closed systems.
20.1.4 Potato gun pneumatic launcher (pneumatic, Greek pneuma wind)
Be careful! This experiment, often called a spud gun, can be dangerous if participants are not wearing safety goggles and if the metal tube is not kept in a vertical position.
Use 2 m of 3 cm internal diameter metal piping and a 5 m wooden pole < internal diameter of the metal piping.
Secure the wooden pole vertically in the ground with the length of the metal piping above the ground.
Push each end of the metal piping into a potato so that the cavity of the metal tube is completely blocked.
Hold the metal tube vertically above the wooden pole.
Hold the metal tube with both hands and push down quickly.
The potato plug flies out of the other end of the metal tube.
The increase of pressure on the air between the potato plugs causes expansion of the gases between them.
This application of Boyle's law is used in many more powerful devices, e.g. dry ice gun, that have no place in a school science laboratory.
20.1.5 Pressure effect on gas volume of syringe
See diagram 20.2.1: Pressure effect on gas volume of syringe.
Use a calibrated syringe mounted on a block of wood and with a platform securely attached to the top of the plunger.
Measure the masses of platform and plunger, the outer diameter of the plunger or the inner diameter of a syringe.
Put light oil on the plunger to lubricate it.
Lift the plunger, record the original position of it.
Seal the outlet with a piece of rubber tube.
Put weights on the platform and record volume of air in the cylinder using the scale on the syringe.
Change the weights on the platform, record the volume of the air in the syringe under different case, but maintain the temperature constant in this process.
Calculate the air pressure.
The pressure acted on air in the syringe = atmospheric pressure + the pressure produced by weights of plunger and platform + the pressure produced by weights added on the platform.
Observe and test according to measured volume and pressure calculated.
As the temperature is constant and the gas has a definite mass, when its pressure increases its volume decreases, and vice versa.
The product of pressure and volume of the gas remains the same, i.e. PV = C.
Finally, graph the relationship between volume and pressure of air in the syringe.
20.1.6 Scuba diving and Boyle's law
Swimmers descending more than one metre experience increase in pressure caused by the weight of the overlying water.
The resulting decrease in volume of air spaces causes mild to acute pain in the middle ear, collapse of the auditory tubes and the forcing inwards of the tympanic membranes.
The swimmer can overcome this discomfort by closing the mouth, pinching the nose and exhaling.
This action increases the pressure in the nasopharynx, forces air through the auditory tubes to the middle ear to return the tympanic membrane to its normal position.
When the swimmer returns to the surface, the pressure drops and the air in the middle ear expands, forcing its way out through the auditory tube and into the nasopharynx.
So the swimmer back on the surface normally feels no discomfort unless a blockage occurs in these passages.
Scuba divers breathe air under the same pressure as the level in the water.
Descending more than 10 metres doubles the pressure on a diver.
If the diver takes a full breath of air at this depth and quickly returns to the surface, the volume of air in the lungs will double.
This expansion may cause a tear in the lung wall and even air bubbles in the blood vessels leading to heart attack or stroke.
So scuba divers avoid holding their breath and must exhale when returning to the surface to avoid this over expansion syndrome.
20.1.7 Volume and pressure of air, Boyle's law
See diagram 4.238: Volume and pressure of air.
1. Use a rubber stopper that just fits inside a measuring cylinder or large syringe.
Attach it to the lower end of a wooden rod.
Fit a lid to the upper end of the rod to act as a scale pan.
Lubricate the piston so formed with some petroleum jelly or heavy engine oil.
Use the piston to trap air in the container, put different weights on the pan and measure the volume of air inside the glass cylinder for each weight.
Note that the volume is in inverse proportion to the pressure.
At constant temperature as the volume, V, decreases the pressure, P, increases, so P × V = a constant.
This relationship is called Boyle's law.
20.2.1 Bicycle pump
Pump up a bicycle tyre and feel the increase in temperature.
Measure the decrease in volume of air in the pump during a pump stroke.
The temperature difference is very small for one stroke, but if you keep pumping vigorously you can feel the difference in temperature.
The air in the sealed off pump is compressed quickly, so work is done on the air.
With little time for heat transfer to occur, so Q is = approximately 0, and the change in the gas is approximately an adiabatic process.
The rise in temperature indicates the increase in internal energy, is indicated by the feel of the pump, or by using a thermocouple inside the pump.
Energy is conserved (first law of thermodynamics) and the work done is converted to thermal (internal) energy of the air in the pump.
20.2.2 Carbon dioxide gas cylinder
Open the valve and observer the formation of dry ice.
20.2.3 Use oil instead of mercury for school Charles's law experiments
After Geoff Snowdon, The Australian Science Teachers Journal, Vol. 33 No. 2
Coloured oil can be put into a 30 centimetre length of capillary tubing by using the following procedure:
Leave both ends open.
Heat the tube strongly at one third the length.
Dip an end into the oil.
The oil rises into the tube.
Manipulate the tube to get a 5 cm length of oil.
20.2.4 Cold air is heavier than warm air, inverted paper bag balance
See diagram 37.118: Balanced flasks.
1. Use two identical paper bags or balloons that are the same size.
Inflate each by blowing into them.
Tie the openings closed tightly with string.
Tie the end of the string into a loop and suspend the bags from the end of a balanced rod.
Move the loops along the rod until the inflated bags exactly balance.
Gently heat the air beneath one bag with a small candle.
The bag containing the heated air moves up and the bag containing the cooler air moves down.
Move the candle under the other bag to see the same result.
The bags are sealed and so the mass of gas is unchanged when heating or cooling takes place.
This experiment shows Archimedes' principle in action, not mass change.
2. Open two same size paper bags.
Attach identical pieces of string to the bottom of each bag with an identical pieces of adhesive tape.
Make a loop in the other end of each piece of string.
Put the loops over each end of a balanced rod.
Adjust the positions of the loops until the rod is horizontal.
Heat the air below one paper bag.
The end of the rod supporting that paper bag rises.
Leave the balance to stand without heating a bag.
The rod becomes horizontal again.
Heat the air below the other bag.
The other end of the rod rises.
This experiment shows that a volume of warm air weighs slightly less than a volume of cool air.
However, the experiment does not give any information about the weight of a volume of air.
The flame under the paper bag heats the air in it and it expands, following Charles's law.
Some heated air spills out of the paper bag leaving less air and less dense air in the paper bag.
The air in the heated paper bag weighs less than the air it displaces
By Archimedes' principle there is an upthrust greater than its weight that causes the paper bag to rise.
When you remove the flame, the warm air in the paper bag cools and contracts drawing in air at atmospheric pressure.
The weight of a paper bag full of air and the bag crunched together, with all the air squeezed out, is the same.
If air in a hot air balloon is heated, it expands and becomes lighter and the balloon is pushed up.
The air left in the balloon is less dense than the surrounding atmospheric air.
Seal an end or heat to seal.
20.2.5 Expansion indicator
See diagram 20.1.1.3: Expansion indicator.
Use a piece of thick cardboard on a table as a base.
Paste another piece of cardboard vertically at the side of the base and mark it as a scale.
Stretch tight a rubber film over the mouth of a bottle to air proof the bottle.
Flatten one end of a drinking straw then paste it at the middle of the rubber film.
Cut the other end of the drinking straw into a sharp needle to act as an indicator.
Place the bottle on the base.
Adjust the position of the bottle so that the indicator points to half way up the scale.
Observe the movement of the indicator during the day.
When heated the air in the bottle expands to press the rubber film so that the indicator moves up.
20.2.6 Gas-filled bulb and U-tube manometer
Connect a glass bulb containing air or other gases to one arm of a manometer.
Place your hand over the bulb and observe the change in levels of the liquid in the manometer.
20.2.7 Heat air and cool air
See diagram 20.1.2: Heat conical flask.
Use a 100 mL conical flask, a rubber stopper, a N-shaped capillary of 250 mL length and a straight capillary longer than the height of the flask,
Use the equipment with a 400 mL beaker of coloured water.
Add ink to the water.
20.2.8 Heat flask with hands
See diagram 2.1.3: Heat flask with hands.
1. Use a flask fitted with a one-hole stopper and glass tube that extends into the flask.
Put a small amount of oil in the glass tube to trap air in the flask.
Hold the flask in your hands.
The oil moves up the tube, because the heat from your hands causes the trapped air to expand.
If you look carefully note that the oil first moves down, because the heat from your hands first causes the glass of the flask to expand.
When you cool the flask under the tap, the oil moves down.
2. Use a small bottle or flask fitted with a stopper and inserted glass tube that extends into the bottle.
Put a small amount of oil in the glass tube to trap air in the flask.
Hold the flask in your hands.
The oil moves up the tube, because the heat from your hands has expanded the air.
The bottle will change size first before you heat, or cooled the gas, because the glass of the bottle will expand.
Cool the flask.
The oil moves down.
Seal the flask with the rubber stopper.
On the stopper insert the N-shaped capillary.
Insert the other end of the N-shaped capillary into the coloured water at the beaker.
Cover the flask with your hands to heating the air in the flask.
Observe the end of the capillary under the coloured water.
Leave your hands off the flask then hold the capillary.
Observe the end of the capillary under the coloured water again.
While you heat the air in the flask, its volume expands and pressure increases.
So air bubbles appear at the end of the capillary until the pressure inside the flask is equal to the outside pressure.
The amount of the air in the flask decreases at the process.
While the air in the flask becomes cold, the air pressure decreases to less than the outside pressure so that the coloured water in the beaker under the atmosphere pressure, enters the capillary to contract the air volume to make the inside and outside pressures balance.
3. Fit a hard glass test-tube with a one-hole stopper that has a length of glass tubing through it.
Invert the test-tube so that the end of the tubing is in a container of water.
Clamp the test-tube in an inverted position so that you can heat it with a burner.
Heat the test-tube and note the bubbles from the end of the tube in the container of water.
Heat has caused the air to expand.
Cool the test-tube by pouring cold water over it.
Water moves up the glass tubing as the cooling air contracts.
4. Fit a toy balloon over the neck of a small flask.
Put the flask in a container of water.
Heat the water.
The balloon expands as the heated air in the flask expands.
Partially inflate a balloon and tie the neck tightly.
Leave it in a warm place or in the sunlight.
The balloon becomes fully inflated as the air inside expands when heated.
20.2.9 Heated air expands in a balloon
See diagram 20.1.1: Heated air expands.
1. Fit a toy balloon over the neck of a small flask.
Put the flask in a container of water.
Heat the water.
The balloon expands as the heated air in the flask expands.
Partially inflate a balloon and tie the neck tightly.
Leave it in a warm place or in the sunlight.
The balloon becomes fully inflated as the air inside expands when heated.
Fit a hard glass test-tube with a one hole stopper with glass tubing through it.
Invert the test-tube so that the end of the tubing is in a beaker of water.
Clamp the test-tube in that inverted position and heat it with a Bunsen burner.
Heat the test-tube and observe the bubbles from the end of the tube in the beaker of water.
Heat has caused the air to expand.
Cool the test-tube by pouring cold water over it.
Water moves up the glass tubing as the cooling air contracts.
2. Use a flask fitted with a one-hole stopper and glass tube that extends into the flask.
Put a small amount of oil in the glass tube to trap air in the flask.
Hold the flask in your hands.
The oil moves up the tube, because the heat from your hands causes the trapped air to expand.
If you look carefully note that the oil first moves down, because the heat from your hands first causes the glass of the flask to expand.
When you cool the flask under the tap, the oil moves down.
3. Fit a hard glass test tube with a one-hole stopper that has a length of glass tubing through it.
Invert the test tube so that the end of the tubing is in a container of water.
Clamp the test tube in an inverted position so that you can heat it with a burner.
Heat the test tube and note the bubbles from the end of the tube in the container of water.
Heat has caused the air to expand.
Cool the test tube by pouring cold water over it.
Water moves up the glass tubing as the cooling air contracts.
20.2.10 Hot air balloons
As the temperature of air in the balloon increases, its volume increases to inflate and lift the balloon.
A buoyant force, upthrust, is caused by the weight of air displaced by Archimedes' principle.
The density of air inside the balloon is less than the density of the surrounding air.
In 1783, J. A. C. Charles, in Paris tested a balloon four metres in diameter containing hydrogen.
This experiment lead to Charles's law in 1787, later to be improved as Gay-Lussac's law in 1809 by J. L. Gay-Lussac.
20.3.1 Boltzmann constant / Ideal gas constant
The Boltzmann constant (Ludwig Boltzmann, 1844-1906, Austria), kB, relates the energy of particles to temperature, as the gas constant, R, divided by the Avogadro, NA.
The SI value is 1.380648 x 10-23 J / K.
It relates temperature increase per particle.
The ideal gas constant (universal gas constant), R is 8.3144621 J K-1 mol-1.
It relates temperature increase per mole.
20.3.2 Dalton's law of partial pressures
The total pressure of a mixture of gases or vapours in a closed container is equal to the sum of the partial pressures of each gas or vapour, i.e. the sum of the pressures if each gas or vapour alone occupied the space in the closed container.
So each gas or vapour exerts its own pressure regardless of the presence of any other gas or vapour.
When a gas is collected over water, the water molecules in the water vapour contribute to the total pressure over the water.
Total pressure = pressure of gas produced + pressure of water vapour.
20.3.3 Gas pressure
Gases are made up of molecules moving randomly at high speeds, in straight lines, in all directions.
Pressure is a measure of the force per unit area on a surface.
Pressure = Force / Area where force is in newton, N, and area in square metres, m2.
1 newton / metre2 = 1 pascal, Pa.
The pressure of a gas = number of collisions per second per unit area × the average impulse per collision for the molecules.
Standard pressure of one atmosphere, Pa = 1.02 × 105 N / m2.
20.3.4 Henry's law and decompression sickness, "the bends"
Henry's law
If the temperature is constant, the mass of gas dissolved at equilibrium is directly proportional to the partial pressure of the gas.
The solubility of a gas depends directly on the gas pressure.
The concentration of dissolved gas depends on the partial pressure of the gas.
If the pressure is doubled the concentration of the dissolved gas double.
If the temperature remains constant increasing the pressure will increase the amount of dissolved gas.
O2 (g) <-->O2 (aq)
Pgas = KC, at constant temperature, where P = pressure, C = concentration and K = the Henry's law constant that is different for every gas, temperature and solvent.
The concentration to pressure ratio is the same when the pressure changes.
C1 / P1 = C2 / P2.
Decompression sickness:
It occurs when a person experiences a sudden change in atmospheric pressure.
Nitrogen has a low solubility in body fluids at normal atmospheric pressure, but at higher than normal pressure, additional nitrogen molecules diffuse across the alveolar surfaces of the lungs and into the bloodstream and tissues.
If the atmospheric pressure then decreases slowly, the excess nitrogen can diffuse out of the tissues, into the blood, and across the alveolar surfaces without discomfort to the person.
However, if the atmospheric pressure decreases suddenly, nitrogen leaves solution and forms bubbles of nitrogen gas in the blood, tissues, and body fluids.
Small bubbles fuse to form larger bubbles that twist tissues and produce severe pain in joint capsules and causing the person to bend over, "the bends".
Treatment by recompression forces the nitrogen in the tissues back into solution to alleviate the problem.
Then a gradual reduction in pressure allows the nitrogen to diffuse slowly out of the tissues without forming bubbles.
The bends still occurs to scuba divers who have dived too deep or stayed too long at depth, to construction crews working in pressurized surroundings and even to passengers experiencing a sudden loss of cabin pressure in a commercial aeroplane, where a pressure experienced at 2000 metres height is normally maintained.
20.3.5 Pressure law, constant volume
The pressure law states for a gas with constant volume the pressure is proportional to the Kelvin temperature of the gas.
The volume occupied by one mole of gas at STP = 2.24 × 10-2 m3.
See diagram: Constant volume bulb.
A constant volume bulb is used to measure the pressure in the sealed sphere.
Hold the bulb in the hand or immerse it in ice water and note change in pressure shown by gauge.
Read the pressure when the bulb is immersed in boiling water, in air at room temperature, and put different gases into the bulb.
20.3.6 Standard atmosphere
A standard atmosphere (International Standard Atmosphere, ISA), atm, is a hypothetical atmosphere used as a basis of comparing altimeters.
At mean sea level, 1 atm is a pressure of 101.325 Nm-2 (101, 325 pascals), equivalent to the pressure exerted by a column of mercury 760 cm high at 0oC.
The density of a gas is inversely proportional to the absolute temperature provided the pressure remain constant, ρ 2 / ρ 1 = T1 /T2, where ρ = density and T = absolute temperature.
The density of a gas is directly proportional to the pressure, if the temperature remain constant.
The standard atmosphere (USA, 1976 ) has mean conditions at sea level as follows:
Pressure 101325 Pa, Temperature 288.15 K (15oC), Density, ρ, 1.225 kg / m3, Standard gravity, g = 9.90665 m /s2, R = 8.31432 JK-1 mol-1,
and Composition: N2 (78.084%), O2 (20.9476%), Ar (0.934%),
CO2 (0.0314%), Ne (0.001818%), He (0.000524%), CH4 (0.0002%).
However, the international standard atmosphere (ISA) used in international weather data has conditions at sea level pressure 101325 Pa (1 atmosphere) at 15oC with lapse rate -6.5oC / km and a range of different conditions for specific layers of the atmosphere.
20.3.7 Standard temperature and pressure, STP, density of gases
States of matter, at STP (Standard Temperature and Pressure) (0oC and 1 atmosphere pressure):
solid (s), or liquid (l), or gas (g), or aqueous solution, (dissolved in water), (aq).
STP refers to the standard conditions used in calculations of the effects of changing temperature and pressure.
IUPAC, set the standard of STP in 1982 as temperature 273.15 K (0oC, 32oF), absolute pressure 100, 000 Pa (1bar, 0.98692 atm)
Since 1982, STP has been defined as a temperature of 273.15 K (0 °C, 32 °F) and an absolute pressure of exactly 105 Pa (100 kPa, 1 bar).
Formerly it was as follows, and is still commonly used as the SATP, standard ambient temperature and pressure:
T = 0oC or 273.15 K, P = 760 mm Hg, 1 atm, or 101, 325 pascals, Pa (101.325 Nm-2).
The combined gas equations can be used to find the volume of a gas at STP, i.e. at 0oC and 760 mm Hg pressure.
At STP, one mole of gas occupies 22.4 litres, L.
20.3.8 Universal gas equation, Ideal gas constant
The universal gas equation combines the three gas laws as PV = nRT, where P = absolute pressure in Pascals, Pa, V =
volume gas in cubic metres, m3, n = amount of gas in moles.
R is the universal gas constant (gas constant, molar gas constant, ideal gas constant, universal molar gas constant) =
8.314 JK-1 mol-1 (8.314510).
T = thermodynamic temperature, K (0oC + 273).
For the Boltzmann constant KB or K, N = number of particles, not moles, pV = NKBT
For specific gas constant, Rspecific = R / M (molar mass of the gas)
20.4.1 Thermodynamics, isothermal change and adiabatic change
"Air Investigations Lab Pack", clouds, refrigeration, air conditioning, (commercial)
See diagram 20.4.0: Thermodynamics, Isothermals and adiabatic expansion, Carnot cycle.
See diagram 20.4.1: Thermodynamics, Pistons.
See 2.0.5: Conic sections, hyperbola.
Thermodynamics is about how energy changes from one form to another, the direction of heat flow and how energy does work.
Isothermal means having the same temperature throughout a process.
Internal energy of the system
An isolated system contains a certain quantity of energy called the internal energy of the system = total kinetic energy and potential energy of all the atoms and molecules in the system.
Thermodynamics is about how energy changes from one form to another, the direction of heat flow and how energy does work that can be transferred as heat.
Internal energy does not include chemical energy or nuclear energy.
The value of the internal energy of a system can be changed by the following:
1. transfer of mass, 2. transfer of heat, 3. work done on or by the system.
In an isothermal change the temperature remains constant, and PV = a constant.
On a pressure / volume graph, an isothermal change is shown as a rectangular hyperbola.
20.4.1.1 First law of thermodynamics
Heat can be changed into mechanical energy and mechanical energy can be changed into heat energy, but the total energy of the system remains constant, i.e. the law of conservation of energy always holds true.
20.4.1.2 Second law of thermodynamics
Heat cannot pass from a body at lower temperature to a body at high temperature.
Heat always flows from hot bodies to cold bodies.
A machine unaided by an external agent cannot transfer heat from a body at lower temperature to a body at higher temperature.
Experiment
Ice cubes in boiling water
Heat a pot of water until it is boiling steadily.
Add several ice cubes to the pot.
The boiling action stops almost immediately as heat is transferred from the burner to the lower temperature ice rather than to the higher temperature water.
When all the ice is melted the boiling action starts again.
20.4.1.3 Third law of thermodynamics
The entropy of a substance approaches zero as is temperature approaches absolute zero.
Entropy measure the unavailability of the energy of a system to do work.
In any closed system an irreversible change is associated with an increase in entropy.
For an adiabatic process no heat transfer occurs and the entropy remains constant during the process.
Increase in entropy is another way of stating the second law of thermodynamics.
(Adiabatic means neither gain nor loss of heat.)
20.4.4 Adiabatic change, thermodynamics
Adiabatic means neither gain nor loss of heat.
In an adiabatic change, no heat is received or lost from the surroundings.
If a piston is raised in a cylinder containing a gas, the volume of the cylinder increases and the temperature of the gas falls as work is done against the rising piston.
In an adiabatic change, no heat is received from or lost to the surroundings, so there is a transfer of energy into or out of the system in the form of work only,
e.g. very rapid expansion of a gas or any process in an insulated container as in a piston.
Adiabatic change can occur in a container with thermally-insulated walls or where the change occurs in a very short period of time process happens in an extremely short time, so there is little heat exchange with the surroundings.
In diesel engines, adiabatic heating occurs during the compression stroke with compression ratios > 20: 1 to raise the temperature enough to ignite the fuel.
On a pressure / volume graph an adiabatic change is always steeper than a rectangular hyperbola, because adiabatic expansion is accompanied by a fall in temperature.
20.4.5 Adiabatic compression
It occurs when you compress a gas quickly, or insulate a gas from surroundings, the temperature rises, used in igniting the fuel in a diesel engine.
Experiment
Light a match head.
Push down hard on a piston in a close fitting cylinder to ignite a match head at the bottom of the cylinder.
Show the generation of heat during compression of air.
Place scrapings from a match head within the combustion chamber of a fire syringe kit.
Force the piston rapidly down to compress the air adiabatically and raise its temperature to approximately 1000 K to ignite the match head scrapings.
A small piece of tissue paper may be ignited instead of the match head scrapings, but the demonstration is difficult to do.
In a typical compression the volume decreases by a factor of 20.
Both nitrogen and oxygen, the primary constituents of air have adiabatic index, γ (gamma) =1.4. T = 994oK.
20.4.6 Adiabatic cooling
Adiabatic cooling occurs when a gas expands quickly, or when a gas is insulated from the surroundings.
The gas does work and the temperature drops, as in refrigeration.
At room temperature, all gases except hydrogen, helium and neon cool upon expansion by the Joule-Thomson process.
Clouds form when air rises and becomes saturated in response to adiabatic cooling.
Experiments
1. Pressurize a one gallon jar with a bicycle pump until the cork blows.
Measure the temperature of adiabatic heating and cooling.
Expansion of air in a cylinder moves a piston back and forth.
Use a thermocouple to measure the temperature of adiabatic heating and cooling.
2. Expansion chamber.
Make a temperature detector to insert into a flask that will be warmed
and cooled by compression and expansion of air in the flask.
3. Use a thermocouple to measure the temperature
change as cools on expansion and heats on expansion.
4. Release air is from a pneumatic tyre.
The outlet air is cooler than the tyre and the valve stem is be cold to the touch.
20.4.7 Adiabatic processes
In an isothermal change the temperature remains constant.
In an adiabatic change no heat is received from or lost to the surroundings, so there is a transfer of energy into or out of the system in the form of work only, e.g. very rapid expansion of a gas or any process in an insulated container, e.g. a piston.
Adiabatic change can occur in a container with thermally-insulated walls or where the change occurs in a very short period of time process happens in an extremely short time, so there is little heat exchange with the surroundings.
In diesel engines, adiabatic heating occurs during the compression stroke with compression ratios > 20: 1 to raise the temperature enough to ignite the fuel.
Experiments
1. Light a match head
Push down hard on a piston in a close fitting cylinder to ignite a match head at the bottom of the cylinder.
Show the generation of heat during compression of air.
Place scrapings from a match head within the combustion chamber of a fire syringe kit.
Force the piston rapidly down to compress the air adiabatically and raise its temperature to approximately 1000 K to ignite the match head scrapings.
A small piece of tissue paper may be ignited instead of the match head scrapings, but the demonstration is difficult to do.
In a typical compression the volume decreases by a factor of 20.
Both nitrogen and oxygen, the primary constituents of air have adiabatic index, γ (gamma) =1.4. T = 994oK.
2. Expansion cloud chamber
Put some smoke and alcohol in a stoppered flask and shake.
When the stopper is released a fog forms.
3. Apply pressure to a flask containing saturated water vapour then release the pressure.
No changes are observed.
Repeat the experiment by adding smoke to the water vapour.
On releasing the pressure a cloud forms in the flask.
4. Adiabatic cooling
* Pressurize a one gallon jar with a bicycle pump until the cork blows.
Measure the temperature of adiabatic heating and cooling.
* An air cylinder moves a piston back and forth.
Use a thermocouple to measure the temperature of adiabatic heating and cooling.
5. Expansion chamber
Make a temperature detector to insert into a flask that will be warmed and cooled by compression and expansion of air in the flask.
6. Joule-Kelvin coefficients
Use a thermocouple to measure the temperature change as cools on expansion and heats on expansion.
20.4.8 Cincinnati flask
See diagram 20.4.06: Cincinnati Flask.
Experiment
Heat water to boiling in a Cincinnati flask.
Remove heat, stopper and invert.
Boiling continues again as ice is added to the dimple in the flask.
The ultimate in discrepant events.
Show students that you can boil water by cooling it!
Fill the specially designed flask one-third full with water.
Bring the water to a boil and then remove the flask from the heat.
Place a rubber stopper fitted with a thermometer in the neck of the flask.
Invert the flask in a support ring and add crushed ice to the concave bottom of the flask.
The water now begins to rapidly boil again.
As the boiling continues, the temperature steadily drops to within 15-20 of room temperature.
All materials included are reusable.
20.4.9 Crookes' radiometer
See diagram 20.4.2: Crookes' radiometer.
The Crookes radiometer was invented by Sir William Crookes in 1873, but he mistakenly thought it measured the pressure of light.
It consists of a paddle wheel of vertical mica paddles with alternate surfaces white or black connected by a vertical spindle in a partly evacuated glass container.
It spins with the white sides approaching the source of radiation.
Increasing the radiation increases the speed of turning.
Maintaining the level of radiation, but decreasing the temperature reverses the direction of rotation.
Gas molecules move from the cold white side to the warmer black side causing the cold white side to move forward.
Gas molecules hitting the edges of the warmer sides bounce off the paddle with increasing speed causing the slight increase in temperature of the black sides, so these rebound molecules increase in speed.
Crookes' radiometer is still sold in novelty shops as a "light mill".
The radiometer, a sensitive detector of thermal radiation, has four light mica vanes black on one side and silver on the other, on a common pivot, so they can rotate freely in a horizontal plane.
The glass container is evacuated to a pressure of about 10 mm of mercury.
At this pressure there are still many gas molecules to interact with the vanes and their mean free path is large.
Experiment
This radiometer spins in the opposite direction from what theory predicts: the white side moves forward.
Turn on the light and watch it spin.
20.4.10 Expansion cloud chamber
See diagram 20.4.05: Expansion cloud chamber.
Experiments
1. Put some smoke and alcohol in a stoppered flask and shake.
When the stopper is released a fog forms.
Pressurize a jug of saturated water vapour with and without smoke particles.
When the pressure is released, a cloud will form in the jug
with the smoke particles.
2. Apply pressure to a flask containing saturated
water vapour then release the pressure.
No changes are observed.
Repeat the experiment by adding smoke to the water vapour.
On releasing the pressure a cloud forms in the flask.
20.4.11 Heat cycles, Carnot cycle, Stirling engine
Stirling Engine
The working of an ideal reversible engine is shown as the Carnot cycle.
A gas is contained in a cylinder with a conducting base and nonconducting walls and frictionless piston.
Stage 1: A constant heat source, temperature T1, heats the conducting base and the load on the piston is decreased.
Heat is taken in.
Isothermal expansion of the gas at temperature T1 occurs.
Stage 2: The heat source is removed, the conducting base of the cylinder is insulated and the load on the piston is decreased.
Adiabatic expansion of the gas occurs as the temperature of the gas falls to T2.
Work is done by the gas.
Stage 3: The conducting base of the cylinder is no longer insulated, it is heated by a constant heat source, temperature T2 and the load on the piston is increased.
Heat is given out.
Isothermal compression of the gas at temperature T2 occurs.
Work is being done on the gas.
Stage 4: The heat source is removed, the conducting base of the cylinder is insulated and the load on the piston is increased.
Adiabatic compression of the gas occurs until the temperature returns to T1.
Work is done on the gas
No engine can be more efficient than the theoretical reversible engine
working between the same temperature limits (T2 - T1).
Experiment
The hot air chamber of the Stirling engine is heated by an alcohol burner.
Light the burner.
If the engine is cold, it takes several minutes before it is hot enough to run.
Start the engine by turning the fly-wheel in the proper direction.
20.4.12 Heat transfer and laws of thermodynamics
1. Heat is a form of energy.
The unit of work and energy is the joule, J (i.e. newton.metre) (James Prescott Joule 1818 - 1889).
Heat transfer by conduction, convection and radiation, coefficient of expansion, the joule / calorie
* First law of thermodynamics
When other forms of energy are converted to heat or when heat is converted to other forms of energy there is no loss of total energy.
* Second law of thermodynamics
Heat always flows from hot bodies to cold bodies.
2. Heat is a form of energy measured in joules, J.
Heat transfer is the process of transfer of energy from an object to another, or from a part of an object to another one.
Heat energy can be transferred by conduction, convention, and radiation.
The natural flow of heat is from higher temperature towards lower temperature.
So heat energy spreads out from concentrations at high temperature.
When you apply heat to one end of a solid conductor, the particles at that end, e.g. atoms and molecules vibrate more rapidly.
This energy is passed from particle to particle through the material by conduction.
All metals are good conductors of heat, but many liquids and gases are poor conductors.
Liquids and gases can transfer heat by convection when hot fluid rises and is replaced by colder surrounding fluid.
Heat can be transferred through space as electromagnetic radiation.
Rough or black surfaces are good absorbers and good emitters of radiation whereas polished or white surfaces are not.
22.4.13 Heat transfer coefficient, 1 / thermal insulation
Heat transfer usually by convection between a fluid and a solid or by phase change h = q / A × δT
h = heat transfer coefficient, W / (m2K), watts per meter2 Kelvin
q = heat flow, in or out, J / S = W
A = surface area of heat transfer
δT = temperature difference of solid surface and fluid area (T1 -T2)
Heat transfer coefficients
Aluminium 237, Brass 110, Copper 398, Glass 0.96, Ice 2.18
Convective Heat Transfer Coefficients
Free Convection, Air : 5 - 25 W / (m2K)
Free Convection, Water: 20 - 100 W / (m2K)
Rate of heat transfer, R = K × area (T1-T2) / thickness
22.4.14 Heat transfer by Dewar flask (vacuum flask, "Thermos" flask)
See: Thermos flask, vacuum flask, Dewar flask, (Commercial).
See diagram 23.05: Dewar flask.
A Dewar flask, James Dewar, 1842-1927, Scotland, is a double-walled
vessel, the space between the walls being a near vacuum.
Heat cannot be conveyed through the two walls by conduction and the air cannot get into direct contact with the inner wall, so that heat is not conveyed away by convection.
The outer side and the inner side of the thin walls of the double glass
flask have silver surfaces to reduce the loss or gain of heat by
radiation.
The near vacuum between the double walls and the cork or rubber stopper
reduces conduction and prevent evaporation.
The cork or plastic stopper prevents heat transfer by convection at
the top and the loss of heat by conduction through the bad
conductor cork or plastic.
Also, the case reduces convection.
Experiment
1. Draw a cooling graph for a vacuum flask.
Almost fill a vacuum flask with boiling water.
Note the time and temperature of the hot water every half hour until
it is cold.
Draw a graph of your results.
When the contents are hot, heat is lost at a greater rate so the temperature-time
graph is a swooping curve rather than a straight line.
Thus heat losses are faster when the difference in temperature between
the hot object and the surroundings is greatest.
Put an equal quantity of ice cold water in a second vacuum flask it
and graph the rate of warming.
A vacuum flask keeps heat from getting out from hot things and stops
heat from getting in to cold contents.
2. Keep a broken thermos flask in a safe place in the laboratory store room so that the walls of the thermos flask may be examined.
22.4.15 Reduce heat loss with heat insulation materials
Use identical cans of water, one wrapped with insulation.
Do NOT use asbestos or any product containing asbestos.
1. Use four large tin cans of equal size and four smaller tin cans of equal size.
Inside the first large can put a small can on two corks in a large can as the control.
Select types of insulating material, e.g. sawdust, cork pieces, newspaper, plastic.
Put a small can inside each large can.
Pack one type of insulating material under and around each of the smaller cans.
Put a cardboard cover on each large can.
Make a hole in each cover for a thermometer.
Fill each small can to the same depth with water that is nearly boiling.
Record the initial temperature of the water in each can.
Record the temperature of the water in each can at five minute intervals.
Draw cooling curve graphs by plotting temperature against time for each tin can.
Note which material is the best insulator.
| Material |
Initial temp. |
after 5 minutes |
after10 minutes |
after15 minutes |
after 20 minutes |
| control (air) |
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| sawdust |
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| cork |
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| newspaper |
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| plastic |
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2. Test the heat insulation properties of common materials
See diagram 23.1.5: Four big and four small beakers.
Use 4 big beakers and 4 small beakers.
Put a small beaker into each big beaker.
Put 3 kinds of heat insulators, e.g. polyester plastic, paper and shredded wood, in the space between a big beaker and a small beaker.
The fourth large beaker contains a small plastic stopper and the small beaker so the beakers are separated mostly by air as a control.
Pour the same volume of hot water into each small beaker.
Put a thermometer in each small beaker.
Record the temperature in each small beaker at one minute intervals for 10 minutes.
Plot a graph of temperature against time on one sheet of graph paper for all beakers.
3. Be careful! To avoid scalding, prepare a sponge to absorb overflowing water.
Use 5 plastic fruit juice bottles.
Punch a hole in each lid to insert a thermometer through it.
Select a heat insulation material, e.g. paper, cloth, plastic cloth, sponge.
Cut the materials into a shape that you can wrap around the bottles.
Pour hot water into the bottles, close the lids tightly, and insert the thermometers.
Wrap bottles with three layers of heat insulation materials and attach the outer layers with adhesive tape.
Record the temperature in each bottle in equal time intervals.
Draw a temperature / time graph.
The horizontal axis is for time.
The vertical axis is for temperature of water.
20.4.16 Isothermal change, internal energy, thermodynamics, Diesel engine
Isothermal change
Diesel engine 1. An isolated system contains a certain quantity of energy called the internal energy of the system = total kinetic energy and potential energy of all the atoms and molecules in the system that can be transferred as heat.
Internal energy does not include chemical energy or nuclear energy.
The value of the internal energy of a system can be changed by the following: 1. transfer of mass, 2. transfer of heat, 3. work done on or by the system.
In an isothermal change, the temperature remains constant, and PV = a constant.
On a pressure / volume graph an isothermal change is shown as a rectangular hyperbola.
(Isothermal means having the same temperature throughout a process.
The temperature remains constant.)
2. Isothermal change and adiabatic change, Diesel engine
"Fire Syringe", adiabatic heating in diesel engine, starts fire with compressed air (toy product)
For an ideal gas, i.e. no attractive forces between its molecules, the volume decreases with temperature down to -273oC, called "absolute zero" or 0o Kelvin, K.
For monatomic gases, the molar heat capacity cp = 12.5 joule / mole K, so you need 12.5 joules to raise the temperature of a mole of a monatomic gas by 1K.
In an isothermal change the temperature remains constant.
In an adiabatic change no heat is received from or lost to the surroundings.
Adiabatic expansion occurs when a gas expands quickly, or when a gas
is insulated from the surroundings.
The gas does work and the temperature drops, as in refrigeration.
Adiabatic compression occurs when you compress a gas quickly, or insulate
a gas from the surroundings.
The temperature rises, as is used in igniting the fuel in a diesel engine.
20.4.17 Maxwell's Demon
See diagram 20.4.14: Maxwell's demon opens the door for the fast molecules O to pass.
Maxwell's Demon, an idea created by James Clerk Maxwell (1831-1874), seems to contradict the second law of thermodynamics.
A box filled with a gas is separated into two halves by a partition with a door in it.
A "demon" opens the door to allow faster than average gas molecules to pass to the left side of the box and slower than average molecules to pass to the right side of the box.
So the left side of the box becomes hotter than the right side and this separation would allow a heat engine to run as heat flows from the hot side to the cold side.
Experiment
The Maxwell Demon demonstrator consists of a sealed flask containing 5 white and 5 black spheres.
The balls in the stem of the flask are in the unmixed condition.
Inverting the flask mixes the balls.
If the flask is to be inverted again, it is unlikely that the original arrangement of balls would return.
With 5 black and 5 white balls the odds against reproducing the initial array are 252 to 1.
Rotate the flask in a stem up position to impart a swirling motion to the spheres.
Turn the flask to the stem down position and slowly reduce the rotation rate and the balls will spiral back into the neck.
With a little practice the initial array is restored.
The action of the "demon" is caused by the fact that the black spheres are constructed so as to rapidly lose momentum.
20.5.1 Expansion of air in a flask or bottle
See diagram 20.1.3: Heat flask with hands.
1. Air is trapped in the flask by means of a small bead of oil in the glass tubing.
Gentle heating with the hand will produce a sufficient temperature rise to cause the oil drop to move up the tubing.
Then plunge the flask first into cold and then into warm (not hot) water.
In place of the flasks, glass test-tubes and corks with capillary tubing could be used.
2. Use a flask fitted with a one-hole stopper and glass tube that extends into he flask.
Put a small amount of oil in the glass tube to trap air in the flask.
Hold the flask in your hands.
The oil moves up the tube, because the heat from your hands causes the trapped air to expand.
If you look carefully note that the oil first moves down, because the heat from your hands first causes the glass of the flask to expand.
When you cool the flask under the tap, the oil moves down.
3. Empty a cold bottle of soft drink, put a small coin on top of the opening, and firmly grab the bottle with both hands.
The coin will start to dance up and down, because the cold air in the bottle has become warmed by the hands and it expands to push up the coin.
When the coin rises some air escapes so the con drops to later rise again with further expansion of air.
20.5.2 Expansion of air in inverted test-tube
See diagram 2.111: Expansion of air in inverted test-tube.
Fit a hard glass test-tube with a one-hole stopper that has a length of glass tubing through it.
Invert the test-tube so that the end of the tubing is in a container of water.
Clamp the test-tube in an inverted position so that you can heat it with a burner.
Heat the test-tube and note the bubbles from the end of the tubing in the container of water.
Heat has caused the air to expand.
Cool the test-tube by pouring cold water over it.
Water moves up the glass tubing as the cooling air contracts.
20.5.3 Expansion of air in a balloon
1. Fit a toy balloon over the neck of a small flask.
Put the flask in a container of water.
Heat the water.
The balloon expands as the heated air in the flask expands.
2. Partially inflate a balloon and tie the neck tightly.
Leave it in a warm place or in the sunlight.
The balloon becomes fully inflated as the air inside expands when heated.