Heat, Thermometers, Bunsen burner, Specific heat.

School Science Lessons
(UNPh22)
2024-09-17a

Heat
Contents
22.1.0 Bunsen burners
22.3.0 Construct heat sources
22.2.0 Heat energy, thermal properties of matter
22.4.0 Mechanical equivalent of heat
22.5.0 Heat measuring devices
22.7.0 Thermometers and temperature

22.1.0 Bunsen burners
Experiments
22.1.1 Bunsen burner flame
22.1.2 Bunsen burner flame and candle flame
22.1.3 Bunsen burner flame can melt copper wire
22.1.4 Bunsen burner gas
22.1.5 Bunsen burner safety
22.1.6 Compare flame of a Bunsen burner and a candle.
22.1.7 Construct a Bunsen burner
22.1.8 Hottest part of the Bunsen burner flame
22.1.9 Lighting a Bunsen burner
22.1.10 Study a Bunsen burner
22.1.11 Study a Bunsen burner flame
22.1.6 Compare flame of a Bunsen burner and a candle.

22.3.0 Construct heat sources
22.3.1 Construct an air oven
22.3.2 Construct a beverage can charcoal burner
22.3.3 Construct a metal can heater

22.2.0 Heat energy, thermal properties of matter
22.2.1 Calorie
22.2.2 Heat as a form of energy
22.2.3 Heat and temperature
22.2.4 Joule, J
Experiments
22.2.6 Heat absorbed depends on mass
22.2.7 Heat dissipation
22.2.8 Heat has no weight
22.2.9 Heat solid sphere and hollow sphere
22.2.11 Movement in hot water and cold water
22.2.12 Pressure affects the boiling point

22.4.0 Mechanical equivalent of heat, Joule's equivalent
22.4.1 Mechanical equivalent of heat, Joule's equivalentt
22.4.2 Mechanical equivalent of heat, J, by an electrical method
Experiments
22.4.3 Dropping lead shot
22.4.4 Flint and steel
22.4.5 Friction ignition, bow and stick, fire maker, drill and dowel
22.4.6 Hammer on lead
22.4.7 Heat by bending
22.4.8 Waterfall

22.5.0 Heat measuring devices
22.5.1 Heat capacity (thermal capacity) C, of a calorimeter, (Experiment)
22.5.2 Molar heat capacity, C_m
22.5.3 Specific heat capacity, shc
Experiments
22.5.4 Candle burns below water level
22.5.5 Heat capacity (water equivalent)
22.5.6 Heat of combustion, bomb calorimeter, energy values of food
22.5.7 Heat capacity of a metal, e.g. lead shot
22.5.8 Heat capacity of a metal with a Styrofoam cup calorimeter
22.5.9 Melting wax calorimeter
22.5.10 Mix heated water
22.5.11 Rate of heat transfer
22.5.12 Simple calorimeter
22.5.13 Specific heat calorimeter
22.5.14 Specific heat of aluminium by electrical method
22.5.15 Specific heat of canola oil after calibrating the setup with water.
22.5.16 Specific heat of water by electrical method
22.5.17 Teapot physics
22.5.18 Ice calorimeter

22.7.0 Thermometers and temperature
22.7.1 Beurer, express thermometer
22.7.2 Thermochromic substances
22.7.3 Thermometers
Experiments
22.7.4 Air thermometer
22.7.5 Expansion of liquid in a thermometer
22.7.6 Galileo's thermometer
22.7.7 Make a spirit thermometer
22.7.8 Temperature rise and quantity of heat intake
22.7.9 Temperature rise calculation
22.7.10 Temperature sense
22.7.11 Test a liquid in glass thermometer
22.7.12 Test a thermometer
22.7.13 Thermometer, 0^oC to 100^oC range
22.7.14 Thermometer, make a thermometer
22.7.15 Thermometer test, calibrate a thermometer
22.7.16 Temperature of water at maximum density, 4^oC

22.1.1 Bunsen burner flame
See diagram 3.1.1: Bunsen burner.
See diagram 3.1.2.2: Bunsen burner flame.
See diagram 3.1.2.3: Burning the gas in a cone of flame.
1. A flame is the region where combustion occurs.
The colour of the flame depends on the temperature and the substance burning.
Hydrocarbon flames are either blue or yellow.
A blue flame is a not luminous and occurs, because of complete burning of hydrocarbons with plenty of oxygen gas.
The flame does not leave any residue or any other gases.
A yellow flame occurs when there is insufficient oxygen gas.
It is a luminous flame.
The temperature is lower than the blue flame and leaves black soot and other residues.
2. A Bunsen burner has a gas jet at the base that draws air in through the air hole, because of the Bernoulli effect.
It was invented by German chemist Robert Bunsen or his assistant, to improve the efficiency of combustion by combining flammable gas from a jet with air before ignition to give a very hot flame.
This premixed flame has a different structure to the diffusion flame of the candle.
The blue part of the flame in inside the flame.
The flame is conical, because of the shape of the rim of the burner.
The premixed flame burns efficiently with not much yellow flame or production of soot.
3. Control the amount of air by opening and closing the air hole.
Close the air hole.
Turn on the gas.
Light the gas.
The flame is yellow.
Hold a piece of wire in different parts of the flame to discover which part is the hottest.
Hold a splint in different parts of the flame.
The splint can be set alight in all positions in the yellow flame.
Hold a test tube just above the flame.
Carbon is deposited on the glass.
Test whether the unburned carbon causes the yellow colour of the flame by sprinkling powdered charcoal (carbon) into the flame.
Open the air hole.
Mixing the air with the gas allows the gas to burn more rapidly and completely.
The flame has an outer cone of mainly blue flame and a colourless inner cone.
The outer cone has a thin colourless area outside the blue flame.
Hold a splint so that it passes through the inner cone of the flame.
The middle of the splint does not burn, because the inner cone is mainly unburned gas.
Hold a piece of wire in different parts of the flame to discover which part of the flame is the hottest.
The tip of the colourless part around the outer cone is the hottest part of the flame.
The temperature of different parts of a Bunsen burner flame can be measured with a thermocouple.
The hottest part is close to the centre, just above the inner blue cone where oxygen is fed to the heart of the flame through the Bunsen burner chimney.
If the air hole is closed, the flame colour is similar to a candle flame.
4. Test the inner cone of the flame
Put one end of a piece of glass tubing in the inner cone of the flame.
Ignite the unburnt gases that come out of the other end of the tube.
Bunsen burners on sale:
Bunsen burner, suits PVC tube
Bunsen burner, L.P. gas
Bunsen burner, natural gas
Semi micro Bunsen burner with flame retention collar, rotatable air regulator, and gas inlet tube, 100 mm height, natural gas

22.1.2 Bunsen burner flame and candle flame
See diagram 3.1.2.4: Bunsen burner flame, Candle flame.
See diagram 3.1.2.6: Candle burner.
Just above the wick of a burning candle is a dark region of unburned gas.
Above and around it is a yellow region containing incandescent particles of carbon undergoing combustion to form carbon dioxide.
Put the candle flame under an evaporating basin.
Note the deposits of carbon, soot, because of insufficient oxygen to complete combustion.

22.1.3 Bunsen burner flame can melt copper wire
See diagram 3.1.4.1: Bunsen burner flame.
Bunsen burner flame can melt lead, MP= 327^oC, and zinc, MP = 419^oC.
Many people think the temperature of a flame cannot exceed 500^oC.
Copper melts at 1 085^oC.
It is not usually considered possible to melt copper with a Bunsen burner flame.
The temperature of different parts of a non-luminous flame, a blue flame, with the air holes fully open, vary.
The hottest part of the flame is at the tip of the central cone.
The central core of the flame contains a mixture of unburnt gas with air.
The intense blue region surrounding the central core is the main zone of combustion
This is where gaseous hydrocarbon fuel reacts with the oxygen to form short-lived gases.
The lighter blue outer flame is where these short-lived gases are completely oxidized to carbon dioxide and water.
Copper metal reacts readily with oxygen from air when heated strongly, forming a coating of black copper oxide, CuO.
Under reducing conditions, black copper oxide is reduced readily to metallic copper.
When heated strongly, but below its melting temperature, copper glows with a bright red heat.
To show that the maximum temperature is reached, use pieces of copper wire with three different thickness found by stripping the insulation from electrical flex.
1. Light a Bunsen burner, turn the flame to maximum height, and open the air holes so that the flame is completely blue.
Hold a piece of thick copper wire with tongs and probe the flame with the wire:
1. starting from the bottom 2. around the sides and tip of the central cone, and 3. around the outer blue flame.
At each place, record the appearances of the copper, e.g. black, orange, red hot, or tending to melt.
2. Repeat the process with a thinner piece of copper, then with a very thin piece of copper wire.
3. Reduce the flow of gas and repeat the procedures with a smaller blue flame.
The flame has six separate zones: Zone 1 is the core of unburnt gas and air at the base of the flame.
Zone 2 is the bright blue region of intense combustion surrounding the core:
zone 2A is the tip of the central core, zone 2B is the region at the sides of the central core.
Zone 3 is the outer region of the flame where combustion becomes complete, zone 3A is at the top of the flame, zone 3B is the outer part of the sides of the flame.
Zone 4 is the region just outside the flame.
At each zone observe the appearance of the copper, e.g. black, orange, red hot, tending to melt, three different thickness of wire.
You can melt fine copper wire in the flame, but not thick copper wire.

22.1.4 Bunsen burner gas
See diagram 3.1.4.1: Bunsen burner flame.
See diagram 3.1.4.2: Bunsen burners.
Bunsen burner gas is usually natural gas, i.e. mostly methane CH4, or LPG, bottled gas, mostly propane, C3H8.
In a laboratory, the pilot light should burn with a 90% blue flame.
If the flame is yellow, the gas may be contaminated with condensates.
Do not use such a gas, but immediately inform the local gas authority.
Previously, laboratories used town gas, based on coal gas, containing equal volumes of methane and hydrogen, carbon monoxide, and hydrogen sulfide safety smell.
Heating values of fuels: town gas 88 MJ / kg, natural gas 55.6 MJ / kg, LPG gas 49 MJ / kg, diesel fuel oil 38 MJ / L,
kerosene 36.7 MJ / L, coke or coal 27 MJ / kg, dry split wood 12.5 MJ / kg.
CH4 (g) + 2O2 (g) --> CO2 (g) + 2H2O (g) + heat

22.1.5 Bunsen burner safety
See diagram 3.1.4.2: Bunsen burners.
Be careful! Do not turn the gas on without lighting the Bunsen burner.
Gas forms an explosive mixture in air.
1. Combustion is the burning in oxygen of a substance to produce heat and sometimes light energy.
A flame appears during combustion when a gas has such a high temperature that it emits heat and light.
A flame appears only where the burning gas and oxygen are in contact.
2. The Bunsen burner consists of a: 1. barrel, 2. an air regulator, (sleeve with a hole), 3. jet, (air mixture needle valve), 4. base, 5. gas inlet opening.
3. Adjust the flame by opening or closing the gas tap.
When the air regulator is open, the gas burns with a noisy blue flame that may be nearly invisible in strong light.
If the flame rises up from the burner, turn down the gas supply.
4. When not using the Bunsen burner, either turn off the gas or close the air regulator to give a safety flame.
The flame is yellow, because of the incandescence of carbon particles.
It is not as hot as the blue flame and leaves black soot deposits on glassware.
5. Regularly inspect gas fittings on the benches and hoses connecting Bunsen burners to gas turrets to make sure that connections are free of leaks.
With paired gas outlets make sure that only the gas tap connected to the Bunsen burner is turned on.
6. Tests for leaks by dipping the part in soapy water.
Be careful! Do not use a lighted match.
7. Heat flammable liquids in water baths using electrical hot plates, not Bunsen burners.
8. Use the Bunsen burner only in a draught free area.
Allow the Bunsen burner to cool before you move or store it.
9. Do not heat low melting point objects, e.g. plastics, solder, lead, over the barrel of the burner.
Melted pieces may fall inside the barrel.
Hold the burner at an angle.
If a match is blown out, turn gas off, then light the Bunsen burner again.
10. Turn the gas off first at the gas tap, then at the cylinder or main supply tap.
When not in use, turn off Bunsen burners to limit the production of carbon monoxide.

22.1.6 Compare flame of a Bunsen burner and a candle.
See diagram 3.1.2.4: Bunsen burner flame, Candle flame.
1. Bunsen burner flame
A flame is the region where combustion occurs.
The colour of the flame depends on the temperature and the substance burning.
Hydrocarbon flames are either blue or yellow.
A blue flame is not luminous.
It occurs, because of complete burning of hydrocarbons with plenty of oxygen gas.
The blue flame does not leave any residue or any other gases.
A yellow flame occurs when there is insufficient oxygen gas.
It is a luminous flame.
The temperature is lower than the blue flame and it leaves black soot and other residues.
2. Candle flame A candle contains wax made from petrochemicals.
The wick is lighted, and the flame melts the wax.
The evaporated wax rises and catches fire.
As the vapours rise higher, they stay longer in the hot regions of the flame and start burning completely with oxygen gas.
A candle flame has three regions.
The inner zone appears black, contains unburned wax vapours and is the least hot region of the candle.
The middle zone is where the wax vapours start burning giving a yellowish flame of partially burnt gases.
This is because of insufficient gases for complete combustion.
The flame in the luminous region, but not very hot.
The outer zone is where the wax vapours have enough oxygen gas to burn completely.
The flame appears blue and the temperature is very high, up to 1 100^oC or more.

22.1.7 Construct a Bunsen burner
See diagram 3.1.4.2: Bunsen burners.
Construct a Bunsen burner from scrap.
Use brass tubes about 1 cm in diameter.
Melt pieces of scrap lead, then pour into a shoe polish can to form a heavy base for the Bunsen burner.
Drill holes vertically and horizontally to fit the brass tubes.
The brass tubes must be slightly tapered and hammered into the lead.
The gas supply tube should extend about 2 cm into the base, but the jet tube must just enter the horizontal hole.
When the tubes have has been tested for size, fill it with a lead plug, cast in the tube.
Pour the lead in around a centrally-placed greasy sewing needle which, when extracted, leaves the jet hole.
During the casting, hold the tube in a shallow hole drilled in a wooden block with the needle precisely centred in the tube.
Use brass tubes of a suitable diameter for the barrel and the brass collar in which matching air holes are to be cut.
A tin tube may be substituted for a brass tube.
If the barrel is not already a good fit on the tube, seal it in position with epoxy glue.
To make the matching air holes, first to run the tubes on to tapered dowel rods held in a vice.
Flatten the tubes with a file, then use a 0.5 cm drill.
Trim the holes to shape with a round file, using it to smooth the inner surface of the collar so that it turns easily on the barrel.
Solder a ring of copper wire on the barrel, just above the collar.
This prevents loss of the air adjuster.
With double air holes, i.e. drilled right through the tubes, overdoing the air supply may be possible.
Then the burner may "strike back" in use.
It should not be allowed to burn for long at the lead jet or this may soften and close the hole.

22.1.8 Hottest part of the Bunsen burner flame
See diagram 3.1.4.1: Bunsen burner flame.
Add 3 cm of water and small boiling chips to 3 test tubes.
Measure the time taken to boil the water when the bottom of a test tube is held at the top of:
1. yellow flame (air hole closed), 2. non-luminous (dark blue) flame (air hole open), 3. light blue flame (air hole open).
The test tube held at 3. boils the soonest.

22.1.9 Lighting a Bunsen burner
1. Close the air regulator, light a match, hold the match flame at the side of the barrel opening, turn the gas tap on, raise the match flame to light the gas.
The gas burns with a visible yellow flame, a quiet safety flame.
Hold a test-tube just above the flame.
Note the carbon (soot, carbon black) that deposits on the glass.
To test whether unburned carbon gives the yellow colour to the flame, sprinkle powdered charcoal on the flame and compare the yellow colours.
2. Start to open the air regulator until the gas burns with a medium blue flame with a light blue inner cone and a pale violet outer flame with a bushy appearance.
The flame has an outer oxidizing zone where combustion is complete, a middle reducing zone, and an inner unburned gases zone surrounded by a blue cone.
This flame is the most useful for heating.
Fully open the gas regulator until you get a roaring blue flame with a light blue triangle in the centre of the blue cone.
3. Open the air regulator.
Keep turning down the gas supply.
The gas "blows back", "strikes back".
The gas is burning inside the barrel.
Turn the gas fully on and strike the gas supply rubber tube with a sharp blow from the side of your hand.
If the flame does not reappear, immediately turn the gas off and leave to cool, because the barrel may be hot.
Then light the Bunsen burner again.

22.1.10 Study a Bunsen burner
See diagram 2.0.0: Bunsen burner flame.
See diagram 2.1.1: Bunsen flame and candle flame.
See diagram 3.2.0.1: Alcohol lamp, spirit burner.
See diagram 3.1.2.3: Burn gas from the cone of the flame.
Find the hottest part of the Bunsen flame.
1. Close the air hole and turn the gas tap full on.
Light the gas and hold a piece of wire in different parts of the flame, moving it from the bottom to the top to find the hottest part of the flame.
Open the air hole.
Hold the wire in the flame, moving from the bottom to the top to find the hottest part in this flame.
Compare the two flames to find which has the hottest point.
2. Close the air hole.
Hold a test-tube with its bottom end just above the flame.
Carbon may be deposited on the glass.
Test whether unburned carbon which gives the yellow colour to the flame by sprinkling powdered charcoal into the flame to see whether it gives the same effect.
3. Open the air hole again.
Note whether carbon id deposited on a test-tube held in this flame.
Air mixing with the gas helps it to burn more rapidly and efficiently.
4. Find what is in the cooler inner cone by holding a splint of wood in the flame so that it passes through the inner cone.
Note which part of the splint burns.
5. Place a tube with one end in the inner cone as shown in the diagram.
The gas that comes out of the other end of the tube can be ignited.
6. Investigate a candle flame and the flame of a spirit lamp in a similar way.
Find the hottest part of the flame.
Look for unburned carbon particles in the flame.
Look for an inner cone of unburned gases.

22.1.11 Study a Bunsen burner flame
See diagram 3.1.2.2: Bunsen burner flame.
See diagram 3.1.2.3: Burn gas from the cone of the flame.
1. Hold the end of a glass tube in the centre of the cone.
You can light the gas coming out of the other end of the glass tube.
2. Hold a piece of wire in different parts of each kind of flame, moving it from the bottom to the top.
Find the hottest flame and the hottest place in each flame with a piece of nichrome wire or iron wire stuck into a cork for a handle.
The approximate temperatures and colours for the wire are as follows:
2.1 <500^oC, wire gives no light, flame is non-luminous,
2.2 500^oC to 950^oC, wire becomes red, then dark red, then bright red (red-hot)
2.3 950^oC to 1350^oC, wire becomes yellow-red then becomes white,
2.4 >1350^oC, wire becomes white (white-hot).
The safety flame has a similar temperature in different parts about 300^oC.
It is never used for heating.
The medium blue flame has the hottest point at the tip of the blue cone at about 500^oC.
The roaring blue flame has the hottest point at the tip of the cone at about 700^oC.
3. Close the air regulator.
Use a wood splint or a taper to test that parts of the flame support ignition.
The wood splint match is set alight in all positions in the yellow flame where no air mixes with the gas.
Repeat the experiment with the air regulator open.
A cone of mixed air and gas exists in the centre of the cone where the gas is not burning.
4. Turn off the gas.
Push a pin at right angles through a match just below the chemical on the end of the match.
Use the pin to hang the match in the barrel with the chemical end just above the rim.
Open the air regulator and light the gas again.
The match does not ignite inside the cone.
Move the match to the outer cone of the blue flame.
The match ignites.
5. Close the air regulator and light the gas.
Hold a piece of copper wire gauze with tongs 3 cm above the top of the barrel.
Hold a lighted match above the gauze.
The gas ignites above the gauze.
Lower the gauze until the flame passes through it.
Repeat the experiment with an open air regulator.
Light the gas and lower a copper wire gauze down on the flame.
The flame remains below the wire gauze as the gauze becomes red-hot.
Heat is removed from the gas air mixture by the copper gauze.

22.2.1 Calorie
The CGS (cgs) unit, the calorie, is the amount of heat required to raise the temperature of 1 gram of water by 1^oC at 15^oC (room temperature), i.e. 1 K.
Nowadays the SI unit the joule, J, is used.
1 calorie (cal) = 4.184 J, commonly, 4.2 joules.
1 joule = 10⁷ ergs = 0.2388 calorie.
Nowadays the SI unit the joule, J, is used.
You may see "kilocalories", 1000 calories, in nutritional information about weight loss.
In some "calorie counter" books, 1000 calories is a "Calorie", so in their tables 1 "Calorie" = 4.2 kilojoules.
Experiment
Suspend a metal can containing 50 mL water and a thermometer over a small Bunsen burner flame or a candle.
Record the initial temperature.
Heat it for two minutes, constantly stirring, and record the final temperature in degrees Celsius, ^oC.
Empty the water and repeat the experiment with 100, 150, 200 mL water, using the same flame.
Assume 1 mL (1 cm³) water = 1g.
Find the product of mass of water x by rise in temperature.
The same heat is given out by the flame to each mass of water (100, 150, 200 mL).
So a convenient unit of amount of heat would be the amount of heat absorbed by 1 g water rising in temperature by 1^o C.
This unit is the calorie or gram calorie.
Calorific value of fuel
"Calorific value" refers to the number of joules of energy released when 1 g of a fuel burns completely.
A 1^oC change in temperature of 1 mL of water requires 4.2 J.
Experiment
Hang a small metal can from a stand.
Pour 100 mL of cold water into the can.
Record the initial temperature, t₁.
Put a small piece of candle on a tin lid and weigh them, w₁.
Put the candle and tin lid under the can of water.
Light the candle.
Stir the water with a thermometer as the temperature rises.
When the temperature reaches 60^oC, t₂, blow out the flame.
Weigh the tin lid and candle again, w₂.
The calorific value of the fuel = 100 X 4.2 X (t₂- t₁) / (w₂- w₁).
However, the calorific value of fuels is usually expressed in megajoules per kilogram, MJ kg^-1, e.g. petrol 45, natural gas 40, coal 35, ethanol 30, dry wood 15.
Nutritional information usually expresses calorific value in kilojoules per gram, kJ g^-1, e.g. fat 40, cheese 30, sugar 16, potatoes 5.

22.2.2 Heat as a form of energy
1. The distinction between internal thermal energy, heat energy and temperature.
Heating is a process by which internal energy transfers occurs as the result of a temperature difference.
Heat is the transfer of internal energy from a hot object to a cold object.
Internal energy is the sum total of all the kinetic and potential energy of all the particles (atoms or molecules or ions) in an object.
2. Heat is a means by which energy can be transferred.
The numerical value for the heat is the amount of energy transferred.
Internal energy is the energy associated with the total kinetic and potential energy of all of the molecules of the object.
If energy is transferred to an object its internal energy rises.
The internal energy would also increase if work were done on the object.
The SI unit of heat energy is the joule (i.e. newton.metre).
The extent to which an object will transfer or absorb heat (hotness or coldness) is measured by temperature, related to the average kinetic energy of its molecules.
The process by which the energy is transferred as heat as one of the following three: convection conduction and radiation.
3. Heat is NOT the energy an object contains.
Heat is the energy that flows from a hot object to a cold object.
Heat energy can flow indefinitely between two objects provided a temperature difference is maintained between them.
The direction of heat flow between two objects depends on their temperatures, not on how much internal energy they each contain.
If a 50 g block of steel at 30^oC placed in contact with a 200 g block at 25^oC,
Then heat will transfer from the 50 g block to the 200 g block, even though the larger block contains much greater internal energy than does the smaller block.
4. The amount of internal energy contained in an object depends on the following:
* The mass of the object, the greater the mass, the greater is the number of particles, the greater the amount of internal energy,
* The temperature of the object, higher the temperature, the greater is the amount of internal energy,
* The type of substance making up the object.
5. Objects do not contain heat, they contain internal energy.
Internal energy is the sum total of all the kinetic and potential energy of all the particles (atoms or molecules or ions) in an object.
So heat is not the energy an object contains.
Heat is the energy that flows from a hot object to a cold object.
The direction of heat flow between two objects depends on their temperatures, not on how much internal energy they each contain.
The direction of heat flow is determined by the temperature difference, not on how much internal energy each block contains.
The term thermal energy refers to the internal energy present in a system by virtue of its temperature.

22.2.3 Heat and temperature
The joule, J, is the SI unit of work and energy.
A joule is equal to the amount of work done when the point of application of a force of one newton moves one metre in the direction of the force.
The c.g.s. unit, the calorie, is the amount of heat required to raise the temperature of 1 gram of water by 1^oC at 15^oC.
Nowadays the SI unit the joule, J, is used.
1 calorie (cal) = 4.184 J, commonly, 4.2 joules.
You may see "kilocalories", 1000 calories, in nutritional information about weight loss.
In some "calorie counter" books, 1000 calories is a "Calorie", so in their tables 1 "Calorie" = 4.2 kilojoules.
Experiment
1. The calorie
Suspend a metal can containing 50 mL water and a thermometer over a small Bunsen burner flame or a candle.
Record the initial temperature.
Heat it for two minutes, constantly stirring, and record the final temperature in degrees Celsius, ^oC.
Empty the water and repeat the experiment with 100, 150, 200 mL water, using the same flame.
Assume 1 mL (1 cm³) water = 1g.
Find the product of mass of water X by rise in temperature.
As the same heat is given out by the flame to each mass of water (100, 150, 200 mL) a convenient unit of amount of heat would be the amount of heat absorbed by 1 g water rising in temperature by 1^oC.
This unit is the calorie.
2. Rate of heat transfer
Use two identical thermos flasks containing:
(a) 300 g water at 40oC and
(b) 100 g ice + 200 g water at 0oC.
Which reaches room temperature 20^oC first?
The loss or gain of heat is greatest when the difference in temperature between the contents of the flask and the surroundings is greatest.
In thermos flask (a), the temperature difference with room temperature is continually decreasing from its original temperature of 40oC = 20oC.
In thermos flask (b) the temperature difference remains the same at 0oC when latent heat is absorbed to convert the ice to water.
More than half the total heat is received when the temperature difference remains at 20oC.
So (a) reaches room temperature first.
3. Add water at 90^oC to tea in a teapot and leave it to stand for 5 minutes.
If you want the tea to be as cool as possible before drinking it, do you add milk immediately after pouring out the tea or just before drinking it?
The rate of loss of heat depends on the temperature difference between the body and the surroundings.
If the milk were added immediately after pouring the temperature of the tea would fall and the rate of loss of heat would be less than if the milk had not been added.
So the milk should be added just before drinking the tea.

22.2.4 Joule, J
The joule, J, is the SI unit of work and energy.
(James Joule, 1818-1889, England)
A joule is equal to the amount of work done when the point of application of a force of one newton moves one metre in the direction of the force.
A joule may be referred to as a newton.metre.
1 Joule of work done = 1 newton force moves object through 1 metre
Work done = force distance in direction of force, W = Fs
Work and Energy
Work = force x distance (displacement), joule (newton.metre).
Work done on an object changes its energy that may be stored as potential energy or cause change in speed, kinetic energy.
When a wheel is moved by a force, the work done = displacement x component of the force in the direction of the displacement.

22.2.6 Heat absorbed depends on mass
Experiment
Place a large iron bolt and a small nail in a beaker filled with boiling water.
Fill other two beakers with equal masses of cold water at same temperature put a thermometer each.
Note that the amount of water is better to just immerse iron bolt.
To do this you can test a suitable the amount of water by bolt before experiment.
In the process of the experiment let large iron bolt and small nail are in boiling water for a while then take them out of boiling water.
Put large one into a beaker small one into another beaker rapidly.
Observe the temperature of water rises record the temperature in each beaker after temperature is stable.
The different temperature shows the different amount of heat they have.
The objects in different masses at the same temperature absorb an amount of heat that depends on their mass.

22.2.7 Heat dissipation
Experiment
Fill a small and a large round bottom flask with hot water at the same temperature.
Insert thermometers in both flasks and note the decrease in temperature as heat is dissipated.
Heat dissipation is a function of the area of the surface.
Heat content is a function of he volume of the unit.
The area increases as the square of the radius, but the volume increases as the cube of the radius.

22.2.8 Heat has no weight
The recognition of heat by people in history had experienced a long and tortuous way.
So called "heat mass" in general shows a unclear recognition to the character of heat.
Heat is a form of energy has no weight.
However, some students are probably think that heating a beaker will somehow make it lighter.
A few may believe that putting "heat in " somehow makes the beaker heavier! Hang a spoon in each end of the cross arm in a stand thread tie to the spoon must be strong enough.
Adjust the position of spoon to get them in exactly balance.
Record these two positions.
Then remove the spoons.
Heat one of the spoons by lifting it over an alcohol burner.
Put another spoon into cold storage in a refrigerator or cold water (if use refrigerator it is better not to put spoon into freezer).
Place the two spoons back to the original position exactly note to affirm the cooled spoon is dry in advance and the position must be the original one.
To ensure not to confuse them, it is better to select the two spoons that have different sizes and shapes or use different colour thread to tie them.
If every step is done the whole system is still in balance.
To avoid marking in cross arm to mark the original position the cross arm can be replaced by a meter.

22.2.9 Heat solid sphere and hollow sphere
Experiment
Put a solid metal sphere and a hollow metal sphere with the same external volume in a beaker of water and heat the water.
The expanded external volumes are still the same.
Put the two spheres in identical beakers and add the same volume of water heated to the same temperature.
The hollow sphere expands more than the solid sphere, because its mass is less.

22.2.11 Movement in hot water and cold water
Experiment
1. Prepare two sources of hot water and cold water with ice floating in it.
Hold a needle, point down, and, exactly vertical, above two paper cups.
Push the needle down through the bottom of the two paper caps to make exactly same size holes.
Pour the same volume of hot water or cold water into the paper cups.
Watch the drips of water from the holes in the paper cups.
The hot water leaks faster than the cold water.
The molecules of the hot water are moving faster than the molecules in the cold water so they can move past each other more quickly and more quickly move through the hole in the bottom of the paper cup.
2. Prepare two identical beakers containing the same volume of hot water or cold water.
Put two drops of ink or colouring, e.g. cochineal, into each beaker.
Observe the time taken for the drops to spread throughout the water.
The colouring in the hot water spreads faster, because the water molecules are moving faster around each other.

22.2.12 Pressure affects the boiling point
Use a flask with a one-hole stopper with a thermometer in it.
Fill a flask 3/4 full with water, add boiling chips, heat until boiling, read the temperature, about 100^oC.
Invert the flask and hold it under a cold stream of water.
The water boils again.
Read the temperature, less than 100^oC.
Again invert the flask and hold it under a cold stream of water.
The water boils again.
Read the temperature, less than before.
When the flask is cooled off with the cold water, water vapour in the flask above the water condenses on the colder surface of the flask, a partial
vacuum forms and the pressure inside the flask decreases.
Water boils at 100^oC at standard pressure, i.e. 760 mm mercury or the pressure at sea level.
If the pressure is lower, the boiling point is lower.
So making a good cup of tea on a high mountain is difficult, because the water boils at such a low temperature.

22.3.1 Construct an air oven
See diagram 23.33: Air oven.
A large metal can be used as an air oven.
A hole through the lid fitted with a cork holds a thermometer, and the saucer or dish rests on a wire gauze bridge placed inside the tin.

22.3.2 Construct a beverage can charcoal burner
See diagram 23.1.6: Simple heating devices.
Experiment
Prepare a large tin can of diameter at least 10 cm.
Draw 6 small windows of regular triangle uniformly on the side of the can and an angle of each triangle just up just towards the bottom of the can.
Each window locates at the centre of the side of the can.
For each triangular window cut off the two lines of the angle downward then bend the triangular sheet back.
The 6 triangular sheets form a stand in the can.
Charcoal may be placed on the stand.
Polish the windows with a file and drill some air holes.

22.3.3 Construct a metal can heater
See diagram 23.1.3: Metal can heater.
A heater can be made from an old oil tin.
Water is placed in the tin and heated from below.
Iron wire is wrapped round a test tube and twisted to form a handle.
The substance to be heated is placed in the test tube

22.4.1 Mechanical equivalent of heat, Joule's equivalent
Before SI units were adopted, Joule's equivalent (James Prescott Joule, 1818-1889), was the conversion coefficient between mechanical work and heat.
4.186 J = 1 calorie, i.e. J = 4.2 joules per calorie.
A typical experiment was to stir water with a paddle moved by a falling mass.
For example, a 1000 g mass falling 83.8 metres rotates a paddle in a calorimeter (paddle + calorimeter = 250g), specific heat 0.10, causes a rise in temperature of 0.28^oC.
Loss of energy by falling mass = mgh = 1000 x 980 x 8380 ergs = 8.21 x 10⁹ ergs.
Heat taken in by calorimeter and paddle = mass x specific heat x change in temperature = mst = (250 x 0.10 x 0.28) = 7 calories, + heat taken in by water = mst = (675 x 1 x 0.28) = 189 calories.
So total heat taken in = 7 + 189 = 196 calories.
W = JH, 8.21 x 10⁹ = J x 196, so J = 4.19 x 10⁷ ergs per calorie = 4.2 joules per calorie.
However, in SI units all forms of energy are expressed in joules, so J =1, and the specific heat capacity of water = 4.186 (4.2) kJkg^-1K^-1.
(heat capacity = joules per kelvin, specific heat capacity = joules per kilogram per degree kelvin).

22.4.2 Mechanical equivalent of heat, J, by an electrical method
See diagram 32.2.64: Mechanical equivalent of heat.
The heat energy expended by a current of I amps flowing under a potential difference V volts for t seconds = VIt / J,
where J = a constant called the mechanical equivalent of heat.
If heat lost from the calorimeter to the surroundings is small, the heat energy supplied by the coil = heat energy received by the calorimeter + contents.
Weigh the calorimeter + stirrer, m1.
Weigh the calorimeter + stirrer + enough water to cover the heating coil, m2.
Adjust the rheostat or power supply to a current of 3 amps.
Open the switch, stir the water and note the initial temperature T1^oC.
Close the switch and record the time.
Record the current I amps through the coil.
Record the potential difference V volts across the coil.
Allow the current to flow, still stirring, until the water temperature has risen 10^oC.
Open the switch and record the highest steady temperature, T2^oC.
Record the time of flow of the current, t seconds.
The specific heat of water = 4.2 kg^-1K^-1(or ^oC).
The specific heat calorimeter and stirrer, usually copper = s.
Calculate J using the following equation: Heat energy supplied = heat received by water + heat received by calorimeter and stirrer.
VIt / J = ([m2 -m1]s_water[T2 - T1]) + (m1 x s x [T2 -T1]).

22.4.3 Dropping lead shot
See diagram 22.10.2: Dropping lead shot.
Experiment
1. Stick the temperature probe (not thermometer!) into the bag of lead shot.
Note its initial temperature.
Drop the bag from the height of 2 meters 10-20 times.
Note the temperature of the lead again.
The temp should rise 2-3 degrees Celsius.
The specific heat of lead is 0.031 cal / (g C).
2. Put 1 kg of lead shot in a mailing tube, cardboard cylinder, invert 10 times and measure the rise in temperature rise.

22.4.4 Flint and steel
Experiment
Make sparks fly from flint rubbing against steel or a grindstone.

22.4.5 Friction ignition, bow and stick, fire maker, drill and dowel
Experiment
Make a fire with a bow and stick.
Hold an electric drill with a hardwood dowel in the chuck against a wood block.

22.4.6 Hammer on lead
See diagram 4.2: Transfer kinetic energy to heat energy.
Experiment
Hit a lead block with a heavy hammer and measure the temperature rise.

22.4.7 Heat by bending
Experiment
Keep bending an iron wire and measure the rise in temperature.

22.4.8 Waterfall
If 10 kg water falls 150 metres and all the energy converted to heat (silent waterfall!), and g = 9.8 m / s²
potential energy of water = mgh = 10 x 9.8 x 150 = 14700 joule,
J H = mc(t2-t1), where c = 4200 J / kg^oC = 10 x 4200 x (t2 - t1)
14700 = 42000 (t2-t1), so difference in temperature (t2 -t1) = 14700 / 42000 = 0.35^oC.

22.5.1 Heat capacity (thermal capacity) C, of a calorimeter
Heat capacity (thermal capacity) is the capacity of a body to store heat and is measured by the quantity of heat required to raise its temperature by one degree.
The SI unit is J kg^-1oC^-1.
A calorimeter is an insulated vessel usually containing water and used to measure the thermal quantities of a process, heat changes.
Heat capacity, C_p, is the ratio: heat given to an object / rise in temperature of the object and is expressed in joules per kelvin, J K^-1.
The heat capacity of the calorimeter itself is usually measured as the amount of heat needed to raise the temperature of the calorimeter by 1 K.
This value is usually found by experiment that involves transferring a known amount of heat into it and measuring its temperature increase.
This experiment is done before measuring the heat capacity of an unknown substance.
For example if the temperature of a calorimeter increases by 0.2 K when 8.0 J of electrical energy is used to heat it, the heat capacity of the calorimeter, C = 8.0 / 0.15 = 53.3 J / K.
The heat capacity of the calorimeter is usually compared with the heat capacity of an amount of water.
The heat capacity of one mole of water, C_{p, m} = 18 g mol^-1 x 1 cal g^-1K^-1 x 4.184 J cal^-1 = 75.312 J mol^-1K^-1at 25^oC.
Experiment
Measure the calorimeter constant, B
Use two Styrofoam cups, one inside the other, as a calorimeter.
Weigh the two Styrofoam cups empty, m1, add 70 mL of water at room temperature and weigh the two Styrofoam cups again, m1 + 70, and record the temperature of the water in the calorimeter, T1.
*** Boil water in a beaker, record the temperature, T2, and pour 30 mL of the boiling water into the calorimeter.
Be careful!
Gently stir the water with a thermometer with the bulb 3 cm above the bottom and record the highest temperature,
T3 and mass M1 + 100. B x (T3 - T1) + 70 x (T3 - T1) + 70 x (100 - T3) = 0

22.5.2 Molar heat capacity, C_m
The molar heat capacity is the amount of heat needed to increase the temperature of one mole of a substance by one degree kelvin.
It is expressed in joules per moles per degrees Kelvin.
For example, the molar heat capacity of lead is 26.65 joule per mole kelvin, which means that it takes 26.65 joules of heat to raise 1 mole of lead by 1 K.
The SI units for molar heat capacity are joules/(mole.kelvin).
The molar heat capacity of liquid water = 75.3538.
Dulong and Petit's law states that relative atomic mass x specific heat = constant, approximately 25 J mol^-1K^-1.
Molar heat capacity of a solid element, C_m = relative atomic mass x specific heat capacity = approximately 25 J mol^-1K^-1, (6.0 cal mol^-1K^-1) = 3R (where R = universal gas constant 8.314 J mol^-1K^-1) [3X 8.314 = 24.942].
If specific heat is expressed as gram heat capacity expressed as J g^-1K^-1, and specific heat of iron = 0.473 J g^-1K^-1, then the ratio: molar heat capacity / specific heat = 25 J mol^-1K^-1 / 0.473 J g^-1K^-1 = 52.85 = approximately 55.847, the molar mass of iron.
The molar heat capacity of non-metal compounds of metallic salts is about 60-80% of the molar capacity of heavy metals.
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.

22.5.3 Specific heat capacity, shc
Specific heat capacity is the heat required to raise the temperature of the unit mass of a given substance by a given amount, e.g. 1^oC, under specified conditions.
1. The amount of heat absorbed by object per unit of mass as the temperature rises is called the specific heat capacity, symbol "c" or "shc".
Specific heat is usually measured as the amount of heat required to raise the temperature of 1 kilogram of the substance by 1^oC.
Its unit is joule / kg^oC.
For water, shc = 4200 joule / kg^oC (or 4.1813 J g^-1K^-1 at 25^oC).
For copper, shc = 385 joule / kg^oC, [385 J / (kg K)].
These values are the same if specific heat capacity is expressed using the Celsius scale or the Kelvin scale, because the one degree interval has the same magnitude on each scale.
2. The heat required to raise the temperature of 15 kg mass of copper from 15^oC to 25^oC = mass x (t2 -t1) x specific heat = 15 kg x (25 - 15)^oC x 380 joule / kg^oC = 15 x 10 x 380 = 57000 joule, J (j).
3. Specific heat is the character of the material itself.
The c values of a substance may be different in different states.
The c values of the same gases are quoted as specific heat at constant volume (c_v) for when only its internal energy is increased, or specific heat at constant pressure (c_p), which requires more heat, because the gas expands.
For solids and liquids, the difference between specific heat values is very small.
4. Specific heat capacity is the ratio of the thermal capacity of the substance to the thermal capacity of water at 15^oC so is numerically equal to heat capacity (thermal capacity) with the SI unit J kg^-1oC^-1.

22.5.4 Candle burns below water level
Experiment
Cut a candle to be slightly longer than the depth of a large beaker.
Melt the base of the cut candle and fix it to the base of the beaker.
Fill the beaker with water up to the rim of the candle.
Light the candle and observe it burning down to below the level of the water as it forms a wax wall around the flame.
Water has a high heat capacity that allows it to absorb heat from the candle wax, so it does not melt and evaporate, but instead forms a wall around the flame.

22.5.5 Heat capacity (water equivalent)
The heat capacity (water equivalent) of a body is the quantity of heat the body absorbs when its temperature is raised through one degree Celsius.
This quantity may also be defined as the mass of water that requires the same quantity of heat energy to raise its temperature through one degree Celsius as the body itself requires.
So the water equivalent of 1 kg of iron is 1 kg.
If m = mass of the body, and s = specific heat of the material in the body, then heat capacity (water equivalent) = ms.
The heat capacity of water at 100^oC is 4.22 J / gK
Experiment
Put 1 litre of water in a beaker and 1 kg aluminium + water in another beaker and heat on the same hot plate then measure the temperature in each beaker.
Heat two beakers one with 1 Kg water and the other with 5 Kg water and 5 Kg lead at the same rate.
Heat two beakers on a single hot plate each contains the same mass of either water or oil, water and oil.
Heat an iron plate and a beaker of water with the same mass on identical Bunsen burners then dip your hand in the water and sprinkle it on the iron plate where it will sile.

22.5.6 Heat of combustion, bomb calorimeter, energy values of food
See diagram 23.5.7: Bomb calorimeter.
Determine the specific heat capacity of canola oil after calibrating the setup with water.
Use a bomb calorimeter to show heating value of foods and fuel.
dH = heat energy released at constant pressure, dE = heat energy released at constant volume, dH = dE + [d(PV)] = dE + dn RT, (from ideal gas laws).
Measuring heat produced at constant volume, qv = C dT (temperature change).
Here qv = change in internal energy dE, so dE = qv = CdT.
The heat capacity of a calorimeter can be calculated by burning a known weight of a standard substance, e.g. benzoic acid, dH = -3227 kJ mol^-1.
Experiments
1. Use equal weights of dried food, bread, puffed rice, nuts.
Put 20 mL of water in a test-tube attached to a stand.
Push the blunt end a needle into a cork then stick the sharp end into the food sample.
Record the temperature of the water in the test-tube.
Set alight the food with a Bunsen burner then immediately hold the burning food under the test-tube for two minutes.
Record the temperature.
Repeat the experiment with different kinds of foods.
Note which foods released the most energy when burning
2. In a science laboratory a "bomb calorimeter" is used to burn food sample and calculate the energy stored in the food sample from the increase in temperature of water around the bomb calorimeter.
A 10.5 g sample of sucrose was burnt in excess oxygen in a bomb calorimeter.
The calorimeter contained 1.25 kg water and the temperature of the calorimeter and its contents increased from 20.00˚C to 43.34˚C.
In another experiment, it was found that the calorimeter (water excluded) had a heat capacity of 2.2kJ/˚C, (this is not a calibration factor).
What quantity of heat was released by the combustion of 1.0 mol of sucrose?
(heat capacity of water = 4.18 J/g/˚C).
Water: q = m x C x ΔT = 1250 x 4.18 x 23.34 = 121951.5 J = 121.9515 kJ
Calorimeter: q = 2.2 x ΔT = 2.2 x 23.34 = 51.348 kJ
Total heat released: 173.2995 kJ from 10.5 g sucrose
C₁₂H₂₂O₁₁ 12 x C + 22 x H + 11 x O
12 x 12.0 + 22 x 1.01 + 11 x 16.0 = 342.22 g / mol sucrose
10.5 g x 342.22 = 3.068 x 10² mol sucrose
173.2995 kJ from 0.03068 mol
173.2995 kJ x 0.03068 mol = 5648 kJ / mol
(or 5.65 x 103 kJ / mol to 3 significant figures matching data)
3. Saltpetre, KNO3, is used together with common salt, NaCl, to make a salt solution, brine, to pump into pork to produce ham and bacon.
When NaCl is dissolved in water, very little temperature change occurs, but when KNO3 is dissolved in water, a significant temperature change
occurs in accordance with the equation:
KNO3 (s) --> K⁺(aq) + NO3 (aq), where ΔH = +35 kJ
Is it possible to produce ice by floating a small sample of pure water in an alfoil container of negligible heat capacity in a large container of water
at 20˚C by adding KNO3 and stirring it into solution?
The solution of KNO3 is 37g / 100 mL at 25˚C, but this reduces as the temperature drops near 0˚C, to only 12.5g / 100 mL.
Water freezes at 0˚C, but salty solutions freeze at lower temperatures.
Would it be possible to freeze a small sample of pure water floating in a bath of 1 L of water at 20˚C by dissolving a large amount of KNO3?
If temperature of one litre of water is lowered from 20˚C to 0˚C, q = m x C x ΔT = 1000 x 4.18 x 20 = 83.6 kJ heat needs to be absorbed (assuming
solution has same specific heat capacity as pure water).
KNO3 (s) --> K⁺ (aq) + NO3^- (aq), ΔH = + 35 kJ / mol
83.6 x 35 = 2.39 moles KNO3 (s) needs to be dissolved to lower the temperature this much
KNO3: 1 x K + 1 x N + 3 x O
1 x 39.1 + 1 x 14.0 + 3 x 16.0 = 101.1 g / mol
The solubility of KNO3 is 370 g / L at 25˚C and 125 g / L at 0˚C
2.39 mol KNO3 has a mass of 101.1 x 2.39 = 241.5 g
This mass would dissolve at 25˚C, but the solution would become saturated and KNO3 would precipitate from the solution if it reached 0˚C.
If dissolving is endothermic, then its reverse, precipitation must be exothermic.
Saturation would be reached at some temperature above 0˚C and there would be no further cooling.
So it would not be possible to freeze a small sample of pure water floating in a bath of 1L of water at 20˚C by dissolving a large amount of KNO3 in it.
4. Heat of combustion of five different alcohols - methanol, ethanol, propanol, pentanol and octanol.
The molar heat goes up linearly with the increase in molecular size as an extra C and 2 H are added.
However, the heat per gram is much better for longer alcohols, than the short ones, due to the short ones having a much higher relative mass of oxygen (already more oxidized).

22.5.7 Heat apacity of a metal, e.g. lead shot
Put a test-tube containing lead shot in boiling water.
Repeat the above experiment, but add the hot lead shot to the calorimeter instead of the 30 mL of boiling water. (See: *** above)
B x (T3 - T1) + 70 x (T3 - T1) + 70 x (100 - T3) = 0
The heat capacity of a metal in joules per gram per kelvin x molar mass of the metal in grams per mole = the constant, Cp x M =
25 J mol^-1 K^-1 (Dulong and Petit's law).
Heat capacity of metals expressed as joules per gram per degree Kelvin, J.g^-1.K^-1
Mg 1.04, Al 0.904, Fe 0.473, Ni 0.444, Cu 0.387, Zn 0.386, Ag 0.236, Sb 0.207, Au 0.129, Pb 0.128
Experiment
Put 100 ml tap water in a 250 mL beaker and place on a hotplate.
Tie cotton thread around two 50 g brass masses so they can be lifted.
Put the brass masses in the beaker leaving some thread hanging over the side.
Heat the water to boiling.
Put 50 mL (or 50 g) of water from the tap into a foam cup.
Put a thermometer in the foam cup, record the temperature, and leave the thermometer in the foam cup.
When the water in the beaker has boiled, use metal tongs to pick up the ends of the cotton.
Hold the brass mass in the air until all the water evaporates from it, then quickly put the brass mass in the water in the foam cup without splashing.
Stir the water with the thermometer until the temperature reading remains constant.

22.5.6.5.The specific heat of a metal
oC.
Specific heat is a unique property of each substance.
Units of Cm are J kg^-1 K^-1. Equipment: Calorimeter, Laboratory boiler or 250-mL Pyrex beaker in which water can be heated, Metal mass
Figure 1. Diagram of a calorimeter.
Heat exchange takes place between the hot metal and the water in the calorimeter.
If the mass of the metal, its change in temperature, and its heat loss are known, its specific heat can be calculated.
Objectives: During this investigation you will employ the law of conservation of energy as a means of measuring the specific heat of a metal.
Procedure:
1 Record all initial measurements required.
2 Fill a beaker about half full of water and heat it.
3 While waiting for the water to boil, measure and record the mass of the calorimeter cup.
Fill the cup about half full of cold water.
Calculate the mass of the water.
Put the cup in the calorimeter jacket and cover it.
4 Measure the mass of the piece of metal you are using.
Using string, lower the metal mass into the boiling water.
Let the metal remain in the boiling water for about 5 minutes.
5 Measure and record the temperature of the cool water and cup.
6 Measure and record the temperature of the boiling water.
This will also be the temperature of the metal you will place in the boiling water.
7 Remove the metal from the beaker and quickly lower it into the cool water in the calorimeter cup.
Replace the cover at once.
Stir gently with the thermometer.
When the water reaches a constant temperature, record this temperature as the final temperature of the system.
8 Determine the change in temperature of the metal (deltaTm) and the change in the temperature of the cool water and the cup (deltaTcw).
9 If time permits repeat the experiment.
Results: Table 1: Data Trial 1 Trial 2 Trial 3.
Specific heat of calorimeter (J kg-1 K-1) Trial 1 Trial 2 Trial 3.
385 385 385.
Mass of metal calorimeter cup (kg).
Mass of entire calorimeter cup (kg).
Mass of cup plus cool water (kg).
Mass of metal (kg).
Initial temperature of cup and cool water (oC).
Temperature of hot metal (^oC).
Final temperature of system (^oC).
Table 2: Calculations.
Mass of water (kg).
deltaT metal (^oC).
deltaT cool water and cup (^oC).

22.5.8 Heat capacity of a metal with a Styrofoam cup calorimeter
A calorimeter made from two Styrofoam coffee cups is a constant pressure calorimeter in that measures the change in enthalpy of a reaction with the atmospheric pressure remaining constant.
Cp = [W x dH / (M x dT)], where dH = enthalpy of solution, dT = change of temperature, w = weight of solute, m = molecular weight of solute.
In calorimetry, dH is the heat energy released at constant pressure and dE is the energy released at constant volume.
dH = dE + [d(PV)], where The enthalpy increase, dH, is the heat added to a system at constant pressure.
dH = mCpdT where dH = change in enthalpy, m = mass of substance, Cp = heat capacity at constant pressure (J g^-1K^-1), dT = temperature change.
Styrofoam cup + cool water + hot water = 0 dHcal + dH_cw + dH_hw = 0
mCpdT + m_cwCpdT_cw + m_hwCpdT_hw = 0
Let B, the calorimeter constant for the Styrofoam cup = MCp J /K
The calorimeter constant describes how the calorimeter responds to added heat.
So BdT_cw + m_cwCpdT_cw + m_hwCpdT_hw = 0

22.5.18 Ice calorimeter
Heat different metals of the same mass to the same temp and lower into funnels filled with crushed ice then collect the melted water in graduated cylinders.

22.5.9 Melting wax calorimeter
See diagram 24.1.1: Melting wax calorimeter.
Experiment
Prepare equal masses of aluminium, steel and lead.
Tie a thread to each metal and put each into a copper beaker.
Pour the same volume of boiling water on the metals.
Use a smooth wooden board.
Its size can meet the needs of all metals arranged one by one leaving space between each.
Cover a thick layer of paraffin wax evenly on the board in advance.
Lift the thread tied to the metal out of the boiling water, put it rapidly on the board.
The hot metal produces a concavity on the layer of paraffin wax.
After the paraffin wax stops melting, compare the width and depth of different concavities that can show that different metal has different specific heat.
From this you can know roughly the values of specific heat of different metals.

22.5.10 Mix heated water
Experiment
Mix different masses of hot and cold water and compare the final temperature to the calculated value.
If the temperature of a 10 kg mass of copper, specific heat capacity 400 J / kg^oC, rises from 20^oC to 35^oC, heat received by copper mass: H = mc(t2-t1) = 10 x 400 x (35 -20) = 60, 000 joules, J.

22.5.11 Rate of heat transfer
Experiment
Use two identical thermos flasks containing:
(a) 300 g water at 40^oC and
(b) 100 g ice + 200 g water at 0^oC.
Which reaches room temperature 20^oC first?
The loss or gain of heat is greatest when the difference in temperature between the contents of the flask and the surroundings is greatest.
In thermos flask (a), the temperature difference with room temperature is continually decreasing from its original temperature difference of 40^o - 20^o = 20^oC.
In thermos flask (b) the temperature difference remains the same at 0^o - 20^o = 20^oC, when latent heat is absorbed to convert the ice to water.
More than half the total heat is received when the temperature difference remains at 20^oC.
So (a) reaches room temperature first.

22.5.12 Simple calorimeter
See diagram 22.1.8: Simple calorimeter.
1. Use small soup tins that fit loosely into a jam jar.
If the top of the soup tin is cut off cleanly with a rotary type opener, it serves as an excellent calorimeter.
Attach a stout rubber band round the soup tin or use Styrofoam from drink cups to suspend the soup tin inside the jam jar.
Fill the space between them with glass wool or crumpled paper.

22.5.15 Specific heat of canola oil after calibrating the setup with water.
Experiment
Weigh the beaker and the beaker plus 100 mL water + masses, and record them on the data sheet.
Set the Bunsen burner, retort stand with clamp, thermometer, tripod, and beaker with water up as per diagram 1.
The thermometer must not touch the bottom of the beaker.
Heat the water over a blue flame, taking measurements every 30 seconds.
Repeat steps 1-3 with canola oil.
Be careful!.
If there is water present in the oil it will cause the oil to bubble and spit.
Be sure to dry your beaker thoroughly.

22.7.12 Test a thermometer
1. Thermometer scales are marked at two fixed points, i.e. boiling water temperature and the temperature of melting ice.
Use a thermometer and place it in steam immediately above the surface of water boiling in a flask.
Leave it there for several minutes and note how closely it reads 100^oC.
2. Remove the thermometer from the steam, allow it to cool for a few moments, and then place it in a jar of melting ice.
Now observe how nearly it reads 0^oC. 3. If you live at a high altitude, the temperature of boiling water may be below 100^oC, because of the reduced atmospheric pressure.
The thermometer will read exactly 100^oC only at sea level or where the barometer reading is 760 mm of mercury.

22.5.13 Specific heat calorimeter
Experiment
Heat known masses of lead and copper are heated and put into calorimeters with a known mass of water then calculate specific heats of metals from initial and final temperatures.
If 1 kg aluminium in boiling water (100^oC) put in 0.5 kg water at 10^oC and temperature of water rises to 38.23^oC.
Specific heat of water =4200 J / kg^oC
Heat lost by aluminium = mass x specific heat x (final temperature - initial temperature) = 1 x c x (100 -38.23)
Heat gained by water = mass x specific heat x (final temperature - initial temperature) = 0.5 x 4200 x (38.23 -10)
heat lost = heat gained, so 1 x c x 61.77 = 2100 x 28.23, c = 959.7 J / kg^oC (specific heat of aluminium = 960 J / kg^oC)

22.5.14 Specific heat of aluminium by electrical method
1. Use a solid aluminium cylinder with two holes drilled into it and weighing 1 kg.
Put a thermometer in one hole and an electric immersion heater, e.g. 12 V power supply, into the other hole.
Record temperature when steady and time taken.
Power of immersion heater, J x time immersion heater switched on, seconds, s = mass, m (1 Kg) x c x (t2 -t1).
2. Connect an immersion heater to the 12 volt supply in series with an ammeter and a rheostat.
The immersion heaters are 12 volt, 60 watts.
Adjust the rheostat to a current of about 4 amps.
Insert the immersion heater in the aluminium block and the thermometer in its hole in the block.
Use paraffin oil in the thermometer hole to ensure good thermal contact with the block.
Wait for five minutes then record the temperature of the block.
Close the switch.
Start the clock and note the temperature change.
Read the temperature every half minute and draw a graph.
Connect the immersion heater to the 12 volt supply in series with an ammeter and a rheostat.
The immersion heaters are 12 volt, 60 watts.
Adjust the rheostat to give a current of about 4 amps.
Switch on, start the clock and note the temperature change.
Two methods can be used: 1. Take the temperature every half minute and plot a graph.
2. Take the total temperature rise over the known heating period.
Continue recording temperature after the switch is opened.
If the potential difference across the heater gave a current through it of 31 amps, and the temperature rose from 24.5^oC to 37.8^oC in six minutes when the electricity supply was turned off, the temperature may continue to rise to 40.1^oC after 8.5 minutes.
Close the switch, start the clock, record the temperature, open the switch after 10^oC rise, record the time for which the heater was in operation, record the temperature after a further four minutes has elapsed.
(temperature rise, ^oC x specific heat) / time, seconds = (current, amps x potential difference, volts) / J.

22.5.16 Specific heat of water by electrical method
Note the very high and unexplained specific heat of water, 4.1823 J g^-1 K^-1 (4.2 joules per gram per kelvin) at 25^oC.
The higher specific heat means that water is very suitable for use in central heating systems or cooling engines.
Water has about five times the specific heat capacity of land, which keeps islands cooler in summer and hotter in winter.
So continents have greater temperature variations than islands.
The exceptionally high heat capacity of water slow temperature changes, allows heat to be transported around the world by ocean currents and influences climate change.
Experiments
1. Examine the electric jug or immersion heater for a power rating in watts.
If power is not shown on an immersion heater, put immersion heater in water, connect to its power supply with an ammeter in the circuit, e.g. 12 volts DC, and note current used, e.g. 4 amps, power, P = VI = volts x amps = 12 x 4 = 48 watts.
You can also use an ohmmeter to measure the resistance of the heating element and find the current drawn by the element using Ohm's Law, V = IR.
Then calculate the power rating of the element.
2. Measure 1 litre of water.
The heating element must be completely immersed in the water.
3. Measure the temperature of the water, t1
4. Switch on the electric power and record the time.
5. After a period of time when the temperature of the water has increased, but before the water boils, i.e. below 100^oC, switch off the power and record the time and the temperature of the water, t2. 6. Q = mc(t2 - t1), where Q is the energy absorbed, m is the mass of the water, "c" is the specific heat of water and t1 and t2 are the initial and final temperatures of the water.
Assume that the water absorbs all of the energy output from the heating element.
Thus, P = Q / t where P is the power rating of the element, t is the time taken to heat the water.
So P = mc(t2 - t1) / t.
Specific heat of water, c = 4186 J / kg^oC.
If a 3 kW immersion heater raises the temperature of 60 kg water from 10^oC to 60^oC in 70 minutes.
Heat from immersion heater = heat gained by water.
3000 joules x 70 x 60 seconds = 60 kg x c x (60-10)^oC, c = 4200 J / kg^oC
6. Salt water has a lower specific heat capacity than pure water so it will heat up faster, but the salt increases the boiling point of water.
Cooks who add a pinch of salt when cooking pasta causes the water to heat faster, but only by an insignificant time, but increase to time to boiling, but again only by an insignificant time.
The addition of salt does affect the taste of the pasta.
Use the above method to find the specific heat of 1. Sea water, 2. Water with added salt at different concentrations.

22.5.17 Teapot physics
Experiment
Add water at 90^oC to tea in a teapot and leave it to stand for 5 minutes.
If you want the tea to be as cool as possible before drinking it, do you add milk immediately after pouring out the tea or just before drinking it?
The rate of loss of heat depends on the temperature difference between the body and the surroundings.
If the milk were added immediately after pouring the temperature of the tea would fall and the rate of loss of heat would be less than if the milk had not been added.
So the milk should be added just before drinking the tea.

22.7.1 Beurer express thermometer
This thermometer is powered by an alkali manganese battery.
This thermometer has a flexible measuring tip to provide greater comfort and safety during measurement, particularly in infants and persons who are asleep or have a reduced level of consciousness.
Using the thermometer in strong electromagnetic fields such as next to a mobile phone, can cause a malfunction.
Portable and mobile HF communication systems may interfere with this unit.
Temperature measurement in the anus (rectal) is the most reliable and most accurate method.
Guide the tip of the thermometer carefully 2 to 3 cm into the anus.
The measurement time with this thermometer at this site is only about 10 seconds.
The end of the measuring time is indicated by an signal tone.
For measurement in the mouth cavity (oral measurement), guide the tip of the thermometer carefully into one of the heat pouches beneath the tongue, to the left or the right of the root of the tongue.
For measurement in the armpit (axillary measurement), this method of measurement is relatively inaccurate, so it cannot be recommended from a medical point of view.
Range of measurement: 32^oC to 42.9^oC
Accuracy of measurement: 0.1^oC in a water bath with a temperature of 35.5^oC to 42.0^oC
Ambient temperature: 10^oC to 40^oC, with 30% to 85% relative humidity

22.7.2 Thermochromic substances
Substances that show thermochromism have a reversible range of colour or shade when heated or cooled.
"Moody Colour Changing Putty", temperature-sensitive putty (toy product)
"Thermochromic paper", heat-sensitive colour-changing periodic table, thermochromic paper, liquid crystal film (toy product)

22.7.3 Thermometers
Thermometers, alcohol, -10^oC to 110^oC, 250 mm long
Thermometers, alcohol, -10^oC to 110^oC, 300 mm long
Thermometers, alcohol, -15^oC to 50^oC, classroom, wall, hanging type, not used in direct sunlight
Thermometer stand, wooden, holds 20 thermometers, fits 200 mm to 300 mm laboratory thermometers
Digital thermometer, indoor / outdoor, with minimum / maximum capacity, -50^oC to 70^oC
Digital fever thermometer, water resistant, speed read, display in Celsius, memory recall of last reading,
includes 10 disposable probe covers, replaceable battery
Digital thermometer, Omron MC510 Gentle Temp electronic ear thermometer, LCD display,
measures body temperature in 10 seconds, backlit screen, light weight
Probe covers for electronic ear thermometers
Clinical thermometer, mercury in glass, alcohol thermometer
Maximum and minimum thermometer, Six's thermometer, mercury in glass
Thermometers
See diagram 23.7.01: Science experiments thermometer.
See diagram 23.7.02: Clinical thermometer.
See diagram 23.7.03: Wall thermometer.
A thermometer uses changes in a selected property to measure changes in temperature:
change of length in thermostats,
change in volume in mercury thermometers,
change in pressure in gas thermometers,
change in resistance of wires,
change in the emf in thermocouples.
Calibration of thermometers
To calibrate a thermometer in ^oC the value of the selected property is first found at the ice point and mark this 0^oC.
Then the value of the property is found at the steam point and mark this 100^oC.
Divide the change in property between 0^oC and 100^oC into 100 equal parts each equivalent to a change in temperature of 1^oC.
The International Temperature Scale (1990) extends from 272.5 C to approximately 1, 085 C.
Thermometers containing mercury are being phased out in schools.
In hospitals, mercury thermometers are being replaced by electronic digital thermometers, ear thermometers.

22.7.4 Air thermometer
See diagram 23.4.9: Air thermometer.
Experiment
Select a small medicine bottle with a cork stopper.
Bore a hole through the cork with a cork borer to take a piece of thin glass tubing 30 cm long.
Put ink in the bottle to cover the end of the tubing, which should be touching the bottom of the bottle.
Warm the bottle with your hands to make the air in the bottle expand, press down on the liquid in the bottle and force it up the glass tube.

22.7.6 Galileo's thermometer
Use a set of glass spheroid buoys of varying density in a glass cylinder arranged so the lowest floating ball represents the temperature.

22.7.10 Temperature sense 1. Check if your temperature sense is reliable.
Use containers of hot water, warm water and cold water.
Put both hands in the warm water.
The hands feel the same temperature.
Put one hand in the hot water and the other hand in the cold water.
Quickly dry your hands and put them both into the warm water again.
The two hands do not feel the same temperature.
This experiment shows that your temperature sense is not always reliable.
2. Fill three plastic dishes with the same amount of water at different temperatures, 10^oC, 20^oC, 30^oC, arranged from left to right with temperatures from low to high.
Place the right hand in the 10^oC water and the left hand in the 30^oC water.
After two minutes, place both hands in the 20^oC water.
Does the 20^o water feel warm or cold?
The hand feels warmer when it moves from lower temperature to higher temperature.
The hand feel colder when it moves from higher temperature to lower temperature.
So "warm" and "cold" have not a definite standard.
To show the degree of warm and cold of an object quantitatively, it is necessary to introduce an important quantity, temperature, and make equipment that can measure the temperature, the thermometer.

22.7.5 Expansion of liquid in a thermometer
Experiments
1. Fill a flask with coloured water.
Insert a one hole stopper carrying a 30 cm length of glass tubing until the water rises 5 cm in the tubing.
Put the flask in a beaker of water.
Heat the beaker and observe the water level in the tubing.
The water rises in the tubing.
However, if you carefully observe the water level in the tubing when the heating begins, you will see that it falls slightly and then begins to rise!
It falls, because the glass in the flask starts to expand before the water inside.
When the heat energy reaches the water it expands.
So the expansion of liquid you see in a thermometer is really the expansion of liquid less the expansion of the glass tube.
2. Put some crushed ice in a Erlenmeyer flask add a little water.
The depth of mixture in the flask is about one inch.
Spin the flask to mix water and ice.
Insert a thermometer into the mixture.
As the column of liquid in thermometer is stable mark on thermometer beside the scale marked 0^oC.
Put the flask on wire gauze on a tripod heat it with a burner.
As water in the flask boils mark again as 100^oC.
Note to maintain two marks at the same vertical line.
The freezing and boiling point of water are primary standards for calibrating thermometer.
Draw a scale of thermometer between 0^oC and 100^oC.
Divide it into 100 same spaces.
Draw a longer line every 10 points written as 10^oC, 20^oC.
Draw a shorter line every 5 points written as 5^oC, 15^oC, 25^oC.
Compare the scale you have done to original one.
Perhaps they are quite different.
The main reason is the scale lines of 100^oC are not the same is, because the atmospheric pressure is not the standard as you do the experiment and the boiling point of water may not exactly the 100^oC.
To draw line and calibrate easily you can use a sheet of co-ordinate paper stuck to the thermometer in advance.

22.7.7 Make a spirit thermometer
This activity is too dangerous for schools especially when it comes to heat-sealing the tube.
To make a simple alcohol thermometer accurate enough for most purposes, use 20-30 cm of glass tubing of about 5 mm external diameter with about 1 mm bore.
A bulb of about 1.5 cm external diameter is first blown in one end of the tubing.
The tube is inverted and the open end is placed in alcohol.
The bulb is alternately heated and cooled.
After each cooling the alcohol drawn into the tube is shaken into the bulb.
In this way the thermometer is filled with alcohol, care to be taken to remove any air bubbles.
The bulb of the thermometer is then placed in water at 60^oC, which is slightly below the boiling point of alcohol.
The excess alcohol is removed from the top of the tube as it oozes out.
Now seal the open end with a hot flame.
Caution! - great care must be taken when sealing the tube.

22.7.8 Temperature rise and quantity of heat intake
See diagram 4.1: Temperature rise and quantity of heat intake.z
1. Put a large iron bolt and a nut for the bolt in a container of boiling water to bring them to the same temperature.
Put equal volumes of water in two containers with each volume enough to immerse the bolt.
Put the hot bolt in one container and the hot nut in the other container.
Record the temperature of the water in each container after the same period.
The difference in temperature change of the water in the two containers is, because of the different amounts of heat stored in the iron bolt and the iron nut.
2. Check if your temperature sense is reliable.
Use containers of hot water, warm water and cold water.
Put both hands in the warm water.
The hands feel the same temperature.
Put one hand in the hot water and the other hand in the cold water.
Quickly dry your hands and put them both into the warm water again.
The two hands do not feel the same temperature.
Is your temperature sense reliable?
This may be a silly experiment, but it shows that your temperature sense is not always reliable.

22.7.9 Temperature rise calculation
Press report:
“The second and more recent study period, from February to July 2020, showed the average temperature under the shade structure during the peak heat hours was 31.4C, compared with 32.1C outside of it - a difference of 0.7C, or around 2 per cent.”
Ricard Wardling comments as follows:
If the author made the calculations in Kelvin, it would say “304.4 K, compared with 305.1 K outside of it - a difference of around 0.7 K, or around 0.2 per cent”.

22.7.11 Test a liquid in glass thermometer
Use a thermometer with a scale, e.g. a mercury or alcohol thermometer, -10^oC to 110^oC.
Also, use a tall flask containing coloured water fitted with a one-hole stopper and glass tube extending into the bottle.
Attach a blank scale to the glass tube.
A thermometer scale has two fixed points, the lower fixed point and the upper fixed point.
Put the bulb of a thermometer in crushed ice that is melting.
Check that the temperature is 0^oC on the calibrated thermometer.
Mark the lower fixed point on the blank scale.
Put a thermometer in steam immediately above the surface of boiling water.
Check that the temperature is 100^oC on the calibrated thermometer.
Mark the lower fixed point on the blank scale.
Divide the distance between the upper and lower fixed point to obtain 100 marks representing a temperature difference of 1^oC.
If you do the experiment on a mountain at a high altitude, the temperature of boiling water will be below 100^oC, because of the reduced atmospheric pressure.
If you do the experiment in a submerged submarine, the temperature of boiling water may be above 100^oC, because of the increased pressure with depth.
The thermometer in the boiling water reads exactly 100^oC only at sea level or where the barometer reading is 760 mm of mercury.

22.7.13 Thermometer, 0^oC to 100^oC range
Experiment
First put thermometer into crushed ice after a few minutes observe if the liquid column in the thermometer is exactly at 0^oC.
Take out the thermometer from the ice wait for a while at the temperature of surroundings.
Put the thermometer on the surface of boiling water the measuring bulb is near the heat steam above the surface of water.
Wait for a few minutes observe if the liquid column in the thermometer is at the scale of 100^oC.
If you live in the region of higher height above sea level the boiling point of water there will be lower than 100^oC due to the lower atmospheric pressure.
On the contrary, in the region of lower height above sea level the boiling point of water will be higher than 100^oC.
Only in height of sea level the boiling point of water is just 100^oC the atmospheric pressure there is 760 mm Hg.

22.7.14 Thermometer, make a thermometer
Experiment
Prepare a small glass bottle with a rubber stopper perforated to allow a thin glass tube to pass through it.
Use a thermometer to mark 0^oC and 100^oC on the glass tube, or select two suitable temperature points.
Measure the length between 0^oC to 100^oC on the glass tube divide the length into 10 evenly and mark on the tube.
(i.e. mark 9 points with the same distance between them).
Divide by 10 again between every space and mark it so every space is 1^oC.
In regions above sea level correct the fixed point of 100^oC according to the boiling temperature there.

22.7.15 Thermometer test, calibrate a thermometer
There are two marked points on the scale of a thermometer.
If a thermometer can measure the temperature from 0^oC to 100^oC the two fixed points are at 0^oC and 100^oC.
0^oC is the melting point of ice and 100^oC is the boiling point of water.
Other scales of temperature are all defined refer to two fixed points.
The melting point and boiling point depend on atmospheric pressure so the melting point of ice and the boiling point of water are at standard atmosphere.
Experiment
1. Use a thermometer with a scale, e.g. thermometer -10^oC to 110^oC mercury or alcohol, or a tall flask containing coloured water fitted with a one hole stopper and glass tube that extends into the bottle.
Mark thermometer scales at two fixed points.
The lower fixed point is the temperature of melting ice.
Put the bulb of a thermometer in crushed ice that is melting.
Leave it there for some minutes.
Check that the temperature is 0^oC on the calibrated thermometer, or make a mark on the blank scale or on the glass tube attached to the small flask.
The upper fixed point is the temperature of boiling water.
Put a thermometer in steam immediately above the surface of boiling water.
Leave it there for some minutes.
Check that the thermometer reads 100^oC on the calibrated thermometer, or make a mark on the blank scale or on the glass tube attached to the small flask.
Divide the distance between the upper and lower fixed point to obtain marks representing a temperature difference of 1^oC.
If you do the experiment on a mountain at high altitudes the temperature of boiling water may be below 100^oC, because of the reduced atmospheric pressure.
If you do the experiment in a submerged submarine the temperature of boiling water may be above 100^oC.
The thermometer in the boiling water reads exactly 100^oC only at sea level or where the barometer reading is 760 mm of mercury.

22.7.16 Temperature of water at maximum density, 4^oC
Experiments
1. Fill a bottle with water and put the top on securely.
Wrap the bottle in a cloth, to prevent the shattered glass from falling.
Put the bottle into the freezing compartment of a refrigerator.
After 24 hours, remove the bottle and examine it.
The bottle may be cracked, because of pressure from the expanding ice.
2. Put a large piece of ice into a glass of water.
Arrange two thermometers so that they measure the temperatures near the top and the bottom of the water.
The water cooled by the ice falls to the bottom.
This fall in temperature continues until the water at the bottom of the glass reaches a temperature of 4^oC.
The water stays at this temperature for a long time, the colder water remaining higher up near the ice.
So water at 4^oC is denser than the water at 0^oC.
This curious behaviour of water is of great practical significance in nature, and explains why a pond freezes from the surface downwards while the bottom seldom falls below 4^oC.
3. To study the expansion of freezing water, use two identical drinking cups.
Fill the first cup with tap water at room temperature so that the water heaps up to form a meniscus.
Put the second cup in the freezing compartment of the refrigerator then add extra water to the cup to get the highest possible meniscus.
When the water in the cup is frozen, compare the meniscus of the frozen water with the meniscus at room temperature.
The frozen water heaped up, because it had expanded.
Water has a maximum density at 4^oC.
When water cools from room temperature to 4^oC, it contracts in volume.
When water cools from 4^oC to 0^oC, it expands in volume.
At 4^oC the density of water is 1000 kg m^-3 (1 g per cc).
At 0^oC the density of water is 999.87 kg m^-3 and the density of ice is 918 kg m^-3, so ice floats on water.
4. Fill a large jar with cold water and weigh it accurately on a balance.
Empty the jar.
Fill the jar with exactly the same volume of hot water and weigh.
You will observe that the jar of warm water weighs less.
Volume for volume, cold water is heavier than warm water; so when water is heated convection current are set up, the warm water being lifted, because of buoyancy, by the cold surrounding water.
Hot water is less dense than cold water, and this is the cause of convection current in a liquid.
5. To study the expansion of freezing water, use two identical drinking cups.
Fill the first cup with tap water at room temperature so that the water heaps up to form a meniscus.
Put the second cup in the freezing compartment of the refrigerator then add extra water to the cup to get the highest possible meniscus.
When the water in the cup is frozen, compare the meniscus of the frozen water with the meniscus at room temperature.
The frozen water heaped up, because it had expanded.
Water has a maximum density at 4^oC.
When water cools from room temperature to 4^oC, it contracts in volume.
When water cools from 4^oC to 0^oC, it expands in volume.
At 4^oC the density of water is 1 000 kg m^-3 (1 g per cc).
At 0^oC the density of water is 999.87 kg m^-3 and the density of ice is 918 kg m^-3, so ice floats on water.