School Science Lessons
Change of state, latent heat,
humidity, heat engines, laws of thermodynamics
2009-11-10
Please send comments to: J.Elfick@uq.edu.au
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Table of contents
24.0.0 Change of state
24.1.0 Phase changes, liquid / solid, melting
point
24.2.0 Phase changes, liquid / gas, boiling
point
24.2.1.0 Latent
heats
of fusion and vaporization
6.3.1.5
Temperature, Fahrenheit scale, Celsius scale, Kelvin scale
24.3.0
Cooling by evaporation
24.3.0.1 Sublimation
24.4.0 Dew point and humidity
24.7.0 Heat engines, steam turbines,
refrigerator
24.0.0 Change
of state
3.8
Pressure affects the boiling point
4.1
Temperature rise and
quantity of heat intake
4.2
Transfer kinetic energy to
heat energy
24.1.0
Phase changes
liquid / solid, melting point and freezing point
3.2
Melting point of naphthalene
3.3
Melting point of naphthalene with a capillary tube
3.4
Impurities affect the melting point of a substance
3.43 Solid,
liquid and gas in a
plastic drink bottle (Primary)
4.39 Melt
different solids
(Primary)
4.41 Ice
experiments (Primary)
7.4.0
Melting point, m. p. of solids
7.4.1
Melting point and cooling curve
of stearic acid
7.4.2
Melting point of different
substances
7.4.3
Melting point of ice and
freezing point, f.p. of water
7.4.3.1
Temperature at which ice
melts.
7.4.3.2
Temperature at which ice and
salt mixture freezes
7.4.3.3
Lift ice cube with match
stick
24.1.01 Nucleation
24.1.02 Vapour pressure of pure water and
solution of a non-volatile solvent
24.1.03 Raoult's law
24.1.04 Depression of freezing point and
elevation of boiling point
24.1.05 Critical point, critical temperature
24.1.1 Melting point of ice or
freezing point of water
24.1.2 Lift ice cube with salt
24.1.3 Ice melting by pressure and refreezing,
regelation
24.1.6 Bitumen foaming and bitumen decay
24.1.7 Heat of solution
24.1.8 Heat of crystallization
24.1.9 Water crystals in soap film
24.1.10 Crystal growth on the overhead projector
24.1.11 Metglas
24.1.12 Wood's metal
24.1.13 Hot water freezes faster than
cold water, the Mpemba effect
24.1.14 Molal freezing point constant of
cyclohexane solvent
24.1.15 Molar mass of solute from depression of
freezing point
24.2.0 Phase
changes, liquid / gas,
boiling point
3.5
Boiling point of water
3.5.1
Boiling point of sodium chloride solution
3.6
Boiling point of inflammable liquids
3.7
Volatility of different liquids
3.8
Pressure affects the boiling point
4.239 Pressure affects boiling
point of water
24.2.1.2 Boil a kettle with a spout
24.2.2 Heat of condensation of water
24.2.4 Pressure and boiling point of
water
24.2.5 Temperature of steam above
water boiling
24.2.6 Set fire to paper with a jet of
steam
24.2.7 Supercooled water
24.2.8 Broken bottle, ice bomb
(Dangerous experiment not suitable for schools!)
24.2.9 Liquefying dry ice
7.5.0 Boiling
point, b.p. of liquids
7.5.0.1
Elevation of boiling point,
ebullioscopic constant, kB,
7.5.1 Measure boiling point of water
3.5.1
Boiling point of sodium chloride solution
7.5.3
Boiling point of mixture of two
liquids, water and alcohol
7.5.4 Boiling points of inflammable
liquids, ethanol
7.5.4.1 Boiling points of
inflammable liquids, acetone
3.7
Volatility of different liquids
4.40 Heat and cool water (Primary)
24.2.1.0 Latent
heats
of fusion and vaporization
4.10
Heat energy to change solid
to liquid,
melting point, latent heat of fusion
4.11
Heat
energy to change liquid to
vapour, boiling point, latent heat of vaporization
24.1.4
Latent heat, change of state, melting
and boiling
24.1.5
Latent heat of fusion of ice to
water
24.2.1
Latent heat of steam
24.2.3
Latent heat of vaporization of
water
24.3.0
Cooling by evaporation
4.38
Liquids in the sun (Primary)
24.3.1 Water "lost" by evaporation
24.3.2 Factors which affect the rate of
evaporation
24.3.3 Evaporate dichloromethane
24.3.4 Evaporate from the blackboard
24.3.5 Cryophorous
24.3.6 Freezing by evaporation
24.3.7 Drinking bird, drinking duck, dippy bird,
dunking bird
24.3.8 Cooling by evaporation, ether
or ethyl chloride
24.3.9 Water bag, Coolgardie safe
24.3.10 Evaporation from a microscope slide
24.4.0 Dew point and humidity
5.22
Rain gauge (Primary)
6.17
Measure relative humidity (Primary)
24.4.2
Dew point measurement
24.4.3
Condensation nuclei
37.8.0
Measure relative humidity
37.8.1
Wet and dry bulb hygrometer (wet and dry bulb thermometers)
37.8.2
Sling psychrometer
37.8.3
Dew point hygrometer
37.8.4
Relative humidity table, depression of the wet bulb
37.5
Make a rain gauge, precipitation gauge
37.6
Make a hair hygrometer
37.15
Atmospheric moisture
37.16
Weigh water "lost" by evaporation
37.17
Moisture evaporates from soil
37.18
Moisture comes from plants
37.19 Moisture comes from other plants
37.20
Moisture from breathing
37.21
Surface area affects evaporation
37.22
Temperature affects the rate of evaporation
37.23
Moving air affects the rate of evaporation
37.24
Moisture in the air affects the rate of evaporation
37.25
Moisture condenses on cool surfaces
37.26
Study the water cycle
37.27
Make a rain cycle
37.29
Make cloud in a bottle
37.30
Study snowflakes
4.41 Ice experiments (Primary)
37.38.1
Warm front
37.38.2
Cold front
37.38.3
Tornadoes
37.38.4
Hurricanes
24.7.0 Heat engines,
steam turbines,
refrigerator
24.7.1 Steam turbines
24.7.2 Refrigerator, ice chest,
portable ice boxes, expresso coffee
machine
24.0.0 Change of state
State of matter, solid (s),
liquid (l), gas (g),
aqueous solution (dissolved in water) (aq)
Change of State, melting point and freezing point, boiling point and
air pressure, evaporation and condensation, cooling by evaporation
PVT Surfaces (for an ideal gas, the pressure, volume, temperature 3
dimensional plot on a graph, 3-dimensional models)
Change of state
The three states of matter are solid, liquid and gas. When
a substance changes from one state of matter to another, it
has undergone a change of state (change of phase). Change of state
always occurs with a change of heat energy that flows into or out of
the material but no change of temperature occurs. During change of
liquid to solid, freezing, heat leaves the liquid as it freezes. During
change of liquid to gas, vaporization (boiling and evaporation), heat
goes into the liquid as it vaporizes. During change of gas to liquid,
condensation, heat leaves the gas as it condenses. During change of
solid to gas, sublimation, heat goes into the solid as it sublimates.
The heat energy exchanges during a change of state are changes in
potential energy, not kinetic energy. The heat energy that comes into
ice during a phase change breaks the bonds between the molecules of
water. The molecules are now at a higher potential energy state, but
not, on the average, moving any faster, so their average kinetic
energy remains the same, and the temperature remains the same.
24.1.0 Phase
changes, liquid / solid, melting point and freezing point
Change of
state, melting point and freezing point
The melting point is the temperature at which a substance changes from
solid form to liquid form. A pure substance under standard conditions
of pressure has a definite melting point. If you supply heat to a solid
at its melting point, the temperature does not change until the melting
process is complete. The melting point of ice is 0oC.
Boiling is rapid conversion of a liquid into vapour that takes place
when the liquid reaches a certain temperature. It involves the
formation of vapour bubbles within the body of a liquid, whereas
evaporation occurs only at the surface. Boiling point, for any given
liquid, is the temperature at which the application of heat raises the
temperature of the liquid no further, but converts it to vapour. The
boiling point of water under normal pressure is 100oC. The
lower the pressure, the lower the boiling point and vice versa.
Freezing is the change of state from liquid to solid that occurs at the
freezing point of a substance. For a given substance, freezing occurs
at a definite temperature, called its freezing point, which is
invariable under similar conditions of pressure. The temperature
remains at this point until all the liquid is frozen.
24.1.01 Nucleation
The phenomenon of nucleation refers to the start of many types of new
processes, including crystallization, bubble formation in a saturate
liquid, liquid droplets in a saturated vapour, fracture.
The places where nucleation occurs are called
nucleation sites, e.g. suspended particles, dust, cracks in the inner
surface of a container, iodine crystals used for cloud
seeding, spaces in solid materials. Such nucleation is
called heterogeneous nucleation. However when nucleation occurs without
nucleation sites it is called homogeneous nucleation, e.g. in
supercooled liquids.
24.1.02 Vapour pressure
of pure water and solution of a non-volatile solvent
See diagram 24.1.02: Vapour Pressure | See
also: Table of saturated vapour pressure over water, Psvp
Put a beaker of deionized water under a strong container, e.g. a bell
jar and evacuate the air with a strong pump. Measure the pressure in
the container with an aneroid barometer. At 25oC,
the air
pressure falls from atmospheric pressure to the vapour pressure of pure
water, 23.8 mm Hg (3.169 kPa) as water molecules evaporate from the
water surface in the beaker. Repeat the experiment with a 1 molal
solution in water of ethylene glycol (ethane-1,2-diol, antifreeze).
This lowering of vapour pressure is a colligative property, i.e. a
property that depends on the relative number of particles and not their
chemical properties.
24.1.03 Raoult's law
The vapour pressure of a solvent above a solution, P1 = Xsolvent
X P10, where Xsolvent = the mole fraction of the
solvent in the solution and P10= the vapour pressure of the
pure solvent.
Xsolvent = number of moles of solvent / total number of
moles. For example, if 0.1 moles of sugar are dissolved in
10 moles of water, Xsolvent = 10 / 10.1 =
0.99.
Raoult's law is colligative because the solute added must not react
with the solvent so the solution should behave as an ideal solution. In
general Raoult's law only applies to very dilute solutions. and is
based on the principle that when a solute is added to a solvent
there is a decrease in the relative number of solvent molecules at the
surface of the liquid to evaporate.
24.1.04 Depression of
freezing point and elevation of boiling point
See 7.5.0.1: Elevation of
boiling point,
ebullioscopic constant, kB
See diagram 24.1.04: Freezing point
depression and boiling point elevation | See
diagram 24.1.04.1: Elevation of boiling point
This phase diagram shows that when a non-volatile solute is dissolved
in a solvent the solvent vapour pressure is lowered. The vapour
pressure
curve for the solution is lower than the vapour pressure curve for pure
water but not parallel to it. The vapour pressure curve of the solution
intersects the sublimation pressure curve at a lower temperature than
the vapour pressure curve. The freezing points curves for the solution
and water begin at the triple points, the intersections of the
sublimation pressure curve and the vapour pressure curve, where solid,
liquid and gas phases can simultaneously exist. So the freezing point
curve for the solution is at a lower temperature than freezing point
curve for water. Freezing point depression can be used to calculate the
molar mass of a substance. The freezing point depression is
proportional to the sum of the concentrations of all the different
dissolved particles so it can be used to measure biological
concentrations, e.g. urine.
A solution has no boiling point because the solvent tends to evaporate
more than the solute, so the concentration of the solution increases.
However, a constant boiling mixture (azeotrope)
is an exception to Raoult's law where the composition of the
vapour is the same as the composition of the boiling liquid, e.g. a
mixture of 4% water and 96% ethanol. So a mixture of water and alcohol
can never be converted to absolute alcohol by boiling.
Freezing point depression, dTf = i X kf X m,
where dTf
= difference between freezing point of solvent and freezing point of
solution, as a positive value, and kf = molal freezing point
depression constant, e.g. water = 1.86oC kg /mole.
Boiling point elevation, dTb = i X kb X m, where
dTb = difference between boiling point of solvent and
boiling point of solution, as a positive value, and kb
= molal boiling point elevation constant, e.g. water = 0.52 oC
kg /mole
m = molality (molal concentration) mol kg-1. (Molality =
moles of solute / kg of solvent.)
i = Van't Hoff factor. (If not ionized particles, i =1, e.g.
sucrose. If particles of two ions, i = 2, e.g. NaCl. If
particles of 3 ions, i = 3, e.g. MgCl2.)
Each solvent has a specific value for kb and kf, e.g. for benzene, kb =
2.53oC kg /mole and kf = 5.10oC
kg /mole. The colligative properties of a solution include depression
of freezing point, elevation of boiling point and osmosis. The effects
of alcohols on the temperature range decrease with the increase of
molecular weight because the size of colligative properties depends on
the number of particles. So if equal weights of methanol (32.04 g /
mole) and 1-butanol (74.12 g / mole) methanol would have more
particles. While boiling point elevation can be explained by decrease
of relative number of solvent molecules to evaporate, freezing point
depression cannot be explained by saying that solute particles "get in
the way of the solvent particles" but rather by the difference in
the chemical potential of the solvent particles. Depression of
freezing point is used when salting ice for making ice cream at home
and salting roads in cold climate winters. Elevation of boiling
point is used when adding ethylene glycol ("antifreeze") to the water
of car radiators when "winterizing" the car.
24.1.05 Critical point,
critical temperature
See diagram 24.1.05: Critical point of carbon
dioxide
For water, the upper limit of the vaporization curve is at 374oC
and 218 atm, called the critical point. Above the critical temperature,
a gas cannot exist in the liquid state no matter what pressure is
applied. Helium has the lowest critical temperature, -268oC.
The critical point for carbon dioxide is 31oC and 73 atm. At
the critical point the liquid and gas phases have the same density.
At the critical point the the physical properties of the liquid and
gaseous phases are the same. At the critical point (critical state)
the liquid state of the matter ceases to exist. When a liquid is
heated, its density decreases and the pressure and density of its
vapour increases until the critical temperature is reached where the
two densities are equal and the phase boundary between liquid and gas
disappears. The heat of vaporization becomes zero. The critical point
of water is at 647K (374°C) and 22.064 MPa (3200 PSIA or 218
atmospheres, atm.).
24.1.1 Melting point of ice or freezing point of
water
You need a refrigerator with the temperature in the freezer about -5oC.
Cover a 250 mL beaker with insulating material, half fill it with water
and add small pieces of ice. After 10 minutes, record the temperature
of mixture of water and ice with a thermometer. Take out some water
from the beaker, add smaller pieces of ice just taken from the freezer.
Add layers of ice and salt until they half fill the beaker. Put a
test-tube with half full of water into the beaker so that it becomes
surrounded by ice. Put a thermometer in the test-tube. Compare the
melting temperature of ice in the beaker with the freezing temperature
of water in the test-tube. Let the ice in the beaker melt completely
and let the water in the tube freeze to ice completely. The readings of
the two thermometers both reach the same value below 0oC.
Observe the liquid water and solid water exist together.
24.1.2 Lift an ice cube with salt
See diagram 24.2.7
Ice melts and freezes again by the action of salt. Put an ice cube in a
full cup of cold water so that it floats. Let a thread lie on the ice
cube and sprinkle salt on it, see the diagram. The aim of sprinkling
salt is to wet the thread by melting the ice near the thread. Wait
until the melted ice near the thread condenses again. Lift one end of
the thread to check if water on it has frozen completely. If it is,
lift two ends of the thread to hold the ice cube from the water. As the
salt liquid and pure water have different melting points and freezing
points, the salt can lift an ice cube.
24.1.3 Ice melting by pressure and refreezing,
regelation
See diagram 24.1.3: Melting by pressure
Increase of pressure increases the
freezing point of substances but increase of pressure on substances
that expand on freezing lowers the freezing point. So increased
pressure on ice turns it to liquid but when increased pressure stops
the ice freezes again without change in temperature, regelation.
1. When you apply pressure to ice you lower the freezing point of it.
Freeze a rectangular ice
cube in a freezer. Put it on a wooden board. Tie 500 g weights to each
end of a thin nylon thread. Spread the thread across the upper surface
of the ice and hang the weights at each side. Observe the ice melting
below the thread as it cuts slowly into the ice. If the temperature is
low enough as the ice freezes, the ice above the thread can be frozen
again. Then you can lift the ice cube by nylon thread. Substances that
expand on freezing show a lowering melting point under pressure.
2. Push your finger down on the middle of an ice cube. Some ice at
the upper surface melts due to increased temperature and pressure.
Sprinkle salt on the ice cube and press it again. The ice cube melts
slower because salt lowers the melting point of ice.
3. Squeeze crushed ice in a thick walled cylinder to form a solid
block.
4. Cut through a block of ice with a wire loop that has a heavy mass
hanging from each end. Copper wire cuts through faster than iron or
cotton thread. However, the copper wire should be the same temperature
as the ices otherwise it is transferring heat and melting the ice. The
ice under a skaters ice skates quickly melts under
the skate then refreezes again when the skater has passed.
However, some people claim that pressure from ice skates is not enough
to
melt ice, except when the temperature is a fraction of a degree below 0oC.
In this view, skates move easily over ice because a very thin layer of
water already on the ice lowers the friction and melting from
frictional heating give more water lubricant.
24.1.4 Latent heat, change of state, melting
and boiling
1. A change in the physical state (solid, liquid, or gas) of a
material is called a change of state, e.g. melting, boiling,
evaporation,
solidification and condensation.
2. Latent heat of fusion: The amount of heat per unit mass that has to
be removed to freeze a substance is a constant for any given substance,
and is called the latent heat of fusion. When a pure substance changes
from solid to liquid the temperature does not rise but remains at the
melting point until all the solid has melted. The heat required to melt
one kilogram of the substance is called the latent heat of fusion Water
= 3.33 X 105 joule/kg. The exceptionally high latent heat
of fusion of water, 333 kj kg-1, prevents water temperature
from changing quickly at temperatures near 0oC above or
below, because of the extra energy needed to freeze or melt water.
3. Latent heat of vaporization: The heat required to evaporate one
kilogram of the substance is called the latent heat of vaporization
e.g. Water = 22.6 X 105 joule/kg. When a pure substance
changes from liquid to vapour the temperature does not rise but remains
at the boiling point until all the liquid has boiled away.
A pure substances absorbs heat energy to change from solid to liquid at
the constant temperature of the melting point. The heat required to
melt 1 kilogram of the substance is called the latent heat of fusion (lf).
The lf of water = 3.33 x 105 joule/kg. A pure
substances absorbs heat energy to change from liquid to vapour at the
constant temperature of the boiling point. The heat required to
evaporate 1 kilogram of the substance is called the latent heat of
vaporization (lv). The lv of water = 22.6 x 105
joule/kg. The latent heat of evaporation of water is the highest of all
substances, 2260 kj kg-1, which reduces water loss and heat
loss to the atmosphere.
4. The assembled state of molecules or atoms in matter is called phase.
The transformation of the matter from a state to another is called
phase change. The amount of heat absorbed or emitted by an object of
unit mass in the process of phase change is called latent heat in phase
change. The heat of vaporization, the heat of solution and the heat of
sublimation are all latent heat in phase change. Under the same
pressure the same kind of matters are transformed from liquid state to
steam state, the heat of vaporization needed is equal to that from
steam state back to liquid state. As the pressure and temperature are
in the condition of matter which is in such a state that the solid,
liquid and steam, three phases exist together, there is a relation of
value: the heat of sublimation = the heat of solution + the heat of
vaporization.
24.1.5 Latent heat of fusion of ice to water
1. Weigh a 500 mL beaker. Half fill with crushed ice and weigh again.
Record the temperature of the ice. Heat the beaker and ice for a minute
and again note its temperature. Repeat until all the ice has been
melted for several minutes. Record your observations as: State of
matter, Time in minutes, Temperature in "C. Draw a graph of your
results by plotting temperature against time and note the shape of the
graph. When there is no ice but only ice water, weigh the container and
contents. There should be no noticeable change in mass because few
particles from the original ice or water have escaped.
2. Continue as in the last experiment heating more strongly and note
the temperature every minute until the water boils. Again there would
be no noticeable change in mass until a considerable number of
molecules have been lost as steam. Again record your results and make a
graph of temperature against time. Note the rise in temperature as heat
is supplied.
3. The heat of solution of ice can be measured by
measuring how much ice can be melted by a piece of hot solid covered on
the ice. Measure the mass of a piece of metal, its specific heat has
known. Hang it by a thread to let it go into water and heat water until
boiling. Then put it into a funnel as quickly as possible, cover pieces
of ice on it. Collect the melted water to a measuring cylinder through
the funnel. Record the volume and calculate the mass of melted water.
Note the funnel should be big enough, and the pieces of ice are more
enough, in other words, there should be left a few pieces of ice which
cannot be melted. Calculate the amount of heat emitted by metal in the
process of the temperature dropping to 0oC, the ratio of it
to the mass of water melted is the heat of solution of ice.
4. Plot the cooling curve of ethanamide,
acetamide
24.1.6 Bitumen foaming
and bitumen decay
The behaviour of bitumen during foaming is primarily due to physical
processes although chemistry does also play a role. When a cold water
droplet at ambient temperature makes contact with the bitumen at 170oC
to 180o̊C, the following chain of events occurs:
1. The bitumen exchanges energy with the surface of the water droplet
heating the droplet to a temperature of 100̊C and cooling the bitumen.
2. The transferred energy of the bitumen exceeds the latent heat of
steam resulting in explosive expansion and the generation of
steam.
3. Steam bubbles are forced into the continuous phase of bitumen
under pressure.
4. With emission from the bitumen spray nozzle the encapsulated steam
expands until a thin film of slightly cooler bitumen holds the bubble
intact through its surface tension.
5. During expansion, the surface tension of the bitumen film
counteracts the ever diminishing steam pressure until a state of
equilibrium is reached. Due to the low thermal conductivity of bitumen
and water, the bubble can remain stable for a few seconds.
6. So many bitumen bubbles form to together produce foamed bitumen.
7. As the colloidal mass cools to ambient temperature, the steam in the
bubbles condenses causing bubbles to collapse and the bitumen foam to
“decay”.
24.1.7 Heat of solution
Observe cooling when sodium thiosulfate (hypo), ammonium chloride or
ammonium nitrate are added to water. Observe heating when sulfuric acid
is added to water. Equal weights of water and ammonium nitrate will
lead to dissolved in water.
24.1.8 Heat of crystallization
1. Prepare a supersaturated solution of sodium acetate or sodium
sulfate and drop in a crystal to trigger crystallization A thermocouple
will show the change in temperature. Observe the change in temperature
of a sodium thiosulfate solution as it crystallizes.
2. Observe heating
when a flask of supercooled hypo solution crystallizes, or sodium
acetate or sodium hyposulfate crystallizes.
3. Examine a sodium acetate self-heating pack. Reusable heat packs
are used for medical purposes to provide heat externally to parts of
the body. The commercial heat pack consists of a sodium acetate
solution and a copper disc sealed in a plastic envelope. The patient
can bend the disc by squeezing it through the heat pack to cause
friction between the disc and the sodium acetate to change the liquid
solution to become solid in an exothermic reaction. The reaction
generates heat in seconds up to 54oC. The crystals in the
pack are returned to the liquid state by immersing the pack in boiling
water for about 10 minutes.
24.1.9 Water crystals in soap film
A ring with a soap film is cooled in a chamber surrounded by dry ice on
the overhead projector Water crystals form water crystals in soap film.
24.1.10 Crystal growth on the overhead projector
Melt together tartaric acid and benzoic acid and observe the crystal
growth on cooling between crossed Polaroids on the overhead projector.
24.1.11 Metglas
Metglas is a metal that has been quenched from liquid to solid without
crystallization. It has interesting electrical and magnetic properties.
24.1.12 Wood's metal
This fusible alloy contains one part of cadmium, two parts of tin, four
parts of lead, seven parts of bismuth and melts between 66oC
and 71oC.
24.1.13 Hot water freezes faster than cold water,
the Mpemba effect
Hot water, e.g. 90oC does appear to freeze faster than the
same amount of cold water, e.g. 18oC. This phenomenon was
observer by Aristotle 4th century BC and is named after Erasto Mpemba a
schoolboy in Tanzania who drew attention to it. The
initially hot water has less of the apparently frozen ice solid because
it contains trapped liquid water. The initially cold water freezes at a
lower temperature to a solid ice with less included liquid water.
The initial lower temperature causes intensive nucleation and a
faster crystal growth rate. At a freezing temperature of -6o
C the initially hot water apparently freezes first but eventually
the initially cold water completely freezes before the initially
hot water. Initially cold water will have the maximum
concentration of such 20-face clusterings. that do not easily allow the
rearrangement of water molecules for formation of hexagonal ice
crystals. Initially hot water has lost its ordered clustering, so if
the
cooling time is short, clustering will not occur fully before
freezing.
24.1.14 Molal freezing
point constant of cyclohexane solvent
Find
the freezing point of cyclohexane. Fit a test-tube with a stopper,
thermometer and stirrer. Weigh the test-tube and stopper, add 10 mL of
cyclohexane, C6H12,
and weigh again. Put the test-tube in a water and ice mixture and
record the temperature that cystals form. Add 0.5 g of
p-dichlorobenzene,
C6H4Cl2, then warm and stir the
mixture to aid dissolving. Put the
test-tube into a water, ice and salt mixture and record the
temperatue when the first crystals appear. Calculate the molal freezing
point constant for cyclohexane, the change in freezing point caused by
1 mol of p-dichlorobenzene per kilogram of cyclohexane.
The freezing point depression, dtf = i X Kf X
m, where freezing point
depression = (freezing point of pure solvent - freezing point of
solution), m = molality of particles in the solution (molality = moles
of solute / weight of solvent, kg), and Kf = the molal
freezing point depression constant for the solvent = 20.2oC
kg / m for cyclohexane. Repeat the experiment with cyclohexanol, C6H11OH,
as the solvent.
24.1.15 Molar mass of
solute from depression of freezing point, molar mass of alcohols
Measure the freezing point of known mass, kg, of a solvent, t1,
with known freezing point constant, Kf,
Add known mass of a solute, kg, with known Van't Hoff factor, e.g. i =1
if solute is molecular
Measure the freezing point of the solution, (solute in solvent) t2
freezing point depression, dtf = iKfm, so molality = (t1 -t2) / Kf
and molality of the solution = moles of solute / kg of solvent.
Molecular weight = mass of solute / moles of solute
1. Add 10 g of a known alcohol to 100 g of water. Put the mixture in a
container of dry ice from a mobile ice cream seller. Be careful! Dry
ice is frozen carbon dioxide. Do not touch it or lick it because
it will cause severe burns. Record the freezoing point of the mixture.
Repeat the experiment with other knwn alcohols. Draw a graph to show
the relationship between molecylar weightof alcohol and depression of
freezing point.
dtf = Kf m (for alcohols the Van't Hoff factor, i = 1)
2. Record the freezing point of 100 g of water. Record the freezing
point of a mixture of 30 g of an unknown alcohol in 100 g of
water.
m = dtf / Kf, where m = moles of solute / kg solvent, Kf
water = 1.86 oC kg / m
Molecular weight of unknown alcohol solute = weight of unknown
alcohol / m, moles of solute.
24.2.1 Latent heat of steam
When liquid becomes steam, its energy increases and needs more energy
provided by outside. The boiling is the phenomenon of violent
vaporization, at the moment of boiling the temperature of the liquid
remains. Pour a certain amount of liquid into a beaker. Heat the beaker
in a place where there is not wind flow and note to stir the liquid by
a thermometer. Record the increased temperature of the liquid every 10
seconds. Take time as horizontal axis, increased temperature as
vertical axis, mark the values in experiment recording and graph a
time / temperature graph, connect recording points to form a smooth
line.
The points in the graph shows the average increased temperature of the
liquid in every unit time. Use the mass of the liquid multiply its
specific heat, again multiply the increased temperature in every unit
time use ed from the above graph, you can calculate the amount of heat
absorbed by liquid every minute. As the amount of heat absorbed by
beaker is far less than that of liquid absorbed, so the former is
omitted in calculation. Heat the liquid continuously until it boils.
After that although heat it but the temperature of the liquid no longer
increases and produces the steam, i.e. some liquids become gases. After
boiling 10 minutes, stop heating and cool the liquid until its
temperature becomes back to that of indoors. Measure the weight of the
liquid again. Calculate the loss mass of the liquid that is also the
mass becomes steam. Use the total amount of heat absorbed in 10 minutes
by liquid divides the loss mass of the liquid, that is the amount of
heat needed for liquid every unit mass to becomes steam, called the
latent heat of the steam of the liquid.
24.2.1.2 Boil a kettle
with a spout
Heat water in a teak kettle with a spout. Look carefully just outside
the spout of a boiling tea kettle. You see nothing because water
vapour is colourless. However, you see steam short distance away
from
the spout you see condensed vapour with drops big enough to be seen.
24.2.2 Heat of condensation of water
See diagram 24.3.3
If an object, specific heat s, mass m, temperature (t1),
placed in steam. Collect all liquid produced by condensation. The
amount of heat emitted by condensation of steam are used to increase
the temperature of the object. Measure the mass of the liquid m1 and
the temperature t2 increasingly, from: ms (t2- t1)
= lambda x m So lambda = m (t2- t1) / mL. Use a
test-tube diameter of 4 cm, height 20 cm, with a rubber stopper. Punch
a hole at the side of the centre of the stopper. Insert a thermometer.
Use another small test-tube with length is less than 10 cm. Use a
copper wire with diameter approx. 1 mm. Wrap it around the upper part
of the small test-tube to allow the test-tube to be inserted freely.
Stretch it to the lower part of the tube. Wrap one circle again. Bend
it upward to make a small hook. The frame should allow the small
test-tube be at the centre of the large test-tube and not touch the
bottom. Hang a 4 cm iron nail with a thread. The other end of the
thread is taped to the centre of the bottom of stopper with white tape.
The length of the thread should make the nail higher than mouth of the
small test-tube after putting it into the large test-tube. The
thermometer neither touches the wall of the large test-tube nor touches
the nail and small test-tube. Fix the whole apparatus to a stand. Add
suitable amount of water into the large test-tube. Dry the small
test-tube, then measure its mass m1. Put small test-tube
into the copper wire frame, then put it into the large test-tube.
Measure the mass of nail m, then tie it to the thread and tape it to
the rubber stopper. Heat the water in the large test-tube, drive the
air in the tube completely. Absorb the possible little water in the
small test-tube by blotting paper. Insert the thermometer into rubber
stopper, record the temperature in the room t1. Then hang
the stopper right above the large test-tube. After the nail is stable,
cover the mouth of the tube slowly downward. Check if the nail is just
opposite to the small test-tube. Observe that water drops into the
small test-tube from the nail. Record the temperature t2 as
water does not drop completely. Take out the small test-tube, absorb
water outside, then measure the mass, m2. Calculate lambda =
ms(t2 - t1) / (m2 -m1)
24.2.3 Latent
heat of vaporization of
water
See diagram 24.2.3: Temperature / energy
(time) heating curve for water | See diagram
24.3.4
Use a beaker, a test-tube, a thermometer and an alcohol burner. Pour
water into the beaker and test-tube, 150 mL in beaker and 50 mL in a
test-tube. Put them on a supporter, adjust the height to allow the
test-tube is near the bottom of beaker. Measure the temperature of
water in beaker, then in test-tube by a thermometer which is held by a
clip. Heat the water in beaker over the alcohol burner, measure the
temperatures in beaker and test-tube each continuously. Especially
after the water boils, see if the temperature in beaker is equal to
that in test-tube? If water in test-tube boils? Keep heating until
water in beaker has completely been dry. Observe if water in test-tube
can boil finally? The temperatures in the beaker and test-tube read
from thermometer in the whole process of boiling may neither be 100oC,
this is due to thermometer itself exists error and the atmospheric
pressure is not just standard pressure. However, the temperatures
inside and
outside the tube are always equal to each other. It is they are in a
state of heat equilibrium, the conduction of heat has been stopped that
the water in conduct tube cannot use the heat of vaporization to boil.
24.2.4 Pressure and boiling point of water
See diagram 24.4.1
Use a thermometer which measuring range is over 100oC, a
pressure cooker and a aluminium cooker with a lid on which there is a
hole. Add water in both cookers. Cover the pressure cooker, take off
the valve used to reduce the pressure. Cover the aluminium cooker,
tighten the lid by several clips. Heat the two cookers over a roaring
fire until water in two cookers boils violent. Wear a pair of gloves,
hold the top of the thermometer, put the measuring bulb of thermometer
into the steam outlet on pressure cooker first, then into the hole on
the lid of aluminium cooker. The thermometer stays in steam until the
liquid column in thermometer does not rise up. Record from the
thermometer in two cases. The readings you have read may have
difference from boiling point of water in the cooker, but the
difference between two readings shows the difference of temperatures
caused by pressure in the cooker. Pour flask with half full of water,
cover a stopper which is inserted with a thermometer. Put it on a
frame, heat it over an alcohol burner until the water boils, record the
temperature at the moment. Then remove the flask from the heat source
by using a test-tube clip. Place the flask over an empty container.
Wait for a while, as the readings in thermometer begins to drop (do not
open the lid), pour cold water on the flask, observe that the water in
the flask boils again. Record the temperature again. As the steam
condenses while it becomes cold, the pressure in the flask reduces, so
water boils again. Pour a flask with half full of water, insert a
stopper with two holes. One hole is for inserting thermometer, the
other is for inserting a tube used for conducting out. Insert a
thermometer into water in the flask. Insert a short glass tube, the end
downward must be far from the surface of water, connect the upper end
of it a pump apparatus with a rubber tube. Fix the flask stable with a
pad, then begin to pump air out, Observe that with the pressure in the
flask decreases, first you observe some air bulbs rise up from water,
then water begins to boil. Record the temperature of water in the
flask. If temperature of water is very high, pump only by using a large
syringe. This time, to consider the second pumping, you should
add a rubber tube screw clamping apparatus on rubber tube as shown in
the diagram. After pump screw the clamping tightly, then take out the
syringe from rubber tube. Push out air inside the syringe (i.e. push
the inner cylinder to the bottom of the outer one), connect the rubber
tube again. Unscrew the clamping, pump secondly. If temperature of
water is merely that in the room, or to avoid the trouble of
pumping secondly, the pump apparatus is better to be the pump or air
pump.
24.2.5 Temperature of steam above water boiling
See diagram 24.4.2
Do this experiment behind a safety screen. This experiment can be very
dangerous because the glass flask could explode! Heat the flask with
the pinch clip open. When boiling commences at 100oC, close
the clip just long enough to see the temperature start to rise, then
open the clip and turn off the gas.
Temperature of boiling water or rather the temperature of the steam
above water boiling under pressure greater than in the normal
atmosphere. The steam is not allowed to escape and the thermometer
reading rises. In a pressure cooker, the pressure of the steam above
the boiling water can be twice normal atmospheric pressure and the
boiling point can be 120oC. The high pressure cooks the food
more quickly.
24.2.6 Set fire to paper with a jet of steam
See diagram 24.4.3
The steam from the boiling can passes through the copper pipe with the
Bunsen burners under it and becomes very hot.
24.2.7 Supercooled water
Cool a small test-tube of water in a Peltier device or dry ice/ alcohol
bath and use a thermocouple to record the temperature. Shake to freeze
and the temperature will rise. A Peltier device is a very small solid
state device that functions as heat pump made of layers of ceramic
plates and bismuth telluride. Apply d.c. current to move heat from one
side to the other so the cold side can cool a small electronic device.
They contain no moving parts and no Freon refrigerant.
24.2.8 Broken bottle, ice bomb
(Dangerous experiment not suitable for schools!)
1. Completely fill a small bottle with water, wrap it in two old tea
towels and put it in the freezer. After a few days remove the broken
bottle.
2. Completely fill a cast iron bomb with water, put in a
dry ice / acetone mix at -77oC, then put this in a strong
box. Be careful! When the water in the bomb freezes, the cast iron bomb
explodes showing that the volume of ice is greater than the volume of
an equal mass of liquid water and that freezing water produces great
pressure.
24.2.9 Liquefying dry ice
Press down on a piston on dry ice in a clear tube until at 5
atmospheres liquidation occurs. Put some carbon dioxide in a small
transparent syringe and squeeze to liquefy.
24.3.0 Cooling by evaporation
Kinetic theory,
compare evaporation and boiling, condensation and sublimation
Evaporation is the process in which a liquid turns to a vapour without
its temperature reaching boiling point. At any time, a proportion of
its molecules will be fast enough (have enough kinetic energy) to
escape from the attractive intermolecular forces at the liquid surface
and into the atmosphere. As the mean kinetic energy of the molecules of
the liquid rises, the number of the molecules of the liquid possessing
enough energy to escape rises. So the rate of evaporation rises with
increased temperature. Condensation is the conversion of a vapour to a
liquid as it loses heat. This is frequently achieved by letting the
vapour come into contact with a cold surface.
24.3.0.1 Sublimation
Sublimation is the change
of state from solid to vapour, or from vapour to solid, which occurs
without passing through an intermediate liquid state. Under normal
temperature and pressure, only ice, dry ice, camphor, sulfur, phosphor
may sublimate. In the formation of white frost, for example, water
vapour in the atmosphere passes directly from the vapour state to the
solid crystals of ice that make up frost.
24.3.1 Water "lost" by evaporation
Understand evaporation is a ubiquitous phenomenon and study the factors
related to evaporation sheep. Wet a towel then screw it until no water
drops. Hand it on a hanger and make it flat. Use a stick and a piece of
rope. Tie the rope to the middle of the stick. Hang the stick up with
the rope. Prepare a weight and two pieces of short rope. Tie the weight
and the towel to separate two sides of the stick. Adjust their
positions to make the stick balance. Leave them aside without change in
poison for one hour. What happened? Where did the water of the towel
go? Add else weight on the hanger to make the stick balance again. Take
the added weight on a beam to measure its mass. The mass is just that
of lost water. Repeat the experiment but place an electric radiator
near the towel. Repeat the experiment again but let the towel pile
together when hang it on the hanger. The evaporation speed is related
to the temperature and the evaporation area. The higher the
temperature, faster the evaporation speed, the larger the evaporation
area, faster the evaporation speed. On the contrary the evaporation
speed is slower.
24.3.2 Factors which affect the rate of
evaporation
See diagram 24.2.2
Fill a large flat dish with half full of water. Use another container,
the diameter of which is smaller than that of dish, height of which is
higher than that of dish, fill the same amount of water in it. Mark the
level of water, arrange them side by side in a place where the
temperature and case of air flowing are the same. The next day, observe
in which container the level of water is lower. If the level of water
varies a little, you may wait for another day.
24.3.3 Evaporate dichloromethane
Float a small test-tube of dichloromethane in a larger beaker of water.
Inset a glass tube into the dichloromethane and blow down the tube. The
volatile dichloromethane evaporates at room temperature, loses latent
heat of vaporization and some touching the small test-tube freezes.
24.3.4 Evaporate from the blackboard
1. Use a moist sponge or cloth wet two areas of equal size some
distance apart on a blackboard surface. Then fan one area with a piece
of paper board, leave another to evaporate without fanning. Observe
which area evaporates first.
2. Fasten a piece of cloth over a wooden frame that is about 30 cm2
and 3 cm thick. Wet the cloth, but do not drop water. Next make two wet
areas on blackboard surface. Cover one wet area with a frame with wet
cloth, leave the other area open. After a few moments, remove the frame
and observe which of the two areas evaporates faster.
24.3.5 Cryophorous
A cryophorus shows the freezing of water by its own evaporation. It has
two glass bulbs, connected by a glass and contains only water and water
vapour, and no air. The water is in one of the bulbs is cooled below
0°C in a ice / salt mixture or an alcohol / dry ice mixture, the
equilibrium shifts to produce more water vapour. This cools the water
at the top of the tube and it quickly freezes.
24.3.6 Freezing by evaporation
Freeze water in a watch glass over a dish of sulfuric acid in a bell
jar.
24.3.7 Drinking bird, drinking duck, dippy bird,
dunking bird
See diagram 24.3.7 | See
diagram 24.3.7: Drinking bird, photograph of commercial product
1.
The drinking bird is a small heat engine that converts a temperature
difference into cyclical motion. It is not a toy, and may be
adjusted to
work properly by moving the sleeve up or down the connecting tube.
However, it contains toxic and hazardous liquids, so
only the teacher, not the students, should make any adjustments to it.
Some commercial versions of drinking bird do not work well below
a room temperature of 25oC!
2. The drinking bird consists of two glass bulbs for head and body
connected
by a glass connecting tube that dips down into a volatile liquid in the
body, usually
methylene chloride (dichloromethane, CH2Cl2,
"methelene chloride"). Previously ether, Freon and other volatile
liquids were used. A sleeve attached to a pivot, fulcrum, clasps the
glass tube. Each
end of the pivot is inserted into holes at the top of the legs. The
head bulb contains vapour and is covered with
absorbent material, e.g. felt, to form a head covering and a beak.
3. Put a container full of cold
water in front of the drinking duck such that the height of the sides
of the container are just lower than the height of the pivot. Hold your
warm fingers around the body bulb to increase the vapour pressure and
force liquid up the connecting tube into the head bulb. The centre of
gravity lowers so the head tips forward. The beak and head becomes
immersed in the water which the absorbent material absorbs. The lower
end of the connecting tube tube rises above the level of liquid in the
body. Bubbles of vapour pass up through the connecting tube to
equalize the vapour pressure in the head bulb and body bulb. Liquid is
displaced down the
connecting tube to re-enter the body. The centre of gravity rises and
the head
of the drinking bird
rises again. The body bulb + liquid is heavier than the head bulb.
4. As the water in the absorbent material evaporates, vapour in the
head loses latent heat of vaporization, becomes cooler, and condenses.
The vapour pressure
in the head bulb decreases relative to the vapour pressure in the body
bulb. Liquid rises up the connecting tube
and enters the head bulb. The head becomes heavier than the body. The
drinking duck tips down and the
absorbent material in the beak absorbs water from the container. Water
starts to evaporate from the wet absorbent material and the cycle
starts again.
5. A similar apparatus called a "hand boiler" or "libido detector" is
held in the palm of the hand.
24.3.8 Cooling by evaporation, ether or ethyl
chloride
Use ether or ethyl chloride is used to freeze water in a small dish or
cool a thermometer.
24.3.9 Water bag, Coolgardie safe
People living in the bush keep water in a canvass water bag hung under
a tree in the shade or attached to the bumper car of a vehicle. Water
seeps through the canvass and evaporates from its outside surface to
keep the water cool.
24.3.10 Evaporation
from a microscope slide
Put a drop of a volatile liquid, e.g. propanone, on a microscope slide.
Observe the drop and note the time for the drop to evaporate. Note the
comparative time for evaporation under different conditions, e.g. drop
spread out, drop on warmed microscope slide, drop cooled by fanning,
drop warmed by your blowing on it.
24.4.1 Sling psychrometer to find relative
humidity.
Two thermometers one with a wet wick are mounted on a device swung
around the head.
24.4.2 Dew point measurement
Evaporating alcohol cools a shiny surface until dew forms.
24.4.3 Condensation nuclei
An extinguished match is held in the steam from a tea kettle. Place
moistened cotton in a bell jar and evacuate until fog forms.
24.7.1 Steam turbines
See diagram 24.4.1
Attach a sharpened pencil with eraser end to the clamp which attached
to the ring stand. Cover a test-tube downward on top of the pencil.
Then cut a circular piece of cardboard about 6 cm in diameter, punch a
hole exactly the same to test-tube in diameter to fit it
tightly around the test-tube. It is better to tape it to the test-tube.
Use 10 tooth sticks, glue them around the edge of the cardboard. Note
to let each of them half out of the cardboard. Cut 10 piece of
aluminium foil each takes a shape of square. Fold each one in half to
get a shape of rectangular, tape each on the tooth stick which should
be in centre of the folded surface and their inclined direction should
be the same. Pour water into an Erlenmeyer flask, cover a rubber
stopper with a bent glass tube tightly. Put the flask on the tripod and
heat it over an alcohol burner. At the same time, adjust the position
of the tube clip to make aluminium foil blades are just opposite to the
open end of the bent glass tube and at the same level. As the water in
flask boils and the steam has produced, observe what happens.
24.7.2 Refrigerator, ice chest,
portable ice boxes, expresso coffee
machine
See diagram 24.7.2
1. In a refrigerator, food loses heat by conduction, convection and
radiation to the air, shelves and inner cabinet. Volatile refrigerant
fluids in a coil around the freezer evaporate and so take in latent
heat to cause cooling. The vapour from this evaporation is removed by
an electric pump to the black heat exchanger coil with cooling fins
behind the refrigerator which radiates heat to the surrounding air.
When the electric pump removes vapour from the coil around the freezer
it also reduces the pressure in the coil and lowers the boiling point
and induces more evaporation. In the heat exchanger coil the vapour
compresses, loses latent heat of vaporization, and is returned as a
liquid to the coil around the freezer. The electric pump can be
adjusted to switch on and off to control the temperature of the
refrigerator. Industrial refrigerators use ammonia and Freon-22
refrigerants and other gases, e.g. carbon dioxide. Domestic
refrigerators use less toxic refrigerants, e.g. Freon 11.
2. Ice
chests and portable ice boxes, e.g. "Esky" cooler, are
insulated boxes containing an ice block or ice in small cubes. Heat is
lost from food and drink by conduction if touching the ice, convection
currents and radiation into the ice.
3. Expresso coffee machines make steam that can be passed through
water containing ground coffee or milk. The steam condenses quickly
liberating latent heat to brew coffee quickly and froth the milk for
cappuccino or other milk coffees.