Chemistry, boiling point, melting point, solubility.

School Science Lessons
(UNChem1)
2024-08-30
Contents
3.1.0 Boiling point
3.2.1 Heat substances that decompose and lose mass when heated
3.3.0 Melting point
3.5.0 Shrinking volume
3.6.0 Solubility and solutions
3.7.0 Thermal decomposition

3.1.0 Boiling point, BP
3.1.1 Boiling point of liquids
3.1.2 Boiling point of inflammable liquids
3.1.3 Boiling point of sodium chloride solution
3.1.4 Boiling point of two liquids, water and alcohol
3.1.5 Boiling point of water
3.1.6 Elevation of boiling points, ebullioscopy constant, kB
3.1.7 Leidenfrost effect
3.1.8 Volatility of different liquids

3.3.0 Melting point
3.3.1 Ice melts, de-icers
3.3.2 Impurities affect the melting point of a substance
3.3.3 Melting point, MP, of solids
Experiments
3.3.4 Melting point and cooling curve of stearic acid
3.3.5 Melting point experiments, Octadecan-1-ol
3.3.6 Melting point and pressure
3.3.7 Melting point of 1,4-dichlorobenzene
3.3.8 Melting point of ice and freezing point of water
3.3.9 Melting point of naphthalene
3.3.10 Melting points of naphthalene with a capillary tube
3.3.11 Melting point of substances, (candle wax, urea, cetyl alcohol)
3.3.12 Temperature at which ice melts
3.3.13 Temperature at which ice and salt mixture freezes.

3.5.0 Shrinking volume
3.5.1 Container holds more
3.5.2 Container not leaking
3.5.3 Shrinking volume of flask
3.5.4 Shrinking mixture of liquids, lost volume

3.6.0 Solubility and solutions
3.6.1 Heat of solution
3.6.2 "Magnetic" sugar cube dissolves
3.6.3 Miscible liquids
3.6.4 Solubility and agitation
3.6.5 Solubility and particle size
3.6.6 Solubility and solvents
3.6.7 Solubility and temperature, solubility of salts in water
3.6.8 Solubility in water at a given temperature
3.6.9 Solubility in water of different salts

3.7.0 Thermal decomposition
3.7.2 Decomposition of boric acid
3.7.3 Decomposition of carbonates
3.7.4 Decomposition of chlorates
3.7.5 Decomposition of chlorides
3.7.6 Decomposition of dichromates
3.7.7 Decomposition of ferricyanides
3.7.8 Decomposition of hydrogencarbonates, bicarbonates
3.7.9 Decomposition of hydrates, hydrated salts
3.7.10 Decomposition of hydroxides
3.7.11 Decomposition of manganates
3.7.12 Decomposition of metals, metallic salts
3.7.13 Decomposition of nitrates
3.7.14 Decomposition of oxalic acid
3.7.15 Decomposition of oxides
3.7.16 Decomposition of phosphates
3.7.17 Decomposition of sulfates
3.7.18 Decomposition of sulfites

3.1.1 Boiling point of liquids
Evaporation occurs only at the surface of liquids, but may occur at all temperatures.
When boiling occurs, bubbles form inside the liquid and the liquid boils at a definite temperature, the boiling point depending on the pressure.
When boiling occurs the temperature of the liquid remains constant until all the liquid has evaporated.
The liquid boils when the saturation pressure is equal to the pressure acting on the surface of the liquid.
Air bubbles form first as small nuclei increasing in size with temperature.
When bubbles of vapour form in a boiling liquid, the vapour pressure of the gas in the bubbles is greater than atmospheric pressure.
When the chemical bonds between liquid molecules are strong, only a few molecules can break the bonds to become a vapour.
Smaller molecules usually have lower boiling points.
The boiling point of a liquid is the temperature at which the liquid boils when exposed to the atmosphere.
So the boiling point of a liquid is the temperature at which the vapour pressure of the liquid equals the pressure of the atmosphere, 1 atmosphere.
At 100oC, the vapour pressure of pure water is one atmosphere (101.325 KPa, kNm-2).
The boiling point varies with pressure, so water boils at <100oC on high mountains.

3.1.2 Boiling point of inflammable liquids
See diagram 3.6: Boiling point of inflammable liquids.
A rubber band, B capillary tube, C test-tube, D inflammable liquid, E thermometer
1. Do not use a Bunsen burner to find the boiling point of inflammable liquids, e.g. ethanol (B P 78.4oC) and acetone (B P 56oC).
Use an electric hot plate or use the following method.
Pour 2 cm of the inflammable liquid into a test-tube in an empty container.
Place a thermometer in the test-tube with its bulb in the liquid.
Boil water in an electric jug or on an electrical hot plate.
Pour the hot water into the container so that the level is higher than the inflammable liquid in the test-tube.
Stir the inflammable liquid gently with the thermometer and read thermometer when the inflammable liquid boils.
It is not good practice to stir liquids with thermometers!
2. Use a very small test-tube or seal one end of a piece of glass tubing, 8 cm length and 3 cm external diameter.
Put the inflammable liquid into this test-tube.
Put a capillary tube, sealed at one end, into the inflammable liquid with the sealed end up and the open end down in the inflammable liquid.
Use a rubber band to attach the test-tube containing inflammable and capillary tube to the bulb of a thermometer.
Hold the apparatus in a container of water and heat gently with an electric hot plate.
When the temperature rises, bubbles slowly come out of the capillary tube.
At the boiling point the bubbles suddenly come out as a steady stream.
Read the temperature.
Let the water cool and read the temperature again when the steady stream of bubbles stops.
Calculate the boiling point as the average of the two readings.

3.1.3 Boiling point of sodium chloride solution
1. A solution of sodium chloride in water boils at a higher temperature and has a lower freezing point than pure water.
Use freezing points and boiling points to find the purity of substances.
Use three test-tubes containing the same volume of water.
Add some sodium chloride to the second test-tube.
Keep adding sodium chloride to the third test-tube until no more dissolves to produce a saturated solution at that temperature.
Join the test-tubes with an elastic band.
Heat the test-tubes equally over a Bunsen burner.
The first test-tube containing only water boils first.
The second test-tube containing some sodium chloride boils next.
The third test-tube containing the saturated solution of sodium chloride boils last.
However, it is reported that the addition of 20 g of salt to 5 litres of water increases the boiling temperature to only 100.04oC.!
2. Put a beaker containing demineralized water in a broad pan containing a concentrated salt solution.
Slowly heat the broad pan and note that the demineralized water boils first.

3.1.4 Boiling point of a mixture of two liquids, water and alcohol
Table 7.5.3
Liquids Solution 1 Solution 2 Solution 3
1. Water 25% 50% 75%
2. MS 75% 50% 25%
MS = Methylated spirit
Use the above method to compare the boiling points of mixtures of water and methylated spirit in the listed proportions
BE CAREFUL! Alcohol is Highly flammable!
Use an electric hotplate, NOT a Bunsen burner. Table 7.5.4

Liquids	Solution 1	Solution 2	Solution 3
1. Water	25%	50%	75%
2. methylated spirit	75%	50%	25%

3.1.5 Boiling point of water
See diagram 3.5: Boiling point of water.
1. Pour water into a test-tube.
Hold a thermometer with its bulb just under the water.
Heat very slowly by moving the test-tube in and out of the flame or add boiling chips, anti-bumping granules.
Heat the water gently until it boils.
Record the temperature.
Note the same temperature in all parts of the test-tube.
Note any change in the reading if the thermometer touches the bottom of the test-tube.
The water must cover the bulb of the thermometer and the bulb must not touch the sides of the test-tube.
2. Show that the boiling point of water does not depend on the size of the container.
Repeat the experiment with a large container.
Heat the water quickly.
The water first starts to boil near the bottom and sides of the container.
Note the temperature in different parts of the container.
Note any change in the reading if the thermometer touches the bottom of the container.
The boiling point is the same in small and large containers.

3.1.6 Elevation of boiling points, ebullioscopy constant, kB
The boiling point of a liquid is raised if substances are dissolved in it.
The elevation is proportional to the number of particles or molecules or ions dissolved in the liquid.
The elevation in oC = the molal concentration of the solute in the liquid × a constant.
The value of the constant (ebullioscopy constant, kB) depends on the solvent.
So if kB is known, the molecular weight of the solute can be calculated.
Addition of substances to water used for cooking has little effect on cooking time, because adding a tablespoon (20 g), of sodium chloride (common salt), to 5 litres of water increases the boiling point of pure water by only 0.04oC.
The presence of a solute both raises the boiling point and lowers the freezing point of the solvent.

3.1.7 Leidenfrost effect, water drops on hot stove
An insulating layer of vapour may support liquids hotter than the liquid's boiling point.
Let drops of water fall on a very hot stove top.
The drops do not immediately evaporate or spread into a puddle, but jump around supported by a vapour layer below them caused by the high heat that evaporated some of the liquid.
The liquid floats on this vapour cushion and boils without bubbling.
If the hot surface suddenly cools, the vapour layer can collapse and the water bubbles explosively.
In chemical factories, there is a risk of explosion if water touches very hot metal, but this may be controlled if the metal has a rough texture.
Similarly, liquid nitrogen may mover erratically over the floor.
Foolish experimenters have demonstrated the Leidenfrost effect by dipping their fingers in molten lead with temperature > 450oC
Do not try to demonstrate this molten lead experiment in the laboratory!

3.1.8 Volatility of different liquids, perfume
Evaporation is the movement of particles from the surface of a liquid to the gas state, when below the boiling point.
Volatile liquids evaporate readily at room temperatures.
1. Select liquids from the laboratory, e.g. water, methylated spirits, gasoline, mineral turpentine, kerosene (paraffin oil), household oil, machine oil, car oil, vinegar, vanilla essence, eucalyptus oil, glycerine.
Wet a 5 cm piece of absorbent paper with a liquid.
Write the name of the test liquid in pencil.
Attach the piece of paper to a horizontal string.
Examine the paper every ten minutes, every two hours, and each day.
2. Repeat the experiment with perfumes.
Smell the paper every ten minutes, every two hours and each day.
Some perfumes soon disappear, but others last for days.
Record the relative "person-attracting" capacity for each perfume!

3.2.1 Heat substances that decompose and lose mass when heated
See diagram 3.30: Heat copper sulfate crystals.
The phrase "Substances that do not decompose when heated" refers to substances that remain stable after heating constantly with a Bunsen burner flame.
However, all compounds breakdown when heated to a high enough temperature.
1. Most carbonates decompose to form carbon dioxide and a metallic oxide.
Sodium hydrogen carbonate, NaHCO3 (sodium bicarbonate), begins to lose carbon dioxide at 50oC to form sodium carbonate.
A solution of a sodium hydrogen carbonate begins to lose carbon dioxide at 20oC.
2. Hydroxides decompose to form water and a metallic oxide.
3. Nitrates decompose to form oxygen, nitrogen dioxide and a metallic oxide, except potassium nitrate and sodium nitrate that form the nitrite and oxygen.
Lead nitrate decomposes at 470oC.
4. Nearly all oxides are stable, e.g. zinc oxide, ZnO, M P above 1, 800oC.
5. Some sulfates decompose to form sulfur trioxide and metallic oxide.
6. Metal compounds higher in the activity series are usually more stable than compounds of metals lower in the activity series.
7. Some salts first lose water of crystallization and then become stable.
8. The salts that remain stable when heated constantly with a Bunsen burner flame are calcium sulfate, potassium chloride, potassium sulfate, sodium carbonate, sodium chloride, and sodium sulfate.
Ammonium oxalate (NH4COO)2, and ammonium dichromate (NH4)2Cr2O7, decompose before melting.
Ammonium sulfate (NH4)2SO4, decomposes above 280oC.
9. Boric acid, H3BO3, loses water until it decomposes to the anhydride, B2O3.
10. Oxalic acid begins to sublime at 100oC, becomes anhydrous at 189oC and when heated rapidly decomposes into carbon dioxide, carbon monoxide, formic acid and water.
11. Potassium chlorate, KClO3, decomposes above 368oC into potassium perchlorate and oxygen.
Potassium ferricyanide, K2Fe(CN)6, decomposesbefore melting.
12. Monosodium orthophosphate, NaH2PO4.H2O, and disodium orthophosphate (disodium hydrogen phosphate (V), Na2HPO4.12H2O, lose water of crystallization.
Experiments
1. Heat copper sulfate crystals.
Put 4 cm of crushed blue copper (II) sulfate crystals in a dry test-tube fitted with a one-hole stopper and delivery tube.
Heat the dry test-tube and crystals gently.
Note whether vapour collects on the cooler parts of the dry test-tube and whether any liquid collects in the receiving test-tube.
Observe any change of colour of the crystals from blue to white.
Identity the liquid in the receiving test-tube by measuring the boiling point.
When all the copper (II) sulfate crystals have become white and the dry test-tube has cooled, pour the liquid in the receiving test-tube back on the white crystals.
Note whether the blue colour of the crystals is restored and if any heat is given off.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous copper (II) sulfate + water.
2. Prepare test-tubes containing 1 cm of (a) iodine crystals (b) sodium hydrogen carbonate granules or crystals (c) silica sand (d) zinc oxide.
Fix a cotton wool plug in the mouth of each test-tube to prevent loss of solid during heating, then weigh each test-tube.
Heat each test-tube and cotton wool plug thoroughly and weigh it again.
Note any change in weight, because of the loss of water of crystallization.
3. Put black, shiny crystals of iodine in an evaporating dish.
Cover the dish with a piece of filter paper and stand a filter funnel upside down on the dish.
Heat the dish gently.
Purple vapours rise through the filter paper.
As they cool in the filter funnel, shiny black crystals of iodine form again.
4. Heat sodium hydrogen carbonate crystals.
The crystals lose water and carbon dioxide, and at 100oC are converted to sodium carbonate.
5. Silica sand consists of pieces of silicon (IV) oxide (SiO2) crystals.
Heat sand in a crucible.
The sand particles may break up physically, but do not break up chemically.
6. Heat zinc oxide in a crucible.
The colour changes from white to yellow, but no change in weight occurs.
The substance does not decompose and does not gain anything from the air or lose anything to the air.

3.3.1 Ice melts, de-icers
1. Calcium chloride, can melts ice down to -31.7oC, and gives off heat as it dissolves to melt the ice quicker than other de-icers, but leaves a slimy residue, corrosive to metal, damages vegetation if over-application.
2. Magnesium chloride is less corrosive and safer on concrete and plants, fast-acting and more effective at de-icing than rock salt, over-application can damage plants and keep pavement wet. 3. Sodium chloride, rock salt, least expensive and very efficient, melts ice down to -6.7oC, dries icy surfaces, not harmful to concrete, but damages vegetation, corrosive to metal and may leach into soil and contaminate groundwater.
However, saltwater freezes at -18oC, so salt is an ineffective de-icer below this temperature.
4. Potassium chloride, more expensive than other ice melt products, but works well as 50/50 with rock salt.
It can melt ice down to -25oC, but can cause plant injury.
5. Urea, melts ice to -9oC, but may harm vegetation with excess nitrogen or may add useful nitrogen fertilizer to plants, frequently used by gardeners, but potential of runoff into water sources.
6. Calcium magnesium acetate, [CaMg2(CH3COO)6], (CMA), from dolomite limestone and acetic acid, not effective below -6.7oC, only slight affect on plants and concrete, does not form a brine-like salts, prevents snow particles from sticking together on road surface, prevents re-freezing and leave a slush.
When the freezing point of water is lowered by creating a brine, freeze/thaw cycles and expansion of freezing water can damage concrete.
However, calcium magnesium acetate has relatively few environmental impacts, because it readily biodegrades and does not persist in the environment.
7. Potassium acetate is a high-performing de-icer, which is biodegradable and non-corrosive.
8. Beet Juice is quicker and less toxic for melting ice on roads and it has less of an impact on the environment, because it is a natural material.
It can lower the melting point of water to -29oC.

3.3.2 Impurities affect the melting point of a substance
Mix stearic acid with the naphthalene to make the naphthalene impure.
Note changes in the melting point.
Impurities lower the melting point.

3.3.3 Melting point, MP, of solids
"DigiMelt is a modern, low cost, digital melting point apparatus designed specifically for the student lab.", (Commercial)
The melting point (MP or fusing point), is the temperature at which a solid starts to liquefy.
The melting point and freezing point of a pure substance are the same temperature.
Melting (fusion), is the "solid to liquid" type of phase change.
Other phase changes include change from liquid to solid (freezing), solid to gas (sublimation), liquid to gas (evaporation), gas to liquid (condensation).
Pure substances melt at constant temperature.
Impurities lower the melting point.
Impure substances, e.g. alloys, melt over a range of temperature.
The melting point graph for a pure substance is horizontal as the solid melts.
The melting point graph for an impure substance has an inclined line as the solid melts.
Melting points and melting behaviours can be used to identify a substance and decide if it is pure.
Melting is also called fusion and the melting point can be called the fusing point.
The melting point of glass is 1400-1600oC, depending on the composition of the glass.

3.3.4 Melting point and cooling curve of stearic acid
See diagram 3.2.2: Melting point and cooling curve of stearic acid
Stearic acid melts at 69oC.
Stearic acid (octadecanoic acid), is a safer alternative to using naphthalene for the following experiments:
1. Put 2 cm of the acid in a test-tube with a thermometer.
Put the test-tube in a beaker containing water.
Heat the beaker tube gently until the acid just melts.
Note the time and remove the test-tube from the beaker.
Note the temperature every 30 seconds as the substance cools.
The acid solidifies again at the same temperature.
Heat the test-tube again to free thermometer.
Draw the cooling curve of the acid.
Plot temperature on the vertical axis and time on the horizontal axis.
When the acid is changing from liquid to solid, the curve is horizontal.
2. For a more accurate method to measure melting points, heat one end of a 10 cm capillary tube to seal it and put octadecanoic acid in the capillary tube.
Attach the capillary tube to a thermometer with a rubber band.
Put thermometer in a beaker of water.
Heat the water while stirring with thermometer.
Record the temperature at which the acid melts.
Leave the capillary tube to cool.
Record the temperature at which the acid solidifies.

3.3.5 Melting point experiments, Octadecan-1-ol
Octadecan-1-ol, octadecyl alcohol, 1-octadecanol, Stearyl alcohol, C18H38O, long-chain primary fatty alcohol
Use octadecan-1-ol for melting point curve experiments.

3.3.6 Melting point and pressure
See diagram 24.1.3: Melting point and pressure.
See diagram 24.1.3a: Regelation.
Regelation is when ice melts under pressure and refreezing when pressure is reduced.
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.
Experiments
1. Regelation
Cut through a block of ice with a copper wire loop that has a heavy mass hanging from each end.
Copper wire cuts through faster than iron or cotton thread.
At the start of the experiment the copper wire should be the same temperature as the ice, otherwise it is just transferring heat and melting the ice.
When the copper wire exerts pressure on the ice below it the ice melts, because its melting point is now lower.
The ice under pressure is still at 0oC, which is above its melting point.
So the wire sinks down through the melt water.
However, above the wire the pressure decreases to near atmospheric pressure.
The melt water still at 0oC refreezes. because the melting point is again 0oC.
The refreezing water above the wire loses latent heat of fusion, which is conducted down to the copper wire and increases the melting process.
The copper wire keeps passing down through the whole block of ice, cutting it on two and joining the cut regions together again.
2. 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 freeze again, regelation.
Then you can lift the ice cube by nylon thread.
Substances that expand on freezing show a lowering melting point under pressure.
3. Tie the ends of a 50 cm length of iron wire or nylon cord to two 10 cm lengths of cut broomstick to act as handles.
Lean over a block of ice, grip the handles and use them to push the wire or cord down onto it.
Move the wire from side to side and with a downward motion, as if you are sawing on the ice.
Eventually you can move the wired or cord down through the whole block of ice, but it remains as one block of ice.
4. Push your finger down on the middle of an ice cube.
Some ice at the upper surface melts, because of the 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.
5. Squeeze crushed ice in a thick walled cylinder to form a solid block.
In a snowball fight, throw loose snow only.
If you squeeze the snow in your gloved hand, you create a rocky piece of ice that is dangerous to throw.
6. 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.
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.

3.3.7 Melting point of 1,4-dichlorobenzene
1,4-dichlorobenzene, C6H4Cl2, is used as a men's toilet deodorant
It occurs as perfumed crystals in the urinal.

3.3.8 Melting point of ice and freezing point of water
"Ice Melting Blocks", endothermic heat flow (toy product).
For pure substances, the melting point (MP) = freezing point (FP).
The melting point of ice at 1 atmosphere of pressure is 0°C (32 °F; 273 K ) = the freezing point of water.
For water, MP and FP = 0oC.
However, freezing mixtures of ice and salt have temperatures below 0oC.
The freezing point of water in motor car radiators is lowered by adding antifreeze solutions, e.g. ethylene glycol (ethane-1,2-diol), that does not freeze above -20.6o.
Freezing points can be used for detecting water in milk or other adulteration.
2. Use 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.

3.3.9 Melting point of naphthalene
See diagram 3.2: Approximate melting point of naphthalene.
Put 2 cm of naphthalene flakes in a test-tube.
Hold a thermometer with its bulb in the naphthalene.
Use a small flame to heat the test-tube gently and watch the thermometer reading.
To find the melting range, note the temperature when the naphthalene melts.
Leave to cool and note the temperature when the naphthalene solidifies.
To find the melting point, calculate the average of these two values.
The melting point of pure naphthalene is 80.5oC.

3.31.0 Melting point of naphthalene with a capillary tube
See diagram 3.3: Melting point of naphthalene with a capillary tube.
For a more accurate way of finding the melting point
Make a capillary tube by drawing out a glass tube over a hot flame.
Put a very small amount of naphthalene in a capillary tube sealed at one end.
Attach a thermometer to the capillary tube, a sealed end down.
Put the thermometer and capillary tube in a container of water and slowly heat the water while stirring with the thermometer and capillary tube.
Do not let water enter the capillary tube.
To find the melting range, note the temperature when the naphthalene melts, leave to cool, and note the temperature when the naphthalene solidifies.
To find the melting point, calculate the average of these two values.
Repeat the experiment with stearic acid (M P 69oC), palmitic acid (M P 63oC), 1.4-dichlorobenzene (deodorizer) (M P 53oC), paraffin wax (M P 45oC to 65oC), butter, soap.

3.3.11Melting point of substances, (candle wax, urea, cetyl alcohol)
Repeat the above experiments with the following substances:
1. Candle wax, MP = 45oC to 65oC,
Most candle waxes melt at about 60oC.
Do not melt candle wax over direct heat, because the vapour may ignite.
If it ignites, smother the flames with a lid, fire blanket, sodium carbonate powder, or moist towel, do not use water.
Melt candle wax in a heat resistant container in gently boiling water, in an electric frying pan or over a hot plate.
2. Urea, MP = 133oC Grind the urea to a fine powder mix in a clean, dry mortar and pestle.
Wash the mortar and pestle with soap and water, rinse with tap water, heat in a fume hood to clean and dry. Cinnamic Acid, an oxidation product of cinnamon oil, has a similar melting poit to urea, 132-133 oC.
3. Cetyl alcohol, MP 49.3oC
Cetyl alcohol, hexadecan-1-ol , palmityl alcohol, iH3(CH2)15OH, at room temperature, occurs as a waxy white solid or flakes.

3.3.12 Temperature at which ice melts
Half fill a beaker with tap water.
Note the temperature of the water after five minutes.
Put pieces of ice in the water.
Note the temperature every five minutes.
The temperature drops to zero and remains at zero while ice remains floating in the water.
Wait until all the ice melts.
The temperature rises again until it reaches room temperature.

3.3.13 Temperature at which an ice and salt mixture freezes
At 0oC, the molecules in pure water form strong bonds to form ice.
If sodium and chlorine are between water molecules, it is harder for these bonds to form.
Sea water contains about 35 grams of salt per litre and freezes at -1.8oC.
Mix crushed ice with salt.
Note the temperature after five minutes.
The temperature of the ice and salt mixture is below zero, e.g. -20oC, if 1: 3 ratio of sodium chloride to ice.

3.5.1 Container holds more
Fill a large beaker with marbles and note the top level of the marbles.
Slowly add sand to the beaker while tapping to make the sand settle between the marbles.
Tip out the sand water and measure its volume.
Replace the sand in the beaker of marbles.
Add water to the marbles and sand.
Tip out the water and measure its volume.

3.5.2 Container not leaking
Fill a measuring cylinder by 1 / 3 with water softener pellets (Calgon, sodium hexametaphosphate) or with table salt (sodium chloride).
Add water until full and mark the liquid level with a grease pencil or rubber band.
Leave it to stand for 5 minutes and note the liquid level.
The drop in water level is not, because water is absorbed by the salt.
Pour out some water before all the salt dissolves and refill it with fresh water.
The water level drops again, because the crystals of the dissolving salt breaks down into ions that can slip in between the water molecules and make the total volume decrease.
Repeat the experiment with other salts that dissolve in water and sucrose.
The sugar molecule is large and does not ionize when dissolved in water, so that the water level will not drop.

3.5.3 Shrinking volume of flask
1. Fill a small, narrow-necked flask with water to a level in the neck and mark this level.
Add sodium chloride to the water with continual shaking until the solution is saturated and no more dissolves.
Note the new level of the liquid.
The volume of the solution is only slightly greater than the original volume of the water.
2. Close one end of a glass delivery tube about 30 cm long.
Fix it upright, half fill it with water and mark the level of the water.
Slowly add alcohol to fill the delivery tube.
The water and the alcohol fill equal lengths in the tube.
Shake the tube thoroughly to mix the water and alcohol.
The new level of the solution in the tube shows a slight decrease in total volume.

3.5.4 Shrinking mixture of liquids, lost volume
See diagram 10.3.1: Lost volume.
1. Sodium chloride is an ionic solid that exists as ions.
The ionic structure breaks down in solution.
These ions can slip between the water molecules and make the total volume decrease.
Other salts that dissolve in water have the same property.
1.1 Put sodium chloride crystals or another soluble salt in a test-tube.
Add water until full.
The water level drops slightly as the crystals dissolve.
1.2 Put about 30 mL sodium chloride crystals into a measuring cylinder.
Add water until the measuring cylinder is exactly full.
After a few minutes, note that as the sodium chloride dissolves, the liquid level drops.
2. Repeat the experiment using sucrose crystals instead of sodium chloride crystals.
The water level does not drop, because the sugar molecule is larger than water molecules and does not form ions.
The sucrose crystals dissolve to form sucrose molecules that are much bigger than the sodium ions or chloride ions.
3. Half fill a test-tube with water.
While holding this test-tube at an angle, pour ethanol slowly from a beaker until the test-tube is full.
Hold the test-tube by placing your thumb on the mouth of the tube so that no air bubble is trapped.
The test-tube seems full.
Invert the test-tube several times while keeping a thumb on the opening.
Do not release the pressure.
The liquid level becomes lower.
The alcohol or water did not evaporate and no liquid spilt, because of inverting the test-tube.
By inverting the test-tube, mixing of water and alcohol occurs, and the alcohol molecules slip between the water molecules in the spaces between the molecules, thus making the total volume of the mixture become less.
The spaces between the molecules cannot be seen by the naked eye.
4. Repeat the experiment with methanol or rubbing alcohol (isopropanol).
Not all combinations of liquids give the strange shrinking illusion, but many water / alcohol mixtures do.
5. Make a quantitative estimate of shrinking using two identical, smaller, measuring cylinders, and a larger measuring cylinder with a volume at least the combined volume of the smaller, measuring cylinders.
Leave the smaller measuring cylinder to stand for some time in a place of no vibration.
Put 3 drops of a blue food colouring in the water.
Use the same procedure to put 3 drops of a yellow food colouring in the methylated spirits.
Observe the spread of the two food colourings.
Put the contents of the first measuring cylinder into the large measuring cylinder followed by the contents of the second measuring calendar.
Observe the colour change as the food colourings mix and note the final total level of the solution.
Use a pipette to bring the solution up to the total volume of the smaller measuring cylinders.
Calculate the % shrinkage, about 2 vols 1 / vol 2 = 1.9.
The ethanol also reduces the surface tension of the water.
6. Repeat the above experiment at different ambient temperatures.

3.6.1 Heat of solution
Dissolve some equal quantities of sodium hydroxide, potassium hydroxide, ammonium chloride and urea in separate test-tubes half full of water.
Feel the test-tubes and note any change in the temperature.
Sodium hydroxide and potassium hydroxides dissolve in water with an increase in temperature.
Ammonium chloride and urea absorb heat from their surroundings when dissolving in water.

3.6.2 "Magnetic" sugar cube dissolves
Fill a large dish with water.
Wait until the water is absolutely still then lower a matchstick into the centre of the water.
Carefully dip a sugar cube in the water near the edge of the dish.
The matchstick moves towards the dissolving sugar cube.
When the sugar dissolves in the surface water, the surface water becomes heavier and falls to be replaced by surface water flowing towards the sugar cube, carrying the matchstick with it.

3.6.3 Miscible liquids
Put 10 mL of water in three test-tubes.
Add 1 mL of: 1. methylated spirits, 2. glycerine, and 3. kerosene.
Shake each test-tube.
Miscible liquids can mix in all proportions.
Alcohol and water are miscible.
Glycerine and water are miscible.
Kerosene and water are not miscible, immiscible.

3.6.4 Solubility and agitation
Prepare two equal samples of cane sugar.
Put one sample of cane sugar into a test-tube half full of water.
Add a stopper and shake the test-tube until all the sugar dissolves.
Put the other sample of cane sugar into a test-tube.
Very slowly add the same volume of water as in the first test-tube.
Leave to stand.
Most of the sugar has not dissolved but, if left to stand for long enough, all the cane sugar will dissolve as in the first test-tube.

3.6.5 Solubility and particle size
Show that small particles dissolve faster than large particles.
1. Add coarse table salt to a first test-tube half filled with water.
Add the same quantity of fine table salt to a second test-tube that contains the same volume of water.
Shake both test-tubes equally and simultaneously.
Note the amount of undissolved table salt left in each test-tube.
2H">2. Use two equal samples of large crystals of copper (II) sulfate.
Grind one sample into a fine powder.
Put both samples into equal quantities of water in separate test-tubes and shake.
Compare the rates at that the different samples dissolve and cause the water to change colour.

3.6.6 Solubility and solvents
1. Fill two test-tubes one third full with water and another with methylated spirits.
To each test-tube add 1 g sodium chloride, attach a stopper and shake.
Sodium chloride dissolves readily in water, but not so readily in methylated spirits.
2. Add sodium chloride crystals to a dilute solution of sugar in water and note whether the crystals dissolve.
Drop crystals of potassium dichromate into the solution.
Note whether the solution changes colour.
Colour change shows that potassium dichromate is also dissolving.
The presence of one dissolved substance does not prevent other substances dissolving in the solution.
Unless the concentrations are high, one solute does not affect the solubility of other solutes in the solution.

3.6.7 Solubility and temperature, solubility of salts in water
The solubility of a potassium dichromate in 100 g of water varies with temperatures:
0oC to 5 g,
10oC to 7 g,
20oC to 12 g,
30oC to 20 g,
40oC to 26 g,
50oC to 34g,
60oC to 43 g,
70oC to 52 g,
80oC to 61g,
90oC to 70 g,
100oC to 80 g.
Experiment
Show that a saturated solution contains less dissolved solid at a lower temperature than at a higher temperature.
Make a 50 mL saturated solution of potassium dichromate or potassium nitrate at 60oC.
Pour the clear solution into a clean container and keep the temperature of this container at 40oC until crystals stop forming.
Pour the clear solution from this container into another clean container.
Do not pour crystals into the container.
Leave to cool and note more crystals forming as the solution cools.

3.6.8 Solubility in water at a given temperature
Add sodium hydrogen carbonate (sodium bicarbonate) to 100 g of water in a container while stirring.
Keep adding sodium carbonate until no more solute will dissolve.
Decant the clear saturated solution.
Read the temperature of the saturated solution, i.e. room temperature.
Weigh a clean evaporating dish, w1.
Add some clear saturated solution and weigh again, w2.
Carefully evaporate the solution in the evaporating dish to dryness and weigh again, w3.
The mass of the dissolved sodium hydrogen carbonate = w3 to w1.
The mass of water = w2 to w1 to w3.
Calculate the solubility of the sodium hydrogen carbonate as weight in grams dissolved in water at room temperature - (w3 to w1) / (w2 to w1 to w3).
Repeat the experiment using water at a higher temperature.
The solubility of sodium bicarbonate in 100 g of water varies with temperature:

Temperature	Dissolves	Temperature	Dissolves
0oC	to 6.9 g	30oC	to 11.1 g
10oC	to 8.15 g	40oC	to 12.7 g
20oC	to 9.6 g	50oC	to 14.5 g
25oC	to 10.4 g	60oC	to 16.4 g

3.6.9 Solubility in water of different salts
In this website, the word "solution" refers to substances dissolved in water, i.e. aqueous solutions.
A solvent is a liquid that dissolves another substance, the solute, to form a solution.
The three ways to increase the rate at which a solid dissolves in water are as follows:
1. grinding the solid until finely divided,
2. shaking the solution while the solid is dissolving, and
3. heating the solution.
Experiment
Try to dissolve 5 g of different salts each in 15 mL of water in a test-tube.
Attach a stopper and shake vigorously.
Solubility is a characteristic of a particular substance.
Classify each salt as soluble or slightly soluble or insoluble.
The solubility of a salt can be expressed as the number of grams able to dissolve in 100 g of water at 20oC

Chemical	Dissolves	Chemical	Dissolves
ammonium chloride	37.2 g	potassium chloride	34.0 g
barium chloride	35.7 g	potassium iodide	144.0 g
calcium chloride	42.7 g	sodium bicarbonate	9.6 g
copper (II) sulfate	20.7 g	sodium chloride	36.0 g
lead nitrate	54.4 g	sodium hydroxide	109.0 g
magnesium sulfate	25.2 g	sodium nitrate	87.5 g

Heat gently then strongly.
Use tongs to raise the lid.
The magnesium darkens before it melts.
When the magnesium starts to burn, put the lid back on the crucible and remove the burner.
Every few seconds raise the lid slightly to let more air enter.
Do not let white magnesium oxide smoke escape.
When the magnesium does not burn after you raise the lid, remove the lid and heat the crucible strongly.
Hold the lid ready in case the magnesium starts to burn again.
Let the crucible cool.
Again weigh the crucible + lid + contents = W2. Note W2 > W1.
The formation of magnesium oxide causes the increase in weight.
3. However, magnesium has density 1.74 g / cm3 and melting point 650oC, but magnesium oxide has density 3.58 g /cm3 and melting point 2 800oC, because the Mg2+-- O2- chemical bond is stronger than the Mg--Mg bond.

3.7.2 Decomposition of boric acid
Boric acid, H3BO3, loses water until it decomposes to the anhydride, B2O3.

3.7.3 Decomposition of carbonates
All the Group 2 carbonates undergo thermal decomposition to give the metal oxide and carbon dioxide gas.
Most carbonates decompose to form carbon dioxide and a metallic oxide.
Sodium carbonate and potassium carbonate do not decompose when heated to a high temperature.
CaCO3 (s) --> CaO (s) + CO2 (g)
white calcium carbonate --> white calcium oxide + carbon dioxide
Sodium hydrogen carbonate, NaHCO3, (sodium bicarbonate), begins to lose carbon dioxide at 50oC to form sodium carbonate.
A solution of a sodium hydrogen carbonate begins to lose carbon dioxide at 20oC.
MgCO3 (s) --> MgO (s) + CO2 (g)
white --> white
PbCO3 (s) --> PbO (s) + CO2 (g)
white --> yellow
ZnCO3 (s) --> ZnO (s) + CO2 (g)
white --> white (cold) or yellow (hot)
Ammonium carbonate may decomposes completely without heating when exposed to the air, to give ammonia, water and carbon dioxide.
(NH4)2CO3 (s) --> 2NH3 (g) + H2O (vapour) + CO2 (g)
colourless

3.7.4 Decomposition of chlorates
Potassium chlorate, KClO3, decomposes above 368oC into potassium perchlorate and oxygen gas.

3.7.5 Decomposition of chlorides
Sodium and magnesium chlorides are solids with high melting points.
The other chlorides are liquids or low melting point solids.

3.7.6 Decomposition of dichromates
Ammonium dichromate decomposes on heating
(NH4)2Cr2O7 + heat --> Cr2O3 + N2 + 4H2O

3.7.7 Decomposition of ferricyanides
Potassium ferricyanide, K2Fe(CN)6, decomposes before melting.

3.7.8 Decomposition of hydrogencarbonates, bicarbonates
Hydrogencarbonates decompose to form the metal carbonate, water and carbon dioxide.
Calcium bicarbonate and magnesium bicarbonate can exist only as a solution that on heating form the metal carbonate, water and carbon dioxide.
Sodium hydrogen carbonate, NaHCO3, (sodium bicarbonate), begins to lose carbon dioxide at 50oC to form sodium carbonate.
A solution of a sodium hydrogen carbonate begins to lose carbon dioxide at 20oC.
Heat sodium hydrogen carbonate crystals.
The crystals lose water and carbon dioxide, and at 100oC are converted to sodium carbonate.
2NaHCO3 (s) --> Na2CO3 (s) + CO2 (g) + H2O (vapour)
colourless --> colourless
Ca(HCO3)2 (aq) --> CaCO3 (s) + CO2 (g) + H2O (vapour)
Mg(HCO3)2 (aq) --> MgCO3 (s) + CO2 (g) + H2O (vapour)
2KHCO3 (s) --> K2CO3 (s) + CO2 (g) + H2O (vapour)

3.7.9 Decomposition of hydrates, hydrated salts
A hydrated salt is a crystalline salt molecule, attached to one or more water molecules.
A salt molecule, not bound to a water molecule may be called an anhydrate.
Hydrated salts have one or more water molecules within the crystalline structure of the salt molecule.
Heated hydrated salts may first lose their water of crystallization, and then become stable anhydrous powders.
In doing so, they lose their former crystalline shape and may lose their colour.
Experiments
Prepare test-tubes containing 1 cm of 1. iodine crystals 2. sodium hydrogen carbonate granules or crystals 3. silica sand 4. zinc oxide.
Fix a cotton wool plug in the mouth of each test-tube to prevent loss of solid during heating, then weigh each test-tube.
Heat each test-tube and cotton wool plug thoroughly and weigh it again.
Note any change in weight, because of the loss of water of crystallization.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous copper (II) sulfate + water.
CuSO4.5H2O (s) --> CuSO4 (s) + 5H2O (vapour)
blue --> grey white
Na2CO3.10H2O --> Na2CO3 (s) + 10H2O (vapour)
colourless --> white

3.7.10 Decomposition of hydroxides
Hydroxides of very active metals are stable when heated, e.g. sodium hydroxide, potassium hydroxide.
Hydroxides of less active metals decompose with strong heat to form water and the metallic oxide.
Mg(OH)2 (s) --> MgO (s) + H2O (g)

3.7.11 Decomposition of manganates
Potassium permanganate decomposes into potassium manganate, manganese dioxide and oxygen gas.
2KMnO4 --> K2MnO4 + MnO2 + O2

3.7.12 Decomposition of metals, metallic salts
Metal compounds higher in the activity series are usually more stable than compounds of metals lower in the activity series.
The salts that remain stable when heated constantly with a Bunsen burner flame are calcium sulfate, potassium chloride, potassium sulfate, sodium carbonate, sodium chloride, and sodium sulfate.
Ammonium oxalate (NH4COO)2, and ammonium dichromate, (NH4)2Cr2O7, decompose before melting.
Ammonium sulfate (NH4)2SO4, decomposes above 280oC.

3.7.13 Decomposition of nitrates
All the Group 2 nitrates undergo thermal decomposition to give the metal oxide, nitrogen dioxide and oxygen.
Nitrates decompose to form oxygen gas, nitrogen dioxide and a metallic oxide, except potassium nitrate and sodium nitrate that form the nitrite and oxygen.
Lead nitrate decomposes at 470oC.
2Ca(NO3)2 (s) --> 2CaO + 4 NO2 (g) + O2 (g)
colourless --> white
2Cu(NO3)2 (s) --> 2CuO + 4NO2 (g) + O2 (g)
blue --> black
2Pb(NO3)2 (s) --> 2PbO + 4NO2 (g) + O2 (g)
colourless --> yellow
Lead nitrate decomposes at 470oC.
2Zn(NO3)2 (s) --> 2ZnO + 4NO2 (g) + O2 (g)
colourless --> white (cold), yellow (hot)
Potassium nitrate and sodium nitrate first melt and then decompose to give the metal nitrite and oxygen gas.
Potassium nitrate melts at 336oC.
2KNO3 (s) --> 2KNO2 (s) + O2 (g)
colourless --> colourless
Sodium nitrate melts as 316oC.
2NaNO3 (s) --> 2NaNO2 (s) + O2 (g)
colourless --> yellow
Silver nitrate decomposes to give the metal, nitrogen dioxide and oxygen gas.
2AgNO3 (s) --> 2Ag (s) + 2NO2 (g) + O2 (g)
colourless --> silver
Ammonium nitrate decomposes to form water vapour and nitrous oxide, N2O (laughing gas), so the ammonium nitrate disappears.
NH4NO3 (s) --> N2O (g) + H2O (g)
colourless

3.7.14 Decomposition of oxalic acid
Oxalic acid begins to sublime at 100oC, becomes anhydrous at 189oC and when heated rapidly decomposes into carbon dioxide, carbon monoxide, formic acid and water.

3.7.15 Decomposition of oxides
Oxides of most metals are stable.
Oxides of potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead and copper do not decompose above 1, 800oC.
Black-grey silver oxide decomposes into the metal and oxygen gas.
2Ag2O (s) --> 4Ag (s) + O2 (g)
silver oxide --> silver + oxygen
Heat zinc oxide in a crucible.
Zinc oxide becomes yellow when hot and white when cold, but no change in weight occurs.
The substance does not decompose and does not gain anything from the air or lose anything to the air.
Zinc oxide has melting point above 1, 800oC.
ZnO (s) <--> ZnO (s)
white (cool) yellow (hot)
Thermal decomposition of higher oxides of lead
2PbO2 (s) --> 2PbO (s) + O2 (g)
brown lead dioxide --> yellow lead oxide + oxygen gas
2Pb3O4 (s) --> 6PbO (s) + O2 (g)
red trilead tetroxide --> yellow lead oxide + oxygen gas

3.7.16 Decomposition of phosphates
Monosodium orthophosphate, NaH2PO4.H2O, and disodium orthophosphate [disodium hydrogen phosphate (V)] Na2HPO4.12H2O, both lose water of crystallization.
10KClO3 <--> 6KClO4 + 4KCl + 3O2

3.7.17 Decomposition of sulfates
Sulfates if heated very strongly may decompose to form the metallic oxide, sulfur dioxide and oxygen gas.
Some sulfates decompose to form sulfur trioxide and metallic oxide.
Put 4 cm of crushed blue copper (II) sulfate crystals in a dry test-tube fitted with a one-hole stopper and delivery tube.
Heat the dry test-tube and crystals gently.
Note whether vapour collects on the cooler parts of the dry test-tube and whether any liquid collects in the receiving test-tube.
Note any change of colour of the crystals from blue to white.
Identity the liquid in the receiving test-tube by measuring the boiling point.
When all the copper (II) sulfate crystals have become white and the dry test-tube has cooled, pour the liquid in the receiving test-tube back on the white crystals.
Note whether the blue colour of the crystals is restored and if any heat is given off.
2CuSO4 (s) --> 2CuO (s) + 2SO2 (g) + O2 (g)
grey white --> black
2PbSO4 (s) --> 2PbO (s) + 2SO2 (g) + O2 (g)
white --> yellow
2ZnSO4 (s) --> 2ZnO (s) + 2SO2 (g) + O2 (g)
white --> white (cold) yellow (hot)

3.7.18 Decomposition of sulfites
Sulfites mostly decompose into the metal oxide and sulfur dioxide.
Sulfites of sodium and potassium do not decompose when heated.
CaSO3 (s) --> CaO (s) + SO2 (g)
white --> white
MgSO3 (s) --> MgO (s) + SO2 (g)
white --> white
ZnSO3 (s) --> ZnO (s) + SO2 (g)
white --> white (cold) yellow (hot)