School Science Lessons
2024-06-06
Colloids and crystalloids Emulsions, Silicon, Solubility, Suspensions and precipitates
(topic 07a)
Table of contents
7.8.0 Colloids and crystalloids
7.5.0 Emulsions
7.4.0 Silicon, Silica
7.7.0 Solubility and solutions
7.6.0 Suspensions and precipitates
7.8.0 Colloids and crystalloids
7.8.1 Colloids and crystalloids
7.8.2 Aerosols
7.8.3 Carrageenans
7.8.4 Gellan gum
7.8.5 Gels
7.8.1a Hydrocolloids
Hydrocolloids, Viscosity of non-Newtonian fluids:
7.8.17 Hydrophilic substances
7.8.6 Sols
Experiments
7.8.7 Prepare bean curd, (tofu, soya bean), (Experiment)
7.8.8 Prepare gelatine gel, (Experiment)
7.8.9 Prepare ferric hydroxide colloid, (Experiment)
7.8.10 Foams, (Experiment)
7.8.11 Prepare gels in the home kitchen, (Experiment)
7.8.12 Prepare metallic salts gels, (Experiment)
7.8.13 Prepare silver chloride precipitate used in photography, (Experiment)
7.8.14 Prepare sulfur in methylated spirits, (colloidal sulfur), (Experiment)
7.8.15 Prepare sunbeam mists, (Experiment)
7.8.16 Tyndall effect, sunbeams, (Experiments)
7.5.0 Emulsions
7.5.1 Emulsions
7.5.2 Emulsions with a microscope
7.5.3 Bile salts as an emulsifying agent
7.5.4 Prepare face cream emulsion
7.5.5 Lotions
7.5.6 Margarine
7.5.7 Prepare mayonnaise and salad dressing emulsions
7.5.8 Milk
7.5.9 Ouzo effect, microemulsions
7.5.10 Soap as an emulsifying agent
7.5.11 Temporary and permanent emulsions
7.4.0 Silicon, Silica
7.4.1 Silicon compounds, glass
7.4.2 Prepare silica gel, (Experiment)
7.4.4 Prepare silicon glass, (Experiment)
7.4.5 Prepare silly putty
7.4.6 Silicon tetrachloride with water
7.4.7 Silicone rubber
7.4.8 Soda-lime glass
7.4.9 Water glass
7.7.0 Solubility and solutions
7.7.1 Solubility, solubility rules, solubility equilibrium
7.7.1a Solubility equilibrium, solubility product, Ksp
7.7.2 Solutions, Rates of solution, (Experiment)
7.7.3 Evaporation of sea water
7.7.4 Separate soluble from insoluble substances
7.7.5 Solubility of copper (II) sulfate and particle size
7.7.6 Solubility of different salts
7.7.7 Solubility of gases and temperature, carbon dioxide
7.7.8 Solubility of potassium dichromate and temperature
7.7.9 Solubility of sodium chloride and iodine in different solvents
7.7.10 Solubility of sodium chloride and potassium dichromate in sugar solution
7.7.11 Solubility of sodium chloride in water at room temperature
7.7.12 Solubility of sucrose (cane sugar), syrup
7.7.13 Temperature affects solubility of sucrose
7.7.15 Weight of solids dissolved in tap water
7.7.16 Volume of gas dissolved in tap water.
7.6.0 Suspensions and precipitates
7.6.01 Suspensions and precipitates
7.6.1 Clay soils in water
7.6.2 Clay suspensions in a centrifuge
7.6.3 Aluminium sulfate with clay suspensions
7.6.4 Prepare aluminium hydroxide precipitate
7.6.5 Prepare magnesium carbonate precipitate
7.6.6 Salts with clay suspension
7.4.1 Silicon compounds, glass
Silicon is a metalloid, because it has physical properties of metals and chemical properties of non-metals.
Silicon is a semiconductor.
Silicon does not exist free in nature, but occurs mainly as silicon (IV) oxide (SiO2), in silica sand, sandstone, clay, quartz and opal.
Silicates occur in most rocks and glass.
Portland cement is a mixture of calcium and aluminium silicates.
In the silicone oils and greases, the silicon atoms form polymers containing a chain of silicon and oxygen atoms with carbon and hydrogen atoms attached to the chain.
Silicones repel water.
Silica glass, an amorphous solid, is more like a "super cooled liquid" than a crystal, because, unlike crystalline substances, it does not have a sharp melting point.
However, some chemists say that glass is a disordered solid not a supercooled liquid.
7.4.2 Prepare silica gel
Some gels can lose water by heating to leave a rigid gel.
Silica gel is an amorphous form of hydrated silica, very hygroscopic, and is used to protect delicate machinery from rusting.
Experiments
1. Prepare silica gel Add sodium silicate to water.
Add drops of phenolphthalein.
Add 3 M hydrochloric acid until the red colour disappears.
The solution sets as a gel.
Heat the gel to remove moisture.
The gel can absorb water again.
2. Prepare silicate gardens
Mix one part of sodium silicate (IV), (Na2SiO3), with four parts of water to make water glass.
Gently add crystals of salts to the solution without mixing to make chemical "flowers" grow:
* Chlorides: Co, Fe, Cu, Ni and Pb,
* Sulfates: Al, Fe, Cu, and Ni,
* Nitrates: Co, Fe, Cu and Ni.
Put sand 1 cm deep in a 500 mL jar.
Make a 1: 1 mixture of sodium silicate and water (water glass), and pour it onto the sand to almost fill the jar.
Leave the jar to stand undisturbed for a day.
Drop in crystals of metal salts, e.g. metal hydroxides, iron sulfate, copper (II) sulfate, alum, Epsom salts.
Observe crystals forming "shoots".
Some shoots are directed up by small bubbles.
The metal hydroxide skin formed around the crystal is permeable only to water and not the salt.
The water diffuses in to balance the concentrations each side of the skin until the skin bursts the skin then forms again further from the crystal.
7.4.4 Prepare silicon glass
1. Pick up sodium carbonate in a nichrome wire loop.
Dip the loop into powdered silica and heat over a burner to form a transparent bead of glass.
2. Prepare silicon glass in a furnace
Prepare glass in a crucible by heating a glass mixture in a furnace or over a Meker burner with a hot wide flame.
Glass mixture A: 17 g clean sand, 4.4 g sodium carbonate, 5.2 g disodium tetraborate (III)-10-water (borax).
Glass mixture B: 6 g clean sand, 2 g sodium carbonate, 1 g calcium carbonate.
3. Prepare coloured glass
Add a metal oxide to the glass mixture.
Heat the end of a glass rod to red heat.
Dip it into a powdered metallic oxide and heat until the oxide fuses into the glass.
4. Use salts to colour glass:
* For amethyst, use manganese (IV) oxide.
* For green, use black copper (II) oxide.
* For ruby, use red copper (I) oxide.
* For white, use tin oxide.
5. Mix a little silica with an equal quantity of calcium carbonate and about twice as much anhydrous sodium carbonate.
Grind them to a powder in a mortar with a pestle.
Make a loop in a platinum wire.
Heat the loop and place it in the mixture.
Reheat the wire with the adhering mixture until the mixture fuses, then cool it.
6. Describe the bead.
Hit it with a hammer to se if it is t brittle,
Does it dissolve in water?
7.4.5 Prepare silly putty
Dilatant compound, "Silly Putty", silicone, (toy product)
Silly putty, silicone, bouncing putty (Dow Corning 3179 dilatant compound), "Tricky Putty", (toy product)
Super ball, (toy product), is made from polybutadiene with small amounts of sulfur to reinforce the material and serve as a vulcanizing agent.
The ball is moulded under very high pressure and temperature and is said to have a 92% resiliency, about three times the resiliency of a tennis ball.
It can continue to bounce for about a minute after being dropped from a short height.
The silicone polymer in silly putty, polyborosiloxane, have covalent bonds within the molecules, but hydrogen bonds between the molecules.
The hydrogen bonds are easily broken.
A silicone is chains of (OH-Si-O-Si-O-Si-OH), with two methyl groups, CH 3 on each of the Si atoms.
However, in "Silly putty", boron atoms that can cross link weakly with oxygen atoms in other chains replace some of the silicon atoms.
Experiments
1. Mix 55% Elmer's glue solution and 16% sodium borate in a 4: 1 ratio.
Ingredients: 65% dimethyl siloxane, hydroxy terminated, polymers with boric acid, 17% silica, quartz crystalline, 9% thixotrol ST, 4% polydimethylsiloxane,
1% decamethyl cyclopentasiloxane, 1% glycerine, 1% titanium dioxide.
2. Apply small amounts of stress are slowly to the putty only a few bonds are broken and the putty "flows" and stretches a great distance.
Apply larger amounts of stress quickly, many hydrogen bonds are broken, and the putty breaks or tears.
Roll it into a ball that you can bounce.
Press it onto a pencil drawing so that it lifts off the pencil marks so you can see the drawing on the surface of the silly putty.
7.4.6 Silicon tetrachloride with water
Silicon tetrachloride, a chlorosilane, is an epoxy resin hardener.
Highly toxic by all routes, with irritating fumes.
Stoppers of glass reagent bottles containing it may stick and the bottles may explode, so use very small containers.
It reacts vigorously with sodium and other reactive metals.
Silicon tetrachloride reacts vigorously with water to form insoluble silicon dioxide and an acidic hydrogen chloride solution.
The solution gives off fumes of hydrogen chloride and dissociates to form hydrochloric acid solution.
SiCl 4 (s) + 2H2O (l) --> SiO2 (s) + 4HCl (aq)
HCl (aq) --> H + (aq) + Cl - (aq).
7.4.7 Silicone rubber
Silicone rubbers, e.g. PVMQ. a low-temperature resistant rubber cross-linked with vinyl groups, and used for gaskets, because it is temperature resistant.
Polysiloxanes, [repeat unit: -Si(RR ' )O-, where R = organic group, e.g. methyl], are also used for aquarium seals and building seals.
7.4.8 Soda-lime glass
Main types of glass
1. Soda-lime glass
Round glass jars may contain: SiO2, CaO, K2O, MgO, TiO2
Flat window glass may contain: SiO2, CaO, Al2O3, TiO2
2. Borosilicate glass, Pyrex
Addition of borate allows the formation of a glass that melts at a lower temperate than silica, and expands less on heating than soda glass.
Also, it is more plastic over a wider temperature range, e.g. Pyrex and glass wool.
So borosilicate glass has a very low coefficient of thermal expansion and a softening temperature above 800 oC.
The composition may be 70% silica, 10% boron oxide and some sodium, potassium and calcium oxides.
The chemical composition of the Pyrex used in laboratory glassware may be different from the Pyrex used in kitchenware.
Pyrex is sold as, Mixing bowls, Pyrex, glass 1 litre, Casserole dish, freezer and dishwasher safe, 2 litre.
3. Fused quartz glass may contain: SiO2, used in some camera lenses.
7.4.9 Water glass
1. Water glass is a colloidal solution of sodium or potassium silicate in water that solidifies on exposure to the air.
It is used in chemical gardens, egg preserving, paper sizing, fire resistant paint and fresco painting.
Sodium carbonate is heated with sand to produce sodium silicate, the water soluble water glass used is an inorganic builder in detergents, and for fireproofing.
Na2CO3 (s) + SiO2 (s) --> Na2SiO3 + CO (g)
Na2CO3 + SiO2 --> Na2O∙SiO2 + CO2
sodium carbonate + silicon dioxide --> sodium silicate (water glass) + carbon dioxide
Water glass is an aqueous solution of potassium silicate or sodium silicate, which solidifies in the air and seals the pores in the egg shell to preserve the eggs.
Commercial water glass is sold both as thick syrup liquid and as a powder. Experiment
Dissolve 1 part of water glass in 10 parts of boiled water.
Pour the solution over the eggs packed in a suitable container and store them in a cool place.
Do not wash the eggs before preserving them, because this removes the natural mucilaginous coating on the outside of the shell.
To prevent the shells of eggs preserved in water glass from cracking when boiled, puncture the blunt end of the egg with a pin before putting it into the water.
2. When sodium oxide Na2O 15%, silicon dioxide SiO2 70%, and calcium oxide CaO 10% are heated together up to 1000 oC, insoluble silica glass forms.
All the crystalline order of the added minerals has been lost.
In silica glass, (soda lime silica glass, crown glass), each silicon atom is surrounded by four oxygen atoms as a tetrahedron, each linked to other tetrahedra.
When ionic oxides are added in the glass making melt they get between the Si-O-Si bridges and weaken it.
For example: | transition temperatures silica glass Tg = about 1200 oC | Pyrex Tg = 550 oC | window glass Tg = 550 oC |.
3. The glass in high quality wine glasses (lead crystal), contains lead, which gives the glass a ringing sound, higher refractive index and more brilliance.
Cobalt gives blue glass, chromium gives green glass, and copper gives red or blue-green glass.
Boron oxide, B2O3, gives shockproof borosilicate glass, "Pyrex", that is resistant to all chemicals except hydrofluoric acid, HF.
Flint glass, lead glass, has no colour unlike crown glass that has a slight green to yellow colour due to iron impurity.
4. The 2 distinct constituents of glass are as follows:
4.1 The network former, i.e. the non-metal as an oxide is usually silicon, but it can be boron, aluminium, or phosphorus.
4.2 The network modifiers, e.g. sodium, potassium, calcium, and magnesium.
Glass may crystallize over a period of many years and then become more brittle, but some glass has remained uncrystallized for 4000 years in Egypt.
Sodium sesquicarbonate Na2CO3.NaHCO3.H2O occurs as a mineral.
4.3 A famous urban legend: glass in panes of old cathedrals and very old buildings is thicker at the bottom, because the glass has "flowed" down over along periods.
The most likely explanation is that at a time when the thickness of panes of glass varied at that stage of technology, glaziers always inserted the thicker sides down.
5. Collect drinking glasses, including wine glasses, of roughly the same size.
Strike each glass and listen to the ring to identify the existence of modifiers in the glass.
7.5.1 Emulsions
An emulsion is liquid dispersed in another liquid, e.g. mayonnaise, cream, milk.
An emulsion is the suspension of one liquid as fine droplets in another with which it does not mix.
Emulsions are colloidal systems with both the dispersed phase and the continuous phase are liquids, e.g. oil in water.
The droplets of a liquid remain suspended in another liquid.
Emulsions may be cloudy or opaque.
Emulsions are like suspensions, because they settle on standing.
A solid emulsion is a liquid dispersed in another solid, e.g. butter.
Emulsions: milk, mayonnaise, ice cream, kerosene in water, suntan lotion, hand cream, photography emulsions, butter (water in oil), cream (oil in water), Emulsions: margerine, milk, mayonnaise, ice cream, kerosene in water, suntan lotion, hand cream, butter (water in oil), cream (oil in water)
Emulsion: Continuous phase / outer phase, continuous "outside" liquid around second liquid having discontinuous droplets.
Emulsifying agents
Emulsifying agents are used to keep the phases dispersed so that the droplets remain suspended and the emulsion remains stabilized.
When you shake two immiscible liquids, droplets of one liquid are dispersed in the other.
In temporary emulsions, e.g. kerosene in water or oil in water, there is no attraction between the two liquids.
The two phases will disperse in each other if shaken together, but will separate on standing.
However, by adding an emulsifying agent, e.g. soap, the two liquids, kerosene and water, will remain dispersed within each other.
7.5.2 Emulsions with a microscope
Put a blob of an emulsion on the end of a microscope slide.
Put a drop of paraffin oil on one side of the blob and put a drop of water on the other side.
Stir a little of the emulsion into each liquid.
Smooth mixing occurs only when the liquid forms the continuous phase.
7.5.3 Bile salts as an emulsifying agent
The liver secretes bile that contains bile salts, sodium salts of the hydroxy steroid bile acids, e.g. taurocholic acid and glycocholic acid.
Bile salts are emulsifying agents and reducers of surface tension.
They allow fat particles to remain suspended long enough for enzymes to digest them before absorption into the bloodstream.
Bile salts are sold as sodium tauroglycocholate.
7.5.4 Prepare face cream emulsion
See diagram 16.2.2 : Emulsifiers used in cosmetics
Heat disodium tetraborate (III) in water until dissolved.
Heat a mixture of medicinal paraffin (propan-2-yl tetradecanoate), beeswax and petroleum jelly on a hot plate.
Pour the hot disodium tetraborate (III) solution into the mixture of hot oils and stir.
When cool, add perfume and colour.
7.5.5 Lotions
Lotions, cosmetic creams and ointments are mostly emulsions of oils dispersed in water.
Also, hair wave lotion and water glass are used to preserve eggs.
7.5.6 Margarine
Margarine is an emulsion of water, flavours, colours, and vitamins in a semi-solid fat, usually vegetable oils.
However, vegetable oils are liquids at room temperature, so hydrogenation is used to break double bonds, saturate the chains with hydrogen and increase the MP.
7.5.7 Prepare mayonnaise and salad dressing emulsions
Mayonnaise is a mixture of olive oil dispersed in vinegar and stabilized by egg yolk or mustard.
Beat 15 ml of edible oil with one egg yolk and water to produce a weak emulsion
Beat 120 ml of oil with one egg yolk and water to give a strong emulsion
Place egg yolks, lemon juice or vinegar and salt in a large glass bowl.
Use a balloon whisk to combine the mixture until it begins to thicken.
Pour drops of an edible oil into the egg yolk mixture and whisk until well combined.
Whisk the oil in gradually, because if it is added too quickly the oil may not emulsify and the mayonnaise could separate or curdle.
7.5.8Milk
Milk is a poor emulsion.
The forces of cohesion between the emulsified cream droplets are greater than the forces of adhesion between the milk and the cream, so cream floats above the milk.
In homogenized milk, the droplets have been broken into smaller particles and dispersed to form one phase.
7.5.9 Ouzo effect, microemulsions
When water is added to clear solutions of some alcoholic drinks, e.g. ouzo, pastis, sambuca, the drinks become a milky white microemulsio.
This milkiness is caused by the terpenes, e.g. Anethole, used for flavouring are soluble in alcohol, but not in water.
The added water creates a suspension of terpenes that have left the alcohol solution.
7.5.10 Soap as an emulsifying agent
Soap in water have molecules that have both lyophilic and lyophobic parts so is called an association colloid.
Soap molecules (sodium stearate), can cause droplets of fat to become negatively charged.
These droplets remain suspended in the water and can be become washed away during cleaning.
Soap is not soluble in saltwater.
Experiment
Add drops of vegetable oil or paraffin oil to: 1. water, 2. soap solution, 3. gelatine solution.
Leave the solutions to stand and note the separation.
The oil separates first in 1, then in 2, and then in 3.
If a stable emulsion forms in 3. the liquids may not separate.
Association colloids, e.g. soap in water have a dispersed phase part lyophobic and part lyophilic.
7.5.11 Temporary emulsions and permanent emulsions
1. Add 5 drops of kerosene to 5 mL of water and shake the test-tube.
Leave to stand and observe the separation into two layers.
Add liquid detergent, shake, and leave to stand.
Observe the permanent emulsion.
2. Add 5 drops of oil to 5 mL of water and shake the test-tube.
Leave to stand and observe the separation into two layers.
Add liquid detergent, shake, and leave to stand.
Observe the permanent emulsion.
3. Mix 5 mL of oil and 5 mL vinegar.
Leave to stand and observe the separation into two layers.
Add mustard or egg yolk, shake, and leave to stand.
Observe the salad oil permanent emulsion.
7.6.01 Suspensions and precipitates
Suspensions are heterogeneous systems of particles containing many molecules in a liquid, e.g. clay particles in water.
Observe the suspension particles.
Suspension particles eventually settle on standing.
The smaller the size of the suspension particles the longer the time to settle.
Under a microscope the suspension particles show Brownian movement.
A suspension has large solid solute particles that can be seen, settle out on standing and do not pass through filters or permeable membranes.
Examples suspensions include. milk of magnesia, Mg(OH)2, calamine lotion and muddy water.
The two or more components of a suspension are easily visible.
Precipitates are solids that form in a solution.
Precipitates form when the particles of dissolved substances join, and fall down, to leave the solution.
A precipitate may appear as a cloudy suspension in solution or as coagulated lumps.
It may be white or coloured.
7.6.1 Clay soils in water
Clay soils needs deep digging with the addition of gypsum and compost.
Fill to 3/4 a measuring cylinder with mixture of clay soil and water.
Shake the mixture then leave it to settle.
Note which particles settle and in what order.
For particles that do not float the smaller the particles the longer they take to settle.
Filter the milky coloured liquid.
The filter paper stops some particles, but the filtrate is still cloudy, because clay particles can pass through the filter paper.
Some particles remain in suspension for a long time, because they are small and water molecules hit them on all sides, caused by Brownian movement.
7.6.2 Clay suspensions in a centrifuge
1. Make a low-cost centrifuge from a meat grinder.
Put a test-tube of clay suspension in a sling and whirl it around the head. BE CAREFUL!
Another method is to attach the test-tube to the spoke of an upturned bicycle then turn the pedals with your hand.
During the whirling, the heavier particles move to the closed end of the tube and separate partly from the liquid.
2. Shake an emulsion in water then centrifuge it.
The liquids separate faster.
3. Centrifuge copper (II) sulfate solution.
No separation occurs.
7.6.3 Aluminium sulfate with clay suspensions
Half fill two measuring cylinders with a clay suspension or muddy water.
1. Add aluminium potassium sulfate [potassium alum, Al2(SO4)3.K2(SO4).24H2O] to a measuring cylinder.
2. Add aluminium sulfate [Al2(SO4)3.18H2O] and sodium carbonate to a measuring cylinder.
Shake both measuring cylinders and leave to stand.
Compare the clarity of the water in the two measuring cylinders.
Solid and liquid aluminium sulfate is used to treat and clarify wastewater, industrial effluent, and potable waters.
It is also used in the paper, food, dairy, oil, textile and chemical industries.
However, the alleged association of potable water containing aluminium sulfate with the incidence of Alzheimer's disease has discontinued its use in some cities.
7.6.4 Prepare aluminium hydroxide precipitate
1. Add 1 cm depth of potassium alum solution [Al2(SO4)3.K2(SO4).24H2O] to equal volume of dilute ammonia solution or dilute washing soda solution, [Na2CO3.10H2O].
A white jelly-like precipitate of aluminium hydroxide forms.
Add dilute citric acid or sulfuric acid and the precipitate disappears.
2. Put in a test-tube 1 cm depth of water and one drop of a common dye or ink, e.g. cochineal (pink cake colouring), Congo red, black ink, red ink.
Add an equal amount of dilute ammonia solution.
Fill the rest of the test-tube with potassium alum solution and leave to stand.
Aluminium hydroxide precipitate forms as a coloured jelly leaving the liquid above it colourless.
7.6.5 Prepare magnesium carbonate precipitate
Mix equal volumes of magnesium sulfate solution with sodium carbonate solution.
Filter the mixture so that magnesium carbonate remains on the filter paper and the sodium (Na + ) and sulfate (SO 4 2- ) spectator ions are washed through the filter paper.
Add deionized water to the filter paper to wash through any remaining spectator ions.
Leave the filter paper to dry then collect the magnesium carbonate.
MgSO4 (aq) + Na2CO3 (aq) --> MgCO3 (s) + Na2SO4 (aq).
7.6.6 Salts with clay suspension
In many hot countries, salt crystallizes from pans built on clay beds near the mouths of rivers.
Divide the clay filtrate into two test-tubes.
Keep test-tube as a control.
To the other test-tube, add drops of sodium chloride solution.
The filtrate becomes clear.
The effect occurs when a clay suspension in a river meets the salts in sea water.
7.7.1 Solubility, solubility rules, solubility equilibrium
1. Solubility
The solubility of a solute in a solvent is the number of moles or grams of the solute that can dissolve in a volume.
It is usually 100 g of the solvent (1 dm3 of water), at room temperature, usually 20 oC.
The solubility of a substance is the weight of solute that can dissolve in a solvent at a particular temperature.
For example, the solubility of sodium chloride is 36 g /100 g of water at 20 oC.
The solubility of gases decreases as the temperature rises.
When a more concentrated solution is diluted with a solvent,
C 1 V 1 = C 2 V 2, where C 1 = original concentration, V 1 = original volume, C 2 = final concentration, V 2 = final volume.
2. Solubility rules
Rules 1.
1. Most "group 1 elements" compounds, and all NH4+ compounds are soluble.
2. All nitrates, acetates, and chlorates are soluble.
3. All binary compounds of the Cl, Br and I with metals are soluble, except compounds of Ag, Hg(I), and Pb, but Pb compounds are soluble in hot water.
4. All sulfates are soluble, except sulfates of Ba, Sr, Ca Pb, Ag and Hg (I), but Pb, Ag and Hg sulfates are slightly soluble.
5. Carbonates, hydroxides, oxides, silicates, and phosphates are insoluble if not compounds of "group 1 elements" and NH4 + ions.
6. All sulfides are insoluble except sulfides of Ca, Ba, Sr, Mg, Na, K and NH4 +.
Rules 2.
All sodium, potassium and ammonium salts are soluble.
All nitrates are soluble.
All acetates are soluble.
All chlorides are soluble except silver chloride and lead chloride.
Lead chloride is slightly soluble in cold water and is more soluble in hot water.
All carbonates are insoluble except lead sulfate and barium sulfate.
Calcium sulfate is only slightly soluble.
All carbonates are insoluble sodium, potassium and ammonium carbonate.
Rules 3.
1. All ethanoates (acetates), are soluble, but the Ag + salt is slightly soluble.
2. All carbonates are insoluble, except the Na +, K + and NH4 + salt.
3. All chlorides are soluble, except the Ag + and Hg + salt.
The Pb 2+ salt is slightly soluble, but more soluble in hot water.
4. All hydroxides are insoluble, except the Na +, K + and NH 4 + salt.
The Mg 2+ and Ca 2+ salts are slightly soluble.
5. All nitrates are soluble.
6. All phosphates are insoluble, except the Na +, K +, NH 4 + salts and some acid phosphates.
7. All common sodium, potassium and ammonium salts are soluble.
8. All sulfides are insoluble, except the Na +, K +, NH 4 +, Mg 2+, Ca 2+ and Ba 2+ salts.
9. All sulfates are soluble, except the Ba 2+, Pb 2+, Ca 2+ and Hg 2+ salts.
The Ag 2+ salt is slightly soluble.
10. All salts of silver are insoluble, except silver nitrate and silver chlorate.
Silver ethanoate (silver acetate), and silver sulfate are slightly soluble.
7.7.1a Solubility equilibrium, solubility product, Ksp
When the forward process continues as the same rate as the reverse process in a reversible chemical reaction, the system is at equilibrium and its properties will not change, e.g. colour.
A dissolving solid and its solution reach equilibrium when the rate of crystallization equals the rate of dissolving.
Silver chloride has a low solubility in water.
AgCl (s) <=> Ag + (aq) + Cl - (aq)
Let [Ag+] = concentration of silver ions and [Cl-] = concentration of chloride ions.
The solubility at fixed temperature is constant.
Therefore, constant = [Ag+] × [Cl-], called the solubility product, K sp for a saturated solution of silver chloride at room temperature.
It is the equilibrium constant for that solid solution system.
Some solubility products, Ksp:
Table 7.7.2
Compound Ksp
1. AgCl2 2 × 10-10
2. CaSO4 3 × 10-4
3. BaSO4 2 × 10-9

If the solubility product is greater than K sp the a precipitate will form until the product of the ion concentrations equals K sp is large.
Then at equilibrium, the concentration of products is much greater than the concentration of the reactants.
7.7.2 Solutions, Rates of solution
A solution is a homogeneous system (no boundaries), that consists of a solute dissolved in a solvent.
The particles are usually molecules or ions.
An alloy, e.g. steel, is an example of solutions of solids dissolved in solids.
A saturated solution is a solution in which no more solute can dissolve at that temperature.
In a solution, the solid particles are very small, cannot be seen, and do not settle.
A solution is uniform in appearance, and clear or coloured.
The solute may be regained from the solvent by evaporation and can pass through fine filters and permeable membranes.
Solutions may contain more than one solute
The presence of one dissolved substance does not prevent other substances dissolving in the solution.
As general rule, unless the concentrations are high, one solute does not affect the solubility of others.
Rates of solution
Increase the rate at which some solid dissolves in water in three ways:
1. Grind the solid into smaller pieces.
Take two equal quantities of large crystals of copper (II) sulfate-5-water.
Grind one quantity into a fine powder.
Put both samples into equal quantities of water in separate test-tubes and shake.
Compare the rates at which the different samples dissolve.
2. Shake the solution while the solid is dissolving.
Put equal quantities of sugar into separate equal quantities of water in two test-tubes.
Shake one tube and leave the other to stand.
Compare the rates at which the samples dissolve.
3. Heat the solution.
Put equal quantities of potassium nitrate in equal quantities of water in two test-tubes.
Shake both test-tubes, holding one of them over a flame.
Compare the rates at which the samples dissolve.
7.7.3 Evaporation of sea water
1. When sea water in a rock pool evaporates, white crystals form along the edge of the pool.
If all the sea water evaporates, the final solid will contain about 0.5% carbonates, 3% gypsum, and 96.5% sodium chloride.
2. However, when the sea water starts to evaporate and crystals start to form, calcium is the only cation near saturation.
3. As water evaporates from the sea water, the order of precipitation is mainly:
* Calcium carbonate, calcite, and also calcium magnesium carbonate, dolomite,
* Calcium sulfate dihydrate, gypsum,
* Sodium chloride when the volume of the sea water is about 10% of the original,
* Potassium and magnesium salts.
4. If carbon dioxide is dissolved in the sea water, there is no extra precipitation of calcium carbonate.
The pH of sea water is 7.8 and carbon dioxide is present mainly as bicarbonate (HC03 -, about 90%).
Dissolving more carbon dioxide lowers the pH and the equilibrium shifts to converting more carbonate to bicarbonate.
5. On a ship, fresh water can be produced form seawater, by using heat from the engine to heat the water in a chamber which has a partial vacuum,
i.e. pressure lower than atmospheric pressure.
The evaporated sea water is left to cool to room temperature.
This process is expensive, because ot the cost of running evaporation pumps.
7.7.4 Separate soluble from insoluble substances
Shake a mixture of sand and sodium chloride in a test-tube containing water.
Filter the mixture into an evaporating basin.
Heat the filtrate to form sodium chloride crystals.
Sodium chloride, the solute, dissolves in water, the solvent, to form a solution.
Sand is insoluble in water.
Repeat the experiment with a mixture of ammonium chloride and sulfur.
The sulfur is insoluble in water.
The ammonium chloride is soluble in water and can be recovered by filtration and evaporation of the filtrate.
7.7.5 Solubility of copper (II) sulfate and particle size
Separate large crystals from small crystals of copper (II) sulfate.
Add 5 g of each to a test-tube containing the same amount of water.
Shake both test-tubes equally and simultaneously.
Note the amount of undissolved salt left in each tube after the same number of shakes.
Small particles dissolve faster than large particles.
7.7.6 Solubility of different salts
1. Test if a salt is soluble in water.
Select salts: Ammonium chloride, barium chloride, barium sulfate, calcium sulfate, copper nitrate, copper (II) carbonate, copper (II) sulfate, lead (II) nitrate,
potassium nitrate, potassium chloride, potassium sulfate, sodium chloride, sodium ethanoate (acetate), sodium sulfate, sodium carbonate.
Put 5 g of each salt in a test-tube.
Note the room temperature.
Add 10 mL of water and stir or shake vigorously.
Note whether the temperature of the mixture changes.
Classify each salt as soluble or slightly soluble or insoluble.
Check whether the results agree with the solubility rules.
2. Shake powdered blackboard chalk with water in a test-tube.
Filter the mixture and collect the filtrate.
Evaporate the water by heating the basin over a beaker of boiling water.
Observe the inside of the evaporating basin.
If you see any residue, part of the solid did dissolve.
3. Put 20 mL of tap water and deionized water or demineralized water in clean evaporating basins and evaporate each to dryness.
Observe the inside of each evaporating basin for any residue.
A residue indicates that the water contains dissolved solids.
4 Shake a small quantity (on a little finger nail), of each of the following salts with 10 mL deionized water or demineralized water:
Ammonium chloride, sodium acetate, sodium sulfate, sodium carbonate, barium chloride, barium nitrate, barium sulfate, copper nitrate, copper (II) sulfate,
copper carbonate, lead chloride, lead nitrate, lead sulfate, lead carbonate, calcium nitrate, calcium sulfate.
If the salt dissolves, note any change in the temperature of the mixture.
Classify each salt as soluble or slightly soluble or insoluble.
5. Mix the following pairs of substances in small quantities and observe whether a solution forms:
Sodium chloride and kerosene,
Olive oil and water,
Methylated spirit and water,
Petrol and olive oil,
Petrol and kerosene,
Methyl alcohol and copper (II) sulfate crystals,
Ethanol and copper (II) sulfate crystals,
Kerosene and petroleum jelly.
6. Half fill three identical beakers with water.
Add teaspoons measures of a different substance to each beaker until no more can dissolve, e.g. sugar, table salt, sodium bicarbonate.
Note how many teaspoons of the substance can be added to the water until the solution becomes saturated and no more of the substance can dissolve.
You can probably dissolve 20 spoonfuls of sugar in a cup of coffee before the coffee solution becomes saturated and undrinkable.
7. Solubility of different salts and temperature
Most ionic substances increase in solubility with increases of temperature.
After completing the experiment, keep the recrystallized salts in special jars and record each solubility on the label of the jar.
Record as solubility in 100 g of water, to nearest gram and o C.
Table 7.7.5
Chemical 20 o C 30 o C 40 o C 50 o C
Sodium nitrate 88 g 95 g 102 g 109 g
Potassium nitrate 31 g 46 g 62 g 82 g
Potassium chloride 34 g 36 g 39 g 42 g
Sodium chloride 35 g 36 g 37 g 38 g

Experiment
Plot solubility curves
See diagram 3.7.6 : Solubility and temperature, solubility curves
Measure the solubility of different salts at different temperatures, e.g. sodium nitrate, potassium nitrate and potassium chloride.
Use water at different temperatures.
Construct a table to compare the amount of each salt dissolved at different temperatures.
Draw a graph showing the results and the experimental results above.
Plot the solubility expressed as the number of grams of solute dissolved in 100 mL water along the vertical axis.
Plot the temperatures expressed in o C along the horizontal axis.
7.7.7 Solubility of gases and temperature, carbon dioxide
Use two identical plastic bottles or aluminium beverage cans of aerated water, fizzy drink, cola.
Leave the two drinks on the table for some time until you are sure that the contents are at the same room temperature.
Put one drink in the refrigerator, not in the r.
The next day take the drink out of the refrigerator and quickly open both drinks without any shaking.
Note which drink releases the most carbon dioxide gas as fizzy bubbles.
The warmer drink releases the most gas, because more carbon dioxide remains dissolved in the cooler solution.
7.7.8 Solubility of potassium dichromate and temperature
Dissolve potassium dichromate in 50 mL of water at 60 o C, until no more dissolves.
The solution is saturated.
Pour the clear solution into a second beaker.
Let the temperature drop slowly to 40 o C.
Crystals form in the second beaker.
Pour the clear solution from this beaker into a third beaker.
Leave it cooling to room temperature to form more crystals.
The experiment shows that a saturated solution contains fewer dissolved solids at a low temperature than at a higher temperature.
7.7.9 Solubility of sodium chloride and iodine in different solvents
Half fill two test-tubes with water and methylated spirit.
Add equal weight of sodium chloride.
The sodium chloride dissolves more in water than in the methylated spirit.
Repeat the experiment with iodine crystals.
The iodine crystals are more soluble in the methylated spirit than in the water.
7.7.10 Solubility of sodium chloride and potassium dichromate in sugar solution
Dissolve some sugar in a small quantity of water.
Add sodium chloride crystals to the solution.
Note whether it also dissolves.
Drop pieces of potassium dichromate into the solution and shake it.
The colour change of the solution shows that potassium dichromate is dissolving.
7.7.11 Solubility of sodium chloride in water at room temperature
Dissolve sodium chloride in 25 mL of water at room temperature.
Stir until no more dissolves.
The solution is saturated.
Filter the solution to remove undissolved salt.
Record the temperature of the saturated solution.
Weigh an evaporating dish (W1).
Pour the saturated solution into the dish and heat the dish slowly to evaporate the solution to dryness.
Cool the dish and weigh again (W2).
Mass of the salt dissolved = (W2 - W1).
Mass of water evaporated = mass of 25 mL.
The solubility of the salt = (W2 - W1) / 25 × 100 g per 100 mL water at room temperature.
7.7.12 Solubility of sucrose (cane sugar), syrup
Add sucrose to a test-tube of water at room temperature until no more dissolves after stirring.
Record the weight of the sucrose dissolved.
Sucrose, cane sugar, is very soluble in water.
Syrups are made of 850 parts of sugar and 150 parts of water.
Syrups generally act only as a medium in cooking, preparations and medicines.
7.7.13 Temperature affects solubility of sucrose
Repeat the experiment at 10 o C intervals until 70 o C and record the weight of sucrose dissolved.
Heat the solution to 100 o C.
Pour it into an evaporating basin and leave to cool.
Observe the cooling solution.
7.7.15 Weight of solids dissolved in tap water
Evaporate to dryness equal volumes of tap water and distilled water.
Observe the inside surface of each evaporating basin.
The residues in the basin show that tap water contains dissolved solids.
Weigh the residues.
7.7.16 Volume of gas dissolved in tap water
Hold a round bottom flask under water.
Insert a stopper with a delivery tube so that water completely fills the whole apparatus.
Note the original volume of water.
Put the end of the delivery tube into a test-tube full of water.
Heat the flask to boiling to collect dissolved gases in the test-tube.
Leave the apparatus to cool to room temperature.
Measure the volume of gases expelled from the water.
7.8.1 Colloids and crystalloids, Hydrocolloids
Colloids are a kind of mixture where one substance is dispersed throughout another substance.
So two phases exist, the dispersed internal phase and the continuous dispersion medium phase.
Colloids are glue-like amorphous substances, e.g. gelatine or starch, with particles bigger than most molecules,
i.e. 5 × 10 -9 to 5 × 10 -6 m (5 to 5000 nanometres, nm).
The particles are too large to pass through a membrane, but too small to be observed under a microscope.
Colloids have a dispersed phase of particles scattered through a continuous phase, the medium.
Colloids differ from crystalloids, e.g. inorganic salts that can pass through a membrane.
Colloids are a type of mixture with properties between that of true solutions and suspensions.
Colloids do not separate on standing, but can pass through filters.
Colloidal particles are uniformly distributed in solution, cannot be seen under the microscope, but are still too large to pass through membranes.
Colloids can scatter light and the common white colour of colloids is, because of the reflection by the particles of all colours of the spectrum.
The Scottish scientist Thomas Graham was the first to study colloids in 1861.
Table 7.8.1.0 Types of colloids
Medium / Phase
 Solid dispersed
phase		 Liquid dispersed
phase		 Gas dispersed
phase
Solid continuous medium Solid sols: alloys, coloured ruby glass, gemstones, paper, minerals,
gold in glass		 Liquid sols: paints, Fe2O3, clay,
chocolate drink,
pigmented ink		 Solid aerosols:
iodine vapour, dust, cement, ammonium chloride, smoke
Liquid continuous medium Gels: celluloid, gelatine, mud, pearl, silica gel, opal (Hydrocolloids: agar, gelatine, pectin, jelly, jam, glue) Emulsions: milk, mayonnaise, ice cream, kerosene in water, suntan lotion, hand cream, photography emulsions, butter (water in oil), cream (oil in water), Liquid aerosols:
fog, agricultural sprays,
hair sprays, clouds, visible steam
Gas
continuous medium		 Solid foams: rubber, pumice, Styrofoam, bread, cake, marshmallow, plaster, aerogel Foams: lather, froth, soap suds,
whipped cream,
shaving cream		 No colloids:
(All gases can mix together.)

CDS, complex disperse system, has mixture of solids, liquids and gases (air), e.g. ice cream.
Experiments
Note the properties of common colloids, e.g. mayonnaise, cod liver oil, glue, fog, smoke, aerosols, soups, human tissue, ice cream, fondants, marshmallows,
beaten egg white, face cream, milk, salad dressing (olive oil and vinegar), gravy, soap solution, salad cream, furniture cream, hand lotion, hair cream, ointment,
oil in water garden spray (white oil), creamy milk (unstable emulsion), cod liver oil, polyvinyl acetate paint.
Hydrocolloids
Hydrocolloids are gels where the liquid continuous medium is water so are gels or liquid sols.
Most food gels are hydrocolloids.
Hydrocolloids affect the texture (viscosity), of food.
Wounds may be covered with hydrocolloid dressings to improve healing.
7.8.17 Hydrophilic substances
Hydrophilic / hydrophobic substances
1. Hydrophilic substances are "water-loving" substances, polar molecule materials that mix with water, are attracted to water and may dissolve in water to form hydrogen bonds, e.g. glucose, sugars.
They have an affinity for water and are readily absorbed or wetted by water.
Hydrophilic colloids readily form hydrosols or remain as hydrosols.
2. Hydrophobic substances are "water-hating" substances, non-polar molecule materials, often oily, do not mix with water or they repel water, e.g. oils, proteins, greases, clays.
Hydrophobic colloids do not form or remain as hydrosols.
Soap molecules have one end polar and the other end non-polar, so they can attach to oils yet dissolve in water.
The aversion to water of a person suffering from the disease rabies is called "hydrophobia".
3. The hydrophobic force prevents the dispersion of oil in water, so to make oil and water emulsions, e.g. French salad dressing.
A salad dressing contain a vegetable oil, an acidifying agent, e.g. vinegar or lemon juice, egg yolk and starch.
Emulsifiers contain one region that is hydrophobic. (oil loving). and the other half is hydrophilic, (water loving).
The best emulsifying ingredients for salad dressings and vinaigrettes are egg yolks, mustard, mayonnaise, honey, and mashed avocado.
The components have to be shaken with the addition of stabilizing agents.
Salad dressings
A droplet of oil put in water forms a circle and does not mix with water.
Only continuous shaking of oil in water would keep them "mixed", but they will separate when the shaking stops.
Vinegars contains water and acetic acid.
Oil and water do not mix, so you would have to be continually shaking a mixture of oil and water to keep it at all “mixed”.
The only way of making oil and vinegar mix together for salad dressings and vinaigrettes, and not separate, is to make an emulsion by using an emulsifying agent.
One end of the molecules in emulsifying agents is hydrophobic, while the other end is hydrophilic.
The emulsifying agents pulling the molecules of oil and vinegar together, to form a uniform mixture.
Egg yolks contain lecithin and are one of the most commonly used emulsifiers for salad dressings.
Some cooks avoid using raw eggs from the fear of salmonella, so use pasteurized eggs.
(The use of raw eggs for restaurant cooking is illegal in some countries.)
Mix the emulsifying agent with vinegar first, with continuous whisking, then slowly add the oil.
Mayonnaise is a good emulsifier, it contains egg yolks.
Dijon mustard is the most popular type of mustard used for dressings and vinaigrettes.
The simplest vinaigrette is made with Dijon mustard, red wine vinegar, and olive oil.
One of the easiest vinaigrettes to make with honey is to mix it with balsamic vinegar, olive oil, and seasonings.
4. Non-polar molecules, do not have a strong separation of charge.
The positive and negative charges in a non-polar molecule are more or less distributed evenly.
Fats and oils are examples of non-polar molecules.
Non-polar molecules mix well with other non-polar molecules, but they do not mix with polar molecules.
So oil and grease mix well together, but neither mixes with water.
Non-polar molecules are often said to be hydrophobic, or "water fearing".
Lycopodium powder, made from the dried spores of club mosses, contain a lot of fat, which is hydrophobic.
When you sprinkle lycopodium powder on the surface of water, then stick your finger through the powder into the water, the powder coats your finger and prevents the water from touching your skin, so when you pull out our finger, it feels dry.
7.8.2 Aerosols
See diagram 13.8.2.0 : Aerosol can
Aerosols (fogs), are colloidal dispersions of particles of liquids or solids suspended in another gas, e.g. fly spray, mist, fog, clouds, smoke.
The term "aerosols" also refers to substances packed under pressure in a container so that the substances can be released as a fine spray.
The containers with contents are sometimes called "aerosols", e.g. a domestic fly spray.
An aerosol spray can is a pressurized container with an attached spray mechanism.
It is used to produce small amounts of product in a finely divided form, e.g. anti-perspirants, insecticides, medicines (nasal sprays), paint.
The pressure inside the aerosol can is much greater than atmosphere pressure.
Propellant
A propellant is dissolved in the product to push it out when the aerosol can is opened.
The propellants are usually alkanes, e.g. butane.
Chlorofluorocarbons, CFC gases, are no longer used as propellants, because they may increase the greenhouse effect.
Definition from the "Draft Australian criteria for the classification of hazardous chemicals": Aerosols
Aerosols (aerosol dispensers), means any non-refillable receptacles made of metal, glass or plastics,
containing a gas compressed, liquefied or dissolved under pressure, with or without a liquid, paste or powder,
and fitted with a release device allowing the contents to be ejected as solid or liquid particles in suspension in a gas, as a foam, paste or powder,
or in a liquid state or in a gaseous state.
7.8.3 Carrageenans
Sulfated polysaccharide hydrocolloid produced from red seaweed, Rhodophyceae, Chondrus crispis, Eucheuma spinosum, Ucheuma cottonii, Gigartina sp.
Vegetable gum, food additive E407, used for gelling and thickening, can form firm gels with potassium ions,
and can form elastic gels and thixotropic fluids with calcium ions
Cform viscous non-gelling solutions are free flowing jellies or firm gels.
Trade names: Genulacta, Gelgenuvisco.
7.8.4 Gellan gum
Water soluble tetrasaccharide hydrocolloid from two glucose residues, multi-functional gelling agent, in jellies, stabilizer in soya milk for soy protein suspension.
Fod additive E418 emulsifier, stabilizer, thickener, stable at up to 120 o C, alternative to agar, formed by fermentation of Sphingomonas elodea.
Trade names: Gelzan, Kelcogel, Phytadel, Gelrite.
7.8.5 Gels
1. Gels are colloids in which liquids are dispersed in solids to form a jelly.
They are semi-rigid systems, between the liquid state and the solid state, that consist of random networks of colloidal fibres or crystals, with liquid in the spaces.
Gels are colloidal systems that have set.
Some gels can lose one component by heating to leave a solid, e.g. silica gel.
Hydrogels have insoluble chains of polymers in a network that can contain very high concentrations of water.
They can be used for breast implants, water holding granules for dry soils, burn dressings, baby nappies (diapers), sanitary napkins, contact lenses and water sensors.
Environmentally sensitive hydrogels can be used to detect specific concentrations of substances or temperature or pH then release the substances they are carrying.
Hydrogels include acrylate polymers, polyvinyl alcohol, sodium polyacrylate and natural hydrogels, e.g. methyl cellulose and may have thixotropy,
i.e. become fluid when disturbed, but solidify at rest, e.g. hair gels.
2. Gels are colloids with a three dimensional structure of a network of linked molecules, e.g. gelatine, jelly, silica gel.
Gels easily change their shape under pressure, but may flow under high pressure.
Gels form when colloidal solutions are left standing to allow a 3 dimensional network to form.
Collagen is denatured by heat to form separate molecules called gelatine that form an amorphous network when cooled.
When gelatine is dissolved in water, it forms chemical bonds with the water and acts as a semi-solid or gel, jelly.
The gel can "dissolve" when heated and then form again when cooled again.
When the proteins in egg white are denatured by heating, they form a permanent gel that does not "dissolve" when heated.
7.8.6 Sols
1. Sols are colloidal solutions, suspensions of solid or liquid particles of colloidal dimensions in a liquid.
So sols are dispersions of small groups of molecules in a medium.
Sols remain dispersed, because Brownian movement prevents the groups of molecules precipitating under the influence of gravity.
2. Aerosols are more like solutions and are stable, e.g. starch in water, have large dispersed particles with an affinity with the medium.
Hydrophilic sols are the water loving sols.
Lyophilic sols are easily dispersed in certain mediums and may be redispersed after coagulation.
The particles of some protein solutions have a chemical shell of water molecules around them This prevents them flocculating unless a strong salt solution is added to precipitate the colloid.
4. Lyophobic sols (hydrophobic sols, solvent-hating sols), are difficult to disperse into an unstable solution and cannot be reformed after coagulation.
They can form precipitates, because they have a dispersed phase with no affinity with the medium, e.g. silver chloride precipitates in photography.
The particles keep apart, because of their electrical charge, but eventually they precipitate.
5. Xanthan gum, vegetable gum food additive, E415, comes from corn sugar (maize), fermentation, e.g. maize, by a strain of Xanthomonas campestris.
7.8.7 Prepare bean curd, (tofu, soya bean)
Weigh 50 g of moth free and mildew free soybeans and put them in a 500 mL beaker.
Add 300 mL water to soak for 24 hours to make the soybeans fully swell.
Replace the water if the atmospheric temperature is high.
Then pour out the water.
Grind the soaked soybeans in 200 mL water by using a household grinder.
To make soybean milk, transfer the soybean slurry to a filter fitted with two pieces of filter cloth, and filter by suction.
Wash the filter cake many times with 100 mL water to extract the soybean milk fully from the bean dregs.
The filtrate obtained is concentrated soybean milk.
Pour the concentrated soybean milk (or alternatively use industrial concentrated soybean milk in bags), into a clean 500 mL beaker and heat to about 80 o C.
Add saturated gypsum aqueous solution to the hot soybean milk.
Stir constantly until white wads appear.
Stop heating and let the soybean milk stand for five minutes.
Solid lumps start to separate out of the bean milk.
After standing for about 20 minutes, filter the solidified lumps from the bean milk.
Gather the lumps and shape them into a cube folded up in the filter cloth.
Place the cube on a clean plate and press it by putting a small beaker containing cold water on it.
About 30 minutes later, a cake of bean curd forms.
The bean curd will be whiter and more tender if concentrated soybean milk is used.
To preserve the freshly made bean curd from deterioration for a few days, soak it in 2-5% table salt aqueous solution and keep it in a cool, shady place.
7.8.8 Prepare gelatine gel
Water soluble protein from boiling collagen, e.g. horse hooves.
1. Weigh a teaspoon of gelatine and dissolve it in 100 mL of hot water.
Cool it to form a jelly.
2. Repeat the experiment to find the smallest concentration of gelatine needed to make a firm jelly.
7.8.9 Prepare ferric hydroxide colloid
Ferric hydroxide, iron (III) hydroxide, Fe(OH)3
Dilute a ferric chloride solution until it appears pale yellow in colour.
Pour some of this solution into a test-tube then place the test-tube in a beaker of warm water until the solution turns brown.
Do not heat the solution to boiling.
The colloid formed is hydrated iron oxide, Fe2O3.
Keep the colloid for later use.
Fe 3+ (aq) + 2H2O (l) --> (colloidal) + 3H + (aq).
7.8.10 Foams
Foams are solutions of gases in solids or liquids.
Foams, e.g. foam rubber, expanded polystyrene, have been stabilized with surfactants (surface acting agents), e.g. detergents.
A foam is gas dispersed in a liquid, e.g. foam on beer, whipped cream, fire extinguisher foam.
A solid foam is gas dispersed in a solid, e.g. marshmallow, foam rubber mattress, expanded plastics (foamed plastics), e.g. Styrofoam.
A foam is a sort of cellular solid with a gas matrix.
Solid foams are often highly flammable.
Experiments
1. Put bubble bath solution into a container of water and beat strongly until some foam forms.
Here a gas is dispersed in a liquid.
2. Beat or whisk some cream (emulsion), until it becomes whipped cream foam.
3. Dissolve a packet of household jelly crystals in one cup of hot water.
Cool and when the jelly is partially set, beat until frothy.
Whip the mixture with one cup of chilled evaporated milk until the mixture is stiff.
Add 2 tablespoons of lemon juice.
Whip the mixture again.
Fold this into the jelly gradually.
Pour into pie crust and refrigerate for 3 hours.
You can eat this foam colloid!
7.8.11 Prepare gels in the home kitchen
1. In the home kitchen, dissolve a sachet of household gelatine in 50 mL hot water.
Pour into a dish and leave to cool until a firm jelly forms.
2. In the kitchen, repeat using agar or jelly crystals instead of gelatine.
7.8.12 Prepare metallic salts gels
1. Dissolve 19 g calcium chloride in 25 mL water.
Dissolve 28 g potassium carbonate in 25 mL water, then pour into the calcium chloride solution and stir vigorously.
A gel forms of CaCO 3.nH 2 O.
2. Make a gel by adding a solution of magnesium sulfate (26 g in 100 mL water),
to potassium hydroxide solution (107 g in 100 mL water), or sodium hydroxide solution (42 g in 100 mL water).
Add 30 mL calcium acetate solution (35 g /L).
The gel formed, called "Sterno", is flammable and is used for outdoor stoves.
7.8.13 Prepare silver chloride precipitate used in photography
Mix in a test-tube 4 cc of 2% silver nitrate solution with 4 cc of 5% sodium chloride solution.
A white precipitate of silver chloride forms.
Wrap a photographic negative around the tube and expose to sunlight for an hour.
Remove the negative.
See the image from the negative on the walls of the test-tube, because the light causes reduction of the silver in the silver chloride to black silver.
NaCl (aq) + AgNO3 (aq) --> AgCl (s) + NaNO3 (aq)
7.8.14 Prepare sulfur in methylated spirits, (colloidal sulfur)
Mix sulfur powder (flowers of sulfur), with methylated spirits.
Shake the mixture then filter.
Continuously turn the mixture and pour the solution into a beaker of water.
A weakly opalescent colloidal sulfur solution forms.
7.8.15 Prepare sunbeam mists
1. Pour concentrated hydrochloric acid and concentrated ammonia solution into two watch glasses.
Be careful!
Move the watch glasses close to each other and observe the smoke of ammonium chloride that forms.
2. Put a burning splint of wood on a tile and shine a beam at the smoke, then observe what happens.
3. In a darkened laboratory, shine a projector beam at the mist as it emerges from the nozzle of an aerosol spray can.
4. Use a projector beam or an electric torch to see colloidal dust, including chalk dust and "sunbeams", (particles in the air).
7.8.16 Tyndall effect, sunbeams
The Tyndall effect (John Tyndall, 1820-1893, UK), is caused by reflection of light by very small particles in suspension.
It can be seen from the dust in the air when sunlight comes in through a window to form "sunbeams",
or when light comes down through holes in clouds or the visible headlight beams during foggy nights or on dusty roads.
A laser pointer can show the Tyndall effect in diluted milk or colloidal silver as the light beam passing through the liquid.
The Tyndall effect increases with concentration, but it is not used to measure the concentration of a sol, but only whether a colloid is present or not.
Shine light in one side of a box with a scattering solution, e.g. colloidal silver, smoke, and see the scattered light out in a perpendicular direction.
The Tyndall effect refers to the translucence of some colloids caused by scattering of light by colloid particles.
When a beam of light passes through some suspensions and colloids containing particles with diameters < 1/20 the wavelength of light,
the scattered light appears mainly blue, e.g. tobacco smoke suspension.
Direct a beam of light through the solution.
If you can see the solute particles, the solution is colloidal, because of the scattering of light.
Experiments
Fill 250 mL beaker with yellow potassium chromate solution (K 2 CrO 4 ).
This is a true solution so it will not scatter light.
Prepare the following test solutions in 250 mL beakers:
1. Copper (II) sulfate solution,
2. Starch solution,
3. 2 mL olive oil,
4. 10 mL water, 4 drops detergent, shake and put in a large test-tube,
5. Weak black tea,
6. Instant coffee solution,
7. Detergent in a beaker of water,
8. Toothpaste shaken in a beaker of water,
Put each beaker of test solution next the beaker of potassium chromate solution and pass light from a projector through both solutions.
Observe light passing through both solutions to detect which solutions are colloids.