School Science Lessons
Please send comments to:

Table of contents
Boiling point, bp
7.8.0 Colloids and crystalloids
7.3.0 Metals, non-metals, transition metals
7.4.0 Melting point, m.p.
3.0.0 Particles
7.8.1 Prepare colloids and crystalloids Prepare silicon compounds, glass
7.2.3 Silicon compounds
3.9.0 Solubility and solutions
7.6.0 Suspensions and precipitates

7.5.0 Boiling point, bp Boil an egg
7.9.12 Boiling chips, anti-bumping chips
2.0 Boiling point of elements, (See: Elements) Boiling point of Helium
3.6 Boiling point of inflammable liquids, ethanol, acetone
7.5.0 Boiling point of liquids
7.5.3 Boiling point of mixture of two liquids, water and alcohol
3.5.1 Boiling point of sodium chloride solution
3.5.0 Boiling point of water
7.9.7 Constant boiling mixture, azeotrope
7.5.1 Elevation of boiling points, ebullioscopic constant, kB
24.1.04 Freezing point depression and boiling point elevation
7.5.2 Leidenfrost effect
22.4.0 Melting point and boiling point
24.2.0 Phase changes, liquid / gas, boiling point
24.2.4 Pressure and boiling point of water
3.8 Pressure of the atmosphere affects the boiling point
13.7.13 Simulated boiling
3.7 Volatility of different liquids

7.8.0 Colloids and crystalloids
"Aquadag", (colloidal graphite solution), carbon, C:
Aerosol, sunbeam mist:
Bile salts as an emulsifying agent:
Colloids and crystalloids, sols, emulsions, gels, aerosols, foams:
Colloids and crystalloids, Prepare mayonnaise:
Colloidal graphite, graphite, (graphite solution), Aquadag
Colloidal nature of egg white: 9.172
Colloidal solution of starch: 9.180
Colloids, diffusion, semipermeable membrane, osmosis, plasmolysis: 9.12.0
Colloids in food:
Emulsifying agents (for pesticides):
Emulsions with a microscope:
Ferric hydroxide, iron (III) hydroxide, colloid:
Foaming agents, synthetic detergents, (syndets):
Foams, polymer foam: 3.5.1
Gellan gum:
Gels in the home kitchen:
Metallic salts gels, calcium carbonate gel, calcium acetate gel:
Separate a colloid from a crystalloid by dialysis: 9.181
Silver chloride precipitate in photography:
Size of colloidal particles:
Soap as an emulsifying agent:
Sodium thiosulfate with dilute hydrochloric acid, concentration and rate of reaction: 3.92
Sols, lyophilic sols, lyophobic sols, xanthan gum:
Sulfur in methylated spirit colloid:
Temporary emulsions and permanent emulsions, kerosene, detergent:
Tyndall effect:
Vegetable gums, food additives:
Viscosity of non-Newtonian fluids:
7.8.1 Prepare colloids and crystalloids
Prepare bean curd, (tofu, soya bean):
Prepare face cream emulsion:
Prepare gelatine gel:
Prepare silica gel:

7.3.0 Metals, non-metals, transition metals
2.6 Free element metals
Metals, non-metals, transition metals
7.3.1 Properties of metals
7.3.2 Properties of non-metals
2.43 Metals (Primary) Tests for metals with flame tests, metals and their compounds Tests for metals with borax beads, metals in metallic salts and minerals
1.12 Transition elements, Transition metals

7.4.0 Melting point, m.p. Ice melts, de-icers
7.4.0 Melting point, m. p., of solids
3.2 Melting point of naphthalene
3.3 Melting point of naphthalene with a capillary tube
3.4 Impurities affect the melting point of a substance
24.1.2 Lift an ice cube with salt
2.9 Melt different substances (Primary)
7.4.1 Melting point and cooling curve of stearic acid Melting point of substances, candle wax, cetyl alcohol Melting point of 1,4-dichlorobenzene, C6H4Cl22, men's toilet deodorant Melting point of ice and freezing point, (fp), of water, antifreeze Temperature at which ice melts Temperature at which ice and salt mixture freezes

3.0.0 Particles
3.55 Brownian movement
3.58 Clay soil suspension
3.56 Particles of matter and dilution
3.57 Size of a molecule

7.2.3 Silicon compounds, glass
Silicon compounds, glass
35.14.0 Quartz Prepare silicon compounds, glass Silicon tetrachloride with water
35.14.1 Silicates group, polysilicates, polysilicon
7.2.6 Silly putty, silicone, bouncing putty (Dow Corning 3179 dilatant compound) "Tricky Putty"
7.2.7 Silicone rubbers, PVMQ Prepare silicon compounds, glass
7.2.4 Prepare silicon glass Prepare silicon glass in a furnace Prepare coloured glass
7.2.5 Prepare silicate gardens

3.9.0 Solubility and solutions
6. Equipment and solutions
7.7.14 Fractional crystallization of sea water Group tests to identify cations, prepare a solution for group analysis
3.17 Heat of solution Invisible ink, Prepare ink, disappearing, secret writing, sympathetic ink
3.72 Magnesium displaces copper from solution of copper ions
3.17.1 "Magnetic" sugar cube dissolves
3.16 Miscible liquids, methylated spirit, glycerine, kerosene
3.1.0 Prepare crystals, (List)
8.0 Prepare salt solutions
5.1.1 Prepare molar solutions
5.3.0 Prepare stock solutions, standard solutions
5.4.0 Prepare solutions of known concentration Prepare toffee candy
10.3.1 Separate sand and salt mixture
7.7.4 Separate soluble from insoluble substances
10.12.1 Separate by solvent extraction, extract oil from peanuts, (groundnuts)
5.2.0 Serial dilutions, Prepare serial dilutions, different percentage concentrations
10.4.1 Shake different liquids in water
3.71.2 Sodium chloride solution with copper (II) sulfate solution
19.1.1 Solid acids, solubility
7.6.7 Solubility
3.14 Solubility and agitation, cane sugar
3.12 Solubility and particle size, sodium chloride, copper (II) sulfate
3.13 Solubility and solvents, sodium chloride, methylated spirit
7.7.6 Solubility and temperature, plot solubility curves
3.10 Solubility and temperature, solubility of salts in water
3.71.5 Solubility, Blackboard chalk, (school chalk) (See 1.)
7.7.2 Solubility equilibrium, solubility product, Ksp
3.11 Solubility in water at a given temperature
3.9 Solubility in water of different salts
3.34.3 Solubility of acidic oxide, carbon dioxide, in water, acidity of soda water, fizzy drinks
12.7.2 Solubility of alkalis
7.7.8 Solubility of copper (II) sulfate and particle size
7.6.8 Solubility of different salts
7.7.5 Solubility of different salts and temperature
7.7.1 Solubility of gases and temperature, carbon dioxide Solubility of potassium dichromate and temperature Solubility of sodium chloride and potassium dichromate in sugar solution Solubility of sodium chloride and iodine in different solvents Solubility of sodium chloride in water at room temperature
7.7.9 Solubility of sucrose (cane sugar), syrup
7.7.3 Solubility rules
5.04 Solutions for herbal remedies, infusion, decoction, tincture, poultice, ointment
5.0 Sodium hydrogen phosphate, disodium hydrogen phosphate, standard buffer solutions:
7.6.6 Solution Solution, Aqueous solutions, states of matter
5.0.0 Solutions and mixtures:
7.7.10 Solutions that contain more than one solute Solvent for Sudan III Temperature affects solubility of sucrose
3.71.4 Test if precipitate forms when solutions are added to lead (II) nitrate
3.71.3 Test if precipitate forms when solutions of salts are mixed
3.15 Volume of solutions
7.7.13 Weight of solids dissolved in tap water Volume of gas dissolved in tap water

7.6.0 Suspensions and precipitates
7.6.0 Suspensions and precipitates
7.6.1 Clay soils in water
7.6.3 Clay suspensions in a centrifuge Limewater with clay suspension Potassium alum or aluminium sulfate with clay suspensions
7.6.4 Prepare aluminium hydroxide precipitate
7.6.5 Prepare magnesium carbonate precipitate
7.6.2 Salts with clay suspension
3.58 Tests for clay soil suspension

7.2.3 Silicon compounds, glass
See 9.30: Egg preservation (See 6.)
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 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.
1. Sodium carbonate is heated with sand to produce sodium silicate, the water soluble water glass used s an inorganic builder in detergents, for preserving eggs and for fireproofing materials.
Na2CO3 (s) + SiO2 (s) --> Na2SiO3 + CO2 (g)
sodium carbonate + silicon dioxide --> sodium silicate (water glass) + carbon dioxide
2. When sodium oxide Na2O 15%, silicon dioxide SiO2 70%, and calcium oxide CaO 10% and other oxides are heated together with temperatures up to 1000oC, insoluble silica glass forms in which 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 and each of these is 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. as shown in the transition glass temperatures: silica glass Tg = about 1200oC, Pyrex Tg = 550oC, window glass Tg = 550oC.
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 has the opinion that glass in panes of  old cathedrals and even early American 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.

7.2.4 Prepare silicon glass
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. 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 Prepare coloured glass
1. Add a metal oxide to the glass mixture.
2. 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. Use salts to colour glass: 1. amethyst use manganese (IV) oxide, 2. green use black copper (II) oxide, 3. ruby use red copper (I) oxide, 4. white use tin oxide.
3. 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.
Cool it. What does the bead look like? Hit it with a hammer. Is it brittle? Does it dissolve in water?

7.2.5 Prepare silicate gardens
1. 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:
1.1 Chlorides: Co, Fe, Cu, Ni and Pb
1.2 Sulfates: Al, Fe, Cu, and Ni
1.3 Nitrates: Co, Fe, Cu and Ni.
2. 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 shoos 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. 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 that gives off fumes of hydrogen chloride and dissociates to form hydrochloric acid solution
SiCl4 (s) + 2H2O (l) --> SiO2 (s) + 4HCl (aq)
HCl (aq) --> H+ (aq) + Cl- (aq)

7.2.6 Silly putty, silicone, bouncing putty (Dow Corning 3179 dilatant compound) "Tricky Putty"
See 3.4.04: Super ball | See 3.4.11: Prepare slime balls, "Silly putty", silicone polymer to amuse children
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, CH3 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.
1. 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.
2. 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.

7.2.7 Silicone rubber
Silicone rubbers, e.g. PVMQ a low temperature resistant rubber cross-linked with vinyl groups and used for gaskets because temperature resistant, polysiloxanes, [repeat unit: -Si(RR')O-, where R = organic group, e.g. methyl], also used for aquarium seals and building seals.

7.3.0 Metals, non-metals, transition metals
1. Metals are elements that form positive ions. Metals may also form complex ions with a non-metal, e.g. chromate ion, CrO4 2-.
2. Non-metals are elements that form negative ions. Non-metals may also combine with other non-metals, e.g. NO3-. The first 105 elements in the periodic table contain 84 metals and 21 non-metals (or metalloids). Metalloids have both metal and non-metal e.g. B, Si, As, Te and Ge, Sb, Bi.
3. Transition elements include:
First series: Cr, Mn, Fe, Co, Ni, Cu (and perhaps Zn),
Second series: Mo and Ag, and
Third series: Pt, Au (and perhaps Hg).
Transition elements are hard, dense metals with high melting points and boiling points, form coloured ions and compounds and have more than one valence (oxidation number), e.g. copper (Cu1+ and Cu2+) and chromium (Cr2+ and Cr3+). The ion Cr2+ is a strong reducing agent and forms blue salts in solution. Cr3+ salts are green in solution. CrO42- salts are yellow in solution. Cr2O72- is a strong oxidizing agent with orange salts in solution, e.g. K2Cr2O7.
3. Transition metals form complex ions and often have catalytic activity, e.g. Fe in the complex haemoglobin molecule, and Fe in the manufacture of ammonia Observe common metals and non-metals, e.g. aluminium foil, calcium granules, carbon as charcoal, copper foil or wire, iodine crystals, lead foil or lead shot, magnesium ribbon, nickel sheets or plating, red phosphorus, potassium (in liquid paraffin) sodium (in liquid paraffin) sulfur powder or "flowers", tin foil or tin-plating, zinc foil or granulated zinc.
List the metallic or non-metallic properties: you can observe.

7.3.1 Properties of metals
1. Are usually opaque, hard solids with high density and have shiny, silvery lustre, when cut or scratched.
2. Have strong metallic bonds that make them malleable, ductile and easy to bend. They have high melting points, boiling points and density. You can hammer metals into new shapes because they are malleable. Metals can be drawn into a wire because they are ductile.
3. Have freely moving electrons and thus are good conductors of electricity and heat.
4. May form oxides that turn moist litmus blue. For example, calcium oxide dissolves in water to form the alkali calcium hydroxide. Some metals form a surface layer of oxide that prevents more oxidation, e.g. Al, Cr, Mg and Zn.
5. Reactive metals have stable compounds. Stable metals have compounds that easily decompose.
7.3.2 Properties of non-metals
1. May be monatomic, e.g. Ne, He, polyatomic, e.g. F2, white phosphorus P, or network solids, e.g. diamond, red phosphorus.
2. May be solids, liquids or gases.
3. Non-metals that are solids, they are usually dull and brittle, and have low density.
4. Are usually poor conductors of electricity and heat.
5. Usually form oxides that turn moist litmus red. For example, carbon dioxide dissolves in water to form the weak acid, carbonic acid (H2CO3).
H2O (l) <--> H+ (aq) + OH- (aq)
2H+ (aq) + CO32- (aq) <--> H2CO3 (aq) carbonic acid
CO2 + H2O <--> H3O+ + HCO3-
HCO3- + H2O <--> H3O+ + CO32-
6. Exceptions to the properties: of non-metals include graphite carbon that conducts electricity, silicon that has the physical properties: of a metal.

7.4.0 Melting point of solids
See 3.2: Melting point of naphthalene
The melting point (m.p., 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- 1600 °C and it depends on the composition of the glass.

7.4.1 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. Melting point of substances, cetyl alcohol
Repeat the above experiments with substances, e.g. candle wax, m.p. = 45oC to 65oC, urea, m.p. = 133oC and hexadecan-1-ol (cetyl alcohol) [CH3(CH2)14CH2OH].
Cetyl alcohol, palmityl alcohol, hexadecanol, 1-hexadecanol, CH3(CH2)14CH2OH, hexadecan-1-ol,  Toxic by all routes, highly flammable. Cetyl alcohol forms the ester cetyl palmitate (C15H31.COO.C16H33) m.p. 42oC to 47oC, that is the main component of spermaceti, the white wax found in the head of the sperm whale. The wax was used to make cosmetics and ointments but nowadays we use other chemicals as substitutes to "Save the Whales"! Melting point of 1,4-dichlorobenzene, C6H4Cl2, men's toilet deodorant
Substitute hexadecan-1-ol or octadecan-1-ol for melting point curve experiments. Melting point of ice and freezing point of water
E33 Ice Melting Blocks, endothermic heat flow, "Prof Bunsen", (commercial website)
1. For pure substances, the melting point, (mp) = freezing point, (fp). 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 -20.6oC, also methanol. 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. 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. Temperature at which an ice and salt mixture freezes
See 19.1.16: Table salt and rock salt
At 0°C, 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 -1.8°C. Mix crushed ice with salt. Note the temperature after five minutes. The temperature of the ice and salt mixture is below zero, e.g. -20°C, if 1:3 ratio of sodium chloride to ice. Ice melts, de-icers
1. Calcium chloride, melts ice to -25˚F, and off heat as it dissolves to melt the ice quicker,  but leaves a slimy residue,  corrosive to metal, damages vegetation.
2. Magnesium chloride is less corrosive and safer on concrete and plants.
3. Sodium chloride, rock salt, is least expensive and very efficient, melts ice to 20˚F, dries icy surfaces, not as harmful to concrete as other products but damages vegetation and corrosive to metal.
4. Potassium chloride,  more expensive than other ice melt products but works well as 50/50 with rock salt, melts ice to 12˚F, can cause plant injury.
5. Urea,  melts ice to 15˚F, but may harm vegetation.
6. Calcium magnesium acetate (CMA), from dolomite  limestone and acetic acid, not effective below below 20˚F, 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.

7.5.0 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, so the boiling point should be high and the vapour pressure should be low. 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.
7.5.1 Elevation of boiling points, ebullioscopic 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 (ebullioscopic 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. Boil an egg
Put two same size eggs in the same volume of water in same size saucepans. Add 1 teaspoon of salt to one saucepan. While heating  the saucepans for four minutes for semi-firm yolks and hard whites, insert thermometers and compare the temperatures of the water. If the cooking eggs are stirred clockwise the yolks are supposed to remain in the centre. The differences in temperature are negligible. After 4 minutes  take out the eggs, open them and compare the taste and firmness of the contents. However a cook advises: The best way to make easy-to-peel hard-boiled eggs with pure yellow centres is to place the eggs in a pot of cold water and add a teaspoon of salt. Place the pot on the stove and bring to a boil. Cover the pot and remove from heat and let sit for 13 minutes. Drain the eggs and immediately place in an ice bath to avoid green yolks until completely cooled. The green colour of yolks is caused by iron in the yolk combining with the sulfur in the white of the egg. The longer the egg is cooked the more likely a green ring is seen in the yolk from the formation of green-grey ferrous sulfide and hydrogen sulfide gas. To test whether an egg is cooked spin the egg on its side on the table. An uncooked egg wobbles when spun. Some cooks add salt to the water to denature the protein and solidify any of the white that may leak from a crack in the egg shell so that the egg does not shoot out a streamer of white. Very fresh eggs are more acidic and more difficult to peel.

7.5.2 Leidenfrost effect
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 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 experimenter 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!

7.5.3 Boiling point of a mixture of two liquids, water and alcohol
Use the above method to compare the boiling points of mixtures of water and methylated spirit in proportions:
BE CAREFUL! Alcohol is Highly flammable!
Table 7.5.3
Liquids Solution 1 Solution 2 Solution 3
1. Water 25% 50% 75%
2. methylated spirit 75% 50% 25%

7.6.0 Suspensions and precipitates
1. 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, e.g. milk of magnesia, Mg(OH)2, calamine lotion and muddy water. The two or more components of a suspension are easily visible.
2. 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.
3. When water is added to clear solutions of some alcoholic drinks, e.g. ouzo, pastis, sambuca, the drinks become milky white because the terpenes 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.

7.6.1 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 because of Brownian movement.

7.6.2 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 one as a control. To the other 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. Limewater with clay suspension
Repeat the experiment by adding drops of limewater. Potassium alum or aluminium sulfate with clay suspensions
Half fill two measuring cylinders with a clay suspension or muddy water. Add aluminium potassium sulfate [potassium alum, Al2(SO4)3.K2(SO4).24H2O] or add aluminium sulfate [Al2(SO4)3.18H2O, water filter powder] and sodium carbonate to one 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 in the potable water, paper, food, dairy, oil, textile and chemical industries. However, because of the alleged association of potable water containing aluminium sulfate with the incidence of Alzheimer's disease, its use has been discontinued in some cities.

7.6.3 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.
7.6.4 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 (SO42-) 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 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, clear or coloured, the solute may be regained from the solvent by evaporation and can pass through fine filters and permeable membranes.

7.6.7 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, usually 100 g of the solvent (1 dm3 of water) at room temperature, usually 20oC. 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 20oC. The solubility of gases decreases as the temperature rises. When a more concentrated solution is diluted with a solvent,
C1V1 = C2V2, where C1= original concentration, V1 = original volume, C2= final concentration, V2 = final volume.

7.6.8 Solubility of different salts
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.

7.7.1 Solubility of gases and temperature, carbon dioxide
See 3.34.3 Solubility of acidic oxide carbon dioxide in water
Use two identical plastic bottles or aluminium drink-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 freezer. 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.

7.7.2 Solubility equilibrium, solubility product
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, Ksp 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 Ksp the a precipitate will form until the product of the ion concentrations equals Ksp is large, then at equilibrium, the concentration of products is much greater than the concentration of the reactants.

7.7.3 Solubility rules
Solubility 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+.
Solubility 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.
Solubility 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 Pb2+ salt is slightly soluble, but more soluble in hot water.
4. All hydroxides are insoluble, except the Na+, K+ and NH4+ salt. The Mg2+ and Ca2+ salts are slightly soluble.
5. All nitrates are soluble.
6. All phosphates are insoluble, except the Na+, K+, NH4+ salts and some acid phosphates.
7. All common sodium, potassium and ammonium salts are soluble.
8. All sulfides are insoluble, except the Na+, K+, NH4+, Mg2+, Ca2+ and Ba2+ salts.
9. All sulfates are soluble, except the Ba2+, Pb2+, Ca2+ and Hg2+ salts. The Ag2+ 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.
1. Test if a salt is soluble in water. Select salts from the laboratory, e.g. 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, ethyl alcohol and copper (II) sulfate crystals, kerosene and petroleum jelly.
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 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 oC.
Table 7.7.5
Chemical 20oC 30oC 40oC 50oC
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 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.
7.7.6 Solubility and temperature, 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 oC along the horizontal axis. Solubility of potassium dichromate and temperature
Dissolve potassium dichromate in 50 mL of water at 60oC, until no more dissolves. The solution is saturated. Pour the clear solution into a second beaker. Let the temperature drop slowly to 40oC. 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.
7.7.8 Solubility of copper (II) sulfate and particle size
See 3.12: Solubility and particle size, sodium chloride, copper (II) sulfate
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.
7.7.9 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. Temperature affects solubility of sucrose
Repeat the experiment at 10oC intervals until 70oC and record the weight of sucrose dissolved. Heat the solution to 100oC. Pour it into an evaporating basin and leave to cool. Observe the cooling solution. Prepare toffee candy
Use the solubility of sucrose to make toffee candy. Prepare toffee or candy by heating the sucrose solution with milk or butter until it boils. Heat to 125oC. Pour it into paper cups and leave to cool. Constant stirring and temperature control is essential. The actors in movies who appear to fall through a sheet of glass without hurting themselves are in fact falling through a sheet of a special kind of clear toffee called "sugar glass". Prepare crystal rope
Make a saturated solution of sucrose (cane sugar). Tie a weight to a cotton thread and suspend it so that the weight and most of the thread are in the solution. As the solution cools, crystals form on the thread and weight. Take out the thread and weight. Hold the weight and set fire to the cotton. A rope of sugar crystals remains.
7.7.10 Solutions that 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. 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. 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. Solvent for Sudan III
Put drops of Sudan III dye in a test-tube. Add water and shake the test-tube. The dye does not dissolve. Add a little oil or melted butter and shake the test-tube. The dye dissolves in fats and in oils.
7.7.12 Rates of solution
The rate at which some solid dissolves in water can be increased 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.
7.7.13 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. Volume of gas dissolved in tap water
See 3.25: Separate gases dissolved in a water sample
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.
7.7.14 Fractional crystallization of sea water
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. However, when the sea water starts to evaporate and crystals start to form, calcium is the only cation near saturation. As water evaporates from the sea water, the order of precipitation is mainly calcium carbonate, calcite, and also calcium magnesium carbonate, dolomite, then calcium sulfate dihydrate, gypsum, then sodium chloride when the volume of the sea water is about 10% of the original, then potassium and magnesium salts. If you dissolve more carbon dioxide in the sea water, you would not precipitate more 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. Colloids and crystalloids
Colloids, sols, emulsions, gels, aerosols, foams
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 × 10-6m, (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 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, particulate air pollution
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. 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. Prepare mayonnaise
1. Beat 15 ml of edible oil with one egg yolk and water to produce a weak emulsion
2. Beat 120 ml of oil with one egg yolk and water to give a strong emulsion
3. Place egg yolks, lemon juice 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. Tyndall effect
Tyndall's experiment, "Scientrific", (commercial website)
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.
Fill 250 mL beaker with yellow potassium chromate solution (K2CrO4). 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. Size of colloidal particles
1. Use three 50 mL, beakers. Pour 20 mL copper (II) sulfate solution into Beaker 1. Add 20 mL skim milk solution to Beaker 2 and Beaker 3. Add 5 mL 3 M acetic acid to Beaker 2. to form the precipitate casein. Filter the contents of each beaker into three large test-tubes 1, 2 and 3. Use two 20 cm lengths of dialysis tubing, soak for ten minutes in distilled water to soften and tie up one end of each length. Pour the copper (II) sulfate solution from Beaker 1 into one set of dialysis tubing. Pour the skim milk from Beaker
2 into the other set of dialysis tubing. Tie the ends of both sets of tubing, rinse with deionized water and put into two separate beakers of water. Observe what happens after some time.
2. Repeat the last experiment using a mixture of starch and glucose solution inside the dialysis tubing. After 30 minutes test the water outside the dialysis tubing for water for glucose and starch. The next day, repeat the test for glucose and starch. Filter the starch and glucose solutions and observe whether the filter papers contain any residues. State your conclusions the sizes of colloids compared with the true solutions in these experiments. Aerosol
See diagram Aerosol can
1. Aerosols (fogs) are colloidal dispersions of particles of liquids or solids suspended in another gas, e.g. fly spray, mist, fog, clouds, smoke. So "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.
2. An aerosol spray can is a pressurized container with an attached spray mechanism is used to produce small amounts of product in a finely divided form, e.g. anti-perspirants, insecticides, medicines (nasal sprays), paint. A propellant is dissolved in the product to push it out when the aerosol can is opened. The pressure inside the aerosol can is much greater than atmosphere pressure. 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 (aerosol dispensers) means any non-refillable receptacles made of metal, glass or plastics and 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".
3. Prepare sunbeam mists. 1. Pour concentrated hydrochloric acid and concentrated ammonia solution into two watch glasses.
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. 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). Emulsions
An emulsion is liquid dispersed in another liquid, e.g. mayonnaise, cream, milk. A solid emulsion is a liquid dispersed in another solid, e.g. butter. An emulsion is the suspension of one liquid as fine droplets in another with which it does not mix. Emulsions are colloids with dispersed and continuous phases both liquids, e.g. oil in water. 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. By adding an emulsifying agent, e.g. soap, the two liquids, kerosene and water, will remain dispersed within each other.
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. Emulsifying agents are used to keep the phases dispersed so that the droplets remain suspended and the emulsion remains stabilized.
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 the cream floats above the milk. In homogenized milk, the droplets have been broken into smaller particles and dispersed to form one phase.
Mayonnaise is a mixture of olive oil dispersed in vinegar and stabilized by egg yolk or mustard. Margarine is an emulsion of water, flavours, colours, and vitamins in a semi-solid fat. Lotions, cosmetic creams and ointments are mostly emulsions of oils dispersed in water. Also, hair wave lotion and water glass used to preserve eggs. 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. 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. Temporary emulsions and permanent emulsions, kerosene, detergent
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. 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 salt water.
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. 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 so that fat particles remain suspended long enough for enzymes to digest them before absorption into the bloodstream. Bile salts are sold commercially as sodium tauroglycocholate. 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 commercial concentrated soybean milk in bags) into a clean 500 mL beaker and heat to about 80oC. 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 commercial 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. 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, upholstery, expanded plastics (foamed plastics), e.g. Styrofoam and other polymer rigid foam articles. A foam is a sort of cellular solid with a gas matrix. Solid foams are often highly flammable.
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! Gels
See Prepare sodium polyacrylate gels, acrylic sodium salt polymer, ASAP
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. 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.
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. Prepare gelatine gel
Water soluble protein from boiling collagen, e.g. horse hooves. Found in photographic emulsions, food jellies, adhesives.
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. 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. Metallic salts gels, calcium carbonate gel, calcium acetate gel
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 CaCO3.nH2O. 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. 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. Gellan gum
Water soluble tetrasaccharide hydrocolloid from two glucose residues, multi-functional gelling agent, in jellies, stabilizer in soya milk for soy protein suspension, food additive E418 emulsifier, stabilizer, thickener, stable at up to 120oC, alternative to agar, formed by fermentation of Sphingomonas elodea.
Trade names: Gelzan, Kelcogel, Phytadel, Gelrite. Carrageenans
Sulfated polysaccharide hydrocolloid produced from red seaweed, Rhodophyceae, Chondrus crispis, Eucheuma spinosum, Eucheuma cottonii, Gigartina, vegetable gum, food additive E407, used for gelling and thickening, can form firm gels with potassium ions, can form elastic gels and thixotropic fluids with calcium ions, and can form viscous non-gelling solutions, free flowing jellies or firm gels. Trade names: Genulacta, Gelgenuvisco. 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 solutions of solid or liquid particles in a gas, e.g. smoke, fog.
3. Lyophilic sols, (solvent loving sols) 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 that prevents them flocculating unless a strong salt solution is added to precipitate the colloid.
4. Lyophobic sols, (hydrophobic sols, solvent hating sols, water 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, from corn sugar (maize) fermentation, e.g. maize, by Xanthomonas campestris, forms stable viscous sols. Silver chloride precipitate 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. You can 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) 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.
Fe3+ (aq) + 2H2O (l) --> (colloidal) + 3H+ (aq) 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. Keep the solution.