topic05

School Science Lessons
Topic 5 Solutions and mixtures, mole, dilute acids with alkalis, acid-base indicators, alloys
Updated 2009-10-19
Please send comments to: J.Elfick@uq.edu.au
See: Interesting web sites

Table of contents
5.00 Solutions and mixtures
5.1.0 Mole, amount of substance
5.2.0 Prepare stock solutions, standard solutions
5.3.0 Dilute acids with alkalis
5.3.4 Acid-base indicators
5.4.0 Prepare solutions of known concentration
5.5.0 Alloys

5.00 Solutions and mixtures
2.41 Mix and dissolve (Primary)
2.42 Mix liquids (Primary)
3.80 Reactions that give out heat energy, exothermic reactions
3.81 Reactions that take in heat energy, endothermic reactions
3.82 Heat of a neutralization reaction
3.83 Heat of reaction when metals displace copper
5.01 Dilute solution
5.01.1 Concentrated solutions
5 01.2 Supersaturated solution of sodium ethanoate-3-water
5.01.4 Solutions used in making herbal remedies
3.17.1 "Magnetic" sugar cube dissolves
14.1.0 Exothermic reactions, reactions that give out heat energy
14.2.0 Endothermic reactions, reactions that take in heat energy

5.1.0 Mole, amount of substance
5.1.0.1 Avogadro's hypothesis, (Avogadro's principle), Avogadro's number box
5.1.1 Prepare molar solutions
5.1.2 Concentration of a sulfuric acid solution
5.1.3 Molar volume
5.1.4 Mass / volume relationships
5.1.5 Relative atomic mass of magnesium

5.4.0 Prepare solutions of known concentration
5.4.01 Concentrations and volumes
5.4.02 Preparation instructions
5.4.03 Molarity
5.4.04 Molality
5.4.05 Normal solution
5.4.06 Normal saline, physiological saline
5.4.07 Percentage solutions
5.4.08 Specific gravity concentrations
5.4.09 Calculation of concentration after dilution

13.1.6, Molar volume of oxygen prepared with hydrogen peroxide
23.5.02, Molar heat capacity, Cm
24.1.14, Molal freezing point constant of cyclohexane solvent
24.1.15, Molar mass of solute from depression of freezing point
3.32.1, Molar mass of air, composition of the atmosphere and greenhouse gases

5.3.0 Dilute acids with alkalis
5.3.1 Dilute acids
5.3.2 Dilute bases
5.3.2.1 Alkalis, potassium hydroxide (caustic potash) sodium hydroxide (caustic soda)
5.3.2.2 Ammonia solution, (10% ammonia solution)
5.3.2.3 Limewater, (lime water), ionization of calcium hydroxide
5.3.3 Bromine water

Acid-base indicators
5.53.01 pH and acid-base indicators, acidity and alkalinity, ionization of water
5.53.1 Test acid-base indicators
5.53.2 Prepare acid-base plant extract indicators
5.53.3 Dissociation constant, Ka
3.31.3 Tests for water with cobalt (II) chloride
5.3.4.1 Litmus paper, prepare litmus solution
5.3.4.2 Universal indicator
5.3.4.3 Cobalt (II) chloride paper
5.3.4.4 Heat sensitive paper, cobalt (II) chloride, ammonium chloride, (sal ammoniac)
12.8.0, Acid-base neutralization, acid with base forms a salt and water
12.9.4, Rainbow reactions, t-butyl chloride (2-chloro-2-methylpropane) with sodium hydroxide
19.1.5, Acid-base indicators in the home, in plants, grape juice, red cabbage

5.4.10 Series dilution
5.4.1 Ammonium molybdate
5.4.2 Calcium hydroxide, Ca(OH)₂
5.4.3 di-Potassium hydrogen orthophosphate, K₂HPO₄
5.4.4 di-Sodium hydrogen phosphate, Na₂HPO₄.l2H₂O
5.4.5 Ethanoic acid, (acetic Acid), CH₃COOH
5.4.6 Hydrochloric acid, HCl
5.4.7 Hydrogen peroxide, H₂O₂
5.4.8 Iodine solution, I₂
5.4.9 Sodium hydrosulfite, Na₂S₂O₄.2H₂O
5.4.10 Sodium hydroxide, NaOH
5.4.11 Starch solution
5.4.12 Sulfuric acid, H₂SO4
5.4.13 Nitric acid, HNO₃
5.4.14 Oxalic acid (ethanedioic acid), (COOH)₂.2H₂O
5.4.16 Sodium chlorate (V), NaClO₃
5.4.17 Sodium dihydrogen phosphate, NaH₂PO₄.2H₂O
5.4.18 Tin (II) chloride, SnCl₂.2H₂O

5.5.0 Alloys
5.5.0.0 Eutectic mixture
5.5.01 Low melting point alloys
5.5.02 Higher melting point alloys and parts by weight
5.5.03 Copper-zinc alloys, brass
5.5.04 Copper-tin alloys, bronze
5.5.05 Copper-aluminium alloys, bronze
5.5.06 Copper-nickel alloys
5.5.07 Tin-lead alloys
5.5.08 Alloys of "noble metals", Au, Ag, Pt, Pd
5.5.09 Cast iron alloy, steel
5.5.09.1 Paper clips and rusting
5.5.1 Alloy collection
5.5.2 Copper in brass

5.53.1.0 Test acid-base indicators
5.53.1.1 Bromothymol blue solution, acid-base indicator
5.3.4.1 Litmus, acid-base indicator
5.53.1.3 Methyl orange, acid-base indicator
5.53.1.4 Methyl red, acid-base indicator
5.53.1.5 Phenolphthalein, acid-base indicator
5.3.4.2 Universal indicator, acid-base indicator
5.53.1.7 Rose petal, acid-base indicator
5.53.1.8 Berry juices as acid-base indicators
5.53.1.9 Vegetable juices as acid-base indicators

5.00 Solutions and mixtures
The term solution refers to a homogeneous mixture of two or more components in the same phase, i.e. with no boundaries, where the atoms or molecules are interspersed, e.g. salt water, diamond. So a mixture of gases can also be referred to as a solution. However, the term solution commonly refers to an aqueous solution, a solute dissolved in the solvent water In this document the term solution refers to an aqueous solution unless otherwise indicated, e.g. an ethanol solution. Water has a high dielectric constant, insulator, so it is a ready solvent for ionic substances. The ions of a solute may interact with the molecules of a solvent, solvation. If two liquids may mix as molecules (miscible). Solid solutions occur in some alloys. A heterogeneous mixture has different phases in the same system, e.g. chalk dust in water, or has distinct substances where the atom or molecules are not interspersed, e.g. iron filings and sulfur.

5.01 Dilute solution
In this document a "dilute solution" means a 2-M solution unless otherwise specified. A "concentrated " acid or any other substance means "as supplied by commercial suppliers", e.g. concentrated hydrochloric acid is 36% w / w, unless otherwise specified.

5.01.1 Concentrated solutions
Concentration (formerly molarity) is the amount of substance dissolved per unit volume, symbol c, has unit mol litre^-1, mol per litre, mol l^-1, mol dm^-3. In this document, litre is shown as "litre" and not as "l". The "mass concentration" can be expressed as g cm^-3,and similar expressions. The "molal concentration" or "molality" can be expressed as mol kg^-1. For example, 1.00 molal KCl solution is made by dissolving 74.55 g of KCl in 1.00 kg of water. A saturated solution contains the maximum amount of solute at that temperature. In a saturated solution, the rate of loss of solute particles leaving the solution is equal to the rate of solute particles entering the solution, so the dissolved substance is in equilibrium with the undissolved substance. As the maximum equilibrium concentration depends on temperature, a saturated solution can become an unstable supersaturated solution by slow cooling. If you add a small crystal of the solute to a supersaturated solution, the excess solute will crystallize out of the solution.

5.01.2 Supersaturated solution of sodium ethanoate-3-water
Dissolve 125 g sodium ethanoate-3-water, CH₃CO₂Na.3H₂O, in 12.5 mL water. Heat to form a clear solution, cover with a watch glass and leave to cool. Hold a watch glass in the palm of your hand, pour in some solution then add a few crystals of sodium ethanoate-3-water. The supersaturated solution immediately crystallizes and you can feel heat given out. The exothermic property of the crystallization of saturated solutions is used in commercial "heat packs".

5.01.4 Solutions used in making herbal remedies
1. Herbal infusions of dried or fresh herbs, herbal teas, herbal tea bags , e.g. chamolile, fennel, lemon balm, vervain
2. Syrups, infusions or decoctions with honey or unrefined sugar
3. Decoctions of rough material. e.g. barks, berries , roots. The material in cold water is heated to boiling then simmered for over 20 minutes, e.g. dried liquorice root, Glycyrrhiza uralensis. However, valerian root, Valeriana officinalis, is macerated then steeped in cold water overnight.
4. Tinctures are made from herbs steeped usually in a 25% alcohol and 75% water mixture. The alcohol helps to extract the active principle of the herbal remedy and keeps the tincture preserved. The tincture may be stored for years in dark glass bottles. The alcohol may be ethanol or an alcoholic beverage , e.g. vodka or rum. A 1: 5 tincture is made with 100 g of herbs and 400 mL of alcohol / water mixture, e.g. cinnamon stick tincture, Cinnamomumzeylanicum. Tinctures may be boiled before use to produce alcohol-reduced tinctures for children and pregnant women.
5. Tonic wines are made from herbs steeped in wine, e.g. Korean ginseng, elecampane tonic wine, Inula helenium. The tonic wine can be kept for long periods if kept topped up with wine.
6. Hot infused oils are usually for external use, e.g. rosemary, comfrey, stinging nettle Urtica dioica
7. Cold infused oils, a slow process to produce massage oils and oils for compresses and poultices, e.g. St. John's wort Hypericum perforatum, pot marigold Calendula officinalis, melilot, Melilot officinalis.
8. Creams and ointments are a mixture of water with fats and oils. , e.g. an infused oil with beeswax and lanolin
9. Lotion and emulsions are water-based mixtures to relieve irritation or inflammation., e.g. chickweed lotion, Stellaria media

5.1.0 Mole, amount of substance
1. The mole concept and stoichiometry enable the determination of quantities in chemical processes. The mole, defined arbitrarily using the isotope carbon-12, is the basic quantity in stoichiometric calculations. Every chemical reaction can be represented by a balanced equation, whose coefficients indicate both the number of reacting particles and the reacting quantities in moles. Balanced equation can be used when determining whether reagents are limiting or in excess. The use of molarity for expressing concentration allows easy interconversions between volume of solution and moles of solute. The ideal gas equation may be used to relate the volume of a gas at defined temperature and pressure to its quantity in moles. One mole of any chemical compound has a mass equal to its relative molecular mass expressed in grams. One mole of any substance contains the same number of atoms or molecules. The number of particles in a mole is 6.02 X 10²³(Avogadro's constant, Avogadro's number).
2. The mole, symbol mol, is the SI unit for amount of substance. A mole represents how much substance that contains as many atoms or molecules (elementary units) as there are atoms in 0.012 kg of the carbon isotope carbon-12. A mole of a substance is the amount of that substance whose weight is equal to the molecular or formula weight. The molecular weight of H₂ = 2, so 1 mole of H₂ weighs 2 g. The molecular weight of CO₂ = 44, so 1 mole of CO₂ weighs 44 g. A mole of any substance has the same number of molecules, Avogadro's constant (formerly Avogadro's number) 6.022 X 10²³. Strictly speaking, to get a mole of a substance, weigh out its relative atomic mass or relative molecular mass in grams.
3. A proposed new definition of amount of substance, mole: A mole is such that the Avogadro constant is exactly 6.0221415 X 10²³ per mole.
4. A mole contains 6.023 x 10²³ single units, e.g. atoms, molecules, electrons.

5.1.0.1 Avogadro's hypothesis, (Avogadro's principle), Avogadro's number box
Equal volumes of all gases contain the same numbers of molecules, under identical conditions of temperature and pressure. So one mole of any substance contains the same number of particles. One mole of any gas, under identical conditions of temperature and pressure, has the same fixed volume, the molar volume (molecular volume) of a gas, 22.414 litres at s.t.p. (standard temperature and pressure) T = 273.15 K, P = 1 atmosphere (atm).
Avogadro's number box
A cube with sides of 28.2 cm has a volume of 22.4 litres at s.t.p.
A 22.4 litre box represents the volume of one mole at s.t.p.
Observe mole samples of carbon, iron, copper, zinc

5.1.1 Prepare molar solutions
1. State volumes in millilitres (mL) and litres (L). One millilitre (mL) is equivalent to one cubic centimetre (cc or cm³) for all practical purposes. State mass in grams (g). State molar solution in moles (M).
2. Molar solution
A molar solution, 1 M, contains one mole of the substance per litre of the solution. The energy change when 1 mole of solute dissolves in the solvent is called the heat of solution. Exothermic dissolving processes of a solid in a solvent are associated with high solubility. Endothermic dissolving processes of a solid in a solvent are associated with low solubility. A substance is "soluble" in a solvent if it dissolves to give a concentration > 0.1 M. The dissolving process of a gas in a liquid is exothermic. Solubility of gases in liquids decreases with temperature. When an ionic solid dissolves in water the cations and anions separate. Usually, ionic solids, e.g. sodium chloride NaCl, are soluble in water and non-ionic substances are insoluble. However, some ionic solids, e.g. silver iodide, AgI, is insoluble in water. A precipitate is a solid produced in solution. Differences in solubility can be used to separate mixtures of ions.
A molar solution (1 M solution) contains one mole of the solute dissolved in 1 litre of water.
3. Make a molar solution of MgSO₄. Calculate the total of the relative atomic masses of all atoms and express the total in grams:
Relative atomic mass Mg = 24.3
Relative atomic mass S = 32.1
Relative atomic mass O = 16.0 X 4 = 64.0
Total = 120.4 Relative molecular mass
Weigh 120.4 g of MgSO₄ and dissolve it in 1 litre of water. When making a molar solution, dissolve all the substance in less than one litre of deionized water, then add more deionized water until the volume is exactly 1 litre (making up to 1 litre.)
4. Make a molar solution of magnesium chloride crystals
If a substance contains water of hydration water of crystallization, e.g. magnesium chloride crystals (MgCl₂.6H₂O) the weight of the water is included in the weight of one mole:
Mole of atoms of Mg = 1 X 24.3 = 24.3 g
Moles of atoms of Cl = 2 X 35.4 = 70.8 g
Moles of atoms of H = 12 X 1.0 = 12.0 g
Moles of atoms of O = 6 X 16.0 = 96.0 g
Total = 203.1 Relative molecular mass
One mole of MgCl₂.6H₂O weighs 203.1 g. Weigh the solute to the nearest gram.

5.1.2 Concentration of a sulfuric acid solution
Measure a fixed volume, 25.0 mL of sodium hydroxide solution of known concentration, 0.10 M. Measure the volume of added sulfuric acid solution until the reaction is just complete, e.g. 27.5 mL.
H₂SO4 + 2NaOH --> Na₂SO4 + 2H₂O
1 mole + 2 moles
Calculate the concentration of sulfuric acid:
1000 mL of 0.10 M NaOH contains 0.10 moles
25.0 mL of 0.01 M NaOH contains 0.10 x 25 / 1000 moles
So number of moles NaOH = 0.0025
From the equation, 2 moles NaOH react with 1 mole H₂SO₄
0.0025 moles NaOH react with 0.0025 / 2 mole H₂SO₄
So number of moles H₂SO₄ = 0.00125
27.5 mL H₂SO₄ contains 0.00125 X moles
1000 mL H₂SO₄ contains 0.00125 x 1000 / 27.5 moles = 0.045
No. of moles per litre = 0.045
Concentration of H₂SO₄ solution = 0.045 M

5.1.3 Molar volume.
The molar volume of a substance is the volume occupied by 1 mole of it. One mole of any substance contains Avogadro's number of particles. Equal numbers of gas molecules occupy equal volumes. So the molar volumes of all gases are the same at the same temperature and pressure. One mole of any gas at s.t.p. occupies 22.4 litres (22,400 mL) (molar volume of gas at s.t.p., or gram molecular volume, G.M.V., at s.t.p.)
s.t.p. = 0^oC (273.15 K) and 760 mm Hg (101325 pascals, Pa).
The molar volume varies with temperature and pressure in accordance with the combined gas equation, P1 X V1 / T1 = P2 X V2 / T2. Find the number of moles of gas present by converting the volume measured under experimental conditions to the volume at s.t.p., then compare with the molar volume.
How many moles of gas are present in 320 mL methane at 27^oC and 600 mm pressure?
P1 x V1 / T1 = P2 x V2 / T2
600 x 320 / 300 = 760 x V2 / 273
Volume at s.t.p., V2 = 273 x 600 x 320 / 760 x 300 = 229.89 = approx. 230 mL
230 mL of the gas contains 230 / 22400 moles at s.t.p. = 0.01027 moles

5.1.4 Mass / volume relationships
Find the volume occupied by 1.60g of oxygen gas at s.t.p.
1 mole of oxygen gas occupies 22.4 litres at s.t.p.
32g of oxygen gas occupies 22.4 litres at s.t.p.
1.60 g of oxygen gas occupies 1.60 x 22.4 / 32 = 1.12 litres at s.t.p.

5.1.5 Relative atomic mass of magnesium
See diagram: 3.1.5: Magnesium ribbon
The molar volume of most gases at 0^oC and 1 atmosphere is 22.4 litres. The molar volume of most gases at 25^oC and 1 atmosphere is 24.4 litres.
In this experiment, the volume of hydrogen gas produced and mass of magnesium reacting with dilute hydrochloric acid are used to calculate the mass of magnesium that would be needed to produce one mole of hydrogen molecules, the relative atomic mass of magnesium.
1. Clean 4 cm of magnesium ribbon with fine emery paper and cut off a 3.5 cm length, weighing about 0.03 g. Use a top pan balance, accurate to +/- 0.001 g.
2. Pour 25 mL of 2 M hydrochloric acid into a 50 cm² burette. Very carefully pour 25 mL of water on top of the hydrochloric acid, leaving a space between the liquid and the top of the burette. The two solutions should not mix much.
3. Push the length of magnesium ribbon by the middle to be just inside into the open end of the burette. Curl the magnesium ribbon around so that it stays in place like a spring under tension.
4. Pour water into a beaker, close the top opening of the burette with your finger and quickly invert the burette so that the lower end is under the water
5. Clamp the burette to a burette stand and quickly note the inverted burette reading on the scale before the magnesium starts reacting with the acid.
6. When all the magnesium has reacted and no more gas bubble form, note the inverted burette reading again. Calculate the difference in burette readings, about 30.5 cm³.
7. Mg (s) + 2HCl (l) --> MgCl₂ (l) + H₂ (g)
So 1 mole of magnesium produces 1 mole of hydrogen molecules, i.e. 24.4 litres = 24,400 cm³ of hydrogen gas.
If 0.03 g of magnesium produces 30.5 cm³ of hydrogen gas, the mass of magnesium needed to produce 24,400 cm³ of hydrogen gas = 24,400 X 0.03 / 30.5 = 24 g. So the relative atomic mass of magnesium is 24.
8. You can use the gas equations to convert the volume of gas collected at room temperature and actual atmospheric pressure to conditions under standard
temperature and pressure. However, the hydrogen gas is mixed with water vapour so you would have to subtract the vapour pressure of water at that room temperature.

5.2.0 Prepare stock solutions, standard solutions
Prepare solutions of solutes dissolved in water: Weigh the calculated quantity of solid on a watch glass. Tip the solid into a clean, dry glass beaker. Use a wash bottle of deionized water to rinse the watch glass into the beaker. Stir the mixture until the solid has dissolved. If heat is used, the solution should be cooled before dilution. Pour the solution into a measuring cylinder or volumetric flask. Use a wash bottle to rinse the beaker several times with deionized water. Then pour the washings into the graduated container. Carefully add deionized water up to the desired level for the final volume of solution. Use a small quantity of the solution to rinse out a cleaned stock bottle. Pour the solution through a clean, dry filter funnel into the stock bottle. Use a stopper, then label the solution.
A standard solution is a solution with accurately known concentration.

5.3.0 Dilute acids with alkalis
Acids and alkalis are dangerous materials, especially when they are concentrated. Keep nearby, calcium carbonate or sodium hydrogen carbonate or whiting (for acids) and solid citric acid (for alkalis). Use safety glasses and nitrile chemical-resistant protective gloves when carrying out dilution of concentrated acids or alkalis, but be sure that the outside surfaces of gloves and bottles are dry to avoid accidental slipping of bottles when handled. Any acid or alkali that contacts the skin should be immediately washed off with copious quantities of water. Then apply dilute sodium bicarbonate solution to the affected area (for acid burns) or very dilute ethanoic acid (acetic acid) or vinegar solution or boracic acid (for alkali burns) to neutralize traces of the acid or alkali. Any spills of acids or alkalis should be diluted as above, before mopping up. For large spills, solid neutralizers such as solid sodium bicarbonate or whiting (for acids) and solid citric acid (for alkalis) should be used. Glacial ethanoic acid (acetic acid) concentrated hydrochloric acid, concentrated nitric acid and ammonia produce poisonous or extremely irritating gases or vapours. So use the fume cupboard when removing stoppers from bottles of these acids.

5.3.1 Dilute acids
See 7.0 Preparation instructions for acids and bases
Use commercially available 2 M acids.
Be Careful! Teachers who have not previously diluted concentrated to make dilute acids should ask an experienced teacher to help them!
Carefully add the concentrated acid in small quantities to 750 mL deionized water with continuous stirring.
BE CAREFUL! ALWAYS ADD ACID TO WATER!
When the solution is cool, add the rest of the water to total volume of one. Large quantities of heat are produced when you mix concentrated sulfuric acid with water. Dangerous spitting that is hazardous to eyes, skin and clothing, can result from the addition of water to the acid or from the too rapid addition of the acid to water. Allow the solution to cool before transferring to stock bottles.

5.3.2 Dilute bases
See 7.0 Preparation instructions for acids and bases
Be careful! Always use safety glasses and nitrile chemical-resistant gloves when handling strong bases!
Dilute concentrated bases as shown in the preparation Instructions. Add strong bases to water in a beaker in a sink. When the solution is cool, dilute with more water. Seal all bottles containing solutions of alkalis with plastic screw tops or rubber stoppers because alkalis attack glass stoppers making them hard to dislodge.

5.3.2.1 Alkalis: Potassium hydroxide (caustic potash) sodium hydroxide (caustic soda)
These alkalis are usually supplied as solid pellets in bottles that must be kept airtight because the pellets are hygroscopic. If the pellets stick together in a solid mass in the bottle use a stainless steel spatula to remove the mass then seal the bottle as quickly as possible. You must also weigh these solids quickly because of their tendency to absorb water vapour from the air. Dissolve the pellets in cold water with stirring, using a glass beaker, not a plastic container. As the solution generates a large amount of heat, continuously stir it until the solid has dissolved. Then allow the solution to cool before transferring it to stock bottles. Further dilution to the final desired volume in the stock bottle may be necessary. Stopper the bottle firmly and invert several times to mix.

5.3.2.2 Ammonia solution (10% ammonia solution)
Ammonia solution produces dangerous irritating ammonia fumes. You should carefully unstopper bottles of ammonia solution in a fume cupboard to prevent damage to eyes from the fumes. Dilute by adding ammonia solution to water in a fume cupboard.

5.3.2.3 Limewater (lime water) ionization of calcium hydroxide
See diagram 3.34.1: Limewater test for carbon dioxide
Ionization of calcium hydroxide, Kb = 3.5 X 10^-2Ca(OH)² <--> CaOH⁺ + OH^- CaOH⁺ <--> Ca²⁺ + OH^-
Limewater (calcium hydroxide solution) is a weak alkali made up by adding slaked lime (calcium hydroxide solid) to deionized water in a large stock bottle. Shake vigorously and leave to stand. Calcium hydroxide solid is only slightly soluble in water. When the white solid has settled as a fine white sediment, carefully siphon off or pour off the clear limewater above the sediment without disturbing the sediment. To replenish the limewater, add more deionized water to the sediment in the stock bottle, shake and leave to settle. The settling process to produce clear limewater may take several days. The clear limewater prepared as above may be diluted 1:2 as needed.

5.3.3 Bromine water
1. Cool 1 mL ampoule of bromine water Br₂ in a refrigerator, break with forceps under 200 mL water.BE CAREFUL! Use safety glasses and nitrile chemical-resistant gloves
2. Put the ampoule inside a test-tube of an appropriate size. Fix a short piece of wide bore rubber tubing to the end of the test-tube. Seal the other end of the rubber tubing with a wide bore glass rod. Move the ampoule inside the test-tube so that the top is inside the rubber tubing. Break the ampoule by squeezing the rubber tubing with pliers. Put in water and carefully cut the rubber tubing to release the bromine.

5.3.4.1 Litmus paper, prepare litmus solution, test acid-base indicator
1. Litmus paper contains several dyes, including the very sensitive purple-red dye azolitmin, red pH 4.5, blue pH 8.3. Use red books pH 5 and blue books pH 8.
2. Prepare litmus solution. Boil 10 g crushed litmus powder in 500 mL water for five minutes. Leave to stand, then filter the solution and store in a bottle. Add drops of nitric acid until a purple colour appears. Then filter and store in a bottle but keep the solution exposed to the air. Use a fresh solution before testing pH. Litmus solution can from blue to red. Put a finger width of litmus powder in the test-tube, and add water. Shake to make the powder dissolve. Add a finger width of tartaric acid powder until the colour changes to red. Sodium carbonate can change this red liquid to blue. Add a finger width of sodium carbonate to the red liquid until the colour changes back to blue. Shake the test-tube to help the mixing.
3. Prepare litmus solution. Grind 250 g of granular litmus and put it in a flask with 500 mL of 40% ethanol. Heat and boil the solution for one minute. Decant the liquid to storage leaving a residue in the flask. Add 500 mL of 40% ethanol to the residue. Heat and boil the solution for one minute then add it to the stored decanted liquid. Centrifuge the solution and adjust the volume of the supernatant to 1000 mL with 40% ethanol. Add M hydrochloric acid drop by drop until the solution becomes purple. Test the solution by boiling 10 mL of deionized water, leave to cool, add one drop of the litmus solution. Mix the drop with the water and the water should become mauve in colour. For laboratory use, make a 2.5% solution on the litmus indicator in deionized water.

5.3.4.2 Universal indicator
1. Prepare universal indicator. Dissolve the following in 500mL ethanol: 0.0250 g thymol blue, 0.0625g methyl red, 0.5 g phenolphthalein, 0.25 g bromothymol blue. Dilute this solution to 1 litre with deionized water. Add drops of 0.05 M sodium hydroxide until mixture is green.
2. Universal indicator test paper, (FLAM 1142), is mixture of acid-base indicators that causes a colour change for each change in pH value over a wide range. Note the colour chart on the bottle or package: Red pH 1-3 (strong acid solution), Orange pH 4-5 (weak acid) (Pink pH 4), Yellow pH 6 (weaker acid), Green or pale green pH 7 (neutral), Blue or green-blue pH 8 (very weak base), Indigo pH 9-10 (weak base) (Blue pH 9) (Blue-violet pH 10), Violet pH 11 to 14 (very basic solution)
3. Use 2 drops of Universal Indicator to 10 mL of test solution. Test the pH value of the following substances: baking soda solution, demineralized water, dill pickle juice, distilled water, household ammonia, liquid soap, pineapple juice, sodium bicarbonate solution, tap water, lemon juice, limewater, sodium hydroxide solution, vinegar, washing soda, "Windex" window cleaning solution.
4. Dissolve in 500 mL ethanol: 0.0250 g thymol blue, 0.0625 g methyl red, 0.5000 g phenolphthalein, 0.2500 g bromothymol blue. Dilute this solution to 1 litre with deionized water. Add drops of 0.05 M sodium hydroxide until the mixture is green.
5. Test colours of universal indicator
Slowly neutralize limewater containing universal indicator, by adding acid drop by drop. Describe what you see. The starting colour is blue. As citric acid is added the colour changes because the acid keeps weakening the alkali, neutralizes it exactly (pale green colour), and thereafter gradually builds up its own strength.

red	orange	yellow	pale green	green	blue	violet
strong acid	weak acid	weaker acid	neutral	weaker acid	weak acid	strong alkali

6. Make a very dilute solution of citric acid by adding 7 g of citric acid to one litre of water. Put two test-tubes of limewater in a beaker. Put a piece of universal indicator paper in each beaker. Stir until an inky blue solution forms then remove the universal indicator paper. Add citric acid to the beaker, drop by drop. Note that although much acid has to be added to form the different colours, it is the last drop that causes one colour to change into another. If you miss a colour by adding the acid too quickly, add limewater to the beaker to restore the blue colour and start again. To make the colours more easily seen, put the beaker on a white tile. Citric acid is not a strong acid.

5.3.4.3 Cobalt (II) chloride paper
Use this chemical to test for the presence of water. Dissolve 5 g of cobalt (II) chloride in 100 mL deionized water. Cut strips of absorbent paper 5 cm x 1 cm and soak in the cobalt (II) chloride solution. Remove strips, drain and spread flat on trays. Place trays in an oven at 100^oC until the strips are blue. Put strips in a bottle containing dry silica gel (blue in colour) or anhydrous calcium chloride. Keep the bottle sealed, preferably in a desiccator. If the paper turns pink, heat it again as described above until it turns blue again. Do not handle the paper with fingers as moisture from the skin will affect it.

5.3.4.4 Heat sensitive paper, cobalt (II) chloride, ammonium chloride (sal ammoniac)
Add cobalt (II) chloride solution to ammonium chloride solution (sal ammoniac). Dilute the solution until it is pale pink. Soak paper in the solution and leave to dry. The paper turns bright green colour when heated.

5.53.01 pH and acid-base indicators, acidity and alkalinity, ionization of water
The pH tests use an indicator which changes color with changes in the concentration of hydrogen ions, or the acidity of the solution. The pH scale (Soren Peter Sorensen 1868 - 1939) is a scale of acidity and alkalinity that runs from pH 0, most acid, to pH 14, most alkaline. A neutral solution has pH = 7, an acid solution has pH < 7, and a basic or alkaline solution has pH >7. The term "pH" stands for "power of hydrogen" and measures the concentration of hydrogen ions in water. The pH scale is logarithmic so a pH 4 solution is ten times more acidic than a pH 5 solution.
pH = -log₁₀(H⁺), where (H⁺) = concentration of hydrogen ions. (OH^-) = concentration of hydroxyl ions. For water (H⁺)(OH^-) = 1 X 10^-14 at 25^oC. Pure water is neutral where (H⁺) = (OH^-) = 1 X 10^-7M, i.e. at pH 7. For acid solutions (H⁺) is greater than (OH^-), so pH is less than 7 (0 to 7). For alkaline solutions (H⁺) is less than (OH^-), so pH is greater than 7 (7 to 14).
To convert pH to hydrogen ion molar concentration, [H₃O⁺] = Antilog(-pH), so if pH = 2.55, [H₃O⁺] = Antilog(-2.55) = 2.8 X 10^-3 M.
Ionization of water
2H₂O <--> H₃O⁺ + OH^- (25^oC) Ka =1.00 X 10^-14, pKa = 14.00
2H₂O <--> H₃O⁺ + OH^- (0^oC) Ka = 0.11 X 10^-14, pKa = 14.94

5.53.1.0 Test acid-base indicators
See 5.3.4: Acid-base indicators
Acid-base indicators change colour in acidic or basic solutions. They may be weak acids that dissociate and change colour in alkaline solutions.
Test the following indicators with acids and bases, e.g. dilute HCl, lemon juice or vinegar, ammonia solution, dilute sodium hydroxide solution, limewater, tap water, demineralized water.

5.53.1.1 Bromothymol blue solution
pH 6.0 yellow to pH 7.6 blue, in 20% alcohol solution. Dissolve 0.5 g of bromothymol blue in 500 mL of water. Add a drop of ammonia solution to turn the solution deep blue in colour.

5.53.1.3 Methyl orange
It is best for solutions with concentration > M/5. Mix 1 g of commercial methyl orange powder with water. Use 2 drops for each 25 mL of solution in a titration. Methyl orange solution can change from orange to red
Put a finger width of the methyl orange powder in the test-tube and add water to half fill the test-tube. Shake to make the powder dissolve. Add tartaric acid powder until the colour changes to red. Sodium carbonate change this red liquid to orange. Add a finger width of sodium carbonate to the red liquid until the colour changes back to orange again. Shake the test-tube to help the mixing.

5.53.1.4 Methyl red
1. It is a sensitive indicator for titrating weak organic bases and ammonia. Dissolve 1 g of commercial powder in 500 mL of 60% alcohol. Use 2 drops for 25 mL of liquid in a titration.
2. Dissolve 0.04 g of methyl red in 40 mL of ethanol and make up to 100 mL with water.

5.53.1.5 Phenolphthalein
(FLAM +13^oC 1170) pH 10 red, with excess alkali colourless again
1. Add 5 g to 500 mL of ethanol, add 500 mL water. Stir.
2. Dissolve 1 g of commercial powder in 500 mL of 50% alcohol. Add drops of this phenolphthalein solution to 100 mL of 0.5 mol / litre sodium hydroxide solution until a deep pink colour appears. Divide this solution into 3 test-tubes. Leave the first test-tube as a control. Add drops of HCl to the second test-tube until the pink colour disappears. Add 3 pellets of solid sodium hydroxide to the third test-tube. Shake to dissolve. The pink colour reappears.
3. Add colourless phenolphthalein indicator to limewater. The liquid turns pink. Blow into the liquid through a drinking straw. The pink colour disappears and the liquid becomes cloudy.

5.53.1.7 Rose petal acid-base indicator
Boil red rose petals in some water until the petals have almost lost their colour and a pink solution forms. Test this pink solution with acids and bases.

5.53.1.8 Berry juices as acid-base indicators
Test the juices from stewed blackberries, blackcurrants, and raspberries. Also, mix a spoonful of fruit jam with warm water then filter it to get a colourless liquid. Test the solution with acids and bases.

5.53.1.9 Vegetable juices as acid-base indicators
Test the green water juices from boiled cabbage, boiled beetroot and other juices.

5.53.2 Prepare acid-base plant extract indicators
See appendix C: Antacids
1. Use plant extracts to "indicate" whether a substance is acidic or basic. Select brightly coloured flowers or leaves, e.g. rose, Bougainvillaea, hibiscus, geranium, red carnation (light red with acid and bright green with alkali), sweet pea, snapdragon, pansy, tulip, willow herb. The colours are usually caused by anthocyanin water soluble pigments that change colour with change in pH. Boil a fresh unboiled beetroot, red cabbage, tomato skins (colourless in acid and deep yellow in alkali), blackberry or blackberry jam, damson, elderberry. Squeeze or grind the plant material with a mortar and pestle with a mixture of 2 mL of acetone and 2 mL of methylated spirit. Filter the solution, collect the filtrate, and label the indicator, e.g. "rose extract". Rose extract colours may be scarlet-pink at pH 1, pale pink at pH 3, green at pH 4, yellow-brown at pH 7 and orange at pH 12. Use universal indicator solution to test the plant indicators. Indicators made from plants are mostly red with acids but yellow, green or purple with alkalis.
2. Test common substances and note the colour change of the plant extract indicator, e.g. ammonia solution, antacid tablet solution, baking soda solution, bleaching powder solution, coconut milk, coffee grounds, fertilizer solution, fruit juice, lemon juice, lemonade, limewater, red cabbage juice, saliva, soap solution, sugar solution, vinegar, tap water, tea bag in hot water, whitewash. Estimate the range of pH tested by the plant extract indicators.
3. Soak cut pieces of red cabbage leaf in boiling water for 30 minutes then remove them. Pour cabbage water into 3.1 water, it stays violet, 3.2 white vinegar, it turns red, 3.3 baking soda solution or ammonia solution, it turns green.
4. Boil shredded red cabbage for 15 minutes then squeeze out the juice. Fry an egg. When the "white" of the egg is about to change from colourless to white as the protein albumin denatures add some red cabbage juice. The "white" of the egg turns green.
Also, use the liquid from a container of pickled cabbage. Fresh grape juice turns red in acid lemonade and blue in alkaline dishwater.
5. Put spots of plant extract indicators on absorbent paper and leave to dry. Put one drop of lemon juice on each spot and note the colour change. Note the colours given by sodium bicarbonate solution, washing soda, limewater and a dilute solution of sodium hydroxide. These are alkaline, basic, substances. Note whether they all give the same colour. Plant extracts can act as indicators to test whether a substance is acidic or basic.
6. Add a few drops of sodium bicarbonate solution to 1 mL of flower extract indicator in a test-tube. Then add lemon juice and note any colour change.
7. Repeat the experiment with limewater and indicator followed by dilute hydrochloric acid. Note any colour change. Note whether the original colour returns after by adding more limewater. Note how many times the indicator colour can change before the test-tube is full.
8. Betacyanin pigments cause the red colour of beetroot. These are acid/base indicators have optical stability at pH 4 to 5 but are structurally unstable at extremes of pH. So the red colour in urine after eating beetroot depends on urine pH and the pigments not being broken down by digestion processes. Eating excess beetroot as in borscht soup usually causes red or pink urine.
9. Wave a bluebell flower closely over an ant nest. The angry ants rush out to squirt formic acid on the bluebell flower and the blue pigment in the petals turns red.

5.53.3 Dissociation constant, Ka
The dissociation constant is the equilibrium constant of a reversible dissociation including the ionization reactions of acids and bases in water. It is also called the acid dissociation constant or acidity constant.
pKa = -log₁₀ (1/Ka)
Ionization reactions at 25^oC
pKa, the larger the value the weaker the acid

12.9.4 Rainbow reactions, t-Butyl chloride (2-chloro-2-methylpropane) with sodium hydroxide
Make a pH 12 solution by adding 10 drops of 0.1 M NaOH to 100 mL water, in a 250 mL beaker. Add universal indicator to produce a distinct colour. Start with universal indicator. Use a second 250 mL beaker to mix by pouring the solution back and forth between the two beakers or put a magnetic bar into the solution and start the stirrer motor at a fast rate. Add 15 drops of t-butyl chloride (2-chloro-2-methylpropane) to the solution and begin mixing. Observe any colour changes. After 40 seconds add universal indicator and observe any colour changes. The full range of colour changes (purple, blue, cyan, emerald-green, lime-green, yellow, orange, orange-red, take about two minutes. The changes in the middle are more rapid than the changes at either extreme. Use different indicators to show different colour changes and different induction times:

Indicator	Colour Change	Induction Time	Indicator	Colour Change	Induction Time
Methyl red	Yellow	Initial	Phenolphthalein	Pink	Initial
"	Orange	40 seconds	"	Pale pink	40 seconds
"	Green	45 seconds	"	Colourless	45 seconds
Bromothymol blue	Blue	Initial	Bromophenol blue	Blue	Initial
"	Green	40 seconds	"	Green	75 seconds
"	Yellow	45 seconds	"	Yellow	80 seconds
Thymol blue	Blue	Initial	m-Cresol purple	Violet	Initial
"	Green	50 seconds	"	Red	52 seconds
"	Yellow	52 seconds	"	Yellow	54 seconds

Different formulations of universal indicator may give differing times and colour changes. For a wide range universal indicator, double the amounts of each reactant. The reaction is an "S_N1" reaction, i.e. a nucleophilic substitution reaction, in which the chlorine radical is replaced by an hydroxyl radical. As H⁺ions are produced in solution in the reaction, the OH^- are gradually neutralized as the reaction proceeds, and eventually excess H+ are produced. Thus, the pH of the solution progressively falls because of reaction.
(CH₃)₃-C-Cl + H₂O --> (CH₃)₃-C-OH + HCl

2. Prepare two solutions: 0.1 M 2-chloro-2 methylpropane (t-butyl chloride) in ethanol (1 g per 100 mL) and 0.01 M sodium hydroxide. Put 5 mL 0.1 M C₄H₉Cl in a test-tube. In another test-tube put 5 mL 0.1 M NaOH, 10 mL water and two drops of any one of the following indicators. Mix the solutions back and forth once and observe for the colour change that occurs after an induction period. With equal volumes 0.01 M sodium hydroxide and 0.1 M C₄H₉Cl the colour changes with universal indicator were: Purple to blue (on mixing) blue to green (after 12 seconds) green to yellow (after 15 seconds) yellow to orange (after 25 seconds total). Cooling the solutions greatly slows the reaction, increasing the induction period, e.g. with iced water, the methyl red change took more than 50 seconds.

Indicator	Colour Change	Induction Time
Methyl red	Yellow to red	6 seconds
Bromothymol blue	Blue to yellow	5 seconds
Thymol blue	Blue to yellow	10 seconds
Phenolphthalein	Red to colourless	11 seconds
Universal indicator	Purple to blue Blue to green Green to pink	Instantly 8 seconds 10 seconds

5.4.0 Prepare solutions of known concentration
See 8.0 Preparation instructions for salt solutions
When preparing solutions the number of molecules of water of crystallization shown on the bottle should be checked against the table. If different, you will need a new calculation. The formula after each substance is for its commercial form. In making up solutions, deionized water (or demineralized water) should be used. If none is available, tank water is a suitable substitute. Unless otherwise stated, the amount of chemical shown in the table or calculated should be dissolved in about one quarter of the volume of deionized water needed and then diluted to the needed volume. The symbol, M, stands for molarity, i.e. the number of moles of solute per litre of solution. Usually, the masses are not exact fractions of the formula mass. Where concentrations are needed other than those shown in preparation Instructions, calculate the quantity of solute or solvent needed. Some solutions, e.g. Ca(OH)₂, limewater, react with the carbon dioxide of the air, so the solutions should be kept in sealed containers.

5.4.01 Concentrations and volumes
Concentration is the quantity of dissolved substance, solute, per unit quantity of solvent in a solution, or concentration is the number of ions or molecules of a substance, in a given volume of solvent.
Concentration, c (formerly molarity) is expressed as moles per litre, mol l^-1, mol dm^-3
Concentration can also be expressed as moles per cubic metre or moles per cubic decimetre
Mass concentration, rho = g cm^-3, kg l^-1, kg dm^-3
Molal concentration, molality, is expressed as mol kg^-1
1 litre = 1 cubic decimetre (1 dm³) = 1000 cubic centimetres (1000 cm³)

5.4.02 Preparation instructions
See 7.0: Preparation instructions for acids and bases
If a solution more concentrated than in preparation Instructions is needed, dissolve more solute in one litre of water. For a 4 M sodium hydroxide solution, the table 6.1 shows that 80 g sodium hydroxide is needed to make 1 litre of 2 M solution. So dissolve (4 / 2)M X 80 g = 160 g sodium hydroxide in 1 litre of water to make a 4 M solution.
If a solution less concentrated than in preparation Instructions is needed, dissolve less solute in one litre of water. For a 0.01 M sodium iodide solution, the table shows that 15 g of sodium iodide is needed to make 1 litre of 0.1 M solution. So dissolve (0.01 / 0.1) M X 15 g = 1.5 g sodium iodide in 1 litre of water to make a 0.1 M solution.
A stock bottle or "Winchester" bottle may have a volume of 2.4 litres. To make 2.4 litres of 2 M Sulfuric, 55 mL of concentrated acid is needed to make 1 litre of 1 M solution. So use (2.4 / 1) mL X (2 / 1) M X 55 mL = 264 mL concentrated acid to make 2.4 litres of 2 M solution

5.4.03 Molarity
1. The molarity of a solution is the number of moles of solute per litre of the solution. A molar solution, 1 M, contains 1 mole of solute per litre of solution. 1 M = 1 gram molecular weight of solute / litre of solution, so 1 M NaCl solution contains (23 + 35.5) 58.5 g of NaCl in 1 litre of solution. 1 M HCl = (1 + 35.5) 36.5 g / L.
2. Molarity is the concentration of the solution expressed as the number of moles of the dissolved substance per dm³ (litre) of solution. A molar solution has a concentration of one mole per dm³ (litre).

5.4.04 Molality
1. The molality of a solution is the number of moles of solute per kg of solvent. 1 molal = moles per Kg water
2. Molality is the concentration of a solution expressed as number of moles of the dissolved substance per kilogram of solvent.

5.4.05 Normal solution
1. A normal solution contains 1 gram equivalent weight of solute per litre of solution. An equivalent weight = molecular weight / valence, so 1N NaCl contains 58.5 g NaCl / litre, 1N HCl contains 36.5 g HCl / litre, 1N H₂SO₄ contains 98 / 2 = 49 g H₂SO₄ / litre. So molarity X valence = normality. A 2 mol per litre solution of sulfuric acid, 2M, is a 4 N solution. The term normal solution is obsolescent and is no longer used or taught in chemistry teaching but you may see it being used in the chemical industry.
2. Normality was a concentration unit formerly used for acid, bases, oxidizing agents and reducing agents, based on the concentration of H⁺ and OH^- in a solution. A normal solution has a concentration of one gram equivalent per dm³ (litre). One litre of a normal solution contains the weight in grams of the solute that is equivalent to 1 g of replaceable hydrogen gas. So 25 mL of N HCl would react exactly with 25 mL of N NaOH or 250 mL of N / 10 NaOH. For monobasic acids, the equivalent is numerically the same as the molecular weight. For dibasic acids, the equivalent is numerically equal to half the molecular weight. So a 2M (2 moles per litre) solution of H₂SO₄ = 4N solution of H₂SO₄.

5.4.06 Normal saline, physiological saline
This solution is used in medicine as a diluent, plasma substitute in intravenous drips because it is isotonic with human blood plasma, although it may have a slightly higher osmality. It contains 0.9% NaCl + 5% d-glucose, dextrose, which is isotonic with 0.9% NaCl and is used to decrease the concentration of sodium.

5.4.07 Percentage solutions
1. Percentage W / W solution, percentage of weight of solute in the total weight of solution, number of grams of solute in 100 grams of solution. To make a 10% (W / W) NaCl solution, dissolve 10 g NaCl in 100 g of solution.
2. Percentage W / V solution, percentage of weight of solute in the total volume of solution, number of grams of solute in 100 mL of solution. To make a 10% (W / V) NaCl solution, dissolve 10 g of NaCl in 100 mL of solution. (This is the most common percentage solution.)
3. Percentage V / V solution, percentage of volume of solute in the total volume of solution, number of millilitres of solute in 100 mL of solution. To make a 10% (V / V) methanol solution add 10 mL of methanol to 100 mL of solution, assume water unless otherwise stated.

5.4.08 Specific gravity concentrations
Concentrated acids and other liquid reagents may have concentration expressed as the specific gravity of the solution as the number of grams per millilitre. If specific gravity of HCl = 1.1885, then each mL of solution contains 1.1885 g of HCl, i.e. 1188.5 g / L = 1188.5 / 36.5 = 32.56 M. If other substances are present in the solution the percentage assay may be stated, e.g. 38% HCl means that 38% of 1188.5 is HCl
1188.5 X 0.38 = 451.63 g / L = 45.163 g / 100 mL
A 38% of HCl at 20^oC has specific gravity 1.1885 and contains 451.6 g HCl per litre

5.4.09 Calculation of concentration after dilution
V1C1 = V2C2, where V = volume, C = concentration as percentage or M or N) 1 = the more concentrated solution, 2 = the less concentrated, new dilute, solution

5.4.10 Series dilution
Prepare different percentage concentrations

100% soln	Prepare 100 mL of saturated solution, then filter.
0.5% soln	0.5 mL of saturated solution, add water to 100 mL
0.1% soln	20 mL of 0.5% solution, add water to 100 mL
0.05% soln	50 mL of 0.1% solution, add water to 100 mL
0.01% soln	20 mL of 0.05% solution, add water to 100 mL
0.005% soln	50 mL of 0.01% solution, add water to 100 mL
0.001% soln	20 mL of 0.005% solution, add water to 100 mL

5.4.1 Ammonium molybdate (NH₄)₆Mo₇O₂₄.4H₂O	Add 45 g to water containing 120 mL 10% ammonia. Add 120 g NH₄NO₃then dilute.
5.4.2 Calcium hydroxide Ca(OH)₂ Saturated (limewater) Ca(OH)₂	Add 10 g to water, shake, let settle, decant clear liquid.
Calcium hydroxide Ca(OH)₂0.02 M	Add 1.48 g to water, add excess, filter off precipitate.
5.4.3 diPotassium hydrogen orthophosphate K₂HPO₄ 0.1 M	Add 17.4 g to water.

5.4.4 di-Sodium hydrogen phosphate Na₂HPO₄.l2H₂O 0.1 M

Add 35.8 g of Na₂HPO₄.l2H₂O to water.

Na₂HPO₄.2H₂O 0.1M	Add 17.8 g Na₂HPO₄.2H₂O to water.
Ethanoic Acid (Acetic Acid) CH₃COOH 17M	As supplied
5.4.5 Ethanoic Acid (Acetic Acid) CH₃COOH 2 M (approx.)	Dilute 120 mL concentrated (glacial) or use 360 mL 33% acid.
Ethanoic Acid (Acetic acid) CH₃COOH 2 M	Dissolve 117 mL of 17.15 M acid (99% w / w 1.048 g / mL).
Hydrochloric Acid HCl concentrated 10 M	As supplied
5.4.6 Hydrochloric acid HCl 2 M	Dissolve 173 mL of 11.55 M acid (36% w / w 1.17 g / mL).
5.4.7 Hydrogen peroxide H₂O₂For laboratory use	To 20 volume solutions (6%) add twice the volume of water.

5.4.8 Iodine solution I₂Alcoholic	Add flakes to methylated spirit.
Iodine solution I₂/ I₂KI Aqueous I₂/ I₂KI	Add l g iodine crystals and 5 g potassium iodide to 50 mL water. Dilute to 100 mL.
5.4.9 Sodium hydrosulfite Na₂S₂O₄.2H₂O	For 100 mL solution for use as oxygen gas absorber, add 16 g Na₂S₂O₄.2H₂O + 13 g NaOH to 100 mL water. Add 4 g B-anthraquinone sulfonate to improve the reagent.
Sodium hydroxide NaOH 2M (approximate)	Add 80 g to water in a beaker in a sink. When cool, dilute with water. Store in a bottle with a rubber stopper. Use safety glasses and nitrile chemical-resistant gloves.
5.4.10 Sodium hydroxide NaOH 2 M	Add 81.6 g (98% NaOH) to water in a beaker in a sink. When cool, dilute with water. Store in a bottle with a rubber stopper. Use safety glasses and nitrile chemical-resistant gloves.
Sodium hydroxide (for CO₂ absorption)	Add 330 g to water.
5.4.11 Starch solution, 1%	Add 10 g starch to cold water to make a paste. Then dilute to 100 mL with boiling water. Let it boil, stir then leave to stand.
Sulfuric acid concentrated H₂SO₄18 M	As supplied.
5.4.12 Sulfuric acid 2 M	Add 113 mL of 17.75 M acid (97% w/w 1.83 g / mL) slowly to water with stirring. Use safety glasses and nitrile chemical-resistant gloves.
Nitric acid HNO₃ concentrated 16 M	As supplied
Nitric acid HNO₃2 M	Dilute 125 mL of concentrated acid. Use safety and nitrile chemical-resistant gloves!
5.4.13 Nitric acid HNO₃2 M	Dissolve 130 mL of 15.40 M acid (69% w/w 1.41 g / mL).
5.4.14 Ethanedioic (Oxalic) acid (COOH)₂.2H₂O 0.1 M	Add 12.6 g of crystals to water.
5.4.15 Phenolphthalein indicator	Add 5 g to 500 mL of ethanol, add 500 mL water. Stir.
5.4.16 Sodium chlorate (V) NaClO₃0.1 M	Dilute 10% solution with equal volume water.
5.4.17 Sodium dihydrogen phosphate NaH₂PO₄.2H₂O 0.1 M	Add 15.6 g to water
5.4.18 Tin (II) chloride SnCl₂.2H₂O 0.1 M	Add 22.6 g to 100 mL concentrated hydrochloric acid, then dilute with water. Add pieces of tin.

5.5.0 Alloys
Alloys have metallic properties and are composed of two or more elements. Alloys can be compounds, solid solutions, e.g. gold copper alloys and alum crystals, or just mixtures of the components. Alloys are used for coins, type metal, heating elements, construction, machinery, e.g. brass taps, bronze coins or ship propellers, duralumin aircraft parts, stainless steel knives. Light metal alloys, e.g. duralumin, 94.5% aluminium, 4% copper, 1% manganese, 0.5% magnesium have strength and lightness for construction of aircraft and racing cars. Coins and jewellery are made of a precious metal mixed with another metal to make them harder and not wear easily. Iron alloys are tougher than pure iron so are used to make machines, e.g. stainless steel is made from iron chromium nickel alloys. Magnetic alloys, e.g. "Alnico" (iron, aluminium, nickel and cobalt) are used to make permanent magnets for refrigerator doors and radio speakers. However, electromagnets must be turned on or off so a different iron nickel alloy is used. One piece of alloy metal may contain several solid phases because of the limited solubility of one metal in another. Metals may be suddenly cooled, quenched, to prevent phase equilibrium being established.

5.5.0.0 Eutectic mixture
An eutectic mixture is a mixture of two substances at the composition yielding the lowest melting point. A dystectic mixture is a mixture of two substances at the composition yielding the highest melting point. The melting point of the alloy is lower than the melting points of any of the metals it contains. The eutectic temperature is the lowest temperature at which both solid components of a mixture are in equilibrium with the liquid phase.
Eutectic temperature and % composition by weight
183^oC: Sn 63.0% Pb37.0%
198^oC: Sn 91.0% Zn 9.0%
221^oC: Sn 96.5% Ag 3.5%
227^oC: Sn 99.2% Cu 0.8%
When the temperature of a mixture of 63.0% tin and 37.0% lead cools to 183^oC, the liquid mixture freezes sharply. So above 183^oC the mixture is liquid and below 183^oC the mixture is solid. An eutectic mixture has a sharp melting point as if it were a pure substance. This mixture (electricians' solder) has a comparatively low melting point, conducts electricity, "wets" copper wire and sets very quickly so components are not overheated.
When the temperature of mixture of 30% tin and 70% lead (plumbers' solder) cools to 250^oC, the liquid mixture starts to solidify but it does not freeze sharply. It finally solidfies at 183^oC, giving the plunber time to ensure a good setting of a joint.

5.5.01 Low melting point alloys
Use a Bunsen burner. These alloys have a melting point lower than any of the constituent metals, e.g. Wood's metal, has melting point 80^oC, but the lowest melting point of the constituents, tin, has melting point 232^oC. A sprinkler stop valve fire control system may have a metal plug in the fire sprinkler bulb made of a bismuth alloy that melts at 155^oC to open the sprinkler and put out the fire. Similarly electrical fuse wire in a circuit melts when heated to a certain temperature when excessive electric current flows through it. Common solder remains in a semi-melted state at around 200^oC so it can be managed easily. For alloys containing bismuth and lead, first melt them together then add the other ingredients. The temperature should not be higher than necessary to prevent excess oxidation.
The parts shown are by weight.

Alloy	Lead	Tin	Bismuth	Cadmium
Electrical fuse	8.5	2.5	1.3	0
Solder	1	1	0	0
Wood's metal	4	2	7	1

5.5.02 Higher melting point alloys and parts by weight
Use a burner for solder (melting point 250^oC) but use a furnace for the others. Melt the copper first then add the other metals. Brass alloys is very resistant to corrosion so it is commonly used for taps. Bronzes are very hard and are used for bearings. There are many specialist bronzes, e.g. phosphor bronze, aluminium bronze, bell metal for casting bells, gunmetal for gears. Copper alloys have a special designation, e.g. CZ copper zinc brass, PB phosphor bronze, CT copper tin, CN copper nickel, LG leaded gunmetal.

Alloy	Copper	Tin	Zinc
Bronze	80	5	15
Brass, malleable	58	0	42
Brass, casting	72	4	24

5.5.03 Copper-zinc alloys, brass
1. 5% Zn alloy, gilding metal, used for cheap jewellery, low value coins, 2. 10% Zn, commercial brass, 3. 15% Zn, red brass, 4. 30% Zn, used for cartridge and shell cases, 5. 2% Zn + 88% Cu + 10% Sn, admiralty gun metal, was used for guns

5.5.04 Copper tin alloys, bronze
1. 10% Sn, carrots bronze, 2. P added to give stronger phosphor bronze

5.5.05 Copper-aluminium alloys, bronze
10% aluminium, aluminium bronze (looks like 18 carat gold!) Cu Al Ni alloys used in Australian $1 and $2 coins and in English one pound and two pound coins. The Australian one dollar "kangaroos" coin is 92% copper, 6% aluminium, 2% nickel. The United Kingdom one pound coin is 70% copper, 24.5% zinc, 5.5% nickel.

5.5.06 Copper-nickel alloys
As copper nickel sheathing it releases copper, and loses weight, to protect against barnacle growth on ships, 70% copper, 30% nickel used in marine condensers, CU NI Zn alloy, German silver, nickel silver, monel metal, used for plumbing, a base for silver platted middle value coins. The United States "nickel" (0.05 US dollars) is 75% copper, 2% nickel.

5.5.07 Tin-lead alloys
1. 60% Sn 40% Pb, electrician's solder, has a sharp eutectic freezing point lower than either tin or lead. 2. 30% Sn 70% Pb, plumbers solder, freezes slowly. 3. 35% Pb Sn, pewter for dishes

5.5.08 Alloys of "noble metals" Au, Ag, Pt, Pd
Purity of gold is measured in carats (ct). 24 carat gold is pure gold. 14 carat gold is 14 parts pure gold and 14 parts copper or other metal. 18 carat gold is 18 parts of gold in 24 parts of alloy, usually copper. The Australian uncirculated $100 coin is 91.67% Au, 4.17% Ag and 4.17% Cu. Sterling silver is an alloy of 92.5% silver and 7.5% copper. (For diamonds, which are not alloys but pure carbon, the term "carat" means 0.2 grams.)

5.5.09 Cast iron alloy, steel
When iron ore is heated in a carbon fire two reactions occur. Cast iron contains about 4% carbon
3Fe₂O₃ + 11C --> 2Fe₃C + 9CO (g)
Fe₂O₃ + 3C --> 2Fe + 3CO (g)
Fe₃C is called cementite. Impurities are removed by adding a limestone flux that forms a glassy slag that can be converted to insulation fibres, rock wool.
Pure iron can be made by hammering to form wrought iron, with < 0.25% carbon: Fe₃C + FeO --> 4Fe + CO (g).
Alloy cast iron contains some additional combination of Ni, Cr, Cu, Mo. to obtain the high temperature form austenite or the low temperature ferritic form. If the austenite is cooled very quickly by quenching, it forms the very hard martensite steel. Tempering between 220^oC and 450^oC oxidizes the carbon in the steel to soften it and make it more ductile.
Plain carbon steel contains up to 2% carbon and up 0.8% manganese, 0.3% silicon and 0.5% sulfur and phosphorus. Low alloy steels contain in addition up to 5% Mn, Ni, Cr, Va, Mo. Stainless steel contains 12-5% Cr to produce the stainless chromium oxide film on the surface to prevent corrosion. So stainless steel must be kept clean to maximum availability of oxygen to the chromium atoms.
Case-hardening produces a hard surface layer in steel, either by heating in a carbon rich medium followed by quenching and tempering, or by rapidly heating the surface of a high carbon steel above the ferrite / austenite transformation temperature, 550^oC, followed by quenching and tempering.

5.5.09.1 Paper clips and rusting
Observe a paper clip used to clip together an old pile of paper. Note any rust marks. Rusting starts where the paper clip is closest to the paper because there is the least exposure to oxygen gas to allow the chromium layer to produce protective oxides.

5.5.1 Alloy collection
Make a collection of alloys found in the home, school laboratory and workshop.

5.5.2 Copper in brass
Add dilute nitric acid to brass, pure copper and iron filings and compare the reactions. Yellow brass may be 70% copper and 30% zinc.