topic05

School Science Lessons
Topic 5 Acid-base indicators, alloys, mole, prepare solutions of known concentration and standard solutions, series dilution, solutions and mixtures
2012-01-26 SP
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites

Table of contents
5.6.0 Acid-base indicators
5.5.0 Alloys
5.7.0 Concentration, Molarity
5.1.0 Mole, amount of substance
5.4.0 Prepare solutions of known concentration
5.3.0 Prepare stock solutions, standard solutions
5.2.0 Series dilution
5.0.0 Solutions and mixtures

5.6.0 Acid-base indicators
Order online: Disappearing Ink, acid / base indicator, thymolphthalein
1.0.0 Acid-base indicators (list)
12.8.0 Acid-base neutralization, acid with base forms a salt and water
19.1.5 Acid-base indicators in the home
5.6.8 Berry juices as acid-base indicators
5.6.3 Bromothymol blue solution
5.6.5.2 Cresol red
5.6.11 Dissociation constant, Ka
5.3.7 Litmus paper, prepare litmus solution, test acid-base indicator
5.6.4 Methyl orange
5.6.5 Methyl red
5.6.5.1 Methyl violet
Methyl yellow
5.6.1 pH meters and acid-base indicators, acidity and alkalinity, ionization of water
5.6.6 Phenolphthalein
5.6.10 Prepare acid-base plant extract indicators
5.6.12 Rainbow reactions, t-butyl chloride (2-chloro-2-methylpropane) with sodium hydroxide
5.6.7 Rose petal acid-base indicator
19.1.3 Solid acids, add sodium carbonate
19.1.2 Solid acids, pH
19.1.1 Solid acids, solubility
12.3.1 Taste of acids, solid acids in the home
5.6.2 Tests for acid-base indicators
5.3.8 Universal indicator
1.0 Universal indicator, Quantity of indicator per 10 L
5.6.9 Vegetable juices as acid-base indicators

5.5.0 Alloys
5.5.0 Alloys
5.5.01 Amalgams
5.5.2 Alloy collection
5.5.11 Alloys of "noble metals", Au, Ag, Pt, Pd
5.5.12 Cast iron alloy, steel, wrought iron
5.5.02 Coins
5.5.8 Copper-aluminium alloys, bronze
5.5.9 Copper-nickel alloys
5.5.7 Copper-tin alloys, bronze
5.5.6 Copper-zinc alloys, brass
5.5.3 Copper in brass alloy
5.3.14 Corrosion of alloys, restore bronze coin
5.5.1 Eutectic mixture
3.64 Heat treatment of steel needles, annealing, quenching, tempering
5.5.5 Higher melting point alloys and parts by weight
5.5.4 Low melting point alloys
3.63 Melting point of metals and alloys
5.5.13 Paper clips and rusting
3.61 Tin-lead alloys, make with a casting mould
3.62 Tin-lead alloys and pure metals, hardness
5.5.10 Tin-lead alloys, different constituents

5.5.02 Coins
Alloys of "noble metals", Au, Ag, Pt, Pd: 5.5.11
Coin cells, electricity from two coins: 33.3.3
Coins on a slope (Primary): 6.11
Conduction of heat by a coin on paper: 4.21
Copper-aluminium alloys, bronze: 5.5.8
Copper-zinc alloys, brass: 5.5.6
Copper-nickel alloys: 5.5.9
Electricity from two coins: 4.51
Inertia tricks (coin on card trick): 4.13
Restore bronze coins, corrosion of alloys: 15.3.14
Silver mirror tests for aldehydes, Tollens' tests for acetaldehydes: 16.3.5

5.7.0 Concentration, Molarity: 5.1.7
Concentration and rate of reaction: 3.7.1
Concentrations and volumes: 5.1.5
Concentration calculation after dilution: 5.1.13
Concentration of a sulfuric acid solution: 5.1.2
Concentration of hydrogen peroxide: 3.49a

5.1.0 Mole, amount of substance
5.1.0 Mole, amount of substance
3.32.1 Composition of the atmosphere and greenhouse gases, Molar mass of air
5.1.0.1 Avogadro's hypothesis, (Avogadro's law, 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 Concentrations and volumes
5.1.6 Preparation instructions
7.0 Preparation instructions for acids and bases
4.0 Preparation instructions for reagents
8.0 Preparation instructions for salt solutions
5.1.7 Molarity, concentration
5.1.8 Molality
5.1.9 Normal solution, normality
5.1.10 Normal saline, physiological saline
5.1.11 Percentage solutions
5.1.12 Specific gravity (relative density) concentrations
5.1.13 Concentration calculation after dilution
5.1.14 Relative atomic mass of magnesium
3.3.3 Measure the size of the stearic acid molecule
24.1.14 Molal freezing point constant of cyclohexane solvent
22.5.02 Molar heat capacity, C_m
13.1.6a Molar volume of oxygen prepared with hydrogen peroxide
24.1.15 Molar mass of solute from depression of freezing point

5.4.0 Prepare solutions of known concentration
5.4.0 Prepare solutions of known concentration
1.6 Iodine solution, test for starch, biology solution
5.4.1 Ammonium molybdate
5.4.2 Calcium hydroxide, Ca(OH)₂
5.4.3 Dipotassium 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
5.4.9 Sodium hydrosulfite, Na₂S₂O₄.2H₂O
5.4.10 Sodium hydroxide, NaOH
5.4.11 Starch solution
5.4.12 Sulfuric acid, H₂SO4
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

5.3.0 Prepare stock solutions, standard solutions
5.3.0 Prepare stock solutions, standard solutions
5.3.0.1 Dilute acids with alkalis
5.3.1 Prepare dilute acids
5.3.2 Prepare dilute bases
5.3.3 Alkalis: Potassium hydroxide (caustic potash) sodium hydroxide (caustic soda)
5.3.4 Ammonia solution (10% ammonia solution)
5.3.7 Litmus paper, prepare litmus solution, test acid-base indicator
12.19.9.1.1 Prepare bromine water
5.3.5 Prepare limewater, ionization of calcium hydroxide
5.3.8 Universal indicator
5.3.9 Cobalt (II) chloride paper
5.3.10 Heat-sensitive paper, cobalt (II) chloride, ammonium chloride (sal ammoniac)

5.0.0 Solutions and mixtures
5.0.0 Solutions and mixtures
2.41 Mix and dissolve (Primary)
2.42 Mix liquids (Primary)
3.17.1 "Magnetic" sugar cube dissolves
5.01 Dilute solutions
5.02 Concentrated solutions
5.03 Supersaturated solution of sodium ethanoate-3-water, "heat pack"
5.04 Solutions used in making herbal remedies, infusion, decoction, tincture
14.1.0 Exothermic reactions, reactions that give out heat energy
14.2.0 Endothermic reactions, reactions that take in heat energy

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. 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 solutions
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.

5.02 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 a small crystal of the solute is added to to a supersaturated solution, the excess solute will crystallize out of the solution.

5.03 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. Feel the heat given out. The exothermic property of the crystallization of saturated solutions is used in commercial "heat packs".

5.04 Solutions used in making herbal remedies
The active ingredient needed from the plant must be harvested, stored and preserved. Soon after harvest spread the plants to dry and inactivate the enzymes. Dry roots and bark in the sun but not plants where the active ingredient is a volatile oil, e.g. lemon balm.
1. Herbal infusion of leaves and flowers, tisane, tea of dried or fresh herbs, herbal teas, herbal tea bags, e.g. camomile, fennel, lemon balm, vervain
Pour boiling water over the herbs, apply a lid, then leave to steep for about ten minutes. The lid prevents evaporation of volatile oils. Strain the liquid then drink hot or cold. The infusion can be reheated but not boiled. A common dosage is 100 mL three times a day before meals. A single dose may contain 5 g herbs to 100 mL water. To make them more palatable sweeten infusions with honey.
2. Syrups, infusions or decoctions with honey or unrefined sugar, e.g. liquorice, Glycyrrhiza glabra, marshmallow Althaea officinalis, wild cherry Prunus serotina. Add two parts by weight of white cane sugar to one part of infusion or decoction. Sip 10 mL very three hours.
3. Decoctions of rough material. e.g. barks, berries, roots, woody parts. The material is first crushed, cut or broken into small pieces. For example add 30 g chopped liquorice root, Glycyrrhiza uralensis or dandelion Taraxacum officinalis to 600 mL cold water, heat to boiling, then simmer for over 20 minutes. Add water to make up to original volume then take as an infusion. However, valerian root, Valeriana officinalis, is macerated then steeped in cold water overnight. Store decoctions in a refrigerator and use them within a day.
4. Tinctures are made from about 50 g of finely cut or powdered herbs steeped in 500 mL of 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. brandy, gin, run, vodka. A 1: 5 tincture is made with 100 g of herbs and 400 mL of alcohol / water mixture, e.g. cinnamon stick tincture, Cinnamomum zeylanicum. Tinctures may be boiled before use to produce alcohol reduced tinctures for children and pregnant women. 5 g of herbs in infusion is equal to about 5 mL (one teaspoon) of tincture.
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. Herbs are used externally as a poultice, compress and oil. To make a poultice place crushed herbs in a cloth, e.g. cheesecloth, fold the cloth around the herbs and soak the bundle in hot water. Squeeze out the hot water and then apply the poultice to the affected part of the body. Replace the poultice before it has cooled completely. A compress is soft cloth soaked in an infusion or decoction and usually applied very hot. Oils for ear drops or massage are made from finely ground herbs in a light vegetable oil, e.g. olive, grape seed, safflower oil. Keep the oil in a tightly lidded jar. Use 50 g of herbs in 500 mL of oil. Shake the mixture every day for two weeks then strain and seal the oil.
7. Hot infused oils are usually for external use, e.g. rosemary, comfrey, stinging nettle Urtica dioica
8. 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.
9. Creams and ointments are a mixture of water with fats and oil, e.g. an infused oil with beeswax and lanolin. To make a soothing oil, heat 480 g almond oil, 120 g white beeswax and 120 g anhydrous lanolin in a saucepan on a water bath until they have all melted. Heat 400 mL of a fresh infusion or 200 mL of a tincture to the same temperature and add to the saucepan. Transfer the saucepan to a hot plate and simmer until all the water is boiled off. Leave the mixture to cool and store in a screw top glass jar.
10. Lotion and emulsions are water-based mixtures to relieve irritation or inflammation, e.g. chickweed lotion, Stellaria media.
11. Plasters, compresses and poultices are used to apply substances directly to the skin by using cotton bandages soaked in the mixture or applying tea bags.

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) is the number of atoms or molecules in one mole of a substance, (Amedeo Avogadro 1776-1856), N_A = 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 mole is the number of carbon-12 atoms whose mass equals 12 grams.
4. 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.
5. A mole contains 6.023 x 10²³ single units, e.g. atoms, molecules, electrons.

5.1.0.1 Avogadro's hypothesis, (Avogadro's law, Avogadro's principle), Avogadro's number box
[Avogadro's number, relative atomic mass, relative molecular mass and relative formula mass, molar mass, molar volume, molarity, empirical and molecular formulas, percentage composition]
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 STP. (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 STP.
A 22.4 litre box represents the volume of one mole at STP.
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, are 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 STP occupies 22.4 litres (22,400 mL) (molar volume of gas at STP, or gram molecular volume, G.M.V., at STP)
STP = 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 STP, 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 STP, V2 = 273 x 600 x 320 / 760 x 300 = 229.89 = approx. 230 mL
230 mL of the gas contains 230 / 22400 moles at STP = 0.01027 moles

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

5.1.5 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)

5.1.6 Preparation instructions
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

5.1.7 Molarity, concentration
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).
3. 1 M (1 molar solution) = 1 mol / L = 1 mol / dm³ = 1 mol dm^-3 = 1000 mol / m³
4. molar concentration, c_i = amount of constituent, n_i / volume of the solution, V
5. SI unit for molality = mol / kg.

5.1.8 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.

5.1.9 Normal solution, normality
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 it is still 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.

5.1.10 Normal saline, physiological saline
This solution is used in medicine as a 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.

5.1.11 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.

5.1.12 Specific gravity (relative density) 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

5.1.13 Concentration calculation 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.

5.1.14 Relative atomic mass of magnesium
See diagram 5.1.14: Relative atomic mass of magnesium
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 (3.5 mm standard 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. The liquid level in the burette must start on the graduated scale. If it is not on the scale turn the tap on and off quickly to let the level drop to be on the scale.
6. When all the magnesium has reacted with the downwards diffusing acid and no more gas bubble form because all the magnesium has reacted, 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. 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 subtract the vapour pressure of water at that room temperature.

5.2.0 Series dilutions
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

5.3.0 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.

5.3.0.1 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 Prepare dilute acids
See 7.0 Prepare 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 concentrated sulfuric acid is mixed 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 Prepare 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.

5.3.3 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.

5.3.4 Ammonia solution (10% ammonia solution)
Ammonia solution produces dangerous irritating ammonia fumes. 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.

5.3.5 Prepare limewater, 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^-
1. 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.
2. Prepare limewater. Slaked lime dissolves to a slight extent in water, and the solution is called limewater. Shake a teaspoonful of slaked lime with two thirds of a test-tube of water for a minute or two. Filter the milky liquid. The filtrate is limewater. Test it with a piece of red litmus paper. The paper will turn blue, showing that limewater is an alkali. Blow into the lime-water through a glass test-tube or drinking-straw. The carbon dioxide gas in the breath will turn the limewater milky. To keep a supply of lime-water always available put a tablespoon of slaked lime into a large plastic bottle and fill up the bottle with water. The clear limewater can be poured off into a smaller bottle as is needed, the large bottle being replenished with water each time.

5.3.7 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.

5.3.8 Universal indicator
1. Prepare universal indicator. Dissolve the following in 500 mL 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 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). Observe the colour of the indicator in coloured solutions either by holding the test-tube up to the light or by dropping the indicator on the top of the solution and observing the colour at the interface.
3. Use 2 drops of Universal Indicator to 10 mL of test solution. Test the pH value of the following substances: Chlorox solution (0.5 M ammonia solution), baking soda solution, bicarbonate of soda (sodium bicarbonate solution), demineralized water or deionized water, dill pickle juice, dilute sulfuric acid solution 0.5 M, distilled water, household ammonia, liquid soap, pineapple juice, tap water, lemon juice (citric acid solution), limewater (0.02 M calcium hydroxide solution), milk of magnesia (magnesium oxide), salt solution (sodium chloride solution), soda water (carbonic acid solution), sodium hydroxide solution, vinegar (0.05 M ethanoic acid solution), 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 the reaction. 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 a colour is missed 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.

5.3.9 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.

5.3.10 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.

5.4.0 Prepare solutions of known concentration
See 8.0 Prepare 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, a new calculation is needed. 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.

5.4.1.0 Prepare solutions of known concentration

5.4.1 Ammonium molybdate (NH₄)₆Mo₇O₂₄.4H₂O	Add 45 g to water containing 120 mL 10% ammonia. Add 120 g NH₄NO3 then 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.
5.4.5 Ethanoic Acid (Acetic Acid) CH₃COOH 17M	As supplied
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).
5.4.6 Hydrochloric Acid HCl concentrated 10 M	As supplied
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.
5.4.9 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.
5.4.10 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.
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.
5.4.12 Sulfuric acid concentrated H₂SO₄18 M	As supplied.
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.
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
1. (Latin: alligere, bind together) 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.
2. Alloys are a compound, solution or mixture of two or more metals. A few alloys contain a non-metal, e.g. steel. The first known alloy was bronze, a mixture of copper and tin, used to make utensils and swords, spears and shields about 6000 years ago. Later came brass, copper and zinc, steel, iron and carbon and many other alloys. Usually alloys have characteristics different from the component metals. For example they usually melt at lower temperature and are harder than the constituent metals. The fusion of metals to form alloys is often done under a flux that may promote liquefaction, prevent volatilization and unnecessary exposure to the air.

5.5.01 Amalgams
See: Mercury and dental amalgam
Amalgams are alloys of mercury with most other metals. Silver / mercury amalgams were commonly used in dentistry but nowadays are replaced by hard plastic. Previously the backs of mirrors and glass balls were coated with 30 parts mercury / 70 parts tin amalgams. All amalgams are decomposed by heat into mercury vapour and the metal residue. Copper amalgam, Viennese metal, was used for cementing metals. Zinc amalgam was used in electric batteries. Gold amalgams were used for gilding. Lead amalgams were used to solder metals. Nowadays, amalgams are seldom used for fear of mercury poisoning when the amalgams decompose.

5.5.1 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 percentage 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 solidifies at 183^oC, giving the plumber time to ensure a good setting of a joint.

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

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

5.5.4 Lower 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

5.5.5 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 are very resistant to corrosion so are 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

5.5.6 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

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

5.5.8 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.

5.5.9 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.

5.5.10 Tin-lead alloys, different constituents
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

5.5.11 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.)

5.5.12 Cast iron alloy, steel, wrought iron
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.

5.5.13 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.

5.6.1 pH meters and acid-base indicators, acidity and alkalinity, ionization of water
The pH tests use an indicator which changes colour with changes in the concentration of hydrogen ions, or the acidity of the solution. The pH scale (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
As the pH scale has no upper or lower limit it can be lower than 0 or higher than 14, but if higher, a base must be dissolved in an aqueous solution to bind with hydrogen ions. Some pH meters have a range of pH 1.0 to pH 15.0, with an extended range resolution of 0.1 pH and accuracy of +/- 0.2 pH, in operating temperature 0 to 50^oC. A pH meter is actually a voltmeter that uses a hydronium ion selective glass electrode to measure ion concentrations in the vicinity of the probe.
Commercial pH meter, digital, range 1.0 to 15.0 pH, extended range resolution 0.1 pH, accuracy +/- 0.2 pH, operating temperature 0 to 50^oC, waterproof
Commercial soil pH meter, Soil tester, measure soil pH and moisture

5.6.2 Tests for acid-base indicators
See: 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.

5.6.3 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.

5.6.4 Methyl orange
pH < 3.1 red to pH > 4.4 yellow-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.
1. Use as a 0.1% solution in 20% ethanol
2. 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.

5.6.5 Methyl red
pH <4.2 red to pH > 6.3 yellow
1. It is a sensitive indicator for titration of 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.

5.6.5.1 Methyl violet, C.I. basic violet 1
pH 0.2 yellow to pH 3.0 violet
Use as 0.1 % solution microscopy stain

5.6.5.2 Cresol red, C₂₁H₁₂NaO₅S, acid-base indicator, o-Cresolsulfonephthalein, triarylmethane dye
Cresol red used to monitor the pH in aquaria, Irritates eyes, do not ingest, Quantity of indicator per 10 mL: 2.8
m-cresol, 3-hydroxytoluene, Highly toxic by all routes, lung irritant vapour, highly corrosive to skin
m-cresol, Solution < 12%, Not hazardous
Cresols: 16.1.3.2.

5.6.6 Phenolphthalein
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.

5.6.7 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.

5.6.8 Berry juices as acid-base indicators
Test the juices from stewed blackberry, blackcurrant, and raspberry. 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.

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

5.6.10 Prepare acid-base plant extract indicators
Order online: Goldenrod Indicator Paper, acid-base indicator, tumeric
Order online: Red Cabbage Colour Crystals, acid-base indicator
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 "egg white" of the egg is about to change from colourless to white as the protein albumen denatures, add some red cabbage juice. The "egg 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. Add lemon juice to a cup of black tea and the colour becomes less intense. The red-brown pigments in black tea, thearubigin polyphenols, are weak acids with coloured anions. The hydrogen ions from lemon juice restrict the the ionization so the tea has a lighter colour.
10. 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.

5.6.11 Dissociation constant, Ka
1. Dissociation refers to the break up of a molecule into smaller molecules, atoms or ions. In a buffer solution of the salt of a weak acid with a weak acid, the dissociation of the weak acid is negligible but the salt dissociated completely into ions.
2. 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
For pKa, the larger the value the weaker the acid.

5.6.12 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