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5.0 Solutions, mixtures, alloys, concentration, mole

Table of contents

5.5.0 Alloys

5.7.0 Concentration, Molarity

5.1.0 Mole, amount of substance

5.0.0 Solutions and mixtures

See: Chemistry (Commercial)
5.5.0 Alloys
5.5.11 Alloys of noble metals and coinage metals
5.5.2 Amalgams
Babbitt's metal, alloy for bearings (Sn 5-90%, Sb 7-10%, Cu 1.5-6%, Pb 5-48%)
5.5.12 Cast iron alloy, steel, wrought iron
Stainless steel
5.5.8 Copper-aluminium alloys, bronze
5.5.14 Copper-chromium alloys, cupaloy
5.5.9 Copper-nickel alloys
5.5.7 Copper-tin alloys, bronze
5.5.6 Copper-zinc alloys, brass
5.5.3 Copper, brass and bronze alloys
15.3.14 Corrosion of alloys, restore bronze coins
5.5.1 Eutectic mixture
German silver, silvery cheap alloy of copper, zinc and nickel.
Manganin, resistance alloy, 86% Cu, 12% Mn, 2% Ni, for low temperatures because almost zero temperature coefficient of resistance
Misch metal, mixed metal alloy, Ce, La, Nd, cigarette lighter flint, deoxidizer
Pinchbeck, cheap alloy resembling gold, 83.6% copper, 16.4% zinc
5.5.10 Tin-lead alloys, different constituents

5.5.5 Higher melting point alloys and parts by weight
5.5.4 Lower melting point alloys
5.5.15 Nitinol memory wire
5.5.13 Paper clips and rusting

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

5.1.0 Mole, amount of substance
5.1.0 Mole, amount of substance
5.1.01 Atomic mass, atomic weight Avogadro's hypothesis, Avogadro's number box
5.1.5 Concentrations and volumes
5.1.4 Mass / volume relationships
5.1.8 Molarity
22.5.02 Molar heat capacity, Cm
5.1.03 Molar mass, MM
37.42.1 Molar mass of air, Composition of the atmosphere and greenhouse gases
5.1.3 Molar volume
5.1.7 Molarity, concentration
5.1.02 Molecular mass, molecular weight, MW
5.1.10 Normal saline, physiological saline
5.1.9 Normal solution, normality
5.1.11 Percentage solutions, calculations
5.1.6 Preparation instructions, calculations
5.1.12 Specific gravity (relative density) concentrations, calculations

5.1.2 Concentration of a sulfuric acid solution
24.1.14 Molal freezing point constant of cyclohexane solvent
13.1.6 Molar volume of oxygen prepared with hydrogen peroxide
24.1.15 Molar mass of solute from depression of freezing point
5.1.1 Prepare molar solutions
5.1.14 Relative atomic mass of magnesium
3.3.3 Size of stearic acid molecule

5.0.0 Solutions and mixtures
5.02 Concentrated solutions
5.01 Dilute solutions
5.04 Solutions for herbal remedies
Experiments Conductivity of solutions
14.2.0 Endothermic reactions, reactions that take in heat energy
14.1.0 Exothermic reactions, reactions that give out heat energy
3.17.1 "Magnetic" sugar cube dissolves
4.29 Mix liquids with water (Primary)
4.28 Mixing and dissolving (Primary)
5.3.3 Prepare alkalis
5.3.4 Prepare ammonia solution
5.3.1 Prepare dilute acids
5.3.2 Prepare dilute bases
5.3.5 Prepare lime water
5.3.6 Prepare standard sodium carbonate solutions
10.3.2 Shrinking mixture of liquids
5.04 Solutions for herbal remedies
5.04.0 Types of solutions for herbal remedies
5.04.1 Infusions, herbal infusions
5.04.2 Syrups, herbal syrups
5.04.3 Decoctions, herbal decoctions
5.04.5 Tonic wines, herbal tonic wines
5.04.4 Tinctures, herbal tinctures
5.04.6 Poultices, herbal compresses
5.04.7 Infused oils, hot herbal infused oils
5.04.8 Creams and ointments, herbal creams and ointments
5.04.9 Excipients, herbal excipients
3.8.0 Prepare herbal tinctures

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. 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 chemical 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 a supersaturated solution, the excess solute will crystallize out of the solution.

5.04.0 Types of solutions for herbal remedies
1. infusion, 2. decoction, 3. tincture, 4. poultice, 5. ointment
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.

5.04.1 Infusions
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.
Use an earthenware pot with a tight fitting lid.
Do not use aluminium containers.
Pour boiling water over the herbs, apply a lid, then leave to steep for up to 4 hours.
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.
Suitable herbs for infusions:
| Nettle
| Oat straw, oatstraw (Avena sativa)
| Red clover (Trifolium pratense)
| Russian comfrey |

5.04.2 Syrups
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 and sip 10 mL very three hours.

5.04.3 Decoctions
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.
Strain the mixture and discard the plant matter.
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.

5.04.4 Tinctures
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 medicine and keeps the tincture preserved.
Store the mixture in a screw top jar, leave in a warm place and shake the jar every day for two weeks.
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
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.

5.04.5 Tonic wines
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.

5.04.6 Poultices
Herbs are used externally as a poultice, compress and oil.
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.
A poultice is prepared from 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.
The bruised leaves or roots are applied directly to the skin.
Dried herbs should first be made into a paste using hot water or apple cider vinegar.
Replace the poultice before it has cooled completely.
A compress is prepared with soft cloth soaked in an infusion or decoction and usually applied very hot.
A compress does no contain the original plant matter.
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.

5.04.7 Infused oils
Hot infused oils are usually for external use, e.g. rosemary, comfrey, stinging nettle Urtica dioica.
Cold infused oils are prepared by 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.

5.04.8 Creams and ointments
See: Water Bath (Commercial)
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.
Lotion and emulsions are water-based mixtures to relieve irritation or inflammation, e.g. chickweed lotion, Stellaria media.

5.04.9 Excipients
Excipients are inactive substances combined by dissolving or mixing with active ingredients in pharmacy to add bulk, to improve taking
of a medicine, allow more accurate doses and prolong the effect of the active ingredient.

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 stoichiometry calculations.
The mole is the amount of substance in a system which contains as many elementary entities as there are in 0.012 kilogram of carbon
The elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of particles.
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 equations 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 × 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), NA = 6.022 × 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.
A proposed new definition of amount of substance, mole:
A mole is such that the Avogadro constant is exactly 6.0 221 415 × 1023 per mole.
A mole contains 6.023 × 1023 single units, e.g. atoms, molecules, electrons.

5.1.01 Atomic mass, atomic weight
The atomic mass of an atom is arbitrarily defined relative to the mass of the isotope carbon-12.
The relative atomic mass of an element is the ratio of the average mass of the element to 1/2 of the mass of one atom of the isotope
However, the "atomic mass unit", abbreviated "amu", is an archaic unit.
Atomic mass is also called atomic weight.
Tables of standard atomic weights published by the International Union of Pure and Applied Chemistry (IUPAC), apply to materials
used in the laboratory.

5.1.02 Molecular mass, molecular weight
Molecular mass, formerly molecular weight, is the mass of one mole of that material.
The molecular mass (m), is the mass of a molecule, kg.
Molecular (molecular weight), is the mass of one molecule of a substance and is expressed in the unified atomic mass units (u).
(1 u is equal to 1/12 the mass of one atom of carbon-12)
The relative molecular mass for an element or compound is the ratio of the average mass of molecules of the substance to 1/12 the
mass of one atom of C-12.
Molecular mass is also called molecular weight, MW.

5.1.03 Molar mass, MM
Molar mass, MM (molar weight), is the mass of one mole of a substance and is expressed in g/ mol.
The molar mass of a substance will contain 6.02 x 1023 molecules, Avogadro's number.
Molar mass was formerly called gram-molecular weight. Avogadro's hypothesis
Avogadro's hypothesis (Avogadro's law, 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 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
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).
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.
The dissolving process of a gas in a liquid is exothermic
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.
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.

1. Prepare a molar solution of magnesium sulfate.
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 × 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).

2. Prepare 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 × 24.3 = 24.3 g
Moles of atoms of Cl = 2 × 35.4 = 70.8 g
Moles of atoms of H = 12 × 1.0 = 12.0 g
Moles of atoms of O = 6 × 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 × 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 × moles
1000 mL H2SO4 contains 0.00125 × 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.
1. 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 = 0oC (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 × V1 / T1 = P2 × 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 27oC and 600 mm pressure?
P1 × V1 / T1 = P2 × V2 / T2
600 × 320 / 300 = 760 × V2 / 273
Volume at STP, V2 = 273 × 600 × 320 / 760 × 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 × 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, ρρ = 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, calculation
| Prepare
| Prepare solutions of known concentration (Table)
1. 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, Table shows that 80 g sodium hydroxide is needed to make 1 litre of 2 M solution.
So dissolve (4 / 2)M × 80 g = 160 g sodium hydroxide in 1 litre of water to make a 4 M solution.

2. 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, 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 × 15 g = 1.5 g sodium iodide in 1 litre of water to make a 0.1 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
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, ci = amount of constituent, ni / volume of the solution, V.

5.1.8 Molality
The molality of a solution is the number of moles of solute per kg of solvent.
1 molal = moles per Kg water.
Molality is the concentration of a solution expressed as number of moles of the dissolved substance per kilogram of solvent.
SI unit for molality = mol / kg.

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 × 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

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
Normal saline, physiological saline, 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

5.1.11 Percentage solutions, calculations
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
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
To make a 10% (W / V), NaCl solution, dissolve 10 g of NaCl in 100 mL of solution, 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
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
Specific gravity, SG (relative density, RD) concentrations
See: Specific gravity (Commercial)
See: Specific gravity, Multiple reagent strips
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
For example:.
38% HCl means that 38% of 1188.5 is HCl
1188.5 × 0.38 = 451.63 g / L = 45.163 g / 100 mL
A 38% solution of HCl at 20oC has specific gravity 1.1885 and contains 451.63 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 0oC and 1 atmosphere is 22.4 litres.
The molar volume of most gases at 25oC 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.
(The relative atomic mass (atomic weight, standard atomic weight), is the ratio of the average mass of one atom of an element to one
twelfth of the mass of an atom of carbon-12).

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 × 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. 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
7.0 Prepare acids and bases
Use 2 M acids from chemical suppliers.
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.
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
7.0 Prepare 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 Prepare alkalis
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.
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 Prepare ammonia solution
Ammonia solution (10% 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 lime water
Ionization of calcium hydroxide
See diagram 9.154: Lime water test for carbon dioxide in the breath
Add 200 g of calcium hydroxide to 2.5 litres of water.
Calcium hydroxide has a low solubility in water so shake it vigorously at intervals for some time and finally allow it to settle.
Decant off the clear lime water solution from above the undissolved solids.
Top up the vessel with deionized water and repeat the process.
Ionization of calcium hydroxide, Kb = 3.5 × 10-2
Ca(OH)2 <--> CaOH+ + OH-
CaOH+ <--> Ca2+ + OH-.

1. Lime water (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 lime water above the sediment
without disturbing the sediment.
To replenish the lime water, add more deionized water to the sediment in the stock bottle, shake and leave to settle.
The settling process to produce clear lime water may take several days.
The clear lime water prepared as above may be diluted 1:2 as needed.

2. Prepare lime water.
Slaked lime dissolves to a slight extent in water, and the solution is called lime water.
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 lime water.
Test it with a piece of red litmus paper.
The paper will turn blue, showing that lime water 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 lime water 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
The clear lime water can be poured off into a smaller bottle as is needed, the large bottle being replenished with water each time.

5.3.6 Prepare standard sodium carbonate solutions
Dispose of the wastes into labelled waste bottles for different kinds of chemicals.
1. Record the weight of an empty weighing bottle.
Use a spatula to add about 6.63 g of anhydrous sodium carbonate powder to the empty weighing bottle.
Record the weight of the weighing bottle + anhydrous sodium carbonate powder
If weight > 6.63 g, add more anhydrous sodium carbonate powder until weight is just more than 6.63 g.
1.1 Record mass of sodium carbonate powder =. g
2. Tip out sodium carbonate powder into a beaker.
Record weight of the weighing bottle (+ remaining sodium carbonate powder).
2.1 Record mass of any remaining sodium carbonate powder in weighing bottle = g.
Do not try to remove the last traces of sodium carbonate clinging to the weighing bottle.
Mass of sodium carbonate in beaker = g (1.1 - 2.1).

3. Prepare 250 mL of sodium carbonate solution from the powder in the beaker.
Add about 100 mL of distilled water or deionized water to the beaker.
Stir with a glass rod until all the powder has dissolved
Pour the solution through a filter funnel into a 250 volumetric flask.
Rinse the beaker with distilled water from a plastic wash bottle.
Wash the glass rod and inside of the funnel.
Repeat steps the above procedure.
Add more distilled water from the plastic wash bottle to the volumetric flask, to about 2 cm below the graduation mark.
Add more distilled water from a dropper until the lower meniscus of the solution just rests on the graduation mark.
Put the stopper in the mouth of the flask.
Press the stopper into position and invert the flask several times to mix the contents well.

4. Calculate the molarity of the standard sodium carbonate solution prepared.
Mass of Na2CO3 used = x g
See Table 2: Molar mass of Na2CO3 = 106 g mol-1 (23.00 x 2 + 12.00 + 16.00 x 3).
Volume of sodium carbonate solution prepared = 250 mL.
Molarity of the standard sodium carbonate solution prepared = number of moles of Na2CO3 / volume of sodium carbonate solution.
= (x / 106) / 250 = M.

5. Prepare a standard sodium carbonate solution by dilution
5.1 Prepare a washed pipette.
Use a pipette filler to suck water into a pipette to fill part of the bulb.
Hold the pipette horizontally and rotate it slowly so that water washes the inner wall up to the graduation mark.
Let the washing water to run out into the sink.
Repeat the above procedure to prepare a washed pipette using the prepared sodium carbonate solution instead of water.
5.2 Use the pipette filler and the washed pipette to transfer 25.0 mL of the prepared sodium carbonate solution into a clean 250 mL
volumetric flask.
Add distilled water from a plastic wash bottle into the volumetric flask, to a level about 2 cm below the graduation mark.
Add distilled water from a dropper until the lower meniscus of the solution just rests on the graduation mark.
Put the stopper in the volumetric flask.
Press the stopper into position then invert the volumetric flask several times to mix the contents well.
5.3 25.0 mL of the M sodium carbonate solution is diluted to a volume of 250 mL.
So, the molarity of the diluted sodium carbonate solution = ( X ) M =.M
5.4 Attach a label to the volumetric flask, e.g. Name, Class, Na2CO3 (aq), Molarity, Date, and store the volumetric flask in a safe place.

6.0 Find the molarity of a dilute hydrochloric acid solution using the prepared standard sodium carbonate solution
6.1 Repeat 5.1
6.2 Repeat 5.2
6.3 Repeat 5.1 to wash a burette with 20 mL of water then with the dilute hydrochloric acid solution.

5.3.10 Tests for heat, heat-sensitive paper
Add cobalt (II) chloride solution to ammonium chloride solution (sal ammoniac).
Dilute this solution until it is pale pink.
Soak absorbent paper in the solution and leave to dry and become almost colourless.
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 form provided by chemical suppliers.
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, lime water, react with the carbon dioxide of the air, so the solutions should be kept in sealed containers.

5.5.0 Alloys
1. Alloys (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
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.
Make a collection of alloys found in the home, school laboratory and workshop.

5.5.1 Eutectic mixture
1. 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.

2. Eutectic temperature and percentage composition by weight
183oC: Sn 63.0% Pb 37.0%
198oC: Sn 91.0% Zn 9.0%
221oC: Sn 96.5% Ag 3.5%
227oC: Sn 99.2% Cu 0.8%.

3. When the temperature of a mixture of 63.0% tin and 37.0% lead cools to 183oC, the liquid mixture freezes sharply.
So above 183oC the mixture is liquid and below 183oC 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 250oC, the liquid mixture starts to solidify
but it does not freeze sharply.
It finally solidifies at 183oC, giving the plumber time to ensure a good setting of a joint.

5.5.2 Amalgams
See 3.0.0: 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.

5.5.3 Copper, brass and bronze alloys
Formerly, brass was any alloy of copper with tin or zinc and even other metals, but nowadays brass is only the yellow alloy of copper
and zinc.
Bronze is a brown alloy of copper with up to 30% tin and other metals.
The hardness and workability of brass make it a valuable manufacturing material.
Alpha brasses have < 40% zinc, are malleable and can be worked cold.
Beta brasses have > 40% zinc, must be worked hot, but are harder and stronger.
Since the Bronze Age, brass was made by melting copper with calamine, basic zinc carbonate, ZnCO3.2ZnO.3H2O.
In the King James version of the Bible, 2 Chronicles, Chapter 4, Verse 16: 16: "The pots also, and the shovels, and the flesh hooks
and all their instruments, did Huram his father make to king Solomon for the house of the LORD of bright 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.

5.5.4 Lower melting point alloys
Use a Bunsen burner to find the melting points.
These alloys have a melting point lower than any of the constituent metals, e.g. Wood's metal, has melting point 80oC, but the lowest
melting point of the constituents, tin, has melting point 232oC.
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 155oC
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 200oC 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.
Lower melting point alloys with the parts shown are by weight:
Electrical fuse: 8.5 Pb, 2.5 Sn, 1.3 Bi, 0 Cd
Wood's metal: 4 Pb, 2 Sn, 7 Bi, 1 Cd
Lower melting alloys may be produced by using a Bunsen burner.
Where both bismuth and lead occur together in an alloy, the bismuth and lead are melted together, and then the other ingredients added.
The temperature should not be higher than necessary to prevent excess oxidation.
The higher melting alloys. e.g. bronze and brass, are produced in a furnace with the copper melted first and the other metals added.

5.5.5 Higher melting point alloys and parts by weight
Use a burner for solder (melting point 250oC), 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.

5.5.6 Copper-zinc alloys, brass
1. 5% Zn alloy, gilding metal, used for cheap jewellery, low value coins
2. 10% Zn, brass as sold by chemical suppliers
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
Copper- tin alloys, tin bronzes, have corrosion resistance.
Tin bronzes are stronger and more ductile than red and semi-red brasses.
They have high wear resistance and low friction coefficient against steel.
The room temperature phase transformations are slow and usually do not occur, therefore these alloys are single phase alloys.
The tin bronzes are used in bearings, gears, piston rings, valves, fittings, coins, bells.
1. 10% Sn, carrots bronze
2. P added to give stronger phosphor bronze.

5.5.8 Copper-aluminium alloys, bronze
The aluminum-copper alloys usually contain 2 to 10% copper.
The copper provides substantial increases in strength and facilitates precipitation hardening.
Copper can also reduce ductility and corrosion resistance of aluminium.
However, these alloys can be the most difficult aluminium alloys to weld because of their susceptibility to solidification cracking.
These alloys are some of the highest strength heat treatable aluminium alloys.
The most common applications are aerospace, military vehicles and rocket fins.
The 10% aluminium alloy, 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
The addition of nickel to copper improves strength and corrosion resistance but good ductility is retained.
Copper-nickel alloys have excellent resistance to marine corrosion and biofouling.
The two main alloys are 90/10 and 70/30.
The 70/30 is stronger and has greater resistance to sea water flow, but the 90/10 will provide good service for most applications and
being less expensive tends to be more widely used.
Both alloys contain small but important additions of iron and manganese that have been chosen to provide the best combination of
resistance to flowing sea water and to overall corrosion.
Copper-nickel alloys are widely used for marine applications due to their excellent resistance to seawater corrosion, high inherent
resistance to biofouling and good fabricability.
They have provided reliable service for several decades whilst offering effective solutions to today's technological challenges.
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 plated middle value coins.
The United States "nickel" (0.05 US dollars), is 75% copper, 2% nickel.

5.5.10 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.

5.5.11 Alloys of noble metals and coinage metals
Carat: 3.2.2
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% gold, 4.17% silver and 4.17% copper
Sterling silver is an alloy of 92.5% silver and 7.5% copper.
Diamonds are not alloys but pure carbon.
However, but the term "carat" is used, to mean 0.2 grams.

5.5.12 Cast iron alloy, steel, wrought iron
Cast iron contains about 4% carbon.
When iron ore is heated in a carbon fire two reactions occur.
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 220oC and 450oC 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
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, 50oC
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.

5.5.14 Copper-chromium alloys
Chromium copper alloys containing 0.6 to 1.2% Cr are used for their high strength, corrosion resistance and electrical conductivity.
They are "age-hardenable", i.e. a change in properties occurs at high temperature due to the precipitation of chromium out of the solid
The strength of fully-aged chromium copper is nearly twice that of pure copper and its conductivity remains high at 85% IACS
i.e. 85% of pure copper.
These high strength alloys retain their strength at elevated temperatures.
The corrosion resistance of chromium copper alloys is better than that of pure copper, because chromium improves the chemical
properties of the protective oxide film.
A time capsule made of specially-created "cupaloy" alloy made of 99.4% copper, 0.5% chromium, and 0.1% silver to last for 5000

5.5.15 Nitinol memory wire
Nitinol Memory Wire, solid state chemistry, Ni Ti alloys
This crystalline metal changes phase at about 50oC.
Bend it, then drop it into hot water and it returns to its original shape.
The wire can be set into different shapes by heating with a candle flame.