School Science Lessons
Topic 7 Chemical changes and physical changes, elements and compounds
2012-01-27 SP
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites
Table of contents
7.1.0 Chemical changes and physical changes
7.9.0 Chemistry terminology
7.2.1 Classify substances, pure substances, mixtures, solutions
7.8.0 Colloids and crystalloids
19.2.0 Composition of food
7.2.2 Elements and compounds
5.5.4 Low melting point alloys
7.3.0 Metals and non-metals
7.4.0 Melting points of solids
3.0.0 Particles
9.0 Prefixes and suffixes
6.3.5 Prefixes, SI prefixes
7.2.0 Pure substances and impure substances, elements and compounds
7.2.3 Silicon compounds
3.9.0 Solubility
7.7.0 Solutions
7.6.0 Suspensions and precipitates
7.1.0 Chemical changes and physical changes
7.1.0 Chemical changes and physical changes
7.1.1 Chemical changes, burn magnesium
7.1.1.1 Chemical changes, burn steel wool, combustion of iron wool
7.1.2 Chemical changes, heat organic substances
7.1.3 Chemical changes, heat metals in chlorine
7.1.4 Physical changes, magnetize and demagnetize iron wire
7.1.5 Physical changes, prepare forms of sulfur, allotropes of sulfur
3.70 Reactions between two elements, iron with sulfur
7.2.2 Elements and compounds
7.2.2 Elements and compounds, descriptions of common elements
3.0 Atoms
6.0 Chemical bonds
7.0 Compounds
8.0 Direct union of elements to form compounds
2.0 Elements
7.2.2a Elements experiments
5.0 Ions
1.0 Matter
4.0 Molecules
7.2.2 Elements and compounds, descriptions of common elements
[The properties listed below are based on information from IUPAC, International Union of Pure and Applied Chemistry.]
Al, Aluminium properties: 7.2.2.1
Alums: 7.2.2.1a
Aluminium chemical reactions: 12.1
Sb, Antimony properties: 7.2.2.2
Antimony chemical reactions: 12.2.1
Ar, Argon properties: 7.2.2.3
As, Arsenic properties: 7.2.2.4
Arsenic chemical reactions: 12.2.2
Wood treated with copper chrome arsenate (CCA): 12.2.2.1
Ba, Barium properties: 7.2.2.5
Barium chemical reactions: 12.2.3
Bi, Bismuth properties: 7.2.2.6
Bismuth chemical reactions: 12.2.4
B, Boron properties: 7.2.2.7
Br, Bromine properties: 7.2.2.8
Bromine experiments: 12.19.9
Cd, Cadmium properties: 7.2.2.9
Cadmium chemical reactions: 12.3.1
Ca, Calcium properties: 7.2.2.10
Calcium chemical reactions: 12.4.1
C, Carbon properties: 7.2.2.11
C, Charcoal blocks: 7.2.2.11a
Carbon chemical reactions: 16.1.0
Cl, Chlorine properties: 7.2.2.12
Chlorine experiments: 12.19.8
Cr, Chromium properties: 7.2.2.13
Chromium chemical reactions: 12.5
Co, Cobalt properties: 7.2.2.14
Cobalt chemical reactions: 12.6.0
Cu, Copper properties: 7.2.2.15
Copper chemical reactions: 12.7
F, Fluorine properties: 7.2.2.16
Fluorine experiments: 12.19.7
Au, Gold properties: 7.2.2.17
He, Helium properties: 7.2.2.18
H, Hydrogen gas properties: 7.2.2.19
I, Iodine properties: 7.2.2.20
Iodine experiments: 12.19.6
Fe, Iron properties: 7.2.2.21
Iron chemical reactions: 12.8
Kr, Krypton properties: 7.2.2.22
Pb, Lead properties: 7.2.2.23
Lead chemical reactions: 12.9
Lead paint: 7.2.2.23a
Tetraethyl lead: 7.2.2.23b
Li, Lithium properties: 7.2.2.24
Lithium chemical reactions: 12.9A.1
Mg, Magnesium properties: 7.2.2.25
Magnesium chemical reactions: 12.10.1
Mn, Manganese properties: 7.2.2.26
Manganese chemical reactions: 12.17
Hg, Mercury properties: 7.2.2.27
Mercury, Reactions of mercury compounds: 12.08b.1
Mercury and dental amalgam
Mo, Molybdenum properties: 7.2.2.28
Ne, Neon properties: 7.2.2.29
Ni, Nickel properties: 7.2.2.30
Nickel, Reactions of nickel compounds: 12.09.1
N, Nitrogen properties: 7.2.2.31
Nitrogen experiments: 12.11
Reactions of the nitrites, NO2-: 12.11.1
Reactions of the nitrates, NO3-: 12.11.2
Ammonia and the ammonium ion, NH3,NH4+: 12.11.3
O, Oxygen gas properties: 7.2.2.32
Oxygen experiments: 12.12
Ozone, O3, Prepare ozone: 3.50
P, Phosphorus properties: 7.2.2.33
Red phosphorus: 7.2.2.33a
White phosphorus: 7.2.2.33b
Phosphorus chemical reactions: 12.13
Phosphorescence, luminescence, chemiluminescence, fluorescence: 7.2.2.33.1
Detergent phosphates: 7.2.2.33.2
Pt, Platinum properties: 7.2.2.34
K, Potassium properties: 7.2.2.35
Potassium chemical reactions: 12.14
Ra, Radium properties: 7.2.2.36
Rn, Radon properties: 7.2.2.37
Se, Selenium properties: 7.2.2.38
Si, Silicon properties: 7.2.2.39
Silicon chemical reactions: 12.15
Silicon solar panel: 7.2.2.39.1
Ag, Silver properties: 7.2.2.40
Silver chemical reactions: 12.16
Na, Sodium properties: 7.2.2.41
Sodium chemical reactions: 3.73
Sr, Strontium properties: 7.2.2.42
Strontium, Reactions of strontium compounds: 12.17.1
Strontium chemical reactions: 12.16a.1
S, Sulfur properties: 7.2.2.43
Sulfur chemical reactions: 12.18
Sn, Tin properties: 7.2.2.44
Tin chemical reactions: 12.20
W, Tungsten properties: 7.2.2.45
U, Uranium properties: 7.2.2.46
Xe, Xenon properties: 7.2.2.47
Zn, Zinc properties: 7.2.2.48
Zinc, Reactions of zinc and its compounds: 12.21.1
Zr, Zirconium properties: 7.2.2.49
7.4.0 Melting point, m. p., of solids
7.4.0 Melting point, m. p., of solids
3.2 Melting point of naphthalene
3.3 Melting point of naphthalene with a capillary tube
3.4 Impurities affect the melting point of a substance
24.1.2 Lift an ice cube with salt
2.9 Melt different substances (Primary)
7.4.1 Melting point and cooling curve of stearic acid
7.4.2.0 Melting point of substances
7.4.2.1 Melting point of 1,4-dichlorobenzene, C6H4Cl2
7.4.3.0 Melting point of ice and freezing point, f.p., of water, antifreeze
7.4.3.1 Temperature at which ice melts
7.4.3.2 Temperature at which ice and salt mixture freezes
7.3.0 Metals and non-metals
7.3.0 Metals and non-metals
7.3.1 Properties of metals
7.3.2 Properties of non-metals
2.43 Metals (Primary)
12.11.3.2 Tests for metals with flame tests, metals and their compounds
12.11.3.1 Tests for metals with borax beads, metals in metallic salts and minerals
7.1.5 Physical changes, prepare forms of sulfur
7.1.5 Physical changes, prepare forms of sulfur, allotropes of sulfur
7.1.5.1 Prepare monoclinic sulfur from powdered sulfur (flowers of sulfur)
7.1.5.2 Prepare monoclinic crystals from roll sulfur
7.1.5.3 Prepare plastic sulfur then rhombic sulfur
12.18.1 Prepare forms of sulfur
7.1.0 Chemical changes and physical changes
1. In a chemical change, one or more substances changes into one or more new substances, e.g. hydrogen gas combines with oxygen gas to form water. In this document the expression used is as follows: "hydrogen gas with oxygen gas forms water". The new substances, the products, have properties different from the original substances, the reactants. For example burning wood forms black ash containing carbon, white ash containing mineral salts, carbon dioxide gas and water vapour. New substances form and the change cannot be reversed. The arrow symbol " -->" represents this type of change.
2H2 (g) + O2 (g) --> 2H2O (l)
2. In a physical change, the properties: of a substance change, but the substance is still the same. The change can be reversed. The physical properties: of water change when water freezes to ice. However, the ice is composed of water molecules and the change can be reversed. When electricity passes through a tungsten filament in a light bulb, the filament becomes hot and emits light. It is still tungsten. When you turn off the light, the tungsten filament is the same as before. Physical or chemical changes may be fast or slow, e.g. hammering iron, wearing away rock by wind erosion, an explosion when hydrogen gas burns, rusting of iron.
Classify common examples of changes into chemical change or physical change.
Do the following activities and in each case say whether it is a chemical or physical change: 1. Light a match. (chemical change) 2. Turn on the light. (physical change)
3. Let a nail rust. (a chemical change takes place in air causing a gain in weight) 4. Chew food. (physical change for biting and masticating, and a chemical change for the reactions of starch with salivary amylase)
7.1.1 Chemical changes, burn magnesium
Put a piece of magnesium ribbon in a crucible. Weigh the crucible and lid and magnesium ribbon. Burn the magnesium in the nearly closed crucible so that the ash is not lost. Weigh the crucible and lid and ash. The weight of the ash is greater than the weight of the magnesium. The reaction forms a new substance, magnesium oxide. Mg, density 1.74 g / cm3 oxidizes to MgO, density 3.58 g / cm3, so when oxygen attaches to magnesium the volume decreases because of the strength of the (Mg2+-- O2-) bond.
2Mg (s) + O2 (g) --> 2MgO (s)
7.1.1.1 Chemical changes, burn steel wool, combustion of iron wool
Cover the end of a wooden ruler with aluminium foil then note where the ruler balances as a first order lever, see-saw, over a fulcrum. Put the ruler balance on a sink bench. Put 5 g loose steel wool on the aluminium foil end of the ruler then put a weight on the other end of the ruler so that this end is just down. Heat the steel wool with a Bunsen burner flame. The steel wool glows and its side of the balance moves down as the iron becomes iron oxide.
7.1.2 Chemical changes, heat organic substances
1. Often one product is black because carbon forms. Heat the following substances in a hard glass or Pyrex test-tube. Observe any decomposition of the substances. When smoke appears, leave the test-tube to cool. No chemical change occurs when you heat charcoal. For the following substances, a chemical change occurs and the residue in the test-tube is carbon:
1.1 Wood or saw dust, note the brown substances that stick to the sides of the test-tube and the black substance that remain in the bottom, 1.2 sugar, 1.3 bread, 1.4 fruit, 1.5 fat or oil, 1.6 starch or potato or rice, 1.7 paper, 1.8 wool or hair or nail clippings, 1.9 meat.
2. Put a small piece of wood in a Pyrex test-tube. Heat the tube gently then more strongly. The wood decomposes into a black solid and brown liquids and vapours. Stop heating when you see no more dense smoke. Let the test-tube and observe the substances on the sides of the tube. Put any remaining solid substance in another clean test-tube. Heat it gently and then more strongly. You produce no further liquids or gases from this substance. Compare the black substance with a sample of the original wood for colour and can be bent. Place the end of each in a Bunsen flame for a few seconds. Remove and observe them. The black substance does not decompose or burn with a flame but glows red-hot. The black substance is charcoal, i.e. carbon. Heat small quantities of sugar, paper, wool, meat and compare the residue in the test-tubes with the original substance.
7.1.3 Chemical changes, heat metals in chlorine
Be Careful! Chlorine is poisonous. These experiment may not be allowed on your school system. D the experiments in a fume cupboard.
1. Use tongs to heat steel wool in a Bunsen burner flame, then place it in a test-tube containing chlorine gas. Observe whether any heat is given out during the reaction. Iron and chlorine react together to form a brown yellow iron chloride.
2. Using a wire gauze, scrape some powdered antimony from a lump of the element so that it falls into a test-tube of chlorine. The sparks and burning show that a chemical change is taking place. The new white substance is antimony chloride.
7.1.4 Physical changes, magnetize and demagnetize iron wire
When physical changes occur, no new substance forms.
Pull a thick iron wire (fencing wire) through iron filings. The iron does not attract the iron filings. Magnetize the iron by stroking it with a bar magnet. Hold the iron wire in the iron filings. The iron wire now attracts the iron filings. A physical change has occurred. Hammer the iron wire. The iron wire cannot attract the iron filings so strongly. The physical change has been reversed.
7.1.5 Physical changes, prepare forms of sulfur, allotropes of sulfur
Sulfur is a non-metallic element that occurs in several allotropic forms. Allotropes are variations of the same element with bonding and crystal structure. Sulfur occurs in three forms: Form 1: Rhombic sulfur is a light yellow powder. At 100oC it changes to: Form 2: Monoclinic sulfur with a deeper colour. The monoclinic and rhombic forms differ in the arrangement of the S8 molecules. At 160oC, it melts to form a sticky dark brown liquid that can be cooled quickly to form red brown plastic sulfur. Form 3: Plastic sulfur contains long chains of S atoms.
Put sulfur powder in a test-tube. Heat the sulfur extremely gently until it slowly melts to a golden yellow liquid. Continue to heat more strongly until a red gas appears above the liquid. Leave the test-tube to cool. Sulfur forms deposits on the sides of the tube and in the bottom of the tube.
7.1.5.1 Prepare monoclinic sulfur from powdered sulfur (flowers of sulfur)
Monoclinic sulfur has a deep yellow colour, m.p. = 119oC, and density = 1.96.
Heat powdered sulfur extremely gently in an evaporating dish. The sulfur changes to liquid. Add more powdered sulfur. The colour should stay pale yellow. If the sulfur turns dark, you have overheated it. Repeat the experiment with gentler heating. Leave the sulfur to cool, without moving the evaporating dish. Monoclinic crystals form between 96oC and 114oC. After a thin crust forms, punch two holes through the crust with a nail. Pour out the hot sulfur through one hole. Remove the crust and note the monoclinic sulfur crystals on the underside.
7.1.5.2 Prepare monoclinic crystals from roll sulfur
Put small pieces of roll sulfur in a test-tube. Heat the test-tube extremely gently until the sulfur melts, m.p. = 114.5oC. The liquid is lemon yellow. Pour the liquid sulfur into a folded filter paper. A crust forms on the surface. When the crust forms, open the filter paper. Needle-shaped crystals remain on the filter paper.
7.1.5.3 Prepare plastic sulfur then rhombic sulfur
When heated to the melting point, sulfur usually ignites and forms sulfur dioxide gas that may distress people suffering from asthma.
Rhombic sulfur is a yellow powder, m.p. = 119oC, and density = 1.96.
Put roll sulfur in a test-tube. Heat the test-tube slowly. Note the changes at the melting point from a light yellow colour to a red liquid. The red brown sulfur becomes viscous and it does not flow out if you hold the test-tube upside down. Continue heating until the reaction forms a brown black liquid. Continue to heat until the sulfur boils at 445oC. BE CAREFUL! Pour the melted sulfur into a beaker of cold water. The reaction forms strands of amorphous sulfur. When the strands are cool, twist them to show that they are elastic. Later the sulfur hardens because it returns to the rhombic form of sulfur as a ring of eight sulfur atoms in tiny crystals. When heated, the ring breaks open to form long chains.
7.2.0 Pure substances and impure substances, elements and compounds
1. Pure substances contains only one kind of atom or molecule, e.g. iron, sulfur, water, and oxygen. It can consist of elements or compounds. Elements may be metals or non-metals. Impure substances contain more than one kind of substance. They may be mixtures, e.g. sulfur and iron filings, or air, or solutions, e.g. sea water. Mixtures in the home include flour, milk, ink, face powder, and tooth paste. Solutions in the home include fruit drinks (although often an emulsion) lemonade, mineral water. You make toffee or candy with a supersaturated sugar solution. French polish is shellac with methylated spirit solution.
2. An element cannot be broken down into simpler substances by a chemical reaction and all the atoms in it have the same atomic number, the number of protons and electrons is the same, but the number of neutrons may vary. Ninety two naturally occurring elements exist..
3. A compound is composed of two or more elements combined in fixed proportions as a result of a chemical reaction and which cannot be separated into simpler substances by any physical process, e.g. shaking. Most compounds are ionic, e.g. common salt (NaCl) or covalent molecular, e.g. CO2 or H2O, or covalent network, e.g. SiO2. Each compound has a name or formula. Compounds with names ending in "-ide" contains two elements. Compounds with names ending in "-ate" or "-ite" contain oxygen, with more oxygen in the "-ate". Compounds found at home include water, sugar, starch, baking soda, cream of tartar (potassium hydrogen tartrate).
3. Inorganic chemistry is the chemistry of the elements and their compounds, including CO2, CO and carbonates. Organic chemistry is the chemistry of carbon and its compounds.
7.2.1 Classify substances, pure substances, mixtures, solutions
1. Describe materials at home and in the classroom from their observable physical properties: 1. colour (shiny or dull) 2. opaque (or transparent or translucent) 3. shape (or shape of crystals) 4. odour, 5. state (solid, liquid, gas) (change in state when heated or cooled) 6. mass (heavy, light) more dense or less dense than water, 7. taste (sweet, sour, bitter, other) 8. can be magnetized, 9. can conduct heat, 10. can conduct electricity, 11. can absorb liquids, 12. flexibility (can be bent, fragile). Describe the materials used to make items in the classroom and in the home.
2. Describe materials as pure substance, or solution or mixture
If a solution is a homogeneous mixture of a liquid with a gas or a solid, then you would not classify air as a solution, nor brass as an alloy, unless they were in liquid form. However, if a solution is a homogeneous mixture of two or more components in a single phase, and usually refers to a solution in water (aqueous solution) then perhaps you can classify air and solid alloys as solutions.
Substances Pure substance Solution Mixture
1. Ice floats in water + -
 -
2. Tincture of iodine -
 + -
3. Washing powder -
 -
 +
4. Tap water -
 + -
5. Air -
 + or mixture +
6. Brass alloys -
 + or mixture +
1.0 Matter
Matter is composed of atoms which, in turn, contain protons and neutrons in a nucleus, and electrons outside the nucleus. The number of positively charged protons is equal to the number of negatively charged electrons in a neutral atom, and determines all the chemical properties of an atom. Materials may be elements, compounds or mixtures.
2.0 Elements
Elements cannot be broken down into simpler substances by a chemical change. An element is a substance in which all atoms have the same number of protons. Atoms of an element may contain different numbers of neutrons, and are known as isotopes. Every element is assigned a unique chemical symbol. At room temperature and atmospheric pressure, eleven elements are gases: H2, He, N2, O2, F2, Ne, Cl2, Ar, Kr, Xe, Rn. Two elements are liquids at room temperature: Hg and Br. Reactive elements have atoms weakly bound together and have electrons available for bonding, e.g. F. Unreactive elements have atoms joined by strong bonds, e.g. diamond. Elements that exist as separate small molecules have low boiling points and melting points, e.g. He, O2.
3.0 Atom
Atom is the smallest division of an element that can chemically exist alone and have the characteristics of the element. Atomic mass (atomic weight) of an atom is arbitrarily defined relative to the mass of the isotope carbon. The relative atomic mass of an element is the ratio of average mass of atoms of the element to 1 / 12 of the mass of one atom of the isotope C-12. Atomic mass unit, (a.m.u.) is a unit of mass used to express the relative atomic mass. So 1 atomic mass unit = 1 / 12 of the mass of the carbon-12 isotope, i.e. 1.66 X 10-27 kg. The mass of 1 atom of oxygen = 16 a.m.u.
The atomic theory dates from the 5th century B.C. when Greek philosophers, e.g. Democritus, said that matter consists of indivisible indestructible particles. The modern atomic theory started with the hypothetical thinking of John Dalton, England, (1766 - 1844). However, he did not envisage the structure of atoms, i.e. nucleus, electrons and other particles.
4.0 Molecule
Molecule contains two or more atoms joined chemically to form the simplest stable structure of that element or compound. So a hydrogen molecule must contain 2 atoms, H2, because the hydrogen atom, H, cannot exist by itself. A molecule of water, H2O, contains 2 atoms of hydrogen, H, and 1 atom of oxygen, O.
5.0 Ions
Ions form when an atom or group of atoms covalently bound together may gain or lose one or more electrons. Ionic bonding occurs when positive and negative ions are held together in a crystal lattice by electrostatic forces. The structure of metals involves positive ions embedded in a sea of electrons.
6.0 Chemical bonds
Chemical bonds are forces of attraction that the atoms together in a molecule or a crystal. Ionic, electrovalent, bonds form by transfer of electrons and covalent bonds form by sharing of electrons. Hydrogen bonds are weak intermolecular forces of attraction between polar molecules that contain hydrogen, i.e. between the H atom of one molecule and the negative charged atoms in another molecule. In water, there is weak attraction between the hydrogen in a water molecule with the oxygen (in the OH) of another water molecule. Pairs of atoms may be bound together by the sharing of electrons between them in a covalent bond. Two or more atoms bound together by one or more covalent bonds form a molecule, with definite size, shape and arrangement of bonds.
7.0 Compounds
Compounds are composed of two or more elements that are chemically united in fixed proportion. Compounds can be broken down to simple substances. Chemical compounds form when chemical bonds, whether ionic or covalent, are formed between different elements. A chemical compound can be represented by a chemical formula. Forces weaker than covalent bonding exist between molecules. In compounds containing carbon-hydrogen bonds, organic compounds, the carbon atoms bind to one another through single, double or triple covalent bonds to form chains or rings.
Observe samples of iron, carbon, copper, magnesium, mercury in a thermometer and solid sticks of sulfur.
Classify each of them according to the following characteristics: 1. Is it a hard solid, a liquid or a gas at room temperature? 2. If the element is a solid, does it shine and have a lustre? If necessary scratch its surface. 3. Can the element be bent or twisted or does it fracture easily? You may require pliers or a vice to do this. 4. Will the element conduct electricity? Place the element between two electrical contacts in a circuit.
8.0 Direct union of elements to form compounds
Salts
2Na + Cl2 --> 2NaCl
Fe + S --> FeS
12.2.2.1 Heat iron with sulfur
Acids
H2 + Cl2 --> 2HCl
H2 + S --> H2S
7.2.2.1 Aluminium, Al, is not safe for school use so it is not supplied as  powder. Aluminium reagent test kit, No. 1 tablets, No. 2 tablets, AAS Std, leaf, flakes, sheet, "Alfoil" cooking foil, drink-can, ingot, wire, aluminium foil, 300 width X 150 m roll, "Alfoil" disposable containers, rectangular 120 mm X 175 mm, round 120 mm diameter
1. Aluminium, Al, is a silver white metal, with mechanical strength and forms protective oxide layer in air that prevents further oxidation, available as foil, sheet, and wire. The aluminium foil used in home kitchens from grocery stores and the aluminium wire and aluminium sheeting from hardware stores are all pure aluminium. Aluminium is extracted from the ore bauxite (Al2O3) 8% of the earth's crust so most abundant element in Earth's crust, extracted by electrolysis of bauxite dissolved in cryolite, protected by layer of oxide that can be thickened if an anode in electrolytic cell, i.e. anodized then can be coloured, malleable, low density, used in "Alfoil" cooking foil, drink-cans, saucepans, cars, duralumin, alloy of Al + Cu, Mg, for aircraft bodies and struts.
2. Aluminium reacts with dilute HCl or H2SO4 to form H2 and metal ion, with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, with steam to give the oxide and hydrogen gas. Heated powder forms oxide. Excess aluminium may cause short term toxicity, e.g. aluminium sulfate in drinking water, corroded cooking utensils used for cooking acidic foods, alum treatment of water. In 1990, The International Union of Pure and Applied Chemistry (IUPAC) confirmed the use of the spelling "aluminium" but people in the U.S.A. are accustomed to use the spelling "aluminum".
3. Finely-divided aluminium particles (powder or fine turnings) burn in air with an intense white flame if ignited. A significant amount of ultraviolet radiation is emitted, and the flame should not be watched with the naked eye. Unoxidized aluminium powder reacts vigorously with both concentrated acids or alkalis to yield hydrogen gas, which is explosive when mixed with air. The reaction may occur with such vigour that the aluminium particles and acid or alkali may be ejected from the container. Most reactions of aluminium are inhibited by a surface coating of aluminium oxide. This is the reason that such an active metal as aluminium can be used widely for construction and for containers. Aluminium metal reacts violently with halogens (chlorine, bromine, iodine). Reaction of aluminium with liquid bromine is hazardous and should not be attempted.
Atomic number: 13, Relative atomic mass: 26.9815, r.d. 2.70, m.p.: 660oC, b.p.: 2470oC, Specific heat capacity: 900 J kg-1 K-1, E 173
7.2.2.1a Alums
The original "alum" was the hydrated double salt of aluminium and potassium with colourless octahedral crystals and an astringent taste. Later the term "alum" was used to describe similar double sulfates where other elements or radicals replaced aluminium or potassium. Some plants with astringent roots are called alum roots. So an alum is a hydrated aluminium potassium sulfate and related minerals. Aluminium ions form alums that are complex hydrated metal sulfates that contain 12 or 24 H2O, e.g. aluminium potassium sulfate Al2(SO4)3.K2(SO4).24H2O, alum, (potassium alum), used as mordants in dyeing.
Aluminium potassium sulfate is a buffer and firming agent, E522.
Alums:
Ammonium alum, AlNH4(SO4)2.12H2O
Chrome alum, chromium (III) potassium sulfate, potassium chromium sulfate, KCr(SO4)2.12H2O or K2SO4.Cr2(SO4)3.24H2O
Potassium alum, potash alum, hydrated potassium aluminium sulfate, "common alum", Al2(SO4)3.K2(SO4).24H2O, or AlK(SO4)2.12H2O or KAl(SO4)2.12H2O
7.2.2.2 Antimony, Sb, blue white metal, burns in air but no reaction with water or dilute acids, attached by halogens and oxidizing acids, poor conductor of heat and electricity, used in alloys for cable covers, pewter and lead cell accumulator plates, donor impurity in silicon chips, radioactive isotopes to produce neutrons, from stibnite Sb2S3.
Atomic number: 51, Relative atomic mass: 121.75, r.d. 6.68, m.p.: 630.5oC, b.p.: 1750oC, Specific heat capacity: 210 J kg-1 K-1
7.2.2.3 Argon, Ar, is a non-metal inert colourless odourless noble gas at room temperature and pressure Ar, 0.93% of the air, extracted from liquid air. Chemically inactive, no compounds, monatomic gaseous element, used in incandescent light bulbs, fluorescent tubes and lasers, and for welding. Most abundant noble gas, 0.9% of atmosphere by volume.
Atomic number: 18, Relative atomic mass: 39.948, r.d. 1.40 (87 K), m.p.: -189oC, b.p.: -186oC, Specific heat capacity: 519 J kg-1 K-1
7.2.2.4 Arsenic, As, is a metalloid (or non-metal) poisonous found free and in combined many minerals, three allotropes are yellow, black and main allotrope grey arsenic sublimes at 613oC, and n-type dopant of silicon semiconductors, hardens lead alloys. White arsenic, arsenic (III) oxide, As4O6, common in sulfide ore deposits, very toxic, rodenticide. Salvarsan, Erlich's compound 606 (arsphenamine) was the first drug to treat syphilis. The most toxic form of arsenic is As3+ which reacts with enzymes in the body. Agricultural use of arsenic kills plants before concentration is toxic enough for humans.
Atomic number: 33, Relative atomic mass: 74.9216, r.d. 5.72, m.p.: 814oC, b.p.: 613oC (sublimes), Specific heat capacity: 326 J kg-1 K-1
7.2.2.5 Barium, Ba, is an alkaline earth metal, brittle and expensive, used to absorb high energy particles, in minerals barytes, BaSO4 and witherite, BaCO3, forms poisonous compounds, oxidizes in air and reacts with ethanol. Barium sulfate is used for a contrast medium for X-ray examination of intestines. Surface coating of barium hydroxide corrosive to the eyes. Reaction of barium with water produces flammable hydrogen gas. Very difficult to cut.
Atomic number: 56, Relative atomic mass: 137.33, r.d. 3.51, m.p.: 725oC, b.p.: 1640oC, Specific heat capacity: 192 J kg-1 K-1
7.2.2.6 Bismuth, Bi, white metallic crystals, from bismuthinite mineral Bi2S3 and niccolite, cobaltite, but rare in the earth, very diamagnetic, low thermal conductivity, high electrical resistance, burns in air with blue flame, used in low melting point alloys for fire safety equipment, thermocouples, magnetic flux measurement, liquid metal coolant for nuclear reactors, cosmetics and medicines, e.g. bismuth carbonate for peptic ulcers.
Atomic number: 83, Relative atomic mass: 208.98, r.d. 9.78, m.p.: 271.3oC, b.p.: 1560oC, Specific heat capacity: 123 J kg-1 K-1
7.2.2.7 Boron, B, is a non-metal or metalloid, yellow brown network solid, brown amorphous form and black metallic form, has metallic lustre, very hard (9.3 Mohs' scale) and strong semiconductor, found in minerals, e.g. tourmaline, and associated with volcanic activity as borates, used in control rods for nuclear reactors and in green flares.
Atomic number: 5, Relative atomic mass: 10.81, r.d. 2.34 (amorphous form), m.p.: 2300oC, b.p.: 2550oC, Specific heat capacity: 123 X 103 J kg-1 K-1
7.2.2.8 Bromine, Br, (Greek: bromos, stench), is a red-brown, fuming, volatile, poisonous, non-metal liquid between 19oC and 27oC, suffocation odour, vapour irritates eyes and throat, strong oxidizing agent, used for many chemical compounds including "anti-knock" petrol additive, 1,2-dibromoethane, scarce element extracted from sea water, 65 ppm, as bromide ion. Silver bromide import for photography. Bromoform, tribromomethane, CHBr3 used to separate minerals, bromothymol blue indicator, pH 6.0 to 7.6, potassium bromide formerly used as sedative and was supposed top be put in army tea to quieten soldiers' sexual urges, bromochlorodifluoromethane, CHBrClF2, low toxic fire extinguisher for confined spaces. The fat soluble fire retardant PBDE, polybromyldiphenyl ether, in the deca, octa and penta forms has been detected in mothers' milk, fish and the environment. (Poison COR 1744) Br2 (3.6% bromine) r.d. 3.12 gm cm-3, b.p. 58.7oC, solidifies -7oC (swimming pool sanitation, products from bromine, e.g. BCDMH)
Bromine is a dense red-brown liquid with a powerfully irritant vapour.
Bromine water, a solution of bromine in water, is usually available commercially. Handle pure liquid bromine in s mall quantities in a fume cupboard. The liquid is unexpectedly dense, so increasing the chance of containers being dropped by inexperienced people. Breakage of a bottle of bromine outside a fume cupboard will require evacuation of the area until the vapour dissipates. Bromine reacts violently with active metals such as aluminium / magnesium and sodium. Do not allow active metals to contact liquid bromine. Always store the bromine in a cool secure store area.
Atomic number: 35, Relative atomic mass: 79.904, r.d. 3.12, m.p.: -7.2oC, b.p.: 58.78oC, Specific heat capacity: 448 J kg-1 K-1
7.2.2.9 Cadmium, Cd is a soft blue white metal, toxic at low concentrations, a rare element that occurs in the mineral sphalerite, zinc sulfide, is used for cadmium plating against corrosion, in nuclear reactors and in films sensitive to ultraviolet light, reference voltage in a Western Standard cadmium cell, Cd/ Zn alloys in low melting point solders and aluminium solders, Ni-Cd batteries (Nicad) phosphorescent coating of TV tubes. Cadmium looks like zinc but make a crackling sound like tin when bent. Oysters in polluted water make accumulate cadmium. In former Communist countries that practised collection of "night soil" as a fertilizer, cadmium pollution of low lying agricultural soil has occurred. Cadmium is toxic because it competes with Zn and Ca. Large concentrations cause painful bone ailments. and bone porosity. Our body gets rid of excess Cd by deposition in the kidneys and liver.
Atomic number: 48, Relative atomic mass: 112.41, r.d. 8.64, m.p.: 321oC, b.p.: 765oC, Specific heat capacity: 230 J kg-1 K-1
7.2.2.10 Calcium, Ca, is an alkaline earth metal, granules in liquid paraffin. Reacts with dilute HCl or H2SO4 to form H2 and metal ion, occurs mainly as carbonates, e.g. calcium carbonate, CaCO3, gypsum, CaSO4.2H2O, 3.5% of the earth's crust, essential nutrient element for bones, teeth and muscle contraction in animals and middle lamella of plan cells, extracted by electrolysis of fused calcium chloride. Reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide, SO2, or nitrogen dioxide, NO2. Reacts with cold water and reacts with air to form peroxides. Calcium is the most abundant mineral, and the fifth most abundant element mostly in bone tissue. About 1% is used in nerve transmission, muscle contraction and other functions
Atomic number: 20, Relative atomic mass: 40.08, r.d. 1.54 g cm-3| m.p.: 850oC, b.p.: 1487oC, Specific heat capacity: 653 J kg-1 K-1
7.2.2.11 Carbon, C, charcoal, lump, powder, lamp black, acetylene black, wood charcoal, activated carbon / charcoal (coke is left when coal is heated without air, for blowpipe work, charcoal may contain wood ash mainly potassium carbonate, soft "lead" pencils) non-metal network solid, has two natural crystalline forms, diamond and graphite. Diamond has strong covalent bonds in three dimensions to four other atoms, is the hardest substance known, transparent, brittle and non-conductor. It is used as gemstones, diamond dust abrasives and rock boring tools. Graphite has covalent bonds that are strong in one dimension but weak between layers, has density 3.5, and is soft and black. slippery with lustre and is good conductor only along its layers, available as graphite mineral, graphite powder and colloidal graphite "Aquadag", occurs as mineral deposits or is made from petroleum, used in the "lead" of soft "B" pencils, as a lubricant and electrical conductor, in cast iron and as coke for heating oildag, plumbago, "black lead", black-lead, graphite.
Atomic number: 6, Relative atomic mass: 12.011, r.d. 2.25 (graphite), m.p.: 3730oC (sublimes), b.p.: 4830oC, Specific heat capacity 711 J kg-1 K-1 (graphite), 519 J kg-1K-1 (diamond)
7.2.2.11a Charcoal blocks, C
Charcoal blocks are hazardous in one way only: they smoulder and burn for a long time after ignition and may cause fires if put away before they have been properly extinguished. They have caused fires in schools. If charcoal blocks are hastily collected and put in a drawer or cupboard after a laboratory class, they may smoulder for hours or days before either extinguishing themselves or flaring up. At the end of a laboratory class, leave the blocks totally immersed in a bucket of water. An alternative to using charcoal blocks is to mix equal volumes of metal oxide and charcoal powder in a test-tube and heat with a Bunsen burner while the tube is held horizontally. Charcoal can be produced by charring bread with a Bunsen burner in a fume cupboard.
7.2.2.12 Chlorine, Cl, is a non-metal green yellow, poisonous with irritating smell gas at room temperature and pressure Cl2, strong oxidizing agent and used in bleaching powder, disinfectant, 1.9% of sea water, produced from electrolysis of brine, used to make many organic products, e.g. chloroform trichloromethane CHCl3, chlorofluorocarbon refrigerant and aerosol now phased out because of damage to the ozone layer in atmosphere, chlorophenol red pH 4.8 to 6.4 indicator, chloral hydrate sedative, chloric (V) acid, HClO3 and its salts chlorates (V) powerful oxidizing agents and sometimes explosives. Chloride is a constituent of gastric juice which is about 0.03 M HCl. It is also used for controlling the transport of oxygen and carbon dioxide haemoglobin in red blood cells. Adults require a daily minimum, of 750 mg of chloride.
This gas is very toxic. It can react to cause fires or explosions upon contact with turpentine, ether, ammonia gas, illuminating gas, hydrocarbon, hydrogen gas and powdered metals. Dissolves readily in water forming highly corrosive solution. Do not prepare chlorine in open room. Use fume cupboard. Direct combination of chlorine and hydrogen gas in bright light or ignition of the mixture by lighted taper or electric spark. Reactions of chlorine with metals, solid non-metals, hydrocarbon. Use small quantities only.
Atomic number: 17, Relative atomic mass: 35.453, r.d. 1.56 (238 K), m.p.: -101oC, b.p.: -34.7oC, Specific heat capacity: 477 J kg-1 K-1
7.2.2.13 Chromium, Cr, is a reactive transition metal but forms protective oxide layer in air that prevents further oxidation and forms hard alloys with Ni or Fe. Chromium is available as technical grade chromium, is extracted from chromite (Fe(CrO2)2) and is used for chromium plated metal, hard plating Cr2O3, catalysts, in stainless steel and in heat resistant steel. Strong reducing agent Cr2+ salts blue in aqueous solution, Cr3+ salts green in aqueous solution. CrO42- salts yellow, e.g. potassium chromate, K2CrO4, and strong oxidizing agent Cr2O72-, orange, e.g. potassium dichromate K2Cr2O7. Chromium deficiency reduces tolerance to glucose.
Atomic number: 24, Relative atomic mass: 51.996, r.d. 7.19, m.p.: 1890oC, b.p.: 2482oC, Specific heat capacity: 448 J kg-1 K-1
7.2.2.14 Cobalt, Co, is a transition hard, grey metal, forms complex ions, e.g. [Co(H2O)6]2+, magnetic below 1075oC, essential element but toxic in excess, used in alloys radiography, magnets, steel. Cobalt (II) oxide used to colour glass blue. Occurs in the body only as cyanocobalamin, vitamin B12. Cobalt, (German: kobold, goblin of the mines) was associated by miners with with arsenic and sulfur health damaging impurities.
Atomic number: 27, Relative atomic mass: 58.9332, r.d. 8.90, m.p.: 1492oC, b.p.: 2900oC, Specific heat capacity: 435 J kg-1 K-1
7.2.2.15 Copper, Cu (cuprum) is a bright reddish orange, malleable and ductile transition metal, with high electrical and thermal conductivity, obtained form cuprite, Cu2S, available as filings, foil, powder, turnings, copper nails, copper wire, and copper turnings, is extracted from cuprite (Cu2S) and malachite (basic copper (II) carbonate, CuCO3.Cu(OH)2.H2O) and used for coin alloys, electrical wiring and heating vessels. Reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2. No reaction with dilute HCl or H2SO4. or with water. Heated powder forms oxide. Copper deficiency may occur in infants fed only on cow's milk. Copper bracelets my alleviate arthritis. A copper bowl may be preferable for beating cream. Copper poisoning may occur from water standing for a long time in copper pipes or in a copper hot water service. Copper is a cofactor for many enzymes and proteins, and is used in the development of nerve, bone, blood and connective tissue. Copper competes with zinc for entry from the intestines, so an increase in dietary zinc may result in copper deficiency. The recommended daily allowance (RDA) is 1.5 to 3.0 mg. Copper bracelets do not cure arthritis.
Atomic number: 29, Relative atomic mass: 63.546, r.d. 8.92, m.p.: 1083oC, b.p.: 2595oC, Specific heat capacity: 385 J kg-1 K-1
7.2.2.16 Fluorine, F is a non-metal yellow poisonous gas at room temperature and pressure F2, most reactive element known, never found as free element but in many silicate minerals and fluorite, CaF2, most electronegative non-metallic element, strong oxidizing agent, combines with carbon to form low friction fluorocarbon polymers, e.g. PYFE, Teflon, used to make CFCs, chlorofluorocarbons, freon, that damages the ozone layer. Fluorine compounds are added to toothpaste and drinking water, e.g. tin (II) fluoride, sodium monofluorophosphate (MFP) sodium fluoride and amine fluorides. In Australia solid sodium silicofluoride is added to drinking water in some places. However, some natural water sources already contain the fluoride ion. The enamel of teeth are formed from the crystalline mineral hydroxyapatite, Ca10(PO4)6(OH)2. The fluoride ion F- replaces the OH- to form in the teeth Ca10(PO4)6F2. Fluoride is essential for teeth and bones.
Atomic number: 9, Relative atomic mass: 18.9984, r.d. 1.11 (85 K), m.p.: -220oC, b.p.: -188oC, Specific heat capacity: 824 J kg-1 K-1
7.2.2.17 Gold, Au (aurum) is a yellow, lowest metal in the reactivity series so most unreactive transition metal, is found in nature as free element form, available as gold leaf, occurs as the element and is used in coin alloys, electrical components, dentistry alloys, jewellery and as a monetary standard. No reaction with dilute HCl or H2SO4, air, water or concentrated oxidizing acids, e.g. HNO3 or H2SO4. Reacts with aqua regia, "royal water" that dissolves gold the "king of metals" (1 part concentrated HNO3 + 3 parts con. HCl) to form AuCl4. Atomic number: 79, Relative atomic mass: 196.967, r.d. 19.3, m.p.: 1063oC, b.p.: 2970oC, Specific heat capacity: 130 J kg-1 K-1
7.2.2.18 Helium, He, is a non-metal noble gas (inert gas) at room temperature and pressure. Helium has no known compounds. Helium has separate small molecules, 0.0005% of the air, superfluid at 2.2 K, lowest boiling point, obtained from natural gas wells, using in diving gases, balloons (in party balloons but these balloons may travel over the oceans to fall and choke sea animals), welding. large concentration can asphyxiate, change of pitch of voice due to increased velocity of sound. Helium has the lowest critical temperature, -268oC. Packaging gas, propellant gas E939.
Atomic number: 2, Relative atomic mass: 4.00260, r.d. 0.147(4 K), m.p.: -270oC, b.p.: -269oC, Specific heat capacity: 5.19 X 103 J kg-1 K-1
7.2.2.19 Hydrogen, H, is a colourless odourless gas at room temperature and pressure H2, lightest element, burns to form water, most common element in space, natural isotopes hydrogen and deuterium and manufactured isotope radioactive tritium, product of electrolysis of water, used to fix nitrogen and make ammonia in Haber process, reduction of ore oxides, manufacture of HCl, hydrogenation of oils, elemental gas in balloons and possible potential use as hydrogen gas fuel in motor cars.
Atomic number: 1, Relative atomic mass: 1.0079, r.d. 0.070 (20 K), m.p.: -259oC, b.p.: -252oC, Specific heat capacity: 1.43 X 104 J kg-1 K-1
7.2.2.20 Iodine, I, I2, resublimed [COR 1759] is a non-metal forms violet black solid poisonous scales with special smell, the least reactive of the halogens, most common as iodides, insoluble in water but dissolves in ethanol and a solution containing I-, 1% in KI, because it forms I3-, when heated sublimes to form vapour that irritates the eyes, important for function of the thyroid gland, intense blue colour is test for starch, extracted from unpurified Chile saltpetre (caliche) and seaweed, powerful disinfectant when dissolves in ethanol to form tincture of iodine, as radio iodine isotope Iodine-123 used in nuclear medicine, especially thyroid gland disorders. Strong oxidizing agent and antiseptic. As povidone-iodine complex (PVD-iodine) in the antiseptic "Betadine". Iodine is used in the muscles and in the thyroid gland as part of the chemicals thyroxine and triiodothyronine. Thyroxine is a hormone containing iodine, from thyroid gland, controls metabolic rate. Iodine deficiency causes hyperthyroidism and an enlarged thyroid gland. The recommended daily allowance, RDA, is 150 pg. Radioactive labelled iodine used to measure percentage uptake of it by thyroid gland (RAI test). Iodine poses fewer risks in school laboratories than the other halogens, chlorine and bromine. Do not inhale the acrid iodine vapour. Avoid skin contact with solid iodine. Iodine, whether as a solid, in solution or a vapour can temporarily stain the skin. Do not allow iodine to react with concentrated ammonia solution because it forms nitrogen triiodide, NI3NH3, (touch powder), which is violently explosive.
Solid crystalline iodine is toxic by all routes of exposure and a highly irritating vapour comes off from the crystals and iodine solutions. Do not mix iodine with concentrated ammonia solution because the highly explosive nitrogen triiodide may form. The reactions of iodine with acetaldehyde or antimony metal are violently exothermic. Use of solid iodine by students should be limited to 0.2 g, (the size of about 2 rice grains), per experimental activity.
Atomic number: 53, Relative atomic mass: 126.905, r.d. 4.93, m.p.: 113.5oC, b.p.: 184oC, Specific heat capacity: 218 J kg-1 K-1
7.2.2.21 Iron, Fe (ferrum, ferrum reductum) as powder, iron nails, iron wire, iron filings, steel wool, is a magnetic and strong transition metal, 4.5% of the earth's crust, used for making iron and steel and is the most commonly used metal, available as filings, iron nails and wire, extracted from iron ores, e.g. haematite, Fe2O3, reacts with dilute HCl or H2SO4 to form H2 and metal ion, reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with steam to give the oxide and hydrogen gas. Heated powder forms oxide. The complex haemoglobin molecule has an iron atom in the centre. During rusting, metallic ion changes to Fe(OH)3.xH2O. Galvanized iron is Fe with Zn coating, e.g. "tin" roof. Iron is used in the haemoglobin protein that carries oxygen and carbon dioxide in the blood. Pig iron is cast iron, 2-4.3% carbon. Iron is a safe material. However, both iron filings and iron powder (ferrum reductum) are hazardous when mixed with either sulfur, chlorine or bromine, because of the highly exothermic reactions that can occur.
Atomic number: 26, Relative atomic mass: 55.847, r.d. 7.86, m.p.: 1535oC, b.p.: 3000oC, Specific heat capacity: 448 J kg-1 K-1
7.2.2.22 Krypton, Kr, is a non-metal colourless odourless noble gas at room temperature and pressure Kr2, extracted from liquid air, 0.0001% of the air, mixed with other inert gases in fluorescent lamps. Forms few compounds.
Atomic number: 36, Relative atomic mass: 83.80, r.d. 2.16 (121 K), m.p.: -157oC, b.p.: -152oC, Specific heat capacity: 247 J kg-1 K-1
7.2.2.23 Lead, Pb (plumbum) as lead shot, fishing sinkers, roof guttering, foil, powder, filings, strip, is a soft dense unreactive metal, available as lead foil, powder and lead shot, extracted from the ore galena (PbS) used in fishing sinkers, solder, lead glazes and X-ray protective shields, holds the pieces of glass together in stained glass windows, used in bullets, lead shot, fishing sinkers, building construction, lead cell accumulators, solder, pewter, bearings and alloys. Formerly, ladies used lead carbonate to whiten their faces and some may have died from such use. Reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2. No reaction with dilute HCl or H2SO4 or with water. Heated powder forms oxide. Inorganic Pb2+ is an accumulated poison and can replace calcium in bone. A "lead pencil" contains graphite, not lead.
Lead is a metal with a silvery appearance that is resistant to attack by acids because of the formation of a protective oxidized layer on its surface. The metal melts at low temperature and is a good conductor of electricity so it is used in solders. The vapours of molten lead are extremely toxic and the effect of inhaling them is cumulative. Lead salts are toxic by inhalation and can be absorbed through the skin so should be handled with great care. Wash laboratory areas where lead salts have been used with a dilute detergent solution to prevent exposure to any residual dust containing lead. Do not heat lead oxide on a charcoal block.
Atomic number: 82, Relative atomic mass: 207.2, r.d. 11.3 g cm-3| m.p.: 327oC, b.p.: 1744oC, Specific heat capacity: 130 J kg-1 K-1
7.2.2.23a Lead paint
Lead paint formerly contained the white pigment "white lead", basic lead carbonate, Pb(OH)2PBCO3 that is now replaced by titanium dioxide, TiO2. Ingestion and skin absorption of lead causes lead poisoning, especially in young children sucking lead paint peeling off old walls.
"Lead in Paint" from the World Health Organization
No level of exposure to lead is considered safe. The poisonous properties: of lead have been recognized since ancient times, and today it is recognized as one of the twenty leading risk factors contributing to the global burden of disease. Eliminating lead exposure from petrol has been one of the most significant environment health improvements in recent times.
Products containing lead are still widely made and sold across much of the developing world. It is very likely that most of the world's people live in countries where exposure to high lead levels in paint is frequent. Lead in paint is the second largest source of exposure to lead following exposure from petrol. Paint containing lead is used in infrastructure like bridges, industry (car parts) and for marine uses, as well as domestically.
The evidence of neurological damage, especially to children (whose intelligence can be impaired) and to workers in the lead industry is beyond doubt. Adults can suffer renal and cardiovascular damage. Some studies suggest a link to behavioural problems as well. Lead damage is irreversible, and its effects appear to persist into adolescence and adulthood. House dust is the commonest way in which children are harmed by lead in paint. The lead remains a risk for many years after the paint has been used.
7.2.2.23b Tetraethyl lead, lead tetra-ethyl
Octane (C8H18) Octane number 16.1.1h | See 32.5.5.5: Spark plugs, pre-ignition | See 11.11.2: Petrol sniffing
The "anti-knock" additive to petroleum, the lead alkyl lead (IV) tetraethyl, Pb(C2H5)4, lessens pre-ignition, "knocking" by inhibiting combustion reactions and so improving its octane rating. This "leaded petrol" is no longer made or used in Australia and other countries because the compound causes lead poisoning and environmental pollution so is replaced by "unleaded petrol" that contains additional hydrocarbons to improve its octane rating, e.g. methanol and methyl tertiary butyl ether (MTBE). However, some pollution is still produced by unleaded petrol.
7.2.2.24 Lithium, L is a least dense, soft and shiny surface when cut by knife then tarnishes, very reactive alkali metal with acids so stored under oil because reacts with air and water, but least reactive element in group I, red flame test colour, rare element found in some granite pegmatites, used in Al and Mg alloys, batteries and anti-depressant medicines. Reacts with oxygen gas and water, on heating reacts with nitrogen and hydrogen gas. Lithium carbonate is used for a craft flux.
Atomic number: 3, Relative atomic mass: 6.941, r.d. 0.53 g cm-3| m.p.: 180oC, b.p.: 1330oC, Specific heat capacity: 3.39 X 103 J kg-1 K-1
7.2.2.25 Magnesium, Mg, as powder, ribbon, turnings, wire, (photographers' flash bulbs, light bulbs, fire starters, cars, aircraft bodies, alloyed with aluminium), Toxic if ingested, silver white alkaline earth metal, 2% of the earth's crust, forms protective oxide layer in air that prevents further oxidation. Burns in air with intense white light. It is available as powder (FLAM, dangerous), and ribbon (FLAM, safest form for school use) and as turnings (FLAM) low density. It is extracted from sea water, found in magnesite, MgCO3 and dolomite, MgCO3.CaCO3, Reacts with dilute HCl or H2SO4 to form H2 and metal ion, reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with hot water, reacts with halogens, sulfur and nitrogen. Heated powder forms oxide. Magnesium is stored in bones. It is used in many adenosine triphosphate, ATP, reactions. The recommended daily allowance, RDA, is 350 mg for adult males, and 280 mg for adult females. Manganese is a cofactor for many enzymes, but magnesium can usually substitute for it.
Atomic number: 12, Relative atomic mass: 24.305, r.d. 1.74, m.p.: 650oC, b.p.: 1110oC, Specific heat capacity: 1.03 X 103 J kg-1 K-1
7.2.2.26 Manganese, Mn, is a white to reddish colour, hard, brittle transition metal, available as electrolytic flake and manganese (IV) oxide (manganese dioxide) extracted by electrolytic treatment of ores, e.g. pyrolusite (manganese (IV) oxide) used in ferromanganese for alloy steel manufacture. Manganese is a cofactor for many enzymes, but magnesium can usually substitute for it.
Atomic number: 25, Relative atomic mass: 54.9380, r.d. 7.20, m.p.: 1240oC, b.p.: 2100oC, Specific heat capacity: 477 J kg-1 K-1
7.2.2.27 Mercury, Hg (hydrargyrum, formerly "quicksilver") is a liquid metal between 19oC and 27oC, the lowest melting point of all substances, m.p. = -38.9oC, dissolves in most metals to produce amalgams, toxic when inhaled or taken internally or absorbed through the skin, free surface mercury not to be used in schools but can be examined in thermometers, extracted from cinnabar (mercury (II) sulfide), [used in dentistry amalgams, mercury vapour lamps, electrical switches, interferometer lights and detonators, e.g. mercury fulminate, Hg(CNO)2, Reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2. No reaction with dilute HCl or H2SO4 or with water. Heated powder forms oxide. Use mercury in a well-ventilated area that does not have porous surfaces. Do not heat mercury because more toxic mercury vapour is given off. Liquid mercury boils violently at atmospheric pressure. Do not react mercury with chlorine or bromine because highly toxic salts are formed. Mercury reacts with the gold in gold rings to form an amalgam and spoil the ring. Mercury and mercury salts react with aluminium in a violently exothermic reaction. Sodium amalgam, NaHg2, is formed by electrolysis of sodium chloride solution with a mercury cathode. Mercury thermometer should only be used by senior students. Junior students should use alcohol or electronic thermometers. Most mercury compounds are highly toxic if ingested or absorbed through the skin.
Mercury vapour is highly toxic when inhaled and is a cumulative poison. The main concern posed by liquid mercury is spillage, followed by inadequate cleanup. Spilled mercury shatters into many tiny droplets that lodge in cracks in the floor and around skirting boards. The droplets have a large total surface area and the mercury can evaporate slowly, over months or years. Any person working in the area may regularly inhale the mercury vapour, with adverse effects. Liquid mercury is extremely dense so the weight of even a small container may surprise some people who may then drop it. If the pressure measured with a mercury-in-glass manometer suddenly changes because of the careless opening of a tap, the moving column of mercury may break the walls of the manometer and mercury spills everywhere. Avoid the use of mercury by using solid state electronic devices for measuring pressure. Use less dense liquids, even water, in a manometer, to demonstrate changes in pressure. Do not place mercury thermometers in a dish washer for cleaning. Drains in laboratories often contain significant amounts of mercury in the S-bends. Before disposal of broken thermometer debris or mercury-contaminated rubbish from S-bends, mix them with an equal volume of sulfur powder. Fit a tray containing a centimetre of sulfur powder beneath the apparatus because mercury spilled on sulfur powder will not shatter into fine droplets and the sulfur will inactivate the mercury. Pour mercury over a plastic tray containing water so that any spilled mercury will not shatter into droplets, but form a lower layer that can be recovered easily. Periodically check the waste water filter of a laboratory dishwasher for liquid mercury. Store mercury in a sealed container on a plastic tray to catch any spilt mercury.
Mercury spills
Mercury metal gives off a highly toxic vapour that acts as a cumulative poison. Spilt mercury shatters into droplets, which lodge in cracks and crevices to evaporate slowly for years causing long-term exposure to low levels of mercury vapour in the air. So spilt mercury must be immediately treated with sulfur powder or mercury decontaminant, a mixture of calcium hydroxide and sulfur, to form mercury (II) sulfide that is stable and harmless. Then sweep up the mercury and sulfur together. Always keep available 2 kg of sulfur or 2 kg of mercury decontaminant in the laboratory near any equipment containing mercury..
Atomic number: 80, Relative atomic mass: 200.59, r.d. 13.6, m.p.: -38.9oC, b.p.: 357oC, Specific heat capacity: 138 J kg-1 K-1
7.2.2.28 Molybdenum, Mo, silvery solid, transition element, in molybdenite, MoS2, not affected by most acids, used in steel alloys, in two enzyme systems xanthine oxidase and aldehyde oxidase. Hard water can provide some of the daily intake of molybdenum.
7.2.2.29 Neon, Ne, is a non-metal inert noble gas at room temperature and pressure, 0.0018% of the air, with very low reactivity, used for high voltage display lighting, has bright orange colour, "neon" lights, strip lighting, high voltage display lighting and in pilot electrodes (starter electrodes to start up sodium vapour discharge lamps) cold cathode tubes, manufactured by recovery from atmosphere.
Atomic number: 10, Relative atomic mass: 20.179, r.d. 1.20 (27 K), m.p.: -249oC, b.p.: -246oC, Specific heat capacity: 1.03 X 103 J kg-1 K-1
7.2.2.30 Nickel, Ni, as foil, powder, is a transition metal that resists corrosion. It is obtained from ores containing NiS and is available as sheet. It is used in shiny coin alloys, nickel plating, and in "silver" cutlery stamped "EPNS" (electroplated nickel silver) that has a shiny metal protective coating. It reacts with dilute HCl or H2SO4 to form H2 and metal ion and reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2. It reacts with steam to give the oxide and hydrogen gas. Heated powder forms oxide. The name nickel comes from the German kupfernickel (copper goblin). The "nickel" coin in U.S.A is made of a nickel copper alloy.
Atomic number: 28, Relative atomic mass: 58.69, r.d. 8.90 g cm-3| m.p.: 1453oC, b.p.: 2730oC, Specific heat capacity: 439 J kg-1 K-1
7.2.2.31 Nitrogen gas N2 is colourless, odourless, tasteless, neutral and unreactive. Nitrogen does not support combustion. Magnesium and calcium will continue to burn in nitrogen to form nitrides. Nitrogen is manufactured by fractional distillation of air. Air contains about 78.8% of nitrogen. Nitrogen, is a non-metal gas at room temperature and pressure, forms oxides: nitrous oxide N2O, nitric oxide NO, nitrogen dioxide NO2, dinitrogen tetroxide N2O4.
Atomic number: 7, Relative atomic mass: 14.0067, r.d. 0.808 (77 K), m.p.: -210oC, b.p.: -196oC, Specific heat capacity: 1.04 X 103 J kg-1 K-1
7.2.2.32 Oxygen, O, oxygen gas, O2, is a non-metal colourless and odourless gas at room temperature and pressure, 20.95% of the air, 47% of the earth's crust, identified by lighting a glowing splint, supports combustion and is necessary for respiration, made in the laboratory by decomposition of hydrogen peroxide with MnO2 catalyst, manufactured by distillation of liquid air, most abundant element 50% of mass of rocks in earth's crust. Reacts with metals to form basic oxides, reacts with non-metals to form acidic oxides. Oxygen is carried by binding directly to the iron, in its 2+ oxidation state. Monthly loss of iron because of menstruation may average 28 mg. The recommended daily allowance, RDA, is 10 mg for adult males and 15 mg for adult females. Lecithin, diacylphosphatidylcholine is used in cell membranes. The lecithins are mixed esters of glycerol and choline with long chain fatty acids and phosphoric acid.
Atomic number: 8, Relative atomic mass: 15.9994, r.d. 1.15(90 K), m.p.: -218oC, b.p.: -183oC, Specific heat capacity: 916 J kg-1 K-1
7.2.2.33 Phosphorus, P, Phosphorus is a non-metal network solid, occurs mainly as phosphates and in many minerals, e.g. apatite, has two main allotropes, Phosphorus has 2 main allotropes: 1. white phosphorus, (yellow phosphorus) and 2. red phosphorus, see below. Phosphates are important agricultural fertilizers, e.g. NPK. Phosphorus occurs as inorganic calcium phosphate in bones and teeth, in tissue and in the ATP molecule. The recommended daily allowance, RDA, is 1200 mg. Phosphoric acid, H3PO4, behaves as a tribasic acid although the normal salts are much hydrolysed in solution.
7.2.2.33a Red phosphorus, brown phosphorus, P4, is not poisonous, ignites above 300oC, has phosphorus atoms bound in a covalent network so is less reactive than white phosphorus and can be stored in air, used in "safety match" striking surfaces that contain red phosphorus, powdered glass, carbon black, a binder and a neutralizer, e.g. calcium carbonate. Used for "doping" semiconductors. Red phosphorus is relatively harmless compared with the white phosphorus allotrope. However, some school systems do not allow red phosphorus to be used in school science experiments. It is not poisonous when pure. It is deliquescent so keep it in a sealed container. Do not heat red phosphorus in a test-tube because it produces phosphorus vapour that condenses to form white phosphorus. Re phosphorus form violently explosive mixtures with oxidizing agents, e.g. metal nitrates (potassium nitrate, sodium nitrate, potassium permanganate), nitric acid, chlorates, perchlorates, peroxides, peroxy salts. Do not heat phosphorus on a platinum wire because it corrodes the platinum.
P (red) Atomic number: 15, Relative atomic mass: 30.9738, r.d. 2.34, m.p.: 590oC, b.p.: 280oC, Specific heat capacity: 670 J kg-1 K-1
7.2.2.33b White phosphorus (yellow phosphorus), P4, is too reactive to be used in school science teaching so it is not permitted in schools, White phosphorus is waxy, poisonous, spontaneously flammable, reacts with oxygen gas in the air to form P2O5, so it is stored under water. White phosphorus is extremely toxic. On contact with iodine it ignites, with bromine it explodes, and it can be ignited with a hot glass rod. Cover spilt white phosphorus with 0.2 M copper sulfate solution that converts it to harmless copper sulfide.
P (white) Atomic number: 15, Relative atomic mass: 30.9738, r.d. 1.82, m.p.: 44.2oC, b.p.: 44oC, Specific heat capacity: 757 J kg-1 K-1
7.2.2.33.1 Phosphorescence is the green glow from the slow oxidation of white phosphorus. Luminescence is emission of light for any any reason other than a rise in temperature, e.g. excited photons returning to a ground state. Chemiluminescence is luminescence resulting from a chemical change. However, the term phosphorescence is also used to describe a situation when the luminescence persists even though the exciting cause has been removed. Luminescence that does not persist when the exciting cause is removed is called fluorescence, e.g. a fluorescent light.
Order online: Fluorescence Activity Kit, four fluorescent minerals
7.2.2.33.2 Detergent phosphates, as polyphosphates, mainly sodium tripolyphosphate (STPP), are used in detergents for different functions. Without using polyphosphates several different chemicals would be needed to replace them. There are also used in ceramics and are an authorized food ingredient. Phosphates occur in sewage from detergents, human foods, (transferred into human wastes) animal manure and food industry wastes. These phosphates can be recovered and recycled back into fertilizers and the detergent industry. Phosphates are the only recyclable ingredient of detergents.
See 8.1.0: Heat sources, light a match
7.2.2.34 Platinum, Pt, is a soft, ductile transition metal, resists most chemical agents and does not oxidize at high temperature, available as foil and wire, occurs in free elemental form placer deposits or in alloys, used for electrical contacts, electrodes and jewellery. No reaction with dilute HCl or H2SO4, air, water or concentrated oxidizing acids, e.g. HNO3 or H2SO4, reacts with aqua regia (concentrated HNO3 + HCl) to form H2PtCl6.
Atomic number: 78, Relative atomic mass: 195.08, r.d. 21.4, m.p.: 1769oC, b.p.: 4530oC, Specific heat capacity: 134 J kg-1 K-1
7.2.2.35 Potassium, K (kalium) is a soft and shiny white alkali metal, highest metal in the reactivity series so does not occur in the free state, 2.5% of the earth's crust, available as the metal in liquid paraffin and is used widely in chemistry, is the " K" in the common fertilizer NPK, usually as K2O. Reacts with dilute HCl or H2SO4 to form H2 and metal ion, reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with cold water and reacts with air to form peroxides. Potassium has less concentration in extra cellular fluids than in intracellular fluids where it is the primary cation. Potassium ion is lost in sweat along with sodium ion. (in liquid paraffin)
Atomic number: 19, Relative atomic mass: 39.0983, r.d. 0.86, m.p.: 63.7oC, b.p.: 774oC, Specific heat capacity: 753 J kg-1 K-1
7.2.2.36 Radium, Ra, is a radioactive alkaline earth metal, used in some cancer therapy, extracted from pitchblende, uraninite, used as a radioactive resource.
Atomic number: 88, Relative atomic mass: 226.025, r.d. 5.0, m.p.: 700oC, b.p.: 1140oC, Specific heat capacity: 121 J kg-1 K-1
7.2.2.37 Radon, Rn, is a non-metal noble colourless water-soluble gas at room temperature and pressure Rn2, formed by decay of radium-226, forms under granite where it may be a health habit to people living in granite houses, has few compounds, used in radiotherapy.
Atomic number: 86, Relative atomic mass: (222), r.d. 4.4 (211 K), m.p.: -71oC, b.p.: -61.8oC, Specific heat capacity: 92 J kg-1 K-1
7.2.2.38 Selenium, Se, is a non-metal, obtained from sulfide ores, decolorizes glass, semiconductor that conducts electricity when e.m.f. applied, used in photoelectric cells and light meters. Similar properties: to sulfur. The Kjeldahl catalyst is sodium sulfate + selenium. Selenium is a trace mineral used to destroy hydrogen peroxide. The recommended daily allowance, RDA, is 70 µ g for males, and 55 µ g for females.
Atomic number: 34, Relative atomic mass: 78.96, r.d. 4.81, m.p.: 217oC, b.p.: 685oC, Specific heat capacity: 322 J kg-1 K-1
7.2.2.39 Silicon, Si, lump [powder] is a non-metal network solid with many properties: similar to carbon. occurs naturally as a brown powder or grey crystals, 28% of the earth's crust, forms a network solid similar to diamond and has valence 4, the second most abundant element mainly in silicates in rocks, prepared by reduction of silica SiO2 in an electric furnace, used as a semiconductor, conductivity can be increased by increasing its temperature or by adding small quantities of boron or phosphorus (doping). The common form of silicon dioxide is silica, SiO2 as in quartz. A major component of rocks is the silicate ion, SiO44-, as in glass. Silicone greases have a polymer network of silicon and oxygen atoms attached to C and H atoms.
Atomic number: 14, Relative atomic mass: 28.0855, r.d. 2.33, m.p.: 1410oC, b.p.: 2360oC, Specific heat capacity: 711 J kg-1 K-1
7.2.2.39.1 Solar panel
Order online: Solar Racer, toy car moved by solar energy
A crystalline silicon cell used in solar panels for solar power has an efficiency of about 22% for the conversion of light to electricity (16% for the cheaper cells in household solar panels). A concentrated solar power system (CSP) use lenses or mirrors to concentrate solar thermal energy, onto a small area to heat water in a steam turbine and generate electricity. A concentrated photovoltaic system (CPV) focus sunlight onto photovoltaic materials to generate electricity directly.
7.2.2.40 Silver, Ag (argentum) is a silver white transition metal, does not oxidize in air and is the best electrical conductor, available as silver wire, occurs as the element and is used in coin alloys, electrical conductors, photographic emulsions, jewellery and ornaments. Reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with concentrated HNO3 and hot concentrated H2SO4. No reaction with dilute HCl or H2SO4, air, or water. The "silver paper" used in wrapping paper or in chocolates is usually tin foil.
Atomic number: 47, Relative atomic mass: 107.868, r.d. 10.5, m.p.: 961oC, b.p.: 2210oC, Specific heat capacity: 234 J kg-1 K-1
7.2.2.41 Sodium, Na (in liquid paraffin) (natrium),  is a very reactive, soft and shiny alkali metal with a silvery colour, 2.5% of the earth's crust, does not occur in the free state and soon becomes dull in air, available as sodium metal in liquid paraffin, extracted by electrolysis of fused sodium hydroxide, rock salt (halite) and evaporated sea water, used in sodium vapour discharge lamps. Reacts with dilute HCl or H2SO4 to form H2 and metal ion, with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with cold water and reacts with air to form peroxides. Sodium has greater concentration in extra cellular fluids than in intracellular fluids. The level of sodium ion in blood is 310 to 333 mg per 100 mL of serum.
Sodium metal reacts violently with water to form highly corrosive sodium hydroxide and hydrogen gas. The surface of the reacting sodium may become sufficiently hot to ignite the hydrogen and explode the air / hydrogen mixture. Reacts explosively with acids, bromides, iodine, sulfur, polyhalogenated hydrocarbons, e.g. chloroform. The surface of sodium metal is coated with a mixture of the oxide and hydroxide, both of which hygroscopic, causing the surface of sodium metal to absorb moisture from the air, and resulting in further reaction to form hydrogen and sodium hydroxide. Sodium metal, unprotected in a container, may absorb sufficient moisture to form an air / hydrogen mixture and then become hot enough to ignite the mixture, resulting in an explosion. So store sodium under paraffin oil or kerosene. Store sodium on a low shelf to reduce the chance of dropping the container. Cut sodium on a clean dry surface with a clean dry blade, e.g. an all metal scalpel or small knife, not a scalpel with detachable blade. Do not lose small pieces of cut sodium. Dispose of sodium waste by reacting it with ethanol or methylated spirit in a beaker in a fume cupboard. The reaction of sodium with an alcohol is less violent than with water, so it is relatively slow and easy to control, but there is still a risk of fire. The use of flammable solvents in the vicinity of sodium is always a risk.
Atomic number: 11, Relative atomic mass: 22.9898, r.d. 0.97, m.p.: 97.8oC, b.p.: 890oC, Specific heat capacity: 1.23 X 103 J kg-1 K-1
7.2.2.42 Strontium, Sr, silver white metallic element, in minerals celestine SrSO4, strontianite SrCO3 and spring water, similar properties to calcium, salts used for crimson flame in fireworks, after nuclear explosion fallout contains 90Sr that can be absorbed in human bone, strontium chloride in toothpaste for sensitive teeth, e.g. "Sensodyne", harder to cut than sodium, use calcium as cheaper alternative, purchase only small quantity and store in tightly-sealed container, give students only few grains per activity
Atomic number: 38, Relative atomic mass: 87.62, r.d. 2.54, m.p.: 800oC, b.p.: 1300oC, Specific heat capacity: 0.3 J kg-1 K-1
7.2.2.43 Sulfur, S,
See diagram 12.18.1: Sulphur crystals
Sulfur, S, brimstone, sublimed sulfur, flowers of sulfur, sulfur sticks, roll sulfur, is a non-metal occurs as S8 molecule rings. Sulfur is insoluble in water, slightly soluble in ethanol, and soluble in benzene. Above 160o the S8 molecule rings break to form long chains of plastic sulfur, polymeric sulfur, that is not soluble in any solvents. Sulfur is not toxic, but can be dangerous because of its flammability and hazardous because of the formation of sulfur dioxide gas as a product of combustion. Used as example of a non-metal element. The main allotropes are as follows:
1. Alpha sulfur, rhombic sulfur, with yellow octahedral crystals
2. Beta sulfur with monoclinic prismatic pale yellow crystals.
Do not mix sulfur with oxidizing agents, e.g. potassium nitrate or potassium permanganate, or mercury (II) oxide, because the mixtures are explosive. Do not prepare mixtures of sulfur with active metal powders, e.g. aluminium and magnesium, because they are also violently reactive. Sulfur usually ignites when it is heated, forming sulfur dioxide gas. This gas is highly irritant to the lungs and should not be inhaled. Use a fume cupboard for experiments that produce sulfur dioxide. Some asthmatics are particularly sensitive to sulfur dioxide.
Atomic number: 16, Relative atomic mass: 32.06, r.d. 2.07 (Alpha), 1.96 (Beta), m.p.: 113oC (Alpha) 119oC (Beta), b.p.: 445oC (Alpha), Specific heat capacity: 0.71
7.2.2.44 Tin, Sn (stannum) (tin-plated jam tin) pewter is a ductile, malleable metal, resists corrosion, is available as granules, pellets, foil and powder, is extracted from the ore cassiterite (tin (IV) oxide, SnO2) and used in alloys and the shiny protective plating of "tin cans", (tin plate) and aluminium drink-cans, and in alloys and solders. It reacts with dilute HCl or H2SO4 to form H2 and metal ion, reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2. No reaction with water. Heated powder forms oxide. The "silver paper" used in wrapping is usually tin foil.
Atomic number: 50, Relative atomic mass: 118.71, r.d. 7.28 (white), 5.75 (grey), m.p.: 232oC, b.p.: 2270oC, Specific heat capacity: 218 J kg-1K-1
7.2.2.45 Tungsten, W, white-grey solid, transition element (formerly wolfram) light bulb filaments, from wolframite (MnFe)WO4, scheelite CaWO4, not affected by dilute acids, used in steels and lamp filaments, tungsten carbide, WC, black powder, used in cutting tools because has Mohs scale 9.5 The tungsten filament in a light bulb reaches about 2,300oC. It has the highest melting point of all metals.
Atomic number: 74, Relative atomic mass: 183.85, r.d. 19.3, m.p.: 3,422oC, b.p.: 5,660oC, Specific heat capacity: 130 J kg-1 K-1
7.2.2.46 Uranium, U, is a radioactive hard grey metal obtained from pitchblende U3O8, 238U main isotope, 235U used as fuel in nuclear power stations and atomic bombs. Atomic number: 92, Relative atomic mass: 238.029, r.d. 19.1, m.p.: 1130oC, b.p.: 3820oC, Specific heat capacity: 117 J kg-1 K-1
7.2.2.47 Xenon, Xe, is a non-metal noble gas at room temperature and pressure, 0.00001% of the air.
Atomic number: 54, Relative atomic mass: 131.29, r.d. 3.52(165 K), m.p.: -112oC, b.p.: -108oC, Specific heat capacity: 159 J kg-1 K-1
7.2.2.48 Zinc, Zn, granulated, zinc foil, dry cell battery case, is a blue white metal, forms protective oxide layer in air that prevents further oxidation, is available as zinc granules, foil, filings and powder (FLAM) extracted from zinc blende (sphalerite, ZnS) used for dry cell battery casings and in the Daniell cell, galvanized iron, alloys and brass, a micronutrient (trace element) required in a very small quantity by living organisms as salts or compounds, not as the pure element. Reacts with dilute HCl or H2SO4 to form H2 and metal ion, reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with steam to give the oxide and hydrogen gas. Heated powder forms oxide. High level of zinc in the diet is undesirable, e.g. from oysters, and zinc deficiency symptoms occur where people live on unleavened bread made from highly extracted wheat flour and no meat or yeast products in the diet. Zinc is a cofactor for about 20 enzymes, e.g. alcohol dehydrogenase that breaks down ethanol and carboxypeptidase that catalyses the hydrolysis of proteins in the small intestine. The recommended daily allowance, RDA, is 15 mg for males and 12 mg for females.
Zinc is dangerous only as a fine powder, zinc dust, that is flammable and has been used in the past to make rocket fuel by mixing it with finely-divided sulfur. This procedure is not permitted in schools. Zinc / sulfur mixtures are extremely hazardous materials, liable to combust violently on ignition and release large amounts of sulfur dioxide from oxidation of unreacted sulfur. Zinc dust forms hazardous mixtures with iodine and many oxidizing agents, e.g. manganese dioxide, potassium nitrate and potassium permanganate.
Atomic number: 30, Relative atomic mass: 65.39, r.d. 7.14, m.p.: 420oC, b.p.: 907oC, Specific heat capacity: 385 J kg-1 K-1
7.2.2.49 Zirconium, Zr, grey-white solid, transition element, in zircon ZrSiO4, zirconium wire used in photography flash bulbs, camera flash cubes, e.g. "Magicube", nuclear reactor neutron absorber
7.2.2a Elements experiment
Describe each example.
1. Note the state of matter at room temperature, solid, liquid or gas.
2. Note whether the solid has a shiny surface or has a lustre when the surface is clean.
3. Note whether the metal can be bent or twisted with pliers, or whether it fractures.
4. Note whether the element conducts electricity when held between two alligator clips as electrical contacts.
5. Put a piece of the element on a combustion spoon and set it alight with a burner flame. Observe the burning element.
6. Shake the products of the combustion in a test-tube containing water. Test the solution with moist litmus paper.