Topic 2 Bonding 2024 Past Paper Co-Sci 10CSI52 PDF

Topic 2 BONDING 2 3 4  5 WHY COMPOUNDS ARE FORMED Most naturally occurring substances exist as molecules, made up of atoms joined together e.g. CH4, CO2, H2O. Many elements also exists as mol...

Topic 2 BONDING 2 3 4  5 WHY COMPOUNDS ARE FORMED Most naturally occurring substances exist as molecules, made up of atoms joined together e.g. CH4, CO2, H2O. Many elements also exists as molecules e.g. O2, N2, I2. However there are a few elements that exist as simple separate atoms. Read Kognity sections C2.4 and C2.5 (page 44 3rd Edition, page 28 4th Edition), to help you answer the following question. 1. Which group of elements exist as separate atoms? ………………………………………………………………………………………………………… In general, molecules are more stable than the atoms from which they are made. Let us now consider why the bonding of atoms together make them more stable? Consider a simple diatomic molecule i.e. a molecule made of two atoms. When the two atoms come close together to form the molecule it is the electrons around the outside of the atoms that will meet first. The nuclei of the atoms will never touch. Bonding must therefore involve an interaction between the outside electrons from the different atoms. Each atom in the molecule tries to make its electron arrangement more stable as a result of bonding. What then is a stable electron arrangement? We can answer this question by finding out which separate atoms are the most stable. If a group of elements exist as separate atoms it is reasonable to conclude that their electron arrangements cannot be made any more stable. 2. Write down the names and electron arrangements of the first 3 members of the Group which exist as separate single atoms: ELEMENT ELECTRON ARRANGEMENT 3. What is different about the electronic arrangements of these elements compared to other elements? ……………………………………………………………………………………………………………………………………………………………………………………… ……………………………………………………………………………………………………………………………………………………………………………………….  6 When atoms bond together to form molecules they try to increase their stability by changing their electron arrangement until it is the same as that of one of the noble gases. This process is called obtaining a noble gas configuration. Atoms can do this in one of two ways:- i) By gaining or losing electrons (IONIC BONDING) ii) By sharing electrons (COVALENT BONDING) You will study these two types of bonding in future lessons. Before doing this it is worth making a simple generalisation about the type of bonding in different compounds. If the compound contains a metallic element then the bonding will be IONIC and if the compound contains only non-metallic elements then the bonding will be COVALENT. 4a. How many electrons must an atom of chlorine lose to obtain a noble gas configuration? …………………………………………………….. b. How many electrons must an atom of chlorine gain to obtain a noble gas configuration? …………………………………………………….. c. Which do you think chlorine is more likely to do – lose or gain electrons? …………………………………………………….. When Atoms join together (bond) they become more stable. 5a. How many electrons must an atom of sodium lose to obtain an inert gas configuration? …………………………………………………….. b. How many electrons must an atom of sodium gain to obtain an inert gas configuration? …………………………………………………….. c. Which do you think sodium is more likely to do – lose or gain electrons? …………………………………………………….. d. Can an atom of sodium get eight electrons by sharing its electrons. …………………………………………………….. e. Explain your answer to d. ………………………………………………………………………………………………………………………………………………………… …………………………………………………………………………………………………………………………………………………………  7 ATOMS INTO IONS Demonstration Video The reaction of sodium with chlorine Record your observations for this demonstration. Write a balanced equation to represent the reaction. EVIDENCE FOR THE EXISTENCE OF IONS One of the first experiments that gave scientist an idea of what holds particles together was that certain substances could be broken down (decomposed) by electricity. It was Sir Humphry Davy who first discovered this process that was called electrolysis in 1800 and realised the importance of this new type of chemical change. From these experiments he concluded that substances that could conduct electricity when they are dissolved in water or molten must be made up of positive and negative particles (ions). He also suggested that these substances where held together by electrostatic attraction of the positively charged particles to the negatively charged particles. In 1807 he isolated Potassium metal for the first time and three days later he obtained Sodium, Barium, Strontium, Calcium and Magnesium. ELECTROLYSIS OF COPPER CHLORIDE Apparatus Draw a labelled diagram of the apparatus used. What happens to the light bulb / ammeter when the solution of copper chloride is tested (compared to the solid)? ……………………………………………………………………………………………………………………………………………………………………………………… What do you see at the anode (positive electrode). ………………………………………………………………………………………………………………………………………………………………………………………  8 You probably cannot see anything happening at the cathode (negative electrode). But what prediction can you make? ……………………………………………………………………………………………………………………………………………………………………………………… What is seen on this electrode when it is removed from the solution? ……………………………………………………………………………………………………………………………………………………………………………………… Conclusion Explain why the solid copper chloride does not conduct electricity whereas the solution does. ……………………………………………………………………………………………………………………………………………………………………………………… ……………………………………………………………………………………………………………………………………………………………………………………… ……………………………………………………………………………………………………………………………………………………………………………………… Explain, with the help of diagrams and electrode equations what is happening (see page 100 3rd Edition, page 84 4th Edition or Kognity C4.1).  9 IONIC BONDING The simplest way for an atom to obtain an inert gas configuration is to lose or gain electrons. After an atom has lost or gained electrons it can no longer be called an atom because it is no longer neutral. It is called an ion. Ions are formed when atoms lose or gain electrons Generally metals in Groups I, II and III will transfer electrons to non-metals in Groups V, VI and VII. Metals will form positive ions and non-metals will form negative ions. Ref pages 46-49 3rd Edition, pages 30-33 4th Edition or Kognity C2.4 Examples (a) sodium + chlorine  sodium chloride Atoms  Ions Formula = (b) magnesium + chlorine  magnesium chloride Atoms  Ions Formula =  10 Exercise Draw bonding diagrams for the formation of the following compounds. Show only the outer valence electrons in your diagrams. 1. magnesium oxide 2. potassium oxide 3. calcium chloride 4. Now try as many of the following: (a) aluminium chloride; (b)* lithium bromide; (c) magnesium sulphide; (d) lithium nitride; (e) sodium hydride (f) aluminium oxide; (g)* strontium iodide * remember you need only consider the outer electrons.  11 WRITING CHEMICAL FORMULAE Chemical formulae are a shorthand way of representing the ratio of atoms of each element present in a compound. A chemical formula gives 2 important pieces of information : (a) It identifies the different types of atom found in a molecule. (b) It indicates how many of each type of atom there are in the molecule. The following notes and examples will help you to write chemical formulae. You can use the valency tables on the back of this sheet to help but it is best to learn valencies as soon as you can. Rules for Writing Chemical Formulae 1. Find the chemical symbols for the ions present, including their charges (valencies indicate the number of electrons lost or gained when bonding happens) (charges are “superscript”) 2. Now re-write the formulae without the charges but include the number of each ion present. This number should be a “subscript” and written after the symbol(s). 3. The number of each ion present should mean that all charges on the original ions cancel out. 4. If the figure 1 appears in the formula after simplification it is left out. e.g. Pb1O2 becomes PbO2. 5. If the formula of the ions is “polyatomic”, brackets should be written around the whole ion to indicate the number of that whole ion present. Example 1 : Magnesium chloride Example 2 : Sodium Carbonate Symbols : Mg2+ Cl- Symbols : Na + CO32- No. req’d : Mg1 Cl2 No. req’d : Na2 (CO3) 1 Therefore formula: MgCl2 Therefore formula: Na2CO3 Example 3 : Aluminium oxide Example 4 : Zinc hydroxide Symbols : Al3+ O2- Symbols : Zn2+ OH - No. req’d : Al2 O3 No. req’d : Zn1 (OH)2 Therefore formula: Al2O3 Therefore formula: Zn(OH)2  12 VALENCY TABLE Ionic Valency (a) Valencies of Metal Ions (Cations) 1 2 3 Sodium Na+ Calcium Ca2+ Aluminium Al3+ Potassium K+ Magnesium Mg2+ Iron (III) Fe3+ Silver Ag+ Zinc Zn2+ Ammonium NH4+ Copper (ll) Cu2+ Copper (l) Cu+ Iron (II) Fe2+ Hydrogen H+ Lead Pb2+ Barium Ba2+ (b) Valencies of Non-metal Ions and Radicals (Anions) 1 2 3 Chloride Cl- Oxide O2- Phosphate PO43- Bromide Br-- Sulphide S2- Iodide I- Carbonate CO32- Fluoride F- Sulphate SO42- Nitrate NO3- Sulphite SO32- Nitrite NO2- Hydroxide OH- Hydrogencarbonate HCO3- Hydrogensulphate HSO4- Covalency 1 2 3 4 Hydrogen, H Oxygen, O Nitrogen, N Carbon, C Chlorine, Cl Sulphur, S Bromine, Br Iodine, I  13 Have a go at the formula using the ionic valency tables. (1) Calcium oxide (6) Copper (II) nitrate (2) Magnesium oxide (7) Zinc nitrate (3) Calcium chloride (8) Sodium hydroxide (4) Sodium chloride (9) Calcium hydroxide (5) Copper (II) sulphate (10) Iron (II) sulphate (11) Aluminium hydroxide (15) Potassium chloride (12) Ammonium nitrate (16) Ammonium sulphate (13) Silver chloride (17) Iron (III) hydroxide (14) Sodium hydrogen sulphate (18) Calcium phosphate  14 COVALENT BONDING Start this topic by reading Kognity C2.5, (3rd Edition p50-53, 4th Edition p34-37 in your text book). Non-metal atoms are joined together by covalent bonds. Outer electrons of the atoms are shared until every atom in the molecule has a share in an inert gas configuration. Each atom shares one of its own electrons with one from another atom. This is repeated until it has a share of 8 outer electrons. Only atoms with 4 or more electrons in their outer energy level can do this. A covalent bond is the attraction between a shared pair of electrons and the nuclei of the bonded atoms. Hydrogen Show how the electrons are shared in a hydrogen molecule. After sharing electrons, both hydrogen atoms must have complete outer electron energy level. Electron Structure H = 1 Bonding diagram Structural Formula : Molecular Formula : Shape : Collect a box of “atoms” and build a model of hydrogen (white). Fluorine Consider the fluorine molecule, F2, as an example. A fluorine atom has an outer energy level of seven electrons. Two fluorine atoms combine to form a molecule, by sharing two electrons – one from each atom. Each atom then has a share of eight electrons and an electron structure like that of neon, the nearest noble gas. Draw the fluorine atom showing the outer electrons only. Electron arrangement: element F : 2, 7 Fluorine molecule outer electron structure only Structural Formula : Molecular Formula : Shape :  15 Hydrogen Chloride Colours : H - white Cl - green A molecule of hydrogen chloride has one atom of hydrogen covalently bonded to an atom of chlorine. Shown in the diagram how this can be done. Remember in the molecule both atoms have complete outer energy levels. Electron structure: elements H 1; Cl 2, 8, 7 Molecule Structural Formula : Molecular Formula : Shape : You should have worked out by now that the number of bonds an atom can have is the same as the number of electrons it needs to fill it’s outer shell. Oxygen O2 Colours : O - red Electron structure: element O = 2, 6 molecule Structural Formula : Molecular Formula : Shape : Nitrogen N2 Colours : N - blue ; Electron structure: element molecule Structural Formula : Molecular Formula : Shape : How many bonds will be formed between the nitrogen atoms  16 Methane CH4 Carbon has four electrons in the outer energy level. Show how carbon can bond with hydrogen to fill its outer energy level. This compound is methane. Electron structure: elements Molecule Structural Formula : Molecular Formula : Shape : The bonds in covalent molecules arrange themselves around the central atom as far apart as possible. This gives molecules like methane definite shape. Look at the model of methane and draw it. The shape is described as tetrahedral. Atoms in molecules are held tightly together because the negative electrons between the atoms attract the positive nuclei and pull them together. Using the above example as a guide draw a diagram to show how the outer energy level electrons are shared in each of the following compounds. From your diagram determine the molecular and structural formula of each compound. Check you answer. Make models of the molecules. Draw and describe the shape of each molecule. 1. Ammonia (nitrogen hydride) Electron structure: elements Molecular Formula : Structural Formula : Shape :  17 2. Tetrachloromethane (carbon tetrachloride) Electron structure: elements Molecular Formula : Structural Formula: Shape : 3. Carbon dioxide (think carefully about this one) Electron structure: elements Molecular Formula : Structural Formula : Shape : 4. Water H2O Electron structure: elements Molecular Formula : Structural Formula : Shape : For the expert try to draw bonding diagrams of the following molecules: ethene C2H4; ethyne or acetylene C2H2, ethanol or alcohol CH3CH2OH; phosphorous trichloride  18 MOLES OF MOLECULES Compounds are also made up of molecules. But a compound contains atoms of different elements joined together by chemical bonds. The formula for water is H2O. This means that a molecule of water contains two atoms of hydrogen and one atom of oxygen. To work out the mass of one mole of a compound we use the same idea as in Unit 1. The relative masses of all the atoms in the compound are added together. This total is called the relative formula mass or Mr Example One What is the mass of one mole of water, H2O? (H = 1, O = 16) Answer The water molecule contains two atoms of hydrogen and one atom of oxygen. Mr H2O = (2 x 1) + (1 x 16) = 18. Therefore the Molar mass = 18 g/mole. The molar mass of a compound contains one mole of molecules. Hence in 18 g of water there are 6 x 1023 molecules of water. When working out the molar mass of a compound containing ions, such as nitrates, sulphates and carbonates, it is important to include every atom. Example Two What is the molar mass of magnesium nitrate, Mg (NO3)2? (Ar Mg = 24, N = 14, O = 16) Answer Mg = 24 2N = 28 6O = 96 Mr Mg(NO3)2 = 148 Molar mass = 148 g/mole.  19 Exercise 6.1 Find the molar mass of the following. You MUST show your working. (a) copper (II) oxide, CuO (b) sulphur trioxide, SO3 (c) copper (II) sulphide, CuS (d) copper (II) carbonate, CuCO3 (e) zinc nitrate, Zn(NO3)2 (f) ammonium carbonate, (NH4)2CO3 (g) glucose, C6H12O6 To work out how many moles there are in a certain mass of a compound we use exactly the same formula. Remember Molar Mass = g/mole, rearranging Moles = mass molar mass Example Three How many moles are contained in 72 g of water? Answer We have already found the relative molecular mass of water to be 18. Therefore molar mass = 18 g/mole. So, Moles = mass = 72 g = 4 moles molar mass 18 g/mole  20 Exercise 6.2 Find the number of moles contained in the following. You MUST show your working. (a) 40 g copper (II) sulphate, CuSO4 (b) 282 g zinc carbonate, ZnCO3 (c) 60 g sulphur trioxide, SO3 (d) 32 g ammonium carbonate, (NH4)2CO3 𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚 Exercise 6.3 Remember 𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚 = , so 𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚 𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚 Mass (g) = Mole x Molar Mass Find the mass of (a) 0.5 mole calcium carbonate, CaCO3 (b) 0.25 mole copper (II) oxide, CuO (c) 0.1 mole of sulphur trioxide, SO3 (d) 0.25 mole ammonium carbonate, (NH4)2CO3  21 AN INVESTIGATION OF THE PROPERTIES OF SOME COMPOUNDS In this experiment you will investigate the physical properties of an ionic compound, sodium chloride, and a covalent compound, stearic acid. This is an organic solvent which is like nail varnish remover (acetone/propanone). Do the tests described below and record your results in the table overleaf. See p.42-43 & 48-49 (2nd Edition 58-59, 3rd Edition 54- 55 to read more on the effect of bonding on the properties of substances. SAFETY CHECK Property Procedure Discussion a) What does this suggest about Appearance What colour are they? Are they solids, the arrangement of particles in liquids, powders or crystals? these substances? b) What does this tell you about Effect of gentle heat. the melting points and the energy Use ignition tubes needed to separate the molecules? Beaker containing some of the pure substance. 4 volts max. c) Are the solid substances Does the substance conductors? What does this tell conduct electricity? you about the particles in the Ask your teacher to check the solid? apparatus. Is water a good solvent for stearic Solubility in water Add 25cm3 of water to the beaker & stir acid? Why? with a glass rod. Beaker containing some of the pure Does the solution substance. 4 volts max. d) What does the water do to the conduct electricity? solid to allow conduction of electricity? Solubility in propanone. Add a small spatula of the substance to 5 e) Which is the better solvent, Demonstration cm3 of propanone in a small test-tube water or cyclohexane? Does the propanone As above if dissolves Explain your observation solution conduct? Demonstration  22 INVESTIGATING SODIUM CHLORIDE AND STEARIC ACID Property Sodium Chloride Stearic acid Appearance Effect of gentle heating Does the pure solid conduct electricity Solubility in water Solubility in cyclohexane If it is soluble does the solution conduct electricity From what you have discovered in this experiment and the information on pages 42-55 3rd edition, pages 26-39 4th Edition, draw up a table on lined paper of the typical physical properties of ionic compounds and simple covalent compounds, adding detail by answering the discussion questions from the table on the previous page.  22 INVESTIGATING SODIUM CHLORIDE AND STEARIC ACID Property Sodium Chloride Stearic acid Appearance white crystalline solid white powdery solid Effect of gentle heating a lot of heating needed to melted easily melt it Does the pure solid conduct No No electricity Solubility in water Dissolves easily Not soluble Solubility in cyclohexane Not soluble Soluble If it is soluble does the solution Yes No conduct electricity From what you have discovered in this experiment and the information on pages 42-55 3rd edition, pages 26-39 4th Edition, draw up a table on lined paper of the typical physical properties of ionic compounds and simple covalent compounds, adding detail by answering the discussion questions from the table on the previous page. Ionic Covalent  Physical Properties of Ionic and covalent compounds Ionic bonding Covalent bonding Lattice structure of ionic compounds Intermolecular forces – weak forces of +ve & -ve ions in ionic compound don’t form attraction between molecules – determines molecules but form crystals made of whether compound is a solid, liquid or gas. repeating pattern of +ve & -ve ions called a giant ionic lattice. Strong electrostatic forces of attraction in all directions between the oppositely charged ions. E.g. Sodium chloride Property Reason Low melting & Weak intermolecular boiling points forces of attraction Usually liquid, gas between molecules or low melting point solid Don’t conduct No mobile electricity ions/electrons Usually insoluble in water (unless they react) Summary Ionic Compound Covalent Compound Metal + Non-metal Non-metal + Non-metal Transfer electrons (metal to non-metal) Sharing electrons Positive and negative charges No charges Solid at room temperature (25oC) Solid, liquid or gas at room temperature High melting and boiling points Low melting and boiling points Strong attraction between particles Weak attraction between molecules  23 REDOX REACTIONS DEFINING OXIDATION AND REDUCTION Read Kognity C6.3 as an introduction to this topic. (3rd Edition 88-91, 4th Edition 72-75). When substances combine with oxygen they are said to be oxidised. Thus magnesium must be oxidised to make magnesium oxide. When we want to remove oxygen from the compound e.g. oxygen from copper oxide then an opposite process to oxidation is carried out. This is called reduction. The first chemists (alchemists) who extract metals from their oxide ore found that they got very little metal from a large lump of ore. So they called the process ‘reducing’ the metal ore. We still use the same term today but for chemists ‘reducing’ now means removing oxygen. For thousands of years, the method used to reduce metal ores was to heat them with charcoal – a form of carbon. Oxidation is the addition of oxygen to a substance. Reduction is the removal of oxygen from a substance. Look back to pages 9 & 10 of this topic and use the dot and cross diagrams to show the bonding in magnesium oxide and magnesium chloride. A transfer of electrons occurs during oxidation and reduction. Magnesium combines with oxygen to give magnesium oxide. Magnesium atoms have the electronic configuration 2,8,2. Oxygen atoms have the electronic configuration 2, 6. When magnesium combines with oxygen, the magnesium atoms lose two electrons to become magnesium ions Mg2+. At the same time, the oxygen atoms gain two electrons to become oxide ion O2-. In this oxidation reduction reaction, there has been a transfer of electrons. When magnesium reacts with chlorine the magnesium loses the same two electrons it would if it had reacted with oxygen so it has been oxidised only by chlorine instead of oxygen. Oxygen is involved in the reaction to make MgO but not in the reaction to make MgCl2. However, in both reactions, magnesium atoms lose electrons. At the same time, another substance gain electrons. Because of this, chemist have devised another set of definitions for oxidation and reduction:  24 Oxidation Is the Loss of electrons by a substance OIL Reduction Is the Gain of electrons by a substance RIG These definitions apply whether or not oxygen is involved. The relation between the two meanings for oxidation and reduction is shown in the Venn diagram below. When one substance loses electrons, another substance must gain electrons. Again, we see that oxidation and reduction always happen together – we cannot have one, without the other. Oxidation and reduction occur together. One cannot take place without the other. Reactions of this type are usually called redox reactions (reduction-oxidation). Chemists have used the term oxidation to mean the addition of oxygen for about 200 years. Similarly, reduction was thought of as loss of oxygen by a substance. All burning processes are oxidation reactions, whether they are rapid (like fires and explosions) or slow (like rusting and respiration). In all these cases the products include oxides. Carbon dioxide is formed from the oxidation of carbon itself, and by the oxidation of fuels and foodstuffs. Iron oxides are formed during the rusting process. Corrosion. Whenever a metal corrodes, an oxidation reaction is taking place. Oxygen from the air is the commonest oxidising agent and this can be a problem especially with iron. E.g. Steel tank structures must be painted every few years if the iron, a more reactive metal, is to be prevented from rusting. If the corrosion is allowed to continue, the structure will be ruined in just a few years. See Kognity C9.5.1, (3rd Edition 212-213, 4th Edition 168-169)  25 Giant covalent structures There are some non-metals that bond covalently but do not form small molecules. They form giant crystal lattices and are said to have giant covalent structures or network lattices. These structures generally have high melting & boiling points, are insoluble in most solvents and are unreactive. Carbon provides the most common examples of such structures. It is able to arrange its atoms in two different ways (allotropy) to form two very different materials. Diamond Each carbon is surrounded by 4 other carbon atoms in a tetrahedral arrangement to form a giant cage-like lattice. Graphite Each carbon is bonded to three others in layers of hexagonal rings. The layers are held together by longer weaker bonds. A new allotrope was discovered in 1985 and given the name of Buckminsterfullerene or more commonly as ‘bucky balls’. It is possible that this allotrope will have many uses in the near future; e.g. catalysts, drug carriers Uses of graphite and diamond related to structure Diamond Graphite Jewellery electrodes Drill bits & lubricant Glass cutters Pencil leads  26 Metallic bonding This is the third type of bonding. Metal atoms are tightly packed together in a regular arrangement to form a crystal lattice. Metallic bonding can be described as ‘a regular arrangement of positive nuclei in a sea of mobile electrons’ diagram It is the ‘mobile electrons’ that give metals their distinctive properties. These are; Property Reason Electrical conductivity Malleability and Ductility Hardness Do not forget that there may be exceptions to the above properties. e.g. mercury is a liquid. Alloys are mixtures of metals in which the different sized atoms prevent sliding of layers. These different sized atoms prevent sliding of layers, thus making the structures stronger. Crystal structures Crystal structure is the result of the regular arrangement of particles. The particles may be atoms, ions or molecules. Ionic crystals can be illustrated by the arrangement of sodium and chloride ions in the crystal lattice of sodium chloride. Several ionic compounds have the same arrangement e.g. MgO Others have different arrangements. Covalent molecular crystals e.g. Water (see pages 54-55 3rd Edition, pages 38-39 4th Edition)  27 Summary of bonding & properties Structure---> Ionic Covalent molecular Giant covalent Metallic e.g. NaCl CH4 Diamond Cu CuSO4 CO2 and graphite Na atoms particles present ions molecules atoms (positive nuclei & mobile e's) covalent bonds ionic & weak forces between molecules covalent metallic Low Very high mod. high melting point & boiling point High (physical state) gases , volatile solids solids, some liquids solids liquids & solids solubility in water good with poor with none none (a polar solvent) exceptions exceptions solubility in non-polar poor good none none solvents (organic) poor in solid poor good when molten (may react with electrical conductivity & in aqueous water to form ions) none good in solid solution Questions on structure (answer on a separate sheet) 1. Give a short phrase that will explain the following: (a) Metals are good conductors of electricity (b) Metals generally have high densities. (c) Methane is a gas at room temperature. (d) Sodium chloride will only conduct electricity when molten or in solution. 2. Copper is a dense, ductile metal with a high resistance to corrosion. Aluminium bronze is an alloy of copper and aluminium. This alloy has the same resistance to corrosion as copper but has twice the strength. (a) The diagram alongside shows the arrangement of atoms in pure copper metal. Use the diagram to help you explain why copper metal (i) is dense (ii) is ductile. (b) (i) Using the symbols shown, draw a labelled diagram to show the arrangement of atoms in aluminium bronze. (ii) Use your completed diagram to explain why aluminium bronze is much stronger than pure copper. 3. Arrange the following substances in three groups, according to their structure: Group A - giant ionic structure Group B - giant covalent structure Group C - molecular structure Methane sulphur sodium chloride silicon dioxide oxygen Iodine graphite water copper sulphate ethanol Ice ammonia potassium chloride hydrogen chloride nitrogen  28 4. The table below gives information about some properties of certain substances A to G Substance Melting Electrical conductivity Solubility in water point oC solid liquid A -112 Poor Poor Insoluble B 680 Poor Good Soluble C -70 Poor Poor Insoluble D 1495 Good Good Insoluble E 610 Poor Good Soluble F 1610 Poor Poor Insoluble G 660 Good good Insoluble (a) Which of the substances are metals? Give reasons for your choice. (b) Which of the substances are ionic compounds? Give reasons for your choice. (c) Two of the substances have very low melting points, compared with the rest. Explain why these could not be ionic compounds. (d) Two of the substances are molecular. Which ones are they? (e) Which substance is a giant covalent structure? (f) Which substance would you expect to be very hard? (g) Which substances would you expect to be soluble in organic solvents?  29 Another way to work out the formula of an ionic compound 1- Metals 1+ Non-Metals Na, Cl, I, K, Li, Ag, Br, 2- 2+ Cu(I), NH4, NO3, OH H, O, 3- Cu (II) MN S, SO4, Mg, Fe(II), CO3, 3+ Zn, Ca PO4, Ba, M2N N, Al, MN2 P, Fe(III) MN M3N MN3 M3N2 M2N3 MN © A.J. Bleasdale