School Science Lessons
Topic 11 Chemical bonds, ions, atoms, molecules, particles of matter
Updated 2009-09-17
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites
Table of contents
3.01 Chemical bonds
3.1.0 Properties of ionic solutions, conductivity of pure water and solutions
3.2.0 Hydrogen bonding in liquids
3.3.0 Sizes of particles of matter
3.4.0 Movement of ions
3.5.0 Models of atoms, molecules and ions
3.1.0 Properties of ionic solutions, conductivity of pure water and solutions
3.1.1 Measure electrical conductivity of different solutions
3.1.2 Measure conductivity of different concentrations of solutions
2.59 Electrical conductivity of solids
2.59.1 Electrical conductivity of melted solids
2.60 Electrical conductivity of liquids
3.2.0 Hydrogen bonding in liquids
3.2.1 Liquids with different viscosity,  hydrogen bonding
3.3.0 Sizes of particles of matter
3.3.1 Observe colour change of diluted potassium permanganate
3.3.2 Measure the size of an oil molecule in floating oil
3.3.3 Measure the size of the stearic acid molecule
3.3.3.1 Measure the size of a carbon atom in a stearic acid molecule
3.3.4 Observe volume change when substances dissolve
3.3.5 Particle movement, aluminium powder
3.4.0 Movement of ions
2.89 Movement of copper and chromate ions
2.90 Observe movement of coloured ions
3.4.1 Observe movement of ions from potassium permanganate solution
3.4.2 Observe movement of ions from sodium sulfate solution
3.4.3 Observe movement of copper ions in ammonium nitrate solution
3.4.4 Electric writing
3.4.5 Test paper to determine electric polarity
3.5.0 Models of atoms, molecules and ions
3.5.1 Make molecular models
Images of atoms
3.01 Chemical bonds
See diagram 3.01: Water molecules
All chemical bonds are caused by attractive force between positive and negative particles. The three main types of bonding are ionic, covalent and metallic bonding.
1. Ionic bonds, electrovalent bonds, e.g. sodium chloride.
Ionic bonds are electrical links between atoms caused by the distribution of electrons around the nuclei of the bonded atoms.
Ionic solids are usually transparent and have high melting points and boiling points. When an ionic solid, e.g. sodium chloride, NaCl, dissolves in water, the cations and anions separate, e.g. Na+ and Cl-.
2. Covalent bonding, e.g. O2
Non-polar covalent bonds exist between two or more atoms, e.g. H:H. Polar covalent bonds exist as a dipole, e.g. in H-Cl the H end has a small +ve charge and the Cl end has a small -ve charge. Water, H2O, is a polar molecule, a dipole,  because each H has a small +ve charge, H+,  and the O has a small -ve charge, O2-.
3. Hydrogen bonds
The hydrogen ion is the positive ion, H+, a proton, formed when a hydrogen atom loses its electron. The hydrogen ion may become solvated to form the hydrated ion H3O+, the hydronium ion.
Hydrogen bonds are weak bonds caused by the electrostatic attraction between strongly electronegative atoms and a hydrogen atom covalently linked to another electronegative atom. The polarized water molecule is a dipole with each H+ linked to the one O2- by a polarized covalent bond. The hydrogen bonds are inter-molecular bonds between polar molecules containing hydrogen. So, in water, the H+ of one water molecule is attracted to the O2-  of another water  molecule. Hydrogen bonds between water molecules only, cause the force of cohesion and surface tension. Hydrogen bonds between water molecules and other molecules, e.g. the molecules in glass, cause the force of adhesion.
4. Network solids
Network solids have atoms in a lattice of strong covalent bonds, e.g. C (diamond) Si, P (red phosphorus) B, and Ge.
4. Metallic bonding
In solid metals or alloys, the atoms exist in a lattice of positive ions with electrons moving freely between them. Metals are good electrical conductors because the outer electrons are held loosely. Metals are ductile and malleable because the metallic bonds are weak. However, alloys with other metals strengthen the lattice and so improve mechanical strength.
6. Inter-molecular bonds
Inter-molecular bonds exist between molecules and intra-molecular bonds exist between the atoms of a molecule. Inter-molecular bonds hold together solids and liquids, but they are stronger in solids and cause the close packed ordered arrangement of particles. Strong intra-molecular bonds in liquids cause high boiling points. Smaller molecules usually have lower boiling points.
3.1.1 Measure electrical conductivity of different solutions
1. Make a standard electrical conduction apparatus. Connect two carbon rods from used torch batteries to a light bulb and a source of direct current. Dip the carbon rods into the following solutions and record the brightness of the light bulb. Rinse the carbon rods with deionized water after each test. The brightness of the bulb is a measure of the conductivity of the solution.
2. Make more accurate measurements of conduction with an ammeter or a galvanometer. Record which solutions are good, fair, poor conductors, or are not conductors. Test demineralized water, deionized water, mineral water, and tap water. Test 0.2 M solutions of: aqueous ammonia, copper (II) sulfate, ethanoic acid (acetic acid) potassium hydroxide, sodium chloride, sodium hydroxide, and sucrose.
Aqueous ammonia solution, ethanoic acid (acetic acid) and sucrose solutions are poor conductors.
3.1.2 Measure conductivity of different concentrations of solutions
Use different concentrations of acids, bases and salts and compare the conductivity. More dilute solutions are better conductors. Test the conductivity of glacial ethanoic acid (acetic acid). Test again after adding different amounts of water.
3.2.1 Liquids with different viscosity,  hydrogen bonding
The viscosity's are different mainly because of difference in hydrogen bonding.
Pour 100 mL of the different liquids into separate conical flasks, e.g.: glycerol (glycerine, propane-1,2,3-triol) glycol, water, methylated spirit (ethanol) benzene (benzol) concentrated sulfuric acid 18 M. BE CAREFUL!
Swirl each flask equally and note the relative time until the disappearance of each vortex. Use this method to show that a mixture of two organic liquids is more viscous than the viscosity of either pure substance. This is caused by hydrogen bond formation.
3.3.1 Observe colour change of diluted potassium permanganate
Put one crystal of potassium permanganate in a test-tube. Add 1 mL of water. Dissolve the crystal completely by shaking vigorously with the thumb over the end of the test-tube. Then add water to a total volume of 10 mL. This is a "X 10" dilution. Pour this 10 mL of purple solution into a 100 mL beaker and then fill up the beaker with water. This is now a "X 100" dilution. Fill the 10 mL test-tube with this solution and discard the rest. Dilute this again in the beaker to 100 mL. This is now a "X 1000" dilution. Continue diluting the solution. Some purple colour is still visible. This shows that if matter consists of particles the particles must be very small.
3.3.2 Measure the size of an oil molecule in floating oil
See diagram 3.57
When the water has a large enough surface area, the oil spreads out in a layer one molecule thick because it does not form "hills" of molecules. When you know the volume of the oil and the area of the surface, you can calculate the thickness of the monomolecular layer.
Use a thin petroleum distillate or a pure vegetable oil and a flat tray containing water. Lightly sprinkle the surface of the water with talcum powder. Pour oil into a burette and measure the volume of fifty drops of oil. Let another drop of oil to fall on a piece of plastic. Pick up oil from the oil drop with the point of a glass rod and transfer it to the water surface by touching the surface lightly. This oil spreads out pushing out the talcum powder. Estimate the area of the oil by comparing it with a piece of graph paper. Estimate how much oil is removed from the oil drop by using the glass point to pick up equal amounts of oil from the oil drop until it is all removed. When you know the volume of the oil on the water, calculate the thickness of the oil layer. The volume of the oil drop = area of oil picked up by point X number of "pick ups" X depth of floating oil. The depth of floating oil = diameter of the water molecule. The diameter of the water molecule is in the range 10-6 to 10-7 mm.
3.3.3 Measure the size of the stearic acid molecule
Make a 0.3700 g / L to 0.3750 g / L solution of stearic acid in benzene (benzol).
BE CAREFUL! Benzene may be carcinogenic. Use safety glasses and nitrile chemical-resistant gloves! Make only enough solution for use.
Add the solution drop by drop to the surface of water in a Petri dish. The dropper must have an outlet small enough so that there are more than 50 drops to make a total of 1 mL of the benzene solution.
The benzene evaporates and leaves the stearic acid to spread over the entire water surface. The hydroxyl group of the organic acid molecule sticks into the water because of its hydrophilic property. The hydrocarbon portion of the molecule resists entering the water surface because of its hydrophobic property. This behaviour leads to a monomolecular film in which all the stearic acid molecules are compactly and fully arranged on the water surface. The effective section area A cm2 of each stearic acid molecule can be calculated by using the following equation: A = M X S X V / m X NA X Vd X (d-1)
M = molar weight of stearic acid, CH3(CH2)16COOH (284 g / mol) S = total area of monomolecular film, m / V = concentration of stearic acid in benzene solution (0.3700-0.3750 g / L) NA = Avogadro's Constant (L) (6.022 X 1023 mol-1) Vd = volume of a drop of stearic acid solution, d = number of drops of stearic acid solution.
Before each determination, it is necessary to wash the Petri dish by using sodium carbonate solution and then tap water, finally distilled water 2-3 times to clear off stearic acid and base. Repeated determination should use the same Petri dish. Control dropping evenly to make the volume of every drop be the same.
Determine S: Measure the inside diameter of the Petri dish to be used (~200 mm) in three directions with inside callipers (or a straight edge). Take the average value and calculate the area. (pi X r2)
Determine Vd: Use a dropper to add benzene (benzol) to a dry, small graduated cylinder until the volume reaches 1.0 mL, and count the drops. Calculate the Vd.
Determine (d-1): Use another dropper to add the stearic acid in benzene (benzol) solution vertically and gently at a point above the Petri dish and 1-2 cm from the water surface. Wait for a fallen drop to diffuse until oil beads cannot be seen, and then give another drop. Diffusion of the stearic acid solution gradually slows. When a fallen drop cannot diffuse any more in two minutes, a monomolecular film has been formed, looking like a lens. Record the number of drops, d, from which subtract the final drop to obtain the actual drop number needed to form the monomolecular film of stearic acid, i.e. d-1. Repeat the determination and get the average value. Calculate A according to the equation mentioned above. (A = 2.2 X 10-15cm2)
3.3.3.1 Measure the size of a carbon atom in a stearic acid molecule
Fill a 14 cm diameter watch glass with water and measure the diameter of the water surface. Add drops of 0.10 g /litre stearic acid solution in hexane  to  the water until one drop sits on the water, i.e. the drop extra to the monolayer of stearic acid in the water. Calculate the mass of pure stearic acid in the total number of drops of 0.10 g /litre stearic acid solution in hexane in the monolayer after evaporation of the hexane. The density of solid stearic acid is 0.85 g / mL.  The area of the monolayer = pi X radius of water surface2. The thickness of the monolayer, L = volume of stearic acid / area of water surface = length of the stearic acid molecule. If the stearic acid molecule contains 18 closely packed carbon atoms in vertical stacks, the diameter of a carbon atom = L / 18, and the volume of a carbon atom = (L / 18)2. Compare your result with the accurately measured diameter of a carbon atom = 1.54 X 10-10 m.
3.3.4 Observe volume change when substances dissolve
Sodium chloride is an ionic solid that exists as ions. The ionic structure breaks down in solution. These ions can slip between the water molecules and make the total volume decrease. Other salts that dissolve in water have the same property.
Put sodium chloride crystals or another soluble salt in a test-tube. Add water until full. The water level drops slightly as the crystals dissolve. Repeat the experiment with sugar. The water level does not drop because the sugar molecule is larger than water molecules and does not form ions. Half fill a test-tube with water. Hold the test-tube at an angle and completely fill it by adding alcohol. Put the thumb over the mouth of the test-tube so that no air bubble is trapped below it. Invert the test-tube several times keeping the thumb over the opening. The level of the liquid is now lower because spaces exist between the water molecules that the alcohol molecules can enter.
3.3.5 Particle movement, aluminium powder
Add a small amount of aluminium powder to a beaker of the treated water. If you add a few drops of detergent to the water, it will clean the very small pieces of aluminium and prevent them from sticking together. The pieces probably had oils or grease on them. You may need to stir the mixture a little to make the aluminium powder mix freely with the water. Now darken the room as much as possible and shine the light of a film projector through the liquid. Watch the pieces of suspended aluminium powder from the side. If you are patient, as your eyes become accustomed to the darkness you will see that the larger suspended aluminium particles will appear not to change but the smallest ones will twinkle like tiny stars.
3.4.1 Observe movement of ions from potassium permanganate solution
See diagram 3.4.1:  Potassium permanganate ions
1. Bind a strip of damp filter paper to a test-tube with two pieces of copper wire. Do not leave any air bubbles between the filter paper and the wall of the test-tube. Fix the test-tube horizontally on an iron stand. Place a length of cotton thread that you have dipped into potassium permanganate solution round the middle of the filter paper. Connect the two pieces of copper wire to a 12 volts or more electric source of direct current. A colour of purplish red can be seen to leave the cotton thread, moving gradually towards the positive terminal.
2. Cut a 1 cm wide strip of dry filter paper. Draw a pencil mark across the centre of the paper. Moisten with tap water so that it is damp, but not very wet. Make a potassium permanganate solution. Use a fine capillary tube to put the coloured ion solution along the pencil mark. Hold the strip of filter paper between two microscope slides. Attach carbon electrodes across the slides to lead the current through the filter paper. Use 12 volts or more direct current.
3.4.2 Observe movement of ions from sodium sulfate solution
Dissolve sodium sulfate in water and add drops of Universal Indicator. The solution should be green showing that the solution is neutral. Dissolve 1 g of powder agar or agar gel in 100 mL of hot water.
Mix the two solutions then pour into an U-tube so that the arms are half full. When the gel has set, pour dilute sulfuric acid into one arm. and dilute sodium hydroxide into the other arm. Put platinum or carbon electrodes into the solutions. Connect the terminal in the sulfuric acid to the positive terminal of a battery. Connect the terminal in the sodium hydroxide to the negative terminal of a battery. Turn on the electric current and note the colour changes at each arm. The violet colour in the gel below the sodium hydroxide solution is due to movement of hydroxide ions into the gel. The red colour in the gel below the sulfuric acid solution is because of the movement of hydrogen ions into the gel. During electrolysis, there is a flow of ions in opposite directions caused by the electric field.
3.4.3 Observe movement of copper ions in ammonium nitrate solution
1. Dissolve 10 g of ammonium nitrate in 125 mL of water. Add 10 mL of concentrated aqueous ammonia solution. Dip a piece of filter paper in this solution of ammonium nitrate and then twine the filter paper round a test-tube by using two pieces of copper wire. No air bubbles should be left between the filter paper and the wall of the test-tube. Connect the copper wires to a source of direct current and apply a voltage of 12 V. You can observe that a greenish blue colour moved from the positive terminal to the negative one.
2. Dissolve 10 g of ammonium nitrate in 125 mL water Add 10 mL concentrated aqueous ammonia solution. Pour the solution into a Petri dish or beaker. Use a 12 volt battery. Use crocodile clips to attach the positive pole to a piece of copper gauze (mesh) and attach the negative pole to a strip of copper. Put the copper electrode 5 cm apart in the solution. Observe the blue streaks of copper ions moving from the +ve electrode towards the -ve electrode.
3.4.4 Electric writing
See diagram 3.4.4: Electric writing
When an aqueous solution of table salt is electrolysed, the sodium hydroxide solution produced around the cathode can make red litmus paper turn blue. The chlorine gas formed at the anode has a bleaching effect. [Comment: The equation is not really necessary]
2NaCl (aq) + 2H2O ---> 2NaOH (aq) + H2 (g) + Cl2 (g)
Soak a piece of red litmus paper in dilute sodium chloride solution and fix it on a sheet of glass. Connect the terminals of a 6 volt battery to pencil "leads". Touch the paper with the leads. Write in blue with the lead attached to the +ve terminal. Write in white with the lead attached to the negative terminal.
3.4.5 Test paper to determine electric polarity
Moisten filter paper with 1% solution of phenolphthalein in alcohol solution and and leave to dry. Dip the filter paper in 10% potassium chloride solution in deionized water. Wet the filter paper and apply it to the terminals. The negative side turns pink.
3.5.1 Make molecular models
The physical and chemical properties of chemical substances are largely influenced by their shape, e.g. the polarity of a molecule and subsequent intermolecular forces. This can be demonstrated by the construction of chemical models requiring the use of geometric skills. Use models to show the different kinds of bonding between atoms to form molecules and ions.
1. Use polystyrene spheres to represent ions in the sodium chloride crystal lattice. Cut holes in a plate to take the bottom layer of spheres then construct layers of spheres. Join the spheres with small springs to show how the ions may vibrate about mean positions.
2. Use corks to represent atoms or ions and use hooks or pins to represent bonds.
Construct ionic crystals: halite NaCl, calcite (CaCO3).
Construct covalent crystals: diamond C, quartz SiO2, ice H2O, sulfur S8
Construct organic molecules: PVC (polyvinyl chloride, poly-chloroethane) (CH2Cl.CH3)n, polypropylene (-CH3.CH.CH2-)2, polystyrene (-C6H5.CH.CH2-)n, nylon-66, -NH2(CH2)6NH2CO(CH2)CO-}n.
Construct biochemical models: (+) glucose (C6H12O6) sucrose (C12H22O11) starch / cellulose (C6H11O5)n, fat (C57H110O6, tripalmitin) polypeptide (-CH[CH3]CO.NH-)n, DNA (deoxyribonucleic acid).