School Science Lessons
Topic 11 Chemical bonds, hydrogen bonds, van de Waals bonds, Deuterium, ions, atomic models, particles of matter, spectroscopy
2012-01-27 SP
Please send comments to: J.Elfick@uq.edu.au
Table of contents
3.01 Chemical bonds
3.2.0 Hydrogen bonding in liquids
3.6.0 Isotopes
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
3.3.0 Sizes of particles of matter
3.7.0 Spectroscopy
3.01.0 Chemical bonds
3.01.8 Bond energy, (bond strength)
3.01.2 Covalent bonds
3.01.3 Hydrogen bonds
3.01.6 Inter-molecular bonds
3.01.1 Ionic bonds, electrovalent bonds, e.g. sodium chloride.
3.01.5 Metallic bonds
3.01.4 Network solids
3.01.7 van de Waals forces (van de Waals bonds)
3.2.0 Hydrogen bonding in liquids
3.2.1 Liquids with different viscosity, hydrogen bonding
3.6.0 Isotopes
3.6.1 Deuterium
3.4.0 Movement of ions
3.4.4 Electric writing
33.3.9 Ionic migration
3.89 Movement of copper ions and chromate ions
3.4.3 Movement of copper ions in ammonium nitrate solution
3.90 Movement of ions, between microscope slides, Cu2+ ions, CO2+ ions
3.4.1 Movement of ions, potassium permanganate solution
3.4.2 Movement of ions, sodium sulfate solution
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
3.1.0 Properties of ionic solutions, conductivity of pure water and solutions
15.59.2 Electrical conductivity of solids
15.59.3 Electrical conductivity of melted solids, fused solids
15.59.4 Electrical conductivity of liquids
3.1.1 Measure electrical conductivity of different solutions
3.1.2 Measure conductivity of different concentrations of solutions
15.59.1 Substances that conduct electricity
3.3.0 Sizes of particles of matter
3.3.1 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
3.3.5 Particle movement, aluminium powder
3.57 Size of a molecule
3.3.4 Volume change when substances dissolve
3.7.0 Spectroscopy
4.134 Spectroscope, Diffraction grating,
36.101 Spectroscope for materials analysis, shoe box spectroscope
36.101.1 Spectroscope, Atomic absorption spectroscopy
27.1.0 Spectrum, Electromagnetic radiation, visible spectrum
3.7.1 Infra-red spectroscopy, IR
3.7.2 Mass spectrometry
3.7.3 Nuclear magnetic resonance, NMR
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.
3.01.1 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-.
Solids with ionic bonds have high melting points and low electrical conductivity.
However, when molten or in solution their electrical conductivity is high.
3.01.2 Covalent bonds
1. Non-polar covalent bonds exist between two or more, usually identical,
atoms, e.g. H:H. Each atom donates one electron to form a pair of electrons
that are equally shared by the two atoms. Non-polar covalent bonds (coordinate
covalent bonds), in ammonium ion, NH4+, have electrons
shared between the atoms and are not soluble in polar solvents, e.g. water,
but are soluble in non-polar solvents, e.g. hexane.
2. Polar covalent bonds form when atoms of two different elements each
donate one electron to form a shared pair of electrons, but one atom has a
greater share than the other atom. You could say that the shared electrons
spend more time around one atom than the other atom, so this kind of covalent
bond has a sort of ionic character. Polar covalent bonds exist as a dipole,
e.g. the H-Cl bond in hydrogen chloride. The H end has a small +ve charge
and the Cl end has a small -ve charge. Also, water, H2O, is a polar
molecule, a dipole with a definite positively charged end and a definite negatively
charged end, because each H has a small +ve charge, H+, and the
O has a small -ve charge, O2-. Dipole-dipole bonds occur between
polar molecules where the slightly positive end of one molecule is attracted
to the slightly negative end of another molecule. Solids with covalent bonds
within the molecules with van der Waals forces or weak dipole to dipole intermolecular
forces have low melting points and no electrical conductivity
3.01.3 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.
3.01.4 Network solids
Network solids have atoms in a lattice of strong covalent bonds, e.g.
C (diamond) Si, P (red phosphorus) B, and Ge.
3.01.5 Metallic bonds
In solid metals or alloys, the atoms exist in a lattice of positive ions
with electrons moving freely between them. Metals (metallic bonds) have
high melting points and 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.
3.01.6 Intermolecular bonds
Intermolecular bonds exist between molecules and intramolecular bonds
exist between the atoms of a molecule. Intermolecular bonds hold together
solids and liquids, but they are stronger in solids and cause the close packed
ordered arrangement of particles. Strong intramolecular bonds in liquids cause
high boiling points. Smaller molecules usually have lower boiling points.
3.01.7 van de Waals forces
(van de Waals bonds)
Although the water molecule is electrically neutral, the distribution
of charge is not symmetrical causing a net attraction between molecules
from momentary dipoles (cohesion of water molecules) that affects viscosity
and surface tension. Even non-polar molecules have some van der Waals bonding.
although nonpolar liquids have low surface tension and low boiling points.
3.01.8 Bond energy, (bond
strength)
The bond energy needed to break both of the two O-H bonds in water, H-O-H,
is 459 kJ mol-1.
Relative bond strengths
van der Waals bonds strengths X 10 = hydrogen bonds strengths X10 = non-polar
covalent bonds strengths, then covalent bonds strengths, then ionic bonds
strengths, then metallic bonds strengths.
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 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 3.2.57
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 10-15 cm2)
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 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 Movement of ions, 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 Movement of ions, 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 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 3.4.1: 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).
3. Commercial molecular models set, "Molymod", stick and ball, organic set, inorganic set
3.6.1 Deuterium
Deuterium, D, hydrogen isotope having a proton and a neutron in the nucleus
Deuterium oxide, D2O, water with H replaced by D
Deuterium exchange, organic compound ROH + D2O --> ROD +
DHO, used in NMR spectroscopy to identify position of H