School Science Lessons
Topic 11 Please send comments to: J.Elfick@uq.edu.au

11.0 Chemical bonds, conservation of mass
Table of contents
3.01.0 Chemical bonds

3.8.0 Fatty alcohols, lauryl alcohol

12.0.0 Hydrochloric acid, HCl

12.10.2 Prepare salts by different methods

13.0.0 Sulfuric acid
Experiments
12.1.0 Conservation of mass

3.5.1 Molecular models

11.2.0 Movement of ions and particles in solution

11.3.0 Size of molecules and particles

3.4.4 Electric writing, NaCl with litmus paper

3.4.5 Tests for electric polarity

12.1.0 Conservation of mass
12.1.0 Chemical equations and ionic equations, conservation of mass
Experiments
12.1.2 Burn steel wool, change in weight
13.3.3 Burn steel wool and burn iron filings
13.3.2 Burn sulfur in oxygen
12.1.1 Effervescent tablets, health salts, sodium bicarbonate (baking soda)
12.1.3 Copper cycle reactions

3.01.0 Chemical bonds
See: Chemicals (Commercial)
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 Intermolecular bonds
3.01.1 Ionic bonds, electrovalent bonds, e.g. sodium chloride
3.2.1 Liquids with different viscosity, hydrogen bonds
3.01.5 Metallic bonds
3.01.4 Network solids
3.01.7 van de Waals forces (van de Waals bonds)

11.2.0 Movement of ions and particles in solution
Experiments
3.3.5 Movement of suspended aluminium powder
3.4.1 Movement of ions, potassium permanganate solution
3.4.2 Movement of ions, sodium sulfate solution
3.4.3 Movement of copper ions in ammonium nitrate solution

11.3.0 Size of molecules and particles
See: Atoms and Molecules (Commercial)
Experiments
3.3.1 Colour change of diluted potassium permanganate
3.3.2 Size of oil molecule
3.3.3 Size of stearic acid molecule
3.3.3.1 Size of carbon atom in stearic acid molecule
7.8.1.2 Size of colloidal particles
3.91 Size of particles and rate of reaction with balloons
3.3.4 Volume change when substances dissolve

3.01.0 Chemical bonds
See diagram 3.01: Water molecules
1. Chemical bonds are forces of attraction between the atoms together in a molecule or a crystal.
All chemical bonds are caused by attractive force between positive and negative particles.
2. 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, electrovalent bonds, form by transfer of electrons
Ionic bonding occurs when positive and negative ions are held together in a crystal lattice by electrostatic forces.
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. Covalent bonds form by sharing of electrons.
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.
An atom or group of atoms covalently bound together may gain or lose one or more electrons to form ions
2. 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 (co-ordinate 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.
3. 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.
Some people 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.
Dipole-dipole forces are intermolecular forces between molecule dipoles at δ + and δ - areas.
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.
5. Molecules may be saturated, i.e. have only single bonds, --, or unsaturated, i.e. contain multiple bonds, e.g. double bond =, or triple
bond, e.g. carbon monoxide, CO, C≡O.

3.01.3 Hydrogen bonds
1. 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.
In water, each hydrogen nucleus is bound to the central oxygen atom by a pair of electrons shared between them, a covalent
chemical bond.
Two of the six outer shell electrons of oxygen are used, leaving four electrons as two non-bonding pairs.
The four electron pairs around the oxygen are as far apart as possible to reduce repulsion between their negative charges.
However, the two non-bonding pairs stay closer to the oxygen atom and exert a stronger repulsion against the two covalent bonding
pairs, and so pushing the two hydrogen atoms closer together.
These forces cause the H-O-H angle, 104.5o.
The positive and negative charges are not distributed uniformly.
The negative charge is concentrated at the oxygen end of the molecule because of the non-bonding electrons and oxygen's high nuclear
positive charge.
This charge displacement causes an electric dipole.
The partially-positive hydrogen atom on one water molecule is electrostatically attracted to the partial negative oxygen on a neighbouring
molecule by a process is called hydrogen bonding.
That the hydrogen bond is weak.
2. Intermolecular force between strongly electronegative atom, e.g. F, O, N, and a H atom covalently bonded to an electronegative atom.
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.
Electronegativity refers to an atom attracting to itself electrons in covalent bonds.

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 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)
Van der Waals attraction and repulsion forces between atoms and between molecules, are caused by fluctuating polarization of atomic
particles and are quite different from the forces of covalent and ionic bonding.
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 × 10 = hydrogen bonds strengths × 10 = non-polar covalent bonds strengths, then covalent bonds
strengths, then ionic bonds strengths, then metallic bonds strengths.

3.2.1 Liquids with different viscosity, hydrogen bonds
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 Size of oil molecule
See diagram 3.2.57: Oil layer and powder on surface of water
Oil floats on the surface of water as a one molecule thick layer, if allowed to spread out freely.
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.
Knowing the volume of the oil and the area of the surface, calculate the thickness of the monomolecular layer.

1. 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.
Knowing 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 × number of "pick ups" × 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 10-7mm.

2. Oil floats on the surface of water as a one molecule thick layer, if allowed to spread out freely.
Use a tray with a glass bottom.
Put graph paper under the tray.
Put water in the tray.
Select a suitable oil to pour on the water, e.g. 1% oleic acid / methylated spirit solution, petroleum distillate.
Use a burette to find the volume of 100 drops of oil, then calculate the volume of 1 drop of oil.
Sprinkle the surface of the water with a very fine light powder, e.g. talcum powder or Lycopodium powder (the "dragon's breath"
used in Chinese fireworks).
Let 1 drop of oil fall from the burette onto the water.
The oil spreads out over the surface of the water but must not touch the sides of the tray.
The oil on the water pushes the powder aside so you can easily see the area covered by the oil.
Look down on the graph paper to measure the approximate area over which it spreads.
The volume of oil put on the water = area of the oil on the water X thickness of the oil layer.
The approximate dimension of a single molecule of oil is 1 X 10-7 m.

3. Another way to calculate the volume of a drop of oil is to let the drop fall on a piece of flat plastic.
Touch the oil with the point of a glass rod then touch the water surface.
Oil leaves the glass rod and spreads over the water.
Use graph paper to measure the approximate area over which the oil spreads.
To calculate the volume of oil placed on the water, keep using the glass rod to remove successive fractions from the flat plastic until no
oil remains.

3.3.3 Size of stearic acid molecule
See: Models, biochemistry Stearic acid, (Commercial)
Stearic acid has the molecular formula, C18H36O2.
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.
M
ake only enough solution for immediate 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 × S × V / m × NA × Vd × (d-1),
where 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 × 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 calipers (or a straight edge).
Take the average value and calculate the area (π × 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 × 10-15 cm2)

3.3.3.1 Size of carbon atom in 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 = π × 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 the experimental result with the accurately measured diameter of a carbon atom = 1.54 × 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 Movement of suspended aluminium powder
Add a small amount of aluminium powder to a beaker of the treated water.
Add a few drops of detergent to the water to clean the very small pieces of aluminium and prevent them from sticking together.
The pieces probably had oils or grease on them.
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.
As the eyes become accustomed to the darkness, 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 previously 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 12 volt 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 volts.
Note a greenish blue colour moving from the positive terminal to the negative terminal.
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, sodium chloride with litmus paper
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.
2NaCl (aq) + 2H2O ---> 2NaOH (aq) + H2 (g) + Cl2 (g)
Experiment
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 Tests for electric polarity, potassium chloride with phenolphthalein
Moisten filter paper with 1% solution of phenolphthalein in alcohol solution 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 Molecular models
See: Models (Commercial)
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.
2.1 Construct ionic crystals: halite NaCl, calcite (CaCO3).
2.2 Construct covalent crystals: diamond C, quartz SiO2, ice H2O, sulfur S8
2.3 Construct alkanes: methane CH4, ethane C2H6, propane C3H8, butane C4H10, 2-methylpropane C4H10, [CH3.CH.CH3.CH3]
2.4 Construct alkanoic acids: methanoic acid HCOOH, ethanoic acid CH3COOH, propanoic acid C2H5COOH, butanoic acid,
(butyric acid), C3H7COOH, [CH3CH2CH2.COOH], 2-methylpropanoic acid (isobutyric acid) C4H8O2, [(CH3)2-CH-COOH]
2.5 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
2.6 Construct biochemical models: (+) glucose (C6H12O6) sucrose (C12H22O11) starch / cellulose (C6H11O5)n, fat (C57H110O6,
tripalmitin) polypeptide (-CH[CH3]CO.NH-)n, DNA (deoxyribonucleic acid).
See 4.4.0 Structure of DNA and RNA
3. Molecular models set, e.g. "Molymod", stick and ball, organic set, inorganic set (toy product)
3.1 The plastic straw, represents a single covalent bond or a "C=O" double bond.
3.2 The flexible straw, 2 flexible straws represent a "C=C" double bond.
3.3 The 4-pronged tetrahedral carbon centre (black), represents C in "-C-".
3.4 The 3-pronged trigonal planar carbon centre (black), represents C in ">C=O".
3.5 The one-pronged hydrogen centre (white), represents "H".
3.6 The one-pronged oxygen centre (red), represents O in ">C=O".
3.7 The 2-pronged V-shaped oxygen centre (red), represents O in "-O-H".

3.8.0 Fatty alcohols, lauryl alcohol
Fatty alcohols are aliphatic alcohols from natural fats and oils in plants and animals.
They have an even number of carbon atoms.
The -OH group is attached to the terminal carbon.
They are amphipatic (having both polar (water-soluble) and nonpolar (not water-soluble) portions in the structure of the molecule so
they can act as non-ionic surfactants.
Fatty alcohols are used in cosmetic formulations as emulsifiers, emollients and thickeners.
The most commonly used fatty alcohols are lauryl alcohol, cetyl alcohol, stearyl alcohol, cetostearyl alcohol and behenyl alcohol.
Lauryl alcohol, 1-dodecanol, C12H26O, CH3(CH2)10CH2OH, from coconut oil fatty acids, is tasteless and colourless but has a
flowery smell.

List of some fatty alcohols
1-heptacosanol
1-docosanol, C22H46O, behenyl alcohol, saturated fatty alcohol, emollient, emulsifyer, antiviral, topical treatment of herpes simplex
infections, side effects include headache, skin irritation
1-dotriacontanol
1-nonacosanol
2-ethyl hexanol
3-Methyl-3-pentanol
Arachidyl alcohol, 1-eicosanol
Behenyl alcohol, 1-docosanol
Capric alcohol, 1-decanol, decyl alcohol
Capryl alcohol, 1-octanol
Ceryl alcohol, 1-hexacosanol
Cetearyl alcohol
Cetyl alcohol, 1-hexadecanol
Erucyl alcohol, cis-13-docosen-1-ol unsaturated
Ethchlorvynol
Geddyl alcohol, 1-tetratriacontanol
Heneicosyl alcohol, 1-heneicosanol
Heptadecyl alcohol, 1-n-heptadecanol, heptadecanol
Lauryl alcohol, dodecanol, 1-dodecanol
Lignoceryl alcohol, 1-tetracosanol
Montanyl alcohol, cluytyl alcohol, 1-octacosanol
Myricyl alcohol, melissyl alcohol, 1-triacontanol
Myristyl alcohol, 1-tetradecanol
Nonadecyl alcohol, 1-nonadecanol
Palmitoleyl alcohol, cis-9-hexadecen-1-ol, unsaturated
Pelargonic alcohol, 1-nonanol
Pentadecyl alcohol, 1-pentadecanol, pentadecanol
Stearyl alcohol, 1-octadecanol
tert-Amyl alcohol
tert-Butyl alcohol
Tridecyl alcohol, 1-tridecanol, tridecanol, isotridecanol
Undecyl alcohol, 1-undecanol, undecanol, Hendecanol

12.1.0 Chemical equations and ionic equations, conservation of mass
Chemical reactions are influenced by the conditions under which they take place and, being reversible, may reach a state of equilibrium.
Chemical reactions occur at different rates and changing the nature of the reactants, temperature, or concentration, or introducing a
catalyst, may alter these.
Chemical reactions may be reversible.
Reversible chemical reactions may reach a state of dynamic balance known as equilibrium, which, when disturbed, will be re-established.
Specific criteria can be used to classify chemical reactions.
Redox reactions involve a transfer of electrons and a change in oxidation number.
Precipitation reactions result in the appearance of a solid from reactants in aqueous solution.
Acid-base reactions involve transfer of protons from donors to acceptors.
Polymerization reactions produce large molecules with repeating units.
Chemical reactions involve energy changes.
All chemical reactions involve energy transformations.
The spontaneous directions of chemical reactions are towards lower energy and greater randomness.
Chemical reactions, but not nuclear reactions, obey the law of conservation of mass.
The mass of the products is equal to the mass of the reactants.
In a balanced chemical equation, the number of atoms of each element in the reactants is equal to the number of atoms in the product.
Ionic equations show soluble ionic compounds in solution as separate ions and the sum of charges on the product side is equal to the
sum of charges on the reactant side.
Chemical equation: HCl (aq) + NaOH (aq) --> NaCl (aq) + H2O (l)
Ionic equation: H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) --> H2O (l) + Cl- (aq) + Na+ (aq)
The Cl- and Na+ ions are called spectator ions because they are on both sides of the equation and do not react.
Cancel the spectator ions to leave the net ionic equation: H+ + OH- --> H2O

12.1.1 Effervescent tablets, health salts, sodium bicarbonate (baking soda)
See diagram 12.1.1: Conservation of matter | 13.7.7 Prepare carbon dioxide, heat hydrogen carbonates
See 16.7.13: Prepare fruit salts, health salts
When a chemical reaction occurs, matter is neither created nor destroyed.
The mass of the reactants = the mass of the products.
1. Effervescent tablets or "fruit salts" contain sodium hydrogen carbonate, and dry citric acid or tartaric acid.
Put the tablet or fruit salts in water in a test-tube.
Carbon dioxide forms as bubbles and any other substance in the tablet or fruit salts dissolves easily.
Tests for carbon dioxide by holding a test-tube containing lime water at an angle near the mouth of the test-tube containing the
effervescent tablet.
The carbon dioxide given off by effervescence is heavier than air and will roll into the lime water test-tube where the lime water will turn
milky because of the presence of carbon dioxide.
Weigh an effervescent tablet or fruit salts and put into the bottom corner of a small plastic bag.
Twist the bag above the corner and tie around the twist with thin string or wire.
Weigh the plastic bag + string + effervescent substance.
Pour a known amount of water into the open bag then tie string tightly to close the bag so that no liquid or gas can escape.
The weight of the water + bag + string + effervescent substance is now known.
Undo the string around the twisted part of the bag and untwist the bag.
The acid and sodium hydrogen carbonate dissolves in the water and react to produce a salt and carbon dioxide.
Weigh the bag and products of the reaction.
The weight is the same.

2. Place a small amount of water in a plastic cup and place on the pan of the electronic scale.
Cover the top with a piece of paper or metal foil.
Place two Alka-Seltzer tablets (ant-acid tablets), on top of the cover.
Record the initial mass.
Tilt the cover so that the tablets drop into the water and immediately replace the cover so water droplets will not escape.
Record the mass reading until it is constant.
Repeat this experiment with the cup enclosed in a sealed container.

3. Place a small amount of water in a large 2 litre plastic drink bottle.
Break two Alka-Seltzer tablets in pieces that will fit in the bottle.
Weigh the bottle, bottle cap, and Alka-Seltzer tablets together.
Drop the Alka-Seltzer tablet pieces into the bottle and quickly replace the bottle cap tightly and place back on the scale.
Record the mass reading until it is constant.
The results differ from the experiment done in an open container.
Remove the bottle from the scale, loosen the bottle cap, and measure again the weight of the bottle and its cap.

12.1.2 Burn steel wool, change in weight
Roll some steel wool into a ball in your hands and weigh it.
Hold it with tongs over a sheet of paper.
Heat the steel wool until red hot; remove the flame, and blow gently on the red hot steel wool until it stops burning.
When cold, weigh the steel wool and any fragments fallen onto the sheet of paper, on the balance.
When iron burns, the product formed, iron oxide, is heavier than the iron.
4 Fe (s) + 3 O2 (g) --> 2 Fe2O3 (s)

12.1.3 Copper cycle reactions
See diagram 1.13a: Simple fume hood
Step 1. Convert copper metal to copper nitrate
1. Weigh 1.000 g of copper wire.
It must be clean, bright and shiny.
Twist the wire into a flat spiral and put it in a beaker in a fume hood.
Slowly add 4 .0 mL of 16 M nitric acid. Be careful!
Note the brown fumes of nitric oxide (NO) nitrogen dioxide (NO2) and dinitrogen tetroxide (N2O4), i.e. NOx.
When all the copper is dissolved, add 100 mL of deionized water.
Cu (s) + 8HNO3 (aq) --> 3Cu(NO3)2 (aq) + 4H2O (l) + 2NO (g) [dilute nitric acid]
Cu (s) + 4HNO3 (aq) --> Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g) [concentrated nitric acid]
N2O4 <--> 2NO2 [in equilibrium]
Step 2 Convert copper nitrate to copper hydroxide
Add 30.0 mL of 3.0 M sodium hydroxide to the solution while stirring.
f red litmus paper does not turn blue in the solution, add more sodium hydroxide.
Note the precipitate of copper hydroxide, an ionic solid.
Cu(NO3)2 (aq) + 2NaOH (aq) --> Cu(OH)2 (s) + 2NaNO3 (aq)
Step 3 Convert copper hydroxide to copper oxide
Heat the solution on a hot plate while continually stirring to prevent bumping caused by steam bubbles.
Note the precipitate changing to a black solid, copper oxide, CuO.
Carefully decant the liquid, add deionized water and decant again.
Heat the precipitate until it becomes a firm mass.
Cu(OH)2 (s) + heat --> CuO (s) + H2O (l)
Step 4 Convert copper oxide to copper sulfate
Reaction 4: Converting copper oxide to copper sulfate
Add 15 mL of 6.0 M sulfuric acid to the copper oxide while swirling, not stirring, the copper oxide to help it dissolve.
CuO (s) + H2SO4 (aq) --> CuSO4 (aq) + H2O (l)
Step 5 Convert copper sulfate to copper metal
CuSO4 (aq) + Zn (s) --> ZnSO4 (aq) + Cu (s)
In the fume hood add 2.0 g of zinc metal and keep stirring until the solution becomes colourless.
The zinc is oxidized as it reduces the copper.
Add one drop of the solution to 1 mL of concentrated ammonia solution in a test-tube.
If the ammonia turns deep blue, some unreduced copper is till in the solution so the reaction is not finished.
When all the copper is reduced, decant the liquid and add 20.0 mL of 6M hydrochloric acid to dissolve excess zinc.
Note the bubbles of hydrogen gas until the reaction is complete.
Cool the beaker and observe the metallic copper settling on the bottom.
Carefully decant the solution and wash the copper metal with deionized water and decant again.
Transfer the copper to an evaporating dish using a small amount of deionized water.
Decant excess water evaporating dish then wash the precipitate with methanol.
Decant the liquid the gently heat the precipitate on a hot plate.
If you heat the copper precipitate too strongly, it will oxidize to copper oxide.
Transfer the dried copper metal to a preweighed beaker and calculate the mass of recovered copper.
Percentage recovery = mass of recovered copper / initial mass of copper × 100.

12.0.0 Hydrochloric acid, HCl
Hydrochloric acid, ACS reagent, 37% | Hydrogen chloride., ≥99.8% Hydrochloric acid is a solution of hydrogen chloride
in water.
Hydrogen chloride occurs as either a colourless liquid with a an irritating, pungent odour, or a colourless to slightly yellow gas which can
be shipped as a liquefied compressed gas; highly soluble in water.
Hydrochloric acid, HCl (aq). spirits of salts, muriatic acid, clear, colourless, fuming liquid, conc. 12 M, r.d. 1.18 gm cm-3, 12 mol dm-3,
E507, dilute HCl 6 M (solution of hydrogen chloride gas), distinct odour,
(aqua regia: 3 vols concentrated HCl + 1 vol concentrated HNO3).
Concentrated hydrochloric acid releases hydrogen chloride gas if open to the atmosphere.
This gas is highly irritant to the lungs and causes coughing.
Use the concentrated gas in a fume cupboard to avoid inhalation.
If a container of the concentrated acid is broken outside a fume cupboard, evacuate the area until the fumes have dissipated.
Small amounts of hydrogen chloride gas are relatively harmless and a <0.1 M dilute acid has low toxicity.
The human stomach forms hydrochloric acid as part of the digestive process.
Concentrated hydrochloric acid is extremely damaging to the eyes.
Wear a face shield when doing any pouring experiments that might result in splashing of the liquid.
Avoid skin contact with hydrochloric acid.
Wash any area of the body that has been in contact with hydrochloric acid with copious amounts of water.
Dilution of hydrochloric acid with water does releases some heat, but not as much as sulfuric acid.
Always dilute by adding acid to water, not water to acid.
Do not mix hydrochloric acid with formaldehyde solution because highly carcinogenic bis(chloromethyl) ether may form.
Do not store bottles of hydrochloric acid near formaldehyde because the reaction might occur in the air.
Do not dispose of formaldehyde solutions and hydrochloric acid in the same sink system.
Common name: Muriatic acid (25% solution) (in concrete bleach, silverware cleaning solutions, drain cleaners.

13.0.0 Sulfuric acid, H2SO4, oleum, fuming sulfuric acid, vitriol or oil of vitriol is concentrated H2SO4, clear, colourless, heavy oily
liquid, very hygroscopic, conc. 18 M, r.d. about 1.84, 18 mol per mL, r.d. fuming H2SO4: 1.92 gm cm-3 (for accumulators), battery
acid, r.d. 1.25 gm cm-3, E513, dilute H2SO4, 3 M, store in cool place.
Pure sulfuric acid is called concentrated sulfuric acid.
Sulfuric acid reacts violently with water to form hydrated hydrogen ions and hydrated sulfate ions.
So the concept of dilution with water is inappropriate for the concentrated acid because a vigorous chemical reaction occurs that releases
much heat.
Do not add water to concentrated sulfuric acid.
Dilute sulfuric acid by adding the acid slowly to water with continuous stirring.
Wear eye protection when handling concentrated sulfuric acid.
Water is less dense than sulfuric acid and will layer itself on top of the acid.
At the acid / water interface, heat is generated by reaction of acid with water.
This causes the water just above the interface to become very hot.
When a suitable nucleating particle or surface is found, a large steam bubble forms, the middle section of the container boils suddenly
and hot acid / water mixture is ejected from the container, often travelling several metres and causing serious chemical burns.
If, the acid is slowly poured into water with stirring, the same amount of heat is liberated, but the boiling point of the solution is much
higher than 100oC because of the presence of the sulfuric acid.
The temperature of the solution may rise to 150oC, but boiling still does not occur.
Sulfuric acid is highly corrosive to the skin and eyes.
Wash sulfuric acid off the skin immediately with a large volume of water.
A small amount of water simply brings about the acid water reaction and liberation of a great amount of heat, causing the remaining
sulfuric acid to be even more corrosive.
Hot sulfuric acid is a strong dehydrating agent and will cause organic materials, e.g. skin, to turn black because of the removal of water,
leaving a residue of carbon.
Sulfuric acid is not volatile, so neither the pure acid nor its solutions release vapour.
However, strong heating of the acid may cause decomposition and boiling to form an aerosol of sulfuric acid droplets and sulfur dioxide
gas that must not be inhaled.
Sulfuric acid forms a spontaneously explosive mixture with potassium permanganate to form liquid manganese heptoxide.
Similarly sulfuric acid forms a spontaneously explosive mixture with chlorate salts where gaseous chlorine dioxide forms.
Dilute sulfuric acid solutions are far less dangerous than the concentrated acid because they do not dehydrate or reducing materials and
do not form much heat when further diluted.
Use sulfuric acid from motor car batteries if it is first diluted with water to four times its volume (Remember: acid to water!), but sulfuric
acid collected from lead-acid batteries will be contaminated with lead.
Poisonous and corrosive if left on clothes.
In many experiments, a solution of sodium hydrogen sulfate (sodium bisulfate), can be used instead of sulfuric acid.
Sulfuric acid dissolves most metals but not copper.
dissolves metal oxides, including copper oxide, is neutralized by alkalis, e.g. ammonia solution and forms carbon dioxide gas from metal
carbonates and bicarbonates.
Sulfuric acid was formerly manufactured by heating iron (II) sulfate crystals, so it was called "oil of vitriol", a name that is still used for it
today.

13.3.2 Burn sulfur in oxygen
Dip a wire loop into sulfur powder. Ignite the sulfur in a burner flame and then put it into a test-tube of oxygen.
The sulfur burns with a bright blue flame to form the colourless gas sulfur dioxide.
S (s) + O2 (g) --> SO2
sulfur + oxygen --> sulfur dioxide
Some sulfur trioxide may also form in this reaction.

13.3.3 Burn steel wool and burn iron filings
1. Collect oxygen in test-tubes with stoppers.
Store test-tubes in a test-tube rack and remove the stoppers just before inserting the burning element.
Fasten steel wool to wire.
Heat the steel wool in a burner flame.
Put it into a test-tube of oxygen.
The steel wool burns with bright sparkles to form black-grey iron oxide.
4 Fe (s) + 3 O2 (g) --> 2 Fe2O3 (s)
iron + oxygen --> iron oxide
2. Repeat the experiment by placing iron filings in a sieve and shaking it over a Bunsen burner flame.
A shower of sparks occurs as in some fireworks.
3. Hold an iron nail against a grinding wheel.
Friction breaks the iron into little pieces and heats them until they ignite and become white hot, like sparks from a fire.

12.10.2 Prepare salts by different methods
Methods of preparing salts
A salt compound is formed when the hydrogen ions in an acid are replaced by metal ions or ammonium ions.
A salt compound is formed by the combination of an acid radical or positive ions with a basic radical or negative ion.
A salt compound is formed by the replacement of all or part of the hydrogen in an acid by a metallic element.

1.0 Prepare salts by neutralization of soluble acid and base
Acid + Alkali --> Salt + Water
8.0 Neutralization, acid + base --> salt + water:
acid + alkali --> salt + water
H2SO4 (aq) + MgOH2 (aq) --> MgSO4 (aq) + 2H2O (l)
H2SO4 (aq) + 2NaOH (aq) --> Na2SO4 (aq) + 2H2O (l)
HCl (aq) + NaOH (aq) --> NaCl (aq) + H2O (l)
HCl (aq) + KOH (aq) --> KCl (aq) + H2O (l)
HCl (aq) + NH3 (aq) --> NH4Cl (aq)
HNO3 (aq) + NaOH (aq) --> NaNO3 (aq) + 2H2O (l)
HNO3 (aq) + KOH (aq) --> KNO3 (aq) + H2O (l)
HNO3 (aq) + NH3 (aq) --> NH4NO3 (aq)
CH3COOH (aq) + NaOH (aq) --> CH3COONa + H2O (l)

2.0 Prepare salts by acids + metals or metal oxides
2.1 Acid + metal
Acid + metal --> salt + hydrogen gas
(Metals that displace hydrogen from an acid may be called "active metals", e.g. Zn, Fe)
12.3.3 Dilute sulfuric acid with steel wool
6HCl (aq) + 2Al (s)--> 2AlCl3 (aq) + 3H2 (g)
2HCl (aq) + Fe (s) --> FeCl2 (aq) + H2 (g)
2HCl (aq) + Mg (s) --> MgCl2 (aq) + H2 (g)
2HCl (aq) + Zn (s) --> ZnCl2 (aq) + H2 (g)
H2SO4 (aq) + Cu (s) --> CuSO4 (aq) + 2H2O (l) + SO2 (g)
H2SO4 (aq) + Fe -->FeSO4 (aq) + H2 (g)
H2SO4 (aq) + Zn (s) --> ZnSO4 (aq) + H2 (g)
8HNO3 (aq) + 3Cu (s) --> Cu(NO3)2 (aq) + 4H2O + 2NO (g)
8HNO3 (aq) + 3Pb (s) --> 3Pb(NO3)2 (aq) + 4H2O + 2NO (g)
4HNO3 (aq) + Zn (s) --> Zn(NO2)2 +2H2O (l) + 2NO2 (g)

2.2 Acids + metal oxides
12.3.5: Dilute acids with basic oxides, metal oxides, copper (II) oxide
carbonic acid + calcium oxide
2HCl (aq) + FeO (s) --> FeCl2 (aq) + H2O (l)
2HCl (aq) + ZnO (s) --> ZnCl2 (aq) + H2O (l)
H2SO4 (aq) + CuO (s) --> CuSO4 ((aq) + H2O (l)
H2SO4 (aq) + ZnO (s) --> ZnSO4 (aq) + H2O (l)
2HNO3 (aq) + PbO (S) --> Pb(NO3)2 + 2H2O

2.3 Acids + metal hydroxides
2HCl (aq) + Ca(OH)2 (s) --> CaCl2 (aq) + 2H2O (l)
HCl (aq) + KOH (s) --> KCl (aq) + H2O (l)
2HCl (aq) + Mg(OH)2 (s) --> MgCl2 (aq) + 2H2O (l)
HCl (aq) + NaOH (s) --> NaCl (aq) + H2O (l)
H2SO4 (aq) + Mg(OH)2 (s) --> MgSO4 (aq) + 2H2O (l)
H2SO4 (aq) + 2NH4OH (aq) --> (NH4)2SO4 (aq) + 2H2O (l)
H2SO4 (aq) + Zn(OH)2 (s) --> ZnSO4 (aq) + 2H2O (l)
2HNO3 (aq) + Pb(OH)2 (s) --> Pb(NO3)2 (aq) + 2H2O (l)

2.4 Acids + metal carbonates
Acid + metal carbonate --> salt + water + carbon dioxide
12.3.9.0 Dilute acids with carbonates
2HCl (aq) + CaCO3 (s) --> CaCl2 (aq) + H2O (l) + CO2 (g)
2HCl (aq) + Na2CO3 (s) --> 2NaCl (aq) + H2O (l) + CO2 (g)
2HCl (aq) + K2CO3 (aq) --> 2KNO3 (aq) + H2O (l) + CO2 (g)
2HCl (aq) + ZnCO3 (s) --> ZnCl2 (aq) + H2O (l) + CO2 (g)
H2SO4 (aq) + CaCO3 (s) --> CaSO4 (aq) + H2O (l) + CO2 (g) [CaSO4 slightly soluble]
H2SO4 (aq) + CuCO3 (s) --> CuSO4 (aq) + H2O (l) + CO2 (g)
H2SO4 (aq) + MgCO3 (s) --> MgSO4 (aq) + H2O (l) + CO2 (g)
H2SO4 (aq) + ZnCO3 (s) --> ZnSO4 (aq) + H2O (l) + CO2 (g)
2HNO3 (aq) + K2CO3 (aq) --> 2KNO3 (aq) + H2O (l) + CO2 (g)
2HNO3 (aq) + ZnCO3 (s) --> Zn(NO3)2 (aq) + H2O (l) + CO2 (g)
HNO3 (aq) + CuCO3 (s) --> Cu(NO3)2 (aq) + H2O (l) + CO2 (g)
2HNO3 (aq) + MgCO3 (s) --> Mg(NO3)2 (aq) + H2O (l) + CO2 (g)
2HNO3 (aq) + PbCO3 (s) --> PbNO3)2 (aq) + H2O (l) + CO2 (g)
[Then H2SO4 (aq) + PbNO3)2 (aq) --> PbSO4 (s) + 2HNO3 (aq)], [to prepare PbSO4(s)]
CH3COOH (aq) + NaHCO3 (aq) --> CH3COONa + H2O (l) + CO2 (g) [prepare sodium acetate]

3.0 Prepare salts by precipitation reactions
Precipitation reactions, double decomposition reactions, double displacement reactions
Mixing two soluble compounds to prepare an insoluble salt,
Solution 1 + Solution 2 --> Insoluble solid 3 + Solution 4
12.2.1a Double decomposition, metathesis
3.33.2 Prepare insoluble hydroxides
AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)
2AgNO3 (aq) + BaCl2 (aq) --> 2AgCl (s) + Ba(NO3)2
Pb(NO3)2 (aq) + 2KI (aq) --> PbI2 (s) + 2K(NO3) (aq)
Pb(NO3)2 (aq) + 2HCl (aq) --> PbCl2 (s) + 2HNO3 (aq)
Pb(NO3)2 (aq) + 2NaCl (aq) --> PbCl2 (s) + 2NaNO3 (aq)
Pb(NO3)2 (aq) + H2SO4 (aq) --> PbSO4 (s) + 2HNO3 (aq)
Pb(NO3)2 (aq) + Na2SO4 (aq) --> PbSO4 (s) + 2NaNO3) (aq)
CaCl2 (aq) + Na2CO3 (aq) --> CaCO3 (s) + 2NaCl (aq)
BaCl2 (aq) + H2SO4 (aq) --> BaSO4 (s) + 2HCl (aq)
BaCl2 (aq) + Na2SO4 (aq) --> BaSO4 (s) + 2NaCl (aq)
CuSO4 (aq) + 2NaOH (aq) --> Cu(OH)2 (s) +Na2SO4 (aq)
Pb(NO3)2 (aq) + Na2SO4 (aq) --> PbSO4 (s) + 2NaNO3 (aq)
Pb(NO3)2 (aq) + H2SO4 (aq) --> PbSO4 (s) + 2HNO3 (aq)

4.0 Prepare salts by direct union of elements
Direct union of elements to form compounds: 8.0.0
Direct union of two elements, Synthesis: 12.2.2.0
2Al (s) + 3Cl2 (g) -- > 2AlCl3 (s)
2Fe (s) +3Cl2 (g) --> 2FeCl (s)
2Na (s) + Cl2 (g) --> 2NaCl (s)

5.0 Prepare salts by bases + nonmetallic oxides
12.17.3.1 Carbon dioxide with sodium hydroxide solution
CO2 (g) + 2NaOH (aq) --> Na2CO3 (aq) + H2O (l)

6.0 Prepare salts by acids + salts
12.18.5.6 Sulfuric acid with sodium chloride
sulfuric acid + sodium chloride --> sodium sulfate
H2SO4 (aq) + NaCl (s) --> NaHSO4 (s) + HCl (g)
H2SO4 (aq) + MgCl2 (s)--> Mg(HSO4)2 (s) + HCl (g)