Topic 17 Rate of reaction, catalysts, enzymes
Updated 2008-10-01
Please send comments to: J.Elfick@uq.edu.au See also: Interesting
websites See also: History of this
document
Table of contents 17.1.0 Measure rates of reaction 17.2.0 Factors affecting rate of
reaction 17.3.0 Catalysis 17.4.0 Enzymes and biological
catalysts 17.5.0 Chemical equilibrium, the law
of chemical equilibrium (law of mass action) 17.1.0 Measure
rates of reaction 3.91
Size of particles and rate of reaction 3.92
Concentration and rate of reaction 3.93
Temperature and rate of a reaction 3.94
Catalysts and rate of reaction 17.1.1 Count
bubbles, dilute
hydrochloric acid with granulated zinc 17.1.2 Volume of gas, dilute
hydrochloric acid with zinc 17.1.3 Gas burette, dilute
hydrochloric acid with marble chips (calcium carbonate) 17.1.4 Balloons, dilute hydrochloric
acid with marble chips 17.1.5 Height of suds, hydrogen
peroxide with manganese (IV) oxide (manganese dioxide) 17.1.6 Clock reaction,
hydrogen
peroxide with potassium iodide 17.1.7 Clock reaction, the "old Nassau flag"
(orange and black), sodium metabisulfite, mercury (II) iodide 17.2.0 Factors
affecting rates of reaction 17.2.1 Particle
size, dilute
hydrochloric acid with marble chips 17.2.2 Concentration of reactants,
hydrochloric acid with sodium thiosulfate (hypo) 17.2.2.1 Concentration of
reactants, hydrochloric acid with magnesium 17.2.3 Temperature and rate of
reaction, sulfuric acid with iron 17.2.4 Rates of reaction of aspirin 17.3.0 Catalysis 17.3.1 Hydrogen peroxide with
different substances 17.3.2 Catalysts, hydrogen peroxide
with manganese (IV) oxide. 17.3.3 Autocatalysis, ethanedioic
acid-2-water (oxalic acid) with potassium manganate (VII) (potassium
permanganate) 17.3.4 Spontaneous combustion, sugar
with potassium chlorate 17.3.5 Reverse colour change,
potassium iodide with hydrochloric acid 17.3.6 Activated complex, using cobalt (II)
chloride as catalyst 17.3.7 Double autocatalytic reaction
with bromine ions, oscillating reaction 17.3.8 Action of the enzyme catalase
using raw beef liver with hydrogen peroxide 17.3.9 Reactions of sodium
thiosulfate with hydrogen peroxide, ammonium molybdate catalyst 17.3.10 Oxidation of acetone vapour
using copper oxides catalyst 17.3.11 Heat potassium chlorate with
manganese dioxide as a catalyst 17.3.12 Catalytic oxidation of methyl alcohol,
ammonia 17.3.13 Bromine catalyses the
oxidation of sulfur to sulfuric acid. 17.3.14 Decomposition of sodium
hypochlorite using cobalt sulfate catalyst 17.3.15 Autocatalytic hydrolysis of
ethyl acetate 17.4.0 Enzymes and biological catalysts 17.4.01 Commercially available
enzymes 17.4.1 Biological catalysts, breakdown of starch
to sugar 17.4.2 Fermentation using yeast 17.4.3 Bromelain enzyme from
pineapples 19.2.9.1 Jelly using fresh
pineapple and
tinned pineapple
17.5.0
Chemical equilibrium, the law of chemical equilibrium (law of mass
action) 17.5.1 Concentration and temperature,
cobalt (II) chloride-6-water 17.5.2 Common ion effect, sodium
ethanoate and ethanoic acid 17.5.3 Equilibrium between ICl and ICl3 17.5.4.0 Rate of reaction depends
on concentration 17.5.4.1 Magnesium with hydrochloric
acid reaction 17.5.4.2 Potassium iodide with potassium iodate
reaction 17.5.5.0 Law of mass action and
reversible reactions, effect of alteration of concentration 17.5.5.1 Hydrolysis of bismuth
chloride 17.5.5.2 Hydrolysis of antimony
chloride 17.5.5.3 Common ion effect to
precipitate sodium chloride from solution 17.5.5.4 Common ion effect to
precipitate barium chloride from solution. 17.5.5.5 Effect of temperature on
chemical equilibrium, thermal dissociation of ammonium chloride 17.5.6.2 Action of heat on nitrogen
tetroxide 17.5.7.0 Explanation of group analysis 17.6.0 Gravimetric
analysis 17.6.1 Weight of iron in iron (II) ammonium
sulfate
17.6.2 Weight of aluminium in aluminium sulfate 17.6.3 Weight of calcium in marble, calcium
carbonate 17.6.4 Weight of magnesium in magnesium sulfate 17.6.5 Weight of sulfate radical in sodium sulfate 17.6.6 Weight of tin in solder 17.1.0 Measure rates of reaction
Rates of reactions for laboratory experiments should not be so fast
that an explosion occurs. Also, they should not be so slow that you
cannot observe or measure any change in reasonable time. To measure the
speed of a reaction you must measure some change, e.g. time taken for
colour change, pH change, precipitate appears, reagents disappear,
temperature change, volume of gas that forms weight of reactants used.
Chemical reactions occur only if the particles of reactants can collide
with sufficient energy, the activation energy, Ea. Reactions
will become faster if the number of collisions increases and if the
energy of collision increases. You can increase the rate of reaction by
increasing concentration of reactants, light energy, particle size,
pressure, temperature, and with catalysts. 17.1.1 Count bubbles, dilute hydrochloric acid
with zinc See diagram 17.1.1
Put a piece of granulated zinc in dilute hydrochloric acid in a
test-tube. Count the number of bubbles of hydrogen that reach the
surface of the solution every 30 seconds. Draw a graph by plotting the
number of bubbles along the vertical axis and time along the horizontal
axis. The reaction starts quickly then becomes slower until it stops. 17.1.2 Volume of gas, dilute hydrochloric acid
with zinc See diagram 17.1.2
Repeat the experiment with a flask connected to a gas syringe to
measure how much hydrogen gas forms in the reaction. 17.1.3 Gas burette, dilute hydrochloric acid with
marble chips See diagram 17.3.3 | See
also: Saturation vapour pressure over water
Add zinc or marble chips to dilute hydrochloric acid. Collect the
gas in a burette inverted over water. Compare how much gas forms in
unit time for each size of marble chips or zinc Weigh the container and
note the loss in mass every half minute while the reaction goes on. 17.1.4 Balloons, dilute hydrochloric acid with
marble chips See diagram 17.3.2
Add marble chips to dilute hydrochloric acid. Show the relative
production of hydrogen by attaching a previously stretched balloon to
each test-tube. 17.1.5 Height of suds, hydrogen peroxide with
manganese (IV) oxide
Manganese (IV) oxide does not take part in the chemical reaction.
Its function is to provide an increased surface area for the reactants.
Pour 10 mL of hydrogen peroxide, 5 mL of detergent and 5 mL of
glycerine into two identical measuring cylinders, Measuring cylinder 1.
and Measuring cylinder 2. Stir both
solutions.
Measuring cylinder 1. Add crystals of manganese (IV) oxide (MnO2,
manganese dioxide) and stir again. A mass of
bubbles arises to form suds. The glycerine
increases the surface tension of the liquid to delay the collapsing of
the bubbles. The hydrogen peroxide decomposes into water and the oxygen
that forms the bubbles. The suds are higher in the measuring cylinder
containing the manganese (IV) oxide because it acts as a
catalyst.
Measuring cylinder 2. Control without manganese (IV) oxide.
Repeat the experiment with dust or dirt or ashes instead of
manganese (IV) oxide. The heights of the detergent suds are
similar to the
heights using manganese (IV) oxide. This observation shows that
manganese (IV) oxide does not take part in the chemical reaction.
Its function is to provide an increased surface area for the reactants. 17.1.6 Clock reaction, hydrogen peroxide with
potassium iodide
1. Solution A: Add drops of water to 0.2 g soluble starch, pour this
paste into a beaker containing 1 cc boiling water and stir. Pour this
solution into a 1 litre beaker and dilute to 800 mL. Add 30 mL glacial
ethanoic acid (glacial acetic acid, CH3CO2H)
4.1 g sodium ethanoate (sodium acetate, CH3CO2Na)
50 g potassium iodide (KI) and 9.4 g sodium thiosulfate-5-water (Na2S2O3.5H2O).
Dissolve all solutes by stirring and when cooled to room temperature
make up solution to 1 litre. The solution is slightly cloudy. Solution
B: Dilute 500 mL 20 vols hydrogen peroxide to 1 litre.
Put 100 mL of each solution in separate beakers. Pour the contents
of one beaker into another and stir quickly or use a magnetic stirrer.
After about 20 seconds the colourless mixture suddenly turns dark blue.
Repeat the experiment at room temperature but note the time taken for
the colour change after mixing.
2. Repeat the experiment at 10oC above room temperature.
Note the time taken for the colour change after mixing, about double
the time compared to 1.
3. Repeat the experiment at 10oC below room temperature.
Note the time taken for the colour change after mixing, about
half the
time compared to 1.
4. Repeat the experiment at room temperature using half the
concentration of solution B, i.e. 50 mL solution B + 50 mL water. Note
the time taken for the colour change after mixing, about double the
time compared to (1.) because the reaction rate is halved. The amount
of
hydrogen peroxide has been halved.
5. Repeat the experiment at room temperature using half the
concentration of solution A, i.e. 50 mL solution A + 50 mL water. Note
the time taken for the colour change after mixing, the same as for
(1.). The reaction rate has been halved but the amount
of thiosulfate also has been halved so you only need to make half the
amount of iodine to combine with all the thiosulfate. If the sodium
thiosulfate was not in solution A but in a solution C, the result of 5.
would be the same as in 4.
H2O2(aq) + 2I-(aq) + 2H+(aq)
---> I2(aq) + 2H2O(l)
Hydrogen peroxide reacts with iodide ions to form iodine. The ethanoic
acrid with sodium ethanoate buffer the reaction
2S2O32-(aq) + I2(aq)
--->
S4O621(aq) + 2I-(aq)
The iodine reacts with thiosulfate to form tetrathionate ions and
iodide ions return to the solution. As soon as all the thiosulfate is
converted to tetrathionate ions the remaining iodine reacts with the
starch solution to form a blue black colour. 17.1.7 Clock reaction, the "old Nassau flag"
(orange and black), sodium metabisulfite, mercury (II) iodide This experiment should NOT be done
in schools because it uses mercuric (II) chloride, HgCl2!
Sodium metabisulfite reacts with water to form sodium hydrogensulfite -
colourless reaction
Na2S2O5(aq) + H2O(l)
--->
2NaHSO3(aq)
Hydrogensulfite ions reduce iodate (V) ions to iodide ions - colourless
reaction
IO3-(aq) + 3HSO3-(aq)
---> I-(aq) + 3SO42-(aq) + 3H+(aq)
With excess iodide ions, mercury (II) iodide, HgI2,
forms as
an orange precipitates when solubility product of mercury (II)
iodide
is exceeded. When all mercury has precipitated, remaining iodate and
iodide ions react to form iodine and a blue black iodine starch complex
forms.
IO3-(aq) + 5I-(aq) + 6H+(aq)
---> 3I2(aq) + 3H2O(l) 17.2.1 Particle size, dilute hydrochloric acid
with marble chips (calcium carbonate)
The smallest particles show the most vigorous reaction because if you
grind two pieces of substance with the same weight into coarse and fine
sizes, the fine size pieces would have total surface area greater than
the coarse pieces. BE CAREFUL!
Hydrogen gas forms.
Break marble chips or granulated zinc into four sizes with a hammer
Put 2 g of each into four test-tubes: Test-tube 1, original size as
control, Test-tube 2, rice grain size, Test-tube 3, half rice grain
size,
Test-tube 4, coarse powder. Add the same volume of 2 M hydrochloric
acid to each test-tube. Compare the reactions. Rate of reaction of
Test-tube 4 > Test-tube 3 > Test-tube 2 > Test-tube 1. 17.2.2 Concentration of reactants, hydrochloric
acid with sodium thiosulfate See diagram 17.4.1
1. The reaction slowly forms sulfur that makes the solution cloudy.
Measure the rate of reaction by measuring the cloudiness in the
solution. Note the time until you can no longer observe a black cross
on a piece of paper below the beaker. Change the concentration of the
sodium thiosulfate and keep the concentration of hydrochloric acid
constant. Dissolve 20 g sodium thiosulfate in 500 mL of water. Pour 50
mL of this solution into a 100 mL beaker. Put the beaker on a sheet of
paper marked with a black cross. Add 5 mL of 2 M acid and stir the acid
into the solution. Record the time when the cross is no longer visible
through the precipitated sulfur in the solution.
2. Repeat the experiment with a smaller concentration of thiosulfate
as follows:
.
Sodium thiosulfate Solution
deionized water
Beaker 1
30 mL
20 mL
Beaker 2
20 mL
30 mL
Beaker 3
10 mL
40 mL
Stir the solution, then add 5 mL of 2 M acid as before. The time for
the cross to become invisible is greater. The reaction takes longer as
the concentration decreases. Draw a graph to plot concentration of the
thiosulfate solution against time taken for the reaction. Express
concentration values as the volume of the original thiosulfate solution
used, e.g. 50 mL, 30 mL, 20 mL and 10 mL.
Na2S2O3(aq) + 2HCl(aq) ---> H2O(l)
+ SO2(g) + S(s)
17.2.2.1 Concentration of reactants,
hydrochloric acid with magnesium
Pour 50 mL 6 M hydrochloric acid solution into a 100 mL standard
flask. Make up the solution to the mark and mix it well. This is
Solution B, 3 M. Pipette 50 mL Solution B into another 100 mL standard
flask. Make up this solution to the mark and mix it well. This is
Solution C, 1.5 M. Pour the remaining solution into a 100 mL beaker.
Pipette 50 mL solution C into another 100 mL standard flask and make
this solution up to the mark and mix it well. Label this solution D,
0.75 M. Pour the remaining solution into a 100 mL beaker. Pipette 50 mL
of solution D into a 100 mL beaker. Cut off four 0.5 cm pieces of
magnesium ribbon. Put one strip into solution A and record the time of
the reaction until you have dissolved all of the magnesium. Repeat with
a new magnesium strip for each of solutions B, C, and D. 9. Graph the
results with time as the vertical axis and concentration as the
horizontal axis. Note whether a linear relationship exists between acid
concentration and the rate at which magnesium dissolves. 17.2.3 Temperature and rate of reaction, sulfuric
acid with iron
Increase in temperature of 10oC usually doubles the rate of
reaction. The rate increases because collisions between particles are
more frequent and with more energy. Add enough potassium
manganate (VII) solution to dilute sulfuric acid to make it pink. Put
in
a nail and record the time for the solution to lose the pink colour at
room temperature. Repeat the experiment at 10oC above room
temperature. Repeat the experiment at 20oC above room
temperature. Draw a graph to show time taken to decolorize against
temperature. The rate of reaction may double for each 10oC
rise in temperature. 17.2.4 Rates of reaction of aspirin Be careful! Aspirin is a
drug so do not let the students take them!
Soluble aspirins dissolve in water with a well defined end point and
constants interaction time. Use these
tablets to investigate the effects of temperature, stirring, crushing
and
varying the water volume on reaction time. Investigate the temperature
change required to halve the reaction
time between one tablet and known quantity of water at room
temperature. Investigate the number of tablets or fractions of tablets
required to exactly double the
reaction time in a known quantity of water at room temperature. 17.3.0 Catalysis
Catalysts increase the rate of reactions without themselves being
chemically changed. A catalyst can change the rate of a chemical
reaction without itself being permanently changed. They provide an
alternative pathway for the reactions and so decrease the activation
energy needed. Substances that slow the rate of reactions are called
inhibitors. 17.3.1 Hydrogen peroxide with different substances
Some substances can increase the rate of reaction for the decomposition
of hydrogen peroxide.
Put two small equal amounts of hydrogen peroxide solution gently
into four test-tubes. Note the bubbles. Tests for oxygen with a glowing
splint. Add the following substances: Test-tube 1, Nothing added
(control) Test-tube 2, Manganese (IV) oxide, Test-tube 3,
Iron (III)
chloride. Test-tube 4, Copper (II) sulfate. Note the bubbles.
Tests for
oxygen with a glowing splint. 17.3.2 Catalysts, hydrogen peroxide with
manganese (IV) oxide catalyst, catalytic decomposition of
hydrogen peroxide
1. Fill a test-tube with hydrogen peroxide to a depth of about 1 cm
and add a little of manganese dioxide. Tests for oxygen with a glowing
splint. A rapid evolution of oxygen occurs and the manganese dioxide
is not lost.
2H2O2(l) > 2H2O(l) + O2(g)
2. Put 2 cm of 20 vols hydrogen peroxide (20 vol. solution) in two
test-tubes, Test-tube 1 and Test-tube 2.
Test-tube 1. Add 5 drops of sodium hydroxide solution.
Test-tube 2.
Add 5 drops of dilute sulfuric acid solution.
Immerse
both test-tubes in a beaker half full of hot water. Use a glowing
splint to
test for oxygen in the mouth of the test-tubes. Oxygen is found in
Test-tube 1 but not in Test-tube 2.
3. Set up a conical flask fitted with a one hole stopper and
delivery tube that leads into a beaker of water. Invert a closed
burette full of water over the end of the delivery tube. Pour 50 mL of
water in the flask and add 2 mL of 20 vols (6%) hydrogen peroxide
solution. Add 1 g manganese (IV) oxide and immediately insert the
stopper with the delivery tube into the flask. Note how much oxygen
forms every 15 seconds. Plot on a graph how much oxygen forms every 15
seconds against the time of the reaction.
4. Repeat the experiment by adding more hydrogen peroxide solution to
the same test-tube. The manganese (IV) oxide is not "used up"
because
more oxygen forms. Repeat the experiment with 1 g copper (II)
oxide.
Add 2 mL of 20 vols (6%) hydrogen peroxide solution.
5. Repeat the experiment with 1 g zinc oxide. Add 2 mL of 20 vols (6%)
hydrogen peroxide solution. Plot a graph for each experiment.
Manganese (IV) oxide is the better catalyst in these reactions.
Warm some
hydrogen peroxide solution gently in a test-tube and hold a glowing
splinter of wood in the mouth of the test-tube.
6. Place two small equal amounts of hydrogen peroxide in separate
test-tubes. Add some iron chloride to one test-tube and a little
manganese dioxide to the other. Apply the glowing splinter test to any
gas given off. Hydrogen peroxide decomposes to give off oxygen when
heated but will decompose without heating when iron chloride and
manganese dioxide are added to it. The black colour of manganese
dioxide and the brown colour of iron chloride remains after the
reactions, so these chemicals may not have been altered during the
reactions. Iron chloride and manganese dioxide are catalysts or
"chemical accelerators". 17.3.3 Autocatalysis, ethanedioic acid-2-water
(oxalic acid) with potassium manganate (VII)
Solution A: 6 g of ethanedioic acid-2-water (oxalic acid) in 300 mL
of water. Solution B: 100 mL of 0.001 M potassium manganate (VII)
solution. Pour 150 mL of Solution A into two beakers, then add 5 mL of
concentrated sulfuric acid to each beaker. Add 50 mL of the potassium
manganate (VII) solution to each beaker. Add a small crystal of
manganese (II) chloride MnCl2 to one beaker then stir
the
solutions in both beakers The solution containing the crystal of
Manganese (II) Chloride starts to lose colour and becomes
colourless in
about a minute. The other solution does not change in colour for two or
three minutes but, when sufficient Mn2+ ions are present, it
starts to become colourless. The Mn2+ autocatalyses the
solution as follows.
2MnO4-(aq) + 5C2O42+(aq)
+ 16H3O+(aq) ---> 2Mn2+(aq) + 10CO2(g)
+ 24H2O(l) 17.3.4 Spontaneous combustion, sugar with
potassium
chlorate BE CAREFUL! Do this experiment in
a fume cupboard or in the open, behind a glass screen.
The reaction of the concentrated sulfuric acid with the sugar released
heat. The heat then activated the release of oxygen from the potassium
chlorate. The oxygen released by the potassium chlorate further
oxidized the sugar. This further oxidation released so much heat that
the sugar bursts into flames.
Mix sugar or powdered sugar (castor sugar) with an equal amount of
potassium chlorate crystals in an evaporating dish. Push a dent in the
top of the heap of powder. Add one drop of concentrated sulfuric acid.
A spontaneous combustion occurs. BE CAREFUL! 17.3.5 Reverse colour change, potassium iodide
with
hydrochloric acid
Dissolve potassium iodide crystals in water. Add drops of starch
solution and dilute hydrochloric acid. The solution is colourless. Add
drops of dilute hydrogen peroxide solution. The solution turns blue
black. Iodide ions are oxidized to iodine that gives starch a blue
black colour.
2KI(aq) + H2O2(l) ---> 2KOH(aq) + I2(g)
Add drops of dilute sodium thiosulfate solution. The solution turns
colourless. The sodium thiosulfate reduces the iodine back to iodide
ions that are colourless.
I2(g) + Na2S2O3(aq) --->
NaI(aq) + Na2S2O5(aq)
Wait until the blue black colour returns. Add drops of sodium
thiosulfate solution and it disappears again. The first reaction is
still going slowly. The second reaction is much slower.
17.3.6 Visible activated complex using
cobalt (II) chloride as catalyst
Use a 250 mL beaker. Add 20 mL 6% hydrogen peroxide solution to a
solution of 5 g potassium sodium tartrate-4-water (Rochelle salt) in 60
mL water. Heat to 75oC. Stir and observe that gases are
formed and a pink colour forms. Add 5 mL of solution containing 0.2 g
cobalt (II) chloride-6-water. Observe frothing and pink colour. Be careful!
Test gases with limewater. Carbon dioxide is in gases produced by the
reaction. Test gases with a glowing splint, not extinguished because
some oxygen from hydrogen peroxide. Later, observe frothing stops and
pink colour returns. Equation in 2 parts:
C4H4O62-(aq) + 3H2O2(aq)
---> 2HCOO-(aq) + 2CO2(g) + 4H2O(l)
Co2+ ---> Co3+
Pink cobalt ions oxidized to a green activated complex with tartrate
2HCOO-(aq) + 2H2O(aq) ---> 2OH-(aq)
+ 2CO2 + 2H2O
Co3+ ---> Co2+
Basic equation: Green activated complex with tartrate reduced to pink
cobalt ions
C4H4O62-(aq) + 5H2O2(aq)
---> 4CO2(g) + 2OH-(aq) + 6H2O(l) 17.3.7 Double autocatalytic reaction with bromine
ions, oscillating reaction
In this reaction, bromine ions form to give a red colour but some
intermediate product also forms to react with bromine ions to give a
colourless solution.
Use a clean beaker washed in deionized water. Add 75 mL
concentrated sulfuric acid to 750 mL deionized water (NOT tap water!). Be careful!
Leave the hot acid solution long enough to cool to room temperature
slowly. Stir the cooled solution fast enough to form a vortex. Add 9 g
propanedioic acid (CH2[CO2H)2, malonic
acid}. Add 8 g potassium bromate (V) [KBrO3]. Add 1.8 g
manganese (II) sulfate [MnSO4.H2O]. Observe
a red
colour that oscillates from red to colourless, with increasing time
between oscillations.
3CH2(CO2H)2(aq) + 4BrO3-
---> 4Br-(aq) + 9CO2(g) + 6H2O(l) 17.3.8 Action of the enzyme catalase using raw
beef liver with hydrogen peroxide
Hydrogen peroxide, acting as an oxidizing agent, it is toxic to cells,
so it is a useful disinfecting agent that disrupts the metabolism of
bacteria. Our body cells contain an enzyme called catalase that
accelerates the conversion of toxic hydrogen peroxide to water and
oxygen gas. It is a reactive oxygen metabolic by-product that regulates
some oxidative stress-related states related to asthma, inflammatory
arthritis, atherosclerosis, diabetic vasculopathy, osteoporosis, and
some neurodegenerative diseases. BE CAREFUL! This
reaction can be violent and the steam formed may be hot. Use safety glasses and nitrile
chemical-resistant gloves.
Put 10% hydrogen peroxide in a clear plastic container and record
its temperature. Put a 6 mm piece of chopped liver in the hydrogen
peroxide. The mixture starts to bubble and foam. Record the temperature
each minute for 5 minutes. The temperature immediately rises, levels,
then decreases. The hydrogen peroxide is decomposed, oxygen gas is
given off, bubbles form and heat energy is given off.
Put a glowing splint in the test-tube of liver and hydrogen peroxide.
The splint flames in the oxygen. If you boil the liver before the
experiment, no bubbling occurs when hydrogen peroxide is added because
you have destroyed the enzyme catalase in the beef liver, i.e. it is
denatured. Put 6 g manganese dioxide into 100 mL 10% hydrogen peroxide.
Observe bubbling and relights a glowing wooden splint. Repeat the
experiment using freshly cut potato instead of raw beef liver.
2H2O2 (l) ---> H2O (g) + O2(g)
17.3.9 Reactions of sodium thiosulfate with
hydrogen peroxide, ammonium molybdate catalyst
Dissolve together in deionized water: 8.7 g sodium
thiosulfate-5-water, 3.8 g sodium ethanoate-3-water (sodium acetate
tri-hydrate) or 2.3 g anhydrous sodium ethanoate, and 0.5 g sodium
hydroxide. he sodium ethanoate buffers the sodium hydroxide. Make up
the solution to 1 litre and add universal indicator until each solution
is blue. Pour 225 mL of the solution into 3 test-tubes labelled 1.
"With catalyst" 2. "No catalyst" 3. "Control". Dissolve 14 mL of 20
vols hydrogen peroxide in deionized water. Make up to 40 mL and divide
into two 20 mL portions. Add 0.08 g ammonium molybdate to test-tube 1.
"With catalyst", and shake to dissolve. Add the 20 mL portions to
test-tubes 1. "With catalyst" and 2. "No catalyst". Observe colour
changes after 5 minutes: 1. "With catalyst": changes from blue to
green to yellow to orange to orange red 2. "No catalyst": Same colour
changes but slower 3. "Control": No colour change.
Na2S2O3(aq) + 4H2O2(aq)
---> Na2SO4(aq) + H2SO4(aq)
+ 3H2O(l) 17.3.10 Oxidation of acetone vapour using copper
oxides catalyst
1. Heat a copper wire coil or copper coil to red heat in a Bunsen
burner flame then hang it just above a very thin layer of acetone in a
beaker. Note the shimmering colours of the copper surface caused
by the heat being maintained in the copper coin from the heat of the
chemical reactions on its surface, and the colours of black
copper (II) oxide, red copper (I) oxide (red) and pink copper
metal. You can warm the beaker in hot water to produce sufficient
acetone vapour. Be careful! Do not smell the
oxidation products because they contain ketones that may be a health
hazard. The red heat is maintained as long as some acetone remains
because it keeps warm through the heat of the exothermic
reaction. The reaction is safe except that at the top of the beaker
where air dilutes the vapour a flame may occur. If this happens move
the hot copper quickly out and cover the beaker to extinguish the flame.
2. Drop the red-hot copper into a beaker with a 2 mm thin layer of
acetone at the bottom. A sizzling sound caused by the cooling of the
copper ends in a loud crescendo when liquid acetone contacts the copper
to increase the rate of cooling. Be
careful! 17.3.11
Heat potassium chlorate with manganese dioxide as a catalyst
Be careful! In some school systems this experiment is not allowed
because potassium chlorate may explode.
Mix 0.5 g of manganese dioxide with 2 g of potassium chlorate and
put the mixture in Ignition tube 1. Put 0.5 g of manganese dioxide in
Ignition tube 2. Put 2g of potassium chlorate in Ignition tube 3.
Insert the ignition tubes vertically and close together in a sand
tray and place a safety screen between you and the sand tray. Slowly
heat the sand tray. Use a glowing splint to test for oxygen at the
openings of the ignition tubes. Oxygen appears first from Ignition tube
1 and later from the other Ignition tubes. When the reaction is
complete, wash the contents into a
beaker. Stir the contents to dissolve all the potassium chloride and
any remaining potassium chlorate. Filter the mixture, dry the residue
on the filter paper and weigh the manganese dioxide residue to show
that there is no loss in weight. 17.3.12 Catalytic oxidation of methyl alcohol,
ammonia
Be careful! Have ready a piece of cardboard or glass to put over the
beaker if the methyl alcohol ignites.
See diagram 17.3.12
1. Prepare a platinum spiral by winding platinum wire around a glass
rod and leave a length of wire above the spiral. Put 1 cm of methyl
alcohol in a small beaker and warm it gently with an electric heater.
Do NOT use a Bunsen burner. Heat the spiral strongly with the electric
heater and transfer the glowing spiral to the beaker, holding the
length of wire above the spiral so that the spiral is just above the
methyl alcohol. The spiral continues to glow and you can notice the
smell of formaldehyde given off from the reaction
2CH3OH + O2 ---> 2HCHO + 2H2O
2. Repeat the experiment with 880 ammonia solution instead of methyl
alcohol and pass oxygen gas from an oxygen cylinder through the
mixture. Observe brown fumes of nitrogen dioxide or white fumes of
ammonium nitrate and ammonium nitrite.
4NH3 + 7O2 ---> 4NO2 + 6H2O 17.3.13 Bromine catalyses the oxidation of sulfur
to sulfuric acid
Do the experiment in a fume cupboard.
Put 1 cc of flowers of sulfur into two evaporating basins. Add 5 mL
of concentrated nitric acid to each evaporating basin. Be
careful! Add one drop of bromine to Evaporating basin 1. Be careful!.
Warm each evaporating basin for 2 minutes. Pour the solutions into
Test-tube 1 and Test-tube 2. Add hydrochloric acid then barium chloride
solution to each test-tube. A precipitate of barium sulfate occurs only
in Test-tube 1.
In this catalysis intermediate compounds form which are more readily
decomposed.
2S + Br2 ---> S2Br2 sulfur
monobromide
2S2Br2 + 2H2O > SO2 +
4HBr + 3S
SO2 + 2HNO3 > H2SO4 +
2NO2
2HBr + 2HNO3 > 2H2O + 2NO2 +
Br2 17.3.14 Decomposition of sodium hypochlorite
using cobalt sulfate catalyst
Warm a test-tube half full of sodium hypochlorite solution and observe
that no decomposition occurs. Add a two drops of cobalt sulfate
solution and test for oxygen with a glowing splint.
2NaOCl > 2NaCl + O2(g) 17.3.15 Autocatalytic hydrolysis of ethyl acetate
1. Prepare the following solutions: Flask A 100 mL of 0.5 M sulfuric
acid, Flask B 0.5 M hydrochloric acid, Flask C 0.5 M acetic acid. Leave
the 3 flasks in a thermostat at 25oC so that all the
contents are at the same temperature. Titrate 2 mL of each acid in
Flask A, Flask B and Flask C
separately against 0.1 M sodium hydroxide, with phenolphthalein
indicator and record the results.
2. Add 5 mL of ethyl acetate solution to Flask A, Flask B and Flask
C then leave them in a thermostat for 15 minutes. Titrate 2 mL of the
Flask A, Flask B and Flask C solutions (+ ethyl acetate) against 0.1 M
sodium hydroxide solution, with phenolphthalein indicator. Repeat the
titrations every 15 minutes for 2 hours and tabulate the results.
Increased titration with time is because of the formation of acetic
acid by hydrolysis of ethyl acetate. The catalytic action of the
hydrogen ions in the original acid is increased by the hydrogen ions
from the acetic acid produced by the hydrolysis. The rate of
hydrolysis in the presence of a mineral acids in Flask A and Flask B is
higher than in Flask C where the hydrogen ion concentration is
lower. This reaction is an example of autocatalysis.
CH3COOC2H5 + H2O ---> CH3COOH
+ C2H5OH
17.4.01 Commercially available enzymes:
Enzymes are organic catalysts, but are usually used up in the reaction.
An enzyme activity experiment uses a Thunberg tube to estimate
dehydrogenase activity in plant material with DCPIP as indicator.
Enzyme activity is affected by time, temperature, substrate
concentration, enzyme concentration, pH, substrate specificity of
enzyme, heat denaturation, and inhibition of enzyme activity by
chemicals.
Commercially available enzymes: Enzyme biotechnology, e.g. enzymes in
washing powders; amylase, starch ---> maltose; cellulase, leaf
materials ---> various breakdown products; diastase, starch --->
maltose; invertase, sucrose ---> (+) glucose + fructose; lactase,
milk ---> (+)glucose + galactose; pancreatin (amylases, proteases,
lipases); food ---> various breakdown products; pectinase, fruit
pulp or leaves ---> various breakdown products; pepsin, skimmed
milk
---> peptides, protease, skimmed milk ---> peptides + amino
acids (in meat tenderizers); rennin (rennet powder, rennet tablets,
rennilase, junket tablets) milk ---> caseinogen ---> casein;
trypsin, proteins ---> peptides + amino acids; urease, urea --->
ammonia and carbon dioxide. 17.4.1 Biological catalyst, breakdown of starch
to
sugar 9.130
Hydrolysis of starch by salivary amylase (ptyalin) 9.142
Tests for starch, Fehling's test for starch
Salivary
amylase enzyme breaks down
starch into the reducing sugars (+)glucose and maltose. Reducing sugars
do not react with iodine
solution and starch does not react with Fehling's solution. The sugars
reduce copper (II) in Fehling's solution to brick-red
copper (I) oxide.
Pour 10 mL of dilute starch solution into a test-tube. Add 1 mL of
saliva and stir. After 2 minutes use a dropper put 2 drops of the
solution on a white tile. To test for starch, add iodine solution and
note the intensity of the blue black colour. To test for reducing
sugars, add Fehling's No. 1 and No. 2 solutions and heat. Repeat the
experiment every 2 minutes with clean droppers. Note the decreasing
intensity of the blue colour that shows that starch is being used up.
Repeat the experiment. Put 3 drops of the reaction mixture in a
test-tube. Add 3 mL of Fehling's solution. Heat the mixture. Note the
intensity of the brick-red colour increasing with time. 17.4.2 Fermentation using yeast See also 3.38: Carbon dioxide and
fermentation for brewing
During fermentation, enzymes breakdown carbohydrates and other organic
molecules in the absence of oxygen. 17.4.3 Bromelain enzyme from pineapples
Add pineapple juice to milk. The milk protein begins to coagulate
and degrade as it reacts with the bromelain. Also, pineapple juice will
also remove the gelatine emulsion surface on black and white
photographic film! 17.5.0 Chemical Equilibrium
1. Equilibrium exists only in a closed system. No reactants are put in
and no reactants are taken out. A closed system cannot exchange matter
with its surroundings but it may exchange energy with its surroundings.
2. Limestone (calcium carbonate) is heated in a furnace to form
quicklime (calcium oxide) and carbon dioxide. The reverse reaction
cannot occur because the carbon dioxide is sucked out of the furnace.
Also, the calcium oxide is steadily and replaced by limestone. This is
a
steady state system. If the furnace were closed, at high temperature
both decomposition and formation of calcium carbonate would occur. When
these process occur at the same rate then equilibrium exists.
CaCO3(s) ---> CaO(s) + CO2(g)
3. If different concentrations of iron (II) nitrate solution are
added
to potassium thiocyanate solution, at equilibrium the concentration of
the FeSCN2+ solution has a certain blue colour. It is a
property of this equilibrium reaction that does not change. If
sodium thiocyanate NaSCN or iron (III) nitrate Fe(NO3)3
is added to the equilibrium mixture it changes colour.
Fe3+(aq) + SCN -(aq) <---> FeSCN2+(aq)
(d) The following equation shows that both forward and reverse
reactions are going on. It does not show the position of equilibrium.
The position at equilibrium shows whether there are more reactants or
more products at equilibrium. It can shift if reactants or products are
added or removed.
N2O4(g) <---> 2NO2(g)
In the above reaction, if, at equilibrium, N2O4 is
added, the system moves to the right, i.e. some N2O4 changes
to NO2 until equilibrium is reached.
In the above reaction, if, at equilibrium, N2O4 is
removed, the system moves to the left, i.e. some NO2 changes
to N2O4 until equilibrium is reached.
Position at equilibrium can shift if the temperature changes but it
will not shift because solid is added or removed from a system.
Position at equilibrium does not change if a catalyst is added to the
reaction but if the system is not at equilibrium it will reach
equilibrium quicker because of the influence of the catalyst.
(d) Le Chatelier's principle states that if the conditions of a system
at equilibrium are altered, changes will occur in the system towards
counteracting the change in conditions, when a system in equilibrium is
subjected to a change in conditions, it adjusts itself so as to oppose
that change, Henri Le Chatelier, 1850 - 1936.
The law of chemical equilibrium (law of mass action)
For reaction aA + bB <---> eE + fF (lower case = number of
species, upper case = different types of species, e.g. 3 species of
oxygen gas = 3O2) at equilibrium: [E]e X [F]f
/ [A]a X [B]b = K, where K is the equilibrium
constant of that reaction at that temperature. If K is large, at
equilibrium, the concentration of products is much greater than the
concentration of the reactants. If K is small, at equilibrium, the
concentration of products is much smaller than the concentration of the
reactants. 17.5.1 Concentration and temperature, cobalt
(II)
chloride-6-water
Test the effect of concentration and temperature on the equilibrium
position. Dissolve 4 g cobalt (II) chloride-6-water in 40 mL
water. The
solution contains Co(H2O)62+. It is
pink. A dd concentrated hydrochloric acid until total volume is 100 mL.
The solution is violet (between pink and blue). Add more concentrated
hydrochloric acid to the violet solution. The solution contains CoCl42-.
It is blue. Add water to the violet solution. It turns pink. Heat the
solution. It turns blue. Cool the solution with ice water. It turns
pink. Add sodium chloride to the pink solution. It turns blue.
Co(H2O)62+(aq) + 4Cl-(aq)
<---> CoCl42-(aq) + 6H2O (deltaH
+ve) 17.5.2 Common ion effect, sodium ethanoate and
ethanoic acid
Hydrochloric acid is a completely dissociated strong acid. Ethanoic
acid is a less dissociated weak acid.
Prepare 3 solutions and add universal indicator to each: Solution
"A" 100 mL 2 M hydrochloric acid, then add universal indicator until
red; Solution "B" 100 mL 2 M ethanoic acid, then add universal
indicator until orange; Solution "C" Add 13.5 g sodium
ethanoate-3-water to 80 mL 2 M ethanoic acid. Add more ethanoic acid to
100 mL, then add universal indicator until yellow. Put 1 g calcium
carbonate powder + 1 mL liquid detergent into 3 measuring cylinders
containing 100 mL water, measuring cylinders "A1", "B1",
and "C1". Simultaneously, put solution "A" into "A1",
solution "B" into "B1", and solution "C" into "C1".
Observe the measuring cylinders: "A1" Fastest froth
produced, height h, "B1" Second fastest froth produced,
height h, "C1" Slowest froth produced, height h/2. In this
mixture of sodium ethanoate and ethanoic acid, the equilibrium moves
left, decreasing the concentration of hydrogen ions.
CH3CO2H(aq) <---> CH3CO2-(aq)
+ H+(aq) 17.5.3 Equilibrium between ICl and ICl3 See diagram 17.5.3
Put 0.1 g iodine in the U-tube. With the 3-way tap in the CB
position, turn on the filter pump to pass air bubbles through the
sodium hydroxide solution. With the 3-way tap in the AB position, turn
on the filter pump and let drops of concentrated hydrochloric to fall
on 10 g potassium permanganate. As the chlorine passes over the iodine,
first iodine monochloride, ICI, forms as a brown liquid. then iodine
trichloride, ICI3, forms as a yellow solid. Turn the 3-way
tap in the CB position, remove stopper 1 from the U-tube to draw in
air. The yellow solid turns into a brown liquid. Replace stopper 1 in
the U-tube and turn the 3-way tap to the AB position. The brown liquid
turns into a yellow solid. Increasing concentration of chlorine moves
equilibrium to the right, obeying Le Chatelier's principle. With the
3-way tap in the AC position to get rid of the chlorine, tighten the
screw clips. When you dip the U-tube in water just below boiling point,
the yellow solid turns into a brown liquid. When you dip the U-tube in
ice water, the brown liquid turns into a yellow solid. Increase of
temperature moves the equilibrium to the left. Decrease of temperature
moves the equilibrium to the right. Dispose of the chemicals by using
different sinks. Pour the sodium hydroxide solution into a laboratory
sink and wash it down with plenty of water to follow. Wash the contents
of the flask generating the chlorine down the sink in the fume cupboard
with plenty of water to follow.
ICl(l) + Cl2(g) ---> ICl3(s) 17.5.4.1 Magnesium with hydrochloric acid
reaction
A reversible reaction can proceed in either direction by
altering the conditions of the reaction, e.g. 1. altering the relative
concentrations, active mass 2. altering the temperature 3. altering
the pressure.
Measure 10 mL. of concentrated hydrochloric acid into 4 beakers: beaker
A add nothing, beaker B add 10 mL water, beaker C add 30 mL water,
beaker D add 70 mL water. Simultaneously add 6 cm of magnesium ribbon
to each beaker. Note how the rate of reaction in each beaker is
proportional to the concentrations of the acid. 17.5.4.2 Potassium iodide with potassium iodate
reaction
Put 30 mL of solution A containing 5 g of potassium iodate per litre in
a measuring cylinder and dilute to 200 mL with water, then transfer to
a beaker and add drops of starch solution. Put 30 mL of solution B
containing 10 g of hydrated sodium sulfite per litre and 2.5 mL
of 2M sulfuric acid in a measuring cylinder and dilute to 200 mL with
water. Pour this solution B into the beaker containing the diluted
solution A and note the starting time in seconds. Record the time when
a dark blue coloration appears. Repeat the experiment with 5 mL, 7.5
mL, 10 mL, 12.5 mL and 15 mL of 2M sulfuric acid. Plot a graph of
volume of acid to reciprocal of time and note the straight line result.
Reaction 1. is a slow and reaction. Reaction 2. is rapid but does not
occur until
reaction 1. is complete. The presence of iodine as a dark blue
coloration indicates the completion of reaction 1.
1. IO3- + 3HSO3- ---> I-
+ 3HSO4-
2. 5I- + IO3- + 6H+ >
3I2 (s) + 3H2O 17.5.5.0 Law of mass action and reversible
reactions, effect of alteration of concentration
A chemical reaction may stop although some of the reacting substances
remain. If A and B are the reacting substances and C and D are
the resulting substances, an equilibrium occurs with some A and B
remaining unchanged and some C and D formed. Initially A and B react at
a rate which depends on their concentrations. A change in the
concentration of either A or B produces a change in the rate of the
reaction. The rate of the forward action is proportional to the product
of the concentrations of A and B, so rate of reaction of A and B is
proportional to (Concentration of A) X (Concentration of B) = k1
(Concentration of A) X (Concentration
of B). However, as soon as A and B react, their concentrations
decrease,
so the rate of reaction continuously decreases. The reaction between A
and B will have formed some C and D, and the concentrations of A
B will increase and in turn react to form A and B with the rate
of reaction proportional to the product of their concentrations. Rate
of reaction of C and D is proportional to (Concentration of C) X
(Concentration
of D) = k2 (Concentration of C) X (Concentration of D). When there is
apparently
no further action, an equilibrium is reached with the rate of reaction
of A and B forming C and D, equal to the rate of reaction of C and D
forming A and B. At equilibrium, k1 (Concentration of A) X
(Concentration of B) = k2
(Concentration of C) X (Concentration of D). k1 / k2 = the equilibrium
constant, k. So
at equilibrium, the product of the concentrations of C and D, divided
by the product of the concentrations of A and B has a definite value.
Hence, if at equilibrium the concentration of, say, A is increased by
an of more of it, the concentrations of B, C, and D will assume new
values that the value of the expression (K) will remain unchanged. This
will involve the combination of some A and B to form more C and D,
i.e., the previous equilibrium concentration of B will be decreased and
those of C and D will be increased. 17.5.5.1 Hydrolysis of bismuth chloride
Put some bismuth chloride in a test-tube, add 1 mL of water and observe
white bismuth oxychloride forming because equilibrium was reached when
certain concentrations of bismuth oxychloride and hydrochloric acid
form.
BiCl3 + H2O <> BiOCl(s) + 2HCl
[BiOCl] [HCl]2 / [BiCl3] [H2O] = k
(Equilibrium constant)
Add more drops of water and observe the white bismuth oxychloride
reappear.
Add drops of concentrated hydrochloric acid until the white precipitate
disappears because some bismuth oxychloride and some hydrochloric acid
reacted to form more bismuth chloride and water. By these changes the
value of the expression assumed the original mathematical value of K.
17.5.5.2 Hydrolysis of
antimony chloride
Repeat the above experiment using antimony chloride in the place of
bismuth chloride. The reactions and the explanations are similar to
those for bismuth.
SbCl3 + H2O <> SbOCl (s) (antimony
oxychloride) + 2HCl 17.5.5.3 Common ion effect to precipitate sodium
chloride from solution
Add a 3 drops of concentrated hydrochloric acid to a saturated solution
of table salt. Sodium chloride precipitates as white crystals.
Increase in the concentration of the chloride ion favours the backward
reaction with subsequent precipitation of common salt. Increase in the
concentration of the chloride ion favours the backward reaction with
subsequent precipitation of common salt.
NaCl <>Na+ + Cl-
HCl ---> H+ + Cl- 17.5.5.4 Common ion effect to precipitate barium
chloride from solution.
Add a 3 drops of concentrated hydrochloric acid to a saturated solution
of barium chloride. Barium chloride precipitates as white crystals.
Increase in the concentration of the chloride ion favoured the backward
reaction with subsequent precipitation of barium chloride. Increase in
the concentration of the chloride ion favoured the backward reaction
with subsequent precipitation of barium chloride. The increased
concentration of ammonium ion from ammonium chloride
reduces the (OH-) concentration in ammonia solution, NH3(aq)
("ammonium hydroxide")
solution. Add a drop of phenolphthalein to a dilute solution of
ammonium
hydroxide. The solution goes pink. Add solid ammonium chloride drop by
drop until the colour disappears. NH4Cl produces NH4+
ions which increases the speed of the back reaction.
NH4OH <> NH4+ + OH- 17.5.5.5 Effect of temperature on chemical
equilibrium, thermal dissociation of ammonium chloride
Heat 1 mL of ammonium chloride in a dry test-tube damp red litmus paper
fixed inside the mouth. The ammonium chloride sublimes and condenses on
the side of the test-tube, leaving a clear space where the tube is hot.
The damp litmus paper turns blue. The clear space contains the
colourless gases ammonia and hydrogen chloride formed by the
decomposition of the ammonium chloride. Ammonia is lighter than
hydrogen chloride so it diffuses faster and reaches the red litmus
paper first. Recombination to form ammonium chloride occurs in
the cooler part of the test-tube
NH4C1 <> HCl + NH3 (hot >, cold <) 17.5.6.2 Action of heat on nitrogen tetroxide See diagram 17.5.6.2: Heat on nitrogen
tetroxide
Heat a test-tube containing the lead nitrate and also heat the
centre of the long horizontal delivery tube. The colour of the gas in
the hot part of the delivery tube is a darker brown than in the cooler
part of the tube because more brown NO2 molecules are there.
N2O4 (pale yellow) <> 2NO2
(brown) (pale yellow < cooling | heating > brown) 17.5.7.0 Explanation
of group analysis
A saturated solution can remain in equilibrium with undissolved
molecules of the solute. Two equilibria exist: 1. an equilibrium
between the undissolved solute and dissolved molecules, and 2.
an equilibrium between dissolved molecules and ions formed by
dissociation
(XY) <> XY <> X+ + Y-
Equilibrium 1. : (XY) <> XY
Equilibrium 2. : XY <> X+ + Y-
(XY) = undissolved molecules
XY = dissolved but unionized molecules
X+ + Y- = ions
The tendency of the solid to pass into solution depends on its active
mass, solution pressure. If the temperature remains constant, the
active mass remains constant because, by the law of mass action,
concentration
dissolved molecules / concentration undissolved molecules = the
constant, K.
The concentration of dissolved molecules is also a constant.
Also, by the law of mass action, if [concentration X+]
[concentration Y-]
/ [concentration dissolved molecules] = a constant, in a saturated
solution, the product of the concentrations of the ions is a constant.
So if large concentrations of ions are brought together into the same
solution, ions of X+ and Y- will precipitate out
of solution as solid molecules until the concentrations of the
remaining ions in solution have such values that the product of their
concentrations equals the specific constant, the solubility product. Group I
Lead, silver and mercury (I) are precipitated as chlorides by chloride
ions from hydrochloric acid. The concentrations of silver and chloride
ions which can remain in solution are small. So when a solution of a
silver salt containing silver ions is mixed with hydrochloric acid,
most of the silver and chloride ions form molecular silver chloride and
leave the solution as a solid phase until the remaining ions attain
equilibrium.
(concentration silver ions) x (concentration chloride ions) = 1 x 10-10,
the
solubility product.
The solubility product has a constant value, so adding excess
chloride ions reduces the concentration of the silver ions to a
negligible quantity. Only silver chloride, lead chloride and
mercury (I) chloride have low solubility products. So the ions of
other metals remain in solution in the presence of high concentrations
of chloride ions. Group II
Assume hydrogen sulfide is ionized
H2S <> 2H+ + S2-
By the law of mass action (concentration H+)2 x
(concentration S2-)
/ (concentration unionized H2S) = the constant, 1.1 x 10-22.
In a neutral solution, the concentration of sulfide ions is low because
hydrogen sulfide is a weak electrolyte. The concentration of hydrogen
ions is also low. In the acid solution used for Group II, the
concentration of hydrogen ions is increased by the presence of the
strong acid. So to maintain the value of the solubility product
constant, the concentration of the sulfide ion is reduced below its
already small value in neutral solution. However, the amount of sulfide
ions is enough to allow the solubility products of the sulfides of
mercury (II) lead, copper and bismuth to be exceeded. Also, cadmium
sulfide may precipitate if the acid is not too concentrated. So in
Group II all the sulfides of mercury (II) lead, copper, bismuth and
cadmium precipitate.
Solubility products: lead sulfide 4 X 10-28, copper sulfide
8 X 10-45, mercury (II) sulfide 4 x 10-54,
cadmium
sulfide 3.6 X 10-29
manganese sulfide 1.4 x 10-15, zinc sulfide 1.2 X 10-24
The concentration of sulfide ion in acid solution is not enough to
allow the higher solubility products of the sulfides of manganese,
zinc, cobalt or nickel to be reached with any possible concentration of
the metal ion, so these sulfides do not precipitate. They precipitate
later in Group IV, where the precipitating agent is the highly ionized
salt, ammonium sulfide, and the concentration of the sulfide ion from
it is high. Group III
The precipitating agent is ammonia solution.
NH4OH <> NH4+ + OH-
By the law of mass action (concentration NH4+)
(concentration OH-)
/ (concentration unionized NH4OH) = a constant.
Ammonia solution is a weak base, so does not ionize much and the value
of the constant is only about 2 X 10-5.
Most of the ammonia solution will be dissolved but not ionized. The
small hydroxyl ion concentration in a solution of ammonia solution
which is also fairly concentrated with respect to ammonium chloride is
still large enough to cause a precipitation of the hydroxides of ferric
iron, chromium and aluminium, but not great enough to precipitate those
of zinc, manganese, cobalt and nickel. Manganese hydroxide may
precipitate slightly if the concentration of ammonium chloride is not
sufficiently great.
Group IV
In Group II the presence of hydrogen ions from the added acid reduced
the concentration of sulfide ions but this reduced value was enough to
allow the solubility products of the metallic sulfides in the group to
be exceeded. In Group IV, hydrogen sulfide is added to a solution made
alkaline with ammonia solution and so contains excess of hydroxyl ions.
(concentration H+)2 (concentration S2-)
/ (concentration
unionized H2S) = a constant.
The ionic product of water [H+] x [OH-1] = (10-14).
The hydroxyl ions
from the ammonia solution lower the concentration of hydrogen ions
causing an increased concentration of S2- ion. In Group IV,
the metal
sulfides not already precipitated in Group II because of their high
solubility products, are here precipitated. So the ionic concentration
of the sulfide ion is controlled by variation of the concentration of
the ions it is associated with, i.e. hydrogen ions.
Group V
The metals still remaining in solution include barium, strontium,
calcium, magnesium, sodium and potassium. Barium, strontium and calcium
are precipitated as carbonates by the addition of CO32-
ions in
alkaline solution because of the low values of the solubility products
[X2+] [CO32-] = K, between 10-8
and 10-9. The solubility product of
magnesium carbonate is low, 10-5,
and in neutral solution would be
precipitated, but its precipitation is prevented in this group by
the ammonium ions, mainly from the ammonium chloride added
before Group III. With the large NH4+ ion
concentration
from this source, the concentration of CO32- ions
is
reduced below the concentration to precipitate magnesium as
a carbonate.
[NH4+]2 [CO32-]
/ [NH4)2CO3] = K 17.6.1 Weight of iron in
iron (II) ammonium sulfate.
Dissolve 2 g of iron (II) ammonium sulfate in 50 mL of
demineralized
water. Heat to boiling then add 2 mL of concentrated nitric acid and
continue boiling for two minutes. Leave to cool. Add ammonium hydroxide
and stir until precipitation is complete. Test with litmus paper. Heat
again to boiling and filter. Use a wash-bottle to wash any iron (III)
hydroxide from the beaker and stirring rod. Leave the filter paper and
filtrate to dry. Weigh a crucible. Transfer the dry filter paper and
precipitate to the crucible. Ignite the filter paper, leave to cool and
weigh. Heat the crucible until it becomes red, leave to cool and weigh
until again until the weight remains constant.
2Fe2O3 > 4Fe + 3O2
Molecular weight of Fe2O3 = 160
Molecular weight of Fe = 56
So 2 Fe, 2 x 56 is equivalent to 1 Fe2O3, 160
Let weight of iron (II) ammonium sulfate = 1.96 g
Let weight of iron (III) oxide = 0.40 g
So weight of iron in iron (III) oxide = 112 x 0.40 / 160 = 0.28 g
% weight of iron in iron (II) ammonium sulfate = (0.28 / 1.96) X 100 =
14.29%. 17.6.2 Weight of
aluminium in aluminium sulfate.
Dissolve 5 g of aluminium sulfate in 50 mL of demineralized water. Add
drops of dilute sulfuric acid and heat to boiling. Add excess of
ammonium hydroxide and heat to boiling. Filter, leave to
dry, ignite the filter paper in a crucible, leave to cool, then weigh.
Al2SO4 + NH4OH > Al(OH)3
+ NH4SO4
AL(OH)3 > Al2O3 + H2O
Molecular weight of Al2O3 = 102.
Let weight of aluminium sulfate = wa g,
Let weight of aluminium oxide = xb g.
So weight of aluminium in oxide = (54 / 102) x wb g.
% weight of aluminium in aluminium sulfate = [(54 x wb) / 102] / wa x
100%. 17.6.3 Weight of calcium
in marble, calcium carbonate.
Stir 1.5 g of marble powder into 10 mL of demineralized water
then add 5 mL
of concentrated hydrochloric acid. Heat the solution until the marble
dissolves, then add water to make the volume 50 mL. Add ammonium
solution until the
solution is alkaline by litmus test, then heat to boiling. Add 3 g of
crushed ammonium oxalate, stir, and again heat to boiling. Leave the
solution to form a precipitate. Wash the liquid and precipitate
into a filter paper. Wash the precipitate with water until the filtrate
contains no more chloride ions as tested with silver nitrate solution.
Dry the precipitate of calcium oxalate, ignite the filter paper in a
furnace, leave to cool, then weigh the final residue of calcium
oxide.
CaCO3 + 2HCl > CaCl2 + CO2 + H2O
CaCl2 + (NH4)2C2O4
> CaC2O4 + NH4Cl
CaC2O4 + "O" > CaO + 2CO2
(The O is from air)
Molecular weight of CaO = 56
Let weight of calcium carbonate = wa g.
Let weight of calcium oxide = wb g.
So weight of calcium in calcium oxide = (40 / 56) x wb g.
% weight of calcium in marble = (40wb / 56wa) X 100%. 17.6.4 Weight of
magnesium in magnesium sulfate
Dissolve 2 g of magnesium sulfate in 50 mL of demineralized water. Add
10 mL of ammonium chloride solution and then ammonium
hydroxide until alkaline after stirring. If magnesium hydroxide
precipitates add more ammonium chloride. Heat
and add sodium phosphate solution in excess, stir, and leave to settle.
Filter then wash the precipitate with ammonium hydroxide to
remove ammonium chloride. Test presence of chloride with silver nitrate
solution
acidified with nitric acid. Heat the solid gradually then strongly,
leave to cool, then weigh.
MgSO4 + NH4OH + Na2HPO4
---> Mg(NH4)PO4 + Na2SO4
+ H2O
magnesium ammonium phosphate
2Mg(NH4)PO4 ---> Mg2P2O7
+ 2NH3 + H2O
Molecular weight of Mg2P2O7, magnesium
pyrophosphate = 222
Let weight of magnesium sulfate = aw g.
Let weight of magnesium pyrophosphate = bw g.
So weight of magnesium in the pyrophosphate = (48 /222) x bw g.
% magnesium in magnesium sulfate [(48bw /222)] / aw X 100%.
17.6.5 Weight of sulfate
radical in sodium sulfate
Dissolve 3 g of sodium sulfate crystals in 50 mL of demineralized
water. Add 5 mL dilute hydrochloric acid and 5 mL of ammonium chloride
solution, then heat to boiling. Add excess barium chloride solution to
assist precipitation of barium sulfate, then and heat to boiling.
Leave to settle and then decant the liquid into filter paper. Wash the
precipitate onto a filter paper then dry in an oven. Add a drops of
concentrated nitric acid then heat the crucible until it becomes
red. Leave to cool then weigh.
Molecular weight of BaSO4 = 233
Let weight of sodium sulfate crystals = aw g.
Let weight of barium sulfate = bw g.
So weight of sulfate in barium sulfate = (96 / 233) x bw g.
% weight of sulfate in sodium sulfate = [(96 bw / 233) / aw] x 100%. 17.6.6 Weight of tin in
solder.
Make finely divided solder with a rough file. Add 10 mL of concentrated
nitric acid to 1 g of finely divided solder in an evaporating
basin. Heat gently in a fume cupboard. When the reaction stops, add
two drops of nitric acid and heat gently again, until no more nitrogen
dioxide forms. Dilute the contents to 50 mL and filter. Wash the
precipitate with dilute nitric acid. Dry the precipitate, ignite the
filter paper, leave to cool, then weigh the final product. The tin was
first oxidized to meta stannic acid, hydrated tin (IV) oxide with
formula H2Sn5O11 or SnO2.xH2O.
It is used as an opacifying colour in ceramics and an abrasive in the
glass industry.
5Sn + 20HNO3 > H2Sn5O11
+ 20NO2 + 9H2O
H2Sn5O11 > 5SnO2 + H2O
Molecular weight of SnO2 = 151
Let weight of solder = aw g
Let weight of tin (IV) oxide = bw g
So weight of tin in tin (IV) oxide = 119bw / 151 g
% of tin in solder = (119bw /151aw) x 100.