Name: _____________________________
Student number: _____________________________
Laboratory Day: _____________________________
Laboratory number: _____________________________
Seat number: _____________________________
SCHOOL OF CHEMISTRY
UNIVERSITY OF KWAZULU-NATAL
DURBAN CENTRE
Chemical Reactivity
CHEM 120
LABORATORY MANUAL
2nd SEMESTER 2010
1
UNIVERSITY OF KWAZULU-NATAL
DURBAN CENTRE
SCHOOL OF CHEMISTRY
I, the undersigned (please print your full name):
___________________________________________________________________
Student Number: ____________________
do hereby acknowledge having read and understood the documents headed Occupational Health
and Safety and Laboratory Regulations. Furthermore, I accept that contravention of these rules and
regulations may lead to my expulsion from the laboratory class, or classes, with subsequent loss of
my Duly Performed (DP) certificate.
I agree to abide by any additional laboratory regulations or safety rules presented in writing in the
laboratory manuals/books or issued verbally by the lecturer-in-charge, or other responsible member
of staff, during pre-laboratory lectures or in the laboratory.
In addition, I understand that I must attend at least 80% of the scheduled laboratory classes and that
failure to do so, irrespective of the reasons, may result in the loss of my DP certificate.
DATE:
___________________________
SIGNATURE: ___________________________
2
Table of contents
Occupational Health and Safety ..........................................................................................................3
Laboratory Regulations .......................................................................................................................4
General Advice
.............................................................................................................................5
Safety Precautions .............................................................................................................................6
General Fire Orders .............................................................................................................................8
Experiment 1:
Purification of an impure organic compound and determination of melting
points....................................................................................................................9
Experiment 2:
Acetylation of aniline using acetic anhydride ...................................................14
Experiment 3:
Esterification of 1-pentanol with acetic acid .....................................................19
Experiment 4:
Reactions of functional groups ..........................................................................26
Reactions of the cations of the metallic elements: Qualitative analysis ............................................41
Experiment 5:
Reactions of the cations of calcium, aluminium, chromium, tin and lead.........42
Experiment 6:
Reactions of the cations of iron, cobalt, nickel, copper, and zinc .....................47
Experiment 7a:
The pH meter and potentiometric titrations .......................................................54
Experiment 7b:
Equilibria of water, weak acids and bases, and buffer solutions .......................64
Experiment 8:
Solubility product of a slightly soluble salt .......................................................70
Experiment 9:
Freezing point depression ..................................................................................76
Experiment 10a: Determination of the molar mass of copper by electrolysis ..............................86
Experiment 10b: The production of an electric potential by means of oxidation-reduction
reactions .............................................................................................................94
Appendix 1:
Laboratory apparatus .......................................................................................101
Appendix 2
The Laboratory Balance ..................................................................................102
Appendix 3:
Volumetric apparatus .......................................................................................105
Appendix 4:
Experimental errors .........................................................................................110
Appendix 5:
The elements ....................................................................................................118
Appendix 6:
Common Solvents ............................................................................................120
Health and Safety
3
Occupational Health and Safety
You are warned that all substances handled and all operations performed in a laboratory can
be hazardous or potentially hazardous. All substances must be handled with care and
disposed of according to laid down procedures. All operations and manipulations must be
carried out in an organised and attentive manner.
In order to assist you in developing good and safe laboratory techniques, a set of Laboratory
Rules and Regulations is attached. You are required to read these and to acknowledge that
you have read and understood them. Additionally, in the laboratory manuals/books and/or
pre-laboratory lectures your attention will be drawn to the correct and safe handling of
specific chemicals/reagents/solvents, and to the correct/safe manner in which specified
laboratory operations must be carried out. These specific instructions and/or warnings must
never be ignored.
It is a legal requirement that
SAFETY GLASSES,
LABORATORY COATS
and CLOSED SHOES
are worn in the laboratory at all times.
Note:
1.
Sunglasses (normal or prescription) may NOT be worn as a substitute for safety
glasses. Prescription glasses (except sunglasses) are acceptable, but MUST be worn at
ALL TIMES.
2.
Some types of contact lenses should not be worn in the laboratory. Students who wear
contact lenses must check the risk factor with their lens supplier.
3.
In addition to being closed, shoes must be sensible – HIGH HEELED SHOES are
HAZARDOUS.
4.
The School requires students to remove, or to make safe, headgear that is considered
dangerous or a potential hazard.
5.
The School requires students with long hair to tie it back.
Laboratory Regulations
4
Laboratory Regulations
1.
Students shall present themselves ten minutes before the start of each scheduled laboratory
session. Latecomers will be refused entry to the laboratory.
2.
No student is permitted to work in the laboratory outside scheduled laboratory hours.
3.
Students are not allowed to enter the preparation room, which is located along the side of
the laboratory. If reagent bottles need to be refilled, broken apparatus replaced, etc.,
students should request assistance from a demonstrator.
4.
Apparatus and chemicals are NOT to be removed from the laboratory.
5.
You will find the laboratory bench clean upon your arrival, and it should be clean when you
leave the laboratory. Bench tops should be wiped and glassware and other apparatus should
be left clean and dry.
6.
Balances and other expensive equipment must be treated with care and kept clean and tidy
at all times. Do not spill chemicals on the balance pan!
7.
All solids must be discarded in the bins at the outer ends of each bench. Do not throw
matches, paper, or any insoluble chemicals into the sink. Liquids must be discarded into the
ceramic sinks or designated disposal bottles.
8.
All students are required to wear a laboratory coat, and no student will be permitted to work
in the laboratory without one.
9.
All students who do not wear conventional spectacles must wear a pair of safety spectacles.
No student will be permitted to work in the laboratory without eye protection.
10.
All students must wear closed shoes in the laboratory.
11.
All students must have a laboratory towel to dry apparatus and clean the bench top.
12.
No food or drink is allowed in the laboratory. Eating is not permitted in the laboratory.
13.
Cell phones must be switched off whilst you are in the laboratory.
General Advice
5
General Advice
In order to work quickly and accurately, students should carefully plan their work before coming
into the laboratory. A schedule of the experiments to be performed will be posted on the notice
board and students are expected to read the relevant portions of the notes in their laboratory
manuals before their practical session.
The pre-laboratory problems on the green sheets should be completed at home prior to the relevant
laboratory. These exercises are designed to familiarise you with certain aspects of the theory of the
experiment you are to carry out, as well as giving you practice in the calculations involved. It is
thus very important that you complete them before coming to the laboratory.
The laboratory
exercises will contain questions which are very similar to those found in the pre-lab exercises.
These exercises thus serve as preparation for the laboratory exercises, and it is thus in your own
interest to ensure that you have mastered the material. You will not be allowed entry into the
laboratory unless you have completed your green sheet beforehand. The demonstrators will
check that these sheets have been completed satisfactorily.
You must record all your results neatly in ink on the sheets provided. If you forget your laboratory
manual, borrow a friend’s and make a copy of the relevant sheets before coming to the laboratory.
All results sheets for a particular laboratory must be handed in at the end of that session; students
who do not do so will be deemed to have been absent, with possible subsequent DP implications.
All absences from practicals will automatically be graded as 0 unless a suitable written excuse
(medical or other) is furnished. Written excuses should be provided within one week of reattendance, or they will not be accepted.
Please keep in mind that a DP certificate will be refused to any student who has not attended the
required minimum number (80%) of laboratory sessions, irrespective of the reasons for absences.
Safety Precautions
6
Safety Precautions
The chemical laboratory is not a place for horseplay. Do not attempt unauthorised experiments or
practical jokes on your neighbour. Such activities are dangerous and can cause serious injuries.
Report all accidents, cuts, burns, etc. - however minor - to your demonstrator or staff member in
charge. Eyewash stations are located in several places in the laboratory. See that you know where
the nearest one to your bench is located in case of an accident.
Liquids – whether corrosive or not – must be handled with care, and spilling on the bench or floor
should be carefully avoided. Any spillage must be cleaned up at once. If a corrosive liquid, such as
an acid or base, is spilled, call your demonstrator or the staff member in charge.
Reagent bottles must be stoppered immediately after use and returned to their correct place. It is
absolutely forbidden to introduce anything into reagent bottles, and solutions taken from reagent
bottles should never be returned to the bottles. Do not lay the stopper of a reagent bottle on the
desktop – it could become contaminated. The correct procedure for pouring liquids from reagent
bottles is described below.
Hold the stopper in the bottle, and tilt the bottle slightly to
wet the stopper. This lubricates the ground glass and
makes removing the stopper easier.
Moisten the inside of the neck and the lip with the
stopper. This stops the first drops from gushing out when
pouring. Replace the stopper.
Remove the stopper again by turning your hand over and
holding the stopper between two fingers. The neck of the
bottle should touch the edge of the vessel you are pouring
into to prevent liquid from running down the outside of
the bottle.
The stopper must remain firmly held between the fingers
whilst pouring the liquid. Replace the stopper when
enough liquid has been poured out.
Do not heat graduated cylinders or bottles because they can easily break. Heat all other apparatus
gently at first to avoid breakage.
Do not put anything in your mouth while working in the laboratory, nor taste chemicals or
solutions!
Safety Precautions
7
Breakages of expensive items of glassware such as burettes, pipettes, thermometers, graduated
cylinders, etc. will be charged for. Examination results can be withheld at the end of a semester
until such charges have been settled.
Safety Precautions
8
General Fire Orders
These orders should be read in conjunction with any fire fighting instructions that are displayed in
the laboratory.
In the event of a fire:
Alert your demonstrator or staff member in charge (if they haven’t already noticed…) and
obey any instructions that they give you.
On hearing a fire evacuation alarm:
Stop normal work immediately. Make any apparatus safe – turn off Bunsen burners, stirrers,
vacuum pumps, etc.
Unless your demonstrator or staff member in charge has given you any other special
instructions, follow the green emergency exit signs out of the building. Assemble on the
grassed area between J and L blocks.
You should make sure that you know the location of the fire extinguisher in your lab.
Experiment 1
9
Experiment 1: Purification of an impure organic compound and
determination of melting points
AIM
To introduce the elementary technique of removal of impurities; to illustrate the techniques of hot
filtration and crystallisation; and to provide practice in the determination of the melting point of a
solid.
INTRODUCTION
Removal of impurities
An insoluble impurity is easily removed from a solution by means of filtration. Similarly, a soluble
impurity can be removed if it is first extracted from solution by adsorption onto a suitable solid.
Activated charcoal is a very good adsorbent by virtue of its large surface area (200 m2 g-1). It is used in
gas masks to adsorb noxious gases such as CO, CO2, and COCl2, and in the wine and gelatine
industries to deodorise and decolorise.
In this experiment activated charcoal will be used to adsorb a soluble dye from an aqueous solution of
an organic compound. The organic compound is sparingly soluble in cold water (0.5 g per 100 cm3 at
10 °C) and appreciably soluble in hot water (2.7 g per 100 cm3 at 90 °C). Thus it is imperative that
filtration be done on a hot solution in a preheated filter.
Crystallisation
To obtain a uniform crystalline product from solution, three conditions are necessary:
(a)
having a reasonable concentration at a high temperature
(b)
ensuring a slow rate of cooling
(c)
preventing evaporation of solvent.
In this experiment quantities of solute and solvent have been selected to fulfil requirements (a) and
(c).
Melting points
Most organic substances consist of molecules held together by covalent bonds. The crystals are held
together by weak van der Waals attractions between the molecules and thus have low melting points in the range 30 to 360 °C.
The melting point of a pure substance is defined as that temperature at which its crystalline state
co-exists in equilibrium with its liquid state. At this temperature any addition of heat energy will
cause the liquid to increase at the expense of the crystalline solid, and any removal of heat energy
will cause the amount of solid to increase at the expense of the liquid. With very good equipment it
can be shown that there is no change in temperature as solid turns to liquid and vice versa. With the
equipment available in this experiment it will suffice to record the melting point as that temperature
range from when the solid first melts to the point at which the last trace of solid disappears. If the
substance under test is pure, the change from all solid to all liquid will occur within a temperature
change of 0.5 °C. If impure, it does not have a melting point - it has a melting range often
exceeding 5 °C in extent. The melting point of a pure substance is as characteristic of that substance
as is its density, refractive index, boiling point, etc. and thus can:
(i)
prove a substance to be pure or impure
(ii)
identify an unknown substance.
Experiment 1
10
The melting points of all known substances are recorded in books of tables, making identification of an
unknown substance easy. Once a match has been obtained between the melting point of the substance
under test and a listed substance, further confirmation of identity is possible. Some of the substance
under test is mixed with its listed match and the melting point of the mixture is determined. If the
mixture has the same melting point as the substance under test, then the identity of the unknown is
established beyond doubt. If, however, the mixture melts at a lower temperature range, then the
unknown substance has not been correctly identified.
NOTES
Filtration
The separation of solids from suspensions requires a septum of suitable pore size that will allow the
molecules of the solvent and solute to pass whilst retaining particles of insoluble substances.
Normally a piece of filter paper is sufficient, but in organic chemistry it is sometimes necessary to
build up a composite septum of several layers in order to retain amorphous solids. A composite
septum usually consists of a piece of filter paper on which is spread an even layer of filter aid.
Common filter aids are: paper pulp, asbestos fibres, glass wool, diatomaceous earth, etc. In this
experiment a diatomaceous earth called kieselguhr will be used.
Because a composite septum is rather thick, filtration tends to be slow. In order to speed up the rate of
filtration, a vacuum is applied to the underside of the septum as shown in the Figure below.
To vacuum tap The procedure for preparing and using a composite septum is:
(i)
(ii)
(iii)
(iv)
(v)
(vi)
(vii)
(viii)
(ix)
make a slurry of two heaped spatulae of kieselguhr in approximately 15 cm3 of water
place a Buchner funnel in the receiver
connect the side-arm of the receiver to the vacuum pump
open the water tap of the vacuum pump fully and leave running till end of filtration
put a circle of filter paper on the funnel and wet it with a few drops of water; ensure the
paper lies flat and covers the perforations completely
pour the slurry of kieselguhr onto the filter paper and cover with another piece of filter
paper
close the relief valve and suck dry
open the relief valve and empty the receiver
if a hot filtration is to be done, rest the funnel on a beaker of hot water until required.
NOTE: Buchner funnels and Hirsch funnels are expensive and must be handled with care.
Experiment 1
11
Determination of melting point
The apparatus used to determine melting point consists of a heat source around a space holding three
sample tubes (under a magnifying glass), and a thermometer. The heating device consists of a boost
heater which serves to raise the temperature quickly to within 20 °C of the expected melting point, at
which point it is switched off. Thereafter, the infinitely variable heater is switched on and set to raise
the temperature at the desired number of degrees per minute.
M
Melting Point Apparatus
Usually a quick run is done at 10 °C per minute to obtain an approximate melting point. The apparatus
is then cooled by means of the cooling plug to 20 °C below the roughly determined melting point.
Now a second and much slower run is done on a fresh sample at a rate of 2 °C per minute temperature
rise in order to obtain an accurate melting point.
For practice in melting point determination two substances A and B are supplied. This practice can
be done during the period when the filtrate is cooling and crystals are forming.
A small quantity of A is crushed on a watch glass and introduced into a melting point tube to a depth
of about 5 mm. A second tube is similarly charged with B and a third tube with a mixture of
approximately equal proportions of A and B (provided).
All three tubes are inserted in the apparatus and a preliminary run done. A and B will melt at their
respective melting temperatures. The mixture of A and B will melt at a temperature lower than
either A or B and the melting process will be spread over a range of some five degrees, illustrating
the certainty with which purity or identification mentioned in the Introduction can be inferred.
A second and slower run on fresh samples of A and B will allow the two melting points to be
accurately determined. The booster heater in this run must be switched off at a temperature of about 20
°C below the melting point obtained in the preliminary run.
Experiment 1
12
Remember to cool the instrument between runs by making use of the cooling plug as described
above.
EXPERIMENTAL PROCEDURE
1.
Prepare a Buchner funnel as described in the filtration notes.
2.
Empty the vial of sample into a 50 cm3 conical flask, add 25 cm3 of water, add 4 anti-bumping
granules and bring to the boil. Swirl frequently to avoid "bumping".
3.
Remove burner and allow two minutes for the contents to cool. Add approximately
0.1 g of activated charcoal and boil for two minutes, with swirling.
4.
Filter hot through the previously prepared Buchner funnel fitted with an empty, clean receiver.
The filtrate should be clear and colourless. DO NOT WASH. Open the relief valve and close
the water tap.
5.
Preheat a 50 cm3 conical flask in the beaker of hot water. Reheat the filtrate in the Buchner
receiver and transfer to the warm conical flask. Cover with foil and set aside to cool to room
temperature.
6.
Practise the determination of melting point as described in “Determination of melting point”
above.
7.
When the conical flask and contents have reached room temperature transfer to an ice bath for
10 minutes.
8.
Prepare a Hirsch funnel by putting a damp piece of filter paper on the perforated bed and
mounting the funnel on the Buchner receiver. Open the water tap of the vacuum pump fully
and close the relief valve.
9.
Swirl the contents of the conical flask and pour the crystals into the Hirsch funnel. If it is
necessary to rinse the flask in order to get all the crystals onto the funnel, use the filtrate for
this purpose.
10.
Invert the Hirsch funnel over a square of absorbent paper and scrape all the crystals onto
it,
cover with another square of paper and gently squeeze out excess moisture.
11.
Transfer the two squares of paper and crystals to a clean marked 50 cm3 beaker and dry in the
oven for 10 minutes.
12.
Determine the melting point of the crystals.
13.
Weigh a 50 cm3 beaker and ask a demonstrator to verify the reading entered on your report
sheet. Transfer the crystals to the beaker and ask a demonstrator to verify the mass as you
re-weigh it.
Experiment 1
13
Experiment 1: Purification of an impure organic compound and
determination of melting points
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Practical Mark:
Laboratory
Mark:
RESULTS
Mass of vial + compound + lid
_____________________ g
Mass of empty vial + lid
_____________________ g
Mass of pure compound
_____________________ g
Melting point range of pure compound
_____________________ °C
Rough melting point of A
_____________________ °C
Accurate melting point range of A
_____________________ °C
Rough melting point of B
_____________________ °C
Accurate melting point range of B
_____________________ °C
Rough mixed melting point range of A and B
_____________________ °C
N.B.: Demonstrator's initials must be obtained for all weighings.
Experiment 2
14
Experiment 2: Acetylation of aniline using acetic anhydride
AIM
To introduce the procedure of acetylation, to introduce the technique of refluxing and to introduce the
concept of yield in an organic synthesis.
INTRODUCTION
Many inorganic reactions are instantaneous and quantitative. Many organic reactions, however,
proceed slowly, and, unconsumed reactants are still present as impurities in the final equilibrium
mixture.
The acetylation of an amine is a good example of this.
RNH2 + (CH3CO)2O
→ RNHCOCH3 + CH3COOH
At equilibrium, the reaction mixture contains the desired product (acetanilide), as well as an unwanted
by-product (acetic acid), unreacted aniline and unreacted acetic anhydride. The ratio of the mass of
desired product obtained experimentally to the mass obtainable stoichiometrically is a measure of the
yield.
A yield of 60% is typical not only because of the incompleteness of the reaction but also because of
loss of desired product in the various processes needed to remove unwanted products and unconsumed
reactants.
NOTES
Handling of Ground Glass Joints
The apparatus used in this experiment is fitted with ground glass joints. When used correctly, these
joints are leak-proof. Both surfaces must be clean and free of grit. Surfaces must be engaged and
disengaged with a turning motion. When caustic alkalies are used in preparations (e.g. saponification
reactions), the apparatus must be dismantled immediately after conclusion of the experiment and
thoroughly washed to prevent etching of the ground surfaces by the alkali. In this case, a THIN layer
of vacuum grease may be used to prevent the glassware from sticking. All clamps used to support
glass apparatus must be rubber covered or plastic covered, and minimum pressure should be used in
fastening the clamps.
The Reflux Condenser
When used in the vertical position (as shown in the figure), a condenser serves to return volatile
reactants and products to the reaction vessel when the reaction is carried out at elevated temperature.
The cooling water enters at the lower side-arm and leaves at the upper. The flow rate of cooling water
must be set at approximately 1/2 litre per minute and must be checked at frequent intervals. When the
flow rate is correctly set, condensation of vapours will be seen to occur in the lower third of the
condenser.
Engaging Flask and Condenser
The reaction flask is clamped in the fixed clamp of the bracket and the condenser in the movable
clamp. In this way the condenser can be lowered and raised to effect a sliding engagement or
disengagement of the reaction flask.
Experiment 2
15
When inserting the condenser into its clamp, great care must be taken to ensure that the upper side arm
is not pressed against the bracket.
Regulating the Heating
The volume of reagents used is small and heating must be carefully controlled. It is possible to turn the
flame down by partially closing the gas tap, but this is a hazardous practice as the flame is likely to
"strike back". Control is far better exercised by altering the height of the wire gauze above the flame.
The most suitable position is that where only the tip of the flame touches the wire gauze. If, at this
setting, boiling is not sufficiently vigorous, the ring supporting the wire gauze can be lowered a
centimetre or two and then the bracket holding the condenser and reaction flask can be lowered until
the flask again touches the wire gauze. At a suitable position, the vapours will be seen condensing in
the lower third of the condenser. The tubes carrying cooling water must be arranged so that they do not
touch the hot gauze.
EXPERIMENTAL PROCEDURE
1.
Wash the 10 cm3 measuring cylinder, the pear-shaped flask and the condenser with ethanol and
then with ether, IN THE FUME CUPBOARD. Pour the washings into the waste bottle
provided.
2.
Clamp the bracket to the upright by means of the bosshead. Make sure the fixed clamp is
below and the movable clamp is above.
3.
Clamp the condenser in the movable clamp of the bracket as shown in the figure below.
Connect the cooling water tubes and set the flowrate of water. Measure 2 cm3 of aniline into
the measuring cylinder and then pour this into the pear-shaped flask.
4.
Next measure 2.5 cm3 of acetic anhydride into the measuring cylinder. Pour this slowly into
the flask whilst stirring. Measure 2 cm3 of glacial acetic acid and add to the flask as well. Add
3 anti-bumping granules and install the flask in the fixed clamp of the bracket.
5.
Slide the condenser down till engaged in the neck of the reaction flask.
6.
Boil the mixture for 30 minutes. Adjust the boiling rate as described in the notes section.
Whilst boiling carry out steps 7 - 9 below.
7.
8.
9.
Pour approximately 90 cm3 of deionised water into a 100 cm3 beaker and cool in an ice bath.
Prepare a composite septum in a Buchner funnel as described in Experiment 5.
Prepare a Hirsch funnel.
Experiment 2
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
16
After boiling the reaction mixture for 30 minutes, shut off the burner, disengage the condenser,
shut off the condenser cooling water flow, unclip the flask and rest it in the wooden stand until
cool enough to handle.
Pour 10 cm3 of the cold water prepared in 7 above into the measuring cylinder.
Pour the contents of the flask into the remaining 80 cm3 of cold water slowly and with stirring.
Filter through the Hirsch funnel. Use the 10 cm3 of cold water in the measuring cylinder in two
lots to rinse the beaker and to wash the crystals.
Transfer the crystals to a 50 cm3 wide-necked conical flask, add approximately 30 cm3 of
water and boil with frequent swirling till complete redissolution has occurred. Any brown
globules visible at the bottom of the flask are unreacted aniline; this will be removed by the
activated charcoal in step 15.
Remove the burner, let the boiling subside, add 1 spatula of activated charcoal and re-boil
CAREFULLY for approximately one minute.
Insert the prepared, pre-heated Buchner funnel into the receiver and filter the contents of the
flask hot.
Heat the contents of the receiver to redissolve the crystals, transfer the hot solution to a beaker,
cover with foil and set aside to cool. When it has cooled to room temperature, cool the beaker
further in an ice bath for 10 minutes.
Filter off the crystals into a Hirsch funnel, transfer to absorbent paper, blot dry, place in a
marked beaker and dry in the oven for 15 minutes.
Determine the melting point of the product, as explained in Experiment 5.
Transfer the crystals to a pre-weighed 50 cm3 beaker and reweigh. Ask a demonstrator to
verify the weighings and to initial these entries on your report sheet.
Show your sample to a demonstrator for marking.
Experiment 2
17
Experiment 2: Acetylation of aniline using acetic anhydride
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC CALCULATION
C6H5NH2 + (CH3CO)2O
→ C6H5NHCOCH3 + CH3COOH
aniline
acetanilide
Molecular formula
C6H5NH2
C6H5NHCOCH3
Molar mass
_________ g mol-1
_________ g mol-1
Density
1.02 g cm-3
Volume
2.0 cm3
Expected mass of acetanilide 
 ____________ g
2.0 g aniline 1 mol acetanilide
x
x
g mol 1
1 mol aniline
g mol 1
Experiment 2
18
Experiment 2: Acetylation of aniline using acetic anhydride
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Practical Mark:
Laboratory
Mark:
RESULTS
Mass of weighing bottle + amide
=
________________ g
Mass of empty weighing bottle
=
________________ g
Mass of amide
=
________________ g
% Yield
=
________________ %
Melting point of product
________________ °C
Experiment 3
19
Experiment 3: Esterification of 1-pentanol with acetic acid
AIM
To introduce the procedure of esterification using an acid as a catalyst, to introduce the use of the
separating funnel for washing a product to remove unused reagents and catalyst, to introduce the use of
an anhydrous salt for drying a product, and to introduce the technique of distillation to recover a
product in pure form.
INTRODUCTION
Esterification is the reaction of a carboxylic acid and an alcohol with the elimination of water to form
an ester.
RCOOH + HOR → RCOOR + H2O
The rate of reaction is slow, but reaching equilibrium can be speeded up by the application of heat and
by the addition of a catalyst, such as a small quantity of concentrated sulphuric acid.
The equilibrium of the reaction can be shifted to the right, i.e. the yield of ester can be improved, by
increasing the active mass of one of the reactants or by removing the unwanted product, water.
In this experiment a 30% excess of acetic acid will be used. Thus the calculation of percentage yield
must be based on the mass of the alcohol used (alcohol is the limiting reagent).
At equilibrium, the reaction flask will contain:
(a)
pentyl acetate (the wanted product) (also called pentyl ethanoate)
(b)
1-pentanol (unreacted reagent)
(c)
acetic acid (unreacted reagent) (also called ethanoic acid)
(d)
sulfuric acid (catalyst)
By exploiting differences in physical properties, the unwanted constituents can be removed in three
steps:
(i)
The two organic substances, 1-pentanol and ester, are insoluble in water and have a lower
density than water, whereas acetic and sulphuric acids are of higher density than water and
infinitely soluble in water. The 2 phases (soluble in water and insoluble in water) are separated
with a separating funnel which is described below. The removal of the acids is also described
below.
(ii)
Traces of water can be removed by certain anhydrous salts that absorb water but do not absorb
esters. This means of dehydration is described below.
(iii)
1-Pentanol and the ester are of similar density and are soluble in one another. Thus it is
necessary to exploit some other difference in physical properties to separate them. 1-Pentanol
boils at 138 °C whereas the ester boils at 148 °C. Thus distillation is used for the final
purification as described below.
NOTES
Esterification
Care is needed in the addition of concentrated sulfuric acid to the mixture of acetic acid and pentanol.
Esterification takes place whilst the mixture is boiled under reflux exactly as in Experiment 6. The use
of anti-bumping granules here is essential.
Experiment 3
20
The separating funnel
On completion of esterification, the contents of the flask must be cooled before being washed into the
separating funnel. In the funnel, two distinct layers will form: the upper layer of unreacted 1-pentanol
and pentyl acetate (the organic phase) being of a lower density, will float on the lower layer (the
aqueous phase) of dilute sulfuric acid and acetic acid. The interface between the two layers is clearly
visible. By opening the stopcock, the aqueous phase can be drained into a beaker. (It is advisable to
run the aqueous phase into a small beaker so that if the stopcock is not closed in time and some of the
organic phase has passed through the stopcock, it can be returned to the funnel and re-separated.) As
much of the aqueous layer as possible is drained off without loss of organic phase.
Now the organic phase is washed several times with water to remove residual acid. This is carried out
as follows:
Approximately 20 cm3 of water is poured into the funnel, and the funnel is stoppered and shaken to
enable acid absorbed in the organic phase to transfer to the aqueous phase. It is not advisable to shake
too vigorously as this will cause emulsification of the ester in the water and consequent loss of
product. It is best to hold the funnel with the thumb and third finger straddling the body of the funnel
and the index finger holding the stopper in place, and turn the funnel upside down, right side up,
upside down a few times. Then the funnel must be held upside down while the stopcock is opened to
release pressure. After closing the stopcock, the shaking cycle with pressure release is repeated.
After one minute's shaking, the funnel is clamped in the upright position to allow the phases to
separate. Then the aqueous phase is drained off and a second and a third wash can be done. Three
water washes will remove most of the acid, but traces of acid in the ester have to be removed by
neutralisation. In order to neutralise the acid, the organic layer is washed with a saturated solution of
sodium bicarbonate. Reaction between bicarbonate and acid will release carbon dioxide gas, so
pressure release must be done frequently.
The release of gas can be used to decide whether all the acid has been neutralised. When, on turning
the funnel upside down and opening the stopcock, there is no escape of gas then one can conclude that
either all the acid has been neutralised, or all the bicarbonate has been consumed. If the aqueous phase
is now run into a beaker it can be tested with litmus paper. If the red litmus paper turns blue, then there
was sufficient bicarbonate to neutralise the acid. If, however, the red litmus paper does not change to
blue, then there has been insufficient bicarbonate to neutralise the acid and another bicarbonate wash
must be carried out.
To remove the traces of sodium bicarbonate a water wash is used. After draining this aqueous layer,
the funnel must be tapped sharply to allow the last traces of water to sink. Then the last drop or two of
water must be drained. Now the organic phase is ready for dehydration.
Experiment 3
21
Dehydration
Calcium chloride, the well-known dehydrating agent, cannot be used here as it absorbs esters.
Anhydrous sodium sulfate is suitable for use here. The organic phase is POURED from the separating
funnel into a 50 cm3 conical flask and anhydrous sodium sulfate is added whilst swirling vigorously. A
contact time of a couple of minutes must be allowed for dehydration and the flask must be swirled
frequently. A clear solution indicates a dried product.
Fluting a filter paper
Esters tend to swell the fibres of a filter paper and then the paper becomes glued to the surface of the
funnel, making filtration very slow. To overcome this, the filter paper is fluted. This style of folding a
filter paper will be demonstrated during the practical.
After sufficient contact with the anhydrous sodium sulfate, the organic phase is filtered through a
fluted filter paper into a clean, dry round-bottom distillation flask.
Distillation
The reflux apparatus is modified to serve as a distillation apparatus. The bracket is clamped in a near
horizontal position - the immovable clamp end slightly higher than the horizontal. The round-bottom
flask is clamped in the immovable clamp and the condenser in the movable clamp.
A stillhead is fitted into the neck of the round-bottom flask. This holds the thermometer in place and
connects the round-bottom flask with the condenser. It is necessary to tilt the round-bottom flask and
stillhead slightly downward and the condenser slightly upward to engage the condenser and stillhead.
A delivery tube is fitted onto the discharge end of the condenser. This directs the condensate into a
receiver. As before, the cooling water enters at the lower side-arm and exits at the upper side-arm as
shown in the Figure.
The thermometer is held in position by an adaptor that clamps onto the stem of the thermometer. It has
two teflon gaskets which make a leak-proof seal around the stem of the thermometer. Great care must
be taken that these are correctly positioned when the thermometer is inserted into the adaptor. The bulb
of the thermometer is positioned at the point where the vapour leaves the flask. Some of the vapour
will be seen to condense on the bulb of the thermometer. Thus the reading shown by the thermometer
will be the temperature of condensing vapours as is required by the definition of boiling point.
Because of the small quantity of liquid being distilled and the small difference in boiling points (10
°C), the heating of the distillation flask requires a special technique. Furthermore, a constant watch
must be kept on the thermometer so that a switch in receptacle for the condensate can be made at the
right moment.
Heating of the flask is done without the usual wire gauze. The naked flame is passed over the bottom
of the flask so as to induce gentle boiling. The temperature will rise quite rapidly to about 138 °C
where the 1-pentanol will boil off. Then, after a dip, the temperature will rise again. At 140 °C the
beaker receiving the condensate must be changed to a weighed sample tube and the ester collected.
Great care must be taken not to heat the flask to dryness.
The temperature at which the ester distils must be recorded, as well as the ambient pressure.
Experiment 3
22
water out
water in
Distillation apparatus
EXPERIMENTAL PROCEDURE
1.
Set up a reflux apparatus as shown in the Figure in Experiment 2.
2.
Clean and dry with acetone the following: a 10 cm3 measuring cylinder, the pear-shaped flask,
the round-bottom flask, the condenser, the adapter and the delivery tube.
3.
Measure 7.5 cm3 of 1-pentanol into the measuring cylinder and transfer to the pear-shaped
flask. Similarly measure 5.5 cm3 of glacial acetic acid and transfer to the flask. Add 3 antibumping granules. Whilst swirling vigorously, add 5 drops of concentrated sulphuric acid.
4.
Heat under reflux for 45 minutes, shut off the flame and allow to cool.
5.
Rinse the separating funnel with water and make sure the tap does not leak when closed.
6.
Clamp the separating funnel upright, making sure the tap is closed, then pour the contents of
the flask into the funnel. Rinse the flask twice with 5 cm3 of water and add the rinses to the
funnel. Allow the phases to separate and then drain off the aqueous phase observing all the
precautions set out in the description of the use of the separating funnel.
7.
Add approximately 10 cm3 of deionised water to the funnel, stopper and shake as described in
the notes.
8.
After one wash with water, add 10 cm3 of saturated bicarbonate solution. Swirl until the
evolution of gas has stopped, then put the stopper in and shake gently. Be sure to release
pressure frequently. Use the gas evolution to judge whether all the acid has been neutralised as
explained above.
9.
Wash with water once to remove residual bicarbonate.
Experiment 3
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
23
Clamp the separating funnel in the upright position and allow sufficient time for separation of
phases. Run off the aqueous phase until the organic phase has entered the tap.
Pour the organic phase into a 50 cm3 conical flask. Do not run it out, the tap and stem are wet
with water. Add a spatula-tipful of anhydrous sodium sulfate to the conical flask, cover with
foil and set aside for dehydration to occur. Occasional swirling is necessary.
Set up the distillation apparatus shown in the Figure and described in the notes section.
Weigh a clean 50 cm3 beaker and ask a demonstrator to verify the mass.
Filter the ester through a fluted filter paper into a clean dry round-bottom distillation flask. Add
anti-bumping granules and install the flask in the distillation apparatus.
Heat the flask carefully with the naked flame as described above. Use a beaker to receive the
initial condensate.
When the temperature has reached 140 °C replace the beaker with the weighed beaker and
collect the ester. Do not take the flask to dryness.
Note the temperature at which the ester distils and enter this value on your report sheet together
with the ambient pressure.
Reweigh the beaker and ask a demonstrator to verify the mass.
Show your sample to the demonstrator for marking.
Experiment 3
24
Experiment 3: Esterification of 1-pentanol with acetic acid
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
PRE-PRAC CALCULATION
Demonstrator: ________________________
Mark:
H+
CH3COOH + CH3CH2CH2CH2CH2OH
→ CH3COOCH2CH2CH2CH2CH3 + H2O
acetic acid
1-pentanol
pentyl acetate
Molecular formula
C2H4O2
C5H12O
C7H14O2
Molar mass
________ g mol-1
________ g mol-1
________ g mol-1
Density
1.05 g cm-3
0.81 g cm-3
Volume used
11.0 cm3
15.0 cm3
Mass used
11.0 x 1.05 g
15.0 x 0.81 g
11.5 g
12. g
=
expected mass of pentyl acetate (remember which is the limiting reagent)
= _________________ g
Experiment 3
25
Experiment 3: Esterification of 1-pentanol with acetic acid
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Practical Mark:
Laboratory
Mark:
RESULTS
Mass of beaker + ester
=
______________ g
Mass of beaker
=
______________ g
Mass of ester recovered
=
______________ g
% Yield
=
______________ %
Boiling point of sample
=
______________ °C
Ambient pressure
=
______________ Pa
Experiment 4
26
Experiment 4: Reactions of functional groups
AIM
To supplement the theory of the lecture course by performing standard tests which identify functional
groups of organic compounds. Also to emphasise that the rate of a reaction gives additional
information which may enable the experimenter to distinguish between isomers, to illustrate that these
tests can be carried out on small quantities of materials, and to emphasise the advantage of tests being
carried out in logically grouped series instead of uncoordinated single operations.
INTRODUCTION
Primary, secondary and tertiary alcohols
Alcohols contain hydroxyl groups which are replaceable by other groups or atoms. The ease with
which a hydroxyl group is replaced decreases in the order:
tertiary > secondary > primary > methanol
and thus the time required to carry out replacement increases in the order:
tertiary < secondary < primary < methanol
R3COH < R2CHOH < RCH2OH < CH3OH
In this experiment the hydroxyl group will be replaced by a chlorine atom using the Lucas
Reagent (zinc chloride dissolved in hydrochloric acid).
R2CH2OH
HCl
ZnCl2
RCH2Cl
+
H2O
The halide produced is insoluble in the reaction medium and will show as turbidity (cloudiness). The
turbidity will appear earliest in the test tube containing the tertiary alcohol and last in the test tube
containing the primary alcohol.
Nature and strength of halide bonds
Halogen atoms in organic compounds may be ionically bonded or covalently bonded. Those which are
ionically bonded dissociate in solution, giving halide ions which react with silver ions instantly to form
an insoluble silver chloride precipitate.
RNH3+Cl- + Ag+NO3-
→ RNH3+NO3- + AgCl(s)
Those which are covalently bonded take longer to react with silver nitrate, and may even require the
use of additional solvents and the application of heat.
Of the covalently bonded haloalkanes, only a tertiary compound will react with aqueous silver nitrate.
The formation of the silver halide is not instantaneous, it takes some ten seconds to appear.
R3CCl + AgNO3 → R3CNO3 + AgCl(s)
A secondary haloalkane, when heated with ethanolic silver nitrate, will yield a precipitate after about
ten minutes whilst a primary compound may yield only a slight turbidity after heating for half-an-hour
with ethanolic silver nitrate.
Experiment 4
27
Detection of unsaturation of the type >C=C<
Ethylenic double bonds are reactive and will rapidly reduce oxidising agents such as permanganate
ions and bromine. The permanganate ion, in slightly alkaline medium, will hydroxylate the carbon
chain at the double bond:
OH
KMnO4
C C
+ MnO2
C C
H2O
OH
.
Bromine, in dichloromethane, will add across the double bond:
C
C
Br2
Br
C C
Br
.
Proof that these reactions take place is the disappearance of the purple colour of the permanganate ion
or the brown colour of the bromine solution.
Other functional groups are also capable of reducing bromine or permanganate ions, but these
reactions are slow unless heated. Thus in performing the test for this type of unsaturation, the
disappearance of colour must be immediate for positive identification of an ethylenic double bond.
Acidic properties
Many organic compounds show acidic properties. Sodium bicarbonate is a useful reagent for showing
the presence of acidity in organic compounds. Its reaction with an acid that is stronger than carbonic
acid produces carbonic acid, which decomposes to release carbon dioxide. Thus evolution of gas is an
indication of reaction with a substance more acidic than carbonic acid.
RCOOH + NaHCO3 → RCOONa + H2CO3
H2CO3 → H2O + CO2(g)
Identification of aldehydes and ketones
A characteristic common to aldehydes and ketones is their reaction with compounds containing a
primary amino group:
>C=O + H2NNH2 → >C=NNH2
i.e. aldehyde or ketone + hydrazine → hydrazone.
Most suitable for this test is 2,4-dinitrophenylhydrazine. The hydrazone formed is crystalline and has
an intense yellow colour. Thus aldehydes or ketones can be detected at very low concentrations (even
parts per million). Furthermore, being crystalline, the hydrazones have very sharp melting points and
thus the aldehyde or ketone parent compound is easily identified.
It is often mistakenly thought that all carbonyl groups (>CO) give this reaction, however the procedure
carried out in this experiment will show that that the carbonyl group in organic acids does not form a
hydrazone.
Experiment 4
28
Reducing properties
Many organic compounds have reducing properties. The most frequently encountered reducing agent
is the aldehyde group which occurs in most of the "sugars". This mild reducing property is used in
industry for the manufacture of high precision reflectors for optical instruments, and in the pathology
laboratory for the diagnosis of diabetes in humans.
The so-called "silver mirror test" for reducing properties uses Tollens' Reagent, a solution containing
[Ag(NH3)2]+, which deposits a layer of bright silver metal on reduction - hence the name.
RCHO + 2[Ag(NH3)2]+ OH- → RCOONH4 + 3NH3 + H2O + 2Ag(s)
In the pathology laboratory Benedict's Reagent is used. It contains complexed cupric ions which have
an intense blue colour. The cupric ion, on reduction, is converted to red cuprous oxide which is
insoluble. Thus the disappearance of the blue colour and the appearance of a red precipitate indicate
the presence of an aldehyde.
RCHO + Cu(OH)2 → RCOOH + Cu2O(s)
NOTES
Heating
Most of the tests require the mixture of unknown compound and test reagent to be heated. Since
heating a test tube in a Bunsen flame may be hazardous, test tubes are heated in a hot water bath.
Marking test tubes
Several test tubes will be in the hot water bath at any one time. In order to avoid confusion, all test
tubes must be clearly marked and a written record must be kept of the contents of each tube.
Measurement of reagents
Most reagents are supplied in bottles fitted with droppers. This enables one to add the specified
number of drops. Please ensure that the correct dropper is returned to a bottle. If a volume is given in
cm3 and not in drops, it is easily converted on the scale of 20 drops is approximately 1 cm3. The test
tubes provided are approximately 1.2 cm in diameter, thus a depth of 1 cm in the tube is approximately
1 cm3, and 2 cm in the tube is approximately 2 cm3. As these tests are qualitative, there is no need to
measure exact volumes. The same result would be obtained if the volume were 4 cm3, 5 cm3 or 6 cm3.
EXPERIMENTAL PROCEDURE
The following fifteen substances are provided on which the six tests outlined in Points 1 to 6 are to be
practised.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
Acetic acid, glacial
Cinnamic acid - solid
Toluene
1-Chlorobutane
Methylamine hydrochloride - solid
2-Chlorobutane
2-Chloro-2-methylpropane
Butan-1-ol
2-Methylpropan-2-ol
Butan-2-ol
Oxalic acid - solid
Methanol
Experiment 4
13.
14.
15.
29
Benzldehyde
Propanone (acetone)
Glucose
One unknown sample is provided on which tests 1, 3, 4, 5 and 6 are to be done. The conclusions of the
tests and the number of the unknown must be entered on your report sheet.
1.
Primary, secondary and tertiary alcohols
Add about 4 cm3 of Lucas Reagent to each of three test tubes. Add about 1 cm3 of No. 8 (butan-1-ol)
to the first test tube, about 1 cm3 of No. 9 (2-methylpropan-2-ol) to the second and about 1 cm3 of No.
10 (butan-2-ol) to the third test tube.
Mix and allow to stand. Record the appearance of turbidity as first, next and last and identify the
alcohols accordingly.
2.
Nature and strength of halide bonds
Add about 2 cm3 of dilute (4 M) nitric acid and 3 drops of silver nitrate solution into each of four test
tubes.
To the first, add about 0.1 g of No. 5 (methylamine hydrochloride) and mix by shaking. Note that the
precipitate of silver chloride forms immediately.
To the second, add 10 drops of No. 7 (2-chloro-2-methyl propane). Note that the precipitate starts
forming after a few seconds and gradually increases over the next ten seconds.
To the third, add 10 drops of No. 6 (2-chlorobutane) and mix by shaking. Put in the hot water bath.
To the fourth, add 10 drops of No. 4 (1-chlorobutane) and mix by shaking. Put in the hot water bath.
Record the case where the precipitate formed immediately, the case where the precipitate formed after
a few seconds and the two cases where a precipitate did not form.
On the last two cases, do the following additional tests:
Into two test tubes introduce about 2 cm3 of ethanol and about 1 cm3 of silver nitrate solution and then
add:
to the first: 4 drops of No. 4,
to the second: 4 drops of No. 6.
Mix by shaking and put in the hot water bath.
Care must be taken when immersing the two test tubes in the hot water bath. The boiling point of
ethanol is 80 °C, and when a tube containing ethanol is immersed in a bath of boiling water, the
evaporation of the ethanol may be so rapid as to cause an eruption. When ready to do this test, a little
cold water must be added to the hot water bath to bring the temperature down below the boiling point
of water before immersing the two test tubes containing ethanol.
Record the case where a precipitate forms after about 10 minutes. Record the case where only a slight
turbidity appears. Identify the halides according to ease of release of the halogen atom as described in
the introduction.
Experiment 4
3.
Detection of unsaturation of the type >C=C<
3.1
By means of permanganate ion
Into three test tubes introduce the following:
into (i)
about 1 cm3 of water
6 drops of No. 1 (acetic acid) and then sufficient solid sodiumcarbonate
make the solution alkaline to litmus paper
into (ii) 1 cm3 of water
about 0,1 g of No. 2 (cinnamic acid) and then sufficient solid sodium
carbonate to make the solution alkaline to litmus paper.
into (iii)
6 drops of No. 3 (toluene). Do not heat.
30
to
To each add 3 drops of potassium permanganate solution and mix by shaking. Record the case in
which the permanganate colour disappears on your report sheet.
3.2
By means of bromine
Into three test tubes introduce the following:
into (i)
6 drops of No. 1 (acetic acid)
into (ii) about 0.1 g of No. 2 (cinnamic acid)
into (iii)
6 drops of No. 3 (toluene).
To each add 6 drops of bromine solution and mix.
Record the case in which the brown colour of bromine disappears, thus identifying the compound
containing an ethylenic double bond.
4.
Acidic properties
Into three test tubes introduce about 1 cm3 of saturated bicarbonate solution and then add
to (i)
4 drops of No. 1 (acetic acid)
to (ii)
about 0.1 g of No. 11 (oxalic acid)
to (iii)
4 drops of No. 12 (methanol).
Record the cases where gas effervescence occurs, thus identifying acids stronger than carbonic acid.
5.
Formation of hydrazones
Into four test tubes introduce 5 drops of 2,4-dinitrophenylhydrazine and then add:
to (i) 2 drops of No. 1 (acetic acid)
to (ii) 2 drops of No. 13 (benzaldehyde)
to (iii) 2 drops of No. 14 (propanone).
Record the cases where a yellow precipitate forms, identifying the presence of either an aldehyde or a
ketone carbonyl group. Note particularly that the carbonyl group in a carboxyl group does not form a
hydrazone.
6.
The reducing properties of aldehydes
6.1
Benedict's reagent
Into four test tubes introduce about 5 cm3 of Benedict's solution and then add:
to (i) 15 drops of No. 1 (acetic acid)
to (ii) 15 drops of No. 13 (benzaldehyde)
to (iii) 15 drops of No. 14 (propanone)
to (iv) about 0.1 g of No. 15 (glucose).
Experiment 4
31
Mix and heat in the waterbath. Record observations as:
(a)
total disappearance of blue colour and copious brown precipitate
(b)
partial disappearance of blue colour and slight brown precipitate
(c)
no visible change
In one case the reaction is quick and quantitative, hence the total disappearance of the blue colour. In
one case the reaction is much slower and after several minutes, only a reduction in the intensity of the
blue colour can be seen. Nevertheless, loss in blue-colour intensity and/or a slight brown precipitate is
positive proof of the presence of an aldehyde.
6.2
Tollens' reagent
Into four test tubes introduce about 2 cm3 of silver nitrate and 2 drops of sodium hydroxide solution. A
precipitate will form. Add aqueous ammonia dropwise and with shaking until the precipitate is
redissolved. Then add:
to (i) 5 drops of No. 1 (acetic acid)
to (ii) 5 drops of No. 14 (propanone)
to (iii) about 0,1 g of No. 15 (glucose).
Mix by shaking. Dip each tube in the hot water bath for a few seconds.
Record the cases where a silver mirror forms in the tube. [To remove the silver deposit from the tube,
pour in nitric acid and heat in the water bath as soon as the test is completed.]
7.
Analysis of unknown compound
Record the number of the unknown compound on the appropriate page of your report, and then do the
analyses prescribed.
Experiment 4
32
Experiment 4: Reactions of functional groups
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
The following example indicates the style of report expected from your observations.
1.
Test for unsaturation
Reagents:
1.
2.
Sample No. 100
KMnO4/OHBr2 (CH2Cl2)
ETHENE
STRUCTURE
H
H
C
H
C
H
Observations:
1.
Purple colour of KMnO4 discharged. Brown precipitate formed.
2.
Brown colour of Br2 (CH2Cl2) discharged. Turned colourless.
Conclusion:
This compound has an ethylenic bond across which groups can add.
This is a test for unsaturation.
__________________________________________________________________________________
Experiment 4
1.
33
Reactivity of primary, secondary and tertiary alcohols
Reagents:
Sample No. 8
1-BUTANOL
STRUCTURE
Observations:
__________________________________________________________________________________
2-METHYL-2-PROPANOL
STRUCTURE
Sample No. 9
Observations:
__________________________________________________________________________________
2-BUTANOL
STRUCTURE
Sample No. 10
Observations:
__________________________________________________________________________________
Conclusion:
(Indicate whether the alcohol is 1°, 2°, or 3° in your answer.)
The rate of reaction of alcohols with the Lucas reagent is in the order:
Experiment 4
2.
34
Strengths of halide bonds
Reagents:
Sample No. 5
METHYLAMINE HYDROCHLORIDE
STRUCTURE
Observations:
__________________________________________________________________________________
2-CHLORO-2-METHYLPROPANE
STRUCTURE
Sample No. 7
Observations:
__________________________________________________________________________________
2-CHLOROBUTANE
STRUCTURE
Sample No. 6
Observation:
__________________________________________________________________________________
1-CHLOROBUTANE
STRUCTURE
Sample No. 4
Observations:
__________________________________________________________________________________
Conclusion
(Indicate whether halide bond is 1°, 2°, 3°, or ionic in your answer.)
The strength of the halide bond INCREASED in the order:
Experiment 4
3.
35
Test for unsaturation
Reagents:
1.
2.
Sample No. 1
ACETIC ACID
STRUCTURE
Observations:
1.
2.
__________________________________________________________________________________
CINNAMIC ACID
STRUCTURE
Sample No. 2
Observations:
1.
2.
__________________________________________________________________________________
TOLUENE
STRUCTURE
Sample No. 3
Observations:
1.
2.
__________________________________________________________________________________
Conclusion: Of the compounds tested above, typical results for unsaturation were
shown by:
Experiment 4
4.
36
Test for acidity
Reagents:
Sample No. 1
ACETIC ACID
STRUCTURE
Observations:
__________________________________________________________________________________
OXALIC ACID
STRUCTURE
Sample No. 11
Observations:
__________________________________________________________________________________
METHANOL
STRUCTURE
Sample No. 12
Observations:
__________________________________________________________________________________
Conclusion:
The substances with acidic properties (stronger than that of carbonic acid) in order of
DECREASING acidic strength are:
Experiment 4
5.
37
Hydrazone formation
Reagents:
Sample No. 1
ACETIC ACID
STRUCTURE
Observations:
__________________________________________________________________________________
BENZALDEHYDE
STRUCTURE
Sample No. 13
Observations:
__________________________________________________________________________________
PROPANONE
STRUCTURE
Sample No. 14
Observations:
__________________________________________________________________________________
Conclusion:
Of the compounds tested above, those showing typical nucleophilic addition reactions
characteristic of the carbonyl group were:
Experiment 4
6.
38
Reducing properties
Reagents:
1.
2.
Sample No. 1
ACETIC ACID
STRUCTURE
Observations:
1.
2.
__________________________________________________________________________________
BENZALDEHYDE
STRUCTURE
Sample No. 13
Observations:
(Benedict's Solution)
__________________________________________________________________________________
PROPANONE
STRUCTURE
Sample No. 14
Observations:
1.
2.
__________________________________________________________________________________
Experiment 4
Sample No. 15
GLUCOSE
39
STRUCTURE
Observations:
1.
2.
__________________________________________________________________________________
Conclusion:
Of the compounds tested above, those showing good reducing properties were:
Experiment 4
7.
Identification of "unknown" functional group(s)
Sample No.
______________
TEST
OBSERVATION
______________________________________________________________________
(a)
Test for unsaturation
(b)
Test for acidity
(c)
Lucas test
(d)
Test with (2,4-dinitrophenylhydrazine)
(e)
Test with Benedict's reagent
____________________________________________________________________________
Conclusion
The following functional group(s) is/are present:
40
Experiments 5‐6 Introduction 41
Reactions of the cations of the metallic elements: Qualitative
analysis
INTRODUCTION
The aim of the next two experiments (6-7) is to learn about the relationship between the position
of the metal in the periodic table and its chemical properties. To achieve this we will study
methods of separating metallic elements from each other based on differences in solubility of
their various compounds.
For simplicity we will limit ourselves to ten metals, which have been grouped in fours in the
order that they will be studied:
1
2
Calcium, Aluminium, Chromium, Tin, Lead
Iron, Cobalt, Nickel, Copper, Zinc
These metals are either related in vertical groups or in horizontal periods, and they are found
together either in nature or in alloys commonly encountered in industry.
GENERAL PROCEDURES
In each Experiment, set up four test tubes and add 10 drops of each cation solution to a test tube.
Then add 3 drops of the reagent to each. Observe what occurs. Repeat this for each reagent.
The precipitates are sometimes slow to form; always scratch the inside wall of the test tube with
a thin glass rod both to mix the solutions and to encourage crystallization of the insoluble
compound.
If in doubt: ask your demonstrator for help. Also, read your textbook, which contains details of
the properties of the compounds that you will prepare in these experiments.
CENTRIFUGE
You will need to separate the precipitates in some cases. This is done by using the centrifuge on
the window bench. Ask your demonstrator for help.
Experiment 5
Experiment 5:
42
Reactions of the cations of calcium,
aluminium, chromium, tin and lead
AIM
To observe the change in properties of the compounds of the elements as their atomic number
increases, and to compare the elements on the far left and right with the transition metals in the
centre.
Ia
1
H
VIII a
2
Periodic Table of the Elements
II a
III a
IV a
Va
VI a
VII a
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
III b
21
IV b
22
Vb
23
VI b
24
VII b
25
26
VIII b
27
28
Ib
29
II b
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
87
88
89
Fr
Ra
**Ac
*Lanthanides
**Actinides
Experiment 5
43
INTRODUCTION
Commonly found compounds of these metals are:
Calcium (Ca2+): Calcite, limestone, marble, CaCO3; Plaster of Paris, 2(CaSO4).H2O;
Whewellite in kidney and bladder stones, Ca(oxalate).H2O.
Tin (Sn4+): Cassiterite, SnO2.
Lead (Pb2+): Litharge, PbO; Galena, PbS; Red lead, Pb3O4.
Aluminium (Al3+):
KAl(SO4).12H2O.
Bauxite
Al2O3.2H2O;
Corundum
Al2O3;
Alum
(Potassium)
Chromium: Chrome green Cr2O3; Chromite FeCr2O4; Zinc chromate ZnCrO4.
EXPERIMENTAL PROCEDURE
First note the colour of each metal solution. Then test solutions of the metals systematically to
determine which compounds are insoluble or coloured and therefore could be used to identify
the cationic species. Test each metal by adding the following solutions to a solution of the metal:
2
3
4
5
6
7
NaOH
Excess NaOH
NH3(aq)
KI, dilute, then heat
K2SO4
Na2CO3
8 (NH4)2CO3
9 NH3, NH4Cl, Na2HPO4
10 Acid H2S
11 Na oxalate
12 K2CrO4
13 NaOH + H2O2
Record your observations on the report sheet. Give the formulae of the insoluble and/or coloured
species in the appropriate spaces.
Note that redox reactions can occur, e.g. Cr3+ + OH¯ + H2O2 → CrO42¯, and the solution
becomes bright yellow.
To get H2S, add 10 drops of thioacetamide solution and warm gently in a water bath.
You will be provided with two unknown solutions. Identify the cations present in these
solutions. Explain your reasoning in the appropriate space on your report sheet.
Experiment 5
44
COMMENTS ON THE REACTIONS
The oxides and hydroxides are generally insoluble.
2
3
Aluminium, chromium and lead are amphoteric, dissolving in excess OHˉ to give
[M(OH)4]xˉ.
4
The Ksp’s of the group II metal compounds are not exceeded because [OHˉ] is too
low in aqueous NH3.
5
PbCl2 and PbI2 are insoluble in cold water, but dissolve in hot water. Note how
yellow PbI2 dissolves to give a colourless solution. Allow it to cool; PbI2 crystallises
as “golden spangles”.
6
The anhydrous sulfates are insoluble: PbSO4.
7, 8
All form insoluble carbonates in basic medium.
10
Only lead sulphide is insoluble in water (and acid). Note the black colour.
11
Insoluble calcium oxalate allows Ca2+ to be separated conveniently from Mg2+.
12
Note the parallel between the insoluble sulfates and chromates, due to SO42ˉ and
CrO42ˉ having the same shape, size and charge.
13
The oxidation of Cr3+ to bright yellow CrO42ˉ is quite characteristic.
Experiment 5
Experiment 5:
45
Reactions of the cations of calcium,
aluminium, chromium, tin and lead
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
TESTS FOR Ca2+, Al3+, Cr3+, Sn4+, Pb2+
Ca2+
1 Colour of solution
2 NaOH
3 Excess NaOH
4 NH3(aq)
5 KI, dilute, then heat
6 K2SO4
7 Na2CO3
8 (NH4)2CO3
9 NH3, NH4Cl, Na2HPO4
10 Acid H2S
11 Na oxalate
12 K2CrO4
13 NaOH + H2O2
Al3+
Cr3+
Sn4+
Pb2+
Experiment 5
46
ANALYSES OF UNKNOWNS
Unknown 1: Identification number _______
This solution contains ONE of the four cations. Identify it, giving details of your reasoning.
Cation _________________________
Unknown 2: Identification number _______
This solution contains TWO of the four cations. Identify both, giving details of your reasoning.
Cations ____________
Experiment 6
Experiment 6:
47
Reactions of the cations of iron, cobalt, nickel,
copper and zinc
AIM
To observe the properties of these transition metals, some of which are critical to the functioning of
the body whilst some are important as alloys in engineering.
Ia
1
H
VIII a
2
Periodic Table of the Elements
II a
III a
IV a
Va
VI a
VII a
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
III b
21
IV b
22
Vb
23
VI b
24
VII b
25
26
VIII b
27
28
Ib
29
II b
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
Fr
Ra
**Ac
*Lanthanides
**Actinides
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Experiment 6
48
INTRODUCTION
The commonly found compounds of these elements are:
Iron (Fe3+): Haematite Fe2O3, Magnetite Fe3O4, Pyrites FeS2.
Cobalt (Co2+): in Vitamin B12, CoSO4.7H2O as food supplement for cattle.
Nickel (Ni2+): no common compounds; the metal is used for alloying, e.g. in coins.
Copper (Cu2+): blue vitriol, CuSO4.5H2O.
Zinc (Zn2+): ZnO in ointments and sunguard, Sphalerite, Wurzite ZnS, in proteins, and with copper
in brass.
EXPERIMENTAL PROCEDURE
First note the colour of each metal solution. Then test solutions of the metals systematically to
determine which compounds are insoluble or coloured and therefore could be used to identify the
cationic species. Carry out the tests by adding the following solutions to a solution of the metal:
7 KI/CH2Cl2
2 NaOH
8 KSCN
3 NH3(aq)
4 conc. HCl, water, acetone
9 K4[Fe(CN)6]
10 DMG, NH3(aq)
5 H2S, acid
6 H2S, NH3(aq)
Give the formulae of all insoluble and coloured compounds.
Note that redox reactions occur for these transition metals.
To get H2S, add 10 drops of thioacetamide solution, and warm gently in a water bath.
You will provided with two unknown solutions. Identify the cations present in these solutions.
Explain your reasoning in the appropriate space on your report sheet.
COMMENTS ON THE REACTIONS
2
All hydroxides are insoluble, but zinc is amphoteric and forms [Zn(OH)4]2-. Note the
brilliant blue of Co(OH)2.
3
Initially the hydroxides precipitate but these dissolve in excess ammonia with the formation
of ammine complexes: [Cu(NH3)4]2+ and [Ni(NH3)6]2+.
5
CuS is formed in acid medium. CoS, NiS and ZnS form only in base medium.
8
The deep red compound [Fe(SCN)]2+ is characteristic of iron.
9
The deep green [Co{Fe(CN)6}] complex is characteristic.
(Turnbull’s) blue is a classic compound, characteristic of iron.
10
The deep red insoluble [Ni(DMG)2] complex is unique. It is destroyed by dilute acid.
The deep blue Prussian
Experiment 6
Experiment 6:
49
Reactions of the cations of iron, cobalt, nickel,
copper and zinc
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
TESTS FOR Fe3+, Co2+, Ni2+, Cu2+, Zn2+
Fe3+
1 Colour of solution
2 NaOH
3 NH3(aq)
4 (a) conc HCl
(b) add a little water
(c) add acetone
5 H2S, acid
6 H2S, NH3(aq)
7 KI/CH2Cl2 (3 cm3)
8 KSCN
9 K4[Fe(CN)6]
10 DMG, NH3(aq)
Co2+
Ni2+
Cu2+
Ag+
Experiment 6
ANALYSES OF UNKNOWNS
Unknown 1: Identification number _______
This solution contains ONE of the four cations. Identify it, giving details of your reasoning.
Cation _________________________
Unknown 2: Identification number _______
This solution contains TWO of the four cations. Identify both, giving details of your reasoning.
Cations ____________
50
Experiment 12 51
Summary of experiments 5-6
The elements can be grouped as:
Those which are basic, do not form stable coordination ion complexes with ammonia, and
1
are “oxygen-lovers”. Note the presence of lead.
Ia
1
H
VIII a
2
II a
III a
IV a
Va
VI a
VII a
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
III b
21
IV b
22
Vb
23
VI b
24
VII b
25
26
VIII b
27
28
Ib
29
II b
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
Fr
Ra
**Ac
2
Those which are “oxygen-lovers”, but tend to be amphoteric and will form stable
coordination complexes.
Ia
1
H
VIII a
2
II a
III a
IV a
Va
VI a
VII a
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
III b
21
IV b
22
Vb
23
VI b
24
VII b
25
26
VIII b
27
28
Ib
29
II b
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
Ra
**Ac
Fr
*Lanthanides
**Actinides
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
In basicity, the lanthanide 3+ cations fall midway between Ca2+ and Al3+. The lanthanides
are “oxygen-lovers”.
Experiment 12 3
52
Those transition elements to the right of Mn/Fe, which preferentially form coordination
complexes with ammonia (and other N-donor molecules). Note how copper, silver and
nickel are used as coinage metals because they resist corrosion.
Ia
1
H
VIII a
2
II a
III a
IV a
Va
VI a
VII a
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
III b
21
IV b
22
Vb
23
VI b
24
VII b
25
26
VIII b
27
28
Ib
29
II b
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
Fr
Ra
**Ac
4
Those “base” metals which are well known because of the vast range of alloys that they
form. They occur in nature as sulfides. The lower ones, e.g. mercury and lead, exhibit the
“Inert Pair” effect, whereby the favoured oxidation state is TWO LESS than the group
number; e.g. Pb2+ in Group IV.
Ia
1
H
VIII a
2
II a
III a
IV a
Va
VI a
VII a
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
19
20
III b
21
26
VIII b
27
28
Ib
29
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
87
88
89
Fr
Ra
**Ac
IV b
22
Vb
23
VI b
24
VII b
25
II b
30
Experiment 12 53
The Diagonal Relationship is found in the top left-hand corner, where the solution properties of
lithium and scandium are remarkably like those of magnesium. Similarly, the properties of
Be2+ and Ti4+ are remarkably like those of Al3+. This is due to the ratio: (charge on cation) ÷
(radius of cation)2 being similar.
3
4
5
6
7
8
9
Li
Be
B
C
N
O
F
11
12
13
14
15
16
17
Na
Mg
Al
Si
P
S
Cl
19
20
21
22
23
24
25
K
Ca
Sc
Ti
V
Cr
Mn
37
38
39
40
41
42
43
Rb
Sr
Y
Zr
Nb
Mo
Tc
55
56
57
72
73
74
75
Cs
Ba
*La
Hf
Ta
W
Re
Titanium, zirconium and hafnium occur commonly in Nature and are “oxygen-lovers”, e.g.
Ilmenite FeTiO3; Zircon ZrSiO4; Baddeleyite ZrO2. However, the chemistry of zirconium is
almost identical with that of hafnium because the radii of the 4+ cations are identical. This is
the result of the “Lanthanide Contraction”.
Experiment 7a
54
Experiment 7a: The pH meter and potentiometric titrations
AIM
To introduce the use of the pH meter, and the technique of potentiometric titrations.
INTRODUCTION
The theoretical aspects of the pH concept are covered adequately in your textbook and were covered in
the first semester. The following notes are thus restricted to background information on the pH meter
and potentiometric titrations.
The pH meter
Cell potentials can be used to determine the concentration of a specific ion in a solution. This is
possible due to the availability of two types of electrodes:
 reference electrodes, which maintain a constant half-cell potential independent of solution
composition
 indicator electrodes, which respond to the changes in the concentrations of specific ions.
One of the most familiar applications of this relationship is found in the use of pH meters to determine
the [H3O+] of a solution.
A species to which an electrode responds is called an electro-active species. The potential of an
electrode changes when the concentration of the electro-active species for that electrode is varied.
When H3O+ is the electro-active species both the electrode potential and the pH change when the H3O+
concentration is varied. This implies that the pH of a solution can be related to the potential of a cell in
which H3O+ is the electro-active species.
It can be shown that a change of one pH unit (ten-fold [H3O+] change) causes an overall change in the
cell potential of 0.059 V.
The pH meter is an electronic instrument which measures the potential difference between a reference
half-cell and a half-cell whose potential changes when [H3O+] varies. The dial of the meter is often
graduated in both pH and millivolt units. A "glass" electrode is sensitive to changes in [H3O+] and is
therefore used to sense pH. A silver-silver chloride electrode is commonly used as a reference
electrode. The half-cell reaction for this electrode is
E = 0.222 V
AgCl(s) + eˉ → Ag+ + Clˉ
For convenience and ease of use these two electrodes are often mounted on the same electrode
body to form a combination electrode (see figure on next page). The thin glass bulb on these
electrodes is very fragile. Use great caution not to touch the bottom of the beaker with it.
Experiment 7a
55
(+)
(-)
Terminals to pH meter
Saturated aqueous
AgCl and KCl
Ag/AgCl wire
in outer electrode
.
Porous plug
Ag/AgCl wire
in inner electrode
Glass membrane
0.1 M HCl
saturated with
AgCl
Combined pH glass electrode
Use of the pH meter
A demonstrator will instruct you in the use of the pH meter. Be extremely careful when using the
electrodes – they are sensitive (and expensive!). DO NOT TOUCH the pH meter until you have
been shown how to use it.
Measurement of pH
1.
2.
3.
4.
Pour a sample of the solution whose pH is to be measured into a clean and dry 50 cm3 beaker.
Immerse the rinsed and dry electrode assembly in the solution to be measured.
Read the measured value on the instrument.
Rinse and dry the electrode assembly.
Potentiometric titrations
Since the potential of an electrode dipping into a solution of an electrolyte depends on the
concentration of the ions to which the electrode responds, it is possible to use the potential as an
"indicator" in volumetric analysis. The electrode potential depends on the logarithm of the
concentration of ions, and is, therefore, not suitable for obtaining concentration directly with any
accuracy, but the change of potential with concentration during a titration provides an accurate
indication of the equivalence point. Thus, the cell
AgAgCl(s),[Clˉ]acid solutionglass membrane[H+][Clˉ]AgCl(s)Ag
will have a certain e.m.f. depending on the pH of the acid solution. On adding small portions of a
standard solution of alkali to the acid, the e.m.f. of the cell will alter slowly at first, because the change
Experiment 7a
56
in the electrode potential depends on the fraction of hydrogen ion removed. As the amount of alkali
added approaches equivalence to the amount of hydrogen ion in the solution, the fraction of the
hydrogen ion concentration removed by each drop of alkali solution rapidly increases, and there is a
correspondingly rapid change in the e.m.f. Later, as excess of alkali is added, the e.m.f. again shows a
slow change. Consequently, when the e.m.f. of the cell, E, is plotted against the volume, V, of standard
alkali added, a curve of the form shown on the next page is obtained. The end-point of the titration is
midway at the point of inflexion of the curve, EP.
In carrying out the titrations, the titrating liquid or titrant is added, in small quantities at a time, from a
burette into the solution to be titrated (the titrate). The solution is kept well mixed by means of a
stirrer. As the titration approaches the equivalence point, the titrant is added in smaller and smaller
volumes so that the graph in the neighbourhood of the equivalence point is obtained with precision.
Advantages
1.
Potentiometric titrations are applicable to any reactions for which an appropriate electrode is
available; for example, sulphides can be titrated with lead salts, using a lead electrode. Many
titrations for which no colour indicator is available can be carried out by using the
potentiometric method.
2.
The determination is very reliable, since the result depends on a number of independent
readings, not on one subjective judgement of an "end-point" (which may be rather ill-defined
in indicator titrations).
Potentiometric titration curves
The titration curve is the graph obtained by plotting pairs of values of reagent volume (abscissae, x)
and electrode potential or pH value (ordinates, y), as shown in the figure. The curve can be drawn
manually from pairs of values obtained during the titration.
The titration curves give information about the run of the whole titration. Apart from the titration end
point, other data may be obtained from them, for example pKa.
pH
½
½
pHE
Ep
pH = pKa
VE/2
VE
3
V/cm
Experiment 7a
57
The end point of the titration corresponds to a point of inflexion (EP) in the curve, when it is
symmetrical. It can be determined in the following way:
Two parallel tangents are drawn on the curve as shown by the example in the figure. A centre line is
drawn parallel to the two tangents and its point of intersection with the titration curve gives the end
point required. The perpendicular line going through it shows on the x-axis the volume of reagent used
to reach the end point, VE.
It is important to understand that only in some instances will the neutralisation point of an acid-base
reaction correspond to a pH value of exactly 7. The main factor affecting the end-point pH value is the
strength of the reactants concerned, as shown from the following curves:
pH
pH
7
7
25
25
Alkali/cm3
(b) Strong alkali and weak acid
Alkali/cm3
(a) Strong alkali and strong acid
pH
pH
7
7
25
Alkali/cm3
(c) Weak alkali and strong acid
25
Alkali/cm3
(d) Weak alkali and weak acid
The above titration curves show the pH changes as a 0.1 M solution of a strong or weak alkali is added
to 25 cm3 of a 0.1 M solution of a strong or weak acid. Note that because the reagent concentrations
are all the same, and the acids and bases are monovalent, the equivalence point occurs at the same
volume. The precise shapes of the curves depend on the actual strengths of the alkalis and acids
concerned. Slightly different curves are also obtained if solutions with concentrations other than 0.1 M
are used.
Experiment 7a
58
EXPERIMENTAL PROCEDURE
The instructor in charge will demonstrate the use of the pH meter. Refer also to the instructions
on the use of the pH meter given earlier.
1(a)
pH of "unknown" solutions
Measure the pH of the two "unknown" solutions provided and enter the results in Table 1. Also
measure the pH of tap water and enter this in Table 1. Record all pH values to the nearest 0.05
of a pH unit. Measure the pH of de-ionised water. Note how the value drifts. Why does this
occur?
(b)
pH of solutions of typical acids and bases
Use the pH meter to measure the pH of the following:
0.1 M H3PO4
0.1 M NH3
0.1 M H3BO3
0.02 M Ca(OH)2
Arrange them into a list according to decreasing H3O+ concentration (Table 2).
2
Potentiometric titration of 0.1 M HCl against approximately 0.1 M NaOH
Pipette 10.00 cm3 of the 0.1 M acid into the titration vessel (250 cm3 beaker) and add
approximately 100 cm3 of distilled water. Place the titration vessel in position and introduce
the electrode assembly. Stir the solution by means of the magnetic stirrer and record the pH
value. Add the alkali from the burette 1 cm3 at a time, recording pH values and volumes after
each addition. As the equivalence point of the titration is approached (approximately 9.0 cm3),
a sharp increase in the pH will be observed. At this stage add the alkali in 0.2 cm3 portions and
continue until past the end point (approximately 12.0 cm3). Then, once again, add the alkali in
1.0 cm3 portions until a total volume of about 15.0 cm3 has been added.
Plot a curve of pH (ordinate, y) against volume of NaOH added (abscissa, x) and determine the
equivalence point of the titration from the point of inflexion of the curve. From your
equivalence point, determine the concentration of the NaOH solution.
List two indicators that you could use for this titration if you wished to determine the
equivalence point using a colour indicator. (Refer to Table on next page.)
Experiment 7a
59
Indicator
Colour change
(acid to base)
pKIn
pH range
methyl violet
yellow, blue, violet
-
0.2 - 3.0
thymol blue
red to yellow
1.7
1.2 - 2.8
orange IV
red to yellow
-
1.3 - 3.0
methyl orange
red, orange, yellow
3.7
3.1 - 4.4
bromophenol blue
yellow to blue violet
4.0
3.0 - 4.6
congo red
blue to red
-
3.0 - 5.0
bromocresol green
yellow to blue
4.7
3.8 - 5.4
methyl red
red to yellow
5.1
4.6 - 6.2
chlorophenol red
yellow to red
6.0
4.8 - 6.8
bromothymol blue
yellow to blue
7.0
6.0 - 7.6
phenol red
yellow to red
7.9
6.8 - 8.2
thymol blue
yellow to blue
8.9
8.0 - 9.6
phenolphthalein
colourless to red
9.6
8.3 - 10.0
alizarin yellow R
yellow to red
-
10.0 - 12.0
indigo carmine
blue to yellow
-
11.4 - 13.0
trinitrobenzene
colourless to orange
-
12.0 - 14.0
Experiment 7a
60
Experiment 7a: The pH meter and potentiometric titrations
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC PROBLEM
A potentiometric titration of 10.00 cm3 0.1 M HCl against NaOH was carried out. An endpoint volume
of 9.23 cm3 NaOH was obtained. Calculate the molarity of the NaOH solution.
Experiment 7a
61
Experiment 7a: The pH meter and potentiometric titrations
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
1
The determination of the pH of solutions
(a)
Table 1: pH of "unknowns"
Sample
pH Value
Sample
Tapwater
Unknown 1
De-ionised water
Unknown 2
Code
pH Value
Reason for the drift in pH with de-ionised water: ____________________________
________________________________________________________________________
(b)
Table 2: pH of solutions of typical acids and bases
Sample
pH
0.1 M H3BO3
0.1 M H3PO4
0.1 M NH3
0.02 M Ca(OH)2
Order of decreasing H3O+ concentration:
[H3O+]
[OHˉ]
Experiment 7a
2
62
Potentiometric titration of 0.1 M HCl against approximately 0.1 M NaOH:
Vol/cm3 NaOH
pH
Vol/cm3 NaOH
0.00
10.00
1.00
10.20
2.00
10.40
3.00
10.60
4.00
10.80
5.00
11.00
6.00
11.20
7.00
11.40
8.00
11.60
9.00
11.80
9.20
12.00
9.40
13.00
9.60
14.00
9.80
15.00
pH
Volume at end point of titration (from your graph)
__________ cm3 NaOH
Molarity of HCl (given)
__________
Hence molarity of NaOH
__________
(Show your calculations.)
Suitable indicators you could use as an alternative to the pH meter
pH range
indicator
Experiment 7a
pH Volume of titrant added/cm3
63
Experiment 7b
Experiment 7b:
64
Equilibria of water, weak acids and bases, and
buffer solutions
AIM
To study the principles relating to the equilibria which exist between hydronium and hydroxide
ions in pure water, and in aqueous solutions of acids and bases. To investigate the behaviour of
buffer solutions.
INTRODUCTION
The ionisation of water
Very pure water shows a very small but measurable electrical conductivity, thus showing slight
dissociation into ions as follows:
H2O + H2O  H3O+ + OHˉ
hydronium hydroxide
ion
ion
As can be seen above, [H3O+] = [OHˉ] and this is found to be equal to 10-7 mol dm-3.
The equilibrium constant, in terms of concentration, for the above equilibrium is:
K
[H 3O  ][OH  ]
[H 2 O]2
Since for all practical purposes [H2O] is a constant, the expression becomes:
Kw = [H3O+][OHˉ] = 10-14 mol2 dm-6 at 25 °C.
Kw is known as the ionic product of water. The ionic product for water applies to all aqueous
solutions, i.e. the relative concentrations of H3O+ and OHˉ are always interdependent: in other
words as the concentration of either ion is varied, so the other will change accordingly, so that the
product remains at 10-14 mol2 dm-6 at 25 °C.
Weak acid or weak base equilibria
A comparison of experimentally measured [H3O+] or [OHˉ] of a weak acid or a weak base with its
total concentration enables one to calculate the degree of ionisation. This is one measure of the relative
weakness of a given acid or base. For example, if a saturated carbonic acid (H2CO3) solution (about
0.04 M) has a pH of 4 ([H3O+] = 1 x 10-4 mol dm-3) we may calculate, from its primary dissociation
step, the extent to which it has been converted into H3O+ and HCO3ˉ, i.e. its degree of ionisation, as
follows:
H2CO3 + H2O  H3O+ + HCO3ˉ
Actual H 3O  concentration 1x10 4

 0.25x10 2  0.25%
Total concentration
4x10 2
This means that 5 out of every 2 000 H2CO3 molecules are dissociated in aqueous solution.
Experiment 7b
65
Buffer solutions
A buffer solution is one having both a reserve of acidity and a reserve of alkalinity, i.e. a solution
which shows only a small change in pH on addition of small quantities of either acid or base. One
type of buffer solution is a mixture containing a weak base and its salt, e.g. NH3 and NH4Cl.
If a strong acid is added to a mixed solution of NH3 and NH4+ ions we have:
NH3 + H3O+ → NH4+ + H2O
(i.e. the base in the buffer solution reacts with the added acid).
Similarly if a strong base is added we have:
NH4+ + OHˉ → NH3 + H2O
(i.e. the acid in the buffer solution reacts with the added base) and the mixture does not greatly
change its pH in either case.
This is further explained by the equilibrium expressions for the ionisation of ammonia,
NH3 + H2O  NH4+ + OHˉ
[NH 4 ][OH  ]
 K b  1.8x10  3
[NH 3 ]
[OH  ]  K b
[NH 3 ]
[NH 4 ]
Since [NH3] represents the concentration of free base in equilibrium with its ammonium salt
([NH4+]), we can rewrite the above expression as:
[OH  ]  K b
[base]
[salt]
[base]
[salt]
[salt]
pOH  pK b  log
[base]
 log[OH  ]  logK b  log
We see that the [OHˉ] depends upon the ratio of concentrations of the ammonium salt, NH4+, to free
ammonia. As long as there is a considerable amount of both salt and base present, the ratio will not
change greatly upon the addition of small amounts of acid or base. The [OHˉ] will be
correspondingly more or less constant.
Other types of buffer solutions are those containing weak acids and their salts. An example of such
a buffer is a mixture of acetic acid and its salt, sodium acetate. The dissociation constant for acetic
acid is 1.82 x 10-5, and once again it can be shown that:
CH3COOH + H2O  H3O+ + CH3COOˉ
Experiment 7b
Ka 
66
[H 3O  ][CH 3COO  ]
 1.82x10 5
[CH 3COOH]
[H 3O  ]  K a
[CH 3COOH]
[salt]
 Ka

CH 3COO ]
[acid]
[salt]
[acid]
A solution containing equal concentrations of acetic acid and sodium acetate (i.e. a half neutralised
solution of the acid) has the maximum buffer capacity. In such a case pH = pKa. Thus a half
neutralised solution of 0.1 M acetic acid will have [H3O+] = 1.82 x 10-5 mol
dm-3 and hence pH = 4.74.
pH  pK a  log
Addition of a small concentration of H3O+ ions to such a solution will result in the H3O+ ions
combining with CH3COO- ions to form undissociated CH3COOH as follows:
H3O+ + CH3COOˉ
 CH3COOH + H2O
Likewise, if a small concentration of OH- ions is added, these will combine with H3O+ ions arising
from the dissociation of acetic acid, and result in unionised water. The equilibrium will be
disturbed, and more acetic acid will dissociate to replace the H3O+ removed by the base. In either
case, the concentration of acid or salt will not be appreciably changed, and so the pH of the solution
will essentially not be affected.
EXPERIMENTAL PROCEDURE
Since acetic acid and ammonia have about the same value for their respective ionisation constants,
a solution of ammonium acetate will be practically neutral. Place 5 cm3 of 1 M CH3COONH4 in a
15 cm3 test tube; place 5 cm3 of deionised water in a second 15 cm3 test tube. Add 2 drops of
methyl orange indicator to each.
Fill your 10 cm3 graduated cylinder to the mark with 1 M HCl. Now add a drop of the HCl to the 5
cm3 sample of water. Add more HCl if needed, until the methyl orange turns red, i.e. until the water
has been changed to about 10-3 M in H3O+. What volume of the 1 M HCl was needed to do this?
Now determine what volume of 1 M HCl is required to produce the same colour change, using the
5 cm3 sample of 1 M CH3COONH4 instead of water.
Similarly, test the buffering action of ammonium acetate against the addition of a base. Prepare two
more 5 cm3 samples of H2O and 1 M CH3COONH4, respectively. To each add 2 drops of alizarin
yellow R indicator.
Fill your 10 cm3 graduated cylinder with the 1 M NaOH.
Add a drop of this, and more if needed, to the water sample until the indicator colour changes
(about 10-2 M OHˉ). Also add the base from the 10 cm3 graduated cylinder, a little at a time, to the
ammonium acetate sample, until the hydroxide ion concentration has been increased to 10-2 M OH-.
Note the respective volumes of 1 M NaOH needed. Explain why the ammonium acetate solution is
able to neutralise both acids and bases.
Experiment 7b
Experiment 7b:
67
Equilibria of water, weak acids and bases, and
buffer solutions
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC PROBLEM
One dm3 of solution was prepared by dissolving 0.25 mol of formic acid, HCOOH, and 0.18 mol of
sodium formate, HCOONa, in water. What was the pH of the solution? Ka for formic acid is 1.7 x
10-4.
Let x = mol dm-3 of acid that ionize, then
Concentration/mol dm-3
HCOOH(aq)
starting
change
equilibrium
Ka =
0.25
-x

H+(aq) + HCOOˉ(aq)
0
=
0.18
= 1.7 x 10-4
x =
pH =
What is the pH of this solution if 50.0 cm3 of 1.00 mol dm-3 sodium hydroxide is added to 1.00 dm3 of
solution?
moles NaOH added =
equation for acid-base reaction:
OHˉ(aq) + HCOOH(aq) 
moles HCOOˉ now in solution =
Experiment 7b
total volume of solution
68
=
[HCOOH] =
[HCOOˉ] =
Let x = mol dm-3 of acid that ionize, then
Concentration/mol dm-3
HCOOH(aq)

H+(aq) + HCOOˉ(aq)
starting
change
equilibrium
-x
Ka =  =  = 1.7 x 10-4
x =
pH =
What was the pH change?
pH =
If instead of adding the NaOH solution to the buffer solution it had been added to 1 dm3 of a solution
containing 0.25 mol HCl, what would have been the corresponding pH change?
Initial pH of HCl solution:
HCl 
[H+] =
pH =
pH of HCl solution after NaOH addition:
NaOH + HCl 
moles NaOH added =
moles HCl left after reaction with NaOH =
moles H+ left =
total volume of solution =
[H+] =
pH =
pH =
Experiment 7b
Experiment 7b:
69
Equilibria of water, weak acids and bases, and
buffer solutions
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
The volume of 1 M HCl needed to increase the acidity to pH 3 for:
5 cm3 water was ______ drops, and for 5 cm3 1 M CH3COONH4 was ______ cm3.
The volume of 1 M NaOH needed to increase the basicity to pH 12 for:
5 cm3 water was ______ drops, and for 5 cm3 1 M CH3COONH4 was ______ cm3.
Suppose the same volume of 1 M HCl which you added to the 5 cm3 of 1 M CH3COONH4 to produce
a pH value of 3 were added to the 5 cm3 of H2O instead. What would the H3O+ concentration of this
solution now be? Determine the pH of this solution. Show calculations.
Explain, by discussion and equations, how ammonium acetate (a strong electrolyte) is able to buffer
the solution against the addition of both acids and bases.
Experiment 8
70
Experiment 8: Solubility product of a slightly soluble salt
AIM
To undertake a quantitative study of chemical equilibrium as applicable to the case of a saturated
solution of a slightly soluble salt, and to determine the solubility product (Ksp) of potassium
hydrogen tartrate (K+HTa-) at room temperature.
INTRODUCTION
For an elementary reversible reaction taking place at a fixed temperature
A+B  C+D
the rate at which A and B react is proportional to the product of their concentrations, that is,
Ratef = kf[A][B]
where kf is the rate constant and the square brackets denote the concentrations in mol dm-3.
The rate of the reverse reaction is given by:
Rater = kr[C][D].
At equilibrium the rates of the forward and reverse reactions will be equal
Ratef = Rater
kf[A][B] = kr[C][D]
[C][D] k f

K.
[A][B] k r
K is the equilibrium constant of the reaction at the given temperature.
For a solid AB(s) dissolving to form ions A+ and B-:
AB(s)  A+(aq) + Bˉ(aq)
the theoretical equilibrium equation is:
[A  ][B  ]
K
.
[AB(s)]
If the solution of AB is saturated, i.e. the solution is in contact with undissolved salt, the concentration
of AB does not change, or
[AB(s)] = constant.
This constant can be incorporated into K to give
Ksp = [A+][Bˉ].
This equation states that the product of the molar concentrations of anion and cation from a binary
electrolyte (a salt that dissolves to give two ions) is a constant at a specific temperature. This
statement will be verified experimentally by using sparingly soluble potassium hydrogen tartrate, the
half-neutralised salt of a dihydroxy dicarboxylic acid, tartaric acid, abbreviated for convenience to
KHTa.
OH H
HOOC
C C
H OH
COOK
Experiment 8
71
In this experiment a saturated solution of KHTa in pure water is made, and the concentration of
HTaˉ(aq) is determined by titration with a standard base. As this is a binary salt, the concentration of
K+ is the same as that of HTa- thus Ksp = [K+][HTaˉ] = . [HTaˉ]2.
Thereafter a series of saturated solutions of KHTa is made in water containing various concentrations
of KCl. As the presence of KCl in the solution increases the concentration of the common ion, K+, the
concentration of HTaˉ must decrease if Ksp is to remain constant. The concentration of HTaˉ can
decrease only if less KHTa dissolves.
Titration will show that the concentration of HTa- is lower. The concentration of K+, however, is now
the sum of the contribution equivalent to HTaˉ and the contribution from the dissociated KCl.
Thus Ksp = [HTaˉ ][K+ex KHTa + K+ex KCl]
EXPERIMENTAL PROCEDURE
In this practical, students work in groups of at least two members. It is thus essential that members of
each group divide the tasks among themselves so as to avoid congestion at the balances.
1.
Prepare 0.01, 0.02 and 0.04 M KCl solutions by pipetting 5, 10 and 20 cm3 of the 0.5 M stock
solution provided into the numbered volumetric flasks and diluting to 250 cm3 with deionised
water.
2.
Rinse the four bottles marked 1-4 with deionised water and then rinse bottles 2, 3 and 4 with
small quantities of 0.01, 0.02 and 0.04 M KCl solution respectively. There is no need to dry the
bottles.
3.
Weigh four portions of KHTa into bottles 1-4. The mass of KHTa need not be weighed
accurately but must not be less than 1.0 g or more than 1.4 g.
4.
Then add
to bottle 1
deionised water
to bottle 2
0.01 M KCl
to bottle 3
0.02 M KCl
to bottle 4
0.04 M KCl
until the bottle is about two thirds full (approximately 80 cm3). The volume of water or KCl
solution need not be measured accurately. Stopper the bottles and shake them for 5 minutes to
ensure saturation. Place the bottles in the water bath provided and leave to stand for at least 20
minutes so that the suspended solid can settle and the solutions equilibrate to the set
temperature.
5.
Set up a burette with the standard base provided (0.02 M NaOH). Rinse the conical titration
flasks with deionised water. Rinse the pipette with a small volume of the solution to be
measured taking care not to disturb the sediment at the bottom of the bottle.
6.
Check that the supernatant liquor is perfectly clear. If so, pipette three 5 cm3 aliquots from the
first bottle, observing all the precautions outlined below.
Precaution:
When pipetting the supernatant liquor from the bottle, great care must be taken not to disturb
the sediment at the bottom. Thus the tip of the pipette must not be more than 1 cm below the
surface and must be kept steady. Special care must be taken that the liquid in the pipette is not
allowed to run back into the bottle and thus stir up the sediment. In case of a mistake, the
contents of the pipette must be discarded in the sink and a fresh attempt made.
Experiment 8
72
7.
Add 2 to 3 drops of phenolphthalein indicator solution to the titration flask and use standard
titration techniques to titrate the KHTa sample. Titrate all three aliquots before pipetting from
the next bottle. If the agreement between any two of the three is within the prescribed range of
0.10 cm3, the next bottle can be done. If the agreement is not within the range, it is most
probably due to the sediment having been disturbed during pipetting. It is then best to re-do
that bottle after all the others have been done so as to allow the maximum time for the
sediment to settle again.
8.
Repeat steps 6 to 7 for bottles 2-4.
9.
Read the thermometer in the water bath where you placed your bottles and enter this
temperature on your report sheet.
10.
Calculate the molarity of HTaˉ from the titrations and the molarity of K+ from the value of
HTaˉ plus the known molarity (0.00, 0.01, 0.02 and 0.04 M KCl). Calculate Ksp and record the
value in the form B.BB x 10-4 mol2 dm-6 or BB.B x 10-4 mol2 dm-6.
Experiment 8
73
Experiment 8: Solubility product of a slightly soluble salt
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC PROBLEM
A student attempted to determine the solubility product of the slightly soluble salt calcium oxalate,
CaC2O4, by titrating a 200 cm3 sample of a saturated solution of this salt with a 0.00135 mol dm-3 HCl
solution. Two replicate titrations were performed and the results obtained are tabled below.
Initial burette reading/cm3
3
Final burette reading/cm
1
2
2.50
13.70
13.14
24.35
Volume delivered/cm3
Use this data to calculate the Ksp for calcium oxalate at 20 °C.
Experiment 8
74
Experiment 8: Solubility product of slightly soluble salts
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
Temperature _________ °C
Express all concentrations as mol dm-3.
N.B.
Molarity of standard NaOH provided = ____________________________ mol dm-3
Solution 1:
1
2
Initial burette reading/cm3
Final burette reading/cm3
average volume
Volume delivered/cm3
=
conc. of HTaˉ =
conc. of K+ =
Ksp =
= _________________mol2 dm-6
cm3
+
Solution 2:
1
2
Initial burette reading/cm3
Final burette reading/cm3
average volume
Volume delivered/cm3
=
conc. of HTaˉ =
conc. of K+ =
Ksp =
= _________________mol2 dm-6
+
cm3
Experiment 8
75
Solution 3:
1
2
Initial burette reading/cm3
Final burette reading/cm3
average volume
Volume delivered/cm3
=
conc. of HTaˉ =
conc. of K+ =
Ksp =
= _________________mol2 dm-6
cm3
+
Solution 4:
1
2
Initial burette reading/cm3
Final burette reading/cm3
average volume
Volume delivered/cm3
=
conc. of HTaˉ =
conc. of K+ =
Ksp =
= _________________mol2 dm-6
+
Average value for Ksp for four solutions = _________________________mol2 dm-6
cm3
Experiment 9
76
Experiment 9: Freezing point depression
AIM
To determine the molar mass of a soluble unknown compound by measuring the depression of the
freezing point of the solvent.
INTRODUCTION
The freezing point of a mixture (solution) of solute and solvent is lower than the freezing point of the
pure solvent. This fact is applied when salt is spread on iced roads in cold weather in order to cause the
ice to melt.
Freezing point depression may be understood from the following considerations. As the temperature of
a liquid is lowered, the average kinetic energy of the molecules decreases and collisions among them
become less vigorous until, at the freezing point, the attractive forces are able to overcome the
disruptive effect of their kinetic motion, and the molecules “stick together”. In a solution, the solute
molecules interfere with the self-organisation of the solvent molecules to form a solid, and the kinetic
motion must be reduced by a further lowering of the temperature in order for the solvent molecules to
form crystals. Thus, the temperature at which the solvent crystallises, i.e. the freezing point of the
solution, is lower than that of the pure solvent.
The extent of the freezing point depression depends on the number of interfering solute particles and is
quantified by the following relationship:
ΔTf  K f m
where Tf is the freezing point depression, m is the molality of the solution and Kf is the freezing point
depression constant of the solvent. This last quantity changes from one solvent to another.
Because the freezing point depression is nearly proportional to the number of solute molecules in
solution, determination of Tf is one of the simplest and most accurate ways of estimating the molar
mass of a non-dissociating covalent solute. From the measured value of Tf and the known value of Kf,
the molality of the solution can be calculated as
 mol solute 
ΔTf (K)
 
m
1
 kg solvent  K f (K kg mol )
(1)
The molality, m, is defined as the number of moles of solute per kilogram of solvent. The number of
moles of solute is simply the mass of solute, wsolute(g), divided by the molar mass Msolute(g mol-1). The
mass of solvent wsolvent(kg) can be obtained from the volume (cm3) of solvent used and its density (g
cm-3) - remember to convert to kg!
The molality, m, is therefore also given by
m
w solute (g)
w solvent (kg)
M solute (g mol 1 )
By equating the above two equations, we obtain that the molar mass of solute is given by
(2)
Experiment 9
M solute (g/mol) 
w solute (g) x K f (K kg mol 1 )
w solvent (kg) x ΔTf (K)
77
(3)
In this experiment you will determine the molar mass of an unknown fatty acid (solute) by measuring
the freezing point of stearic acid (solvent), and the depression of that freezing point when the solute is
dissolved in the stearic acid. The freezing point depression constant for pure stearic acid is 4.50 K kg
mol-1. You will then identify the unknown fatty acid from the list provided in the Table.
Table: Possible unknown fatty acids
Fatty Acid Unknown
Molar Mass/g mol-1
lauric acid
200.32
myristic acid
228.37
palmitic acid
256.24
EXPERIMENTAL PROCEDURE
Prepare an insulating jacket by inserting an 18 x 150 mm test tube, A, in a 25 x 150 mm test tube,
B, with a rubber cone to provide a seal between the two test tubes. Remove the 18 x 150 mm test
tube A, and reserve the 25 x 150 mm test tube, B, and the rubber cone as the insulating jacket (see
Figure A). The insulating jacket prevents premature cooling due to contact with the skin or other
surface.
Figure A. Schematic for the construction of an insulating jacket.
Determination of the freezing point of pure stearic acid
1.
Determine the mass of the 18 x 150 mm test tube removed from the insulating jacket on an
analytical balance.
2.
Fill the test tube approximately 3/4 full, about 9 grams, with stearic acid and reweigh the
test tube and its contents to determine the exact amount of stearic acid employed.
3.
Prepare a hot water bath by filling a 600 ml beaker 3/4 full with geyser water and heating
with a Bunsen burner. The beaker should be supported on a tripod stand with wire gauze.
4.
Immerse the 18 x 150 mm test tube containing the fatty acid sample in the hot water bath to
melt the fatty acid. After the fatty acid sample has completely melted, place the
thermometer in the fatty acid sample and heat until the sample reaches 85 °C. From this
Experiment 9
5.
6.
78
point on, the thermometer is not removed from the fatty acid sample to prevent loss of
material and contamination of bench tops with fatty acids. Remove the test tube from the
water and dry the outside.
Place the 18 x 150 mm test tube containing the fatty acid sample in the previously prepared
insulating jacket. Stirring constantly with the thermometer, record the temperature of the
sample every 30 seconds for 8-10 minutes. Temperatures are collected until the temperature
of the sample remains constant, changing by less than 0.1 °C per reading, for 3 minutes, 6
readings.
Plot a cooling curve of the temperature readings against the time. Two series will be observed.
Draw the best straight line through each series and determine the freezing point from the point
of intersection of the two lines as shown in Figure B.
Figure B. The figure illustrates the data analysis process for determining the freezing point.
7.
Perform a second trial using the same sample. If your two values do not agree within 0.10
°C, determine the freezing point a third time and average the values.
Determination of the molar mass of an unknown compound
8.
Record the number of your unknown.
9.
Weigh about 2 g of the unknown compound into a weighing boat. Now reweigh the boat to the
nearest 0.001 g. Carefully introduce this solid into the stearic acid in the small test tube. Try
to deposit all the sample directly into the bottom of the test tube without allowing any to stick
to the walls at the top. Reweigh the weighing boat to the nearest 0.001 g. Calculate the mass of
sample placed in the test tube.
10.
Repeat steps 4 and 5 above on this fatty acid sample. Repeat the measurement until you are
satisfied that you have a reproducible value.
Clean-up procedure
1.
After you have completed your final trial, use the hot water bath to reheat the test tube and
the fatty acid sample to 85 °C until all of the fatty acid has melted.
2.
Pour all of the melted fatty acid mixture out into a clean waste container labelled
“PRIMARY FATTY ACID WASTE”.
3.
At this point, you must shut off your Bunsen burner due to the flammability of the 2propanol to be used in the next step.
4.
Fill the test tube 3/4 full with 2-propanol, place it in the hot water bath and stir the mixture
with the thermometer to dissolve all of the residual fatty acids deposited on the sides of the
thermometer and test tube walls.
Experiment 9
5.
6.
Once the sample is dissolved, pour the 2-propanol mixture into the waste bucket labelled
“SECONDARY FATTY ACID WASTE”.
Repeat steps 4 and 5 once or twice more until the test tube is completely clean.
Calculations
Each student must perform the calculations individually.
Calculate:
the mass of stearic acid used as solvent;
the freezing point depression for the solution;
the molality of the solution; and
the molar mass of your unknown.
Hence identify your unknown and calculate the percent error in the molar mass determined.
79
Experiment 9
80
Experiment 9: Freezing point depression
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC PROBLEM
Do colligative properties depend on the number of particles dissolved, the identity of the particles
dissolved, or both?
Show, in detail, how equation (3) in the Introduction can be obtained from equation (1).
Experiment 9
81
Experiment 9: Freezing point depression
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
Mass of stearic acid used
Mass of test tube + stearic acid
__________ g
Mass of empty test tube
__________ g
Mass of stearic acid
__________ g
Determination of the freezing point of pure stearic acid
Run 1
time/s
Run 2
temperature/°C
time/s
Run 3
temperature/°C
time/s
temperature/°C
Experiment 9
82
Run 1
C
freezing point of pure
stearic acid ________ °C
T/ °C
t/s
Run 2
C
freezing point of pure
stearic acid ________ °C
T/ °C
t/s
Run 3
C
freezing point of pure
stearic acid ________ °C
T/ °C
t/s
Average freezing point __________ °C
Experiment 9
83
Determination of the freezing point of a solution
Number of unknown used ___________
Solution 1
Accurate mass of weighing boat + unknown before tipping
__________ g
Accurate mass of weighing boat + unknown after tipping
__________ g
Mass of unknown dissolved
__________ g
Run 1
Run 2
time/s
temperature/°C
time/s
Run 3
temperature/°C
time/s
temperature/°C
Run 1
C
freezing point of
solution ________ °C
T/ °C
t/s
Experiment 9
84
Run 2
C
freezing point of
solution ________ °C
T/ °C
t/s
Run 3
C
freezing point of
solution ________ °C
T/ °C
t/s
Average freezing point of solution
__________ °C
Experiment 9
Calculations
freezing point depression of solution
=
=
= __________ °C
molality of solution 1
=
=
= __________ mol kg-1
molar mass of unknown from solution
=
=
= __________ g mol-1
Identity of unknown:
____________________________
Percent error in molar mass determined
=
=
= ___________ %
85
Experiment 10a
Experiment 10a:
86
Determination of the molar mass of copper by
electrolysis
AIM
To introduce Faraday's First Law of Electrolysis, and to determine the molar mass of copper by
electrolysis.
INTRODUCTION
Electrolysis
An electric current flows in a solid conductor by the movement of electrons. In an electrolyte current
flows by the movement of cations and anions through the solution to the cathode (negative electrode)
and to the anode (positive electrode) respectively. Upon arrival at the negative electrode, cations
accept electrons and so are reduced to atoms, e.g.
H+(aq) + eˉ → H
Cu2+(aq) + 2eˉ → Cu
Upon discharge, the atoms can:
(i)
form molecules and escape as a gas
H + H → H2(g)
(ii)
deposit on the cathode
Cu → Cu(s)
Faraday's Laws of Electrolysis
Faraday was the first scientist to quantify chemical change at an electrode. His first law states:
"The mass of an element liberated at an electrode is proportional to the quantity of electric
charge passed".
The quantity of electric charge is measured in coulombs - 1 coulomb is the quantity of charge passing
when a current of 1 ampere flows for a period of 1 second, i.e. 1 C = 1 A s.
Faraday's first law can thus be written:
m  It
(mass liberated  current x time)
The quantity of charge required to liberate 1 mole of a monovalent element is called the Faraday, F.
Thus 1 F is the charge associated with 1 mole of electrons, or, in practical units, 96485 C. It is thus
possible to determine the molar mass of an element by the measurement of time, current and mass.
In this experiment, the mass of copper deposited on an electrode by a measured quantity of charge will
be determined, and from this its molar mass will be calculated.
The equipment available is rather elementary and fluctuations in the current occur continuously. For
this reason a water electrolysis cell is inserted in series with the copper electrolysis cell and the volume
of hydrogen liberated is used as a second measurement of the quantity of charge passed for calculation
of the molar mass of copper.
Experiment 10a
87
EXPERIMENTAL PROCEDURE
1.
Fill the beaker of the water electrolysis cell to within 2 cm of the brim with 1% H2SO4. Fill the
eudiometer tube with 1% H2SO4, close the mouth with your thumb and up-end the tube in the
H2SO4 solution in the beaker. Insert the tube in the clamps on the column.
2.
Remove the cathode from its clamp and insert it in the eudiometer tube. Reclamp it and
position it, as shown in the Figure, not more than 2 cm inside the mouth of the tube. If it is
positioned higher up the eudiometer tube, the resistance will be too high to attain the required
current.
3.
Pour about 200 cm3 of the saturated CuSO4 solution into the beaker of the copper electrolysis
cell.
4.
Clean one of the copper electrodes with the sandpaper provided. Wipe it with a paper towel,
dry it in the hot air oven and weigh it on a fine balance in the balance room.
5.
Insert the two copper electrodes in the perspex lid and position the spacer between them. Rest
the lid on the beaker containing CuSO4. Clamp the black lead onto the prepared electrode by
means of the crocodile clip. Clamp the red lead onto the other electrode and plug both leads
into their respective sockets.
6.
Turn the control knob of the DC supply fully counterclockwise. Make sure the switch of the
DC supply is in the OFF position. Plug in the DC supply to the bench socket and switch on.
7.
As a potential of approximately 8 V is required to drive a current of 150 mA through the
circuit, ensure that the DC power supply reads about 8 V when the ammeter is reads as close to
150 mA as possible.
8.
Move the switch on the DC supply to the ON position, start the timer and turn the “adjust”
knob on the DC power supply clockwise until the ammeter shows 150 mA, and keep it at 150
± 1 mA by adjusting the same control knob.
9.
Run until approximately 50 cm3 of hydrogen has collected in the eudiometer tube.
10.
Switch off and stop the timer. Record the running time.
11.
Remove the copper cathode, rinse it in a stream of deionised water, dry it in the oven and
weigh it.
12.
Calculate the quantity of charge and then the molar mass of copper.
13.
Record the volume of hydrogen collected. Convert this volume to STP.
14.
From the STP value of the volume of hydrogen collected calculate the molar mass of copper.
15.
Do NOT discard the H2SO4 in the large beaker or the CuSO4 in the small beaker.
Experiment 10a
88
variable DC supply
(-)
(+)
150
mA
red
black
red
Anode
Anode
Cathode
2 cm
CuSO4
1% H2SO4
2
Cu + + 2eAnode: Cu
(Cu from the electrode)
2+
Cathode: Cu + 2eCu
2+
(Cu from the solution)
Experiment 10a
89
TABLE: VAPOUR PRESSURE OF WATER IN kPA
Temp/°C
0.0
0.2
0.4
0.6
0.8
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
37
38
39
40
0.6104
0.6566
0.7057
0.7578
0.8133
0.8722
0.9348
1.0015
1.0724
1.1476
1.2275
1.3122
1.4020
1.4971
1.5979
1.7046
1.8174
1.9368
2.0631
2.1964
2.3374
2.4860
2.6429
2.8084
2.9829
3.1667
3.3604
3.5643
3.7789
4.0047
4.2421
4.4915
4.7539
5.0293
5.3305
5.6219
5.9402
6.2740
6.6239
6.9905
7.3747
0.6194
0.6662
0.7158
0.7686
0.8247
0.8844
0.9479
1.0153
1.0871
1.1633
1.2441
1.3298
1.4207
1.5168
1.6188
1.7266
1.8407
1.9615
2.0892
2.2241
2.3665
2.5167
2.6753
2.8425
3.0190
3.2044
3.4003
3.6064
3.8230
4.0513
4.2911
4.5431
4.8079
5.0861
5.3779
5.6844
6.0057
6.3427
6.6958
7.0661
7.4528
0.6285
0.6758
0.7261
0.7795
0.8363
0.8968
0.9610
1.0293
1.1020
1.1790
1.2608
1.3475
1.4395
1.5367
1.6388
1.7490
1.8645
1.9866
2.1156
2.2520
2.3959
2.5478
2.7081
2.8770
3.0555
3.2427
3.4407
3.6490
3.8677
4.0983
4.3404
4.5950
4.8624
5.1432
5.4381
5.7475
6.0717
6.4120
6.7682
7.1422
7.5328
0.6378
0.6857
0.7365
0.7906
0.8482
0.9094
0.9743
1.0435
1.1770
1.1950
1.2777
1.3655
1.4584
1.5569
1.6612
1.7716
1.8883
2.0118
2.1423
2.2801
2.4257
2.5792
2.7413
2.9119
3.0923
3.2814
3.4814
3.6919
3.9129
4.1459
4.3901
4.6474
4.9176
5.2011
5.4988
5.8112
6.1385
6.4820
6.8414
7.2190
7.6128
0.6472
0.6957
0.7471
0.8018
0.8602
0.9220
0.9879
1.0579
1.1323
1.2112
1.2949
1.3837
1.4776
1.5773
1.6828
1.7944
1.9125
2.0374
2.1691
2.3086
2.4557
2.6109
2.7746
2.9473
3.1294
3.3208
3.5226
3.7352
3.9586
4.1938
4.4405
4.7003
4.9732
5.2596
5.5599
5.8756
6.2059
6.5526
6.9155
7.2964
7.6941
Experiment 10a
Experiment 10a:
90
Determination of the molar mass of copper by
electrolysis
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC PROBLEM
In an electroplating cell, 0.485 g of a metal M are plated out from an acidic solution of MO3 in exactly
one hour using a current of 1.50 A. What is the molar mass of M?
Reduction equation:
Charge passed
+
eˉ → M
=
=
Moles of electrons passed
=
=
Moles of M deposited =
=
Molar mass of M
=
=
A water electrolysis cell connected in series to the above cell liberated 754 cm3 of hydrogen when the
temperature was 288 K and the atmospheric pressure was 763.8 mmHg. The hydrogen was collected
over an aqueous solution as in this experiment. The difference in liquid levels inside and outside the
eudiometer tube was measured to be 152 mm. Use this data to calculate the molar mass of M.
Atmospheric pressure
=
=
kPa
Experiment 10a
Calculation of hydrostatic correction:
d/13.6 =
=
mmHg
=
=
kPa
Vapour pressure (VP) of H2O at 288 K =
 Pressure of H2 collected at 288 K = atm. press. - (hyd. corr + VP(H2O))
=
=
p 1 V1
p V
(all at STP)  2 2 (at experimetal conditions)
T1
T2
Volume of H2 at STP:
V1 =
=
1 mole of H2 at STP occupies _________________ dm3
 Moles of H2 liberated
=
=
Reduction equation for H2:
 Moles of M deposited =
=
Molar mass of M =
+ eˉ → H2
91
Experiment 10a
Experiment 10a:
92
Determination of the molar mass of copper by
electrolysis
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
Mass of cathode before electrolysis
= ______________ g
Mass of cathode after electrolysis
= ______________ g
 Mass of copper deposited during electrolysis
= ______________ g
1
Time
= ___________ s
Current
= ___________ A
 Charge passed during electrolysis
=
= ______________ C
Hence molar mass of copper (show your reasoning):
Experiment 10a
2
93
Volume of gas collected
= _____________ cm3
Barometric pressure
= ____________ mmHg
= ____________ kPa
Difference in liquid levels inside and outside
eudiometer tube (d)
= _____________ mm
Equivalent to a height that a mercury column
would have under these conditions (d/13.6)
= _____________ mmHg
 Hydrostatic correction
= _____________ kPa
Temperature of dilute acid
= _____________ °C
Vapour pressure of H2O at this temperature
(from Table above)
= _____________ kPa
 Pressure of H2 at temperature taken above is barometric pressure –
(hydrostatic correction + VP(H2O))
= _____________
= _____________ kPa
 Volume of H2 corrected to STP conditions
=
= _____________ cm3
Hence molar mass of copper (show your reasoning):
(NB : 1 mmHg = 133.3 Pa)
Experiment 10b
Experiment 10b:
94 The production of an electric potential by
means of oxidation-reduction reactions
AIM
To introduce the concept of an electrochemical half-cell, to illustrate the construction of an
electrochemical cell by the combination of two half-cells, to measure the potential difference between
two connected half-cells, to compare measured potential differences with values calculated from the
table of E° potentials, and to show the variation in potential due to change in concentration of the
electrolyte in a half-cell.
INTRODUCTION
The electrochemical cell
All elements show a tendency to become oxidised, i.e. to lose electrons,
M → Mn+ + neˉ.
In some elements, e.g. sodium, the tendency is strong and oxidation is violent; in other elements, e.g.
gold, the tendency is very weak and the extent of oxidation is negligible.
When a metal is immersed in an electrolyte, atoms enter into solution in the electrolyte as cations:
M → Mn+(aq) + neˉ
The electrons remain on the metal and thus the metal acquires a negative charge. Equilibrium is soon
reached - when the number of atoms going into solution as cations is equal to the number of cations
leaving the solution to return to the metal and become united with electrons to form atoms.
When a metal has a strong tendency to become oxidised, equilibrium is reached when the
concentration of cations in the electrolyte is high and thefore electron charge on the metal (the
electrode) is high. The electrode thus has a high negative charge.
When a metal has a weak tendency to become oxidised, equilibrium is reached when the concentration
of cations in the electrolyte is low. The electrode thus has a low negative charge.
Thus if two different metals are immersed in an electrolyte, the electron charges on them will be
different, i.e. a potential difference will exist between the electrodes, and the magnitude of this
difference will be determined by the respective oxidation tendencies of the two metals. If the two
electrodes are connected externally by means of a conductor, electrons will flow from the electrode
having the higher electron density to the electrode having the lower electron density, i.e. from the more
negative electrode to the less negative electrode. The loss of electrons from an electrode will promote
the forward reaction:
M
→ Mn+ + neˉ
and the influx of electrons into the other electrode will promote the reverse reaction
Mn+ + neˉ → M.
An electrochemical cell thus consists of two half-cells, one being the half-cell in which oxidation takes
place and the other being the half-cell in which reduction takes place.
Experiment 10b
95 The notation for such a cell is:
MaMan+  Mbm+Mb
where the solid vertical line denotes a change of phase and the dotted vertical line denotes a
connection between the electrolyte containing Man+ and the electrolyte containing Mbm+. In this
experiment, four metals will be used: zinc, copper, iron and lead. Each will be immersed in a solution
of its own ions as its electrolyte. The connection between electrolytes will be made by means of a salt
bridge (a U tube containing a gel of saturated potassium nitrate), which allows the passage of ions
between the two electrolytes and thus constitutes the internal connection between the two half-cells as
shown in the Figure.
Electrochemical cell with salt bridge
Conventions
Various conventions exist for writing cell notations. In this practical it will be convenient to write the
notation in such a way that electrons in the external circuit flow from left to right. This means that the
half-cell in which oxidation takes place is on the left.
eˉ
2+
ZnZn  Cu2+Cu
I
The half-cell in which oxidation takes place is called the anodic half-cell and its electrode is called the
anode. The other half-cell is called the cathodic half-cell and the electrode is called the cathode.
When the concentration of ions in the electrolyte is unit molarity (1 M), and the temperature of the
electrolyte is 25 °C, the half-cell is said to be in its standard state and its potential is designated by E°.
It is of course impossible to measure the potential of one half-cell; only a difference of potential
between two half-cells can be measured. The measured potentials will be compared with the values
calculated from the list of E° potentials given in the table on page 59. For this practical, it will be
convenient to use the equation:
or
E°cell = E° (reduction half-cell) – E° (oxidation half-cell)
when using notation I above.
E°cell = E°RHE – E°LHE
E° Values
The table of E° values was established by choosing a reference electrode against which to measure.
For this reference electrode the half-reaction
2H+ + 2eˉ → H2
Experiment 10b
96 was arbitrarily chosen and assigned the value 0.00 V. Note that IUPAC convention writes the equation
as reduction despite the fact that equations are often written as oxidations. Thus interpretation of the E°
values (pg. 59) requires knowledge of the conventions.
The algebraic signs + and – must be read as: + means "more readily than" and – means "less readily
than".
Pairing the standard and any other half-reaction, say that of Zn, one has
2H+ + 2eˉ → H2
Zn2+ + 2eˉ → Zn
0.00 V
–0.76 V.
This means zinc ions are less readily reduced than the standard (H+) by a margin of 0.76 V.
Similarly, pairing the reduction of Cu2+ and that of the standard, one has:
2H+ + 2eˉ → H2
Cu2+ + 2eˉ → Cu
0.00 V
+0.34 V,
showing that copper ions are more easily reduced than the standard (H+) by a margin of 0.34 V.
If convenient, the above statements can be changed to:
H2 → 2H+ + 2eˉ
Zn → Zn2+ + 2eˉ
Cu → Cu2+ + 2eˉ
0.00 V
+0.76 V
–0.34 V
and read as: zinc is more readily oxidised than the standard by a margin of 0.76 V. Cu is less readily
oxidised than the standard by a margin of 0.34 V.
Thus in a combination of a zinc half-cell and a copper half-cell, the tendency to become oxidised is far
greater in the zinc half-cell, consequently electrons will flow from the zinc half-cell to the copper halfcell. The transfer of electrons from the zinc half-cell will promote the reaction
Zn → Zn2+ + 2eˉ
II
i.e. encourage oxidation of zinc atoms, while the movement of electrons into the copper half-cell will
reduce the copper ions
Cu2+ + 2eˉ → Cu.
III
As explained above, the convenient notation for the combinations of these two half-cells is:
ZnZn2+  Cu2+Cu
IV
Calculation of cell potential
This change from standard reduction potentials to oxidation potentials is convenient purely for
determining which half-cell will be the reduction half-cell and which will be the oxidation half-cell.
Calculations of cell potential are done using the standard reduction potentials (pg. 59).
Thus for cell IV above:
Experiment 10b
E°cell
=
=
=
97 E° (reduction half-cell) – E° (oxidation half-cell)
0.34 V – (–0.76 V)
1.10 V.
If a combination of a copper half-cell and a silver half-cell is made, the values of Table III show that
Cu is more readily oxidised than Ag. Thus the cell notation will be:
CuCu2+  Ag+Ag
and the cell potential will be
E° = 0.80 V – 0.34 V = 0.46 V.
If a combination of a magnesium half-cell and a zinc half-cell is made, the values on page 59 show
that Mg is more readily oxidised than Zn. Thus the cell notation will be:
MgMg2+  Zn2+ Zn
and the cell potential will be:
E° = –0.76 V – (–2.34 V) = 1.58 V.
Variation of cell potential
It is worthwhile to note that change of concentration of ions in the electrolyte causes changes in the
potential of a half-cell, but that the size and shape of the metal electrode is immaterial.
EXPERIMENTAL PROCEDURE
1.
Clean the metal electrodes with the sandpaper provided. This must be done on the reverse side
of the tiles provided. It is essential that the surfaces of the electrodes be clean and free of oxide.
Lead foil is very soft and its cleaning requires some dexterity.
2.
Pour approximately 50 cm3 of 1.0 M zinc nitrate into one beaker and approximately 50 cm3 of
1.0 M copper nitrate into another. Place them side by side and connect them with the salt
bridge. Put a copper strip into the copper solution and a zinc strip into the zinc solution.
3.
Set the multitester (also known as an AVO meter, because it can measure current (amp),
potential (volt) and resistance (ohm)) to the following:
DC
(meaning direct current)
V
(meaning potential measurement) and
0-2
(meaning range 0-2 V).
4.
Zero the multitester. Some of these multitesters can be zeroed by short-circuiting the leads and
pressing the "zero adjust" button. After zeroing, the "zero adjust" mode must be cancelled by
pressing the button again. Connect the red lead by means of its crocodile clip to the reduction
half-cell electrode and the black lead to the oxidation half-cell electrode.
5.
Record the highest reading obtained in the initial 30 seconds. Rinse the salt bridge with
deionised water and return it to its beaker.
6.
Remove the zinc half-cell and substitute the lead half-cell. Insert the salt bridge, reconnect the
multitester, zero it and record the cell potential.
7.
Remove the lead half-cell and substitute the iron half-cell. Record the cell potential.
8.
Discard the electrolyte from the copper half-cell and replace it with a 0.1 M solution made by
diluting 10 cm3 of the 1.0 M solution to 100 cm3 in a measuring cylinder. Combine this halfcell with the zinc half-cell and measure the potential. Explain why this potential is less than the
potential measured in 5 above.
9.
Switch off the multitester. Rinse the salt bridge and return it to its beaker. Discard the
electrolytes in the sink and flush with water.
Experiment 10b
98 STANDARD REDUCTION POTENTIALS IN VOLTS AT 25 °C
E°
+ 2.85
+ 1.82
-
F2(g) + 2eˉ  2F (aq)
Co3+(aq) + eˉ  Co2+(aq)
MnO4-(aq) + 8H+(aq) + 5eˉ  4H2O + Mn2+(aq)
ClO4-(aq) + 8H+(aq) + 8eˉ  4H2O + Clˉ(aq)
Cl2(g) + 2eˉ  2Clˉ(aq)
Cr2O72-(aq) + 14H+(aq) + 6eˉ  7H2O + 2Cr3+(aq)
MnO2(s) + 4H+(aq) + 2eˉ  2H2O + Mn2+(aq)
O2(g) + 4H+(aq) + 4eˉ  2H2O(l)
Br2(l) + 2eˉ  2Brˉ(aq)
NO3ˉ(aq) + 4H+(aq) + 3eˉ  2H2O +NO(g)
Hg2+(aq) + 2eˉ  Hg(l)
Ag+(aq) + eˉ  Ag(s)
Fe3+(aq) + eˉ  Fe2+(aq)
MnO4ˉ(aq) + 2H2O + 3eˉ  4OHˉ(aq) + MnO2(s)
I2(s) + 2eˉ  2Iˉ(aq)
Cu2+(aq) + 2eˉ  Cu(s)
Sn4+(aq) + 2eˉ  Sn2+(aq)
2H+(aq) + 2eˉ  H2(g)
Pb2+(aq) + 2eˉ  Pb(s)
Sn2+(aq) + 2eˉ  Sn(s)
Ni2+(aq) + 2eˉ  Ni(s)
Cd2+(aq) + 2eˉ  Cd(s)
Fe2+(aq) + 2eˉ  Fe(s)
Zn2+(aq) + 2eˉ  Zn(s)
Mn2+(aq) + 2eˉ  Mn(s)
Al3+(aq) + 3eˉ  Al(s)
Mg2+(aq) + 2eˉ  Mg(s)
Na+(aq) + eˉ  Na(s)
Ca2+(aq) + 2eˉ  Ca(s)
Li+(aq) + eˉ  Li(s)
+ 1.52
+ 1.39
+ 1.36
+ 1.33
+ 1.23
+ 1.23
+ 1.06
+ 0.96
+ 0.85
+ 0.80
+ 0.77
+ 0.59
+ 0.54
+ 0.34
+ 0.15
0.00
– 0.13
– 0.14
– 0.24
– 0.40
– 0.44
– 0.76
– 1.18
– 1.66
– 2.34
– 2.71
– 2.87
– 3.04
Experiment 10b
Experiment 10b:
99 The production of an electric potential by
means of oxidation-reduction reactions
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
PRE-PRAC PROBLEM
Sketch a cell in which a cadmium electrode is in a solution of cadmium nitrate and a silver electrode is
in a solution of silver nitrate. The half-cells are connected by a salt bridge. On your sketch label:
(i)
the anode,
(ii)
the cathode,
(iii)
the direction of electron movement in the external circuit, and
(iv)
the direction of ion movement (cations and anions).
Write
(v)
the electrode half-reactions, and
(vi)
the overall cell reaction.
Calculate E°cell for the above cell.
Experiment 10b
Experiment 10b:
100 The production of an electric potential by
means of oxidation-reduction reactions
Name:
_____________________________
Student no.:
_____________________________
Lab. number:
_____________________________
Seat Number:
_____________________________
Date:
_____________________________
Demonstrator: ________________________
Mark:
RESULTS
Potentials and Components of Electrochemical Cells
Cell
Cell
Potential/V
Anode
(negative
electrode)
Cathode
(positive
electrode)
ZnZn2+  Cu2+Cu
PbPb2+  Cu2+Cu
FeFe2+  Cu2+Cu
Effect of Concentration on the Cell Potential of the Zinc-Copper Cell
Original cell potential/V
Cell potential after dilution of Cu2+ solution/V
Change in cell potential/V
What is your explanation of these results?
Standard cell potentials
calculated from Table
III
Appendix 1
Appendix 1:
101
Laboratory apparatus
Each student workbench is supplied with the following apparatus, which should be returned to its
correct locker at the end of each laboratory session.
Item
Description
Quantity
Erlenmeyer (conical) flasks
Erlenmeyer (conical) flasks
Beakers
Beakers
Measuring cylinder
Measuring cylinder
Volumetric flask
Pyrex filter funnel
Pyrex filter funnel
Weighing bottles
Weighing boat
Pipette filler
Glass rod
Droppers and teats
Buchner funnel
Hirsch funnel
Pasteur pipette
Pyrex filter flask
Spatula
Retort Clamps
Boss heads
Washbottle
Wooden peg
Test tube rack
Test tubes
Test tube brush
Tripod stand
Gauze mats
Nutec Mat
Mini Quickfit clamp
Rubber adaptor for Buchner funnel
Ice bucket
Vacuum pipe
Bunsen burner
250 cm3
50 cm3
100 cm3
50 cm3
50 cm3
10 cm3
250 cm3
45 mm diameter
75 mm diameter
No. 4 vial
(plastic)
Pi-pump type
4mm diameter.
Pasteur
5.5 cm
45 mm
4
1
2
2
1
1
1
1
1
2
1
1
1
2
1
1
1
1
1
1
1
1
1
1
10
1
1
1
1
1
1
1
1
1
100ml
(medium)
Plastic-250 cm3
(small)
(on bench)
Appendix 1
102
Appendix 1
103
Appendix 2
Appendix 2:
104
The laboratory balance
Of all the instruments used by a chemist, the balance is the most important. It is, therefore, essential
that students should learn to weigh accurately and rapidly from the beginning. However, a balance
is an extremely delicate instrument and in order that it should retain its accuracy it is imperative
that it is handled with the utmost care.
The most important attributes of a chemical balance are its sensitivity and rapidity of action, and
the student should appreciate that, since most errors in analysis arise from faulty weighing, it is
impossible to take too much care using the balance.
Using the Balance
The balances used in this laboratory are both extremely expensive and highly sensitive! Please be
CAREFUL when using them!
1. Do not move the balance on the balance bench.
2. Never touch the balance pan with your fingers, or breathe on it.
3. Items to be weighed must always be at room temperature.
4. Never weigh corrosive or volatile substances in an open vessel. A stoppered weighing
bottle should be used for such substances to eliminate the danger of corrosion. Do not
spill chemicals on the balance pan, in the balance or on the balance bench. To minimise
this danger, when weighing a substance, never add or remove any of the substance from
the weighing receptacle while it is on the balance pan. First remove it from the balance
case and then add or remove substance as necessary.
5. Report immediately any accidental spilling of chemicals on the balance to your
demonstrator so that the spill may be cleaned at once.
6. Please leave the balance in the condition in which you would wish to find it, i.e.
perfectly clean.
Appendix 3
Appendix 3:
105
Volumetric apparatus
ERROR OF PARALLAX
Volumetric apparatus should be read with your eye on the same level as the meniscus (curved
surface of the liquid in the vessel). If this is not done, an incorrect reading will be taken. This
source of error is known as an error of parallax.
Always read the bottom of the meniscus.
THE PIPETTE
The laboratory pipette is a glass vessel which delivers a definite volume of liquid under certain
specific conditions. It consists of a cylindrical bulb with tubes at each end. Pipettes come in
different sizes, ranging from 1 to 250 cm3.
The letter B on the bulb signifies that the pipette is a “B” grade pipette, i.e. the manufacturer
guarantees that the volume delivered will be between  0.05 cm3 of the stated volume. “A” grade
pipettes are guaranteed to deliver between  0.01 cm3.
The pipette
Washing and preparing the pipette
The pipette should always be assumed to be dirty, and must be rinsed with both tap water and deionised water, and then three times with the solution to be dispensed.
Appendix 3
106
After rinsing the pipette with both tap water and de-ionised water, dispense ~20 cm3 of the solution
provided into a clean, dry 100 cm3 beaker. Insert the upper end of the pipette into the lower end of
the pipette filler and rotate the pipette carefully to work no more than 0.5 cm of the glass bore into
the rubber sleeve (If you push the pipette further in the filler will not work properly). Do NOT suck
up solutions by mouth; use the pipette filler at all times for filling the pipette with solution.
Use the wheel on the pipette filler to suck solution into the pipette. Remove the pipette from the
solution, detach the filler and immediately place your finger over the end (why?). Hold the pipette
as level as possible, and rotate the pipette on the fingertips to ensure that solution flushes over the
entire inner surface, including a length of 2-3 cm above the graduation mark. Allow the pipette to
drain through the lower end only, and repeat the process twice.
Using the Pipette
In volumetric analysis the procedure of pipetting is the single largest source of error for beginners.
The technique of pipetting requires careful practice, and will be demonstrated to you, in addition to
the notes on the pages that follow. If a thorough study is made of these pages before the laboratory
session, you should be able to follow the demonstration and subsequently be able to use the pipette
correctly and accurately.
After preparing the pipette, discard the remainder of the solution in the 100 cm3 beaker and refill it.
Reassemble the pipette and pipette filler and draw up solution sufficient to fill the pipette to ~2-3
cm above the graduation mark on the pipette (a). Do not allow any liquid to enter the pipette filler;
if you accidentally do so, or you think that the filler has liquid in it already, show your
demonstrator.
Remove the pipette filler and place your finger over the open end of the pipette (b). Wipe the
outside of the pipette with a tissue, being careful not to touch the point of the pipette with the tissue
otherwise solution will be lost by capillary action (c). Keeping the pipette at eye level and in a
vertical position (using both hands), allow the liquid level to fall by lifting your finger slightly, so
that the bottom of the meniscus becomes just level with the graduation mark on the pipette. If a
small drop of solution remains at the pipette tip at this stage, it can be removed by touching the
bottom of the pipette against the side of the beaker (d).
Appendix 3
107
Replace the beaker under the pipette with a conical flask, and, again keeping the pipette vertical,
remove your finger to enable the solution to drain into the flask. Once the solution has finished
draining, touch the bottom of the pipette against the side of the conical flask. The liquid remaining
in the tip of the pipette at this stage must not be blown out into the conical flask, as this extra drop
is taken into account when the pipette is calibrated.
Once pipetting is complete, clamp the pipette on the burette stand to prevent accidental breakage.
Do not leave the pipette on the bench top where it can become contaminated.
THE BURETTE
A burette consists of a tube of uniform bore graduated in cm3 and tenths of a cm3. The volume
delivered can be read on the graduations accurately to the first decimal, i.e. to 0.1 cm3. Readings
are recorded to the nearest 0.02 cm3 by estimation.
The burette
Washing and preparing the burette
A burette should always be assumed to be dirty, and must therefore be rinsed thoroughly before
use, both with tap water and then with de-ionised water. It is then “prepared” by rinsing the inside
with the solution to be used.
After rinsing the burette with both tap water and de-ionised water, dispense ~25 cm3 of the solution
provided into a clean, dry 100 cm3 beaker. Pour about ~8 cm3 of this solution into the burette (make
sure that the burette tap is closed if you don’t want to go home with wet feet!), hold the burette as
level as possible, and rotate on the fingertips to ensure that solution flushes over the entire inner
surface. Drain through the stopcock to make sure that the solution runs freely and is not obstructed.
Repeat twice, draining the final rinse through the open end of the burette.
Discard the remainder of the solution in the 100 cm3 beaker and refill it, then use the beaker to fill
the burette to above the zero mark. Open the stopcock to displace air from the jet and to ensure that
the jet is completely filled with solution. You will be shown how to remove persistent air bubbles.
Finally allow to drain/top up until the meniscus lies between 0-1 cm3. Place the burette in the
burette stand with a tissue under the stopcock while preparing the remainder of the glassware for
Appendix 3
108
the titration; if the stopcock leaks, even very slowly, the wet tissue will alert you and you can
remedy the problem before wasting time performing inaccurate and meaningless titrations.
Using the Burette
After pipetting, the greatest sources of error in volumetric analysis stem from students not being
able to master the techniques of estimating the volume correctly, and adding a single drop from a
burette.
Adding a single drop of solution is a technique
that requires a little practice. For a right-handed
person it is customary to swirl the conical flask
with the right hand whilst operating the tap with
the left hand, as shown in the picture. You will
be shown how to do this - ask your demonstrator
to check that you are doing it correctly.
As mentioned before, only the first decimal place of a volume can be read directly from the burette;
the second decimal place is then estimated to the nearest 0.02 cm3. Being able to estimate correctly
is thus of great importance if your volumetric analyses are to be accurate. Study the sketch below to
give you an idea, and then get your demonstrator to check that you are doing it correctly.
21
21
21
21
21
22
22
22
22
22
21.30
21.34
21.36
21.38
21.40
Appendix 3
109
Placing a white card below the level of the
meniscus, as shown in the diagram alongside,
also aids in taking an accurate reading from a
burette.
The pipette and burette should be clamped
safely when not is use, as shown in this figure.
Appendix 4
Appendix 4:
110
Experimental errors
INTRODUCTION
In nearly all scientific endeavours, measurements are made and the data so obtained used in
calculations to arrive at a result on which a final conclusion is based.
In practice it is seldom possible to make exact measurements. There are unavoidable errors inherent
in the apparatus used, in the methods employed and in the observational powers of the
experimenter. In some instances, errors may not affect the result or the conclusion, but in most
instances the result is open to a degree of doubt and the extent of the error in the result must be
conveyed by following established convention.
An example of the first instance is the result and conclusion based on the following data:
The mass of a consignment of apples is 123 kg. If the mass of a box of apples is 20 kg, how many
boxes are there in the consignment? The result of dividing 123 kg by 20 kg per box is 6.15 boxes,
but the conclusion will be that there are 6 boxes in the consignment.
An example of the second instance is a determination of the STP molar volume of a gas which gave
a value of 22.1 dm3. As the accepted value is 22.4 dm3, it is obvious that errors inherent in the
apparatus, method and operation have (in combination) caused a deviation of 0.3 dm3 from the
accepted value, i.e. an error of
0.3  100
 1.3% .
22.4
SOURCES OF ERROR
In scientific endeavour, errors arise from:




imperfections in apparatus
imperfections in method
variations in the environment
imperfections in technique.
Imperfections in apparatus
Most of the apparatus used in a first course are produced in large batches in order to reduce cost.
Consequently the items are not individually calibrated. However, the manufacturers issue
certificates with each batch indicating the limit of inherent error in any one item of a batch. For
example: a B grade burette certificate may state: Tolerance 50 cm3  0.1 cm3 and a B grade pipette
certificate may state: Tolerance 10 cm3  0.05 cm3. These statements mean that a titration volume
could be in error up to a maximum of 0.2% when using a B grade burette and a 10 cm3 aliquot
pipetted by means of a B grade pipette could be in error up to a maximum of 0.5% due to
imperfections in the shapes of these glass vessels.
Mass measurement apparatus has much lower tolerances than volume measurement apparatus, and,
for all practical purposes, mass measurement errors are negligible in comparison with volume
measurement errors.
Appendix 4
111
Imperfections in method
Imperfections in method likely to be encountered in a first course are: solubility of a precipitate, coprecipitation, decomposition and/or volatilisation during drying in an oven, absorption of moisture
or carbon dioxide from the air, oxidation on exposure to the atmosphere, attack on glass vessels by
caustic substances, etc.
Variations in the environment
Variations in temperature, pressure and humidity of the air in a laboratory occur due to change in
weather conditions, change of the seasons and diurnal fluctuations.
Imperfections in technique
Operator error is usually far greater than all previously mentioned errors combined. Some errors are
unavoidable, for example the perception of indicator colour change is a subjective judgement on the
part of the operator and is bound to differ from the perceptions of other operators. Avoidable errors
are many and often occur because of inattention. The most frequent and serious error in a first
course is incorrect pipetting procedure. Next come the errors due to ineffectual washing of a
precipitate and excessive washing of a precipitate. More serious is physical loss of reagent through
spillage.
In assessing the results of first year students in quantitative analysis, an error of 2% is regarded as
acceptable and may earn full marks. This 2% is made up of 1% operator error and 1% for all other
sources of error combined.
SIGNIFICANT DIGITS
If, in the example above of the determination of STP molar volume, the experimenter reported his
result as 22.100 dm3, he would have been guilty of an absurdity.
The accepted value of 22.414 dm3 has been obtained in advanced research using sophisticated
equipment and working under ideal conditions. According to the conventions of statistics, the
number 22.414 guarantees the value to be closer to 22.414 than it is to 22.413 and to 22.415, i.e. it
lies within the range 22.4135 and 22.4145. (If the uncertainty had been greater, the range would
have been given as, e.g. 22.414  0.001).
When given as 22.414, the uncertainty, and therefore the likely error, is:
0.001  100
 0.00446% . i.e. less than 0.005%.
22.414
In terms of the 2% error allowed in a first course, a result within the range 22.4  22.4 x 0.02 (i.e.
within the range 22.8 and 22.0) would have been acceptable.
Thus even if the experimenter’s calculator showed a final result of say 22.100, the result is still
reported as 22.1 in order not to claim a degree of certainty that does not exist.
To be able to decide on the correct procedure, the student must understand the meaning of the term
significant digits. The number system has ten digits 0, 1, 2… 8, 9. Any given number/figure/value
consists of one or more digits. Thus the statement “12 apples” leaves no doubt as to the quantity.
The number has been obtained by counting, and, except for a blunder, is definite and indisputable.
However, a statement such as “12 km” is the outcome of a measurement and immediately the
question arises how the measurement was made. If the person making the statement had measured
Appendix 4
112
the distance by walking, the reader would interpret the distance as 12  1 km or even 12  2 km,
knowing how fallible the measuring procedure was.
If the measurement had been made by a surveyor using a tellurometer, he would give the distance
as 12.0000 km. The four zeroes following the decimal are not an attempt at pedantism, they are
given deliberately to assure the reader that the surveyor is prepared to guarantee that the value is
closer to 12.0000 than it is to 11.9999 or to 12.0001, i.e. it lies within the range 11.9995 and
12.0005. The surveyor thus implies that his error is not greater than 1 dm.
In statistical language it is said that the number is given to six significant digits, the first five being
beyond dispute and the last digit being uncertain because it is a “best estimate”.
Whereas the zeroes in the surveyor’s 12.000 km are deliberate and therefore significant, the zeroes
in the distance 0.0012 km are not significant because they can be dispensed with by reporting the
distance as 1.210-3 km or 1.2 m or 12 dm. In all three forms there are two significant digits.
However, if this distance were given as 120 cm, it would have three significant digits and thus
claim greater accuracy, i.e. a smaller uncertainty. The value 12 dm implies that the distance lies
within the range 11.5 and 12.5 dm – a variation of 1 dm or 10 cm. The value of 120 cm implies that
the distance lies within the range 119.5 and 120.5 cm – a variation of 1 cm or 0.1 dm.
Similarly, if the distance were given as 1200 mm, i.e. to four significant digits, the variation would
be 1 mm or 0.1 cm or 0.01 dm. Round numbers such as 100, 10 and 1 are considered to have only
one significant digit because the zero’s merely indicate the order of magnitude.
Appendix 4
113
The first rule of significant digits
Measurements, i.e. observed quantities, should
be recorded with one uncertain digit retained.
Figure 1 illustrates this rule. It shows the level
of titrant in a burette. It is clear that the volume
is more than 27.4 but less than 27.5 cm3. The
experimenter has made an estimate of the
interval beyond 27.4 and arrived at the fraction
⅔. He has thus recorded the volume as 27.4 + ⅔
 0.1 = 27.4 + 0.07 = 27.47 cm3 and will use
this value in subsequent computations. The
value 27.47 consists of four significant digits,
the first three indisputable and the last uncertain
because it has been estimated.
The second rule of significant digits
In carrying through a series of computations, one digit beyond the last significant digit must be
retained in order that the last significant digit is not altered in the computational process. This rule
can be illustrated by the following example:
If the experimenter above did two further titrations and obtained say 27.42 and 27.40 cm3 in his
second and third titrations, his calculator would show the mean value as 27.433333 cm3. In further
computations the experimenter must use the value 27.433 cm3 in order to conform to rule 2.
However, if the experimenter obtained say 27.32 and 27.50 cm3 in his second and third titrations,
his mean would still be 27.43333 cm3 but now it would be absurd to use 27.433 cm3 in further
computations as his third digit is already uncertain. He should now use 27.43 cm3 as the mean value
in further computations.
The third rule of significant digits
In a result, there must be as many digits as will give one and only one uncertain digit. This rule has
been re-phrased in a rather simplistic form to read: the result cannot contain more digits than the
factor with the fewest significant digits.
This rule can be illustrated by carrying the example used above to completion. The mean of three
titrations was 27.433 cm3. If the molarity of the titrant was, say 0.1013, and the aliquot 10 cm3
(measured by means of a pipette) the calculation of the unknown concentration is:
27.433 cm 3 x 0.1013 M
unknown molarity 
 0.2778962 M .
10 cm 3
As the volume of the aliquot, 10 cm3, appears to have only one significant digit, the result must be
recorded as 0.2 M according to rule 3. However, the term “10 cm3 pipette” merely indicates an
order of magnitude and does not describe the capacity properly. The manufacturer’s certificate
gives the tolerance of a 10 cm3 pipette as a 0.05 cm3. The lower value, 9.95, has three significant
digits and this fixes the number of significant digits in the result at three.
Appendix 4
114
Thus the result is given as 0.278 M. Had the calculator shown a final value of say 0.27741879 M,
the value reported would be 0.277 M to conform to the rule of rounding off.
The fourth rule of significant digits
The precision of a result must not be diminished by a computation. In the calculation
0.081  141.8  11.4858
the factor 0.081 contains two significant digits and the inclination might be to limit the product to
two significant digits, i.e. to report the product as 11 according to the simplistic version of rule 3
above.
Although the factor 0.081 has only two significant digits, its uncertainty is 1 in 81 whereas the
uncertainty of the product is 1 in 11. Thus the precision of the result has been diminished. In order
to conform to rules 3 and 4, the result must be reported as 11.5. The computations below show that
the first decimal place is the first variable digit,
0.0805 x 141.8 = 11.4149
0.0815 x 141.8 = 11.5567
and a result must always contain one uncertain digit according to rule 3.
By contrast, in the calculation
0.012 x 141.8 = 1.7016
the product would be given to two significant digits, viz. 1.7 because the second digit in 1.7 is
uncertain as shown by the computations below:
0.0115 x 141.8 = 1.6307
0.0125 x 141.8 = 1.7725
Appendix 4
115
Multiplication or division by a factor merely changes the order of magnitude; it does not affect the
precision. Thus if 100.0 g is required to make 1000 cm3 of 1 M solution, then the mass required to
make 200 cm3 of 0.1 M solution is:
100.0 g 
200 cm 3 0.1 M

 2.000 g .
1000 cm 3 1.0 M
Subtraction can reduce the number of significant digits. When the mass of substance is obtained by
difference, then the difference of two balance readings will have fewer significant digits than the
balance readings.
e.g.
mass of vial + substance
mass of vial
mass of substance
12.106 g
11.271 g
0.835 g
5 significant digits
5 significant digits
3 significant digits
The converse is not true, i.e. addition cannot increase the number of significant digits, e.g. 0.603 +
0.731 = 1.33 g.
DRILL PROBLEMS
(You may refer to the answers to some of these problems, after solving them on your own
initiative).
1. Carry out the operations on the data given in each of the following cases to calculate the quantity
called for. Show your method, including the dimensions of measurement. (These units will tell
you which mathematical operation to perform).
(a) Velocity = 50 km h-1, time = 0.5 h, distance = ?
(b) Velocity = 300 000 km s-1, distance = 150 000 000 km, time = ?
(c) Time = 9.3 s, distance = 100 m, velocity = ?
(d) Density Al = 2.70 g cm-3, mass = 2700 g, volume = ?
(e) Mass Hg = 272 g, volume = 20 cm3, density = ?
(f) Mass apples = 120 kg, mass in each box = 20 kg box-1, number of boxes = ?
(g) Mass H2O = 180 g, molar mass = 18 g mol-1, number of moles = ?
2. How many significant digits are there in each of the following numbers?
(a) 3005
(b) 3500
(c) 0.035
(d)
(e)
(f)
0.350
3.050
3.0005
3. Carry out the following operations, recording the answer correctly in accordance with the rules
of significant digits:
(a) Subtract 5.1 from 28.347
(b) Subtract 5.10 from 28.347
(c) Multiply 0.020 by 1.111
(d) Divide 36.02 by (3.0)2
4. A beaker of water has a mass of 1200 grams. How would you write this number so as to avoid
ambiguity, if the mass is known to the nearest:
(a) ten grams
(b) one hundred grams
Appendix 4
116
(c) gram
(d) tenth of a gram.
5. The length of a table is measured as 2 metres, 3 centimetres, and 4 millimetres. Express this
length as:
(a) metres
(b) centimetres
(c) millimetres
(d) kilometres.
How many significant digits in each case?
6. A series of beakers have the following masses: 125.2 g, 90.3 g, 56.2 g and 20.237 g. How should
you record the sum of these masses so as to avoid any incorrect conclusions as to the precision
of measurements?
7. Three determinations of the percentage of chlorine in sodium chloride were 60.1%, 60.5% and
60.3%, averaging 60.3%. The accepted value, based on the atomic masses (Na 22.9979 amu; Cl
35.4571 amu), is 60.650% Cl. What is the percentage error in the analysis, and to how many
significant digits should it be expressed?
8. What is the percentage of uncertainty in measuring 50 cm3 of water in a 50 cm3 graduated
cylinder given that the precision of measurements is 0.2 cm3?
Appendix 4
ANSWERS TO DRILL PROBLEMS
1.
(b)
(d)
(f)
500 s
1000 cm3
6 boxes
2.
(b)
(d)
(f)
2
3
5
3.
(a)
(c)
23.2
0.022
5.
(b)
203.4 cm
(d)
0.002034 km
Four significant digits regardless of the units used.
7.
0.6%
117
Appendix 5
Appendix 5:
Name
Actinium
Aluminium
Americium
Antmony
Argon
Arsenic
Astatine
Barium
Berkelium
Beryllium
Bismuth
Boron
Bromine
Cadmium
Calcium
Californium
Carbon
Cerium
Cesium
Chlorine
Chromium
Cobalt
Copper
Curium
Dysprosium
Einsteinium
Erbium
Europium
Fermium
Fluorine
Francium
Gadolinium
Gallium
Germanium
Gold
Hafnium
Helium
Holmium
Hydrogen
Indium
Iodine
Iridium
Iron
Krypton
Lanthanum
Lawrencium
Lead
Lithium
Lutetium
Magnesium
Manganese
Mendelevium
118 The elements
SYMBOL
Z
Atomic mass/
amu
Ac
Al
Am
Sb
Ar
As
At
Ba
Bk
Be
Bi
B
Br
Cd
Ca
Cf
C
Ce
Cs
Cl
Cr
Co
Cu
Cm
Dy
Es
Er
Eu
Fm
F
Fr
Gd
Ga
Ge
Au
Hf
He
Ho
H
In
I
Ir
Fe
Kr
La
Lr
Pb
Li
Lu
Mg
Mn
Md
89
13
95
51
18
33
85
56
97
4
83
5
35
48
20
98
6
58
55
17
24
27
29
96
66
99
68
63
100
9
87
64
31
32
79
72
2
67
1
49
53
77
26
36
57
103
82
3
71
12
25
101
227.0
26.98
(243)
121.8
39.95
74.92
(210)
137.3
(247)
9.012
209.0
10.81
79.90
112.4
40.08
(251)
12.01
140.1
132.9
35.45
52.00
58.93
63.55
(247)
162.5
(252)
167.3
152.0
(257)
19.00
(223)
157.25
69.73
72.61
197.0
178.5
4.003
164.9
1.008
114.8
126.9
192.2
55.85
83.80
138.9
(260)
207.2
6.941
175.0
24.31
54.94
(258)
Name
Mercury
Molybdenum
Neodymium
Neon
Neptunium
Nickel
Niobium
Nitrogen
Nobelium
Osmium
Oxygen
Palladium
Phosphorus
Platinum
Plutonium
Polonium
Potassium
Praseodymium
Promethium
Protactinium
Radium
Radon
Rhenium
Rhodium
Rubidium
Ruthenium
Samarium
Scandium
Selenium
Silicon
Silver
Sodium
Strontium
Sulphur
Tantalum
Technetium
Tellurium
Terbium
Thallium
Thorium
Thulium
Tin
Titanium
Tungsten
Uranium
Vanadium
Xenon
Ytterbium
Yttrium
Zinc
Zirconium
SYMBOL
Z
Atomic mass/
amu
Hg
Mo
Nd
Ne
Np
Ni
Nb
N
No
Os
O
Pd
P
Pt
Pu
Po
K
Pr
Pm
Pa
Ra
Rd
Re
Rh
Rb
Ru
Sm
Sc
Se
Si
Ag
Na
Sr
S
Ta
Tc
Te
Tb
Tl
Th
Tm
Sn
Ti
W
U
V
Xe
Yb
Y
Zn
Zr
80
42
60
10
93
28
41
7
102
76
8
46
15
78
94
84
19
59
61
91
88
86
75
45
37
44
62
21
34
14
47
11
38
16
73
43
52
65
81
90
69
50
22
74
92
23
54
70
39
30
40
200.6
95.94
144.2
20.18
237.1
58.69
92.91
14.08
(259)
190.2
16.00
106.4
30.97
195.1
(244)
(209)
39.10
140.9
(145)
231.0
226.0
(222)
186.2
102.9
85.47
101.1
150.4
44.96
78.96
28.09
107.9
22.99
87.62
32.07
181.0
(98)
127.6
158.9
204.4
232.0
168.9
118.7
47.88
183.9
238.0
50.94
131.3
173.0
88.91
65.39
91.22
NOTE: Atomic masses in this table are given relative to carbon-12 and limited to four significant figures, although
some atomic masses are known more precisely. For certain radioactive elements the numbers listed (in brackets) are the
mass numbers of the most stable isotopes.
Ia
1
H
1.008
3
VIII a
Periodic Table of the Elements
2
II a
III a
IV a
Va
VI a
VII a
4
5
6
7
8
9
He
4.003
10
Li
Be
B
C
N
O
F
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
12
13
14
15
16
17
18
Na
Mg
22.99
24.31
19
20
III b
21
IV b
22
Vb
23
VI b
24
VII b
25
VIII b
27
26
Ib
29
28
II b
30
Al
Si
P
S
Cl
Ar
26.98
28.07
30.97
32.07
35.45
39.95
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
85.47
87.62
88.91
91.22
92.91
95.94
(98.91)
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
*La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209.0)
(210.0)
(222.0)
87
88
89
Fr
Ra
**Ac
(223.0)
(226.0)
(227.0)
58
*Lanthanides
**Actinides
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
140.1
140.9
144.2
(146.9)
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
(232.0)
(231.0)
(238.0)
(237.1)
(244.1)
(243.1)
(247.1)
(247.1)
(251.1)
(252.1)
(257.1)
(258.1)
(259.1)
(260.1)
Appendix 6
Appendix 6:
Name
120
Common Solvents
Properties*
Formula
(i)
(ii)
(iii)
(iv)
(v)
Disposal
Benzene
C6H6
80
0.9
V
F
I
WB
Dichloromethane
(Methylenechloride)
H2CCl2
40
1.3
V
NF
I
WB
Ethanol (alcohol)
C2H5OH
78
0.8
V
F
M
Sink
Ethoxyethane
(Diethyl ether)
(C2H5)2O
35
0.7
V
F
I
WB
Ethyl ethanoate
(Ethyl acetate)
C2H5COOCCH3
77
0.9
V
NF
I
WB
Hexane
C6H14
70
0.7
V
F
I
WB
Methanol
CH3OH
65
0.8
V
F
M
Sink
Propan-2-one (acetone)
(CH3)2CO
56
0.8
V
F
M
WB
Pyridine
C5H5N
115
1.0
NV
NF
M
WB
Trichloromethane
(Chloroform)
HCCl3
61
1.5
V
NF
I
WB
Water
H2O
100
1.0
NV
NF
-
Sink
*Properties
(i)
Boiling point in °C at 1 atmosphere pressure
(ii)
Density in g cm-3 at 20 - 25 °C
(iii)
Volatility: V = volatile, NV = non-volatile
(iv)
Flammability: F = flammable, NF = non-flammable
(v)
Miscibility with water: M = miscible, I = immiscible
Disposal: WB = waste bottle in fume cupboard
Sink = in sink and flushed down with water