Intro Lab Methods
Chapter 9
Volumetric Analysis
The terms volumetric analysis, titrimetry and titration are used interchangeably to
describe a procedure which analyses chemicals in solution by accurate volume
measurement.
Principles of titration
Titration is used to determine the concentration of an analyte in solution (the
unknown solution) and hence the components of any material which can be made up
into solution may be analysed this way. The method operates by reacting the analyte
in the unknown solution with another solution (the standard solution). The process is
carried out so that reaction between the two can be stopped exactly at the point
when the last trace of analyte has been used up. The volume of both standard and
unknown are used to obtain the required quantitative information.
Schematic outline of the titration process
Practical requirements
To perform titrimetry, the following need to be available:
• standard solutions
• volumetric glassware
• a suitable method of detecting the end-point
• a minimum quantity of analyte
• a standard that reacts appropriately with your analyte
The titration procedure
Burette
Rinse the burette with a small amount of the fill solution as the final step before filling
Overfill with the required solution
Check there are no air bubbles above / below the tap and in the bore of the tap
Wipe any suspended liquid off the burette tip
Zero the burette or record its initial volume. Re-wipe the burette tip
Apparatus used to perform titration
Titration flask
Rinse the flask with distilled water as the final step before adding the aliquot of
solution. It may be left wet.
Rinse a pipette with a small amount of the fill solution as the final step before filling.
Add the aliquot of solution to the flask. (Sometimes an accurately weighed primary
standard solid is used as the reactant, which needs a suitable amount of water
added to dissolve it).
Add a suitable indicator.
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Rough titration
Add the solution from the burette in a slow steady stream to the titration flask while
swirling vigorously.
Observe the colour of the indicator carefully for some change in its shade or hue.
This is a sign of an approaching end-point
Stop the addition when you see the sudden colour change. Record this volume
reading which is a rough end point.
Accurate titration
Add the bulk of the rough titration volume (80-95%) rapidly, with swirling.
Rinse the walls of the flask with distilled water to ensure all reactive species are in
the mixing part of the flask
Complete the addition drop-wise so that the end point is overshot by no more than
one or two drops.
Repeat the titration procedure with fresh materials until you have an acceptable set
of end point volumes.
Replicate determinations in an ideal analysis will all lie within 0.05 mL. The
accumulated errors of the glassware used will produce a range of readings of at
least this magnitude. All novices should aspire to performance which approaches
these standards.
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Practical work 9.1
Practice Titration
Your teacher will demonstrate the practical technique of titration. The technique will
be the same for the remaining practicals. It would be wise to spend a good deal of
time practising before commencing on the remaining practicals.
Procedure:
1. Obtain about 150 mL of the standard 0.1 M acid and unknown alkali solution in
clean, dry labelled beakers. Cover the beakers with a watch-glass
2. Pour about 50 mL of distilled water into each of two 100 mL beakers
3. Add 5 mL of the acid to one beaker and 5 mL of the alkali to the other
4. Add sufficient screened methyl orange (SMO) to produce a colour. Note the
colour of the solutions on your worksheet
5. Mix 10 mL of each of these solutions in a third beaker. Adjust the colour to the
neutral colour by addition of small amounts of the acid or base as required.
6. Prepare a 50 mL burette with the standard acid as shown by the teacher
7. Fill and zero the burette with the acid solution
8. Pipette 25 mL of the unknown alkali into a 250 mL conical flask
9. Add the SMO and titrate as shown by the teacher to the neutral endpoint
10. Record your result
11. Repeat the analysis until you have three titrations within 0.1 mL
Practical work 9.2
Standardisation of approx. 0.1M HCl
1. Pour about 400 mL of distilled water into a 600 mL beaker and carefully, with
stirring,
2. Add the calculated amount of concentrated acid solution. Caution the acid is
corrosive
3. Pour the solution into a clean, labelled 1 L storage bottle and make up to 1L
4. Mix well
5. Weigh accurately, in triplicate, about 0.15g of dried AR grade sodium carbonate
(Na2CO3) into a weigh boat. Record the mass on the worksheet
6. Quantitatively transfer the weighed sodium carbonate into separate labelled
conical flasks.
7. Add SMO to each flask
8. Prepare a burette with your acid solution
9. Titrate until the endpoint is reached. Record the titration value on the worksheet.
10. Repeat for each of the samples
11. Calculate the value of your acid solution and label your bottle
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Practical work 9.3
Standardisation of approx 0.1 M NaOH
1. Weigh out about 4g of sodium hydroxide pellets on a top-pan balance and quickly
transfer the pellets to 1L of polished water. Caution: the sodium hydroxide is
caustic. Stir well
2. Transfer to a labelled storage bottle
3. Weigh out accurately, in triplicate, approximately 0.51g of potassium hydrogen
phthalate (labelled KHP in the desiccators) using an analytical balance. Record
the mass on the worksheet.
4. Quantitatively transfer the KHP to titration flasks
5. Prepare and fill a burette with the sodium hydroxide solution
6. Add phenolphthalein indicator (PP) and titrate to the baby pink end point. Record
the titration value on the worksheet.
7. Repeat for each flask
Cross-standardise your alkali solution with the acid solution you have
previously standardised.
12. Pipette 25 mL aliquots of your HCl into conical flasks
13. Add two drops of phenolphthalein indicator and titrate with the NaOH until the
pink just appears.
14. Record the value on your worksheet and repeat until you have three titrations
within 0.1 mL.
Practical work 9.4
Ethanoic acid content of vinegar
1. Pipette 25 mL of the concentrated vinegar into a 250mL volumetric flask and
make up to the mark with polished water. Shake well. Label
2. Pipette 25 mL aliquots of the diluted vinegar into titration flasks
3. Titrate the vinegar with standard NaOH, using phenolphthalein indicator
4. Repeat until you have three good titration values
Practical work 9.5
Sodium carbonate content of washing soda
1. Accurately weigh about 3.6g of washing soda crystals into a 150 mL beaker
2. Quantitatively transfer to a labelled 250 mL volumetric flask, dilute to the mark
and shake well
3. Titrate 25 mL aliquots of the dilute washing soda with standard acid solution
using screened methyl orange indicator.
4. Repeat until you have three good titration values.
Practical work 9.6
Acid content of commercial fruit juice
1. Prepare and fill a burette with standard 0.1M NaOH.
2. Pipette 25 mL of apple juice into a titration flask and dilute with approx. 75 mL
water.
3. Using phenolphthalein indicator, titrate until the first permanent pink appearance.
4. Repeat until you have three good titration values.
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Student Name:
Practical:
Practice titration
Practical Number: 9.1
Date Performed:
Text book References
Date Submitted:
Results:
1.
Observation of endpoint colour
Colour of solution
using screened M.O.
Colour of solution using
phenolphthalein
Colour of solution
Using M.O.
Acid HCI
Base NaOH
End point
2.
Practice Titration
Concentration of supplied acid
Volume of alkali pipetted into titration flask
Volume of acid needed for titration end point
1
2
3
4
Average volume acid used for end point
3.
Calculation of concentration of supplied alkali
Concentration alkali
=
[Concentration acid
x Volume acid] / Volume base
=
Questions:
1. What are the possible sources of error in the use of a burette?
2. Why should all readings be written down and checked before draining or refilling burettes?
3. Place each of the following pieces of glassware which may be used in a titration in the appropriate
preparation column.
Pipettes, burettes, volumetric flasks, titration conical flasks, beakers and bottles used
to hold standard and other solutions.
Glassware
Used dry
rinsed with solution to
be placed in it
rinsed with distilled
water.
Student Name:
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Practical:
Preparation & Standardisation of your own prepared
approximately 0.1M Hydrochloric acid
Practical Number: 9.2
Date Performed:
Text book References
Date Submitted:
Procedure:
Using the following formula C1V1 = C2V2 calculate the volume of 11.6 M HCl required to
prepare 1 Litre if 0.1M HCl. (10mL)
Show your working in this space:
Results
Standardising your prepared hydrochloric acid using sodium carbonate primary
standard
HCl vs Na2CO3
1
a
Mass Na2CO3 (g)
b
Moles of carbonate
(a ÷ 106)
c
End point Volume (mL)
d
Volume in L
(c ÷ 1000)
e
Moles of your acid
(b x 2)
Concentration of your acid
(e ÷ d)
2
3
4
Average value for molarity of your prepared HCl is
Questions:
1. Explain what is meant by the term quantitative transfer.
2. Why is the sodium carbonate kept in a desiccator?
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Student Name:
Practical:
Preparation & Standardisation of your own ≈ 0.1M NaOH
Practical Number: 9.3
Date Performed:
Date Submitted:
Text book References
Procedure:
Using the following formulas to calculate the mass of NaOH
(FW = 40) required to prepare 1 Litre if 0.1M NaOH. (4g)
[mass = FW x No. of moles] &
[no. of moles = volume (L) x molarity]
Show your working in this space:
Results:
(a)
Standardising your prepared sodium hydroxide using potassium hydrogen
phthalate (KHP) primary standard
1
a
Mass KHP (g)
b
Moles of KHP
2
3
4
(a ÷ 204)
c
End point Volume (mL)
d
Volume in L
(c ÷ 1000)
e
Moles of your alkali
(b x 1)
Concentration of your alkali
(e ÷ d)
Average value for the molarity of your NaOH is
(a)
Cross check your NaOH by titration against your standardised HCl.
This will give you a second answer for NaOH concentration
C2
V2
V1
C1
=
=
=
=
concentration of your HCl from last job
Volumes of HCl used at each end point
Volumes of NaOH used at each end point
Concentration of sodium hydroxide
=
=
=
=
Average value for the molarity of your NaOH is
Questions:
1. How well do the two answers for your prepared sodium hydroxide compare?
2. Why was the NaOH not weighed on an analytical balance?
3. Why are alkaline solutions stored in plastic rather than in glass?
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Student Name:
Practical:
Ethanoic acid content of Vinegar
Practical Number: 9.4
Date Performed:
Date Submitted:
Text book References
Procedure:
Results:
Volume of original vinegar taken =
Volume of diluted vinegar
=
Vol. of diluted vinegar titrated
=
Concentration of NaOH used
=
Titration volumes
(mL)
Average titration value
=
Calculations:
moles of NaOH used
=
=
Concentration NaOH x Average titration / 1000
moles ethanoic acid
=
=
moles of NaOH
Concn ethanoic acid
(diluted)
=
=
moles ethanoic acid / (volume of aliquot/1000)
Concn original ethanoic acid =
=
Concn of diluted vinegar x 10
Mass ethanoic acid per L
=
=
concentration ethanoic acid x 60
Mass ethanoic acid
per 100mL or % (w / v)
=
=
Mass ethanoic acid per L ÷ 10
Percentage of ethanoic acid in vinegar is
% (w/v)
Questions:
1. How would you titrate if brown vinegar had been used in place of white vinegar?
2. How does your answer for the vinegar compared to food authority standards for legal
vinegar?
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Student Name:
Practical:
Sodium Carbonate Content of Washing Soda
Practical Number: 9.5
Date Performed:
Date Submitted:
Text book References
Results:
Mass of washing soda: =
Size of volumetric flask: =
Aliquot volume
=
Concentration of HCl
=
Titration volumes
Average titration value
=
Calculations:
moles of HC1
=
concn of HCl x average titration volume / 1000
moles sodium carbonate
(in diluted aliquot)
=
moles of HCl / 2
molarity of sodium carbonate
(in analytical solution)
=
moles in diluted aliquot ÷ volume of diluted aliquot (L)
mass of sodium carbonate
in 250mL
=
FW x molarity sodium carbonate x 0.25
% sodium carbonate
=
mass sodium carbonate X 100
mass washing soda used
Average value for sodium carbonate content of washing soda is
% (w/w)
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Student Name:
Practical:
Acid content of Commercial Fruit Juice
Practical Number: 9.6
Date Performed:
Date Submitted:
Text book References
Results:
Concentration of NaOH:
=
Volume of juice taken
=
Titration volumes
Average titration value
=
Calculations:
moles of NaOH
=
Concentration NaOH x titration / 1000
moles of malic acid
in aliquot of juice
=
moles NaOH / 2
(only applies to apple juice)
mass of acid in 25 mL
=
moles acid x 134
(only applies to apple juice)
=
Mass of acid in 100 mL
=
% (w/v)
Percentage of malic acid in apple juice is
Questions:
If you were analysing blackcurrant juice would the endpoint be easy to observe?
How could you solve this problem?
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CALCULATION OF TITRATION RESULTS
The double outline box shows your answer all in one go.
The boxes underneath do the calculations step-wise.
TASK 1 STANDARDISE 0.1M HCl Formulae to Your results
use
Concentration of unknown HCl solution =
2 x (mass of Na2CO3 ÷ 106) ÷ (volume of HCl
(mL) ÷ 1000)
Moles of carbonate = mass of Na2CO3 ÷ 106
Moles of acid
= moles of carbonate x
2
Molarity of acid
= moles of acid ÷ (end point
volume (mL) ÷ 1000)
TASK 2
STANDARDISE 0.1M NaOH Formulae
to use
Concentration of unknown NaOH solution = volume
of HCl (mL) x concentration of HCl ÷ volume of
NaOH (mL) at end point
Your results
moles of HCl = volume of HCl (mL) x concentration
of HCl from task 1
moles of alkali = moles of HCl ÷ (end point volume
÷ 1000)
Molarity of NaOH = moles of alkali ÷ end point
volume ÷ 1000
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Intro Lab Methods
Conversion of mass of water to volume of water via its
density at different temperatures
temperature
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
density temperature
999.8395
999.8985
21
999.9399
22
999.9642
23
999.9720
24
999.9638
25
999.9402
26
999.9015
27
999.8482
28
999.7808
29
999.6996
30
999.6051
31
999.4974
32
999.3771
33
999.2444
34
999.0996
35
998.9430
36
998.7749
37
998.5956
38
998.4052
39
998.2041
40
density
997.9925
997.7705
997.5385
997.2965
997.0449
996.7837
996.5132
996.2335
995.9448
995.6473
995.3410
995.0262
994.7030
994.3715
994.0319
993.6842
993.3287
992.9653
992.5943
992.2158
Steps for calculating your true volume
1. Select the temperature which most closely matches your water's temperature
2. Find the density value in the table which corresponds to this temperature
3. For each of your measured masses of water, enter the numerical value for mass into your
calculator
4. divide this value by the density figure from the table
5. multiply by 1000 (to compensate for density being shown as kg / m3)
6. your display will now show your true volume for that mass of water.
7. Enter this value into you manual result sheet.
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