EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
Standardization of Hydrochloric Acid
by Anhydrous Sodium Carbonate
Procedures
Weigh about 1.0 - 1.5 g of anhydrous sodium carbonate powder accurately in a watch glass. Transfer the solid totally into a
250 cm3 beaker where about 50 cm3 distilled water is already filled. Wash the watch glass thoroughly by means of a washing
bottle and transfer all the washings into the 250 cm3 beaker. Add more water to dissolve the remaining solid. Use a glass
rod to stir the solution in order to facilitate the dissolving process.
Transfer the solution carefully to a 250 cm3 volumetric flask by means of a filter funnel and a glass rod. Rinse the beaker,
glass rod and inner surface of funnel with water and transfer all the washings to the volumetric flask. Repeat this process two
or three times. Make up the solution to 250 cm3 in the volumetric flask by adding water just up to the graduation mark.
Stopper the flask and invert it about 30 times to mix the solution thoroughly.
Pipette 25 cm3 of the sodium carbonate solution into a conical flask, add few drops of methyl orange indicator and titrate
against the hydrochloric acid solution. Determine the molarity of hydrochloric acid solution.
Colour change : yellow to reddish orange
Equation :
Na2CO3(aq)
+
2 HCl(aq)
→
2 NaCl(aq)
+
H2O(l)
+
CO2(g)
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EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
Standardization of Hydrochloric Acid
by Disodium Tetraborate (Borax)
Procedures
Weigh about 4.8 g of pure borax (disodium tetraborate) accurately. Dissolve it in distilled water and make up to 250 cm3 in a
volumetric flask. Pipette 25 cm3 of this solution into a conical flask, add a few drops of methyl orange solution and titrate
against the hydrochloric acid provided. Determine the molarity of hydrochloric acid solution. (Molecular mass of borax is
381 g mol-1)
Colour change : yellow to reddish orange
Equation :
Na2B4O7 ⋅ 10 H2O (aq)
+ 2 HCl(aq)
→
2 NaCl(aq)
+
4 H3BO3(aq)
+
5H2O(l)
Remark
Borax is not very soluble in cold water, implying that the dissolving process could be quite slow.
Questions
1.
What is a primary standard ?
2.
Decide which chemical, anhydrous sodium carbonate or borax, is a better primary standard.
Briefly explain your choice.
EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
Standardization of Potassium Permanganate
by an Iron(II) Salt
Procedures
Weigh about 9 - 10 g of ammonium iron(II) sulphate, Fe(NH4)2(SO4)2⋅6H2O, accurately in a watch glass. Transfer the solid
totally into a 250 cm3 beaker where about 100 cm3 dilute H2SO4 is already filled. Add water and stir the solution by a glass
rod to dissolve the solid. Make up the solution to 250 cm3 in the volumetric flask by adding water to the graduation mark.
Pipette 25 cm3 of the iron(II) solution into the conical flask and titrate against the potassium permanganate solution.
Determine the molarity of potassium permanganate solution. (Molecular mass of ammonium iron(II) sulphate is 392 g mol-1)
Colour change : colorless to pink
Equation :
5 Fe2+(aq) + MnO4-(aq) + 8 H+(aq)
→
5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)
Questions
1.
Give a balanced ionic equation if the titration is carried out in an alkaline medium.
(Hint : brown ppt. of MnO2 would be formed instead of colorless Mn2+)
2.
No indicator is employed in this titration.
3.
Explain why iron(II) sulphate, a more common iron(II) salt, is not used as the primary standard.
Explain briefly how the end-point could be determined.
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EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
Standardization of Sodium Thiosulphate
by Potassium Iodate(V)
Procedures
Weigh about 0.7 - 1.1 g of potassium iodate(V), KIO3. Dissolve it in distilled water and make up to 250 cm3 in the
volumetric flask. Pipette 25 cm3 of the iodate(V) solution into a conical flask, add 20 cm3 potassium iodide solution (KI),
10 cm3 dilute H2SO4 solution, 50 cm3 water and titrate against the sodium thiosulphate solution. When the solution changes
to pale yellow, add a few drops of starch solution as indicator and continue the titration until the solution is colorless.
Determine the molarity of sodium thiosulphate solution. (Relative atomic masses : K, 39.1 ; O, 16.0 ; I, 126.9 )
Colour change : reddish brown to pale yellow (before starch is added)
blue-black to colorless (after starch is added)
Equation :
IO3-(aq)
+
5 I-(aq)
I2(aq)
+
+
6 H+(aq) →
2-
2 S2O3 (aq) →
3 I2(aq)
-
2 I (aq)
+
+
3 H2O(l)
S4O62-(aq)
Questions
1.
Briefly explain the colour change throughout the whole course of the titration.
2.
In this experiment, iodine is indirectly generated (in situ) from an iodate(V) salt.
directly employed as the primary standard.
3.
Explain why starch indicator is not added at the beginning of the titration.
Give reason(s) why iodine is not
EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
Calcium Carbonate in Eggshells
Background
In the past, pesticides such as dichlorodiphenyltrichloroethane (DDT) have been used extensively. Their harmful effects on
biological systems are gradually revealed. One effect on wild bird life is the weakening or thinning of the eggshell. The
consequence is the breaking of the eggshell before hatching. Eggshell mainly consists of calcium carbonate. In the present
investigation, we will determine the percentage of calcium carbonate in eggshell by acid/base titration.
Tasks
1.
Design and draw a flowchart to show how you will perform the experiment with the materials provided.
2.
Consult your teacher for the feasibility of your design.
3.
Write down the step-by-step procedure including the precise weight and volume of the reagents used.
4.
Perform the experiment.
5.
Record the observations and collect all the necessary data for calculation.
6.
Prepare a report according to your teacher's instruction.
Questions for Further Thought
1.
HCl and NaOH solutions are not primary standards and need to be standardized. Suggest a primary standard for the
present experiment and explain briefly how to perform the standardization.
2.
Write down the equation for the reaction involved before titration in the present experiment.
3.
Find out the chemical formula and structure of DDT.
4.
Find another toxic chemical which is harmful to the wildlife species. Briefly describe its effect(s) and application(s).
References
1.
http://chem.lapeer.org/Chem1Docs/EggshellTitration.html
2.
http://www.accessexcellence.org/AE/AEC/AEF/1996/tucker_eggshell.html
3.
I. Newton and I. Wyllie, J. Appl. Ecology, 1992, 29, 476.
EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
To Determine the Heat of Neutralization between
a Weak Acid and a Strong Base
You are provided with solid succinic acid [(CH2)2(CO2H)2] , which is a dibasic acid, a solution N of acetic acid and a solution
O of potassium hydroxide. The molarity of solution N may be taken as equal to that of solution O.
You are required to determine:
(a) the molarity of the potassium hydroxide solution O by simple titration, and
(b) the heat of neutralization of acetic acid with potassium hydroxide.
Procedure
(1) Prepare a solution containing between 1.10 g and 1.30 g of succinic acid in 250 cm3. By means of a pipette, transfer
exactly 25 cm3 of the potassium hydroxide solution O into another 250 cm3 volumetric flask and dilute to the mark.
Titrate this diluted potassium hydroxide solution against 25 cm3 succinic acid solution, using phenolphthalein as indicator,
until concordant results are obtained.
(2) Calculate (i) the molarity of the succinic acid solution, and
(ii) the molarity of the original potassium hydroxide solution O.
Second weighing……………=
g
First weighing……………….=
g
Mass of succinic acid taken….=
g
3
Titration of the diluted potassium hydroxide solution against 25 cm of succinic acid
Burette reading (finish)
Burette reading (start)
Volume of solution used
∴_________ cm3 of diluted potassium hydroxide solution ≡ 25 cm3 succinic acid
Calculation
(3) Using measuring cylinders, carefully transfer 50 cm3 of the acetic acid solution N into one plastic cup and 50 cm3 of the
potassium hydroxide solution O into the other. Set aside for about 15 minutes. Place the thermometer provided in the
cup containing solution N and record the temperature at ½-minute intervals, stirring continuously. At exactly 3 ½
minutes pour the entire 50 cm3 of solution O into the stirred solution N and continue to take temperature readings as
before from minute 4 to minute 7.
time/min
½
1
1½
2
2½
3
3½
4
4½
5
5½
6
6½
7
7½
temp/ °C
Plot a graph of temperature against time on the graph paper below (temperature on the vertical axis, time on the horizontal
axis using the scale of 2 cm to 1 minute.
(4) From the graph deduce the maximum rise in temperature. Hence calculate the heat of neutralization of acetic acid with
potassium hydroxide. (It may be assumed that the heat capacity of the final solution is 4.18 J cm3 K-1 and that the heat
capacities of the thermometer and the cup may be neglected.)
(5) By comparing the answer in (Q.4) with the standard heat of neutralization between HCl and NaOH, which is
-57.3 kJ mol-1, comment on any difference in these two values.
EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
To Determine The Heat Of Formation
Of A Metal Carbonate
You are provided with a metal Z, the carbonate of the metal, ZCO3 and 2.0 M hydrochloric acid.
(a) find the heat of reaction between the metal Z and hydrochloric acid, ∆H1,
(b) find the heat of reaction between the carbonate ZCO3 , and hydrochloric acid, ∆H2,
(c) calculate the heat of formation of the carbonate ZCO3, ∆H3.
You are required to :
The following data will be needed:
H2 (g) + ½ O2 (g) → H2O(l)
∆H4 = -286 kJ mol-1
∆H5 = -393 kJ mol-1
C (s) + O2 (g) → CO2(g)
Previous experience of this type of experiment is not required; full details of the method and calculations are given below.
Procedure
(1) Weigh out accurately between 0.090 g and 0.110 g of the metal Z and record the mass below. By means of a burette or
pipette, place 50.0 cm3 of the hydrochloric acid in a plastic cup. Stirring with the thermometer provided, measure the
temperature of the acid every half minute to ±0.2 °C. Record your results below. At exactly 3 ½ minutes add the metal
to the acid and continue taking the temperature of the liquid every half minute from minute 4 to minute 7. Stir the liquid
throughout.
Mass of metal Z =
g
time/min
½
1
1½
2
2½
3
3½
4
4½
5
5½
6
6½
7
7½
temp/ °C
(2) Plot a graph of temperature (vertical axis) against time (horizontal axis; scale 2 cm to 1 minute) and from this find the
temperature rise (∆T1) for the reaction.
°C
∆T1 =
(3) In the space below, calculate the heat of reaction between the metal Z and hydrochloric acid. (∆H1)
(Molecular mass of Z is 40.1; density of water is 1 g cm-3; specific heat capacity of water is 4.18 J g-1K-1)
(4) Weigh accurately a weighing bottle containing between 2.00 g and 2.20 g of the carbonate ZCO3. Record the mass
under Experiment 1 in the table below. Using a burette or a pipette, place 50.0 cm3 of the hydrochloric acid in the plastic
cup and record its temperature to ±0.2 °C. Carefully add the carbonate to the acid. Stirring with the thermometer,
measure the highest temperature attained to ±0.2 °C. Re-weigh the weighing bottle after addition of the carbonate.
Record your results below as Experiment 1.
Repeat this procedure as Experiment 2.
Experiment 1
Experiment 2
Mass of bottle and carbonate /g
Mass of bottle /g
Mass of carbonate used /g
Highest temperature of acid /°C
Initial temperature of acid /°C
Rise in temperature, ∆T2 /°C
(5) In the space below calculate the heat of reaction between the carbonate ZCO3 and hydrochloric acid (∆H2) using the
expression given.
Experiment 1
Experiment 2
∆ H2 =
∆ H2 =
=
=
Mean value of ∆H2 =
_____________
(6) Using the above results, calculate the heat of formation of the carbonate ZCO3, ∆H3.
EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
To Determine The Heat Of Solution
Of Sodium Thiosulphate
The heat of solution ∆Hsoln, of a solute is the enthalpy change that occurs when one mole of the solute dissolves to form an
infinite dilute solution. In this part of the practical you are to devise and carry out a simple experiment to obtain as accurate a
value as possible for ∆Hsoln for hydrated sodium thiosulphate, Na2 S2O3 · 5 H2O. In principle this can be done by measuring
the temperature change of the water in which the sodium thiosulphate dissolves, but you should bear the followings in mind
when designing your experiment.
A. Heat loss must be minimized. In particular, you will need good insulation and the sodium thiosulphate should be
dissolved as quickly as possible. You may be able to think of other ways of reducing heat loss.
B. It will not, of course, be possible to produce an infinitely dilute solution, and your solution will need to be concentrated
enough to produce as accurately measurable temperature rise. You should think carefully about the best concentration to
use.
C. You can assume that the specific heating capacity of sodium thiosulphate solution is the same as that of water,
i.e. 4.18 J g-1 K-1.
D. You will need to bear in mind the apparatus available to you.
1.
Write a brief description of your experiment and tabulate the results.
2.
Work out a value for the enthalpy change of solution of sodium thiosulphate as follows:
a. Use the temperature change of the solution to calculate the enthalpy change that would have occurred in the solution.
b. Knowing the amount of sodium thiosulphate used. Calculate the enthalpy change that would have occurred if one
mole of sodium thiosulphate had been used.
3.
What do you consider to be the major sources of error in your experiment?
4.
How might these errors have been reduced if more sophisticated apparatus had been available?
Project Work: Designing a Coolant System
Campers and others without access to a refrigerator can keep food and drinks cool by lowering the temperature of an insulated
container using chemical means. A convenient method is to add a salt with an endothermic enthalpy change of solution to
water in a plastic bag within the insulated container.
Devise a system that could be used to keep food cool. Here are some points that you might like to think about:
• What salts can be used? Cost and availability need to be borne in mind, as well as suitability of heat of solution.
• What temperature drop is required?
• What quantities of salt and water are needed to achieve the required temperature drop?
• Does the system work in practice?
Discuss your plans with your teacher before starting any practical work.
EXPERIMENT
St. Clare's Girls' School
F.6 Chemistry
DATE :
Determination Of Vitamin C Content
In Processed Fruit Drinks
Introduction
Vitamin C (ascorbic acid) is often added to canned or packaged fruit drinks to enhance their nutritional value. In this
experiment, you are asked to propose a titrimetric method to determine the vitamin C content of a sample of fruit drink
purchased from supermarket.
Tasks
1.
Search for information on methods of determining vitamin C content in fruit drinks. You may use search engines (e.g.
Yahoo or Google), textbooks or reference books.
(a) What are the chemical principles underlying the various methods of determining vitamin C?
(b) Select one method which involves reaction with iodine.
2.
With the possible methods in mind, choose the one which you think is the best and design step-by-step procedures to
investigate quantitatively the sample of fruit drink with regard to its vitamin C content. You may consider the
following when designing your experiments.
(a) What chemicals and apparatus would you use?
(b) What is the chemical structure of vitamin C?
(c) Explain whether vitamin C is stable.
(d) What is the chemical reaction(s) involved in the titration?
(e) Outline calculation steps based on equation stoichiometry.
(f)
Compare experimental result with the vitamin C content indicated on the ingredient label of the package.
3.
For the second part of the investigation, design an experiment to investigate the effect of temperature on the content
of vitamin C in fruit juice.
4.
Before carrying out your investigation, seek advice from your teacher on the details of your proposed experiment.
5.
Carry out your investigation and record all observations.
6.
Write up a report for the investigation. Your report should include:
(a) the underlying chemical principle of your method,
(b) procedures,
(c) tables summarizing titration results,
(d) treatment of data involving calculations,
(e) a conclusion of the investigation and
(f)
comment on the results obtained.
7.
Does your result agree with the figures indicated on the ingredient label? What do you think are the reasons? Try to
design further experiments to find out.