Aims and Objectives of the Experiment
The main objective of this experiment is to determine the concentration of different substances. This is achieved by determining the end point(s) of different titrations using the automated potentiometric titrator. These points are known as equivalence points. The units of measurement are the standard units used in titrations i.e. Molarity (moles/liter) of substances used while the quantities of these substances are measured in mL (milliliter).
Theory and Principles
Potentiometric titration is a volumetric method which measures the potential between two electrodes, namely reference and indicator, with the change in the volume of the reagent that is added, known as the titrant. There are four types of potentiometric titrations including acid-base, redox, precipitation, and complexometric (Chemistry Glossary, 2015).
Potentiometric titrations are preferable as compared to manual titrations because of the added accuracy and precision. These types of titrations are easier to automate as compared to others. The advantage that automated titration systems offer over manual systems is that they are capable of processing greater volumes of samples and the need for the involvement of the one performing the experiment is close to zero.
A titration curve is plotted to study the characteristics of the titration that was carried out. The equivalence point of the titration is marked by the maximum change in the curve. The slope of the curve is given by the derivative, and the endpoint is marked by the volume at which the first derivative has the maximum value.
The titration curve is obtained by plotting the values of the potential on y-axis and the corresponding values, of the volume of titrant, are plotted on the x-axis forming an S-shaped curve. The central part of the plot, in which the curve rises sharply corresponds to the volume at which the end point of the titration occurs. In cases where the potential change has a small value at the end point, for instance in the titration of weak acid with strong base, the location of the end point becomes problematic (SPECTRUM, 2005).
A reference electrode is an electrode with a known potential. This potential remains constant at constant temperature and is independent of the analyte’s composition. The reference electrode is considered to be the left-hand electrode in potentiometric measurements. The types of reference electrodes include Calomel electrodes and silver/silver chloride (Middle East Technical University, n.d.).
The electrode, the potential of which varies with variations in the analyte’s concentration, is known as the indicator electrode. Most of the indicator electrodes used for potentiometric titrations have selective responses. The types of indicator electrodes include metallic indicator electrode and membrane electrodes.
Instrument Used for the Experiment
The instrument used for this experiment was the Mettler Toledo DL 50 Titrator which is shown in Figure 1. The instrument is designated to comprise of a titration stand or in some cases a vessel, the holding stand for interchangeable burettes, a stirrer is used for the titrator vessel which is of propeller type, and the platinum electrodes require an electrode (Mettler Toledo, n.d.). pH and voltage sensors are integrated into the system to support the experiment. The system provides option for the calibration and the pH and voltage sensors may be calibrated using the system. The system provides an option to directly print the results or they may be sent to a computer to print at a later point in time. The system also provides with other options. These include the options for setting the date and time for the experiment, assigning names to the sample/substances or the experiment. The system has its own memory where this data may be saved.
Figure 1: Mettler Toledo DL 50 Titrator
Advantages and Disadvantages of the Potentiometric Method
The potentiometric method offers several advantages over the conventional methods. The potentiometric method can be used for analyte solutions that are colored, turbid or fluorescent. It may also be used in situations where there is no suitable indicator available or the change in color is not evident and tough to recognize. It is used for very dilute solutions. It is also known to be used in titrations of polyprotic acid, mixtures of acids, mixtures of bases, or mixtures of halides (SPECTRUM, 2005).
There are certain disadvantages associated with potentiometry which come along with all the advantages that it has to offer. The disadvantages include the need for frequent calibration, the sensitivity of the method to ionic changes, and others (Woods Hole Oceanographic Institute, n.d.).
Experimental Procedure
First of all, the pH electrode needs to be calibrated using the auto-titrator and the computer running the Lab X software. The second step is to rinse the burette and the dispensing tube with the appropriate titrant. The titrations are carried out as explained in the lab manual, changing the parameters as required using the appropriate method, titrant, and electrode. Once the titrations are complete, the electrodes should be stored in appropriate solutions after rinsing. The experiments that need to be carried out using the above outlined steps include a number of titrations using acids and bases of different concentrations (“Electrochemical Titrations,” n.d.).
Results and Calculations
The calculation of results is completed in the following way.
=
(1) 10 ml NaOH with 0.1 M HCl (1 end point)
NaOH + HCl → KCl + H2O + CO2
The molar ratio between HCl and NaOH is 1:1. Titration of 10 ml 0.1 M NaOH solution requires 9.587 ml of 0.1 M HCl, then the calculation of the concentration of NaOH is possible.
Moles HCl = Moles NaOH
M NaOH × 10ml / 1 = M HCl × R1 / 1
M NaOH × 10 / 1 = 0.1 × 9.374 / 1
M NaOH = 0.09374 M
(2) 10ml KHCO3 with 0.1 M HCl (1 end point)
KHCO3 + HCl → KCl + H2O + CO2
M KHCO3 × 10 ml / 1 = M HCl × R1 / 1
M KHCO3× 10 /1 =0.1 × 10.008 /1
M KHCO3 = 0.10008 M
(3) 10 ml Na2CO3 with 0.1 M HCl (2 end point)
1st step: Na2CO3 + HCl → NaHCo3 + NaCl
2nd step: NaHCO3 + HCl → NaCl + H2O + CO2
Total: Na2CO3 + HCl →2 NaCl + H2O + CO2
M NaHCO3 × 10 ml / 1 = M HCl × R2 / 2
M Na2CO3 × 10 /1 = 0.1 × 8.370/ 2
M Na2CO3= 0.04185 M
(4) 10 ml Na2CO3 + 5 ml KHCO3 with 0.1 M HCl (2 end point)
1st step: Na2CO3 + HCl → NaHCO3 + NaCl
M Na2CO3 × 10 ml / 1 = M HCl × R1 /1
M Na2CO3 × 10/1 = 0.1 × 5.319 / 1
M Na2CO3= 0.05319 M
2nd step: NaHCO3 + HCl →NaCl + H2O + CO2
KHCO3 +HCl→ KCl + H2O + CO2
V HCl for 2nd step = R2- R1
V HCl for NaHCO3 = R1
Therefore, V HCl for KHCO3 = (R2-R1) –R1 =R2 -2R1
R2 -2R1 = 13.428 – 2(5.319) = 2.79
M KHCO3 × 5ml / 1 = M HCl × (R2 -2R1) /1
M KHCO3 × 5 /1 = 0.1 × 2.79/ 1
M KHCO3 = 0.0558 M
(5) 10 ml HCl with 0.1 M NaOH (1 end point)
NaOH + HCl→ NaCl + H2O
HCl ×10 ml / 1 = M NaOH × R1 / 1
M HCl ×10 ×10 /1 = 0.1 × 10.525/ 1
M HCl = 0.10525 M
(6) 10 ml H3BO3 with 0.1 NaOH (1 end point)
H3BO3 + NaOH→ H2O + Na3BO3
H3BO3 + 3 NaOH = 3 H2O + Na3BO3
M H3BO3 × 10 ml / 1 = M NaOH × R1 /1
M H3BO3 × 10/1 = (0.1 × 10.167)/ 3
1.0167 ml/M = 30x (ml)
X = 0.0339
X= the molarity of H3BO3 = 0.0339 M
(7) 5ml H3PO4 with 0.1 M NaOH (2 end points)
1st Step: H3PO4 + NaOH→H2O + Na3PO4
H3PO4 + 3NaOH → Na3PO4 + 3H2O
M (H3PO4)(5ml)1=M(NaOH)(R2)3
M H3PO4 × 5/1 = (0.1 × 10.849)/ 3
1.0849 ml/M = 15x (ml)
X= 0.0723 M
X= the molarity of H3PO4 = 0.0723 M
Discussion
It is pretty much understood that strong acids dissociate completely in water, whereas weak acids seldom dissociate completely. Therefore, the extent of dissociation of acids in water may be calculated using the equilibrium constant for the respective reactions and the volumes of acids and bases added to the solution.
The reactions and their results are arranged in Results and Calculations section. The resulting concentrations are pretty close to the expected values of concentration. It may be observed that, the molarities of the titrants are almost equal to the values of molarity of the substances which were expected.
Some of the reactions were completed in single steps while others took two steps to proceed to completion. In such cases where there are two end points, only the closest one to the equilibrium will be considered. This is due to the fact that the substances containing CO32- are stronger bases as compared to those containing HCO3- and all the carbonate will eventually be converted to bicarbonate.
Conclusion
The use of potentiometric titrations to determine the concentrations of substances by determining the equivalence points. This method offers few benefits along with few disadvantages as well. The use of this technique is favorable in certain situations.
References
Chemistry Glossary. (2015). Potentiometric titration. Retrieved January 24, 2017, from http://glossary.periodni.com/glossary.php?en=potentiometric+titration
Electrochemical Titrations. (n.d.), 1–8 [Lab Manual].
Mettler Toledo. (n.d.). Reference Handbook - DL50/DL53/DL55/DL58 Titrators. mt.com. Retrieved from http://www.mt.com/dam/mt_ext_files/Editorial/Generic/1/Reference_handbook_DL5X_Editorial-Generic_1166537021939_files/51709614_en.pdf
Middle East Technical University. (n.d.). Potentiometric Titrations.
SPECTRUM. (2005). Potentiometry in Analytical Chemistry. Central Department of Chemistry, Tribhuvan University. Retrieved from http://www.nepachemistry.com/2011/03/potentiometry-in-analytical-chemistry.html
Woods Hole Oceanographic Institute. (n.d.). Advantages and disadvantages of current analytical approaches. Retrieved January 25, 2017, from http://www.whoi.edu/fileserver.do?id=54005&pt=2&p=58666
