TITRATIONS
09104040 Payal Niharika 09104047 Rohit Madaan 09104049 Sandeep Middye
ACKNOWLEDGEMENT:
This project was one great opportunity to learn and experience the beauty of practical chemistry. It gave us a deeper insight into the science of titration, its types and mechanism. We would like to express our gratitude for Mrs. Sarla Kumari, Associate Professor, Applied Sciences Department, who guided us throughout the project. We would also like to thank PEC University of Technology for providing us the facilities required to complete the project. Last but not the least; we are very grateful to Almighty God due to whose grace the project stands as it is today.
S.No
Topic
Page No.
1.
ACKNOWLEDGEMENT
2
2.
INTRODUCTION
4
3.
HISTORY
5
4.
PREPARATION OF SAMPLE
5
5.
PROCEDURE
6
6.
GENERAL CALCULATIONS
6
7.
TITRATION CURVES
7
8.
TYPES OF TITRATION
8
9.
BACK TITRATION
16
10.
MEASURING THE ENDPOINT
16
10.
APPLICATIONS OF TITRATIONS
18
11.
BIBLIOGRAPHY
20
INTRODUCTION The word "titration" comes from the Latin word titulus , meaning inscription or title. The French word titre , also from this origin, means rank. Titration, by definition, is the determination of rank or concentration of a solution with respect to water with a pH of 7 (which is the pH of pure H 2O under standard conditions). Titration is a common laboratory method of quantitative chemical analysis that is used to determine the unknown concentration of a known reactant. Because volume measurements play a key role in titration, it is also known as volumetric analysis. A reagent, called the titrant or titrator, of a known concentration (a standard solution) and volume is used to react with a solution of the analyte or titrand, whose concentration is not known. Using a calibrated burette or chemistry pipetting syringe to add the titrant, it is possible to determine the exact amount that has been consumed when the endpoint is reached. The endpoint is the point at which the titration is complete, as determined by an indicator. This is ideally the same volume as the equivalence point — the volume of added titrant at which the number of moles of titrant is equal to the number of moles of analyte, or some multiple thereof (as in polyprotic acids). Many methods can be used to indicate the endpoint of a reaction; titrations often use visual indicators (the reactant mixture changes color). Not every titration requires an indicator. In some cases, either the reactants or the products are strongly colored and can serve as the "indicator". For example, a redox titration using potassium permanganate (pink/purple) as the titrant does not require an indicator. Due to the logarithmic nature of the pH curve, the transitions are, in general, extremely sharp; and, thus, a single drop of titrant just before the endpoint can change the pH significantly — leading to an immediate color change in the indicator. There is a slight difference between the change in indicator color and the actual equivalence point of the titration. This error is referred to as an indicator error , and it is indeterminate.
HISTORY: The origins of volumetric analysis are in late-18th-century French chemistry. Francois Antoine Henri Descroizilles developed the first burette (which looked more like a graduated cylinder) in 1791. Joseph Louis Gay-Lussac developed an improved version of the burette that included a side arm, and coined the terms "pipette" and "burette" in an 1824 paper on the standardization of indigo solutions. However, the major breakthrough in the methodology and popularization of volumetric analysis was due to Karl Friedrich Mohr, acoording to him titration was the "weighing without scale" method, because this process allows determination of the concentration of a sample without using complex instrumentation. He redesigned the burette by placing a clamp and a tip at the bottom, and wrote the first textbook on the topic, Lehrbuch der chemisch-analytischen Titrirmethode (Textbook of analytical-chemical titration methods ), published in 1855.
Preparing a sample for titration In a titration, both titrant and analyte are required to be in a liquid (solution) form. If the sample is not a liquid or solution, the samples must be dissolved. If the analyte is very concentrated in the sample, it might be useful to dilute the sample. Although the vast majority of titrations are carried out in aqueous solution, other solvents such as glacial acetic acid or ethanol (in petrochemistry) are used for special purposes. A measured amount of the sample can be given in the flask and then be dissolved or diluted. The mathematical result of the titration can be calculated directly with the measured amount. Sometimes the sample is dissolved or diluted beforehand, and a measured amount of the solution is used for titration. In this case the dissolving or diluting must be done accurately with a known coefficient because the mathematical result of the titration must be multiplied with this factor.
Burette
Pipette
Many titrations require buffering to maintain a certain pH for the reaction. Therefore, buffer solutions (A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. It has the property that the pH of the solution changes very little when a small amount of strong acid or base is added to it) are added to the reactant solution in the flask to maintain the pH of the solution. Some titrations require "masking" of a certain ion. This can be necessary when two reactants in the sample would react with the titrant and only one of them must be analysed, or when the reaction would be disturbed or inhibited by this ion. In this case another solution is added to the sample, which "masks" the unwanted ion (for instance by a weak binding with it or even forming a solid insoluble substance with it). Some redox reactions may require heating the solution with the sample and titration while the solution is still hot, in order to increase the reaction rate. For instance, the oxidation of certain oxalate solutions requires heating the solution to approximately 60 degrees in order to maintain a reasonable rate of reaction.
PROCEDURE A typical titration begins with a beaker or Erlenmeyer flask (An Erlenmeyer, also known as a conical flask, is a widely used type of laboratory flask which features a flat bottom, a conical body, and a cylindrical neck. It is named after the German chemist Emil Erlenmeyer, who created it in 1861) containing a precise volume of the reactant and a small amount of indicator, placed underneath a burette or buretting syringe containing the reagent. By controlling the amount of reagent added to the reactant, it is possible to detect the point at which the indicator changes color. As long as the indicator has been chosen correctly, this should also be the point where the reactant and reagent neutralize each other, and, by reading the scale on the burette, the volume of reagent can be measured.
Erlenmeyer flask
General Calculations: Using the normality (the normality of a solution is defined as the molar concentration divided by an equivalence factor) equation, we can calculate the strength of the analyte: N1V1=N2V2 Where, N1=normality of the titrant V1=volume of the titrant used N2=normality of the analyte V2=volume of the analyte taken Strength= N2* molecular weight of analyte Usually, the volume is expressed in terms of milliliters (ml) and normality in terms of mol/l or eq/l. The strength is found in terms of gm/l or mg/l.
TITRATION CURVES A titration curve is a curve in the plane whose x -coordinate is the volume of titrant added since the beginning of the titration, and whose y -coordinate is the concentration of the analyte at the corresponding stage of the titration (in an acid-base titration, the y -coordinate is usually the pH of the solution at the corresponding stage, the volume of the analyte is an independent variable). Often it is the case that the titration curve of a titration reflects the nature of the titration quite well; for instance, it reflects the nature of all solutions involved in the titration. In the case of acid-base titrations, titration curves reflect the strength of the corresponding acid and base. For instance, in a strong acid and strong base titration, the titration curve will be relatively smooth, although very steep for points near the equivalence point of the titration. Since in this case, small changes in the volume of the titrant result in large changes of the pH near the equivalence point, an extensive range of indicators would be appropriate (for instance litmus, phenolphthalein or bromothymol blue). On the other hand, if one of the constituents of an acid-base titration is either a weak acid or a weak base, and the other is either a strong acid or a strong base, the titration curve is fairly irregular near the equivalence point (and the pH does not change as much due to the addition of small volumes of titrant). For instance, the titration curve for the titration between oxalic acid (a weak acid) and sodium hydroxide (a strong base) where, the equivalence point occurs at a pH of about 8-10, and thus the analyte is basic at the
equivalence point (more precisely, the sodium salt produced by the reaction hydrolyses in water to produce hydroxide ions). An indicator such as phenolphthalein would be appropriate for this particular titration. The titration curve corresponding to a weak base and strong acid titration is similarly behaved. In this case, indicators such as methyl orange or bromothymol blue are regularly used. On the other hand, titration curves corresponding to acid-base titrations in which the constituents are a weak acid and weak base are quite irregular in nature. Due to the nature of such titrations, no definite indicator may be appropriate, and thus pH meters are often used.
Diagram showing titration curves for acid-base titrations
Types of titration
TITRATION Acid-base
Redox
Complexometric
Zeta-potential
Acid-Base titrations: An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the acid/base with an acid or base of known concentration. This allows for quantitative analysis of the concentration of an unknown acid or base solution. It makes use of the neutralization reaction that occurs between acids and bases and the knowledge of how acids and bases will react if their formulas are known. Acid-base titrations can also be used to find percent purity of chemicals. EQUIPMENT
The key equipments used in a titration are:
Burette White tile - used to see a colour change in the solution Pipette pH indicator (the one used varies depending on the reactants) Erlenmeyer flask/ Conical flask Titrant or titrator (a standard solution of known concentration, a common one is aqueous sodium carbonate) Analyte or titrand (solution of unknown concentration)
METHOD
Before starting the titration a suitable pH indicator must be chosen. The equivalence point of the reaction, the point at which equivalent amounts of the reactants have reacted, will have a pH dependent on the relative strengths of the acid and base used. The pH of the equivalence point can be estimated using the following rules: A strong acid will react with a strong base to form a neutral (pH=7) solution. A strong acid will react with a weak base to form an acidic (pH<7) solution. A weak acid will react with a strong base to form a basic (pH>7) solution. When a weak acid reacts with a weak base, the equivalence point solution will be basic if the base is stronger and acidic if the acid is stronger. If both are of equal strength, then the equivalence pH will be neutral. However, weak acids are not often titrated against weak bases because the colour change shown with the indicator is often quick, and therefore very difficult for the observer to see the change of colour.
The point at which the indicator changes colour is called the end point. A suitable indicator should be chosen, preferably one that will experience a change in colour (an end point) close to the equivalence point of the reaction. First, the burette should be rinsed with the standard solution, the pipette with the unknown solution, and the conical flask with distilled water. Secondly, a known volume of the unknown concentration solution should be taken with the pipette and placed into the conical flask, along with a small amount of the indicator chosen. The burette should always be filled to the top of its scale with the known solution for ease of reading. The known solution should then be allowed out of the burette, into the conical flask. At this stage we want a rough estimate of the amount of this solution it took to neutralize the unknown solution. The solution should be let out of the burette until the indicator changes colour and the value on the burette should be recorded. This is the first (or rough) titre and should be discluded from any calculations. Three more titrations should be performed, this time more accurately, taking into account roughly where the end point will occur. The readings on the burette at the end point should be recorded, and averaged to give a final result. The end point is reached when the indicator just changes colour permanently. This is best achieved by washing a hanging drop from the tip of the burette into the flask right at the end of the titration to achieve a drop that is smaller in volume than what can usually be achieved by just dripping solution off the burette. Acid-base titration is performed with a phenolphthalein indicator, when it is a strong acid strong base titration, a bromothymol blue indicator in weak acid - weak base reactions, and amethyl orange indicator for strong acid - weak base reactions. If the base is off the scale, i.e. a pH of >13.5, and the acid has a pH >5.5, then an Alizarine yellow indicator may be used. On the other hand, if the acid is off the scale, i.e. a pH of <0.5, and the base has a pH <8.5, then a Thymol Blue indicator ( transitions from red to yellow at pH 1.2 – 2.8 and from yellow to blue at pH 8.0 – 9.6) may be used.
Indicator
Color on Acidic Side
Range of Color Change
Color on Basic Side
Methyl Violet
Yellow
0.0 - 1.6
Violet
Bromophenol Blue
Yellow
3.0 - 4.6
Blue
Methyl Orange
Red
3.1 - 4.4
Yellow
Methyl Red
Red
4.4 - 6.2
Yellow
Litmus
Red
5.0 - 8.0
Blue
Bromothymol Blue
Yellow
6.0 - 7.6
Blue
Phenolphthalein
Colorless
8.3 - 10.0
Pink
Alizarin Yellow
Yellow
10.1 - 12.0
Red
REDOX TITRATION Redox titration (also called oxidation-reduction titration) is a type of titration based on a redox reaction between the analyte and titrant. These titrations are based on a redox reaction between an oxidizing agent and a reducing agent. The oxidizing agent (resp. reducing agent) is added to the burette, which was rinsed with the same oxidizing agent. The reducing agent (resp. oxidizing agent) is added to the conical flask, which had been rinsed with distilled water. Like in an acid-base titration, the standard solution is often the one in the conical flask, and the solution whose concentration is to be determined is the one in the burette. The procedure for carrying out redox titrations is similar to that required for carrying out acid-base titrations. Most commonly, a potentiometer or a redox indicator is used to determine the end-point of the titration. For example, when one of constituents of the titration is the oxidizing agent potassium dichromate, the colour change of the solution from orange to green is not definite and, thus, an indicator such as sodium diphenylamine is used. The analysis of wines for their sulfur dioxide content requires the use of iodine as an oxidizing agent. In this case, starch is used as an indicator; a blue starch-iodine complex is formed once an excess of iodine is present, thus signaling the endpoint of the titration. On the other hand, some redox titrations do not require an indicator, due to the intense colour of some of the constituents. For instance, in a titration where the oxidizing agent potassium permanganate (permanganometry) is present, a slight faint persisting pink colour signals the endpoint of the titration, and no particular indicator is, therefore, required. Some of the commonly used oxidizing and reducing agents in the redox titrations are – i.
KMnO4 in presence of dil H 2SO4 -
-
+
2+
MnO4 + 8H + 5e → Mn + 4H2O E°red = +1.52V ii. K2Cr2O7 in dil. H2SO4 is a moderately strong oxidizing agent; oxidizing ability depends strongly on pH, decreasing rapidly as solution becomes more neutral 2-
+
-
Cr2O7 + 14H + 6e iii. Iodine solution -
3+ → 2Cr +
7H2O
–
I2 + 2 e = 2I
i. Mohr’s salt FeSO4.(NH4) 2SO4.6H2O
E°red = +1.33V E°red = +0.54V
Fe2+ → Fe3+ + eii. Oxalic acid H2C2O4.2H2O C2O42- = 2CO2 + 2 eiii. Sodium thiosulphate Na 2S2O3.5H2O 2S2O32S4O62- + 2 e-
E°red = +0.77V E°red = +0.77V E°red = +0.08V
There is no universal oxidizing agent which can be titrated against every reducing agent and vice-versa. Hence, the choice of an oxidizing agent to be used against a particular reducing agent depends upon the reaction conditions and standard reduction potential of the oxidizing agent.
TYPES OF REDOX INDICATORS - Many a times the titrant itself may be so strongly coloured that after the equivalence point, a single drop of the titrant produces an intense colour in the reaction mixture. e.g. potassium permanganate. Such Indicators are called self indicators. Self Indicators generally are strongly coloured as a result of charge transfer transitions in them. - Such indicators are added into the reaction mixtures Such indicators always have reduction potential values lower than the analyte system so that they react with the titrant only when whole of the analyte has been consumed, producing a readily detectable color change. -In case a suitable redox indicator is not available for a given system, an indicator may be employed which will indicate the completion of reaction by physically or chemically reacting with the analyte (not through redox reaction). This reaction between indicator and the analyte may sometimes be an irreversible one and in some cases may even lead to precipitation. In those case indicators are not added to the reaction mixture on the whole, rather used externally on a grooved tile. Such indicators are called external indicators.
COMPLEXOMETRIC TITRATION Complexometric titration (sometimes chelatometry) is a form of volumetric analysis in which the formation of a colored complex is used to indicate the end point of a titration. Complexometric titrations are particularly useful for the determination of a mixture of different metal ions in solution. These titrations are based on the formation of a complex between the analyte and the titrant. The chelating agent EDTA is very commonly used to titrate metal ions in solution. An indicator capable of producing an unambiguous color change is usually used to detect
the end-point of the titration. In general, these titrations require specialized indicators that form weaker complexes with the analyte. A common example is Eriochrome Black T for the titration of calcium and magnesium ions.
ZETA-POTENTIAL TITRATION: Zeta potential titration is a titration of heterogeneous systems, such as colloids, emulsions, etc. Solids in such systems have very high surface area. This type of titration is used to study the zeta potential of these surfaces under different conditions. Zeta potential is a scientific term for electrokinetic potential in colloidal systems. In the colloidal chemistry literature, it is usually denoted using the Greek letter zeta, hence ζ - potential . From a theoretical viewpoint, zeta potential is electric potential in the interfacial double layer (DL) at the location of the slipping plane versus a point in the bulk fluid away from the interface. In other words, zeta potential is the potential difference between the dispersion medium and the stationary layer of fluid attached to the dispersed particle. The Iso-electric point is one such property. The iso-electric point is the pH value at which thezeta potential is approximately zero. At a pH near the iso-electric point (± 2 pH units), colloids are usually unstable; the particles tend to coagulate or flocculate. Such titrations use acids or bases as titration reagents. Another purpose of this titration is determination of the optimum dose of surfactant for achieving stabilization or flocculation of a heterogeneous system. In a zeta-potential titration, the Zeta potential is the indicator. Measurement of the zeta potential can be performed using microelectrophoresis, or electrophoretic light scattering, orelectroacoustic phenomena. The last method makes possible to perform titrations in concentrated systems, with no dilution.
IODOMETRIC TITRATION Iodometry is a method of volumetric chemical analysis, a titration where the appearance or disappearance of elementary iodine indicates the end point. Usual reagents are sodium thiosulfate as titrant, starch as an indicator (it forms blue complex with iodine molecules - though polyvinyl alcohol has started to be used recently as well), and an iodine compound (iodide or iodate, depending on the desired reaction with the sample). The principal reaction is the reduction of iodine to iodide by thiosulfate: I2 + 2 S2O32− → S4O62- + 2 I-
Two possible sources of error can influence the outcome of the iodometric titration:
the air oxidation of acid-iodide solution the volatility of I2.
The first one can be eliminated by adding an excess of sodium carbonate in the reaction vessel. This displaces the oxygen in the vessel by forming carbon dioxide gas, which blows the air away, and being heavier than air, protects the mixture by remaining on top. The other error can be reduced by using an excess of iodide solution which captures liberated iodine to form triiodide ions, I3−.
IODIMETRIC TITRATION In iodimetric titrations, free iodine is used. Since it is difficult to prepare the solution of iodine (iodine sublimates and is less soluble in water) it is dissolved in KI solution. KI+I2
KI3
This solution is first standardized before use. With the standard solution of I2, substances such as sulfite, thiosulphate, arsenite are estimated.
Difference between iodometric and iodimetric titration:
S.no.
Iodometry
Iodimetry
1.
In such titration, iodine is liberated from chemical reactions. KI is used as a reagant
A reducing agent is directly titrated against standard I2 solution KI is basically used for increasing the solubility of I 2
3.
e.g.- determination of copper ions in the copper ore solution
e.g.-determination of thiosulphate, arsenate and sulphite.
Back Titration: Back titration is an analytical chemistry technique that allows the user to find the concentration of a reactant by reacting it with an excess volume of reactant of known concentration. The resulting mixture is then titrated, taking into account the molarity of the excess that was added. This is used as opposed to standard volumetric titration when the substance being analyzed is either too weak to give a valid reaction, or too slow. A back titration is useful if the endpoint of the reverse titration is easier to identify than the endpoint of the normal titration. Back titration is also useful when trying to work out the amount of an acid or base in a non-soluble solid. The laboratory skills required for back titration are more or less similar to that of common acid-base titration.
Measuring the endpoint: Different methods to determine the endpoint include:
pH indicator: This is a substance that changes colour in response to a chemical change. An acid-base indicator (e.g., phenolphthalein) changes colour depending on the pH. Redox indicators are also frequently used. A drop of indicator solution is added to the titration at the start; when the colour changes the endpoint has been reached. A potentiometer can also be used. This is an instrument that measures the electrode potential of the solution. These are used for titrations based on a redox reaction; the potential of the working electrode will suddenly change as the endpoint is reached. pH meter: This is a potentiometer that uses an electrode whose potential depends on the amount of H+ ion present in the solution. (This is an example of an ion-selective electrode.) This allows the pH of the solution to be measured throughout the titration. At the endpoint, there will be a sudden change in the measured pH. It can be more accurate than the indicator method, and is very easily automated. Conductance: The conductivity of a solution depends on the ions that are present in it. During many titrations, the conductivity changes significantly. (For instance, during an acid-base titration, the H+ and OH- ions react to form neutral H2O. This changes the conductivity of the solution.) The total conductance of the solution depends also on the other ions present in the solution (such as counter ions). Not all ions contribute equally to the conductivity; this also depends on the mobility of each ion and on the total concentration of ions (ionic strength). Thus, predicting the change in conductivity is harder than measuring it.
Colour change: In some reactions, the solution changes colour without any added indicator. This is often seen in redox titrations, for instance, when the different oxidation states of the product and reactant produce different colours. Precipitation: If the reaction forms a solid, then a precipitate will form during the titration. A classic example is the reaction between Ag + and Cl- to form the very insoluble salt AgCl. This usually makes it difficult to determine the endpoint precisely. As a result, precipitation titrations often have to be done as "back" titrations (see below). An isothermal titration calorimeter uses the heat produced or consumed by the reaction to determine the endpoint. This is important in biochemical titrations, such as the determination of how substrates bind to enzymes. Thermometric titrimetry is an extraordinarily versatile technique. This is differentiated from calorimetric titrimetry by the fact that the heat of the reaction (as indicated by temperature rise or fall) is not used to determine the amount of analyte in the sample solution. Instead, the endpoint is determined by the rate of temperature change . Spectroscopy can be used to measure the absorption of light by the solution during the titration, if the spectrum of the reactant, titrant or product is known. The relative amounts of the product and reactant can be used to determine the endpoint. Amperometry can be used as a detection technique (amperometric titration). The current due to the oxidation or reduction of either the reactants or products at a working electrode will depend on the concentration of that species in solution. The endpoint can then be detected as a change in the current. This method is most useful when the excess titrant can be reduced, as in the titration of halides with Ag+. (This is handy also in that it ignores precipitates.)
Applications of Titrations:
As applied to biodiesel, titration is the act of determining the acidity of a sample of WVO by the drop wise addition of a known base to the sample while testing with pH paper for the desired pH=8.5 reading. By knowing how much base neutralizes an amount of WVO, we discern how much base to add to the entire batch. Titrations in the petrochemical, polymer science or food industry to define oils, fats, polymers or biodiesel and similar substances. An example procedure for all three can be found here:. Acid-base titrations: Acid value (ASTM D 974 and D 604, DIN EN ISO 2114): The mass in milligrams of potassium hydroxide (KOH) required to neutralize carboxylic
acid groups in one gram of a chemical substance. Titration takes place at low temperatures. This titration is used for example to determine the free fatty acid content.. Saponification value (ASTM D 94 and DIN 51559): The mass in milligrams of KOH required to saponify ester and to neutralize carboxylic acid groups in one gram of a chemical substance. Titration takes place at high temperatures. This method is used for example to get a hint about the average chain length of fatty acids in a fat. Ester value (or "ester index"): This index is not measured, it is calculated: Ester value = Saponification value – Acid value. Amine value (ASTM D 2073): The mass in milligrams of KOH equivalent to the total amine content in one gram of a chemical substance. Hydroxyl number (DIN 53 240-2): The mass in milligrams of KOH required to neutralize the hydroxyl groups in one gram of a chemical substance. It is determined byacetylation using acetic anhydride and titration of the acetic acid and excess anhydride with KOH. Redox titrations: Bromine number: A measure of unsaturation expressed in terms of the mass in grams of bromine that reacts with 100 grams of a chemical substance. Iodine number: A measure of unsaturation expressed in terms of the mass in grams of iodine that reacts with 100 grams of a chemical substance. This method is used for example to indicate the amount of unsaturated fatty acids. Karl Fischer titration: a method to analyze trace amounts of water in a substance in a lab. Nonaqueous_titration: Non-Aqueous Titrations (with solvents other than water).
Titration is used in heaps of industries. It is used in wineries, dairy farms, mining corporations, cleaning material manufaturers, juce makers, food makers, cosmetic industries, health industries, water plants, paint makers and heaps more. Pretty much any industry that relies on something that has a pH uses titration. Usually it's used as a way to make sure that somehting's pH is sutable for human consumption or for human to be close to. However, it is also used to make sure that products, such a cleaning products, remove bacteria. Cleaning products need to be slightly acidic for these products to work so they titrate to get the right molarity
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http://www.ehow.com/about_5065624_definition-titration.html http://www.giangarloscientific.com/analytical/hanna/Titraton2.htm http://en.wikipedia.org/wiki/Titration www.google.com http://en.wikipedia.org/wiki/Titration_curve http://en.wikipedia.org/wiki/Acid-base_titration
