The titration method.
The acid-base titration. The Mohr’s method. The oxidation-reduction titration. Complexometric titration
Titrimetric analysis- is a method of quantitative analysis used to determine unknown concentration of known substance.
The definite volume of the analyte (i.e.; the substance to be determined) is allowed to react with a suitable reagent whose standard solution can be prepared and the volume of the solution consumed for complete reaction is used to find out the concentration of analyte solution.
At this point, it is necessary to know definitions of some useful terms.
In Titrimetric analyses, the solution of accurately known concentration i.e.; standard solution is called the Titrant and the substance to be determined is called Titrand or analyte.
The volume of the titrant added is measured with a special type of a glassware called burette which is graduated and has a stopcock at one extreme end to control the flow of titrant.
The process of determining analyte by adding the small increments of standard solution untill the reaction is just complete ; the reacting ratio of the two being known from stoichiometry or otherwise is called titration.
Eg. Determination of a substance A by adding increments of substance B (almost always as a standardized solution) with provision for some means of recognizing the point at which all of A has reacted, thus allowing the amount of A to be found from the known amount of B added up to this point, the reacting ratio of A and B being known from stoichiometry or otherwise.
-The reverse process—incremental addition of A to B—is seldom applied, except in standardization titrations.
The point in a titration at which the amount of titrant added is chemically equivalent to the amount of substance titrated i. e. The point at which the completion of the reaction occurs is called the equivalence point or theoreticalor stoichiometric end point.
The point at which the completion of a reaction is practically observed is called end point. The point in a titration at which some property of the solution (as, for example, the colour imparted by an indicator) shows a pronounced change, corresponding more or less closely to the equivalence-point.
-The end-point may be represented by the intersection of two lines or curves in the graphical method of end-point determination.
-In an ideal titration, end point and equivalence points should coincide while practically (in reality) this doesn’t happen.
-End point is generally a little ahead of equivalence point.
The completion of the titration is accompanied by some physical change in the reaction mixture which can be identified visually or with the help of some instrumental techniques.
This physical change could be:
a) A colour change produced by the titrant itself (e.g.; pink colour of KMnO4)
b) A colour change produced by an additional reagent called ‘indicator’
c) A sudden change in properties like conductance, pH, e.m.f., shift in Absorption maxima or variation in absorbance etc.
Since in a titration we measure the quantity of one reactant that is required to consume all of another reactant ,we have been concerned about completion of chemical reaction ,achieving chemical equivalence in a reaction, so we put restrictions on system.
-Now what type of systems i.e., reactions can be studied titrimetrically. Definitely not all in the universe. Therefore some criteria have to be set which must be fulfilled for a given set of reactants/reactions so that they can be analysed titrimetrically.
These criteria are:
a.The reaction must be simple and well defined i.e., stoichiometric. This means one should be able to present the reaction with a balanced chemical equation. The reaction must proceed by a definite chemistry. There should be no complicating side reactions.
b. The reaction must approach completion at the equivalence point. In other words, chemical equilibrium favors products.
c. The reaction should be instantaneous or very fast. In some cases a catalyst may be employed to increase the speed of the reaction.
d. There should be a discernible change in some property of the solution when the reaction is complete. This may be a change in the colour of the solution or any other physical property.
e. In case the titrant does not produce a visible change, an indicator should be available which by a change in physical properties (colour change or precipitation) is able to sharply define the end point.
A few rules of thumb for designing a successful titration are:
* The titrant should either be a standard or should be standardized.
* The equivalence point must be stable and well defined and able to be detected.
* The titrant’s and sample’s volume or mass must be accurately known.
* The end point should not be too far from the equivalence point.
In Titrimetric analysis, we often talk about standard solutions. Standard solutionis the one whose concentration is known. The chemicals which are used to prepare these standard solutions are of two kinds: Primary Standardand Secondary Standard.
A Primary Standard substanceis a compound of sufficient purity from which a standard solution can be prepared by direct weighing of a quantity of it followed by dissolution in a defined volume of a solvent. The solution obtained is thus a primary standard solution. A compound should satisfy following criterion to act as a primary standard:
1) It should be pure. In case slight impurities are present then the impurity level should not be too high and its percentage should be known
2) It should be stable upto moderate temperatures required for drying and it should be stable indefinitely at room temperature i.e., it should not be altered in air during weighing. This means it should not be hygroscopic, oxidized by air or affected by CO2. Its composition should be unchanged during storage.
3) The substance should be capable of being analysed for impurities by known reactions.
4) It should have a high relative molecular mass so that weighing errors are minimum or negligible.
5) The substance should be readily soluble under the conditions in which it is employed.
6) The reaction with the standard solution should be stoichiometric and instantaneous. The titration error should be negligible.
A solution prepared from a primary standard substance whose concentration is known from the weight of the substance in a known volume of the solution is called primary standard solution
A standard solution prepared from a primary standard substance whose concentration is known from the weight of that substance in a known volume (or weight) of the solution.
A Secondary standard substanceis a substance whose actual active content is found by comparison against a primary standard through chemical reactions.
Thus a Secondary standard solution is one in which the concentration of the dissolved solute has not been found from the weight of the compound dissolved but by reaction (titration) of a volume of the solution against measured volume of a primary standard solution i.e. its concentration or titre has been obtained by standardization, or which has been prepared from a known weight of a secondary standard substance.
Eg.KMnO4 is not a primary standard but K2Cr2O7 is.
Standardization. The process of finding the concentration or the reacting strength of a solution by titrating with a known amount of the substance which is pure or has a known reaction value.
Titre (titer).The reacting strength of a standard solution, usually expressed as the weight of titrated substance equivalent to 1 ml of the standard solution. One should not confuse it with total volume of the titrant used.
You must know definition of some useful terms:
Titrant or Standard solution – a solution of accurately known concentration.
Titration – the process of determining unknown concentration by adding the small increments of standard solution untill the reaction is just complete.
Burette – kind of laboratory glass for exact measurement of volume of solution used. Burette is graduated and has a burette tap or stopcock at one extreme end to control the flow of titrant.
Equivalence point (synonymous Stoichiometric Point or Theoretical Point). The point in a titration at which the amount of titrant added is chemically equivalent to the amount of substance titrated.
End point. The point at which the completion of a reaction is practically observed. When using an indicator, the end point occurs when enough titrant has been added to change the color of the indicator.
Indicator – a molecule whose conjugate acid or conjugate base has a different color.
Standard substance – a substance used for standardizations primary titrant solution. A compound has satisfy following criterion to act as a standard substance:
It should be pure.
It should be stable up to moderate temperatures required for drying and it should be stable indefinitely at room temperature. This means its composition should be unchanged during storage.
It should have a high molecular mass so that weighing errors are minimum.
The substance should be readily soluble.
The reaction with the standard solution should be stoichiometric and instantaneous. The titration error should be negligible.
Standardization – the process of finding of concentration or the reacting with a known amount of the substance which is pure or has a known reaction value.
Concentrations of Standard solutions
Titrimetric Calculations can be done using two concentration units:
Molarityand Normality
Molarity is defined as no. of moles of solute (or analyte) dissolved per litre of the solution. Molar is abbreviated as M
For relatively small quantities encountered in titrations, where mL are used
Where molarity is expressed in mmol/mL
Now for a reaction where an analyte A reacts with the Titrant T to give the product P
Since the reaction is in 1:1ratio
However the reactions most often are not in 1:1 ratio
thus every mmol of A ~ a/t mmoles of T
or
or
where
t no. of mmol of titrant in balanced chemical equation
a no. of mmol of analyte in balanced chemical equation
Normality
For such cases where reactions are not on 1:1 basis, the calculations are quite often based oormality. Normality of a solution is equal to the number of equivalents of the substance per litre of the solution. Symbol N stands for
Equivalents are based on the same concept as moles but the number of equivalents will depend on the no. of reacting units supplied by each molecule or the number with which it will react.
E.g.;
In 1 mole of HCl, one mole of H+ is present therefore no. of reacting units is 1.
In 1 mole of H2SO4, two moles of H+ are present therefore no. of reacting units is 2.
The no. of equivalents can therefore be calculated from no. of moles as
Equivalent Weight is therefore defined as that weight of a substance (in g) that will furnish one mole of reacting unit.
Thus
Thus no. of equivalents can be calculated as
Normality of a solution is therefore
Relation betweeormality and molarity
equiv = mol x no. of reacting units per molecule (n).
Dividing both sides by L
equiv /L = mol/L x no. of reacting units per molecule (n).
Whereis no. of reacting units per molecule or stoichiometry factor
From here only we can deduce that
Thus our earlier equation
can be written as
Steps in the process of a titration
• primary standard solution preparation
• titrant preparation
• titrant standardization
• analyte titration with the titrant solution
• data analysis
VOLUMETRIC CALCULATIONS
Simple, based upon law of equivalents which states that at the end point or equivalence point, the number of equivalents of the substance titrated is equal to the number of equivalents of the titrating reagent used. Thus if V1 ml of the solution 1 of the normality N1 requires V2 ml of solution 2 of normality N2 for reaction completion (indicated by end point) ,then
This normality formula is the key relation in all volumetric calculations.
Calcultion of normality
In the volumetric exercise V1 ,V2 and one of the normalties are invariably known. The unknown normality is calculated by using formula
Calculation of strength of solution
A development of profound importance in practical analysis was the realisation that titrimetric procedures could be carried out with greater speed and convenience if the concentrations of the two reacting solutions were such that the reaction with the analyte was complete when comparable volumes of sample and titrant solutions had been brought together. More specifically, if volumes V1 and V2 of these solutions were mixed the reaction would be stoichiometric when N1V1 = N2V2 where ‘NX’ the ‘normality’ of the solution designated the number of ‘gram equivalents’ per litre.
The principle involved in all titration methods is to balance a chemical reaction between titrant and titrand.
Types of Titration
Neutralisation (Acid-Base) titration
Precipitation titration
Reduction-Oxidation (Redox) titration
Complexometric titration
Titration is a general class of experiment where a known property of one solution is used to infer an unknown property of another solution. In acid-base chemistry, we often use titration to determine the pH of a certain solution.
A setup for the titration of an acid with a base is shown in :
Figure: A titration setup
We use this instrumentation to calculate the amount of unknown acid in the receiving flask by measuring the amount of base, or titrant, it takes to neutralize the acid. There are two major ways to know when the solution has beeeutralized. The first uses a pH meter in the receiving flask adding base slowly until the pH reads exactly 7. The second method uses an indicator. An indicator is an acid or base whose conjugate acid or conjugate base has a color different from that of the original compound. The color changes when the solution contains a 1:1 mixture of the differently colored forms of the indicator. As you know from the Henderson-Hasselbalch equation, the pH equals the pK a of the indicator at the endpoint of the indicator. Since we know the pH of the solution and the volume of titrant added, we can then deduce how much base was needed to neutralize the unknown sample.
Titration Curves
A titration curve is drawn by plotting data attained during a titration, titrant volume on the x-axis and pH on the y-axis. The titration curve serves to profile the unknown solution. In the shape of the curve lies much chemistry and an interesting summary of what we have learned so far about acids and bases.
The titration of a strong acid with a strong base produces the following titration curve:
Figure: Titration curve of a strong base titrating a strong acid
Note the sharp transition regioear the equivalence point on the . Also remember that the equivalence point for a strong acid-strong base titration curve is exactly 7 because the salt produced does not undergo any hydrolysis reactions.
However, if a strong base is used to titrate a weak acid, the pH at the equivalence point will not be 7. There is a lag in reaching the equivalence point, as some of the weak acid is converted to its conjugate base. You should recognize the pair of a weak acid and its conjugate base as a buffer. In , we see the resultant lag that precedes the equivalence point, called the buffering region. In the buffering region, it takes a large amount of NaOH to produce a small change in the pH of the receiving solution.
Figure: Titration curve of a strong base titrating a weak acid
Because the conjugate base is basic, the pH will be greater than 7 at the equivalence point. You will need to calculate the pH using the Henderson-Hasselbalch equation, and inputting the pK b and concentration of the conjugate base of the weak acid.
The titration of a base with an acid produces a flipped-over version of the titration curve of an acid with a base. pH is decreased upon addition of the acid.
Note that the pH of a solution at the equivalence point has nothing to do with the volume of titrant necessary to reach the equivalence point; it is a property inherent to the composition of the solution. The pH at the equivalence point is calculated in the same manner used to calculate the pH of weak base solutions in Calculating pH’s.
When polyprotic acids are titrated with strong bases, there are multiple equivalence points. The titration curve of a polyprotic acid shows an equivalence point for the each protonation:
Figure: Titration curve of a strong base titrating a polyprotic acid
The titration curve shown above is for a diprotic acid such as H2SO4 and is not unlike two stacked . For a diprotic acid, there are two buffering regions and two equivalence points. This proves the earlier assertion that polyprotic acids lose their protons in a stepwise manner.
Acid-Base Titration
These titrations are based on the neutralization reaction that occurs between an acid and a base, when mixed in solution.
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A neutralization reaction in aqueous solution is a reaction of an acid and a hydroxide base to produce a salt and water Example: molecular: HCl +NaOH →NaCl + HOH ionic: H+ + |
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 a unknown acid or base solution.
An acid-base titration in which a base is titrated with a standard solution of an acid is called Acidimetric
An acid-base titration in which an acid is titrated with a standard solution of an alkali (a base) is called Alkalimetric
Titration techniques
Before starting the titration a suitable pH indicator must be chosen. The burette should be rinsed with the standard solution, the pipette with the unknown solution, and the conical flask with distilled water.
Burette filling instruction
Always use a small funnel to fill a burette
To fill a burette, close the stopcock at the bottom. You may need to lift up the funnel slightly, to allow the solution to flow in freely
Fill the burette past the zero mark
Check the tip of the burette for an air bubble. To remove an air bubble you must lift up tip of burette and then open stopcock. If an air bubble is present during a titration, volume reading may be in error!
Take the funnel out of the burette so that drops of solution from the funnel will not fall into the burette.
When you burette is filled, with no air bubbles, you must level of the liquid to exactly the zero mark. Read the bottom of the meniscus. Be sure your eye is at the level of meniscus, not above or below
After filling burette, 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.
Slowly release known solution from the burette into the conical flask, while swirling the mixture.
The solution should be let out of the burette until the indicator changes colour and value on the burette should be recorded.
From the total volume of known solutioeeded to react the end point, the concentration of the unknown solution can be calculated.
Use formula: N1V1=N2V2
Where:
N1 – normality of solution with known concentration
V1 – volume of solution with known concentration
N2 – normality of solution with unknown concentration
V2 – volume of solution with unknown concentration
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Example: Problem. 30 ml of 0.10N NaOH neutralised 25.0 ml of hydrochloric acid. Determine the concentration of the acid. Solution. N1– normality of NaOH = 0,1 mol-equiv/l V1 – volume of NaOH = 30 ml V2 – volume of HCl = 25 ml N2 – normality of HCl – ?
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Acid-base titration can be used to determine most acids and bases, strong and not too weak, monoprotic and polyprotic.
For example: we can use acid-bace titration to determine concentration of hydrochloric acid, sulphuric acid, acetic acid, as well as bases – like sodium hydroxide, ammonia and so on.
Most commonly used reagents are hydrochloric acid and sodium hydroxide. Solutions of hydrochloric acid are stable; solutions of sodium hydroxide can dissolve glass and absorb carbon dioxide from the air, so they should be not stored for long period time.
There are many standard substances that can be used in acid-base titrations. Those most popular are sodium carbonate (Na2CO3), borax (disodium tetraborate decahydrate Na2B4O7·10H2O).
Acid-base titrations are based on the neutralization reaction. They are sometimes called alkalimetric titrations and general name of the method is alkalimetry, although these are not used as often as just “acid-base titration”.
Acid-base titrations can be used to determine most acids and bases, strong and not too weak, monoprotic and polyprotic. For example we can use acid-base titration to determine concentration of hydrochloric acid, sulfuric acid, acetic acid, as well as bases – like sodium hydroxide, ammonia and so on. In some particular cases, when solution contains mixture of acids or bases of different strengths, it is even possible to determine in one titration composition of a mixture – for example sodium hydroxide and sodium hydrogen carbonate. Using acid-base back titration it is also possible to determine amount of substances that can be easily dissolved in acids, like calcium carbonate. To do so we would add known amount of hydrochloric acid to calcium carbonate and after the solid is dissolved we would titrate excess acid with a strong base.
Most commonly used reagents are hydrochloric acid and sodium hydroxide. Solutions of hydrochloric acid are stable, solutions of sodium hydroxide can dissolve glass and absorb carbon dioxide from the air, so they should be not stored for long periods of time.
There are many standard substances that can be used in acid base titrations. Those most popular are sodium carbonate Na2CO3, borax (disodium tetraborate decahydrate) Na2B4O7·10H2O and potassium hydrogen phthalate KHC8H4O4, often called simply KHP.
Acid – Base indicators
Acid – Base indicators (also known as pH indicators) are substances which change colour with pH. They are usually weak acids or bases, which when dissolved in water dissociate slightly and form ions. Consider an indicator which is a weak acid, with the formula HIn. At equilibrium, the following equilibrium equation is established with its conjugate base:
The acid and its conjugate base have different colours. At low pH values the concentration of H3O+ is high and so the equilibrium position lies to the left. The equilibrium solution has the colour A. At high pH values, the concentration of H3O+ is low – the equilibrium position thus lies to the right and the equilibrium solution has colour B.
Most commonly used indicators in acid-base titration are:
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Acid / Base Indicators |
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Indicator |
Color of acid form |
Transition Range, pH |
Color of basic form |
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methyl orange |
red |
3,0 – 4.4 |
yellow |
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methyl red |
red |
4.8 – 6.0 |
yellow |
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litmus |
red |
4,5 – 8,3 |
blue |
|
phenolphthalein |
colorless |
8.3 – 10.0 |
dark pink |
Type of indicator depends on several factors. One of them is the equivalence point pH. Depending on the titrated substance and titrant used this can vary, usually between 4 and 10. However, even if it is often possible (see list of pH indicators) we are rarely selecting indicator that changes color exactly at the equivalence point, as usually increase of accuracy doesn’t justify additional costs. Thus in practice you will probably use phenolphtalein when NaOH is used as the titrant and methyl orange when titrating with the strong acid.
Different indicators require different methods of preparation. In the following table we have collected some basic information about indicators and their solutions. See more complete list of pH indicators at www.ph-meter.info (note that this list doesn’t contain information about solution preparation).
Please note, that different sources list different methods for indicator preparation. For example for 2,4-dinitro phenol some sources propose to prepare 0.1% ethanol solution, while other suggest using just a saturated water solution.
Colors of some of the indicators listed in the table above:
Colors on the pictures can slightly differ from reality, as hues displayed depend on the monitor calibration. pH values listed on the pictures refer to the actual pH of the solution as prepared, not the pH at which indicator changes color. In lab reality colors will be less saturated, as to be sure colors are well visible we used an excess of indicators.
See also detailed discussion on the pH indicators equilibria and color changes.
Remember to use analytical reagents (AR grade) for standards.
Borax Na2B4O7·10H2O – before weighting keep in a hygrostat over NaBr·2H2O. Equivalent weight is one-half the formula weight (
Sodium carbonate Na2CO3 – heat for 1 hour at 270-
Potassium hydrogen phthalate (KHP, KHC8H4O4) – dry for 2 hours at 120° C. Equivalent weight is
Potassium hydrogen bis(iodate) KH(IO3)2 – dry at
Weak Acid Strong Base Titration
When a weak acid is titrated with a strong base the curve is quite different in two important ways.
· Once the addition of strong base begins the solution is buffered before the equivalence point.
· The solution is basic at the equivalence point because a salt of a weak acid and a strong base undergoes hydrolysis to give a basic solution.
As before we can separate the calculations for this kind of titration into four distinct types corresponding to four regions of the titration curves.
Before any base is added the pH depends on the weak acid alone
After some base has been added but before the equivalence point a series of weak acid/ salt buffer solutions determines the pH.
At the equivalence point, hydrolysis of the anion of the weak acid determines the pH.
Beyond the equivalence point, excess strong base determines the pH.
In this case the titration curve will be as follows.
Weak Base Strong Acid Titration
When we titrate a weak base by a strong acid, we get a titration curve similar to that obtained when a weak acid is titrated by a strong base.
The titration curve of a weak base has an initial point at which the system is a weak base solution followed by a buffer region, then as equivalence point at which the system is a weak acid solution, and finally a region in which strong acid is in excess.
Strong Acid Strong Base Titration
A titration curve is a plot of pH versus the amount of acid or base added. It displays graphically the change in pH as acid or base is added to a solution and shows how pH changes during the course of the titration.
As the base is added the acid is slowly neutralized. At first the change in pH is minimal. This resistance is due to the fact that the flask has a much greater number of H3O+ ions than the
As more and more base is added, more
In this case the titration curve will be as follows.
Weak Acid Weak Base Titration
In the titration of a weak acid and a weak base a titration curve is like
The change in pH at the end point is gradual and indicators will change color gradually. No indicator will give a sharp end point.
The way out of this difficulty is to titrate the weak acid against a strong base and the weak base against a strong acid.
Acid Base Titration Calculations
One of the important application of titration is the acid base titrations. It involves either the estimation of the strength of an acid or a base of unknown strength by titrating it against a base or an acid of known strength respectively. Suitable indicator is used to ascertain the correct end point of the titration.
There are several steps for the Acid base titration calculations. Since the mineral acids are in liquid form it is not possible to make a standard solution which is also called the primary standard solution with them. It is always required to standardize these acids with another titration.
Hence these acids form secondary standards, with which the main titration of finding the strength of a given base is done. In the same way the alkalis like NaOH, KOH are hygroscopic solids and will not give the proper result. These are again unfit to make primary standards. They will be standardized by a primary standard made from organic acid like oxalic acid which is in crystalline form.
Example;Finding the strength of given hydrochloric acid (HCl) using oxalic acid [(COOH)2.2H2O] and NaOH.
Calculations involved in making of Primary standard of Oxalic Acid
Weight of the dish with Oxalic acid before transferring it in to standard flask = ‘a’ grams
1. Weight of the dish with oxalic acid after transferring it in to standard flask = ‘b’ gms.
2. Exact weight of Oxalic acid = (a – b) gms.
Normality of oxalic acid solution dissolved in 250 mL = Mass / Eq.mass.
Oxalic acid Eq.mass = Mol. mass /2. (126/2 = 63) = (a – b) / 63 N
Calculations involved in making of Secondary standard of Sodium hydroxide solution
Volume of Oxalic acid solution taken ( Pipette reading) = Vox mL.
1. Normality of Oxalic Acid = (a – b) / 63 = N ox
2. Volume of NaOH solution ( Burette reading) = VNaOH mL
3. Normality of NaOH = NNaOH = Vox X N ox / VNaOH
Calculations involved in finding the strength of Hydrochloric Acid
Volume of NaOH solution ( Burette reading) = VNaOH mL
1. Normality of NaOH = NNaOH
2. Volume of HCl taken ( Pipette reading ) = V HCl
3. Normality of HCl = N HCl = VNaOH X
4. Strength of HCl = N X Eq.mass ( For HCl Eq.mass= Mol.mass) = NHCl X
Strength is pH can also be calculated by converting the normality in to molarity (In this case it is the same)
Acid Base Titration Problems
Below you could see problems
Solved Examples
Question 1: If 0.5gram of a mixture of K2CO3 and Li2CO3 requires 30ml of 0.25M acid solution for neutralization what is the percentage composition of the mixture.
Solution:
Let x = gram of K2CO3 y = gram of Li2CO3
Then x + y = 0.5
No. of meq of K2CO3 present = xK2CO3/2000 = x0.060
No. of meq of Li2CO3 = yLi2CO3/2000 = y0.0375
No. of meq of acid required = 30 x 0.25
Hence x/0.69 + y/0.0375 = 30 x 0.25
Solving the above simultaneous equations we have
x = 0.247 and y = 0.253
Hence percentage 100x/0.5 = 49.4% K2CO3
and 100y/0.5 = 50.6% Li2CO3
Question 2: A 1.2gram sample of a mixture of (Na2CO3 + NaHCO3) is dissolved and titrated with 0.5N HCl. With phenolphthalein the end point is at 15ml while after further addition of methyl orange a second end point is at 22ml. Calculate the percentage composition of the mixture.
Solution:
15 + 15 = 30ml acid is necessary to neutralize Na2CO3 completely. Total volume needed = 15 + 22 = 37ml that is (37-30) = 7ml acid is needed for neutralizing NaHCO3
Therefore Na2CO3 composition (%) is 30 X 0.5 X 0.0531.2 X 100 = 66.25%
7 X 0.5 X 0.042 X 1001.2 = 24.50% NaHCO3
Precipitation Titration
Precipitation Titration it is a volumetric titration method where the reaction between the titrant and sample solution yield precipitate (low solubility, usually ionic compounds)
The most important precipitating reagent is silver nitrate.
Titrimetric methods based upon silver nitrate are sometimes termed argentometric methods.
Argentometry, where the titrant is a standard AgNO3 solution is the most common precipitation titrimetric method, because
silver precipitates are usually highly insoluble
many species form stoichiometric precipitates with Ag+ (e.g. Cl–, Br–, I–, F–, CN–, SCN–, CrO42-, PO43- etc.)
these precipitates are formed quickly
Titrant is a standardized AgNO3 solution. The titrant needs to be stored in a dark (brown) container.
Argentometry is most often used for determination of chloride ions, but it can be used for other halides (bromide, iodide).
There are 3 techniques of end point determination:
– method of Mohr (indicator: potassium chromate)
– method of Volhard (indicator: ferric salt)
– method of Fajans (indicator: fluorescein)
The most often used Mohr method
Two most important solutions used in argentometric methods are solution of silver nitrate and solution of potassium thiocyanate.
Silver nitrate solution
Silver nitrate solution of known concentration can be prepared using pure solid AgNO3, after drying it (see standard substances used in precipitation titrations section). Most popular solution is that of 0.1M concentration, although for determination of small amount if chlorides more diluted solutions can be used (0.02M). However, use of diluted solutions should be preceded by thorough analysis of possible titration errors. This is especially important in the case of Mohr titration, where some excess of silver must be added before red silver dichromate precipitates and signals end point.
Silver nitrate solutions slowly decompose when exposed to light, so they should be kept in dark bottles.
Potassium thiocyanate solution
Potassium thiocyanate is not used as a standard substance. Its solutions are prepared by dissolving solid KSCN and standardized against solution of silver nitrate of known concentration. Usually used solution is 0.1M.
0.1M silver nitrate standardization against sodium chloride
Silver nitrate solutions of known concentration can be prepared from known mass of dried AgNO3. However, if we don’t have access to the high purity reagent, or if we have a solution of unknown concentration, we can easily standardize it against sodium chloride.
Reaction taking place during titration is
AgNO3 + NaCl → AgCl↓ + NaNO3
Procedure to follow:
Weight exactly about 0.15-
Add about 100 mL of distilled water, dissolve.
Add 1 mL 5% w/w potassium chromate solution.
Titrate with AgNO3 solution till the first color change.
0.1M potassium thiocyanate standardization against silver nitrate solution
Potassium thiocyanate solution has to be standardized, as it is not possible to prepare and dry KSCN pure enough so that it can be used as a standard substance for solution preparation. The easiest method if the standardization require standardized solution of silver nitrate. As KSCN solution is used for back titration of the excess of AgNO3, when we need to standardzie KSCN solution we usually have standardized silver nitrate solution ready.
Reaction taking place is
AgNO3 + KSCN → AgSCN↓ + KNO3
Procedure to follow:
Pipette 25 mL aliquot of about 0.1M AgNO3 solution into 250mL erlenmayer flask.
Add 50 mL of distilled water.
Add 1 mL of 10% FeNH4(SO4)2 solution.
Titrate with potassium thiocyanate till the first visible color change.
Mohr method
Mohr titration is used for determination of halide in a solution.
Potassium chromate can serve as an indicator for the determination of chloride, and bromide ions by reacting with silver ion to form a brick-red silver chromate (Ag2CrO4) precipitate in the equivalence-point region.
Mohr titration has to be performed at a neutral or weak basic solution of pH 7-9 (or 6-10), because silver hydroxide forms at high pH, while the chromate forms H2CrO4 at low pH, reducing the concentration of chromate ions and delaying the formation of the precipitate.
Mohr method of determination of chlorides by titration with silver nitrate is one of the oldest titration methods still in use – it was researched and published by Karl Friedrich Mohr in 1856.
The idea behind is very simple – chlorides are titrated with the silver nitrate solution in the presence of chromate anions. End point is signalled by the appearance of the red silver chromate.
Intense yellow color of chromate may make detection of first signs of formation of red silver chromate precipitation difficult. As some excess of silver must be added before precipitate starts to form, if concentration of titrant is below 0.1M, we may expect singificant positive error. To correct for this error we can determine a blank, titrating a solution of the indicator potassium chromate with standard silver nitrate solution. To make result more realistic we can add small amount of chloride free calcium carbonate to the solution to imitate the white silver precipitate.
Solution during titration should be close to neutral. In low pH silver chromate solubility grows due to the protonation of chromate anions, in high pH silver starts to react with hydroxide anions, precipitating in form of AgOH and Ag2O. Both processes interfere with the determination accuracy.
Exactly the same approach can be used for determination of bromides. Other halides and pseudohalides, like I- and SCN-, behave very similarly in the solution, but their precipitate tends to adsorb chromate anions making end point detection difficult.
Reaction taking place during titration are:
If Ag+ solution is add to a Cl– solution containing of small quantity of CrO4–, then AgCl will firstly precipitated, while Ag2CrO4 has not yet, and concentration Ag+ increases progressively until solubility product of the ions reach the value of Ksp Ag2CrO4 (2,0·10-12) to form brick-red precipitate.
Before titration small amount of sodium or potassium chromate is added to the solution, making it’s slightly yellow colour. During titration, as long as chlorides are present, concentration of Ag+ is too low for silver chromate formation. Near equivalence point concentration of silver cations rapidly grows, allowing precipitation of brick-red silver chromate which signals end point.
Volhard Method
It is not always possible to use Mohr method to determine concentration of chlorides. For example, Mohr method requires neutral solution, but in many cases solution has to be acidic, to prevent precipitation of metal hydroxides (like in the presence of Fe3+). In such cases we can use Volhard method, which is not sensitive to low pH.
In the Volhard method chlorides are first precipitated with excess silver nitrate, then excess silver is titrated with potassium (or sodium) thiocyanate. To detect end point we use Fe3+ cations, which easily react with the thiocyanate, creating distinct wine red complex.
There is a problem though. Silver thiocyanate solutility is slightly lower than solubility of silver chloride, and during titration thiocyanate can replace chlorides in the existing precipitate:
AgCl(s) + SCN– → AgSCN(s) + Cl–
To avoid problems we can filtrate precipitated AgCl before titration. However, there exist much simpler and easier procedure that gives the same result. Before titration we add some small volume of a heavy organic liquid that is not miscible with water (like nitrobenzene, chloroform or carbon tetrachloride). These liquids are better at wetting precipitate than water. Once the precipitate is covered with non polar liquid, it is separated from the water and unable to dissolve.
Precipitate solubility is not a problem during determination of I– and Br–, as both AgBr and AgI have much lower solubilities than AgSCN.
Reaction
There are two reactions, as this is a back titration. First, we precipitate chlorides from the solution:
Ag+ + Cl– → AgCl(s)
Then, during titration, reaction taking place is:
Ag+ + SCN– → AgSCN(s)
Sample size
In back titrations sample size is more difficult to calculate than during normal, direct titrations. For best accuracy excess of silver should be titrated with about 40-45 mL of titrant (assuming – as we usually do – that we are using 50 mL burette). However, that usually means we should use relatively large initial volume of silver solution. Assuming we will start with 50 mL of pipetted silver nitrate and we will titrate excess with about 25 mL of thiocyanate solution, and finally assuming both solutions used are 0.1M, aliquot taken for titration should contain about
Note, that silver nitrate can be added not using single volume pipette, but from burette. If the amount of chlorides is approximately known, this way it is possible to control excess of silver nitrate and volume of the thiocyanate titrant.
End point detection
End point is detected with the use of iron (III) thiocyanate complex, which have very distinct and strong wine color.
Solutions used
To perform titration we will need 0.1M silver nitrate solution to precipitate chlorides, titrant – 0.1M potassium thiocyanate solution, nitric acid (1+1) to acidify solution, ammonium ferric sulfate solution that will be used for end point detection, nitrobenzene, and some amount of distilled water to dilute sample.
Procedure
Pipette aliquot of chlorides solution into 250mL Erlenmeyer flask.
Add 5 mL of 1+1 nitric acid.
Dilute with distilled water to about 100 mL.
Add 50 mL of 0.1M silver nitrate solution.
Add 3 mL of nitrobenzene.
Add 1 mL of iron alum solution.
Shake the content for about 1 minute to flocullate the precipitate.
Titrate with thiocyanate solution till the first color change.
Result calculation
As in every back titration, to calculate amount of substance we have to subtract amount of titrated excess from the initial amount of reactant used. In the case of argentometry calculations are easy, as all substances used react on the 1:1 basis.
First we have to calculate number of moles of silver nitrate initially added to the chlorides sample. Assuming it was 50 mL of
Ag+ + SCN– → AgSCN(s)
Thus amount of excess silver is C×V, and amount of Cl is 0.005-C×V moles.
Sources of errors
Apart from problems listed on the general sources of titration errors page, results of titration can be skewed by the already mentioned replacement of precipitated chlorides by silver thiocyanate. It shouldn’t matter if the procedure was followed carefully.
Reduction-Oxidation (Redox) Titration
A redox titration is based on an oxidation-reduction reaction between analyte and titrant.
Oxidation-reduction reactions or Redox reactions involves oxidation and reduction reaction. In other words; it involves the transfer of electrons between two chemical species.
One compound in reaction loses an electron, this compound is said to be oxidized and at the same time another compound gains an electron and is said to be reduced.
There are specific terms used to describe these chemical species. The compound that is oxidized is called as a reducing agent, while another compound that is reduced is called as the oxidizing agent.
Review about oxidation-reduction reaction
Oxidation Numbers (Oxidation States)
· An oxidatioumber (oxidation state) is the charge an atom would carry if the molecule or ion were completely ionic.
· For a neutral molecule or compound, the sum of the oxidatioumbers (states) for each element in the molecule equals 0.
· For a charged molecule (ion), the sum of the oxidatioumbers (states) for each element in the ion equals the charge on the ion.
Rules for Assigning Oxidation Numbers
a. Fluorine is assigned an oxidatioumber of
b. Oxygen is assigned an oxidation number of
Except in
o Peroxides
o Superoxides
o Oxygen fluorides
c. Hydrogen is assigned an oxidation of +
d. Group 1 elements (Alkali Metals) are assigned an oxidatioumber of +
e. Group 2 elements (Alkaline-earth metals) are assigned an oxidatioumber of +
f. An atom of any element in the
g. Any monatomic ion has an oxidatioumber equal to its charge
h. The sum of the oxidation numbers of all the atoms in formula equals the electrical charge shown with the formula
Reactions in which electrons are transferred from one species to another are known as redox reactions, or, oxidation–reduction reactions.
e.g., 2 Na + Cl2 ¾® 2 NaCl
A redox reaction is made up of two reactions:
reduction — gain of electron(s)
oxidation — loss of electron(s)
Oxidizing agent is the species, which accepts electrons, e.g.,
Reducing agent is the substance, which loses electrons, e.g.,
Writing Redox Equations
In a redox reaction, the number of electrons lost by the species being oxidized must balance the number of electrons gained by the species being reduced.
In a balanced redox reaction equation:
– the number of atoms of each element must be balanced
– the total charge on the ions on the left hand side of the equation will equal the total charge on the ions on the right hand side of the equation.
Redox Reaction
Redox reaction is just like an acid-base reactions. An acid can show its acidic properties in the presence of base only. Like acid-base reaction, redox reactions are concerned with the transfer of electrons between species.
One compound has to give electron and oxidizes and another compound has to accept electron. Hence redox reaction is a combination of oxidation-reduction reaction. Each reaction by itself is known a “half-reaction” and whole reaction is called as redox reaction. For example,
Cu (s) → Cu2+ + 2 e–
This is oxidation half-reaction which implies that copper get oxidized to copper (ii) ion. Here copper is a reducing agent. Another half reaction is reduction half reaction; in which two silver ions accept two electrons to form silver atom. Here Ag+ is the oxidizing agent.
2 Ag+ (aq) + 2 e– → 2 Ag (s)
Now combination of these two half reaction gives redox reaction.
Redox Titration Table
Titration is a laboratory method of quantitative analysis to determine the concentration of given sample by using a known concentration of standard compound. Another term used for titration is titrimetry.
The compound of unknown concentration is called as analyte or titrand, while the standard solution is called as titrant or titrator. Since the titration is completely based on volume measurement, it is also known as volumetric analysis.
Titration is used to determine the concentration of an unknown substance by using standard compound. This method is based on either acid-base reaction or oxidation-reduction (Redox) reaction. Usually potassium permanganate solution is used as a standard solution. Since potassium permanganate ion is a strong oxidizing agent in acidic solution, hence can be used to analyze (by titration) solutions.
The main advantage of permanganate solution in the titration of colorless unknown solutions is that it acts as self-indicating. The color of MnO4– quickly disappears as it is reduced to Mn2+ in the presence of the reducing agent. Finally at the endpoint, all the reducing agent has been used up and next drop of MnO4– solution detected color change. By using the used concentration of the oxidizing agent added, we can figure out the concentration of reducing agent present in the unknown sample.
In place of permanganate solution, dichromate solution can also be used as standard solution. The dark violet color solution of permanganate ion changes to colorless after completely reduced. Hence the end point is very sharp under recommended reaction conditions. Generally for making standard solution of permanganate solution, it has been titrate with oxalic acid.
In case of dichromate solution, a Redox indicator like diphenylamine is used which gives a distinct color after the addition of dichromate ion. Due to extra pour nature, it can directly forms primary standard solution. There are various redox indicators used in redox titration.
Redox Titration table
|
Indicator |
Color of oxidized form |
Color of reduced form |
|
2,2- bipyridine (Ru complex) |
Colorless |
Yellow |
|
Nitrophenanthroline(Fe complex) |
Cyan |
Red |
|
N – Phenylanthranilic acid |
Violet-red |
Colorless |
|
1, 10-Phenanthroline (Fe complx) |
Cyan |
Red |
|
N – Ethoxychrysoidine |
Red |
Yellow |
|
2,2 – BIpyridine (Fe complex) |
Cyan |
Red |
|
5,6 – Dimethylphenanthroline |
Yellow-green |
Red |
|
O-Dianisidine |
Red |
Colorless |
|
Sodium diphenylamine sulphonate |
Red-violet |
Colorless |
|
Diphenylbenzidine |
Violet |
Colorless |
|
Diphenylamine |
Violet |
Colorless |
|
Viologen |
Colorless |
Blue |
The best example of redox titration is the determination of iron by using potassium permanganate or potassium dichromate as standard solution, where dichromate reacts with iron yielding Fe3+ and Cr3+.
The reaction involves oxidation of ferrous ion and reduction of dichromate ion to chromium (iii) ion.
6 Fe2+ = 6 Fe3+ + 6 e (Oxidation half reaction)
(reduction half reaction) 6 e– + 14 H+ + Cr2O72- = 2 Cr3+ + 7 H2O
____________________________________________
6 Fe2+ + 14 H+ + Cr2O72- = 6 Fe3+ + 2Cr3+ + 7 H2O
The redox titration of any substance by using permanganate solution consists of following steps.
Redox Indicators
In the redox titrations, we need a chemical species that can change colour in the potential range corresponding to the sharp change at the end point. A chemical substance, which changes colour when the potential of the solution reaches a definite value, is termed as an oxidation-reduction or redox indicator.
g .Inox + ne → Inred
colour A colour B
A redox indicator may be defined as a substance whose oxidized form is of different colour from that of its reduced form. The oxidation and reduction of the indicator is readily reversible.
Permanganatometry
Potassium permanganate is a very strong oxidizing agent and is employed in the estimation of reducing agents like ferrous salts, oxalic acid, arsenious oxide, etc. The permanganate ion, MnO4-, gets reduced to Mn2+ ion in acidic medium and to MnO2 ieutral and alkaline media.
Titrations involving potassium permanganate are usually carried out in acidic medium. This is due to higher oxidizing power of permanganate ion in acidic medium than ieutral or alkaline medium; secondly, the formation of brown coloured, MnO2 in alkaline medium interferes with the detection of the end point. For acidification of KMnO4 solution, only H2SO4 is suitable whereas the other mineral acids like HCl and HNO3 are not. HCl is not used because some of the KMnO4 will oxidize Cl– ions to chlorine gas according to the following equation and thus interferes in the quantitative estimations of reducing agents.
Nitric acid cannot be used because it is itself a strong oxidizing agent and may oxidize the reducing agent, thereby introducing error. Potassium permanganate is not a primary standard i.e., a standard solution of KMnO4 cannot be prepared by weighing because (i) it caever be obtained in the purest form (99.99%) and is always associated with organic impurities (ii) its normality changes on standing (iii) KMnO4 may react with organic matter present in water in which it is dissolved.
KMnO4 solution can be standardized by titrating with a suitable primary standard solution such as Mohr’s salt, oxalic acid or arsenious oxide etc.
KMnO4 solutioeeds to be added to a known volume of reducing agent containing dilute H2SO4, gradually in small amounts. Rapid addition of KMnO4 results in the formation of hydrated manganese dioxide, MnO2• H2O, which is brown in colour.
The above reaction also occurs if the medium is not sufficiently acidic. Potassium permanganate is such a powerful oxidizing agent that it oxidizes even water, according to the following equation,
An aqueous solution of potassium permanganate, therefore, should be unstable. However, this reaction is extremely slow; hence the permanganate solution attains reasonable stability in the absence of light. Thus, potassium permanganate solution is stored in dark coloured bottles since the above reaction is catalyzed by light.
Permanganate acts as self indicator. Since MnO4– is intense purple while Mn2+ is colourless, the reaction mixture at equivalence point is colourless and even a single drop of the permanganate would impart sufficient pink colour to the solution acting as self indicator.
Preparation of Permanganate Solutions
For the preparation of stock solution of permanganate solution, the weighing amount of the material is dissolved in an appropriate volume of distilled water.
The resulting solution is then heated to boiling for an hour and then filtered. The filtrate solution is used as primary standard solution.
Standardization of Permanganate Solutions
Sodium oxalate or arsenic (III) oxide can be used for the standardization of solution. The equilibrium redox reaction of sodium oxalate and permanganate solution is as follows;
2MnO4– + 5 C2O42- + 16 H+ → 2Mn2+ + 10 CO2 + 8H2O
While primary standard arsenic(III)oxide is soluble in alkali yielding the arsenate.
As2O3 + 4OH – → 2 HAsO32- + H2O
This arsenate further reduces the permanganate in acidic solution and form AsO43- and Mn2+ . Potassium iodide or iodate is used as a catalyst for this reaction.
As O33- → As O43-
5 (H2O + AsO33- → AsO43- + 2 H+ + 2 e)
2(MnO4– + 8 H+ + 5 e → Mn2+ + 4 H2O
________________________________________________________
5 AsO33- + 2 MnO4– + 6 H+ →5 AsO43- + 2 Mn2+ + 3 H2O
Calculations Involving Redox Systems
The equivalent weight of given unknown substance involved in Redox reactions will be,
Equivalent weight = Molecular weight number of einvolved
The relation between equivalent weight and concentration is termed called normality, N, where
N =Number of equivalents L
The relation between normality and molarity is
Normality = Molarity x number of electrons involve in reaction
Some Redox titration involves iodine as an oxidizing agent. If standard iodine is used for the oxidation of a reducing agent (analyte), the method is termed as iodimetry. On the other hand, the indirect presence of iodine in titration in a Redox reaction is called iodometry.
Iodine formed a complex ion (triiodide ion) with iodide ion solution.
I2 + I– = I3–
Here this tri-iodide ion acts as an oxidizing agent and involved in both type of redox titration.
Iodometry
Iodometry is one of the most important redox titration methods. Iodine reacts directly, fast and quantitively with many organic and inorganic substances. Thanks to its relatively low, pH independent redox potential, and reversibility of the iodine/iodide reaction, iodometry can be used both to determine amount of reducing agents (by direct titration with iodine) and of oxidizing agents (by titration of iodine with thiosulfate). In all cases the same simple and reliable method of end point detection, based on blue starch complex, can be used.
Reversible iodine/iodide reaction mentioned above is
2I– ↔ I2 + 2e–
and obviously whether it should be treated as oxidation with iodine or reduction with iodides depends on the other redox system involved.
Second important reaction used excesivelly in iodometry is reduction of iodine with thiosulfate:
2S2O32- + I2 → S4O62- + 2I-
In the case of both reactions it is better to avoid low pH. Thiosulfate is unstable in the presence of acids, and iodides in low pH can be oxidized by air oxygen to iodine. Both processes can be source of titration errors.
Iodine is very weakly soluble in the water, and can be easily lost from the solution due to its volatility. However, in the presence of excess iodides iodine creates I3- ions. This lowers free iodine concentration and such solutions are stable enough to be used in lab practice. Still, we should remember that their shelf life is relatively short (they should be kept tightly closed in dark brown bottles, and standardized every few weeks). Iodine solutions are prepared dissolving elemental iodine directly in the iodides solution. Elemental iodine can be prepared very pure through sublimation, but because of its high volatility it is difficult to weight. Thus use of iodine as a standard substance, although possible, is not easy nor recommended. Iodine solutions can be easily normalized against arsenic (III) oxide (As2O3) or sodium thiosulfate solution.
It is also possible to prepare iodine solutions mixing potassium iodide with potassium iodate in the presence of strong acid:
5I- + IO3- + 6H+ → 3I2 + 3H2O
Potassium iodate is a primary substance, so solution prepared this way can have exactly known concentration. However, this approach is not cost effective and in lab practice it is much better to use iodate as a primary substance to standardize thiosulfate, and then standardize iodine solution against thiosulfate.
Reducing Agents Used in Titrations Involving Iodine
Usually Sodium thiosulfate (Na2S2O3) also known as hypo is used as a reducing agent in redox titration involving iodine.
The reaction between iodine and thiosulfate is as follows;
I2 + 2 S2O32- → 2 I– + S4O62-
Since I2 is present as the tri-iodide in aqueous solutions containing iodide , hence reaction can be written as ;
I3– + 2 S2O32- → 3 I– + S4O62-
Because if the presence of water of hydration, sodium thiosulfate cannot be used as a primary standard. Hence reaction takes place in acidic solution and in presence of excess iodide.
IO3– + 5 I– + 6 H+ → 3 I2 + 3 H2O
Potassium dichromate solution is also used for standardization of thiosulfate in acidic solution, with excess iodide ion;
Cr2O72- + 6 I– + 14 H+ → 2 Cr3+ + 3 I2 + 7 H2O
Indicators Involved in Iodine Method
Since starch can form a complex with iodine, it is one of best and easily available indicator for this redox titration.
Calculations Involved in Iodine Methods
Calculate the molarity of thiosulfate solution if 0.200g of KIO3 required 50.0 mL of the thiosulfate solution in the presence of excess of KI and HCl.
IO3– = 3 I2 IO3– + 5 I– + 6 H+ = 3 I2 + 3 H2O
I2 = 2S2O32- I2 + 2S2O32- = 2I– + S4O62-
IO3– = 6S2O32-
molS2O2−3L = 0.200 KlO3 x molKlO3214gKlO3
6molS2O2−3L x 150.0mLS2O2−3 x 1000mLL = 0.112M
Determination of Iron by Redox Titration
Iron (II) is oxidized by potassium dichromate to iron (III) after complete reduction of iron present. The redox reaction for given reaction will be.
6 Fe2+ + 14 H+ + Cr2O72- → 6 Fe3+ + 2Cr3+ + 7 H2O
Generally iron (ii) salt that is Mohr’s salt is determining iron. The following steps can be used for the determination of iron by using redox titration;
For estimation of the strength of given Mohr’s salt solution by using potassium dichromate solution.
Step 1: Preparation of standard solution of Mohr’s salt (roughly N/40)
Check the weight of a clean dry empty weighing bottle (w1).
· Now put an approximate quantity of Mohr’s salt in the weighing bottle and check the weigh again. Check the weighing readings of weighing bottle with the compound (w2)
· Now transfer the compound into measuring flask and note down the weight of weighing bottle (w3).
· Hence the amount of compound = W3 –W2
· Now dissolve the weighted solid in minimum amount of water and add a few drops of conc. sulphuric acid. Now make up the volume till the 100 ml by adding water and put the
Step 2: Standardisation of K2Cr2O7 solution by titrating with standard Mohr’s salt solution with
Internal indicator Diphenylamine.
· Take 10 mL of Mohr’s salt solution through pipette in the conical flask.
· Add 10 mL of dil. Sulphuric acid to the flask.
· Add 2 drops of the diphenylamine and 2 mL of 1:1 ortho phosphoric acid.
· Titrate the reaction mixture with potassium dichromate solution which is taken in burette till end
· There will be a colour change from light green to blue- violet at that point)
· Repeat the titration for getting three concordant readings.
Step 3: Titration of Standardised K2Cr2O7 soln. with Mohr’s salt solution of unknown strength.
· Repeat the same titration as above given and take unknown sample of Mohr’s salt solution in place of standard Mohr’s salt solution.
Observations and Calculations
Strength (g/L) = Normality x Equivalent weight.
Normality of Mohr’s salt solution to be prepared is 0.025 N
Equivalent weight of Mohr’s salt = Mol. Wt./ no. of equivalents
= 392/ 1
= 392
Hence strength = 0.025 N x 392 = 9.80 g/L
In other words; for 100 mL solution,
Observation Table
1. Weighing observations:
a. Weight of empty bottle; w1(g) = ………………..
b. Weight of empty bottle + Mohr’s salt; w2 (g) = ………………..
c. Weight of weighing bottle after transference of Mohr’s salt to the standard solution.; w3 (g) =……
Amount of Mohr’s salt transferred to 100 ml measuring flask = ( w2 – w3) g
Thus strength of Mohr’s salt solution is = ( w2 – w3) x 10 g/L
And Normality of Mohr’s salt solution is = ( w2 – w3) x 10 N /392
(or Molarity Mohr’s salt solution is = ( w2 – w3) x
2. Titration of K2Cr2O7 soln. Vs Mohr’s salt solution
Volume of Mohr’s salt solution used in each titration = 10 mL
Indicator used = Diphenylamine (2 drops) + 2ml 1:1 H3PO4
Thus applying the normality relation
NMohrVMohr = NdichrVdichr
Ndichr = NMohrVMohr / Vdichr
3. Titration of Standardised K2Cr2O7 soln. Vs Mohr’s salt solute
Volume of Mohr’s salt solution used in each titration = 10 mL
Indicator used = Diphenylamine (2 drops) + 2 ml 1:1 H3PO4
Since;
NMohrVMohr = NdichrVdichr
Hence;
NMohr = NdichrVdichr /VMohr
Strength (g/L) of Mohr’s salt solution (of unknown conc.)
= Normality x Eq. Wt
Or = Molarity x Mol.wt
We can calculate the mass concentration (g/L) of iron in the unknown solution by using its molar mass and the molarity of the ferrous solution.
Mass concentration = mole / L x 56g / mol = g/L
Standardization of Potassium permanganate
The reducing agent in the titration to be discussed is oxalic acid here. The composition of it is H2C2O4·2H2O. Inspite of being a dehydrate it is a good primary standard as its composition is unchanged during storage or weighing.
The reaction between oxalic acid and potassium permanganate can be represented as:
This redox reaction can be split apart in two parts- one showing the oxidation and the other reduction
This titration is carried out in warm conditions (temperature about 60 oC). The reaction at room temperature is slow because of the equilibrium nature of this reaction. CO2 is highly soluble in water and thus heating removes all dissolved carbon dioxide out of the solution.
While noting the burette readings, it should be taken into account that the solution is so intensely coloured that the lower meniscus of the solution may not be clear. Thus for permanganate titrations the upper meniscus in the burette is noted.
Back Titration
As the name implies, back titration is a titration done in reverse manner means a known excess of standard reagent is added to the solution, and the excess is titrated, in place of titrating the original sample.
This type of titration is useful, when the endpoint of the reverse titration is easier to identify compare to the endpoint of the normal titration like with precipitation reactions.
Generally these titrations are useful for slow reaction of the analyte and the titrant or for non-soluble solid analyte.
For example; when a substance or solution of unknown concentration (
After completing the reaction, the resulting solution contains excess of intermediate reactant which is titrated with known volume and concentration of titrant (50 ml of 0.5N NaOH). If subscripts 1 and 2 denotes intermediate reactant and titrant, then
N1V1 = N2V2
0.5 x V = 0.5 x 50
Hence, the volume of excess HCl, V = 50ml
Also,
meq of execss HCl = meq of titrant, NaOH = 0.5 x 50 = 25
For determination of excess volume or meq of intermediate reactant, it allows us to determine the volume or meq of intermediate reactant which reacted withn analyte.
m of chalk = total m of HCl − m of excess HCl
m of chalk = 0.5 x 200 − 25
m of chalk = m of HCl used for chalk = 75
By applying gram equivalent concept to chalk and HCl,
m of chalk = m of HCl used for chalk = 75
E x 1000 = 2 x 1000 (40+12+3X16) =75
g = 7520 = 154 = 3.75gm
It means chalk contained
Redox Titration Problems
Below you could see problems
Solved Examples
Question 1: 100 litres of air at STP is slowly bubbled through 200 ml of 0.03N Ba(OH)2 solution. The BaCO3 formed due to reaction is filtered and few drops of Phenolphthalein is added to the solution rendering it pink. The solution required 25 ml of 0.2N HCl solution when indicator turned colorless. Calculate percentage by volume of CO2 in air.
Solution:
CO2 + Ba(OH)2 ——–> BaCO3 + H2O
Carbon dioxide reacts with Ba(OH)2 and form BaCO3. The excess Ba(OH)2 is back titrated with HCl.
Total meq of Ba(OH)2 solution = 0.03 x 100 = 6
meq of excess Ba(OH)2 solution = 0.1 x 25 = 5
meq of Ba(OH)2 solution used for CO2 = 6 – 5 = 1
g = meq x MO /1000x = 1 x 44 /1000 x 2 =
According to Avogadro’s hypothesis,
volume of CO2 = (22.4 x 0.022) / 44 =
% of CO2 in air = (0.02 x 100) / 50 = 0.04
Question 2: 502g sample of dry CaCO3 and CaCl2 mixture was dissolved in 25.00 mL of
Solution:
During titration 22 ×0.09312 = 2.0486m mole HCl was neutralized. Initially there was 25.00 × 0.925 = 23.12m mole of HCl used, so during CaCO3 dissolution
23.12 – 2.0486 = 21.0714 m mole of acid reacted.
Since calcium carbonate reacts with hydrochloric acid 1:2 ,hence original sample contained 21.0714/2 = 10.535 m mole of CaCO3, or
10.535 / 100 = 1.053 g (here molar mass of CaCO3 = 100.0 g/mol).
So original sample contained
(1.053 /1.502) ×(100%) = 70.10 % CaCO3 and 100.0 – 70.10% = 29.89 % CaCl2.
Question 3: Calculate the molarity of thiosulfate solution if
Solution:
molS2O2−3L = 0.400g KIO3 x molKIO3214gKIO3
6molS2O2−3molKIO3 x 1100.0mLS2O2−3 x 1000mLL
=
Question 4: Calculate the molarity of H2O2 if 125 mL of H2O2 required 25.0 mL of
Solution:
Since,
2MnO4– + 5H2O2 + 6H+ → 2Mn2+ + 5O2 + 8H2O
Hence molH2O2L = 0.1molKMnO41000mLKMnO4 x 25mLKMnO4125mLH2O2=
=5molH2O22molKMnO4 = 1000mLL= 0.05M
Complexometric titration
These titrations are based on complexation reactions.
Most often used reagent is EDTA – EthyleneDiamineTetraAcetic acid. There are also other similar chelating agents (EGTA, CDTA and so on) used. In some of other methods Ag+ is used as a titrant for determining cyanides and Hg2+ as a titrant in Cl- determination.
Changing property of the solution is usually the concentration of the complexed substance, although in some cases it can be much more convenient to express results in terms of titrant concentration. As its concentration changes by many orders of magnitude, and is almost always smaller than 1, we use negative logarithmic scale, similar to that used in pH definition.
In the case of determination of metals detection of the endpoint is mainly based on substances that change color when creating complexes with determined metals. One of these indicators is eriochrome black T, substance that in pH between 7 and 11 is blue when free, and black when forms a complex with metal, other examples are pyrocatechin violet and murexide. It is important that formation constant for these complexes is low enough, so that titrant reacts with complexed ions first.
Indicators used in complexometric titration are to some extent similar to those used in acid-base titrations. Their color changes depending on the concentration of metal ions, just like color of pH indicators changes depending on the H+ concentration. Mechanism of this color change is different, as all complexometric indicators are just complexing agents, changing their color depending on whether they are free in the solution, or ligands in the complex. In most cases they are also weak acids or bases, and quite often their color depends on the solution pH.
For a metal indicator to be useful, several conditions must be met. First of all – its stability constant must be high enough so that the free metal when present in the solution is easily complexed, but lower than the stability of the complex with titrant. Otherwise indicator will be not replaced in the complex and there will be no color change of the solution. Secondly, both indicator complex creation and dissociation reactions must be fast, so that the equilibrium in the solution is achieved almost immediately after titrant addition.
Let’s see how concentration of the free indicator changes in the vicinity of the equivalence point of the calcium titration. Let’s assume we are titrating 0.1M calcium with 0.1M EDTA in the presence of murexide. Complexation constants are
and
What we will calculate is a fraction of free murexide, responsible for violet solution color. We will calculate it for both 99.9 and 100.1% titration.
Probably most popular and universal indicator used in complexometric titrations is Eriochrome Black T. For pH below 6, its color is red, between 7 and 11 – blue, above 12 – yellow-orange. Complexed form is always wine red. Eriochrome Black T solutions are unstable, so it is prepared as a solid, mixed with NaCl (100 mg of indicator ground with
Similar in its properties, but much more stable (solutions can be kept for up to a year) is calmagite. It can be used instead of Eriochrome Black T in most titrations.
Other popular indicators are pyrocatechin violet, murexide and PAN. Also sulfosalicylic acid is used, although it differs from other indicators listed, as it is used only for one cation (Fe3+) and is a one color indicator.
In some cases redox indicators can be used. For example when titrating Fe3+ redox potential of the solution is high in the presence of excess Fe3+, which keeps some of the redox indicators in the oxidized form. After equivalence point concentration of Fe3+ becomes very small and indicator gets reduced. Other interesting approach is used in the case of Al3+ titration. Al3+ can’t be titrated directly with EDTA, as it reacts too slow. To avoid problems back titration is used – and excess EDTA is titrated with Zn2+. Before titration some small amounts of ferrocyanide and ferricyanide are added to the solution, together with diphenylbenzidine. After equivalence point excess Zn2+ precipitates ferrocyanide increasing redox potential of the solution, and diphenylbenzidine gets oxidized.
Finally, there are specific ways of detecting end point for other types of complexometric titrations. For example in the case of cyanides determination with Ag+ solution (Liebig-Dénigès method) reaction that takes place is
Ag+ + 2CN– → Ag(CN)2–
and after equivalence point, when excess Ag+ is added
Ag+ + Ag(CN)2– → 2AgCN
Solid AgCN makes solution turbid and is easy to spot, so there is no need for any other indicator.
EDTA Titration
EDTA titration can be used for direct determination of many metal cations. It reacts directly with Mg, Ca, Zn, Cd, Pb, Cu, Ni, Co, Fe, Bi, Th, Zr and others. With the help of back titration this list can be mad e much longer, as back titration can be used in the cases when the complex is created too slowly (as it happens in the case of Al and Cr), when it is not possible to choose good end point indicator, or when metal could precipitate at as hydroxide at pH required for a direct titration.
Using back titration it is also possible to determine some anions – for example SO42- can be determined by BaSO4 precipitation with the use of BaCl2 and titration of excess barium left in the solution.
Method is not selective – EDTA reacts with almost everything – but careful selection of solution pH allows in some cases to determine one metal in the presence of others. Cations with higher charges (like Bi3+, Fe3+) have much larger stability constants, so they can be titrated at low pH, in the presence of divalent cations (like Ca2+, Mg2+) which will not interfere in this conditions.
Finally, what makes EDTA a convenient reagent is fact, that it always reacts with metals on the 1:1 basis, making calculations easy.
EDTA solution
EDTA solution is not only stable – it can be stored for months – but it can be also prepared without a need of standardization. First, EDTA can be obtained in the form pure enough. Second, after thorough drying its crystallic from has pretty well defined amount of water of crystallization.
Most commonly used solutions are 0.01M (that is 0.01N – regardless of the fact that EDTA has four protons it always reacts with metal cations on a 1:1 base). However, depending on the needs (concentration of metal to be determined) it is possible to prepare and use EDTA solutions of concentrations ranging from 0.1M to 0.001M.
Crystallic EDTA – in the form of either disodium EDTA dihydrate or anhydrous disodium EDTA salt – has to be thoroughly dried out before solutions preparation (see standard substances used in complexometric titrations section). Also note, that solution preparation is time consuming – EDTA dissolves in water very slowly.
0.01M EDTA standardization against metallic magnesium
EDTA can be standardized against many reagents, be it metallic magnesium, calcium carbonate, metallic bismuth and so on. For best results it is good to standardize EDTA solution against the same cation and using the same method as will be later used during sample analysis. Note, that EDTA solution can be prepared without a need for standardization, as EDTA itself can be obtained in form pure enough.
Reaction taking place during titration is
Mg2+ + EDTA4- → MgEDTA2-
Titration is done in pH 10 solution. At this pH about 25% if EDTA is in first protonated form (HEDTA3-), but magnesium complex stability constant is high enough so that we don’t have to worry about.
Procedure to follow:
Weight exactly about 0.24g of metallic Mg ribbon.
Dissolve metal in 25 mL of 1M HCl solution and dilute in volumetric flask to exactly 1L.
Transfer 25 mL of magnesium solution to Erlenmeyer flask.
Add 75 mL of distilled water.
Add 2 mL of pH 10 ammonia buffer solution.
Add a pinch of Eriochrome Black T ground with sodium chloride (100 mg of indicator plus
Titrate with EDTA solution till the color changes to blue.
0.01M EDTA standardization against calcium carbonate
Reaction taking place during titration is
Ca2+ + EDTA4- → CaEDTA2-
Titration is done in about 0.1 sodium hydroxide solution (pH around 13). While this is not our intent when standardizing EDTA solution, at this high pH it is possible to titrate calcium in the presence of magnesium as the latter precipitates in the form of Mg(OH)2.
Procedure to follow:
Weight exactly about 1g of dry calcium carbonate.
Dissolve solid in 25 mL of 1M HCl solution and dilute in volumetric flask to exactly 1L.
Transfer 25 mL of calcium solution to Erlenmeyer flask.
Add 75 mL of distilled water.
Add 10 mL of 1M sodium hydroxide solution.
Add a pinch of murexide ground with sodium chloride (100 mg of indicator plus
Titrate with EDTA solution till the color changes to violet.
To make end point detection easier it is worth to prepare a comparison solution, identical to the titrated one; this way first color change is easier to spot.
Determination of Water Hardness By Complexometric Titration
Hard Water
Hard water is due to metal ions (minerals) that are dissolved in the ground water. These minerals include Ca2+, Mg2+, Fe3+, SO42-, and HCO3–. Our hard water in the southern
The determination of water hardness is a useful test that provides a measure of quality of water for households and industrial uses. Originally, water hardness was defined as the measure of the capacity of the water to precipitate soap. Hard water is not a health hazard. People regularly take calcium supplements. Drinking hard water contributes a small amount of calcium and magnesium toward the total human dietary needs of calcium and magnesium. The National Academy of Science states that consuming extremely hard water could be a major contributor of calcium and magnesium to the diet.
Hard water does cause soap scum, clog pipes and clog boilers.
Soap scum is formed when the calcium ion binds with the soap. This causes an insoluble compound that precipitates to form the scum you see. Soap actually softens hard water by removing the Ca2+ ions from the water.
When hard water is heated, CaCO3 precipitates out, which then clogs pipes and industrial boilers. This leads to malfunction or damage and is expensive to remove.
Water Softeners
If you have hard water you may use a water softener to remove the hardness. Salt is mixed with water. The Na+ ion from the salt replaces the Ca2+ ion, but this causes the water to be too salty for drinking. Water that has been softened should be used only for laundry and bathing.
Types of Hardness
There are two types of water hardness, temporary and permanent.
Temporary Hardness is due to the bicarbonate ion, HCO3–, being present in the water. This type of hardness can be removed by boiling the water to expel the CO2, as indicated by the following equation:
Bicarbonate hardness is classified as temporary hardness.
Permanent hardness is due to the presence of the ions Ca2+, Mg+2, Fe3+ and SO4–. This type of hardness cannot be eliminated by boiling. The water with this type of hardness is said to be permanently hard.
How Hard Is The Water?
The degree of hardness of the water is classified in terms of its calcium carbonate concentration as follows:
|
Hardness rating |
Concentration of Calcium Carbonate (mg/L) |
Concentration of Calcium Carbonate (grains/US gallon) |
|
Soft |
0 to <75 |
0 to <5.2 |
|
Medium hard |
75 to <150 |
5.2 to <10.5 |
|
Hard |
150 to <300 |
10.5 to <21 |
|
Very hard |
300 and greater |
21 and greater |
How to Soften Hard Water
Hard water is high in minerals, usually calcium and magnesium; however, bicarbonate and sulfates levels sometimes contribute to water hardness. Studies have not found any health risks or medical reasons for softening water, but it is an inconvenience nevertheless. Most issues stem from the fact that hard water lowers the effectiveness of all forms of cleaning products and soaps. Hard water often shortens the lifespan of clothing, household appliances and plumbing. There are several ways to soften hard water.
Soften Your Entire Household Water Supply
Determine that your water is hard
· Hard water exists in scattered water supplies throughout the world. Although hard water exists in water sources in all states, the highest concentration of hard water supplies in the
· If you use a public water supply, information about your level of water hardness is available from your local town hall water department.
· This information may also be available from your local town hall if you have private well. Based on your location, your city government may be able to tell you the primary source of water supplies for the area and give you some test results, including water softness levels.
· Bring a water sample to a local water testing lab to test your water hardness or use a commercially sold water hardness test kit if you have a private well and suspect your water is hard.
· Watch the amount of foaming that occurs when you use cleaning products like toothpaste, dish soap, laundry detergent and other household cleaners. If you have to add a lot of soap to the water to work up suds your water is probably hard.
·
Install a mechanical water softener that replaces calcium and magnesium with sodium if your water source is hard.
· Mechanical water softeners prevent lime scale, increase the efficiency of heating your water, and lengthens the life of your clothing and other items that you regularly launder.
· There are several water softeners on the market that range in price, operating costs and effectiveness. You must test the treated water before drinking it.
Install a magnetic water conditioner that alters calcium ions so they are unable to cause lime scale.
· Water conditioners are less expensive and typically cheaper to operate, and the resulting water is safe for drinking.
· Conditioners are not effective on all cases of hard water and do not usually come with softening guarantees. Critics are skeptical that they work at all. The degree of softening varies more than water treated by mechanical softeners.
Soften Water Before Use
· Boil your water before drinking it. You might also want to use boiled water for cleaning your kitchen and bathroom, brushing your teeth, bathing and washing your hair to increase the effectiveness of your cleaning products.
· After boiling water for a few minutes, allow it to cool. Visible lime particles will settle on the surface of the water. Scoop off the top layer of particles and discard them before using the water.
· Alternately, allow the water to sit longer and the particles will fall to the bottom. Scoop out the fresh water carefully so as to not disturb the settled lime particles. Discard the bottom few inches of water where lime particles remain.
·
Soften water using washing soda or lime.
· In the past, households with hard water often softened it by filling large open kegs with water and then adding some washing-soda or lime. The water has to stand for several days and then must be drawn from the top of the kegs.
· This method is not commonly used today because of the amount of time it requires.
·
Add ammonia, borax, lye or washing-soda to the water at the same time you add the soap when doing laundry and other household cleaning.
· These products will not soften the water but are known to prevent the lime from interacting with the soap, helping suds form. Follow package warnings and instructions carefully when using.
· Dissolve
· Dissolve 1/4 tablespoon of lye in one cup of water. Mix the solution into one gallon of water.
·
Put a similar type of filter on your kitchen and bathroom sink faucets to provide softening for water dispensed through the tap. This is the easiest method if you want to filter water for cleaning. Some models have switch-off valves so you can draw water from the tap without running it through the filter if you choose.
Total Water Hardness
Water hardness is a measure of the amount of calcium and magnesium salts dissolved in water. There are no health hazards associated with water hardness, however, hard water causes scale, as well as the reduced lathering of soaps. Hard water should be not used for washing (it reduces effectiveness of detergents) nor in water heaters and kitchen appliances like coffee makers (that can be destroyed by scale). It is also not good for fish tanks. In general, there are many applications where ability to easily determine water hardness is very important.
Complexometric titration is one of the best ways of measuring total water hardness. At pH around 10 EDTA easily reacts with both calcium and magnesium in the same molar ratio (1:1). Stability constant of calcium complex is a little bit higher, so calcium reacts first, magnesium later. Thus, for the end point, we should use the same indicator we use when titrating magnesium – that is Eriochrome Black T. In the case of water that doesn’t contain magnesium at all, to be able to detect end point we should add small amount of magnesium complex MgEDTA2+. Magnesium will be displaced by identical amount of calcium, and it will be titrated later, not changing final result. However, this is a very rare situation.
If solutions contains carbonates, they should be removed as they can interfere with end point detection. To do so we can acidify the solution with hydrochloric acid, boil it, and theeutralize with ammonia. Small excess of ammonia doesn’t hurt, as we finally add ammonia buffer and change of pH by several tenths is not a problem.
Reaction
Reactions taking place during titration are
Ca2+ + EDTA4- → CaEDTA2-
and
Mg2+ + EDTA4- → MgEDTA2-
Sample size
For
End point detection
As it was explained above, calcium is complexed first, so to detect end point we can use indicator used for detection of end point of magnesium titration. That means Eriochrome Black T.
Solutions used
To perform titration we will need titrant –
Procedure
Transfer exactly 50 mL of water to 250 mL Erlenmayer flask.
Acidify the solution with hydrochloric acid.
Bring to boil, cool down.
Alkalize with ammonia.
Filter solution through filter paper.
Add 1 ml of pH 10 ammonia buffer.
Add 3 drops of Eriochrome Black T solution or pinch of Eriochrome Black T ground with NaCl.
Titrate with 0.01M EDTA solution till color changes from violet to blue.
Result calculation
As water hardness is usually reported in terms of mg/L of calcium carbonate (even if water contains both calcium and magnesium), we will use for calculations slightly strange reaction equation:
CaCO3 + EDTA4- → CaEDTA2- + CO32-
That allows direct calculation of calcium carbonate mass for known amount of titrant used.
To calculate water hardness use EBAS – stoichiometry calculator. Download determination of water hardness reaction file, open it with the free trial version of the stoichiometry calculator.
