CHEMICAL ANALYSIS ANALYTICAL TOOL 3 - VOLUMETRIC ANALYSIS (A QUANTITATIVE TECHNIQUE)yr12chemistrycola.wikispaces.com/file/view/Titration... · · 2009-03-23Volumetric analysis is

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CHEMICAL ANALYSIS

ANALYTICAL TOOL 3 - VOLUMETRIC ANALYSIS (A QUANTITATIVE TECHNIQUE)

Volumetric analysis is a technique that is used to determine the concentration of substances in solution. Solids can only be analysed using volumetric analysis techniques if they can be brought into solution. The concentration of the unknown solution whose volume is accurately known is determined by reacting it with a solution whose concentration and volume are accurately known i.e. a standard solution. There are three main types of volumetric analyses: • Acid - Base Titrations

Acid base titrations are used for acid/base reactions which produce a sharp end point. Note: An indicator is always used in acid base titrations.

• Back Titrations

Back titrations are used when a standard titration results in a broad end point. Note: • Back titrations are characterised by a two stage process.

• An indicator is always used in the second stage of the process i.e. the titration.

• Redox Titrations

Redox titrations are used when titrating an oxidant with a reductant. Note: • Indicators are rarely used in such titrations as redox reactions are usually

self indicating.

• There is always an indication of a change in at least one oxidation state during the titration process.


APPARATUS USED IN VOLUMETRIC ANALYSES The equipment used in typical volumetric analyses includes: • The volumetric flask, which is used to make up accurate volumes of solutions eg. the standard solution. • The pipette, which is used to accurately deliver a fixed volume of solution (aliquot). • The burette, which is used to dispense variable volumes (titres) of solutions

accurately.

There are always errors associated with measurements made during experimental work. Typical uncertainties associated with volumetric analysis include: • 20 ml pipette: 0.05 ml± • Burette: 0.02 ml± • 250.0 ml volumetric flask: 0.3 ml± Note: • Although burettes are calibrated in intervals of 0.1 ml , the volume can be estimated to

the nearest 0.02 ml . • If the meniscus lies exactly on a line, it should be recorded to the second

decimal place eg. 15.60 ml


ACID/BASE TITRATIONS METHOD: Step 1: An accurately known volume (aliquot) of the solution whose concentration is to be determined, together with an appropriate indicator is placed into a conical flask using a pipette. This solution is referred to as the “unknown”. Step 2: The standard solution (a solution whose concentration is accurately known) is placed into a calibrated tube known as a burette. This solution is referred to as the “known”. Step 3: An accurately known volume of the standard solution (titre) is delivered until there is a permanent colour change in the conical flask. This process is referred to as a titration.

The burette reading at the point at which stoichiometrically equivalent amounts of acid and base have reacted is referred to as the equivalence point (or stoichiometric point) of the reaction. At this point, neither reactant is present in excess. The reading on the burette at the point at which the indicator permanently changes colour is known as the end point of the reaction. Ideally, the equivalence point and the end point should coincide, but this does not always occur. Indicators must therefore be carefully chosen so that the end point matches the equivalence point as closely as possible. Note: • In some titrations, the “known” is placed into the conical flask and the “unknown” is

placed into the burette. Beware!!! • To minimise errors, titrations should be repeated until three concordant titres are

obtained i.e. three titres that differ by a maximum of 0.10 ml from highest to lowest. QUESTION 1 Explain why the equivalence point and the end point in a titration do not frequently coincide. Solution


QUESTION 2 Which process occurs first in an acid-base titration: the equivalence point or the endpoint? Solution QUESTION 3 The end point was reached when the following volumes of NaOH had been delivered from the burette. Flask 1: ml25.25 Flask 2: ml60.22 Flask 3: ml70.22 Flask 4: ml50.24 Flask 5: ml80.22 Determine the average titre that should be used for the calculations relating to this titration. Solution


PRIMARY STANDARDS

A standard solution is a solution whose concentration is accurately known. Standard solutions may be prepared from primary standards; chemical species that are so pure, that the amount (in mol) of substance can be accurately determined from its mass.

Properties of Ideal Primary Standards • Be readily available in pure form.

• The molecular formula must be known and cannot vary. For example: The species cannot absorb substances from the air, or give off substances to the atmosphere.

• Be easily stored without deteriorating or reacting with the atmosphere.

• Must be soluble under the conditions in which it is to be used.

• Should have a high molecular weight so that weighing errors are minimised.

• Must react rapidly and completely with the analyte.

• Must react stoichiometrically with the analyte.

• Must be selective for the analyte.

PREPARING STANDARD SOLUTIONS

Step 1: Weigh out an accurately known mass of a primary standard.

Step 2: Carefully transfer the weighed mass from the crucible into a volumetric flask. Gently rinse the crucible to ensure that the entire sample is transferred into the flask.

Step 3: Add some distilled water to the volumetric flask and mix the solution carefully.

Step 4: Make the solution up to an accurately known volume.

Step 5: Calculate the concentration.

When a standard solution has been used to determine the concentration of

a solution, we say that the solution has been standadised.


QUESTION 4 Solid sodium hydroxide is not a suitable primary standard, as it readily absorbs moisture from the air (it is deliquescent). It also reacts with 2CO to form carbonates. What procedure would have to be performed before a sodium hydroxide solution prepared from solid NaOH could be used as a standard solution in a titration? Solution


ACID-BASE INDICATORS

An indicator is a weak organic acid that displays different colours at different concentrations of +OH3 in solution. Different indicators display their own characteristic colours across different pH ranges.

COMMON EXAMPLES OF ACID/BASE INDICATORS

(Data Provided on the VCAA Formula Sheet)

Indicator Colour Change

(Acid - Base)

pH Range

Thymol Blue Red to Yellow 1.2 - 2.8 Bromophenol Blue Yellow to Blue 3.0 – 4.6

Methyl Orange Red to Yellow 3.1 – 4.4 Methyl Red Red to Yellow 4.2 – 6.2

Litmus Red to Blue 4.5 – 8.3 Bromothymol Blue Yellow to Blue 6.0 – 7.6

Phenol Red Yellow to Red 6.8 – 8.4 Phenolphthalein Colourless to Red 8.3 – 10.0

Note: Indicators change colour over a pH range,

and not at specific pH values. The ideal indicator is one that changes colour at the pH at which the equivalence point of the reaction occurs. The pH at which the equivalence point occurs depends upon the relative strengths of the reacting species, and depends upon whether the stronger species is acidic or basic in nature. For example:

Acid Base Salt pH at Neutralisation HCl NaOH NaCl 7

2 4H SO KOH 2 4K SO 7

2 3H CO NaOH 2 3Na CO 11

3 4H PO NaOH 3 4Na PO 12

3CH COOH NaOH 3 .CH COO Na 9 HCl 4NH OH 4NH Cl 6


In general: Strong acids and strong bases produce salts with a neutral pH. i.e. 7pH = at 25oC Strong acids and weak bases produce salts with an acidic pH. i.e. 7pH < at 25oC Weak acids and strong bases produce salts with an alkaline pH. i.e. 7pH > at 25oC Weak acids and weak bases produce salts with a neutral pH. i.e. 7pH = at 25oC

As acid-base indicators are very sensitive to small changes in pH, careful choice of indicator will ensure that the end point will closely match the

equivalence point of the reaction.

ACIDITY CONSTANTS, aK , OF SOME WEAK ACIDS

(Data Provided on the VCAA Formula Sheet)

Note: The larger the aK , the stronger the acid species.


QUESTION 5 In an experiment to determine the concentration of ethanoic acid in a sample of vinegar, a standard solution of M100.0 NaOH was titrated against a diluted vinegar solution.

ml00.25 of vinegar was added to a volumetric flask, and water was added to bring the solution to a volume of ml250 . ml00.20 of the diluted vinegar was then transferred into a conical flask containing a few drops of phenolphthalein indicator. It was found that the solution in the conical flask became permanently pink when ml5.22 of NaOH has been delivered from the burette. (a) Calculate the concentration ( M ) of ethanoic acid in the sample of vinegar. (b) Explain why titrations are repeated until three concordant results are obtained. Solution


QUESTION 6 Some students were set the task of determining the concentration of ethanoic acid (acetic acid) in a particular brand of vinegar. An outline of the method they used is given below. 1. A burette is filled with a standard solution of sodium hydroxide. 2. The vinegar is diluted by a factor of 10 in a volumetric flask. A pipette is used to transfer ml00.20 of the diluted vinegar to a conical flask, and a few drops of phenolphthalein indicator is added. 3. The diluted vinegar is titrated with the base. Titrations are repeated until three concordant results are obtained. The equation for the reaction is: )(2)(3)(3)( laqaqaq OHCOONaCHCOOHCHNaOH +→+ (a) Why is the vinegar diluted before titrating? (b) One student’s results are given below. The data shown in the student’s laboratory book was: Concentration of NaOH M11.0 Volume of undiluted vinegar ml00.10 Total volume of diluted vinegar ml00.100 Volume of diluted vinegar used in each titration ml00.20 Average titre of NaOH ml35.15 Based on these results, calculate the concentration, in Lmol / , of ethanoic acid in the undiluted vinegar solution.


(c) What is the approximate pH of the products of this reaction? (d) What are some of the errors involved in such a titration? (e) Why doesn’t the addition of an indicator affect stoichiometric calculations?


QUESTION 7 A g0.15 antacid tablet containing magnesium carbonate was crushed and dissolved to produce a volume of ml00.200 . A ml00.20 aliquot was removed and titrated against a

M00.1 hydrochloric acid solution using methyl orange as the indicator. If ml75.27 of hydrochloric acid was required for the complete neutralisation of magnesium carbonate, calculate the percentage by mass of magnesium carbonate in the antacid tablet.

Solution


NEUTRALISATION CURVES A neutralisation curve describes the changes in pH that occur during an acid-base titration. The general shape of the titration curve will depend upon which species is being added from the burette. For example: If acid is being added from the burette to a base in the conical flask, the pH in the flask will decrease. If base is being added from the burette to an acid in the conical flask, the pH in the flask will increase.

The start and end points on the curve will depend upon the

strengths of the acids and bases. • If an acid is present in the conical flask: The stronger the acid, the lower the initial pH. • If a base is present in the conical flask: The stronger the base, the higher the initial pH.

pH

Volume of Acid

pH

Volume of Base


pH

3

9

20 20.1

IMPORTANT NOTE: To minimise the titration error (the difference in the volume between the endpoint and equivalence point of the reaction), the pH of the solution in the conical flask must change sharply by several units and across a small volume, so that the end point is sharp.

CHOOSING AN INDICATOR FOR A TITRATION

Any indicator that changes colour across the vertical portion of the neutralisation curve will accurately identify the equivalence point of the reaction.

Volume (ml)

As the pH increases markedly across a very small volume, the end

point is sharp.

pH

Volume (ml)

5

7

20 27

As the volume required for a marked change in the pH is high, we say that

the end point is broad.

• Equivalence Point

• Equivalence Point


TITRATION OF A STRONG ACID AGAINST A STRONG BASE

Example: M1.0 HCl and M1.0 NaOH Burette ( NaOH ) Conical Flask ( HCI )

The changes in the pH of the solution in the conical flask

may be represented as follows: • The initial pH is low as the conical flask contains HCl only. • As NaOH is added, the pH increases. The base does not affect the pH of the solution

greatly, until the equivalence point is imminent. • At the equivalence point (•), there is a sharp change in pH of approximately 6 units,

when only one or two drops of NaOH is added. • Any indicator that changes colour across the vertical portion of the graph (between pH

3.5 and 9.5) will accurately identify the equivalence point of the reaction. • As there is a marked pH change at the equivalence point (and across a small volume),

we say that the end point is sharp.


TITRATION OF A WEAK ACID AGAINST A STRONG BASE

Example: M1.0 COOHCH3 and M1.0 NaOH Burette ( NaOH ) Conical Flask ( COOHCH3 )

The changes in the pH of the solution in the conical flask may be represented as follows:

• The initial pH is low as the conical flask contains COOHCH3 only. • The initial pH is higher than that of M1.0 HCl , as COOHCH3 is a weaker acid. • As NaOH is added from the burette, the pH rises quickly ( NaOH is a stronger

species than COOHCH3 ). • The pH rises rapidly around the equivalence point of the reaction. • The equivalence point occurs at a pH greater than 7 units, as the acetate ion produced

functions as a strong base; reacting with water to produce hydroxide ions. • Any indicator that changes colour between pH 6.5 and 11 will accurately identify the

equivalence point. • As there is a marked pH change at the equivalence point (and across a small volume),



TITRATION OF A STRONG ACID AGAINST A WEAK BASE

Example: M1.0 HCI and M1.0 3NH Burette ( 3NH ) Conical Flask ( HCl )

The changes in pH of the solution in the conical flask may be represented as follows:

• The initial pH is low as the conical flask contains HCl only. • The initial pH is lower than that of M1.0 COOHCH3 as HCl is a stronger acid. • As 3NH is added, the pH slowly increases. The change in pH before the equivalence

point is NOT as dramatic as when NaOH is added, as 3NH is a weaker base. • At the equivalence point, there is a sharp change in pH. • The equivalence point occurs at a pH of less than 7 units as a strong acid is produced,

which reacts with water to produce +H ions. • As there is a marked pH change at the equivalence point (and across a small volume),



TITRATION OF A WEAK ACID AGAINST A WEAK BASE

Example: M1.0 COOHCH3 and M1.0 3NH Burette ( 3NH ) Conical Flask ( COOHCH3 )

The changes in pH of the solution in the conical flask may be represented as follows:

• The initial pH is low as the conical flask contains COOHCH3 only. • The pH rises as 3NH is added to the conical flask. • At the equivalence point, there is a GRADUAL change in pH. • The equivalence point occurs at approximately 7 pH units. • No indicator will dramatically change colour across a small volume of base. The end

point is BROAD and will not coincide with the equivalence point of the reaction.

Weak acid/base titrations are generally not performed as a broad end point is obtained. In cases such as these, concentrations of solutions

are determined using back titrations.


QUESTION 8 Benzoic acid, a weak acid, dissociates according to the following equation:

+− +⇔ )()(56)(56 aqaqaq HCOOHCCOOHHC Benzoic acid is used as a preservative, particularly in soft drinks. The diagram below shows the pH changes which occur, when ml0.20 of M100.0 benzoic acid is titrated with

M200.0 sodium hydroxide solution. (a) What volume of sodium hydroxide solution has been added when point C has been reached?


(b) The pH range for the colour change of methyl orange indicator is 3.1 (red) to 4.4 (yellow). For phenolphthalein indicator, this range is 8.3 (colourless) to 10.0 (red). Which indicator is suitable for determining the equivalence point of this titration? Explain your answer. (c) The pipette used was initially rinsed with water and the wet pipette was then used to transfer the benzoic acid solution. What experimental step was left out? What effect did this omission have on the measured concentration of benzoic acid, compared with the actual concentration ( M100.0 )?

(d) ml0.20 of M100.0 hydrochloric acid is titrated with M200.0 sodium hydroxide. (i) How would the pH and volume of solution at the equivalence point compare with the equivalence point of the benzoic acid solution?


(ii) Sketch the graph of the changes in pH that will occur during this titration on the given set of axes.

ERRORS ASSOCIATED WITH INCORRECT EXPERIMENTAL PROCEDURES

EQUIPMENT BASED ERRORS

The burette and pipette must be rinsed with the solutions that are to deliver.

The conical flask must be rinsed with distilled water.

If equipment is incorrectly rinsed immediately prior to use, the calculated

concentration of unknown will not accurately represent the true value.


QUESTION 9 Consider a standard acid base titration where the known is delivered from the burette. If the burette were rinsed with water prior to use, would you expect the calculated concentration of the unknown solution to be higher, lower or equal to the true value? Solution The amount in mole of known in the burette will be ___________________ the expected value. A ___________________ volume will need to be delivered from the burette to neutralise the solution in the conical flask. The calculated mole of known will therefore be ___________________ the expected value. The calculated mole of unknown will therefore be ___________________ the expected value. The calculated concentration of unknown will therefore be ___________________ the expected value.


QUESTION 10 Consider a standard acid base titration where the known is delivered from the burette. If the pipette was rinsed with water prior to use, would you expect the calculated concentration of the unknown solution to be higher, lower or equal to the true value? Solution The amount in mole of unknown in the pipette will be ___________________ the expected value. The amount in mole of unknown in the conical flask will be ___________________ the expected value. A ___________________ volume will need to be delivered from the burette to neutralise the solution in the conical flask. The calculated mole of known will therefore be ___________________ the expected value. The calculated mole of unknown will therefore be ___________________ the expected value. The calculated concentration of unknown will therefore be ___________________ the expected value.


QUESTION 11 Consider a standard acid base titration where the known is delivered from the burette. If the conical flask were rinsed with the solution it was to contain immediately before use, would you expect the calculated concentration of the unknown solution to be higher, lower or equal to the true value? Solution The amount in mole of unknown in the conical flask will be ___________________ the expected value. A ___________________ volume will need to be delivered from the burette to neutralise the solution in the conical flask. The calculated mole of known will therefore be ___________________ the expected value. The calculated mole of unknown will therefore be ___________________ the expected value. The calculated concentration of unknown will therefore be ___________________ the expected value.


ERRORS INVOLVING PRIMARY STANDARDS The molecular formula of the primary standard must be known and cannot vary. For example: The species cannot absorb substances from the air, or give off substances to the atmosphere.

If the species absorbs substances or gives off substances, the calculated mass requirement will be incorrect and therefore, the calculated concentration of unknown will not accurately represent the true value.

QUESTION 12 Consider a standard acid base titration where the known is delivered from the burette. Solid sodium hydroxide is not a suitable primary standard, as it readily absorbs moisture from the air (it is deliquescent). It also reacts with 2CO to form carbonates. If solid sodium hydroxide was used to prepare a standard solution, would you expect the calculated concentration of the unknown solution to be: (a) Lower (b) Equal to (c) Higher than the true results? Explain. Solution The weighed mass will contain___________________ sodium hydroxide than expected. Therefore, the concentration of the standard solution will be ___________________ expected. The amount in mole of known in the burette will be ___________________ expected A ___________________ volume will need to be delivered from the burette to neutralise the solution in the conical flask. The calculated mole of known will therefore be ___________________ the expected value. The calculated mole of unknown will therefore be ___________________ the expected value. The calculated concentration of unknown will therefore be ___________________ the expected value.


QUESTION 13 Consider a standard acid base titration where the known is delivered from the burette. Hydrated 32CONa is not a good primary standard because it is efflorescent i.e. it gives off water to the atmosphere. If hydrated 32CONa was used as a primary standard, would you expect the concentration of the unknown solution to be (a) Equal to (b) Greater than (c) Less than the true value? Explain. Solution The weighed mass will contain ___________________ sodium carbonate than expected. Therefore, the concentration of the standard solution will be ___________________ expected. The amount in mole of known in the burette will be ___________________ expected A ___________________ volume will need to be delivered from the burette to neutralise the solution in the conical flask. The calculated mole of known will therefore be ___________________ the expected value. The calculated mole of unknown will therefore be ___________________ the expected value. The calculated concentration of unknown will therefore be ___________________ the expected value.


ERRORS INVOLVING INDICATOR CHOICE Ideally, the equivalence point and the end point should coincide, but this does not always occur. Indicators must therefore be carefully chosen so that the end point matches the equivalence point as closely as possible. Any indicator that changes colour across the vertical portion of the neutralisation curve will accurately identify the equivalence point of the reaction.

If the indicator is chosen such that the end point and equivalence point do not coincide, the calculated concentration of unknown will not accurately represent the true value.

QUESTION 14 A student carried out an experiment to determine the concentration of an acetic acid solution (concentration: approximately M1.0 ), by titrating it against a M1.0 solution of NaOH . How would the calculated concentration compare to the true value if methyl orange were used as the indicator for this reaction? Note: Methyl Orange changes from red to yellow between pH 3.1 and 4.4 Solution


QUESTION 15 A student carried out an experiment to determine the concentration of a HCl solution (concentration: approximately M1.0 ), by titrating it against a M1.0 solution of 3NH . How would the calculated concentration compare to the true value if phenolphthalein was used as the indicator for this reaction? Note: Phenolphthalein changes from colourless to red between pH 8.3 and 10.0 Solution


EFFECTS OF ERRORS — SUMMARY The following table describes the effects that incorrect techniques will have on the calculated concentration of the unknown solution when the unknown is located in: (a) the conical flask (b) the burette.

Error

Concentration of

Unknown (When unknown is in

the conical flask)

Concentration of Unknown

(When unknown is in the burette)

Burette rinsed with water

Pipette rinsed with water

Conical flask rinsed with solution

End point occurs before equivalence point

End point occurs after equivalence point

Primary standard gives off substances to atmosphere

Primary standard absorbs substances from the atmosphere


QUESTION 16 g427.1 of anhydrous sodium carbonate, 32CONa , was dissolved in a little water in a

300.250 cm volumetric flask, and the solution was then made up to the mark with water. After thorough mixing, 300.20 cm of this solution was pipetted into a flask with a few drops of methyl orange. The solution was yellow. The hydrochloric acid was added slowly from a burette until the colour changed to orange. The volume of acid required was 364.21 cm . (a) Calculate the molarity of the HCl solution.

)(32)()()(32 22 aqaqaqaq COHNaClHClCONa +→+ Known: 32CONa

molmlinn 0135.0106427.1)250( ==

molmlinn 00108.0135.025020)20( =×=

Unknown: HCl

molCONanHCln 00216.0)(2)( 32 =×=

MVnC 09982.0

02164.000216.0

===

V = 250 ml Na2CO3

Remove 20 ml

HCl

Vol = 20 ml

Na2CO3

gm 427.1= 106=M

moln 0135.0106427.1

==


(b) Why is the value obtained a slight overestimate of the true value? (c) A student accidentally made the sodium carbonate solution to ml300 instead of the required ml250 . Would you expect this student's calculated molarity to be higher, less than or equal to the true value? Provide a brief explanation to justify your answer.

(d) Hydrated sodium carbonate is not a good primary standard, as it readily gives off water to the atmosphere. If hydrated sodium carbonate was used as the primary standard in this titration, would you expect the calculated molarity to be higher, less than or equal to the true value? Provide a brief explanation to justify your answer.


(e) If the burette was rinsed with water immediately prior to use, how would the calculated molarity differ from the true value? Provide a brief explanation to justify your answer. (f) If the pipette was rinsed with water immediately prior to use, how would the calculated molarity differ from the true value? Provide a brief explanation to justify your answer. (g) If the conical flask was rinsed with the solution it is to contain immediately prior to use, how would the calculated molarity differ from the true value? Provide a brief explanation to justify your answer.

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CHEMICAL ANALYSIS ANALYTICAL TOOL 3 - VOLUMETRIC ANALYSIS (A QUANTITATIVE TECHNIQUE)yr12chemistrycola.wikispaces.com/file/view/Titration... · · 2009-03-23Volumetric analysis is