Modern Analytical Cheymistry - Chapter 9 pdf

For the reaction of astrong base with a strong acid the only equilibrium reaction of importance is The first task in constructing the titration curve is to calculate the volume of NaOHne

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Titrimetric Methods of Analysis

T itrimetry, in which we measure the volume of a reagent reacting

stoichiometrically with the analyte, first appeared as an analytical

method in the early eighteenth century Unlike gravimetry, titrimetry

initially did not receive wide acceptance as an analytical technique.

Many prominent late-nineteenth century analytical chemists preferred

gravimetry over titrimetry and few of the standard texts from that era

include titrimetric methods By the early twentieth century, however,

titrimetry began to replace gravimetry as the most commonly used

analytical method.

Interestingly, precipitation gravimetry developed in the absence of

a theory of precipitation The relationship between the precipitate’s

mass and the mass of analyte, called a gravimetric factor, was

determined experimentally by taking known masses of analyte (an

external standardization) Gravimetric factors could not be calculated

using the precipitation reaction’s stoichiometry because chemical

formulas and atomic weights were not yet available! Unlike gravimetry,

the growth and acceptance of titrimetry required a deeper

understanding of stoichiometry, thermodynamics, and chemical

equilibria By the early twentieth century the accuracy and precision of

titrimetric methods were comparable to that of gravimetry,

establishing titrimetry as an accepted analytical technique.

9A Overview of Titrimetry

Titrimetric methods are classified into four groups based on the type of reaction

in-volved These groups are acid–base titrations, in which an acidic or basic titrant

re-acts with an analyte that is a base or an acid; complexometric titrations involving ametal–ligand complexation reaction; redox titrations, where the titrant is an oxidiz-ing or reducing agent; and precipitation titrations, in which the analyte and titrantreact to form a precipitate Despite the difference in chemistry, all titrations shareseveral common features, providing the focus for this section

9A.1 Equivalence Points and End Points

For a titration to be accurate we must add a stoichiometrically equivalent amount

of titrant to a solution containing the analyte We call this stoichiometric mixture

the equivalence point Unlike precipitation gravimetry, where the precipitant is

added in excess, determining the exact volume of titrant needed to reach the

equiv-alence point is essential The product of the equivequiv-alence point volume, Veq, and the

titrant’s concentration, CT, gives the moles of titrant reacting with the analyte

reach an end point of our choosing Often this end point is indicated by a change in

the color of a substance added to the solution containing the analyte Such

sub-stances are known as indicators The difference between the end point volume and the equivalence point volume is a determinate method error, often called the titra- tion error If the end point and equivalence point volumes coincide closely, then

the titration error is insignificant and can be safely ignored Clearly, selecting an propriate end point is critical if a titrimetric method is to give accurate results

ap-9A.2 Volume as a Signal*

Almost any chemical reaction can serve as a titrimetric method provided thatthree conditions are met The first condition is that all reactions involving thetitrant and analyte must be of known stoichiometry If this is not the case, thenthe moles of titrant used in reaching the end point cannot tell us how much ana-lyte is in our sample Second, the titration reaction must occur rapidly If we addtitrant at a rate that is faster than the reaction’s rate, then the end point will ex-ceed the equivalence point by a significant amount Finally, a suitable methodmust be available for determining the end point with an acceptable level of accu-racy These are significant limitations and, for this reason, several titration strate-gies are commonly used

A simple example of a titration is an analysis for Ag+using thiocyanate, SCN–,

as a titrant

Ag+(aq) + SCN–(aq) tAgSCN(s)

equivalence point

The point in a titration where

stoichiometrically equivalent amounts of

analyte and titrant react.

end point

The point in a titration where we stop

adding titrant.

indicator

A colored compound whose change in

color signals the end point of a titration.

titration error

The determinate error in a titration due

to the difference between the end point

and the equivalence point.

titrant

The reagent added to a solution

containing the analyte and whose

volume is the signal.

*Instead of measuring the titrant’s volume we also can measure its mass Since the titrant’s density is a measure of its

titrimetry

Any method in which volume is the

signal.

This reaction occurs quickly and is of known stoichiometry A titrant of SCN–is

easily prepared using KSCN To indicate the titration’s end point we add a small

amount of Fe3+to the solution containing the analyte The formation of the

red-colored Fe(SCN)2+complex signals the end point This is an example of a direct

titration since the titrant reacts with the analyte

If the titration reaction is too slow, a suitable indicator is not available, or there

is no useful direct titration reaction, then an indirect analysis may be possible

Sup-pose you wish to determine the concentration of formaldehyde, H2CO, in an

aque-ous solution The oxidation of H2CO by I3

H2CO(aq) + 3OH–(aq) + I3(aq) tHCO2(aq) + 3I–(aq) + 2H2O(l)

is a useful reaction, except that it is too slow for a direct titration If we add a known

amount of I3 , such that it is in excess, we can allow the reaction to go to

comple-tion The I3 remaining can then be titrated with thiosulfate, S2O32–

I3(aq) + 2S2O32–(aq) tS4O62–(aq) + 3I–(aq)

This type of titration is called a back titration.

Calcium ion plays an important role in many aqueous environmental systems

A useful direct analysis takes advantage of its reaction with the ligand

ethylenedi-aminetetraacetic acid (EDTA), which we will represent as Y4–

Ca2+(aq) + Y4–(aq) tCaY2–(aq)

Unfortunately, it often happens that there is no suitable indicator for this direct

titration Reacting Ca2+with an excess of the Mg2+–EDTA complex

Ca2+(aq) + MgY2–(aq) tCaY2–(aq) + Mg2+(aq)

releases an equivalent amount of Mg2+ Titrating the released Mg2+with EDTA

Mg2+(aq) + Y4–(aq) tMgY2–(aq)

gives a suitable end point The amount of Mg2+titrated provides an indirect

mea-sure of the amount of Ca2+in the original sample Since the analyte displaces a

species that is then titrated, we call this a displacement titration.

When a suitable reaction involving the analyte does not exist it may be possible

to generate a species that is easily titrated For example, the sulfur content of coal can

be determined by using a combustion reaction to convert sulfur to sulfur dioxide

S(s) + O2(g)→SO2(g)

Passing the SO2through an aqueous solution of hydrogen peroxide, H2O2,

SO2(g) + H2O2(aq)→H2SO4(aq)

produces sulfuric acid, which we can titrate with NaOH,

H2SO4(aq) + 2OH–(aq) tSO42–(aq) + 2H2O(l)providing an indirect determination of sulfur

9A.3 Titration Curves

To find the end point we monitor some property of the titration reaction that has a

well-defined value at the equivalence point For example, the equivalence point for

a titration of HCl with NaOH occurs at a pH of 7.0 We can find the end point,

back titration

A titration in which a reagent is added to

a solution containing the analyte, and the excess reagent remaining after its reaction with the analyte is determined

by a titration.

displacement titration

A titration in which the analyte displaces

a species, usually from a complex, and the amount of the displaced species is determined by a titration.

A titration curve provides us with a visual picture of how a property, such as

pH, changes as we add titrant (Figure 9.1) We can measure this titration curve perimentally by suspending a pH electrode in the solution containing the analyte,monitoring the pH as titrant is added As we will see later, we can also calculate theexpected titration curve by considering the reactions responsible for the change in

ex-pH However we arrive at the titration curve, we may use it to evaluate an tor’s likely titration error For example, the titration curve in Figure 9.1 shows usthat an end point pH of 6.8 produces a small titration error Stopping the titration

indica-at an end point pH of 11.6, on the other hand, gives an unacceptably large titrindica-ationerror

The titration curve in Figure 9.1 is not unique to an acid–base titration Any titration curve that follows the change in concentration of a species in thetitration reaction (plotted logarithmically) as a function of the volume of titranthas the same general sigmoidal shape Several additional examples are shown inFigure 9.2

Concentration is not the only property that may be used to construct a titrationcurve Other parameters, such as temperature or the absorbance of light, may beused if they show a significant change in value at the equivalence point Many titra-tion reactions, for example, are exothermic As the titrant and analyte react, thetemperature of the system steadily increases Once the titration is complete, furtheradditions of titrant do not produce as exothermic a response, and the change intemperature levels off A typical titration curve of temperature versus volume oftitrant is shown in Figure 9.3 The titration curve contains two linear segments, theintersection of which marks the equivalence point

Volume NaOH (mL)

0.00 0.00

4.00 6.00

8.00

Equivalence point 10.00

12.00 14.00

10.00 20.00 30.00 50.00

2.00

40.00

titration curve

A graph showing the progress of a

titration as a function of the volume of

titrant added.

Figure 9.2

Examples of titration curves for (a) a

complexation titration, (b) a redox

titration, and (c) a precipitation

titration.

9A.4 The Buret

The only essential piece of equipment for an acid–base titration is a means for

deliv-ering the titrant to the solution containing the analyte The most common method

for delivering the titrant is a buret (Figure 9.4) A buret is a long, narrow tube with

graduated markings, and a stopcock for dispensing the titrant Using a buret with a

small internal diameter provides a better defined meniscus, making it easier to read

the buret’s volume precisely Burets are available in a variety of sizes and tolerances

16.0 18.0

10.00 20.00 30.00 50.00

2.0

40.00

14.0 12.0

10.00 20.00 30.00 40.00 50.00 6.0

1.600 1.800

0.200

80

1.400 1.200

Titrations may be automated using a pump to deliver the titrant at a constantflow rate, and a solenoid valve to control the flow (Figure 9.5) The volume oftitrant delivered is determined by multiplying the flow rate by the elapsed time Au-tomated titrations offer the additional advantage of using a microcomputer for datastorage and analysis.

9B Titrations Based on Acid–Base ReactionsThe earliest acid–base titrations involved the determination of the acidity or alka-

linity of solutions, and the purity of carbonates and alkaline earth oxides Before

1800, acid–base titrations were conducted using H2SO4, HCl, and HNO3as acidictitrants, and K2CO3and Na2CO3as basic titrants End points were determinedusing visual indicators such as litmus, which is red in acidic solutions and blue inbasic solutions, or by observing the cessation of CO2effervescence when neutraliz-ing CO32– The accuracy of an acid–base titration was limited by the usefulness ofthe indicator and by the lack of a strong base titrant for the analysis of weak acids.The utility of acid–base titrimetry improved when NaOH was first introduced

as a strong base titrant in 1846 In addition, progress in synthesizing organic dyesled to the development of many new indicators Phenolphthalein was first synthe-sized by Bayer in 1871 and used as a visual indicator for acid–base titrations in

1877 Other indicators, such as methyl orange, soon followed Despite the ing availability of indicators, the absence of a theory of acid–base reactivity made se-lecting a proper indicator difficult

increas-Developments in equilibrium theory in the late nineteenth century led to nificant improvements in the theoretical understanding of acid–base chemistry and,

sig-acid–base titration

A titration in which the reaction between

the analyte and titrant is an acid–base

reaction.

Figure 9.5

Typical instrumentation for performing an automatic titration.

Courtesy of Fisher Scientific.

in turn, of acid–base titrimetry Sørenson’s establishment of the pH scale in 1909

provided a rigorous means for comparing visual indicators The determination of

acid–base dissociation constants made the calculation of theoretical titration curves

possible, as outlined by Bjerrum in 1914 For the first time a rational method

ex-isted for selecting visual indicators, establishing acid–base titrimetry as a useful

al-ternative to gravimetry

9B.1 Acid–Base Titration Curves

In the overview to this chapter we noted that the experimentally determined end

point should coincide with the titration’s equivalence point For an acid–base

titra-tion, the equivalence point is characterized by a pH level that is a function of the

acid–base strengths and concentrations of the analyte and titrant The pH at the end

point, however, may or may not correspond to the pH at the equivalence point To

understand the relationship between end points and equivalence points we must

know how the pH changes during a titration In this section we will learn how to

construct titration curves for several important types of acid–base titrations Our

approach will make use of the equilibrium calculations described in Chapter 6 Wealso will learn how to sketch a good approximation to any titration curve using only

a limited number of simple calculations

Titrating Strong Acids and Strong Bases For our first titration curve let’s considerthe titration of 50.0 mL of 0.100 M HCl with 0.200 M NaOH For the reaction of astrong base with a strong acid the only equilibrium reaction of importance is

The first task in constructing the titration curve is to calculate the volume of NaOHneeded to reach the equivalence point At the equivalence point we know from re-action 9.1 that

Moles HCl = moles NaOHor

MaVa= MbVb

where the subscript ‘a’ indicates the acid, HCl, and the subscript ‘b’ indicates thebase, NaOH The volume of NaOH needed to reach the equivalence point, there-fore, is

Before the equivalence point, HCl is present in excess and the pH is determined bythe concentration of excess HCl Initially the solution is 0.100 M in HCl, which,since HCl is a strong acid, means that the pH is

pH = –log[H3O+] = –log[HCl] = –log(0.100) = 1.00

The equilibrium constant for reaction 9.1 is (Kw)–1, or 1.00×1014 Since this is such

a large value we can treat reaction 9.1 as though it goes to completion After adding10.0 mL of NaOH, therefore, the concentration of excess HCl is

giving a pH of 1.30

At the equivalence point the moles of HCl and the moles of NaOH are equal.Since neither the acid nor the base is in excess, the pH is determined by the dissoci-ation of water

Kw= 1.00×10–14= [H3O+][OH–] = [H3O+]2

[H3O+] = 1.00×10–7MThus, the pH at the equivalence point is 7.00

Finally, for volumes of NaOH greater than the equivalence point volume, the

pH is determined by the concentration of excess OH– For example, after adding30.0 mL of titrant the concentration of OH–is

To find the concentration of H3O+, we use the Kwexpression

giving a pH of 12.10 Table 9.2 and Figure 9.1 show additional results for this

titra-tion curve Calculating the titratitra-tion curve for the titratitra-tion of a strong base with a

strong acid is handled in the same manner, except that the strong base is in excess

before the equivalence point and the strong acid is in excess after the equivalence

point

Titrating a Weak Acid with a Strong Base For this example let’s consider the

titra-tion of 50.0 mL of 0.100 M acetic acid, CH3COOH, with 0.100 M NaOH Again, we

start by calculating the volume of NaOH needed to reach the equivalence point;

Before adding any NaOH the pH is that for a solution of 0.100 M acetic acid Since acetic acid is a weak acid, we calculate the pH using the method outlined inChapter 6.

CH3COOH(aq) + H2O(l) tH3O+(aq) + CH3COO–(aq)

x = [H3O+] = 1.32×10–3

At the beginning of the titration the pH is 2.88

Adding NaOH converts a portion of the acetic acid to its conjugate base

CH3COOH(aq) + OH–(aq) tH2O(l) + CH3COO–(aq) 9.2

Any solution containing comparable amounts of a weak acid, HA, and its conjugateweak base, A–, is a buffer As we learned in Chapter 6, we can calculate the pH of abuffer using the Henderson–Hasselbalch equation

The equilibrium constant for reaction 9.2 is large (K = Ka/Kw= 1.75×109), so wecan treat the reaction as one that goes to completion Before the equivalence point,the concentration of unreacted acetic acid is

and the concentration of acetate is

For example, after adding 10.0 mL of NaOH the concentrations of CH3COOH and

CH3COO–are

giving a pH of

A similar calculation shows that the pH after adding 20.0 mL of NaOH is 4.58

At the equivalence point, the moles of acetic acid initially present and the moles

of NaOH added are identical Since their reaction effectively proceeds to tion, the predominate ion in solution is CH3COO–, which is a weak base To calcu-late the pH we first determine the concentration of CH3COO–

comple-The pH is then calculated as shown in Chapter 6 for a weak base

CH3COO–(aq) + H2O(l) tOH–(aq) + CH3COOH(aq)

x = [OH–] = 5.34×10–6MThe concentration of H3O+, therefore, is 1.87×10–9, or a pH of 8.73

After the equivalence point NaOH is present in excess, and the pH is

deter-mined in the same manner as in the titration of a strong acid with a strong base For

example, after adding 60.0 mL of NaOH, the concentration of OH–is

giving a pH of 11.96 Table 9.3 and Figure 9.6 show additional results for this

titra-tion The calculations for the titration of a weak base with a strong acid are handled

in a similar manner except that the initial pH is determined by the weak base, the

pH at the equivalence point by its conjugate weak acid, and the pH after the

equiva-lence point by the concentration of excess strong acid

4.00 6.00 8.00 10.00 12.00 14.00

2.00

80

Table 9.3 Data for Titration of 50.0 mL of

0.100 M Acetic Acid with 0.100 M NaOH

The approach that we have worked out for the titration of a monoprotic weakacid with a strong base can be extended to reactions involving multiprotic acids orbases and mixtures of acids or bases As the complexity of the titration increases,however, the necessary calculations become more time-consuming Not surpris-ingly, a variety of algebraic1and computer spreadsheet2approaches have been de-scribed to aid in constructing titration curves.

Sketching an Acid–Base Titration Curve To evaluate the relationship between anequivalence point and an end point, we only need to construct a reasonable approx-imation to the titration curve In this section we demonstrate a simple method forsketching any acid–base titration curve Our goal is to sketch the titration curvequickly, using as few calculations as possible

To quickly sketch a titration curve we take advantage of the following tion Except for the initial pH and the pH at the equivalence point, the pH at anypoint of a titration curve is determined by either an excess of strong acid or strongbase, or by a buffer consisting of a weak acid and its conjugate weak base As wehave seen in the preceding sections, calculating the pH of a solution containing ex-cess strong acid or strong base is straightforward

observa-We can easily calculate the pH of a buffer using the Henderson–Hasselbalch tion We can avoid this calculation, however, if we make the following assumption.You may recall that in Chapter 6 we stated that a buffer operates over a pH range ex-tending approximately ±1 pH units on either side of the buffer’s pKa The pH is at the

equa-lower end of this range, pH = pKa– 1, when the weak acid’s concentration is mately ten times greater than that of its conjugate weak base Conversely, the buffer’s

approxi-pH is at its upper limit, approxi-pH = pKa+ 1, when the concentration of weak acid is tentimes less than that of its conjugate weak base When titrating a weak acid or weakbase, therefore, the buffer region spans a range of volumes from approximately 10% ofthe equivalence point volume to approximately 90% of the equivalence point volume.*Our strategy for quickly sketching a titration curve is simple We begin by draw-

ing our axes, placing pH on the y-axis and volume of titrant on the x-axis After

calcu-lating the volume of titrant needed to reach the equivalence point, we draw a vertical

line that intersects the x-axis at this volume Next, we determine the pH for two

vol-umes before the equivalence point and for two volvol-umes after the equivalence point Tosave time we only calculate pH values when the pH is determined by excess strong acid

or strong base For weak acids or bases we use the limits of their buffer region to mate the two points Straight lines are drawn through each pair of points, with eachline intersecting the vertical line representing the equivalence point volume Finally, asmooth curve is drawn connecting the three straight-line segments Example 9.1 illus-trates this approach for the titration of a weak acid with a strong base

esti-EXAMPLE 9.1

Sketch the titration curve for the titration of 50.0 mL of 0.100 M acetic acidwith 0.100 M NaOH This is the same titration for which we previouslycalculated the titration curve (Table 9.3 and Figure 9.6)

SOLUTION

We begin by drawing the axes for the titration curve (Figure 9.7a) We have alreadyshown that the volume of NaOH needed to reach the equivalence point is 50 mL,

so we draw a vertical line intersecting the x-axis at this volume (Figure 9.7b).

*Question 4 in the end-of-chapter problems asks you to consider why these pH limits correspond to approximately

Volume of titrant

Percent titrated

0.00 0.0

4.0 6.0 8.0 10.0 12.0

50.00

50

100.00 2.0

4.0 6.0 8.0 10.0 12.0

50.00

50

100.00 2.0

4.0 6.0 8.0 10.0 12.0

50.00

50

100.00 2.0

14.0

Figure 9.7

How to sketch an acid–base titration curve; see text for explanation.

Figure 9.8

Sketches of titration curves for (a) 50.00 mL

of 0.0500 M diprotic weak acid (pKa1= 3,

pKa2 = 7) with 0.100 M strong base; and

(b) 50.00 mL of a mixture of weak acids

consisting of 0.075 M HA (pKa,HA = 3) and

0.025 M HB (pKa,HB= 7) with 0.100 M

strong base The points used to sketch the

titration curves are indicated by the dots (•).

Equivalence points are indicated by the

arrows.

Before the equivalence point the titrant is the limiting reagent, and the pH

is controlled by a buffer consisting of unreacted acetic acid and its conjugateweak base, acetate The pH limits for the buffer region are plotted by

superimposing the ladder diagram for acetic acid on the y-axis (Figure 9.7c)

and adding the appropriate points at 10% (5.0 mL) and 90% (45.0 mL) of theequivalence point volume

After the equivalence point the pH is controlled by the concentration ofexcess NaOH Again, we have already done this calculation Using values fromTable 9.3, we plot two additional points

An approximate sketch of the titration curve is completed by drawingseparate straight lines through the two points in the buffer region and the twopoints in the excess titrant region (Figure 9.7e) Finally, a smooth curve isdrawn connecting the three straight-line segments (Figure 9.7f)

This approach can be used to sketch titration curves for other acid–base tions including those involving polyprotic weak acids and bases or mixtures of weakacids and bases (Figure 9.8) Figure 9.8a, for example, shows the titration curvewhen titrating a diprotic weak acid, H2A, with a strong base Since the analyte is

4.0 6.0 8.0 10.0 12.0

50.00

50

100.00 2.0

4.0 6.0 8.0 10.0 12.0

50.00

50

100.00 2.0

14.0

diprotic there are two equivalence points, each requiring the same volume of

titrant Before the first equivalence point the pH is controlled by a buffer consisting

of H2A and HA–, and the HA–/A2–buffer determines the pH between the two

equiv-alence points After the second equivequiv-alence point, the pH reflects the concentration

of the excess strong base titrant

Figure 9.8b shows a titration curve for a mixture consisting of two weak acids:

HA and HB Again, there are two equivalence points In this case, however, the

equivalence points do not require the same volume of titrant because the

concen-tration of HA is greater than that for HB Since HA is the stronger of the two weak

acids, it reacts first; thus, the pH before the first equivalence point is controlled by

the HA/A–buffer Between the two equivalence points the pH reflects the titration

of HB and is determined by the HB/B–buffer Finally, after the second equivalence

point, the excess strong base titrant is responsible for the pH

9B.2 Selecting and Evaluating the End Point

Earlier we made an important distinction between an end point and an equivalence

point The difference between these two terms is important and deserves repeating

The equivalence point occurs when stoichiometrically equal amounts of analyte and

titrant react For example, if the analyte is a triprotic weak acid, a titration with

NaOH will have three equivalence points corresponding to the addition of one, two,

and three moles of OH–for each mole of the weak acid An equivalence point,

therefore, is a theoretical not an experimental value

An end point for a titration is determined experimentally and represents the

analyst’s best estimate of the corresponding equivalence point Any difference

be-tween an equivalence point and its end point is a source of determinate error As we

shall see, it is even possible that an equivalence point will not have an associated end

point

Where Is the Equivalence Point? We have already learned how to calculate the

equivalence point for the titration of a strong acid with a strong base, and for the

titration of a weak acid with a strong base We also have learned to sketch a

titra-tion curve with a minimum of calculatitra-tions Can we also locate the equivalence

point without performing any calculations? The answer, as you may have guessed,

is often yes!

It has been shown3 that for most acid–base titrations the inflection point,

which corresponds to the greatest slope in the titration curve, very nearly coincides

with the equivalence point The inflection point actually precedes the equivalence

point, with the error approaching 0.1% for weak acids or weak bases with

dissocia-tion constants smaller than 10–9, or for very dilute solutions Equivalence points

de-termined in this fashion are indicated on the titration curves in Figure 9.8

The principal limitation to using a titration curve to locate the equivalence

point is that an inflection point must be present Sometimes, however, an inflection

point may be missing or difficult to detect Figure 9.9, for example, demonstrates

the influence of the acid dissociation constant, Ka, on the titration curve for a weak

acid with a strong base titrant The inflection point is visible, even if barely so, for

acid dissociation constants larger than 10–9, but is missing when Kais 10–11

Another situation in which an inflection point may be missing or difficult to

detect occurs when the analyte is a multiprotic weak acid or base whose successive

dissociation constants are similar in magnitude To see why this is true let’s

con-sider the titration of a diprotic weak acid, H2A, with NaOH During the titration the

following two reactions occur

Volume of titrant

0.00

(f) (e) (d) (c) (b) (a) 0.0

4.0 6.0 8.0 10.0 12.0

20.00 60.00

2.0

40.00 14.0

Figure 9.10

Titration curves for (a) maleic acid,

pKa1= 1.91, pKa2 = 6.33; (b) malonic acid,

pKa1= 2.85, pKa2= 5.70; (c) succinic acid,

pKa1= 4.21, pKa2 = 5.64 Titration curves are

for 50.00 mL of 0.0500 M acid with 0.100 M

strong base Equivalence points for all three

titrations occur at 25.00 and 50.00 mL of

titrant.

Two distinct inflection points are seen if reaction 9.3 is essentially completebefore reaction 9.4 begins

Figure 9.10 shows titration curves for three diprotic weak acids The

titra-tion curve for maleic acid, for which Ka1is approximately 20,000 times larger

than Ka2, shows two very distinct inflection points Malonic acid, on the otherhand, has acid dissociation constants that differ by a factor of approximately

690 Although malonic acid’s titration curve shows two inflection points, thefirst is not as distinct as that for maleic acid Finally, the titration curve for suc-

cinic acid, for which the two Kavalues differ by a factor of only 27, has only asingle inflection point corresponding to the neutralization of HC4H4O4 to

C4H4O42– In general, separate inflection points are seen when successive aciddissociation constants differ by a factor of at least 500 (a ∆pKaof at least 2.7)

Finding the End Point with a Visual Indicator One interesting group ofweak acids and bases are derivatives of organic dyes Because such com-pounds have at least one conjugate acid–base species that is highly colored,their titration results in a change in both pH and color This change in colorcan serve as a useful means for determining the end point of a titration, pro-vided that it occurs at the titration’s equivalence point

The pH at which an acid–base indicator changes color is determined byits acid dissociation constant For an indicator that is a monoprotic weakacid, HIn, the following dissociation reaction occurs

HIn(aq) + H2O(l) tH3O+(aq) + In–(aq)

for which the equilibrium constant is

9.5

Taking the negative log of each side of equation 9.5, and rearranging to solve for pHgives a familiar equation

9.6

The two forms of the indicator, HIn and In–, have different colors The color of

a solution containing an indicator, therefore, continuously changes as the tration of HIn decreases and the concentration of In–increases If we assume thatboth HIn and In–can be detected with equal ease, then the transition between thetwo colors reaches its midpoint when their concentrations are identical or when the

concen-pH is equal to the indicator’s pKa The equivalence point and the end point

coin-cide, therefore, if an indicator is selected whose pKais equal to the pH at the lence point, and the titration is continued until the indicator’s color is exactlyhalfway between that for HIn and In– Unfortunately, the exact pH at the equiva-lence point is rarely known In addition, detecting the point where the concentra-tions of HIn and In–are equal may be difficult if the change in color is subtle

equiva-We can establish a range of pHs over which the average analyst will observe achange in color if we assume that a solution of the indicator is the color of HInwhenever its concentration is ten times more than that of In–, and the color of In–

14.0

Succinic acid

Figure 9.11

Ladder diagram showing the range of pH levels over which a typical acid–base indicator changes color.

whenever the concentration of HIn is ten times less than that of In–

Substi-tuting these inequalities into equation 9.6

shows that an indicator changes color over a pH range of ±1 units on either

side of its pKa(Figure 9.11) Thus, the indicator will be the color of HIn when

the pH is less than pKa– 1, and the color of In–for pHs greater than pKa+ 1

The pH range of an indicator does not have to be equally distributed on either

side of the indicator’s pKa For some indicators only the weak acid or weak base is

col-ored For other indicators both the weak acid and weak base are colored, but one

form may be easier to see In either case, the pH range is skewed toward those pH

lev-els for which the less colored form of the indicator is present in higher concentration

A list of several common acid–base indicators, along with their pKas, color

changes, and pH ranges, is provided in the top portion of Table 9.4 In some cases,

Table 9.4 Properties of Selected Indicators, Mixed Indicators,

and Screened Indicators for Acid–Base Titrations

bromocresol green and chlorophenol red yellow-green blue-violet 5.4–6.2

Figure 9.12

Titration curve for 50.00 mL of 0.100 M

CH3COOH with 0.100 M NaOH showing the

range of pHs and volumes of titrant over

which the indicators bromothymol blue and

phenolphthalein are expected to change

color.

mixed indicators, which are a mixture of two or more acid–base indicators, provide

a narrower range of pHs over which the color change occurs A few examples ofsuch mixed indicators are included in the middle portion of Table 9.4 Adding aneutral screening dye, such as methylene blue, also has been found to narrow the

pH range over which an indicator changes color (lower portion of Table 9.4) Inthis case, the neutral dye provides a gray color at the midpoint of the indicator’scolor transition

The relatively broad range of pHs over which any indicator changes colorplaces additional limitations on the feasibility of a titration To minimize a determi-nate titration error, an indicator’s entire color transition must lie within the sharptransition in pH occurring near the equivalence point Thus, in Figure 9.12 we seethat phenolphthalein is an appropriate indicator for the titration of 0.1 M aceticacid with 0.1 M NaOH Bromothymol blue, on the other hand, is an inappropriateindicator since its change in color begins before the initial sharp rise in pH and, as aresult, spans a relatively large range of volumes The early change in color increasesthe probability of obtaining inaccurate results, and the range of possible end pointvolumes increases the probability of obtaining imprecise results

The need for the indicator’s color transition to occur in the sharply rising tion of the titration curve justifies our earlier statement that not every equivalencepoint has an end point For example, trying to use a visual indicator to find the firstequivalence point in the titration of succinic acid (see Figure 9.10c) is pointlesssince any difference between the equivalence point and the end point leads to alarge titration error

por-Finding the End Point by Monitoring pH An alternative approach to finding atitration’s end point is to monitor the titration reaction with a suitable sensorwhose signal changes as a function of the analyte’s concentration Plotting the datagives us the resulting titration curve The end point may then be determined fromthe titration curve with only a minimal error

The most obvious sensor for an acid–base titration is a pH electrode.* For ample, Table 9.5 lists values for the pH and volume of titrant obtained during thetitration of a weak acid with NaOH The resulting titration curve, which is called apotentiometric titration curve, is shown in Figure 9.13a The simplest method forfinding the end point is to visually locate the inflection point of the titration curve.This is also the least accurate method, particularly if the titration curve’s slope at theequivalence point is small

Volume of titrant

0.00 0.0

4.0 6.0 8.0 10.0 12.0

2.0

30.00 40.00 50.00 60.00 10.00

14.0

Phenolphthalein Bromothymol blue

Table 9.5 Data for the Titration of a Weak Acid with

0.100 M NaOH

Another method for finding the end point is to plot the first or second

deriva-tive of the titration curve The slope of a titration curve reaches its maximum value

at the inflection point The first derivative of a titration curve, therefore, shows a

separate peak for each end point The first derivative is approximated as ∆pH/∆V,

where ∆pH is the change in pH between successive additions of titrant For

exam-ple, the initial point in the first derivative titration curve for the data in Table 9.5 is

and is plotted at the average of the two volumes (1.00 mL) The remaining data for

the first derivative titration curve are shown in Table 9.5 and plotted in Figure 9.13b

Figure 9.13

Titration curves for a weak acid with

0.100 M NaOH—(a) normal titration

curve; (b) first derivative titration curve;

(c) second derivative titration curve;

(d) Gran plot.

The second derivative of a titration curve may be more useful than the first rivative, since the end point is indicated by its intersection with the volume axis.The second derivative is approximated as ∆(∆pH/∆V)/∆V, or ∆2pH/∆V2 For thetitration data in Table 9.5, the initial point in the second derivative titration curve is

de-and is plotted as the average of the two volumes (2.00 mL) The remainder of thedata for the second derivative titration curve are shown in Table 9.5 and plotted inFigure 9.13c

Derivative methods are particularly well suited for locating end points in protic and multicomponent systems, in which the use of separate visual indicatorsfor each end point is impractical The precision with which the end point may belocated also makes derivative methods attractive for the analysis of samples withpoorly defined normal titration curves

multi-Derivative methods work well only when sufficient data are recorded duringthe sharp rise in pH occurring near the equivalence point This is usually not aproblem when the titration is conducted with an automatic titrator, particularlywhen operated under computer control Manual titrations, however, often containonly a few data points in the equivalence point region, due to the limited range ofvolumes over which the transition in pH occurs Manual titrations are, however, information-rich during the more gently rising portions of the titration curve be-fore and after the equivalence point

Consider again the titration of a monoprotic weak acid, HA, with a strong base

At any point during the titration the weak acid is in equilibrium with H3O+and A–

2

15 20 5

4 6 8

2

15 20 5

–10 –20 –30

14 16 18 20 12

6

4 8 40

80 100 120

60 40 20

13 14 15 11

140

for which

Before the equivalence point, and for volumes of titrant in the titration curve’s

buffer region, the concentrations of HA and A–are given by the following equations

Substituting these equations into the Kaexpression for HA, and rearranging gives

Finally, recognizing that the equivalence point volume is

leaves us with the following equation

Vb×[H3O+] = Ka×Veq– Ka×Vb

For volumes of titrant before the equivalence point, a plot of Vb×[H3O+] versus Vb

is a straight line with an x-intercept equal to the volume of titrant at the end point

and a slope equal to –Ka.* Results for the data in Table 9.5 are shown in Table 9.6

and plotted in Figure 9.13d Plots such as this, which convert a portion of a titration

curve into a straight line, are called Gran plots.

Finding the End Point by Monitoring Temperature The reaction between an acid

and a base is exothermic Heat generated by the reaction increases the temperature

of the titration mixture The progress of the titration, therefore, can be followed by

monitoring the change in temperature

An idealized thermometric titration curve (Figure 9.14a) consists of three

distinct linear regions Before adding titrant, any change in temperature is due to

the cooling or warming of the solution containing the analyte Adding titrant

ini-tiates the exothermic acid–base reaction, resulting in an increase in temperature

This part of a thermometric titration curve is called the titration branch The

temperature continues to rise with each addition of titrant until the equivalencepoint is reached After the equivalence point, any change in temperature is due tothe difference between the temperatures of the analytical solution and the titrant,and the enthalpy of dilution for the excess titrant Actual thermometric titrationcurves (Figure 9.14b) frequently show curvature at the intersection of the titrationbranch and the excess titrant branch due to the incompleteness of the neutraliza-tion reaction, or excessive dilution of the analyte during the titration The latterproblem is minimized by using a titrant that is 10–100 times more concentratedthan the analyte, although this results in a very small end point volume and alarger relative error.

The end point is indicated by the intersection of the titration branch and the excess titrant branch In the idealized thermometric titration curve (see Figure 9.14a) the end point is easily located When the intersection between the two branches is curved, the end point can be found by extrapolation (Figure 9.14b)

Although not commonly used, thermometric titrations have one distinct vantage over methods based on the direct or indirect monitoring of pH As dis-cussed earlier, visual indicators and potentiometric titration curves are limited

ad-by the magnitude of the relevant equilibrium constants For example, the tion of boric acid, H3BO3, for which Ka is 5.8×10–10, yields a poorly definedequivalence point (Figure 9.15a) The enthalpy of neutralization for boric acidwith NaOH, however, is only 23% less than that for a strong acid (–42.7 kJ/mol

titra-Table 9.6 Gran Plot Treatment of the Data

Thermometric titration curves—(a) ideal;

(b) showing curvature at the intersection of

the titration and excess titrant branches.

Equivalence points are indicated by the

V titr

0 (a)

(b)

Volume of titrant

0.00 0.0 4.0 6.0 8.0 10.0 12.0

2.0

3.00 4.00 5.00 1.00

14.0

Volume of titrant

–1 24.980 25.020 25.060

1 2 3 4 5 6 0

25.100

Figure 9.15

Titration curves for 50.00 mL of 0.0100 M

H 3 BO 3 with 0.100 M NaOH determined by

for H3BO3 versus –55.6 kJ/mol for HCl), resulting in a favorable thermometric

titration curve (Figure 9.15b)

9B.3 Titrations in Nonaqueous Solvents

Thus far we have assumed that the acid and base are in an aqueous solution

In-deed, water is the most common solvent in acid–base titrimetry When considering

the utility of a titration, however, the solvent’s influence cannot be ignored

The dissociation, or autoprotolysis constant for a solvent, SH, relates the

con-centration of the protonated solvent, SH2+, to that of the deprotonated solvent, S–

For amphoteric solvents, which can act as both proton donors and proton

accep-tors, the autoprotolysis reaction is

2SH tSH2++ S–

with an equilibrium constant of

Ks= [SH2+][S–]

You should recognize that Kwis just the specific form of Ks for water The pH of a

solution is now seen to be a general statement about the relative abundance of

pro-tonated solvent

pH = –log[SH2+]where the pH of a neutral solvent is given as

Perhaps the most obvious limitation imposed by Ks is the change in pH during

a titration To see why this is so, let’s consider the titration of a 50 mL solution of

10–4M strong acid with equimolar strong base Before the equivalence point, the

pH is determined by the untitrated strong acid, whereas after the equivalence point

the concentration of excess strong base determines the pH In an aqueous solution

the concentration of H3O+when the titration is 90% complete is

corresponding to a pH of 5.3 When the titration is 110% complete, the

concentra-tion of OH–is

or a pOH of 5.3 The pH, therefore, is

pH = pKw– pOH = 14.0 – 5.3 = 8.7The change in pH when the titration passes from 90% to 110% completion is

If the same titration is carried out in a nonaqueous solvent with a Ks

of 1.0×10–20, the pH when the titration is 90% complete is still 5.3.However, the pH when the titration is 110% complete is now

to increase the change in pH when titrating weak acids or bases ure 9.17)

(Fig-Another parameter affecting the feasibility of a titration is the sociation constant of the acid or base being titrated Again, the solventplays an important role In the Brønsted–Lowry view of acid–base be-havior, the strength of an acid or base is a relative measure of the easewith which a proton is transferred from the acid to the solvent, orfrom the solvent to the base For example, the strongest acid that canexist in water is H3O+ The acids HCl and HNO3 are consideredstrong because they are better proton donors than H3O+ Strong acidsessentially donate all their protons to H2O, “leveling” their acid

dis-strength to that of H3O+ In a different solvent HCl and HNO3maynot behave as strong acids

When acetic acid, which is a weak acid, is placed in water, the sociation reaction

dis-CH3COOH(aq) + H2O(l) tH3O+(aq) + CH3COO–(aq)

does not proceed to a significant extent because acetate is a stronger base than waterand the hydronium ion is a stronger acid than acetic acid If acetic acid is placed in asolvent that is a stronger base than water, such as ammonia, then the reaction

CH3COOH + NH3tNH4++ CH3COO–

proceeds to a greater extent In fact, HCl and CH3COOH are both strong acids inammonia

All other things being equal, the strength of a weak acid increases if it is placed

in a solvent that is more basic than water, whereas the strength of a weak base creases if it is placed in a solvent that is more acidic than water In some cases, how-

in-ever, the opposite effect is observed For example, the pKbfor ammonia is 4.76 inwater and 6.40 in the more acidic glacial acetic acid In contradiction to our expec-tations, ammonia is a weaker base in the more acidic solvent A full description of

the solvent’s effect on a weak acid’s pKaor on the pKbof a weak base is beyond thescope of this text You should be aware, however, that titrations that are not feasible

in water may be feasible in a different solvent

9B.4 Representative Method

Although each acid–base titrimetric method has its own unique considerations, thefollowing description of the determination of protein in bread provides an instruc-tive example of a typical procedure

Figure 9.16

Titration curves for 50.00 mL of 10 –4 M HCl

with 10 –4 M NaOH in (a) water,

Kw= l × 10 –14 , and (b) nonaqueous solvent,

Ks= 1 × 10 –20

Figure 9.17

Titration curves for 50.00 mL of 0.100 M

weak acid (pKa = 11) with 0.100 M NaOH in

(a) water, Kw= 1 × 10 –14 ; and

(b) nonaqueous solvent, Ks = 1 × 10 –20 The

titration curve in (b) assumes that the

change in solvent has no effect on the acid

dissociation constant of the weak acid.

leveling

Acids that are better proton donors than

the solvent are leveled to the acid

strength of the protonated solvent; bases

that are better proton acceptors than the

solvent are leveled to the base strength of

the deprotonated solvent.

Representative Methods

—Continued

Method 9.1 Determination of Protein in Bread 4

Description of the Method. This quantitative method of analysis for proteins

is based on a determination of the %w/w N in the sample Since different cereal

proteins have similar amounts of nitrogen, the experimentally determined

%w/w N is multiplied by a factor of 5.7 to give the %w/w protein in the sample (on

average there are 5.7 g of cereal protein for every gram of nitrogen) As described

here, nitrogen is determined by the Kjeldahl method The protein in a sample of

bread is oxidized in hot concentrated H2SO4, converting the nitrogen to NH4 After

making the solution alkaline, converting NH4 to NH3, the ammonia is distilled into

a flask containing a known amount of standard strong acid Finally, the excess

strong acid is determined by a back titration with a standard strong base titrant.

Procedure. Transfer a 2.0-g sample of bread, which has previously been air

dried and ground into a powder, to a suitable digestion flask, along with 0.7 g of

HgO as a catalyst, 10 g of K 2 SO 4 , and 25 mL of concentrated H 2 SO 4 Bring the

solution to a boil, and continue boiling until the solution turns clear, and for at

least an additional 30 min After cooling to below room temperature, add 200 mL

of H2O and 25 mL of 4% w/v K2S to remove the Hg 2+ catalyst Add a few Zn

granules to serve as boiling stones, and 25 g of NaOH Quickly connect the flask to

a distillation apparatus, and distill the NH3into a collecting flask containing a

known amount of standardized HCl The tip of the condenser should be placed

below the surface of the strong acid After the distillation is complete, titrate the

excess strong acid with a standard solution of NaOH, using methyl red as a visual

indicator.

Questions

1 Oxidizing the protein converts the nitrogen to NH 4 Why is the amount of

nitrogen not determined by titrating the NH 4 with a strong base?

There are two reasons for not titrating the ammonium ion First, NH4 is a very

weak acid (Ka = 5.7 × 10 –10 ) that yields a poorly defined end point when titrated with a strong base Second, even if the end point can be determined with acceptable accuracy and precision, the procedure calls for adding a substantial amount of H2SO4 After the oxidation is complete, the amount of excess H2SO4will be much greater than the amount of NH4 that is produced The presence

of two acids that differ greatly in concentration makes for a difficult analysis If the titrant’s concentration is similar to that of H2SO4, then the equivalence point volume for the titration of NH4 may be too small to measure reliably On the other hand, if the concentration of the titrant is similar to that of NH4, the volume needed to neutralize the H 2 SO 4 will be unreasonably large.

2 Ammonia is a volatile compound as evidenced by the strong smell of even

dilute solutions This volatility presents a possible source of determinate error.

Will this determinate error be negative or positive?

The conversion of N to NH3follows the following pathway

N → NH4

NH 4 → NH 3

Any loss of NH3is loss of analyte and a negative determinate error.

The photo in Colorplate 8a shows the indicator’s color change for this titration.

9B.5 Quantitative Applications

Although many quantitative applications of acid–base titrimetry have been replaced

by other analytical methods, there are several important applications that continue

to be listed as standard methods In this section we review the general application ofacid–base titrimetry to the analysis of inorganic and organic compounds, with anemphasis on selected applications in environmental and clinical analysis First,however, we discuss the selection and standardization of acidic and basic titrants

Selecting and Standardizing a Titrant Most common acid–base titrants are notreadily available as primary standards and must be standardized before they can beused in a quantitative analysis Standardization is accomplished by titrating aknown amount of an appropriate acidic or basic primary standard

The majority of titrations involving basic analytes, whether conducted in ous or nonaqueous solvents, use HCl, HClO4, or H2SO4as the titrant Solutions ofthese titrants are usually prepared by diluting a commercially available concentratedstock solution and are stable for extended periods of time Since the concentrations

aque-of concentrated acids are known only approximately,* the titrant’s concentration isdetermined by standardizing against one of the primary standard weak bases listed

in Table 9.7

The most common strong base for titrating acidic analytes in aqueous solutions

is NaOH Sodium hydroxide is available both as a solid and as an approximately50% w/v solution Solutions of NaOH may be standardized against any of the pri-mary weak acid standards listed in Table 9.7 The standardization of NaOH, how-ever, is complicated by potential contamination from the following reaction be-tween CO2and OH–

CO2(g) + 2OH–(aq)→CO32–(aq) + H2O(l) 9.7

When CO2is present, the volume of NaOH used in the titration is greater than thatneeded to neutralize the primary standard because some OH–reacts with the CO2

3 Discuss the steps taken in this procedure to minimize this determinate error.

Three specific steps are taken to minimize the loss of ammonia: (1) the solution

is cooled to below room temperature before adding NaOH; (2) the digestion flask is quickly connected to the distillation apparatus after adding NaOH; and (3) the condenser tip of the distillation apparatus is placed below the surface of the HCl to ensure that the ammonia will react with the HCl before it can be lost through volatilization.

4 How does K 2 S remove Hg 2+ , and why is this important?

Adding sulfide precipitates the Hg 2+ as HgS This is important because NH3forms stable complexes with many metal ions, including Hg 2+ Any NH 3 that is complexed with Hg 2+ will not be collected by distillation, providing another source of determinate error.

Continued from page 297

*The nominal concentrations are 12.1 M HCl, 11.7 M HClO , and 18.0 M H SO

The calculated concentration of OH–, therefore, is too small This is not a problem

when titrations involving NaOH are restricted to an end point pH less than 6

Below this pH any CO32–produced in reaction 9.7 reacts with H3O+to form

car-bonic acid

CO32–(aq) + 2H3O+(aq)→H2CO3(aq) + 2H2O(l) 9.8

Combining reactions 9.7 and 9.8 gives an overall reaction of

CO2(g) + H2O(l)→H2CO3(aq)

which does not include OH– Under these conditions the presence of CO2does not

affect the quantity of OH–used in the titration and, therefore, is not a source of

CO2(g) + OH–(aq)→HCO3(aq)

Under these conditions some OH–is consumed in neutralizing CO2 The result is a

determinate error in the titrant’s concentration If the titrant is used to analyze an

analyte that has the same end point pH as the primary standard used during

stan-dardization, the determinate errors in the standardization and the analysis cancel,

and accurate results may still be obtained

Solid NaOH is always contaminated with carbonate due to its contact with the

atmosphere and cannot be used to prepare carbonate-free solutions of NaOH

So-lutions of carbonate-free NaOH can be prepared from 50% w/v NaOH since

NaCO is very insoluble in concentrated NaOH When CO is absorbed, NaCO

Table 9.7 Selected Primary Standards for the Standardization

of Strong Acid and Strong Base Titrants

Standardization of Acidic Titrants

Na 2 B 4 O 7 Na 2 B 4 O 7 + 2H 3 O + + 3H 2 O → 2Na + + 4H 3 BO 3

Standardization of Basic Titrants

KHC 8 H 4 O 4 KHC 8 H 4 O 4 + OH – → K + + C 8 H 4 O 42–+ H 2 O c

KH(IO 3 ) 2 KH(IO 3 ) 2 + OH – → K + + 2IO 3–+ H 2 O

a The end point for this titration is improved by titrating to the second equivalence point, boiling the

solution to expel CO 2 , and retitrating to the second equivalence point In this case the reaction is

Na 2 CO 3 + 2H 3 O + → CO 2 + 2Na + + 3H 2 O

bTRIS stands for tris-(hydroxymethyl)aminomethane.

c KHC 8 H 4 O 4 is also known as potassium hydrogen phthalate, or KHP.

d Due to its poor solubility in water, benzoic acid is dissolved in a small amount of ethanol before being

diluted with water.

precipitates and settles to the bottom of the container, allowing access to the carbonate-free NaOH Dilution must be done with water that is free from dissolved

CO2 Briefly boiling the water expels CO2and, after cooling, it may be used to pare carbonate-free solutions of NaOH Provided that contact with the atmosphere

pre-is minimized, solutions of carbonate-free NaOH are relatively stable when stored

in polyethylene bottles Standard solutions of sodium hydroxide should not bestored in glass bottles because NaOH reacts with glass to form silicate

Inorganic Analysis Acid–base titrimetry is a standard method for the quantitativeanalysis of many inorganic acids and bases Standard solutions of NaOH can beused in the analysis of inorganic acids such as H3PO4or H3AsO4, whereas standardsolutions of HCl can be used for the analysis of inorganic bases such as Na2CO3.Inorganic acids and bases too weak to be analyzed by an aqueous acid–basetitration can be analyzed by adjusting the solvent or by an indirect analysis For ex-ample, the accuracy in titrating boric acid, H3BO3, with NaOH is limited by boricacid’s small acid dissociation constant of 5.8×10–10 The acid strength of boric acid,however, increases when mannitol is added to the solution because it forms a com-

plex with the borate ion The increase in Kato approximately 1.5×10–4results in asharper end point and a more accurate titration Similarly, the analysis of ammo-nium salts is limited by the small acid dissociation constant of 5.7×10–10for NH4+

In this case, NH4+can be converted to NH3by neutralizing with strong base The

NH3, for which Kbis 1.8×10–5, is then removed by distillation and titrated with astandard strong acid titrant

Inorganic analytes that are neutral in aqueous solutions may still be analyzed ifthey can be converted to an acid or base For example, NO3 can be quantitativelyanalyzed by reducing it to NH3in a strongly alkaline solution using Devarda’s alloy,

a mixture of 50% w/w Cu, 45% w/w Al, and 5% w/w Zn

3NO3 (aq) + 8Al(s) + 5OH–(aq) + 2H2O(l)→8AlO2(aq) + 3NH3(aq)

The NH3is removed by distillation and titrated with HCl Alternatively, NO3 can

be titrated as a weak base in an acidic nonaqueous solvent such as anhydrous aceticacid, using HClO4as a titrant

Acid–base titrimetry continues to be listed as the standard method for the termination of alkalinity, acidity, and free CO2in water and wastewater analysis Al- kalinity is a measure of the acid-neutralizing capacity of a water sample and is as-

de-sumed to arise principally from OH–, HCO3 , and CO32–, although other weakbases, such as phosphate, may contribute to the overall alkalinity Total alkalinity isdetermined by titrating with a standard solution of HCl or H2SO4 to a fixed endpoint at a pH of 4.5, or to the bromocresol green end point Alkalinity is reported asmilligrams CaCO3per liter

When the sources of alkalinity are limited to OH–, HCO3, and CO32–, tions to both a pH of 4.5 (bromocresol green end point) and a pH of 8.3 (phe-nolphthalein or metacresol purple end point) can be used to determine whichspecies are present, as well as their respective concentrations Titration curves for

titra-OH–, HCO3 , and CO32– are shown in Figure 9.18 For a solution containing only

OH–alkalinity, the volumes of strong acid needed to reach the two end points areidentical If a solution contains only HCO3 alkalinity, the volume of strong acidneeded to reach the end point at a pH of 8.3 is zero, whereas that for the pH 4.5 endpoint is greater than zero When the only source of alkalinity is CO32–, the volume

of strong acid needed to reach the end point at a pH of 4.5 is exactly twice thatneeded to reach the end point at a pH of 8.3

alkalinity

A measure of a water’s ability to

neutralize acid.

Figure 9.18

Titration curves for (a) 50.00 mL of 0.100 M NaOH with 0.100 M HCl; (b) 50.00 mL of 0.100 M Na 2 CO 3 with 0.100 M HCl; and (c) 50.00 mL of 0.100 M NaHCO 3 with

M HCl The dashed lines indicate the pH 8.3 and pH 4.5 end points.

Mixtures of OH–and CO32–, or HCO3 and CO32–alkalinities also are possible

Consider, for example, a mixture of OH–and CO32– The volume of strong acid

needed to titrate OH–will be the same whether we titrate to the pH 8.3 or pH 4.5

end point Titrating CO32–to the end point at a pH of 4.5, however, requires twice

as much strong acid as when titrating to the pH 8.3 end point Consequently, when

titrating a mixture of these two ions, the volume of strong acid needed to reach the

pH 4.5 end point is less than twice that needed to reach the end point at a pH of 8.3

For a mixture of HCO3 and CO32–, similar reasoning shows that the volume of

strong acid needed to reach the end point at a pH of 4.5 is more than twice that

need to reach the pH 8.3 end point Solutions containing OH–and HCO3

alkalini-ties are unstable with respect to the formation of CO32–and do not exist Table 9.8

summarizes the relationship between the sources of alkalinity and the volume of

titrant needed to reach the two end points

Acidity is a measure of a water sample’s capacity for neutralizing base and is

conveniently divided into strong acid and weak acid acidity Strong acid acidity is

due to the presence of inorganic acids, such as HCl, HNO3, and H2SO4, and is

com-monly found in industrial effluents and acid mine drainage Weak acid acidity

is usually dominated by the formation of HCO from dissolved CO , but also

Volume of titrant

0.00 0.0 4.0 6.0 8.0 10.0 12.0

70.00

2.0

10.00 20.00 30.00 40.00 50.00 60.00 14.0

Volume of titrant

0.00 0.0 4.0 6.0 8.0 10.0 12.0

140.00

2.0

20.00 40.00 60.00 80.00 100.00 120.00 14.0

0.00 10.00 20.00 30.00 40.00 50.00 60.00 70.00

Volume of titrant

0.0 4.0 6.0 8.0 10.0 12.0

2.0 14.0

includes contributions from hydrolyzable metal ions such as Fe3+, Al3+, and Mn2+.

In addition, weak acid acidity may include a contribution from organic acids.Acidity is determined by titrating with a standard solution of NaOH to fixedend points at pH 3.7 and pH 8.3 These end points are located potentiometrically,using a pH meter, or by using an appropriate indicator (bromophenol blue for

pH 3.7, and metacresol purple or phenolphthalein for pH 8.3) Titrating to a pH of3.7 provides a measure of strong acid acidity,* and titrating to a pH of 8.3 provides

a measure of total acidity Weak acid acidity is given indirectly as the difference tween the total and strong acid acidities Results are expressed as the milligrams ofCaCO3per liter that could be neutralized by the water sample’s acidity An alterna-tive approach for determining strong and weak acidity is to obtain a potentiometrictitration curve and use Gran plot methodology to determine the two equivalencepoints This approach has been used, for example, in determining the forms of acid-ity in atmospheric aerosols.5

be-Water in contact with either the atmosphere or carbonate-bearing sedimentscontains dissolved or free CO2that exists in equilibrium with gaseous CO2and theaqueous carbonate species H2CO3, HCO3, and CO32– The concentration of free

CO2is determined by titrating with a standard solution of NaOH to the phthalein end point, or to a pH of 8.3, with results reported as milligrams CO2perliter This analysis is essentially the same as that for the determination of total acid-ity, and can only be applied to water samples that do not contain any strong acidacidity

phenol-Organic Analysis The use of acid–base titrimetry for the analysis of organic pounds continues to play an important role in pharmaceutical, biochemical, agri-cultural, and environmental laboratories Perhaps the most widely employed

com-acid–base titration is the Kjeldahl analysis for organic nitrogen, described earlier in

Method 9.1 This method continues to be used in the analysis of caffeine and charin in pharmaceutical products, as well as for the analysis of proteins, fertilizers,sludges, and sediments Any nitrogen present in the –3 oxidation state is quantita-tively oxidized to NH4+ Some aromatic heterocyclic compounds, such as pyridine,are difficult to oxidize A catalyst, such as HgO, is used to ensure that oxidation iscomplete Nitrogen in an oxidation state other than –3, such as nitro- and azo-nitrogens, is often oxidized to N2, resulting in a negative determinate error Adding

sac-a reducing sac-agent, such sac-as ssac-alicylic sac-acid, reduces the nitrogen to sac-a –3 oxidsac-ation stsac-ate,

*This is sometimes referred to as “methyl orange acidity” since, at one time, methyl orange was the traditional indicator

Table 9.8 Relationship Between End Point Volumes

and Sources of Alkalinity

An acid–base titrimetric method for

determining the amount of nitrogen in

organic compounds.

eliminating this source of error Other examples of elemental analyses based on the

conversion of the element to an acid or base are outlined in Table 9.9

Several organic functional groups have weak acid or weak base properties that

allow their direct determination by an acid–base titration Carboxylic (—COOH),

sulfonic (—SO3H), and phenolic (—C6H5OH) functional groups are weak acids

that can be successfully titrated in either aqueous or nonaqueous solvents Sodium

hydroxide is the titrant of choice for aqueous solutions Nonaqueous titrations are

often carried out in a basic solvent, such as ethylenediamine, using

tetrabutylam-monium hydroxide, (C4H9)4NOH, as the titrant Aliphatic and aromatic amines are

weak bases that can be titrated using HCl in aqueous solution or HClO4in glacial

acetic acid Other functional groups can be analyzed indirectly by use of a

func-tional group reaction that produces or consumes an acid or base Examples are

shown in Table 9.10

Many pharmaceutical compounds are weak acids or bases that can be analyzed

by an aqueous or nonaqueous acid–base titration; examples include salicylic acid,

phenobarbital, caffeine, and sulfanilamide Amino acids and proteins can be

ana-lyzed in glacial acetic acid, using HClO4as the titrant For example, a procedure for

determining the amount of nutritionally available protein has been developed that

is based on an acid–base titration of lysine residues.6

Table 9.9 Selected Elemental Analyses Based on Acid–Base Titrimetry

aThe acid or base that is eventually titrated is indicated in bold.

Table 9.10 Selected Acid–Base Titrimetric Procedures for Organic Functional Groups Based

on the Production or Consumption of Acid or Base

Functional Group Reaction Producing Acid or Base to Be Titrated a Titration

alcohol b [1] (CH 3 CO) 2 O + ROH → CH 3COOR + CH 3 COOH CH 3 COOH with strong base; ROH is determined

[2] (CH 3 CO) 2 O + H 2 O →2CH 3 COOH

from the difference in the amount of titrant needed to react with a blank consisting only of acetic anhydride, and the amount reacting with the sample.

aThe acid or base that is eventually titrated is indicated in bold.

b The acetylation reaction, [1], is carried out in pyridine to avoid the hydrolysis of acetic anhydride by water After the acetylation is complete, water is added to convert the remaining acetic anhydride to acetic acid, [2].

R C2 — O (aq) + NH OH HCI 2 ⋅ (aq) →

R C 2 — NOH (aq) +HCI(aq) + H O 2 ( ) l

Quantitative Calculations In acid–base titrimetry the quantitative relationship tween the analyte and the titrant is determined by the stoichiometry of the relevantreactions As outlined in Section 2C, stoichiometric calculations may be simplified

be-by focusing on appropriate conservation principles In an acid–base reaction thenumber of protons transferred between the acid and base is conserved; thus

The following example demonstrates the application of this approach in the directanalysis of a single analyte

EXAMPLE 9.2

A 50.00-mL sample of a citrus drink requires 17.62 mL of 0.04166 M NaOH toreach the phenolphthalein end point (Figure 9.19a) Express the sample’sacidity in terms of grams of citric acid, C6H8O7, per 100 mL

SOLUTION

Since citric acid is a triprotic weak acid, we must first decide to whichequivalence point the titration has been carried The three acid dissociationconstants are

2.0 14.0

(a)

(b)

The phenolphthalein end point is basic, occurring at a pH of approximately 8.3

and can be reached only if the titration proceeds to the third equivalence point

(Figure 9.19b); thus, we write

3×moles citric acid = moles NaOHMaking appropriate substitutions for the moles of citric acid and moles of

NaOH gives the following equation

which can be solved for the grams of citric acid

Since this is the grams of citric acid in a 50.00-mL sample, the concentration of

citric acid in the citrus drink is 0.09402 g/100 mL

In an indirect analysis the analyte participates in one or more preliminary

reac-tions that produce or consume acid or base Despite the additional complexity, the

stoichiometry between the analyte and the amount of acid or base produced or

con-sumed may be established by applying the conservation principles outlined in

Sec-tion 2C Example 9.3 illustrates the applicaSec-tion of an indirect analysis in which an

acid is produced

EXAMPLE 9.3

The purity of a pharmaceutical preparation of sulfanilamide, C6H4N2O2S, can

be determined by oxidizing the sulfur to SO2 and bubbling the SO2through

H2O2to produce H2SO4 The acid is then titrated with a standard solution of

NaOH to the bromothymol blue end point, where both of sulfuric acid’s acidic

protons have been neutralized Calculate the purity of the preparation, given

that a 0.5136-g sample required 48.13 mL of 0.1251 M NaOH

SOLUTION

Conservation of protons for the titration reaction requires that

2×moles H2SO4= moles NaOHSince all the sulfur in H2SO4comes from sulfanilamide, we use a conservation

of mass on sulfur to establish the following stoichiometric relationship

Moles C6H4N2O2S = moles H2SO4

Combining the two conservation equations gives a single equation relating the

moles of analyte to the moles of titrant

2×moles C6H4N2O2S = moles NaOHMaking appropriate substitutions for the moles of sulfanilamide and moles of

which can be solved for the grams of sulfanilamide

Thus, the purity of the preparation is

This approach is easily extended to back titrations, as shown in the following example

SOLUTION

In this procedure, the HCl reacts with two different bases; thusMoles HCl = moles HCl reacting with NH3+ moles HCl reacting with NaOHConservation of protons requires that

Moles HCl reacting with NH3= moles NH3

Moles HCl reacting with NaOH = moles NaOH

A conservation of mass on nitrogen gives the following equation

Moles NH3= moles NCombining all four equations gives a final stoichiometric equation of

Moles HCl = moles N + moles NaOHMaking appropriate substitutions for the moles of HCl, N, and NaOH gives

which we solve for the grams of nitrogen

g N = (Ma×Va– Mb×Vb)×AW N(0.1047 M×0.05000 L – 0.1183 M×0.02284 L)×14.01 g/mol = 0.03549 g N

The mass of protein, therefore, is

and the %w/w protein is

Earlier we noted that an acid–base titration may be used to analyze a mixture of

acids or bases by titrating to more than one equivalence point The concentration of

each analyte is determined by accounting for its contribution to the volume of

titrant needed to reach the equivalence points

EXAMPLE 9.5

The alkalinity of natural waters is usually controlled by OH–, CO32–, and

HCO3 , which may be present singularly or in combination Titrating a

100.0-mL sample to a pH of 8.3 requires 18.67 mL of a 0.02812 M solution of

HCl A second 100.0-mL aliquot requires 48.12 mL of the same titrant to

reach a pH of 4.5 Identify the sources of alkalinity and their concentrations

in parts per million

SOLUTION

Since the volume of titrant needed to reach a pH of 4.5 is more than twice that

needed to reach a pH of 8.3, we know, from Table 9.8, that the alkalinity of the

sample is controlled by CO32–and HCO32–

Titrating to a pH of 8.3 neutralizes CO32–to HCO3 , but does not lead to areaction of the titrant with HCO3 (see Figure 9.14b) Thus

The concentration of CO32–, therefore, is

Titrating to the second end point at pH 4.5 neutralizes CO32–to H2CO3, and

HCO3 to H2CO3(see Figures 9.18b,c) The conservation of protons, therefore,

requires that

Moles HCl to pH 4.5 = 2×moles CO32–+ moles HCO3

or

mg COliter

mg

3 2

Solving for the grams of bicarbonate gives

The concentration of HCO3, therefore, is

9B.6 Qualitative Applications

We have already come across one example of the qualitative application ofacid–base titrimetry in assigning the forms of alkalinity in waters (see Example 9.5).This approach is easily extended to other systems For example, the composition ofsolutions containing one or two of the following species

can be determined by titrating with either a strong acid or a strong base to themethyl orange and phenolphthalein end points As outlined in Table 9.11, eachspecies or mixture of species has a unique relationship between the volumes oftitrant needed to reach these two end points

mg HCOliter

Table 9.11 Relationship Between End Point Volumes for Solutions of Phosphate Species

with HCl and NaOH

9B.7 Characterization Applications

Two useful characterization applications involving acid–base titrimetry are the

de-termination of equivalent weight, and the dede-termination of acid–base dissociation

constants

Equivalent Weights Acid–base titrations can be used to characterize the chemical

and physical properties of matter One simple example is the determination of the

equivalent weight* of acids and bases In this method, an accurately weighed sample

of a pure acid or base is titrated to a well-defined equivalence point using a

mono-protic strong acid or strong base If we assume that the titration involves the

trans-fer of n protons, then the moles of titrant needed to reach the equivalence point is

given as

Moles titrant = n×moles analyteand the formula weight is

Since the actual number of protons transferred between the analyte and titrant is

uncertain, we define the analyte’s equivalent weight (EW) as the apparent formula

weight when n = 1 The true formula weight, therefore, is an integer multiple of the

calculated equivalent weight

Thus, if we titrate a monoprotic weak acid with a strong base, the EW and FW are

identical If the weak acid is diprotic, however, and we titrate to its second

equiva-lence point, the FW will be twice as large as the EW

EXAMPLE 9.6

A 0.2521-g sample of an unknown weak acid is titrated with a 0.1005 M

solution of NaOH, requiring 42.68 mL to reach the phenolphthalein end point

Determine the compound’s equivalent weight Which of the following

compounds is most likely to be the unknown weak acid?

ascorbic acid C6H8O6 FW = 176.1 monoprotic

succinic acid C4H6O4 FW = 118.1 diprotic

SOLUTION

The moles of NaOH needed to reach the end point is

Mb×Vb= 0.1005 M×0.04268 L = 4.289×10–3mol NaOHgiving an equivalent weight of

Figure 9.20

Estimating the pKa for a weak acid from its

titration curve with a strong base.

The possible formula weights for the unknown weak acid are

of these values is close to the formula weight for citric acid, eliminating it as apossibility Only succinic acid provides a possible match

Equilibrium Constants Another application of acid–base titrimetry is the nation of equilibrium constants Consider, for example, the titration of a weak acid,

determi-HA, with a strong base The dissociation constant for the weak acid is

9.9

When the concentrations of HA and A– are equal, equation 9.9 reduces to

Ka= [H3O+], or pH = pKa Thus, the pKafor a weak acid can be determined bymeasuring the pH for a solution in which half of the weak acid has been neutralized

On a titration curve, the point of half-neutralization is approximated by the volume

of titrant that is half of that needed to reach the equivalence point As shown in

Fig-ure 9.20, an estimate of the weak acid’s pKacan be obtained directly from the tion curve

titra-This method provides a reasonable estimate of the pKa, provided that the weakacid is neither too strong nor too weak These limitations are easily appreciated byconsidering two limiting cases For the first case let’s assume that the acid is strongenough that it is more than 50% dissociated before the titration begins As a resultthe concentration of HA before the equivalence point is always less than the con-centration of A–, and there is no point along the titration curve where [HA] = [A–]

At the other extreme, if the acid is too weak, the equilibrium constant for the tion reaction

titra-HA(aq) + OH–(aq) tH2O(l) + A–(aq)

may be so small that less than 50% of HA will have reacted at the equivalence point

In this case the concentration of HA before the equivalence point is always greater

2.0

60.00 14.0

than that of A– Determining the pKaby the half-equivalence point method

overes-timates its value if the acid is too strong and underesoveres-timates its value if the acid is

too weak

A second approach for determining the pKaof an acid is to replot the titration

curve in a linear form as a Gran plot For example, earlier we learned that the

titra-tion of a weak acid with a strong base can be plotted in a linear form using the

fol-lowing equation

Vb×[H3O+] = Ka×Veq– Ka×Vb

Plotting Vb×[H3O+] versus Vb, for volumes less than the equivalence point volume

yields a straight line with a slope of –Ka Other linearizations have been developed

that use all the points on a titration curve7or require no assumptions.8 This

ap-proach to determining acidity constants has been used to study the acid–base

prop-erties of humic acids, which are naturally occurring, large-molecular-weight organic

acids with multiple acidic sites In one study, a sample of humic acid was found to

have six titratable sites, three of which were identified as carboxylic acids, two of

which were believed to be secondary or tertiary amines, and one of which was

iden-tified as a phenolic group.9

9B.8 Evaluation of Acid–Base Titrimetry

Scale of Operation In an acid–base titration the volume of titrant needed to

reach the equivalence point is proportional to the absolute amount of analyte

present in the analytical solution Nevertheless, the change in pH at the

equiv-alence point, and thus the utility of an acid–base titration, is a function of the

analyte’s concentration in the solution being titrated

When the sample is available as a solution, the smallest concentration of

analyte that can be readily analyzed is approximately 10–3M (Figure 9.21) If,

for example, the analyte has a gram formula weight of 120 g/mol, then the

lower concentration limit is 120 ppm When the analyte is a solid, it must first

be placed into solution, the volume of which must be sufficient to allow the

titration’s end point to be monitored using a visual indicator or a suitable

probe If we assume a minimum volume of 25 mL, and a lower concentration

limit of 120 ppm, then a sample containing at least 3 mg of analyte is

re-quired Acid–base titrations involving solid or solution samples, therefore, are

generally limited to major and minor analytes (see Figure 3.6 in Chapter 3)

The analysis of gases can be extended to trace analytes by pulling a large

vol-ume of the gas through a suitable collection solution

Efforts have been made to develop methods for conducting acid–base

titrations on a much smaller scale In one experimental design, samples of

20–100 µL were held by capillary action between a flat-surface pH electrode

and a stainless steel rod.10The titrant was added by using the oscillations of

a piezoelectric ceramic device to move an angled glass rod in and out of a

tube connected to a reservoir containing the titrant (see Figure 9.22) Each

time the glass tube was withdrawn an approximately 2-nL microdroplet of

titrant was released The microdroplets were allowed to fall onto the steel

rod containing the sample, with mixing accomplished by spinning the rod

at 120 rpm A total of 450 microdroplets, with a combined volume of

0.81–0.84 µL, was dispensed between each pH measurement In this fashion

a titration curve was constructed This method was used to titrate solutions

of 0.1 M HCl and 0.1 M CH3COOH with 0.1 M NaOH Absolute errors

ranged from a minimum of +0.1% to a maximum of –4.1%, with relative

(a) (b) (c) (d)

(a) (c) (d)

20 40 60 80

Volume of NaOH (mL)

0.00 4.00 6.00 8.00 10.00 12.00

2.00 14.00

Piezoelectric ceramic

Sample

pH meter

Titrant

Figure 9.21

Titration curves for (a) 10 –1 M HCl, (b) 10 –2 M HCl, (c) 10 –3 M HCl, and (d) 10 –4 M HCl In each case the titrant

is an equimolar solution of NaOH.

Figure 9.22

Experimental design for a microdroplet titration apparatus.

Figure 9.23

(a) Experimental set-up for a diffusional

microtitration; (b) close-up showing the tip

of the diffusional microburet in contact with

the drop of sample.

standard deviations from 0.15% to 4.7% The smallest volume of sample that wassuccessfully titrated was 20 µL

More recently, a method has been described in which the acid–base titration isconducted within a single drop of solution.11The titrant is added using a microbu-ret fashioned from a glass capillary micropipet (Figure 9.23) The microburet has a1–2 µm tip filled with an agar gel membrane The tip of the microburet is placedwithin a drop of the sample solution, which is suspended in heptane, and the titrant

is allowed to diffuse into the sample The titration is followed visually using a ored indicator, and the time needed to reach the end point is measured The rate ofthe titrant’s diffusion from the microburet must be determined by calibration.Once calibrated, the end point time can be converted to an end point volume Sam-ples usually consisted of picoliter volumes (10–12L), with the smallest sample being0.7 pL The precision of the titrations was usually about 2%

col-Titrations conducted with microliter or picoliter sample volumes require asmaller absolute amount of analyte For example, diffusional titrations have beensuccessfully conducted on as little as 29 femtomoles (10–15mol) of nitric acid Nev-ertheless, the analyte must still be present in the sample at a major or minor levelfor the titration to be performed accurately and precisely

Accuracy When working with macro–major and macro–minor samples,acid–base titrations can be accomplished with relative errors of 0.1–0.2% The prin-cipal limitation to accuracy is the difference between the end point and the equiva-lence point

Precision The relative precision of an acid–base titration depends primarily on theprecision with which the end point volume can be measured and the precision ofthe end point signal Under optimum conditions, an acid–base titration can be ac-complished with a relative precision of 0.1–0.2% The relative precision can be im-proved by using the largest volume buret that is feasible and ensuring that most ofits capacity is used to reach the end point Smaller volume burets are used when thecost of reagents or waste disposal is of concern or when the titration must be com-pleted quickly to avoid competing chemical reactions Automatic titrators are par-ticularly useful for titrations requiring small volumes of titrant, since the precisionwith which the volume can be measured is significantly better (typically about

±0.05% of the buret’s volume)

The precision of the end point signal depends on the method used to locate theend point and the shape of the titration curve With a visual indicator, the precision

of the end point signal is usually between ±0.03 mL and 0.10 mL End points mined by direct monitoring often can be determined with a greater precision

deter-Sensitivity For an acid–base titration we can write the following general analyticalequation

Volume of titrant = k×moles of analyte

where k, the sensitivity, is determined by the stoichiometric relationship between

analyte and titrant Note that this equation assumes that a blank has been analyzed

to correct the signal for the volume of titrant reacting with the reagents

Consider, for example, the determination of sulfurous acid, H2SO3, by titratingwith NaOH to the first equivalence point Using the conservation of protons, we write

Moles NaOH = moles HSO

microburet