Acid-base indicators 589

In an acid-base titration there will only be observed a pH of 7 at the equivalence point if both titrant and titrand are strong electrolytes. If one is weak. 

Acid-base indicators show differing colours with varying hydrogen ion concentration in a solution. The change in colour occurs in general over a colour change interval of some two pH units. It is necessary to select indicators for particular titrations which show clear colours at pH values close to those known to hold at the equivalence point. 

Indicators are themselves weak organic acids or bases whose undissociated forms differ in colour from the ionic forms due to their considerably different electronic structure and hence absorption spectra. 

The dissociation of an acid indicator molecule in water occurs according to 

In dilute solution the equilibrium constant, K, for this reaction is 

The indicator phenolphthalein is colorless in acidic solution and pink in basic solution. 

There are two common methods for determining the equivalence point of an acid-base 

Use a pH meter (see Fig. 14.7) to monitor the pH and then plot the titration curve. The center of the vertical region of the pH curve indicates the equivalence point (for example, see Figs. 15.1 through 15.5). 

Use an acid-base indicator, which marks the end point of a titration by changing color. Although the equivalence point of a titration, defined by the stoichiometry, is not necessarily the same as the end point (where the indicator changes color), careful selection of the indicator will ensure that the error is negligible. 

The most common acid-base indicators are complex molecules that are themselves weak acids (represented by HIn). They exhibit one color when the proton is attached to the molecule and a different color when the proton is absent. For example, phenolphthalein, a commonly used indicator, is colorless in its HIn form and pink in its In, or basic, form. The actual structures of the two forms of phenolphthalein are shown in Fig. 15.6. 

As pointed out in Chapter 4, an acid-base indicator is useful in determining the equivalence point of an acid-base titration. This is the point at which reaction is complete equivalent quantities of acid and base have reacted. If the indicator is chosen properly, the point at which it changes color (its end point) coincides with the equivalence point To understand how and why an indicator changes color, we need to understand the equilibrium principle involved. 

An acid-base indicator is derived from a weak acid HIn  

It s called an end point because that s when you stop the titration. 

From this expression, it follows that the ratio [HIn]/[In ] and hence the color of an indicator depends on two factors. 

Ka of the indicator. Because K varies from one indicator to another, different indicators change colors at different pHs. A color change occurs when 

Methyl red goes from red at low pH, to orange at about pH 5, to yellow at high pH. 

Bromthymol blue is yellow at low pH, blue at high pH, and green at about pH 7. 

Calculate the pH of a solution obtained by mixing 400. mL of a 0.200 M acetic acid solution and 100. mL of a 0.300 M sodium hydroxide solution. 

Sodium hydroxide, NaOH, is a strong base, so it reacts with acetic acid, CH3COOH, to form sodium acetate, NaCH3COO. If an appreciable amount of excess acetic acid is still present after the sodium hydroxide has reacted, the excess acetic acid and the newly formed sodium acetate solution form a buffer solution. 

We first calculate how much of the weak acid has been neutralized. The numbers of milU-moles of CH3COOH and NaOH mixed are calculated as 

Not enough NaOH is present to neutralize all of the CH3COOH, so NaOH is the limiting reactant. 

Because NaCH3COO is a soluble salt, it provides 30.0 mmol CHsCOO to the solution. This solution contains a significant amount of CH3COOH not yet neutralized and a significant amount of its conjugate base, CH3COO . We recognize this as a buffer solution, so we can use the Henderson-Hasselbalch equation to find the pH. 

The equivalence point is the point at which the acid has been neutralized completely by the added base. The equivalence point in a titration can be determined by monitoring the pH over the course of the titration, or it can be determined using an acid-base indicator. An acid-base indicate is usually a weak raganic acid or base for which the ionized and un-ionized forms are different colors. 

ConsidCT a weak oiganic acid that we will refer to as Hln. To be an effective acid-base indi-cattv, Hln and its conjugate base, ln, must have distinctly different colors. In solution, the acid ionizes to a small extent  

The end point is where the color changes. The equivalence point is where neutralization is complete Experimentally, we use the end point to estimate the equivalence point 

In a sufficiently acidic medium, the ionization of Hln is suppressed according to Le Chatelio s principle, and the preceding equilibrium shifts to the left. In tins case, the color of the. solution will be that of HIil In a basic medium, on the other hand, the equilitaium shifts to the right and the color of the solution will be that of the conjugate base. In.  

The end point of a titration is the point at which the color of the indicator changes. Not all indicates change color at the same pH, however, so the choice of indicator ftx a particular titration depends on the strength of the acid (and the base) used in the titration. To use the end point to determine the equivalence point of a titration, we must select an appropriate indicator. 

The endpoint of an indicator does not occur at a specific pH rather, there is a range of pH over which the color change occnrs. In practice, we select an indicator whose color change occnrs 

To see how molecules such as phenolphthalein function as indicators, consider the following equilibrium for some hypothetical indicator HIn, a weak acid withTf = 1.0 x 10 . 

An acid-base indicator is a substance whose color depends on the pH of the solution to which it is added. Several of the photographs in this and the preceding chapter have shown acid-base indicators in use. The indicator chosen 

Acid-base indicators exist in two forms (1) a weak acid, represented symbolically as HIn and having one color, and (2) its conjugate base, represented as In and having a different color. When just a small amount of indicator is added to a solution, the indicator does not affect the pH of the solution. Instead, the ionization equilibrium of the indicator is itself affected by the prevailing [H3O ] in solution. 

From Le Chatelier s principle, we see that increasing [HsO ] in a solution displaces the equilibrium to the left, increasing the proportion of HIn and hence the add color. Decreasing [H3O ] in a solution displaces the equilibrium to the right, increasing the proportion of In and hence the base color. The color of the solution depends on the relative proportions of the acid and base. The pH of the solution can be related to these relative proportions and to the pK of the indicator by means of an equation similar to equation (17.7). 

An acid-base indicator is usually prepared as a solution (in water, ethanol, or some other solvent). In acid-base titrations, a few drops of the indicator solution are added to the solution being titrated. In other applications, porous paper is impregnated with an indicator solution and dried. When this paper is moistened with the solution being tested, it acquires a color determined by the pH of the solution. This paper is usually called pH test paper. 

Acid Color Intermediate Color Base Color 

The equivalence point, as we have seen, is the point at which the number of moles of OH ions added to a solution is equal to the number of moles of ions originally present. To determine the equivalence point in a titration, then, we must know exactly how much volume of a base to add from a buret to an acid in a flask. One way to achieve this goal is to add a few drops of an acid-base indicator to the acid solution at the start of the titration. You will recall from Chapter 4 that an indicator is usually a weak organic acid or base that has distinctly different colors in its nonionized and ionized forms. These two forms are related to the pH of the solution in which the indicator is dissolved. The end point of a titration occurs when the indicator changes color. However, not all indicators change color at the same pH, so the choice of indicator for a particular titration depends on the nature of the acid and base used in the titration (that is, whether they are strong or weak). By choosing the proper indicator for a titration, we can use the end point to determine the equivalence point, as we will see below. 

Typical indicators change cdor over the pH range given by pH = p/C, 1, where K, is the acid ionization of the indicator. 

Strategy The choice of an indicator for a particular titration is based on the fact that its pH range for color change must overlap the steep portion of the titration curve. Otherwise we eannot use the eolor change to locate the equivalence point. 

The endpoint is defined by the change in color of the indicator. The equivalence point is defined by the reaction stoichiometry. 

Although NH4+ will dissociate, it is such a weak acid that [H+] will be determined simply by the excess H +  

The results of these calculations are shown in Table 8.2. The pH curve is shown in Fig. 8.5. 

We can split the fraction term in Equation (6.51) by employing the laws of logarithms, to yield 

The term Tog10 Ka should remind us of pKd (Equation 6.52), and the term log10[H3O+] will remind us of pH in Equation (6.20), so we rewrite Equation (6.52) as 

Litmus is a naturally occurring substance obtained from lichen. It imparts an intense colour to aqueous solutions. In this sense, the indicator is a dye whose colour is sensitive to the solution pH. 

If the solution is rich in solvated protons (causing the pH to be less than 7) then litmus has an intense red colour. Conversely, a solution rich in hydroxide ions (with a pH greater than 7) causes the litmus to have a blue colour. 

To the practical chemist, the utility of litmus arises from the way its colour changes as a function of pH. Placing a single drop of litmus solution into a beaker of solution allows us an instant test of the acidity (or lack of it). It indicates whether the pH is less than 7 (the litmus is red, so the solution is acidic), or the pH is greater than 7 (the litmus is blue, so the solution is alkaline). Accordingly, we call litmus a pH indicator. 

Range Indicator Lower Color Upper Color 

NH4 /NH3. Greater amounts of protons can be added especially in the range that is most rounded out without the level changing noticeably. However, if the proton reservoir is completely full (equivalence point), a drastic change of proton potential occurs when more protons are added. However, this change slows down in the funnel area of the exponential horn.  

Acid-base pairs with strongly contrasting colors are also interesting. Normally, these are large, water-soluble organic molecules. In tiny amounts, they are used as indicators. By themselves and in equal amounts in a solution, the members of these 

When the proton potential in a solution is raised by adding an excess quantity of the acid of a stronger acidic pair, for example, the indicator base disappears due to 

7 Consequences of Mass Action Acid-Base Reactions 

Experiment 7.3 Acidity effect of mineral water. If the indicator bromocresoi green is put into a bottle of very cold mineral water, the solution turns yellow. When the bottle is opened at room temperature or the content heated, a large portion of the carbon dioxide escapes and the indicator color changes to green and finally to an intense blue. 

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The colorations produced in this reaction arise from the action of nitrous acid on the phenol, giving />-nitrosophenol (I) which then reacts with excess of phenol to form an indophenol (II) which is an acid-base indicator ... 

H H Acid-base indicator recommended for titration of acids... 

Use acid-base indicator solutions. Oxidation causes bleaching of indicator to colorless... 

Harvey, D. T. Statistical Evaluation of Acid/Base Indicators, /. Chem. Educ. 1991, 68, 329-331. 

Standardization—External standards, standard additions, and internal standards are a common feature of many quantitative analyses. Suggested experiments using these standardization methods are found in later chapters. A good project experiment for introducing external standardization, standard additions, and the importance of the sample s matrix is to explore the effect of pH on the quantitative analysis of an acid-base indicator. Using bromothymol blue as an example, external standards can be prepared in a pH 9 buffer and used to analyze samples buffered to different pHs in the range of 6-10. Results can be compared with those obtained using a standard addition. 

Why is the acid-base indicator methyl red added to the solution ... 

The plT at which an acid-base indicator changes color is determined by its acid dissociation constant. For an indicator that is a monoprotic weak acid, ITIn, the following dissociation reaction occurs... 

A list of several common acid-base indicators, along with their piQs, color changes, and pH ranges, is provided in the top portion of Table 9.4. In some cases. 

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

The acidity constant for an acid-base indicator was determined by preparing three solutions, each of which has a total indicator concentration of 5.00 X 10- M. The first solution was made strongly acidic with HCl and has an absorbance of 0.250. The second solution was made strongly basic and has an absorbance of 1.40. The pH of the third solution was measured at 2.91, with an absorbance of 0.662. What is the value of K, for the indicator ... 

Ramsing and colleagues developed an FfA method for acid-base titrations using a carrier stream mixture of 2.0 X f0 M NaOH and the acid-base indicator bromthymol blue. Standard solutions of HCl were injected, and the following values of Af were measured from the resulting fiagrams. 

The indicator method is especially convenient when the pH of a weU-buffered colorless solution must be measured at room temperature with an accuracy no greater than 0.5 pH unit. Under optimum conditions an accuracy of 0.2 pH unit is obtainable. A Hst of representative acid—base indicators is given in Table 2 with the corresponding transformation ranges. A more complete listing, including the theory of the indicator color change and of the salt effect, is also available (1). 

It turns out that in low-viscosity blending the acdual result does depend upon the measuring technique used to measure blend time. Two common techniques, wliich do not exhaust the possibilities in reported studies, are to use an acid-base indicator and inject an acid or base into the system that will result in a color change. One can also put a dye into the tank and measure the time for color to arrive at uniformity. Another system is to put in a conductivity probe and injecl a salt or other electrolyte into the system. With any given impeller type at constant power, the circulation time will increase with the D/T ratio of the impeller. Figure 18-18 shows that both circulation time and blend time decrease as D/T increases. The same is true for impeller speed. As impeller speed is increased with any impeller, blend time and circulation time are decreased (Fig. 18-19). 

Acid-base indicator Acid (or base). Neutralization is complete as determined by color change of indicator. 

Usually special cases of the full scheme are studied so that only one or two relaxation times are observed. Important examples are a solution of an acid-base solute in the presence of an acid-base indicator, and a buffered solution of an acid—base solute. PP... 

Scheme VIII has the form of Scheme II, so the relaxation time is given by Eq. (4-15)—appjirently. However, there is a difference between these two schemes in that L in Scheme VIII is also a participant in an acid-base equilibrium. The proton transfer is much more rapid than is the complex formation, so the acid-base system is considered to be at equilibrium throughout the complex formation. The experiment can be carried out by setting the total ligand concentration comparable to the total metal ion concentration, so that the solution is not buffered. As the base form L of the ligand undergoes coordination, the acid-base equilibrium shifts, thus changing the pH. This pH shift is detected by incorporating an acid-base indicator in the solution.

You are probably familiar with a variety of aqueous solutions that are either acidic or basic (Figure 4.6). Acidic solutions have a sour taste and affect the color of certain organic dyes known as acid-base indicators. For example, litmus turns from blue to red in acidic solution. Basic solutions have a slippery feeling and change the colors of indicators (e.g., red to blue for litmus). 

The objective of the titration is to determine the point at which reaction is complete, called the equivalence point. This is reached when the number of moles of OH- added is exactly equal to the number of moles of acetic acid, HC O originally present To determine this point, a single drop of an acid-base indicator such as phenolphthalein is used. It should change color (colorless to pink) at the equivalence point. 

A less accurate but more colorful way to measure pH uses a universal indicator, which is a mixture of acid-base indicators that shows changes in color at different pH values (Figure 13.5, p. 359). A similar principle is used with pH paper. Strips of this paper are coated with a mixture of pH-sensitive dyes these strips are widely used to test the pH of biological fluids,... 

We will also consider the equilibrium involved when an acid-base indicator is used to estimate pH (Section 14.2). ... 

Red cabbage juice, a natural acid-base indicator. The picture shows (left to right) its colors at pH 1,4,7,10. and 13. 

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