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A-Level Chemistry AQA CAP: Acids

Brief Theoretical Background

Acid Base Interaction

Chemical reactions between the molecules of acids and bases formed the foundation for the acid-base titration method, thanks to which a number of working problems related to the selection of the indicator, determination of the type of interaction, calculation of the equivalence point are solved in professional chemistry. It should be recalled that titration is a quantitative analytical method designed to determine the analyzed component of a substance (in an acid-base titration, it is acid or base, respectively) dissolved in a probe. This method is based on the visual observation of the complete chemical reaction between the analyzed component and the reagent added to the solution with a predetermined concentration.

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Acids

According to the principles of the theory of electrolytic dissociation, developed by Svante Arrhenius, chemical acids are traditionally understood as substances that release hydrogen ion (proton) H+ during dissociation. The second component of the complete molecule is the acid anion, which can be represented by a whole spectrum of atom groups: sulphate, sulfide, nitrate, cyanides, permanganate and others. However, the Arrhenius theory has been replaced by the supplemented Bronsted and Lowry’s theory, which concluded that acids are substances that definitely give off a proton in the interaction:

Acids

Final to date, corrections to the restrictions of the previous definitions made by Gilbert N. Lewis, who claimed that acids are electrolytes that receive electron pairs from the base, in other words, an electron acceptor.

Bases

It should be recognized that the assumptions characteristic of acids were also fair for the bases. According to the Arrhenius theory, the bases are called molecules, which at dissociation in solutions detach the hydroxide ion OH. In the Bronsted and Lowry’s theory, the bases are substances that take isolated protons from acids, and thus are acceptors of protons. On the other hand, the bases, according to Lewis, can be considered any particles capable of giving away an electron pair to form a covalent bond: electron donor.

Acid Base Titration

According to the acid-base titration concept, during this chemical procedure, the reagent is gradually mixed with a known concentration of the analyzed component. It is known that acids and bases characterize the opposite sides of the pH scale, and therefore, this interaction naturally changes this parameter. For instance, if there is an acid in the titration unit as a component in the flask, then the drip addition of base from burette will undoubtedly increase the pH. Since the acid-base reactions are speedy, the chemical equilibrium is also immediate.

Indicators are used professionally to observe changes in the interaction in the chemical environment. Indicators are chemical molecules — usually quite complex organic substances — that can change their color when the pH of the environment changes. Although it is possible to determine the environment using indicators, they cannot show the exact value of the potential of hydrogen, so it is not uncommon for professionals and students to use digital pH meters that show a specific number.

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Objectives

The purpose of this practical work is to investigate the acid-base titration method on the example of the interaction of ethanoic acid with sodium hydroxide. In the course of the work, the author solves such issues as the chemical titration procedure, the construction of a pH calibration plot, the construction of a titration curve (pH), the calculation of Ka, and the comparison of theoretical values with practical ones.

Calculations and Steps

Of primary importance is the chemical acid-base titration procedure to determine practical pH values. It should be specified that the whole titration process should be divided into three consecutive stages: before, during, and after the equivalence point. Thus, each of the stages has a unique formula for calculating the pH value. At the initial moment, even before adding the first drops of sodium hydroxide, the pH of the solution is determined only by dissociation of acetic acid:

A-Level Chemistry AQA CAP: Acids

However, when sodium hydroxide is added to this solution, the interaction will result in an equivalent amount of acetate-ion, which, in a mixture with acetic acid, forms acetate buffer solution.

A-Level Chemistry AQA CAP: Acids

Therefore, the entire branch of the titration curve to the point of equivalence can be calculated utilising the formula used to calculate the pH of buffer solutions formed by weak acid and its highly soluble salt.

A-Level Chemistry AQA CAP: Acids

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When enough amount of sodium hydroxide was added, the point of equivalence came. At this point, acetic acid fully reacted with sodium hydroxide, and in the solution, there were products of their interaction, namely salt and water. At this point, the pH equals:

A-Level Chemistry AQA CAP: Acids

Starting from this stage, the concentration of OH ions increases in proportion to the added volume of alkali, which dissociates in the solution:

A-Level Chemistry AQA CAP: Acids

Subsequently, data on pH values are recorded in Table 3 (see Appendix A).

Calibration Plot

The secondary task is to build a calibration plot, based on which it will be possible to create a curve of acid-base titration. In order to determine the real pH values for the solution, it is necessary to correlate the potential of hydrogen values for the buffer mixture and the sample under analysis. The practical data showing the three measurements are given in Table 1.

Table 1. Information on three pH measurements for the buffer solution and the probe to be analyzed.

Buffer pH probe
4,00 4,78
7,00 7,50
9,20 9,40

Utilizing these values, the author has drawn a calibration plot in coordinates Probe/Buffer. The points were connected by an extrapolation line extended for several periods forward and backward, as illustrated in Figure 1.

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Calibration graph of the pH ratio of the buffer solution and pH of the analyte.
Figure 1. Calibration graph of the pH ratio of the buffer solution and pH of the analyte.

This graph allows determining real pH values obtained not from the calculations but during the experiment. For this purpose, it is necessary to define the point on the horizontal axis, which is fair for the theoretical pH values of the solution presented in the second column of Table 3 (see Appendix A). A detailed procedure for the correlation between the two pH values is shown in Fig. 2. Then, the actual pH values of the analyzed solution are listed in the third column of Table 3.

Demonstration of the use of a calibration graph to determine actual pH values.
Figure 2. Demonstration of the use of a calibration graph to determine actual pH values.

Creating a titration curve

Operating the first and third columns of Table 3, the author of this work plots the acid-base titration curve. In combination with setting the point of equivalence, this curve allows estimating the dynamics of changes in pH values depending on the volume of added sodium hydroxide to weak acetic acid. As presented in Figure 3, the hypothesis of the nonlinear growth of potential of hydrogen is confirmed.

Acid-base titration curve (pH).
Figure 3. Acid-base titration curve (pH).

Thus, when analyzing plot no. 3, it is possible to make two conclusions:

  1. the point of equivalence for this type of titration has approximate coordinates (25.40,8.2).
  2. near the end of the curve, the dynamics of pH values can be noticed, apparently due to the measurement error or undesirable chemical interactions in the solution.

Ka Calculation

In addition, using the titration curve, the acid dissociation constant (Ka) and base constant (Kb) can be determined. In this practical work, the author defined the Ka value for the point of equivalence calculated using the Henderson-Hasselbalch model:

Ka Calculation

Moreover, it should be considered that the calculation of the acid dissociation constant should be done for the half equivalence point, since in the first area of the curve, up to the equivalence point, the pH values have not changed too intensively. Consequently, the calculation of Ka should be done for a point with the 25.40/2 coordinate. Since the acid and base concentrations are equal here, the logarithm is reset to zero:

A-Level Chemistry AQA CAP: Acids

Which means

A-Level Chemistry AQA CAP: Acids
A-Level Chemistry AQA CAP: Acids

Therefore, the acid dissociation constant is equal to:

A-Level Chemistry AQA CAP: Acids

Or alternatively:

A-Level Chemistry AQA CAP: Acids

Comparison With Reference Value

To estimate the error, the reference value of the constant dissociation characteristic of acetic acid should be used. According to the references, for weak acid CH3COOH, the value of Ka is equal to

A-Level Chemistry AQA CAP: Acids

. Then the measurement error should be determined by the formula:

A-Level Chemistry AQA CAP: Acids
A-Level Chemistry AQA CAP: Acids
A-Level Chemistry AQA CAP: Acids

Risk Assessment

Table 2. Five Step Risk Assessment on HSE.

Hazards Error of devices, experimenter error, violation of titration technique, spoiled reagents, errors in calculations.
Subjects of Harm Direct measurement can be broken due to experimental errors or spoiled reagents, while errors in calculations can lead to incorrect indirect data.
Precautions The risks can be eliminated if the correct operation of the titration plant is assessed beforehand and the expiration dates of the reactants are determined. Close examination of the instructions should inhibit the possibility of arithmetical errors.
Conclusions Although there are risks of practical work, they can be eliminated if one carefully studies the components of the experiment and follows the instructions. It should be taken into account that it is practically impossible to get rid of errors completely.
Check For this step it is necessary to repeat titration taking into account the described conclusions.

Appendix A

Table 3. Five steps of risk assessment for HSE.

Volume of NaOH (cm3) Probe recorded pH Actual pH values
0.00 3.05 2.10
2.00 3.70 2.90
4.00 4.05 3.20
6.00 4.26 3.45
8.00 4.45 3.65
10.00 4.58 3.78
12.00 4.68 3.98
14.00 4.84 4.05
16.00 5.00 4.20
18.00 5.18 4.40
20.00 5.35 4.60
22.00 5.58 4.80
24.00 6.01 5.30
24.20 6.05 5.35
24.40 6.15 5.39
24.60 6.21 5.45
24.80 6.32 5.75
25.00 6.51 5.80
25.20 6.67 6.01
25.40 6.96 6.39
25.60 9.39 9.00
25.80 10.70 10.45
26.00 11.20 11.00
28.00 12.09 11.95
30.00 12.00 11.90
32.00 12.56 12.55
34.00 12.26 12.25
36.00 12.93 12.95
38.00 12.67 12.60
40.00 12.56 12.55

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