Acids and Bases
Standardization of NaOH and Determination of Unknown Acid
KEYWORDS: quantitative analysis, titration, buret, endpoint, standardization, half-equivalence point, calorimetric titration, potentiometric titration ABSTRACT: The concentration of sodium hydroxide was determined by colorimetric titration, and the identity of an unknown acid was determined by potentiometric titration. In the first titration, a strong acid standard, potassium hydrogen phthalate (KHP), was used, to determine the concentration of a strong base, sodium hydroxide (NaOH). In order to do so, we prepared NaOH solution, prepared a buret, and standardized this solution by performing a colorimetric titration of the KHP with the solution until a color change was present. The color change was introduced by an indicator known as phenolphthalein, which caused the solution to go from colorless to pink, which marked the endpoint of this titration. As a result, the concentration of the NaOH solution was found to be 0.124 M with a standard deviation of 0.004 M. In the second titration, the goal was to use both an acid’s formula weight and acid dissociation equilibrium constant (pKa) value to determine the identity of this unknown acid. Both values were found by preparing the unknown acid solution, preparing a buret, titrating the acid solution with the strong basic NaOH solution from the previous experiment, and determining the pH at various points using the pH meter. The formula weight of our unknown acid was found to be189.93 g and the pKa was 7.2 and 7.05 using two different methods. We were able to conclude that our proposed identity of our unknown acid was MOPS, which had a formula weight of 209.26 g and pKa of 7.2.
Introduction
Concepts
Through both the standardization of sodium hydroxide and determining unknown acids, the titration process plays a major role. Titration is one of the types of quantitative analysis, which refers to the determination of how much a given component is present in the sample and is then used to determine unknown concentration of various chemicals. It is a very delicate and precise procedure that requires patience and specific observation by which the solution of a specific reactant is judiciously added to another reactant. In order for some solutions to be accurately prepared by mass, it is sometimes necessary to measure the concentration of a diluted solution by using it to titrate to a known amount of acid, which is known as standardization. The concentration of sodium hydroxide is standardized using potassium hydrogen phthalate (KHP). An indicator, which is a weak acid or weak base, is used in titration to determine the endpoint of the reaction as it changes the color of the solution, which is then used to determine the concentration of the unknown. The endpoint occurs when the moles of acid is equivalent to the moles of base. The indicator for the first titration was phenolphthalein. The colorless phenolphthalein is the protonated form, which is the conjugate acid, and the pink phenolphthalein is the deprotonated form, which is the conjugate base.
Chemical Equations
The chemical equation that acts as an acid and base reaction is: NaOH(aq) + HCl NaCl(aq) + H2O(l) (1) Acids are seen as substances that contained hydrogen and bases as substances that contained hydroxide. In a reaction between an acid and a base, hydrogen ions from the acid react with hydrogen ions from the base forming water. The chemical equation that represents the ionization of a weak acid is HA + H2O H3O+ + A- (2) which is also known as a Brønstead-Lowry acid-base interaction. In this reaction, HA represents the free acid, H3O+ and A- represent its dissociation products, and H2O acts as the base.
Mathematical Equations Once one has completed the titration, it is then possible to use the mass of the acid, molecular weight of the base, and the volume of the base to determine the concentration of the base, which is represented by the equation: maMWa=Cb×Vb (3)
One can simply solve for the unknown concentration by plugging in the other known values. Mass of the acid (ma) is expressed in grams, molecular weight of the acid (MWa) is expressed in grams per mole, volume of the base (Vb) is expressed in liters, and the concentration of the base (Cb) is expressed in molarity. The equilibrium constant known as the acid dissociation constant (Ka) is expressed as the equation: Ka=H3O+[A-][HA] (4)
This equation is used to measure how weak an acid is. The dissociation products are H3O and A- and HA is the free acid. The concentration of protonated acid is shown as [HA] and the deprotonated acid is shown as [A-] and these are all expressed in molarity.1 The logarithmic measure of the acid dissociation constant, also known as pKa, is expressed using the equation: pKa= -log(Ka) (5)
The pKa measures how weak or strong an acid is. The lower the pKa, the stronger the acid is.
Rationale
These two titration experiments involved some unknowns that were left up to the experimenter to figure out. Sodium hydroxide (NaOH) absorbs moisture from the air so it is said to be hygroscopic. In order to avoid this from affecting both it’s concentration and it’s mass, the molarity of NaOH must be determined by standardization to avoid excess error. For the first experiment, the NaOH solution was prepared by standardizing and titrating the solution to find the unknown variables to then use the formula that aids in determining the unknown concentration of the base. The second experiment involves preparing the unknown acid solution, titrating this unknown solution, and then using the pH meter to determine the formula weight and the pKa. The next task is to compare these values to a list of proposed identities in order to correctly establish what the name of your unknown acid is.
Materials & Methods
Materials
The chemicals used in this experiment were sodium hydroxide (Fisher Scientific), potassium hydrogen phthalate (Ricca Chemical Company), phenolphthalein (Ricca Chemical Company), and an unknown acid.
Methods
Standardization of NaOH
Preparation of Solution
The first step in this procedure was to prepare 0.1 M of NaOH solution. In order to do this, 1 gram of NaOH was massed out using a Top-loading balance and dissolved in 250 mL of de-ionized (DI) water with manual or magnetic stirring. If NaOH was spilled onto or near the balance, it was transferred to a waste beaker and rinsed into the basic waste. The NaOH solution was stored in a 250 ml plastic Nalgene bottle with cap due to the fact that strong bases would etch glass. After doing so, the bottle was labeled with the chemical, concentration, the experimenter’s name and date using label tape.
Standardization
The next major step in this procedure was to prepare a buret with the NaOH solution. Next, 0.4 g of KHP was massed out using an analytical balance. It is to note that KHP is oven-dried which is why it is stored in the desiccator cabinet to reduce the absorption of water from air. Next, the KHP is transferred to a 125 mL Erlenmeyer flask and added 25 mL of DI water to dissolve the solid. Then, 3 drops of phenolphthalein is added to the KHP solution. Next, KHP is titrated with NaOH until a pale pink color appears. Finally, record the volume reading of NaOH at the end of the titration. Three replicate titrations were completed to ensure accuracy. At the end of the experiment, NaOH solution was stored tightly closed and labeled for any future experiments conducted.
Identification of Unknown Acid
Preparation of Solution
This experiment began by obtaining an unknown sample assignment from a mentor and the unknown color was recorded. Next, 0.4 g of unknown acid was massed out using an analytical balance, which was indicated on the unknown bottle. Then, the unknown sample was transferred to a clean 250-mL beaker. Next, the solid was dissolved in 50 mL of DI water using a stir bar.
Titration
The titration began by setting up and calibrating a pH probe and starting a LoggerPro Experiment. Next, a buret was prepared with the previously standardized NaOH solution. The beaker with the unknown acid was then placed on a stir plate underneath the buret, immersed the pH probe into the beaker, and a small stir bar was added. The pH values were then monitored and recorded using LoggerPro as the titration began by adding little increments of NaOH at a time of until you had passed the endpoint and completed the titration. After both the titration curve graph and the first derivative plot were made and printed out, the titration of the unknown acid was repeated to ensure reproducibility.
Results & Discussion
Standardization of NaOH
Through preparing 0.1 M of NaOH solution, standardizing, and titrating this solution, the goal of figuring out the concentration of NaOH was calculated using equation 1. By knowing the mass of KHP, formula weight of KHP, and volume of base, the concentration of NaOH was determined using this equation. The average molarity of NaOH was 0.124 M and the standard deviation 0.004 M and these variables came from analyzing data from Table 1. A small standard deviation indicates a high level of precision. Some major sources of error in this experiment may have been that one was not able to get the exact amount of KHP while transferring the solute into the 150 mL Erlenmeyer flask, not having exactly 1 g of NaOH due to using the Top-Loading Balance, and finally some trials of the three titrations might have been slightly less or slightly more pink than the other which would cause the final concentration to be not completely accurate.
Identification of Unknown Acid
By preparing and titrating the unknown acid, the goal of identifying the unknown acid from a list of common weak acids was met by comparing the formula weight and the pKa of the unknown acid to other known acids. By looking at Tables 2 and 3, it is apparent that the formula weight of the unknown acid was 189.93 g. By analyzing Figure 1, the pKa was determined to be 7.2 by using the half equivalence point method and 7.05 using the initial pH method by analyzing. By comparing this formula weight and pKa to that of other weak acids, it was able to be concluded that the unknown acid most closely related to MOPS, which was 209.26 g and had a pKa of 7.2 using both the half equivalence method and the initial pH method. The percent error of the formula weight was 9.24%, the pKa using the half equivalence method was 0%, and the pKa using the initial pH method was 2.08%. Since the percent errors were fairly low across the board, reporting that the unknown acid was MOPS was done with high confidence because it was obvious that the findings were fairly accurate.
Figures
Figure 1. Titration Curve Graph. This graph was used to determine that the pKa of the unknown acid was 7.2 using the half equivalence method and 7.05 using the initial pH method.
Figure 2. First Derivative Plot. The peak in this graph states that the equivalence point lies at about 17 mL, which you could then divide by two to determine that the half equivalence point is about 8.5 mL. Using the half-equivalence method, the pH of the unknown acid was determined to be 7.2.
Tables | Mass of KHP (g) | # of moles of KHP (mol) | Initial Buret Reading (mL) | Final Buret Reading (mL) | Volume of NaOH used in titration (mL) | # of moles of NaOH (mol) | Calculated molarity of your NaOH solution (M) | Titration 1 | 0.43 | 0.0021 | 2.49 | 18.79 | 16.30 | 0.0021 | 0.1288 | Titration 2 | 0.37 | 0.00181 | 18.79 | 33.79 | 14.91 | 0.00181 | 0.1214 | Titration 3 | 0.40 | 0.00196 | 33.70 | 49.94 | 16.23 | 0.00196 | 0.121 |
Table 1: This table represents the three replicates that were performed during the standardization of NaOH. These values were used to calculate the unknown concentration of NaOH, which was 0.124 M with a standard deviation of 0.004 M, which was highly precise. | Mass of unknown acid used (g) | Volume of water added to dissolve the unknown (mL) | Molarity of titrant used (M) | Volume of titrant used (initial vol., final vol.) (mL) | Titrant volume at equiv. point (mL) | pH at the equiv. point | Titrant volume at the half equiv. point (mL) | pH at the half equiv. point | Initial pH before any titrant is added | Titration 1 | 0.3886 | 50 | 0.124 ± 0.004 | 3.97, 32.70 | 16.5 | 10.4 | 8.25 | 7.2 | 4.22 | Titration 2 | 0.3975 | 50 | 0.124 ± 0.004 | 1.40, 30.22 | 17.7 | 8.9 | 8.85 | 7.4 | 4.68 |
Table 2: This table represents the values found during the titration of the 2 replicates. Using these values, one was able to conclude that the formula weight was 189.93 g and the pKa was 7.2. | Unknown | Proposed Identity | Percent Error | Color Code Name | Pink | MOPS | N/A | Formula Weight | 189.93 g | 209.26 g | 9.24% | pKa(half equiv. method) | 7.2 | 7.2 | 0% | pKa (initial pH method) | 7.05 | 7.2 | 2.08% |
Table 3: This table represents the final formula weight, pKa values found, and the percent errors. The values of the unknown acid which were that the formula weight was 189,93 g, and a pKa of 7.2 and 7.05 using two different methods were compared to a list of weak acids. The unknown acid best compared with MOPS, which had a formula weight of 209.26 and a pKa of 7.2, which was highly accurate.
References
(1) Neuman, R.C.; Kauzmann, W.; Zipp, A. The Journal of Physical Chemistry. 1973, 77, 2687-2691.
References: (1) Neuman, R.C.; Kauzmann, W.; Zipp, A. The Journal of Physical Chemistry. 1973, 77, 2687-2691.
KEYWORDS: quantitative analysis, titration, buret, endpoint, standardization, half-equivalence point, calorimetric titration, potentiometric titration ABSTRACT: The concentration of sodium hydroxide was determined by colorimetric titration, and the identity of an unknown acid was determined by potentiometric titration. In the first titration, a strong acid standard, potassium hydrogen phthalate (KHP), was used, to determine the concentration of a strong base, sodium hydroxide (NaOH). In order to do so, we prepared NaOH solution, prepared a buret, and standardized this solution by performing a colorimetric titration of the KHP with the solution until a color change was present. The color change was introduced by an indicator known as phenolphthalein, which caused the solution to go from colorless to pink, which marked the endpoint of this titration. As a result, the concentration of the NaOH solution was found to be 0.124 M with a standard deviation of 0.004 M. In the second titration, the goal was to use both an acid’s formula weight and acid dissociation equilibrium constant (pKa) value to determine the identity of this unknown acid. Both values were found by preparing the unknown acid solution, preparing a buret, titrating the acid solution with the strong basic NaOH solution from the previous experiment, and determining the pH at various points using the pH meter. The formula weight of our unknown acid was found to be189.93 g and the pKa was 7.2 and 7.05 using two different methods. We were able to conclude that our proposed identity of our unknown acid was MOPS, which had a formula weight of 209.26 g and pKa of 7.2.
Introduction
Concepts
Through both the standardization of sodium hydroxide and determining unknown acids, the titration process plays a major role. Titration is one of the types of quantitative analysis, which refers to the determination of how much a given component is present in the sample and is then used to determine unknown concentration of various chemicals. It is a very delicate and precise procedure that requires patience and specific observation by which the solution of a specific reactant is judiciously added to another reactant. In order for some solutions to be accurately prepared by mass, it is sometimes necessary to measure the concentration of a diluted solution by using it to titrate to a known amount of acid, which is known as standardization. The concentration of sodium hydroxide is standardized using potassium hydrogen phthalate (KHP). An indicator, which is a weak acid or weak base, is used in titration to determine the endpoint of the reaction as it changes the color of the solution, which is then used to determine the concentration of the unknown. The endpoint occurs when the moles of acid is equivalent to the moles of base. The indicator for the first titration was phenolphthalein. The colorless phenolphthalein is the protonated form, which is the conjugate acid, and the pink phenolphthalein is the deprotonated form, which is the conjugate base.
Chemical Equations
The chemical equation that acts as an acid and base reaction is: NaOH(aq) + HCl NaCl(aq) + H2O(l) (1) Acids are seen as substances that contained hydrogen and bases as substances that contained hydroxide. In a reaction between an acid and a base, hydrogen ions from the acid react with hydrogen ions from the base forming water. The chemical equation that represents the ionization of a weak acid is HA + H2O H3O+ + A- (2) which is also known as a Brønstead-Lowry acid-base interaction. In this reaction, HA represents the free acid, H3O+ and A- represent its dissociation products, and H2O acts as the base.
Mathematical Equations Once one has completed the titration, it is then possible to use the mass of the acid, molecular weight of the base, and the volume of the base to determine the concentration of the base, which is represented by the equation: maMWa=Cb×Vb (3)
One can simply solve for the unknown concentration by plugging in the other known values. Mass of the acid (ma) is expressed in grams, molecular weight of the acid (MWa) is expressed in grams per mole, volume of the base (Vb) is expressed in liters, and the concentration of the base (Cb) is expressed in molarity. The equilibrium constant known as the acid dissociation constant (Ka) is expressed as the equation: Ka=H3O+[A-][HA] (4)
This equation is used to measure how weak an acid is. The dissociation products are H3O and A- and HA is the free acid. The concentration of protonated acid is shown as [HA] and the deprotonated acid is shown as [A-] and these are all expressed in molarity.1 The logarithmic measure of the acid dissociation constant, also known as pKa, is expressed using the equation: pKa= -log(Ka) (5)
The pKa measures how weak or strong an acid is. The lower the pKa, the stronger the acid is.
Rationale
These two titration experiments involved some unknowns that were left up to the experimenter to figure out. Sodium hydroxide (NaOH) absorbs moisture from the air so it is said to be hygroscopic. In order to avoid this from affecting both it’s concentration and it’s mass, the molarity of NaOH must be determined by standardization to avoid excess error. For the first experiment, the NaOH solution was prepared by standardizing and titrating the solution to find the unknown variables to then use the formula that aids in determining the unknown concentration of the base. The second experiment involves preparing the unknown acid solution, titrating this unknown solution, and then using the pH meter to determine the formula weight and the pKa. The next task is to compare these values to a list of proposed identities in order to correctly establish what the name of your unknown acid is.
Materials & Methods
Materials
The chemicals used in this experiment were sodium hydroxide (Fisher Scientific), potassium hydrogen phthalate (Ricca Chemical Company), phenolphthalein (Ricca Chemical Company), and an unknown acid.
Methods
Standardization of NaOH
Preparation of Solution
The first step in this procedure was to prepare 0.1 M of NaOH solution. In order to do this, 1 gram of NaOH was massed out using a Top-loading balance and dissolved in 250 mL of de-ionized (DI) water with manual or magnetic stirring. If NaOH was spilled onto or near the balance, it was transferred to a waste beaker and rinsed into the basic waste. The NaOH solution was stored in a 250 ml plastic Nalgene bottle with cap due to the fact that strong bases would etch glass. After doing so, the bottle was labeled with the chemical, concentration, the experimenter’s name and date using label tape.
Standardization
The next major step in this procedure was to prepare a buret with the NaOH solution. Next, 0.4 g of KHP was massed out using an analytical balance. It is to note that KHP is oven-dried which is why it is stored in the desiccator cabinet to reduce the absorption of water from air. Next, the KHP is transferred to a 125 mL Erlenmeyer flask and added 25 mL of DI water to dissolve the solid. Then, 3 drops of phenolphthalein is added to the KHP solution. Next, KHP is titrated with NaOH until a pale pink color appears. Finally, record the volume reading of NaOH at the end of the titration. Three replicate titrations were completed to ensure accuracy. At the end of the experiment, NaOH solution was stored tightly closed and labeled for any future experiments conducted.
Identification of Unknown Acid
Preparation of Solution
This experiment began by obtaining an unknown sample assignment from a mentor and the unknown color was recorded. Next, 0.4 g of unknown acid was massed out using an analytical balance, which was indicated on the unknown bottle. Then, the unknown sample was transferred to a clean 250-mL beaker. Next, the solid was dissolved in 50 mL of DI water using a stir bar.
Titration
The titration began by setting up and calibrating a pH probe and starting a LoggerPro Experiment. Next, a buret was prepared with the previously standardized NaOH solution. The beaker with the unknown acid was then placed on a stir plate underneath the buret, immersed the pH probe into the beaker, and a small stir bar was added. The pH values were then monitored and recorded using LoggerPro as the titration began by adding little increments of NaOH at a time of until you had passed the endpoint and completed the titration. After both the titration curve graph and the first derivative plot were made and printed out, the titration of the unknown acid was repeated to ensure reproducibility.
Results & Discussion
Standardization of NaOH
Through preparing 0.1 M of NaOH solution, standardizing, and titrating this solution, the goal of figuring out the concentration of NaOH was calculated using equation 1. By knowing the mass of KHP, formula weight of KHP, and volume of base, the concentration of NaOH was determined using this equation. The average molarity of NaOH was 0.124 M and the standard deviation 0.004 M and these variables came from analyzing data from Table 1. A small standard deviation indicates a high level of precision. Some major sources of error in this experiment may have been that one was not able to get the exact amount of KHP while transferring the solute into the 150 mL Erlenmeyer flask, not having exactly 1 g of NaOH due to using the Top-Loading Balance, and finally some trials of the three titrations might have been slightly less or slightly more pink than the other which would cause the final concentration to be not completely accurate.
Identification of Unknown Acid
By preparing and titrating the unknown acid, the goal of identifying the unknown acid from a list of common weak acids was met by comparing the formula weight and the pKa of the unknown acid to other known acids. By looking at Tables 2 and 3, it is apparent that the formula weight of the unknown acid was 189.93 g. By analyzing Figure 1, the pKa was determined to be 7.2 by using the half equivalence point method and 7.05 using the initial pH method by analyzing. By comparing this formula weight and pKa to that of other weak acids, it was able to be concluded that the unknown acid most closely related to MOPS, which was 209.26 g and had a pKa of 7.2 using both the half equivalence method and the initial pH method. The percent error of the formula weight was 9.24%, the pKa using the half equivalence method was 0%, and the pKa using the initial pH method was 2.08%. Since the percent errors were fairly low across the board, reporting that the unknown acid was MOPS was done with high confidence because it was obvious that the findings were fairly accurate.
Figures
Figure 1. Titration Curve Graph. This graph was used to determine that the pKa of the unknown acid was 7.2 using the half equivalence method and 7.05 using the initial pH method.
Figure 2. First Derivative Plot. The peak in this graph states that the equivalence point lies at about 17 mL, which you could then divide by two to determine that the half equivalence point is about 8.5 mL. Using the half-equivalence method, the pH of the unknown acid was determined to be 7.2.
Tables | Mass of KHP (g) | # of moles of KHP (mol) | Initial Buret Reading (mL) | Final Buret Reading (mL) | Volume of NaOH used in titration (mL) | # of moles of NaOH (mol) | Calculated molarity of your NaOH solution (M) | Titration 1 | 0.43 | 0.0021 | 2.49 | 18.79 | 16.30 | 0.0021 | 0.1288 | Titration 2 | 0.37 | 0.00181 | 18.79 | 33.79 | 14.91 | 0.00181 | 0.1214 | Titration 3 | 0.40 | 0.00196 | 33.70 | 49.94 | 16.23 | 0.00196 | 0.121 |
Table 1: This table represents the three replicates that were performed during the standardization of NaOH. These values were used to calculate the unknown concentration of NaOH, which was 0.124 M with a standard deviation of 0.004 M, which was highly precise. | Mass of unknown acid used (g) | Volume of water added to dissolve the unknown (mL) | Molarity of titrant used (M) | Volume of titrant used (initial vol., final vol.) (mL) | Titrant volume at equiv. point (mL) | pH at the equiv. point | Titrant volume at the half equiv. point (mL) | pH at the half equiv. point | Initial pH before any titrant is added | Titration 1 | 0.3886 | 50 | 0.124 ± 0.004 | 3.97, 32.70 | 16.5 | 10.4 | 8.25 | 7.2 | 4.22 | Titration 2 | 0.3975 | 50 | 0.124 ± 0.004 | 1.40, 30.22 | 17.7 | 8.9 | 8.85 | 7.4 | 4.68 |
Table 2: This table represents the values found during the titration of the 2 replicates. Using these values, one was able to conclude that the formula weight was 189.93 g and the pKa was 7.2. | Unknown | Proposed Identity | Percent Error | Color Code Name | Pink | MOPS | N/A | Formula Weight | 189.93 g | 209.26 g | 9.24% | pKa(half equiv. method) | 7.2 | 7.2 | 0% | pKa (initial pH method) | 7.05 | 7.2 | 2.08% |
Table 3: This table represents the final formula weight, pKa values found, and the percent errors. The values of the unknown acid which were that the formula weight was 189,93 g, and a pKa of 7.2 and 7.05 using two different methods were compared to a list of weak acids. The unknown acid best compared with MOPS, which had a formula weight of 209.26 and a pKa of 7.2, which was highly accurate.
References
(1) Neuman, R.C.; Kauzmann, W.; Zipp, A. The Journal of Physical Chemistry. 1973, 77, 2687-2691.
References: (1) Neuman, R.C.; Kauzmann, W.; Zipp, A. The Journal of Physical Chemistry. 1973, 77, 2687-2691.