Chemistry Titration Acid Base Lab Essay

Question: What effect does an indicators pH range have on the end point of the titration of vinegar and 1.00 mol dm-3 of sodium hydroxide solution? Equation: CH3COOH (aq) + NaOH (aq) à ¯ NaCH3COO (aq) + HOH (l) Data Collection Table A: Table Representing the Initial Volume of NaOH in the Burette, the Final Volume of NaOH in the Burette and the Difference between Those Values for Specific Trials When Using Different Indicators. There were a minimum of three trials performed for this lab as there needed to be three of the same difference between the final and initial burette readings of the sodium hydroxide. This is due to the fact that a titration lab requires one to acquire three of the same values for this section to understand the exact amount of base required to reach the endpoint of the reaction. Indicator Initial Volume of NaOH solution (ml)  ±0.05 Final Volume of NaOH solution (ml)  ±0.05 Difference Between the Final and Initial Burette Readings (Volume of NaOH used) (ml)  ±0.1 Qualitative Observations 1. Calculate the Initial Concentration of Acetic Acid Before Diluted With Distilled Water C1V1 = C2V2 C1= C2V2 V1 C1= (0.09mol/dm3)(0.1L) (0.01L) Concentration of Acetic Acid Before Dilution = 0.9mol/dm3 Sample Calculation B: Calculating Percent Uncertainty for the Volume of Sodium Hydroxide Needed for Different Indicators Example: Bromocresol Green =Ãâ€" =25.0% Sample Calculation C: Calculating the Percent Error Percent error )Ãâ€"100 ) Ãâ€"100 =0.034% Sample Calculation D: Propagating Uncertainty for the Volume of Sodium Hydroxide Required for Each Indicator = (0.9 ±0.1) + (1.5 ±0.1) + (0.2 ±0.1) + (2.9 ±0.1) + (0.4 ±0.1) = 5.9ml ±0.1 Graph A: Titration curve of the Amount of Sodium Hydroxide Required to Reach the Endpoint for Different Indicators Tested Graph B: Titrations curve of a strong base and weak acid. Approximately 9.2 is the pH of the equivalence point as seen in the graph. Conclusion This lab required one to determine the different volumes of sodium hydroxide base required to reach the end point of a titration with acetic acid as the analyte when using different indicators. The equation is as followed: CH3COOH (aq) + NaOH (aq) à ¯ NaCH3COO (aq) + HOH (l) It was determined that phenolphthalein was the best indicator for use in this lab. This is due to the fact that the vinegar that was used in class has an acetic acid concentration of 0.87 mol/L which is 5.0% acetic acid (Meheen, n.a). In the lab, phenolphthalein required 0.90 mol/L of the sodium hydroxide solution in order for a color change to take place (endpoint was reached). This value is fairly close the concentration of acetic acid in the vinegar used in the lab, therefore, phenolphthalein was the most accurate of indicators used in this titration lab. Through research it has been determined that phenolphthalein should have, in fact, been the best indicator of use. Phenolphthalein has a pH range between 8.0 to 9.8 which is an appropriate range for a weak acid strong base titration. This is because the equivalence point for this titration will take place at a point of pH approximate to 9 which falls into the pH range for the indicator phenolphthalein. This can be seen in Graph B. This equivalence point will be greater than 8.7 as the weak acid (vinegar) only partially dissociates and releases a little amount of its hydrogen ions, making it a weak acid. On the other hand, sodium hydroxide contains sodium acetate which raises the pH considerably of the vinegar as it is fairly basic. Therefore, the equivalence point will be greater than seven and more specifically approximate to 9. All of the indicators that were used did not have a pH range above 9 except for phenolphthalein. This is the reason why phenolphthalein was the best indicat or. Graph A illustrates the different volumes of sodium hydroxide that was needed to reach the endpoint of the reaction when being added to a solution of vinegar using various indicators. The least amount of sodium hydroxide required to reach the endpoint of the reaction for the indicators used are displayed as followed: methyl orange, bromocresol green, phenolphthalein, bromothymol blue and methyl red. Many indicators were used to determine which one was right for this specific lab. It was necessary to determine which indicator was correct for this lab. This can be explained by the equivalence and endpoint. The equivalence point is the point of the reaction when one has added the correct amount of the base to the acid. However, the endpoint is the point in which the indicator changes color. Error bars are not present on the graphs displayed as titration labs require one to do continue experimentation for a test using a specific indicator until they have obtained a value for the differen ce between the final and initial burette readings (volume of NaOH used) a minimum of three times. The vinegar had a pH or approximately 2.4 making it highly acidic (Rousseau, 2012). Therefore, the equilibrium will be on the left. However, when the sodium hydroxide basic solution was added to the vinegar, the equilibrium started to shift more to the right. For instance, when base was being added to the vinegar solution with the indicator Bromocresol green, the color of the solution turned from yellow to green. When the green color is seen, the end point has been reached however, if the color becomes blue then over-titration has taken place, therefore shifting equilibrium even further to the right. There were few anomalous results that arose during the course of the lab. For instance, for the test using phenolphthalein, one trial had a considerably larger difference between the final and initial burette readings (volume of NaOH used) of 1.4ml. The other trials all had a difference of 0.9ml for using this specific indicator. This can be seen in all tests using different indicators except one represented in data collection table A, Bromocresol green. This proves that there were systematic and/or random errors that took place during the lab. Evaluation There were few errors that took place during the course of experimentation. These errors could have negligibly affected the results obtained throughout the process. One considerable error that took place was making the sodium hydroxide solution. 1gram of white crystalline sodium hydroxide pellets were required in order to create the solution. Unfortunately, while the pellets were being measured on the electronic balance they interacted with moisture from the air. Additionally, the pellets were put in the volumetric flask when water that would subsequently mix with these pellets was being measured in the graduated cylinder. Therefore, there was lot of time before the water was placed in the volumetric flask for the sodium hydroxide pellets to absorb moisture from the air. This is a concern as solid sodium hydroxide has the formula NaOH (s). Additionally in air there is carbon dioxide (CO2). The formula for the reaction between this carbon dioxide and the solid sodium hydroxide is as f ollows: 2 NaOH + CO2 → Na2CO3 + H2O Therefore, the hydroxide from the sodium hydroxide and the oxygen from carbon dioxide combine to from water (H2O) which affects the results that were obtained in the lab (ATSDR, 2011). This is because the sodium hydroxide concentration in the solution that was made to act as the base, would have decreased as the pellets reacted with air to form water. It would have been beneficial to purchase the specific amount (in grams) of sodium hydroxide pellets that was necessary for this experiment (1gram). This would ensure the sodium hydroxide pellets would not react with the carbon dioxide in air for very long before being put in the volumetric flask with water added. This is because there would be no need to measure these pellets on the electronic balance. A very common error with titration labs is that over-titration often occurs. This is the point in which too much titrant is added to the analyte during a trial. Therefore, the reaction passed the endpoint as too much of the basic sodium hydroxide was added to the acidic vinegar solution containing acetic acid. Unfortunately, this took place for most trials. For instance, phenolphthalein reaches its endpoint when it changes from being clear and transparent to becoming a light shade of pink. Though this may be, for all of our trials when using this indicator the color became a hot pink shade proving that over-titration had taken place. An improvement for this component of the lab is simple. An improvement to control this error would be to use a burette with a smaller opening. This is because the endpoint can be overshot quite easily and this would ensure the error would not take place. It would have been best to manage the amount of base entering the beaker as much as possible with greater accuracy. An additional error that took place was how the stopcock was not entirely effective. This is due to the fact that when the stopcock was turned to close the burette opening, little droplets of sodium hydroxide would still pour into the beaker. This means that if the solution had reached the endpoint and additional droplets were leaked from the beaker an inaccurate reading of the sodium hydroxide in the burette would be recorded. This error could have been improved by using a BT50 digital burette (Bibby, 2010). This burette would ensure that a considerably small amount of sodium hydroxide would be discharged from the device at most. A button will be pressed to stop the base from entering the beaker containing the acid electronically. The accuracy of this instrument is within  ±0.2%, and precision is better than 0.1% allowing for the most accurate of results to be obtained. Temperature should have been controlled throughout the process as it does play a role in affecting the data obtained in a titration lab. The volumetric glassware used for the purpose of this lab is calibrated at 20oC and a higher temperature of a solution would result in it holding a greater volume than desired (Atkins). Since the room temperature is approximately 24oC it can be assumed that the temperature of the solutions used in the experiment were greater than 20oC. Therefore, these solutions can expand and the concentration in molarity would decrease. Therefore it would have been beneficial to keep the temperature of the acid and base constant at 20oC. This could have been achieved by using an alcohol thermometer. If the temperature for the acid and base were not 20oC, than it would either need to be placed in the refrigerator if the solution was too hot or would need to be put in a beaker and on a hot plate to reach the desired temperature. The temperature for each indicator that was used was not consistent throughout the lab. This is because the temperature of the room changed throughout experimentation. The door in the classroom was open to get rid of the stench from a lab done in the previous class. Therefore, the temperature of the classroom was gradually getting colder and colder. A consistent temperature for each indicator would ensure that the pH range for a color change would not increase or decrease. This is because with a decrease in temperature there is less ionization that takes place allowing the pKa value to increase and therefore, the color change will occur at a higher pH. To ensure all the temperatures of the indicators were kept consistent throughout the lab it would have been beneficial to use an alcohol thermometer and place it in the indicator solution before conducting the trial (ChemBuddy, 2009). If the temperature changed for one trial, the indicator would either need to be placed in the refrigerator if the solution was too hot or would need to be put in a beaker and on a hot plate to reach the desired temperature. The table below represents how temperature can affect the pH range for a color change for specific indicators. To ensure this does not occur, the lab should be taken out in a room with no windows so the temperature does not change. The table below represents how the endpoint changes with a change in the temperature of an indicator. Indicator Color Change Range 18oC 100 °C Methyl Orange 3.1-4.4 2.5-3.7 Methyl Red 4.4-6.2 4.0-6.0 Phenolphthalein 8.2-9.8 8.1-9.0 Quoted from Chemia analityczna, J.Minczewski Z.Marczenko, PWN, Warszawa 1973. Another error was that there were bubbles in the burette tube. When the sodium hydroxide solution was poured in the burette it was done quite quickly creating many bubbles in the beaker. Unfortunately, my group could not afford rinsing the solution out and pouring in more sodium hydroxide solution because we wanted to ensure we had enough of the solution for the whole lab. Therefore the bubbles caused an inaccurate volume reading. The bubbles occupy a particular amount of volume and this means that the volume in the burette would in fact be less than the amount reported during trials. Since the bubbles all formed near the stopcock it would have been best to pour the sodium hydroxide solution into a waste beaker in hopes that the bubbles would be eliminated before beginning the trial. This is because the pressure of the titrant in a burette will force the bubbles out. Unfortunately, in some cases the bubbles remain near the stopcock. If this is the case, it would be best to use a suction method approach. This involves one to partially open the stopcock allowing the contents of the burette to discharge in the beaker. Then, they will use the pipet bulb to suck air from the bubbles through the nozzle of the burette. This will drive out the air bubbles present in the solution out of the beaker into the suction device. One must ensure that the sodium hydroxide that was poured in the beaker during this process is accounted for by slowing pouring in the amount necessary in the burette for the trial to begin. One experimental error that was unavoidable was the cleaning of the burette before using it. It was necessary to clean the burette using water, as it may have had some residue or contamination from a previous experiment. Additionally this residue could have affected the pH of the sodium hydroxide solution that was to be put in the beaker for the purpose of this titration lab. However, there were some droplets of water left on the sides of the beaker after it was cleansed which would have decreased the concentration of the sodium hydroxide solution that was subsequently poured in this material and lowered the pH level of the sodium hydroxide solution as water is neutral (pH of 7.0). With a lower concentration of sodium hydroxide in the base, the pH would have decreased therefore allowing more of this base to be added to the vinegar in order for the endpoint to be reached in a reaction. This is due to the fact that sodium hydroxide is a highly basic solution and water is neutral. When they are mixed the pH of the solution will be lower than the initial pH of the sodium hydroxide solution. An improvement for the lab could have been to detect the pH of the acid –base titration once it has reached the endpoint using a pH meter. This was one of the limitations in this acid-base titration lab as the indicators that were used did not show a change in color at an exact value of pH but only changed in a specific range of the pH. For example, the indicator Bromocresol green changes the color of the solution of a pH range between 3.8 to 5.4. This range is considerably large resulting in one to obtain results that are not necessarily the most accurate. An improvement for this component of the lab would have been to use a pH meter. A pH meter is a device that that measures the pH of a solution by determining the voltage of the solution by immersing two electrodes in it. Then, the reading device will present the pH value. This would have led for the most reliable pH readings to be obtained and the correct amount of sodium hydroxide to be added to the acidic solution. The lab could have been furthered in several ways. It would have been interesting to test these different indicators and the amount of sodium hydroxide required for the end points to be shown when using a strong acid and base. At the equivalence point the same amounts of hydrogen and hydroxide ions will form water, therefore having a pH of 7. For instance, if hydrochloric acid and sodium hydroxide were used the following reaction would take place: HCl(aq) + NaOH(aq) –> H2O(l) + NaCl(aq) H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) –> H2O(l) + Na+(aq) + Cl-(aq) H+(aq) + OH-(aq) –> H2O(l) The first equation displays the reaction between the strong acid and strong base. The second reaction displays the HCl and NaOH dissociated in their respective ions. The last equation is known as the net ionic equation which eliminates the spectator ions from the previous equation. This proves that water will be produced. The neutralisation reaction will take place because salt will also be formed. This is because the anion from the dissociation of the strong acid and the cation from the dissociation of the strong base will come together to form the salt. The salt is not seen the net ionic equation as it dissociates. For this acid base titration lab it was legitimate to compare the outcomes with other classmates to determine if there were any anomalies between the data that was obtained. This is because it would have provided one with the knowledge of the types of errors that occurred and how much it affected their data. For instance, Sarah and I compared our results with Rachel Hung and Yashna Lakhani’s group. Yashna and Rachel’s groups provided information upon different indicators and some of the ones my group used as well. When comparing with the data of these groups it was determined that all of the results were mostly exactly the same or off by 0.1ml for the difference between the final and initial burette readings. For instance, for phenolphthalein, the value my group obtained for this was 0.9ml however, Yashna’s group got 1.0ml. This could have been due to a human error known as a parallax. A difference of 0.1ml could have taken place as one of our groups were unable to determine the position of the meniscus on the burette and therefore, the wrong readings of sodium hydroxide solution could have be collected. An acid base titration has several uses. One of the main real life uses of this experiment is to mix compounded drugs. A pharmacist will need to mix drugs appropriately in order for them to be in the appropriate pH range for the human body. Antacids are commonly used to help issues concerning heartburn, acid reflux and more. These feelings most likely take place due to excess hydrochloric acid in the stomach which causes an uncomfortable feeling. This subsequently allows a backflow of this acid to go up the esophagus which can make someone feel like their throat is burning. These tablets counteract the acidity as they go to the stomach area and react with the hydrochloric acid. These tablets are basic and change the liquid in the stomach to being not as highly acidic. Common antacid tablets contain Mg(OH)2 and Al(OH)2. A reaction equation by using an antacid tablet is shown as followed: HCl+NaOHà ¯ NaCl+H2O The acid base titration in this case is used to determine the amount of the stomach acid present in one’s body. Therefore, it can be discovered how much antacid will be needed for someone dealing with the issues mentioned above in order to make the hydrochloric acid present in their stomach less concentrated (Cavite, 2010). This will ensure the pH of the stomach will be less acidic. 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