Titration of HCL with NaOH
Practical report - Titration of hydrochloric acid with Sodium Hydroxide
Caution: Hydrochloric acid, as well as Sodium Hydroxide, are both very strong acid/base and harmful to skin and eyes. If any contact to the human body would occur, that section of the body needs to be washed thoroughly with a good amount of water and taken to the emergency room if necessary.
The purpose : The purpose of this practical experiment is to go through the process of neutralization reaction with the acid HCl and Base NaOH.
Background information
The practical was an acid-base neutralization titration in which HCL (acid) and NaOH (base) were used in the experiment. (1, 2) A titration is a chemical technique in which a reagent called a “Titrant” of known concentration also called a standardized solution is used to determine the concentration of an analyte or unknown concentration of a known concentration. Considering the fact that we know what the chemical is, we will know how it will react and thus we can use the reaction to determine the concentration of the solution. (2) In this case, we have an unknown concentration of acid, we can use a known concentration of hydroxide base and this type of action is called a neutralization reaction, where salt and water are products of the reaction.
Acid + Base ⇒ Salt + Water
HCL + NaOH ⇒ NaCL + H2O
We will use BTB, which is a chemical pH indicator that will changes color depending on pH changes to show us when the solution has been fully neutralized. It is good to consider that BTB on its own is a bit acidic which is why when it is added to HCL, the solution turns yellow. (3) The point in which all the acid will be absorbed and no excess base will remain in the solution is called the equivalence point. When the equivalence point will be reached, we will be able to use that state of the solution to determine the initial concentration of acid using a series of calculations. The goal of the titration is to reach as close as possible to the equivalence point by carefully adding the base, which will ensure that the calculated acid concentration is as close to the true value of neutralization reaction as possible. In this case, three titration tries were made until we reached the point of neutralization reaction, where the solution turned green.
Materials
Method
Three titration processes were completed with the final one being a success. Thus meaning that even if the solution of the base is 1%greater than the acid, the experiment becomes a fail and another trial needs to be done. If too much base will be added to the acid, the solution will turn blue, but just the right amount will turn the solution green.[1] The experiment was set up with a stand holding the Burette with its clip, whilst the E-flask was set directly under on a stable table [2] 37 dm3 of NaOH was poured directly into the ±0.1cm3 Burette as well as 15,00 cm3 was measured with a 100 ml ±0.1cm3 measuring cylinder and then poured into the 200 ml ±0.5 cm3 E-flask [3] First trial began with rough titration, where fast drops of NaOH were dropped straight into HCL and we saw a failed trial around 25 ml of NaOH in HCL [4] The blue solution was poured into a separate E-flask and the Burette was filled up to 37 ml of NaOH again [5] Second trial began with a much slower titration with approximately 1 drop each second and the experiment failed again at 25 ml of NaOH again [6] The blue solution was poured into a separate E-flask and the Burette was filled up to 37 ml of NaOH again [7] The third trial was a success with approximately 1 drop each 2 seconds with a more patient titration where the equivalence point was reached when only 24.4 ml of NaOH was left in the Burette.
Result
Calculations :
24,4 ml - 12 ml = 12,4 ml ⇒ Neutral
49 - 12 cm3 = 37 cm3 of NaOH
or
We ended with 24,4 ml of NaOH
37 - 24,4 = 12,4 into the acid HCL
NaOH + HCL ⇒ NaCL + H2O
Mol ratio = 1:1:1:1
V(NaOH) = 0,012 dm3
C(NaOH) = 0,013 mol dm3
n(NaOH) = 0,013 x 0,012 = 1,56 x 10-4
C = n/V ⇒ 0,000156/0,015 = 0,0104 mol dm-3
Conclusion and evaluation
From the practical, the conclusion made is that 12.4 ml of NaOH were needed to neutralize and reach the equivalence point of the acidic 15.0 cm3 HCl. The total value of the third trial was pretty accurate considering the first two trials switched quickly at 25 ml ≤ x, meaning a value less than 25 ml, but still pretty close had to be the point at which the titration curve must be turning making it the equivalence point. This experiment needs a lot of patience, which our group needs to improve on considering the practical was suppose to only consist of one rough titration trial and one slow titration trial. The second trial was a ruthless one with not so many eyes on the E-flask to see at which point the color turns, thus it is the reason to why we had to do a third trial. (4) However, just because we think we reached the equivalence point, that might maybe not be the case. Colour change, especially in a titration process, is very slow and specific that even if three students eyes would focus on it, the color sensitivity of each human differs which can be seen as a limitation. Another point would be that as mentioned in the background information BTB is acidic from the start and if we might have used the wrong amount or better said too much of BTB, then the equivalence point had already shifted from the beginning. It is also important to keep in consideration that BTB is not the best indicator in the market and a PH-meter would have been a better option to get the specific quantitative point of the pH directly on the screen, which would then decrease any additional uncertainties on calculation the limitations of which the Burette and the E-flask possess. However, none of the limitations above matter, because the temperature was not even considered and measured during the titration procedure considering the fact that all indicators are affected by the temperature, thus the color might as well have changed slower/faster for us which probably ruined the final values. An additional limitation would be that we tried the process three times in a row and even though we washed and dried off the E-flask and the Burette maybe some small particles were still remaining in the material (maybe even some particles from other experiments which might not have washed off properly), thus affecting the equivalence point as well as the value of the calculations made above. The Burette might have gotten an air bubble blockage inside with the NaOH, which might have flowed out with NaOH into the HCL making us lose the value of the real volume. Other small limitations such as misreading the volume, swirling the E-flask too much at the point of which the shifting point is affected or eye/sight/angle limitations all are to be considered because all of them can shift the equivalence point and ruin the end results.
Reference :
1. Chemistry Libretexts. Titration. Available from: https://chem.libretexts.org/Core/Analytical_Chemistry/Lab_Techniques/Titration [Accessed 10th May 2017]
2. Spark-notes. Acid-Base Titration. Available from: http://www.sparknotes.com/chemistry/acidsbases/titrations/section1.rhtml [Accessed 11th May 2017]
3. Khan Academy. Acid-Base titration curves. Available from: https://www.khanacademy.org/test-prep/mcat/chemical-processes/titrations-and-solubility-equilibria/a/acid-base-titration-curves [Accessed 12th May 2017]
4. Titrations.info. Titration and titrimetric methods. Available from: http://www.titrations.info/titration-errors [Accessed 20th May 2017]
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