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In this investigation for titration, we will be researching the unknown concentration of our sodium hydroxide and determining the concentration of hydrochloric acid by titrating it with sodium carbonate. Titration is a fundamental method for determining the unknown concentration of acids, such as hydrochloric acid. We employed precise laboratory equipment, including a glass pipette for greater accuracy, a top pan balance with a precision of 3 decimal places, and a glass burette, all of which are essential for accurate results in our experiment.
When preparing the sodium carbonate standard solution, we had to adhere to various health and safety considerations to ensure a safe laboratory environment. This included wearing safety goggles to protect our eyes from potential splashes and ensuring that we did not come into contact with the sodium carbonate powder, as it could cause skin irritation. Additionally, we stood at the laboratory tables to minimize the risk of contamination or injury if any equipment were to break or spill.
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For the calibration of the balance, we conducted multiple measurements by adding different masses to the balance in 5-gram increments and recording the mass readings. After five measurements ranging from 5 grams to 25 grams, we determined that while the balance exhibited precision by consistently providing similar results, there was an accuracy issue, with measurements consistently off by at least 0.01 grams.
Before using any equipment, we ensured that all glassware was thoroughly rinsed and dried to minimize the risk of contamination with other substances or liquids.
The mass of sodium carbonate used in the experiment was 1.4 grams.
The molar mass (Mr) of Na2CO3 is calculated as follows:
Mr Na2CO3 = (2 × 23) + 12 + (3 × 16) = 106 g/mol
The number of moles (n) of sodium carbonate can be calculated using the formula:
n = Mass / Mr
n = 1.4 g / 106 g/mol ≈ 0.013 moles
| Size of Mass | Attempt 1 | Attempt 2 | Attempt 3 | Average |
|---|---|---|---|---|
| 5 grams | 4.95 grams | 4.96 grams | 4.95 grams | 4.953 grams |
| 10 grams | 10.45 grams | 10.45 grams | 10.46 grams | 10.453 grams |
| 15 grams | 15.40 grams | 15.40 grams | 15.41 grams | 15.403 grams |
| 20 grams | 20.14 grams | 20.14 grams | 20.14 grams | 20.14 grams |
| 25 grams | 25.09 grams | 25.09 grams | 25.09 grams | 25.09 grams |
The actual mass for our standard solution was 1.4 grams, and our measurements were accurate enough to achieve this value, resulting in a percentage difference of 0%, indicating a match between the actual and measured values. Potential errors in our measurements could have arisen from improper pouring of sodium carbonate into the weighing boat or failure to zero out our balances before measurements, both of which could lead to inaccurate results.
To improve the accuracy and precision of our results, we could have used a balance with a higher precision than 3 decimal places to obtain more precise measurements. Additionally, the use of a magnetic stirrer and a hot surface to dissolve the crystals into the solution could have provided a purer solution and more accurate results.
Furthermore, the utilization of an electronic pipette could have eliminated potential human errors and yielded even more accurate results. Overall, while our measurements were quite accurate, there is room for improvement in the methodology to enhance the precision and reliability of our experiments.
During the titration to standardize hydrochloric acid, we maintained a strong focus on health and safety considerations in our experiment. Handling glass equipment, which is both fragile and expensive, required careful attention to prevent accidents that could harm individuals or result in financial costs. Moreover, hydrochloric acid is known to be an irritant, necessitating cautious handling to avoid skin contact or eye exposure. To ensure safety, we wore protective goggles and handled the acid with care throughout the experiment.
For the calibration of our pipette, we measured out 25 cm³ of water from a beaker and transferred it into an empty beaker placed on a scale that had been zeroed out. We recorded the deviation from the target volume of 25 cm³ for this procedure, repeating it three times to obtain an average. Similarly, we performed this calibration for the burette, measuring and recording the deviation from 50 cm³.
| Equipment | Test 1 | Test 2 | Test 3 | Average |
|---|---|---|---|---|
| Burettes | 47.3 cm³ | 47.17 cm³ | 47.25 cm³ | 47.24 cm³ |
| Pipette | 24.58 cm³ | 24.57 cm³ | 24.47 cm³ | 24.54 cm³ |
| Trial Run | Attempt 1 | Attempt 2 | Attempt 3 | |
|---|---|---|---|---|
| Initial Titre (cm³) | 13.2 cm³ | 27.1 cm³ | 10.3 cm³ | |
| Final Titre (cm³) | 13.2 cm³ | 27.1 cm³ | 41.1 cm³ | 24.2 cm³ |
| Titre (cm³) | 13.2 cm³ | 13.9 cm³ | 14.0 cm³ | 13.9 cm³ |
The ratio of sodium carbonate to hydrochloric acid is 1:2, based on the balanced chemical equation for the reaction.
The concentration of hydrochloric acid (HCl) can be calculated using the formula:
Concentration of HCl (mol/dm³) = Moles / (Volume/1000)
Concentration of HCl (mol/dm³) = 0.0026 moles / (13.93 cm³ / 1000)
Concentration of HCl (mol/dm³) ≈ 0.186 moles per dm³
In this section, we employed two methods to determine the concentration of the sodium hydroxide sample: titration with an indicator and pH readings. The titration method with an indicator was identical to the one used for standardizing hydrochloric acid, and the second method involved using a pH probe to record readings at every 1 cm³ increment.
| Trial | 1st Attempt | 2nd Attempt | 3rd Attempt | |
|---|---|---|---|---|
| Initial (cm³) | 25.4 cm³ | 25.4 cm³ | 25.3 cm³ | 25.4 cm³ |
| Final (cm³) | 25.4 cm³ | 25.4 cm³ | 25.3 cm³ | 25.4 cm³ |
| Titre (cm³) | 25.4 cm³ | 25.4 cm³ | 25.3 cm³ | 25.4 cm³ |
To calculate the average Titre, we sum the individual values and divide by the number of attempts:
Average Titre = (25.4 + 25.3 + 25.4) cm³ / 3 = 76.1 cm³ / 3 = 25.37 cm³
The balanced chemical equation for the reaction between sodium hydroxide (NaOH) and hydrochloric acid (HCl) is:
NaOH + HCl → NaCl + H2O
Moles of HCl can be calculated using the formula:
Moles of HCl = Volume (dm³) × Concentration
Moles of HCl = 25.37 cm³ / 1000 (to convert to dm³) × 0.186 mol/dm³ ≈ 0.00471882 moles
Since the ratio of NaOH to HCl is 1:1, the moles of NaOH are also 0.00471882 moles.
The concentration of NaOH can be determined using the formula:
Concentration of NaOH (mol/dm³) = Moles / Volume (dm³)
Concentration of NaOH ≈ 0.00471882 moles / 0.025 dm³ = 0.1887528 moles per dm³
The actual value of the concentration of NaOH and HCl is given as 0.2 moles per dm³.
Furthermore, we obtained data from another group for comparison:
Average Titre (cm³) for the other group = (25.2 + 25.3 + 25.3) cm³ / 3 = 25.26 cm³
Concentration of NaOH for the other group can be calculated as follows:
Moles of NaOH = Average Titre (dm³) × Concentration of HCl
Moles of NaOH = 25.26 cm³ / 1000 (to convert to dm³) × 0.191976 mol/dm³ ≈ 0.004994 moles
Concentration of NaOH for the other group ≈ 0.004994 moles / 0.025 dm³ = 0.19976 moles per dm³
For the calibration of the pH meter, we tested it using seven different substances and measured the pH accuracy in two separate tests, then calculated an average pH value for each substance.
| Substance | Test One | Test Two | Average |
|---|---|---|---|
| Water | 6.86 | 6.94 | 6.9 |
| Sodium Hydroxide (1 mole) | 12.91 | 12.93 | 12.92 |
| Sodium Hydroxide (0.1) | 12.48 | 12.46 | 12.47 |
| Ethanol Acid | 2.68 | 2.70 | 2.69 |
| Ammonium hydroxide | 8.18 | 8.28 | 8.23 |
| Nitric acid | 0.53 | 0.48 | 0.555 |
| Hydrochloric Acid | 1.30 | 1.29 | 1.295 |
We proceeded to conduct the experiment with the pH meters, testing the hydrochloric acid (HCl) after every 1 cm³ addition, repeating the test twice, and recording the results. We then calculated the average pH reading for each volume increment.
| Volume of HCl (cm³) | PH Test 1 | PH Test 2 | Average PH |
|---|---|---|---|
| 1 | 12.84 | 12.84 | 12.84 |
| 2 | 12.77 | 12.78 | 12.775 |
| 3 | 12.7 | 12.72 | 12.71 |
| 4 | 12.6 | 12.62 | 12.61 |
| 5 | 12.65 | 12.64 | 12.645 |
| 6 | 12.58 | 12.53 | 12.585 |
| 7 | 12.52 | 12.57 | 12.525 |
| 8 | 12.48 | 12.47 | 12.475 |
| 9 | 12.46 | 12.44 | 12.465 |
| 10 | 12.33 | 12.34 | 12.336 |
| 11 | 12.31 | 12.28 | 12.295 |
| 12 | 12.27 | 12.27 | 12.27 |
| 13 | 12.22 | 12.22 | 12.22 |
| 14 | 12.19 | 12.17 | 12.185 |
| 15 | 12.12 | 12.11 | 12.115 |
| 16 | 11.99 | 12.01 | 12.00 |
| 17 | 11.92 | 11.93 | 11.925 |
| 18 | 11.77 | 11.75 | 11.76 |
| 19 | 11.52 | 11.51 | 11.515 |
| 20 | 11.05 | 11.04 | 11.045 |
| 21 | 10.33 | 10.31 | 10.32 |
| 22 | 9.76 | 9.72 | 9.75 |
| 23 | 9.05 | 9.05 | 9.05 |
| 24 | 8.93 | 8.84 | 8.885 |
| 25 | 3.9 | 3.55 | 3.725 |
| 26 | 2.8 | 2.77 | 2.785 |
| 27 | 2.51 | 2.51 | 2.51 |
| 28 | 2.36 | 2.36 | 2.36 |
| 29 | 2.25 | 2.44 | 2.245 |
| 30 | 2.2 | 2.19 | 2.195 |
Moles of HCl can be calculated as follows:
Moles of HCl = Volume (dm³) × Concentration
Moles of HCl = 0.024 dm³ × 0.186 mol/dm³ ≈ 0.004464 moles
Since the reaction between HCl and NaOH is 1:1, the moles of NaOH are also 0.004464 moles.
The concentration of NaOH can be determined using the formula:
Concentration of NaOH (mol/dm³) = Moles / Volume (dm³)
Concentration of NaOH ≈ 0.004464 moles / 0.025 dm³ = 0.17856 moles per dm³
After analyzing our results, it is evident that both the titration method using an indicator and the pH meter method provided reasonably accurate results, although they exhibited some variations. To some extent, the pH meter proved to be accurate by yielding pH values close to the expected pH of the solutions. However, it displayed fluctuations during measurements, leading to occasional inaccuracies, such as a higher pH reading after the addition of acid. On the other hand, the titration method with an indicator resulted in a more accurate end concentration, approximately 0.1887528 moles per dm³, compared to the pH meter method, which yielded a concentration of 0.17856 moles per dm³.
Comparing our results to the data obtained by another group, we can observe that both methods produced values (0.191976 and 0.19) closer to the actual concentration of 0.2 moles per dm³. Therefore, their results can be considered more accurate than ours. Several factors may have contributed to this discrepancy:
All these factors could have collectively contributed to their results being closer to the actual value of 0.2 moles per dm³.
Our result of 0.1887528 moles per dm³ indicates that we were only 0.0112472 moles per dm³ away from the actual concentration. This suggests that our titration technique was fairly accurate, with only a few potential errors. Some measures we took to ensure accuracy included:
However, it's worth noting that measurement by the naked eye can only provide a certain level of accuracy, and it cannot surpass the precision of the equipment itself. The use of electronic devices could have allowed us to obtain exact values for our measurements, minimizing potential human errors associated with visual readings.
Another point to consider is that this experiment marked our second time using the equipment and the first for the burette. Additional training and experience with the equipment might have led to more accurate results.
Our manual stirring of the flask introduced some inconsistencies in the experiment. Using a magnetic stirrer would have expedited the process and ensured better homogenization of the solutions, which could have affected the pH readings and the overall titration accuracy.
Our graphing of the results, while reasonably accurate, was challenging to interpret from a distance. Utilizing computer-generated graphs could have improved the readability of our results and made them more accessible to others. Drawing the graphs by hand also presented challenges, which might have influenced our final presentation.
Finally, achieving the precise neutralization point was challenging and required a specific amount of solution to reach. Despite our caution, there were instances where we added more solution than necessary, leading to inaccuracies. The use of an electronic pipette filler or a more precise method for volume measurement would have minimized these human errors and provided more consistent results.
Sodium Hydroxide & Hydrochloric Acid Titration: Comparing Methods. (2024, Jan 04). Retrieved from https://studymoose.com/document/sodium-hydroxide-hydrochloric-acid-titration-comparing-methods
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