The purpose of this lab was to understand the difference between a strong acid and a weak acid and how that affects its titration curve. We were also asked to estimate the ionization constant of a weak base with the data we collected. The first step of the experiment was to put 50 drops of acetic acid into a small beaker. We then recorded the pH values after 0,10,20,30,40,45,50,55,60, 70,80,90, and 100 drops of NaOH was put in the beaker. The pH was recorded by using pH paper and dipping it into the solution at each increment and then comparing it to the color chart. This experiment showed that our equivalence point had a pH of 9 and the half equivalence point was at 5.5. In order to find the Ka value, I used the Henderson Hasselbach equation. …show more content…
To find Ka you would raise -5.5 to 10, this gave us a concentration of 3.16*10^-6. The percent error came out to be 82.4%. This percent error seems very high but because we are dealing with very sensitive measurements, I believe it is okay. Various errors could exist in the lab. One error is the possibility of added to many drops of NaOH. If the procedure called for 20 drops but 25 drops was added instead, the half equivalence point would be higher and the whole graph would shift up which ultimately also changes the equivalence point. Another error could be the fact that we use pH paper to qualitatively find the pH. If the color seems to be in between and you estimate up, this could raise up the equivalence and half-equivalence point. To increase the accuracy of the lab, smaller increments on the pH chart should be used. This will give us a better titration curve since some values were the same in the original experiment. The titration curve of a weak acid and strong acid differ because weak acids do not dissociate completely. For example, when HOCl dissociates, it makes OCl. When the conjugate base combines with water, it produces OH- which in turn raises the equivalence
A precipitation reaction can occur when two ionic compounds react and produce an insoluble solid. A precipitate is the result of this reaction. This experiment demonstrates how different compounds, react with each other; specifically relating to the solubility of the compounds involved. The independent variable, will be the changing of the various chemical solutions that were mixed in order to produce different results. Conversely the dependent variable will be the result of the independent variable, these include the precipitates formed, and the changes that can be observed after the experiment has been conducted. The controlled variable will be the measurement of ten droplets per test tube.
1. 3mL of water was added to 2r.g. of KCNS in the test tube and it
One possible source of experimental error could be not having a solid measurement of magnesium hydroxide nor citric acid. This is because we were told to measure out between 5.6g-5.8g for magnesium hydroxide and 14g-21g for citric acid. If accuracy measures how closely a measured value is to the accepted value and or true value, then accuracy may not have been an aspect that was achieved in this lab. Therefore, not having a solid precise measurement and accurate measurement was another source of experimental error.
Because of the misunderstood procedural step from the original procedure, (3. Initially try a 20mL sample of your drink for the preliminary titration. You may need to adjust the amount of sample used with a dilution, if the amount of titrant,NaOH, is not between 20-40mL), and other factors, the percentage error of this experiment was extremely high. This is because more NaOH was added then needed. The volume of NaOH added plays a major role in the concentration of apple juice, as it will directly affect the concentration number.
= = pH 1 2 3 Average Rate of Reaction (cm3/s): 0 - 0. 3 0 0 0 0 0.000 5 0 0 0 0
... Another one of These small errors are the accuracy of the balances - the readings given may not have been entirely accurate. However, this is slight. inaccuracy (0.004g) pales in the face of the other experimental.
Added 30 mL of distilled water and two drops of phenolphthalein indicator to the acid sample. Record the exact mass of the vinegar. Placed sodium hydroxide solution in a clean dry beaker and labelled it. Obtained a burette, burette clamp, meniscus reader, and a pH meter. Emptied distilled water from burette and rinsed with a few milliliters of base.
The Ka of an acid can be calculated in two ways, including (1) the measurement of the pH of a solution containing a know concentration of a weak acid, and (2) measurement of the pH at the
2. While determining the average volume of a 10 mL pipet, it was easy to conclude that there are often errors when measuring due to our 1.28% relative error. There could have been many causes of this error but the most likely would be human error. Error leads to inaccurate measurements. Measurements are not always accurate but when there are more graduations and you can estimate with an extra significant figure, they are more accurate. To minimize error associated in measurements you can check to make sure the scale or balance is calibrated properly and also conduct several trials.
It should start with hypotonic and moves to hypertonic as the concentration of the salt water increases. There are few reasons why the result of the experiment is not reliable. First of all, the standard deviation of the % mass change is so big, which means the experiment was not accurate. According to the Table 3, the % mass change of 3% concentration increases to 18% and decreases to -11%. As each trial shows the big difference each other, the experiment is not reliable.
Draw a graph, using a line of best fit. Table of the results: Total grams of KClO3 g Total volume of distilled water cm3 Temperature at which solid come out of solution oc Solubility of H2O g/100g 2.00 4.00 96 50 2.00 8.00 70 25 2.00 12.00 52 16 2.00 16.00 33 12.5 2.00 20.00 25 10 2.00 24.00 18 8.3 If there was more time I could have repeated this work to make sure this information was reliable but I could not. Analysis Look at the table and the graph above. These results were obtained by the experiment described.
All the above calculations were calculated by a 1 in 10 dilution and then making the sample up to 10ul with sterile H20.0
Carrying out a titration: A conical flask was swilled out with water and a pipette and pipette filler were used to withdraw 25.0cm3 of the sodium carbonate solution from the volumetric flask and transfer it to the conical flask. A burette was first swilled with sulphuric (VI) acid using a clean, dry beaker and a funnel and then filled to below the zero mark. A little of the solution was then run out of the burette into the beaker and the funnel removed. A white tile was then placed underneath the conical flask and a few drops of the indicator methyl orange was added to the sodium carbonate solution.
After creating the second dilution, which gave the phosphate a .00015 molarity, it was recognized by the instructor that the Spec-20 provided had not been properly calibrated. Due to the improper calibration of the Spec-20 and a restraint for time the dilutions had to stop there. Because the Spec-20 had not been properly calibrated, it caused the data for week two to be fairly skewed. In order to obtain proper data, the instructor supplied reference data to be used for the majority of week two. The main issue faced during the second week of lab was having trouble making the dilutions properly.
I am going to be given a sample of sulphuric acid, which is thought to have a concentration between 0.05 and 0.15-mol dm³. The purpose of this experiment is to find the accurate concentration of the sulphuric acid. I will do this by carrying out a titration between sulphuric acid and sodium carbonate solution. Therefore this is an acid-alkali titration (which is the determination of concentration by adding measured amounts of standard reagents to a known volume until the end point is reached). * Sulphuric acid is considered a strong acid, as it is completely in the form of ions in dilute solution.