1. In step 1 HCl, a strong acid, was added one drop at a time to DI water. The addition of HCl to the DI water resulted in the pH going down. The first drop of HCl that was added caused the pH to go from 5.24 down to 4.20. The second drop caused the pH to go to 3.92, followed by the third drop which caused the pH to go to 3.79. The overall addition of HCl caused a 1.45 drop in pH. 2. Step 2 was the same as step 1 but NaCl, a salt of a strong acid-strong base, was used in replace of DI water. The addition of NaCl slightly buffered the pH but not so much so that the pH did not drop at all. This is because NaCl is a neutral salt, it is composed of an alkaline metal and a halide which neither effect H+. Which explains why the pH still went down …show more content…
In step 4 NaOH, a strong base, was added drop by drop to NaCl, a salt of a strong acid-strong base. The addition of NaOH to NaCl resulted in the pH dramatically going up which is the opposite of what happened in step 1. The initial pH of NaCl was 4.71 and when just one drop of NaOH was added the pH went up to 5.14. After all three drops of NaOH were added to NaCl the pH was tested to be 8.21, the highest pH so far. So the NaCl most certainty did not buffer the pH changes. Also step 4 had the largest overall magnitude of change in pH, of 3.5, over double the change in step 1. 5. Step 5 was similar to step 1 but instead of adding HCl to DI water, HCl was added to an acetic acid–acetate ion buffer. By using an acetic acid-acetate ion buffer in replace of DI water when HCl was added caused the pH to barely even change. The initial pH of the acetic acid-acetate ion buffer was 4.75 and when the first drop of HCl was added the pH went down like in step 1 but only .02. After all three drops of HCl were added, the pH was tested to be the exact same as the initial pH. This is not like what happened in step 1 where the pH went down 1.45 by the third drop of HCl. CH3COO-(aq) + H+(aq) ↔ …show more content…
CH3COOH(aq) + OH-(aq) ↔ CH3COO–(aq) + H2O(l) 7. Step 7 was similar to step 1 but instead of adding HCl to DI water, in step 7 HCl was added to an ammonia-ammonium ion buffer. When HCl was added to the ammonia-ammonium ion buffer the pH remained the same which is not like step 1 where the pH went down. The pH of the ammonia-ammonium ion buffer was originally 9.30 and though it fluacted .01 drop 1 and 2, by drop 3 it was once again 9.30. This is not like what happened in step 1 where the pH consistently went down and by drop 3 of HCl the pH had gone down 1.45. NH3(aq) + H+(aq) ↔ NH4+(aq) 8. In step 8, NaOH was added a drop at a time to an ammonia-ammonium ion buffer. By adding NaOH to the ammonia-ammonium ion buffer caused the pH to nearly stay the same, unlike what happened in step 1 where the pH went down after HCl was added. The pH of the ammonia-ammonium ion buffer was initially 9.40 and by the third drop the pH was 9.41, which is only a .01 increase in pH. This is unlike what occurred in step 1 where by the third drop of HCl the pH had increased
Marwah Alabbad Post lab 10/21/15 Question 1: 1. Experiment 1: Number of trails NaOH concentration (M) Volume of HCl solution (mL) Initial volume of NaOH(mL) final volume of NaOH(mL) The volume of NaOH to titrate HCl (mL) Concentration of HCl (M) 1st 0.1023 25.0 10.05 36.12 26.07 0.085 2nd 0.1023 25.0 5.74 31.40 25.66 0.105 3rd 0.1023 25.0 9.84 35.52 25.68 0.105 First trail calculation: 0.02607L× (0.1023mole NaOH/1L)×(1 mol of HCL/1 mol of NaOH)×(1/0.025)= 0.085M of HCl
1mL of Acetic acid was then added to Unknown D and the solution was stirred. Next, 15mL of sodium
The unknown compound was first reacted with an acid. To begin, 0.50 grams of KCl was mixed with 5 mL of water. Then, 1 mL of 6 M H2SO4 was added to the solution. Secondly, the unknown compound was reacted with a base. Exactly 0.50 grams of KCl was mixed with 5 mL of water, and 1 mL of 1 M NaOH was added to the solution next.
The KHP and the acid samples must be dried, because there would still be extra water which would skew the molarity. The calculated molarity of the NaOH would be lower because there would be extra volume in the solution, but still the same amount of moles of NaOH, so the molarity would be less, and thus, will require more titrant in order to
3. Upon adding 20 drops of NaOH, a white precipitate was formed signifying acidic impurity. In the second NaOH mixture, about 20 drops were administered and no precipitate formed indicating that the ample is more pure than before. Data: Weight of flask = 75.10 grams Weight of the flask with solids =
Both of these reactions involves an acid reacting with a sodium bicarbonate, NaHCO3. The pKa value of this compound is around 6.4. looking at the first reaction and comparing the pKa values of the carboxylic acid and the sodium bicarbonate, (4-5, and 6.4 respectively), it can be seen that the sodium bicarbonate will act as a base, due to a higher pKa value, and the reaction will proceed to create a conjugate base and acid. On the other hand, looking at the second equation, it is evident that the phenol when reacted with sodium bicarbonate will act as a base. This is due to sodium bicarbonate having a lower pKa value.
How did the final NaOH molarities compare for each trial? Were they similar? The results of the trials were reproducible in the sense that they all showed consistency in molarities. 2) Why were you able to pour the products down the sink for this lab?
Each buffer was measured in a 100 mL graduated cylinder and contained in a 40 mL beaker. Once the reading of the buffer was stabilized, the program entered into reading 1. The probe was cleaned with distilled water and dried before being placed into the second buffer for reading 2. Once the second calibration was completed the pH probe was cleaned again. Next the probe was placed into the unknown solution.
Firstly, because the NaHCO3 compound was not stored in a sealed container, therefore dust particles could have changed the results, and making the product impure. Also, there are uncertainties associated with the instruments used in this experiment. This, if the products were measured slightly more than should be, this could have affected the concentrations of the solutions, and therefore causing a larger
The equation of the reaction between sodium hydroxide and ethanoic acid is as follows: CH3COOH + NaOH → CH3COONa + H2O We can measure the end point of titration process and we can also measure the amount of reactants. The concentration of ethanoic acid in the vinegar can be determined through stoichiometric calculations, Using the values obtained from the titration, and also the chemical equation as a reference. Phenolphthalein indicator is used in this acid-base titration Equipment and materials:
Its pH is greater than 7 and turns red litmus paper into blue. Acid- base neutralization is done by adding an acid to a base or a base to an acid until the substance has equal hydrogen and hydroxide ions. This is used to determine unknown concentration of a
Introduction Buffer is a solution that resists a change in pH when bases or acid are added. Solutions that are acidic contain high concentrations of hydrogen ions (H+) and have pH values less than seven. Buffer usually consist of a weak acid, and its conjugate base or a weak base and its conjugate acid. The function of buffer is to resist the changes in hydrogen ion concentration as a result of internal and environmental factor. This buffer experiment is important so that we relies the important of buffer in our life.
The drop rate was adjusted to 1 to 2 drops/second. 10.0 mL of the NaOH solution was allowed to drip away into a waste beaker. The exact volume of the sodium hydroxide solution used was determined.
The chemical equation for this experiment is hydrochloric acid + sodium thiosulphate + deionised water (ranging from 25ml to 0ml in 5ml intervals) sodium chloride + deionised water (ranging from 25ml to 0ml in 5ml intervals) + sulphur dioxide + sulphur. As a scientific equation, this would be written out as, NA2S2O3 + 2HCL + H2O (ranging from 25ml to 0ml in
That caused a new initial reading of NaOH on the burette (see Table1 & 2). The drops were caused because the burette was not tightened enough at the bottom to avoid it from being hard to release the basic solution for titrating the acid. The volume of the acid used for each titration was 25ml. The volume of the solution was then calculated by subtracting the initial volume from the final volume. We then calculated the average volume at each temperature.