“Comparative pH titrations of strong and weak acids” Experiment #6 – The aim of the experiment is to titrate a strong acid and a weak acid with a primary standard solution NaOH and finding its pH. The titrant NaOH which is 1M is filled in the burette. Below the titrant, a 250-ml beaker is positioned is such a way that while titrating the NaOH is poured down the beaker which is filled with a solution of 75-ml of DI water and 25-ml of HCL. In order to begin titration, the electrode is put into the beaker such that it doesn’t hit the spinning stir bar. A magnetic stir bar is kept underneath the 250-ml beaker so that the mixture gets dissolved faster. The pH is recorded on the instrument named pH meter which the records the pH …show more content…
Before equivalence point i.e., between 0% to 90% of the reaction, the pH for HCL increases faster than acetic acid to reach the equivalence point. At this phase the solution is acidic. At the equivalence i.e., between 90% and 100% of the reaction, the moles of acid are stoichiometrically equal to the moles of base. The equivalence point for both HCL and acetic acid is 7.67 and 9.19 respectively. The volume at which the equivalence point occurs is 23.25 ml and 24.50 ml respectively for HCL and acetic acid. The volume is different because of the different pH. The phase here changes from acidic to basic and the jumps are sudden for both the acids. Beyond the equivalence point i.e., between 110% and 200% of the reaction, the reaction is almost complete, the base has already reacted with the acid at this point of the reaction and both HCL and acetic acid has nearly a similar pH after the equivalence point because at this phase the pH is controlled by the base since it is in a basic medium. The pH at 0% of the reaction for HCL and acetic acid is 0.70 and 2.72 respectively. At 50% i.e., between the 0% and the equivalence point, the pH for HCL is 3.84
Hydrochloric acid is the clear colourless solutions of hydrogen chloride (HCl) in water, hydrochloric acid is also a highly corrosive substance and a strong mineral acid meaning they are formed from inorganic compounds, hydrochloric acid is a monoprotic acid meaning that it can only ionize one H+ ion. As a result hydrochloric acid can be used in a wide range of industrial practices such as removing rust from steel, ore processing, the production of corn syrup and making of PVC plastics. Hydrochloric acid is made using a very straight forward method which involves dissolving hydrogen chloride (HCl) in water, releasing the H+ cation and Cl- anion. In this aqueous form the H+ ion joins water to form a hydronium ion (H3O+)
A substance called an indicator is added to show the end of the titration. d. Clamp the buret on one side of the buret clamp. Place a white piece of paper labeled "Unknown Acid" under this buret. Drain any remaining pre-rinsed acid solution into a beaker labeled "waste solution". e. Fill this buret with your Unknown acid solution to the zero mark or slightly below it (Not above the zero mark).
In acid-base titration solution without a known molarity is placed in an Erlenmeyer after it’s volume is measured. An indicator is added to the solution most of the time it is phenolphthalein. The solution with a known concentration is placed in burette with a tap in the end. By opening the tap slightly solution in the burette is poured in to the solution in Erlenmeyer drop by drop. After a while the solution in Erlenmeyer forms a color change. This is the turning point for the solution. At the turning point by the volume consumed in burette the molarity of the other solution can be
As shown by Graph 1, the equivalence pH appears to be around 7, and this makes sense since the reaction between the strong acid HCl and the strong base NaOH is simply a direct neutralization reaction since both will strongly dissociate and react with each other. This is further shown by the fact that almost twice as much NaOH had to be used to get to the equivalence point to neutralize the strong HCl than when compared to the amount of NaOH that had to be used to reach the equivalence point with acetic
The whole purpose of this experiment is to determine wether or not the amount of the zinc and or hydrochloric acid effects the out coming percent of the solution after under going chemical reaction.
In the titration experiment, the endpoint was recorded in the experimental data to be at 21.30 mL of NaOH and at a pH of 10.44. However, when all of the data from the table was graphed, the observed endpoint was too high up and on a part of the upper concave down portion of the graph. To ensure that the proper equivalence point was used, a new point had to be extrapolated that was roughly the point at which the graph went from concave up to concave down. This point was at 21.28 mL of NaOH added and pH of 9.20. Dividing both of these points by two, the half equivalence point was found to be at a pH of 5.30 and 10.64 mL of NaOH added. The pH is equal to the pKa here, so the pKa was found to be 5.30. Using data from the equivalence point extrapolated from the graph, the molar mass of the unknown was calculated to by 148 grams per mole. Lastly, because there was only one region of
From looking at the results I can conclude that when the pH was 3 and
more it is in contact with the acid so it will react at a different
and thiosulphate will not be mixed up. The amount of HCl will be 5 cm3
Exactly 10.00 mL of vinegar (Stop and Shop Distilled White Vinegar, All Natural, 5% Acidity) was added to a 125 mL Erlenmeyer flask, using a 50 mL burette. Three drops of phenolphthalein indicator were mixed into the flask. The initial volume of the burette containing the NaOH was recorded to the nearest hundredth milliliter. The flask containing the vinegar was placed under the 50 mL burette and slowly, the NaOH was drained into the flask, with constant mixing. The burette continued to be drained until the acetic acid solution changed from a clear color, to a persistent light pink, indicating the end point of the titration. The final volume of the burette was then recorded to the nearest hundredth milliliter and the difference between the initial and final volumes was calculated; 31.35 mL or 0.03135 L. This volume in liters was multiplied by the concentration of the NaOH solution, 0.2838 M to determine the number of moles of acetic acid solution. This value was then divided by the volume of acetic acid solution originally added to the flask, 10.00 mL or 0.01 L, to obtain the molarity of the acetic acid solution; 0.8903 M. This process was replicated exactly a second and third time. The second titration required 31.39 mL of NaOH to be added, resulting in an acetic acid molarity of 0.8915 M. The third titration required 31.44 mL of
0.1003 M of HCl was diluted tenfold to 0.01003 M. 25cm3 of HCl was pipetted to another 250.0cm3 volumetric flask, then deionized water was added to the volumetric flask until the graduation mark.
...3 occured by an acid attack after a complete removal of Ca(OH)2 at low pH as shown in Equation 2.10.
Initially, before any NaOH is added, the pH of H2C2O4 .2H2O is low because it contains mainly H3O+. The starting pH will, however, be higher for a weak acid, like H2C2O4 .2H2O, than for a strong acid. As NaOH is added, H3O+ is slowly used by OH- because of dissociation of NaOH. The analyte remains acidic but the pH starts to increase as more NaOH is added.
Initially, the concentration of the reagents decreases. As the concentration decreases, the rate of the forward reaction slows down. Meanwhile, the rate of the reverse reaction continually increases. Eventually, the rate of the forward reaction will equal the rate of the reverse reaction, and even though the reaction is still occurring, the change in concentration is insignificant. The equilibrium expression for the general reaction (1) is shown below:Keq is a value that is only dependent on temperature. If A, B, C, and D are mixed at unknown concentrations, the reagents will continuously react until equation (2) is true.
For this experiment we used titration to standardize the exact concentration of NaOH. Titration is the process of carefully adding one solution from a buret to another substance in a flask until all of the substance in the flask has reacted. Standardizing is the process of determining a solutions concentration. When a solution has been standardized it is referred to as a standard solution. To know when a solution is at its end point an indicator is added to acidic solution. An indicator is an organic dye that is added to an acidic solution. The indicator is one color is in the acidic solution and another color in the basic solutions. An end point occurs when the organic dye changes colors to indicate that the reaction is over (Lab Guide pg. 141).