Introduction
The objectives in this laboratory were to explore the titrations of several known acids and bases as well as observing the mole reactions of the chemical reactions that each compound undertook. The four following acids and bases were used to predict when a set amount of reactant would reach equilibrium: Hydrochloric Acid (HCl), H2SO4 (Sulfuric Acid), NaOH (Sodium Hydroxide), and Ba(OH)2 (Barium hydroxide). An unbalanced equation of each reaction was given so the balanced equation was then able to be calculated. These are as follows including the theoretical ratio between the acid and base:
HCl + NaOH → H2O + NaCl 1:1 H2SO4 + 2NaOH → 2H2O + Na2SO4 1:2 H2SO4 + Ba(OH)2 → 2H2O +BaSO4 1:2 2HCl + Ba(OH)2 → 2H2O +BaCl2 1:1 …show more content…
The estimated amount of NaOH required was 16 mL so 14 mL was added quickly, but in the end it took 15.6 mL the first trial and 15.5 mL the second trial to obtain equilibrium. Both trials produced a light pink mixture which indicated that the amount of NaOH added was perfect. The solution of NaOH from the buret and H2SO4 were both discarded into the waste bin and rinsed three times to prepare for the next trial. The same process was used to obtain the 50 mL of NaOH as in the previous experiment, but Ba(OH)2Ba(OH)2 was used instead. HCl was added to the phenolphthalein indicator and deionized water as in the first trial. 4 mL was added quickly and a singular drop was added slowly from there. A reaction took place at exactly 6 mL which was what was predicted. A light pink solution was produced. The second trial went similarly at 6.1 mL. The HCl solution was dumped into the waste bin and rinsed to prepare for the H2SO4 that would follow in the next …show more content…
These objectives were met by observing Hydrochloric Acid (HCl), H2SO4 (Sulfuric Acid), NaOH (Sodium Hydroxide), and Ba(OH)2 (Barium hydroxide) titrate together. By using the exact volume that it took for the acids and bases to react with the help of the phenolphthalein indicator the moles of each were able to be compared and concoct an experimental ratio. The phenolphthalein indicator was used to tell exactly when there was a perfect ratio of acid and base in the solution. This was then compared to the theoretical ratio which allowed a percent of error to be calculated. The experimental ratio for the given reactions in this lab were very similar to the theoretical value. Even in the reaction between Ba(OH)2 and H2SO4 where there was a percent error average of 13.6 this was still a close calculation. In this reaction it was estimated that 2-3 drops of excess H2SO4 were added to the Ba(OH)2 which would have increased the volume ultimately affecting the moles of both compounds. For the remainder of the reactions such as NaOH + H2SO4 the volume was taken one drop before it turned light pink to get the exact amount. This was very precise and only accounted for 2.95%
The objectives in this laboratory were to explore the titrations of several known acids and bases as well as observing the mole reactions of the chemical reactions that each compound undertook. The four following acids and bases were used to predict when a set amount of reactant would reach equilibrium: Hydrochloric Acid (HCl), H2SO4 (Sulfuric Acid), NaOH (Sodium Hydroxide), and Ba(OH)2 (Barium hydroxide). An unbalanced equation of each reaction was given so the balanced equation was then able to be calculated. These are as follows including the theoretical ratio between the acid and base:
HCl + NaOH → H2O + NaCl 1:1 H2SO4 + 2NaOH → 2H2O + Na2SO4 1:2 H2SO4 + Ba(OH)2 → 2H2O +BaSO4 1:2 2HCl + Ba(OH)2 → 2H2O +BaCl2 1:1 …show more content…
The estimated amount of NaOH required was 16 mL so 14 mL was added quickly, but in the end it took 15.6 mL the first trial and 15.5 mL the second trial to obtain equilibrium. Both trials produced a light pink mixture which indicated that the amount of NaOH added was perfect. The solution of NaOH from the buret and H2SO4 were both discarded into the waste bin and rinsed three times to prepare for the next trial. The same process was used to obtain the 50 mL of NaOH as in the previous experiment, but Ba(OH)2Ba(OH)2 was used instead. HCl was added to the phenolphthalein indicator and deionized water as in the first trial. 4 mL was added quickly and a singular drop was added slowly from there. A reaction took place at exactly 6 mL which was what was predicted. A light pink solution was produced. The second trial went similarly at 6.1 mL. The HCl solution was dumped into the waste bin and rinsed to prepare for the H2SO4 that would follow in the next …show more content…
These objectives were met by observing Hydrochloric Acid (HCl), H2SO4 (Sulfuric Acid), NaOH (Sodium Hydroxide), and Ba(OH)2 (Barium hydroxide) titrate together. By using the exact volume that it took for the acids and bases to react with the help of the phenolphthalein indicator the moles of each were able to be compared and concoct an experimental ratio. The phenolphthalein indicator was used to tell exactly when there was a perfect ratio of acid and base in the solution. This was then compared to the theoretical ratio which allowed a percent of error to be calculated. The experimental ratio for the given reactions in this lab were very similar to the theoretical value. Even in the reaction between Ba(OH)2 and H2SO4 where there was a percent error average of 13.6 this was still a close calculation. In this reaction it was estimated that 2-3 drops of excess H2SO4 were added to the Ba(OH)2 which would have increased the volume ultimately affecting the moles of both compounds. For the remainder of the reactions such as NaOH + H2SO4 the volume was taken one drop before it turned light pink to get the exact amount. This was very precise and only accounted for 2.95%