0.4 M Hydrochloric Acid

From the results that were collect throughout the experimental investigation has proved the hypothesis to only be partially right. Multiple tests were made when conducting the experiment, two clear solutions were combined at various temperatures and concentrations. The hypothesis states that by adjusting the concentration of the reactants will cause the reaction to either speed up at a higher concentration or slow down at a lower concentration. In the reaction temperature should have a similar effect on the experiment, in that increasing the temperature will cause an increase of particle movement and cause more collision, thus increasing the reaction rate. Therefore decreasing the temperature will decrease the rate of the reaction. From the results given in Tables 2 and 3 it shown that every time the concentration is halved the time is increased. When the concentrations of both KIO3 and NaHSO3 are decreased the time has increased, some concentrations having a higher increase than others. In each concentration decrease the time is at least doubled from the previous concentration time, which is therefore increasing the rate of the reaction.

The results for the rate of reaction, as seen in Table 2, suggest that as the temperature increases, the rate of reaction increases at a progressively greater rate. When the temperature increases 21℃ from 4℃ to 25℃, the rate of reaction increases from 1.846 s-1 to 6.1100 s-1 while a similar increase in temperature from 25℃ to 65℃ has a much more significant difference in the rate of reaction: 47.85s-1 at 65℃.

As the substrate in increased the initial reaction rate will increase until it is fully submerged and then will not increase any more. The relationship is linear. As the substrate concentration increases, the initial reaction rate increases.

To improve the results from the experiment buffer solutions that were not whole pHs could have been used e.g. pH 4.5, 5.5 etc. This would have provided more reliable results as a wider range of results would have been produced. Using pHs with decimals would also help to more accurately determine the optimum pH as the optimum may have been above or below the pH stated in the hypothesis; 8. In this experiment however the optimum is taken at 8 because the graph does not rise again.

The problem that needed to be solved for this experiment was, “How are organisms able to survive and function with proficiency despite metabolic activities that have the disastrous ability to alter pH from a neutral to an acidic or basic environment?” This question needed to be answered by testing different biological materials, as well as tap water, in order to find out what allows modern day organisms to survive such hazardous conditions. A hypothesis that was formulated before experimenting was, “If type of substance is Tap Water, then the changes in pH will be more drastic because water, considered neutral on the pH scale, doesn’t contain a component that will prevent such a change, unlike the content of a buffer solution.” The independent variable for this experiment was type of substance, evidently because this was the part of the experiment that was constantly being changed from trial to trial. The dependent variable for this experiment was pH of the substance because the pH constantly changed depending on the type of substance was being experimented with. The control of this experiment was the tap water and buffer solutions, as the results of the data from the biological materials were constantly compared to the data from both of these solutions.

In this experiment we tested the effects that enzymes and substrate have on chemical reaction rates, which is the rate at which chemical reactions occur.. This experiment tested how different concentrations of enzyme and substrate affected the light absorption measurements on a spectrophotometer. The experiment also tested how temperature affected the light absorption, and in a separate test, the effect of the enzyme inhibitor hydroxylamine was also tested. In the first test conducted, 3 different concentrations of enzyme, and three different concentrations of substrate were measured in a spectrophotometer. For the enzyme and the substrate, the measurements got higher as the concentrations were higher, but the over measurements of the substrate were smaller than those of the enzyme. In the second test conducted, the medium concentration enzyme was tested under the temperatures; 4°C, 23°C, 37°C, and 60°C. The measurements in this test got higher as the temperature got higher, but did the measurements under 4°C were overall significantly higher than the other temperature measurements. Lastly, the last test conducted showed that the measurements of the substance with 0 and 1 drop of hydroxylamine inhibitor went up, but the measurements of the enzyme with 5 drops of hydroxylamine inhibitor stayed rather low and did not change much. In conclusion, these experiments showed that chemical reaction rates are sped up with higher concentrations of enzyme, substrate,

4. If HNO3 was used in place of HCl, I’d expect the rate of change to be about the same as both are strong acids. However, if acetic acid (CH3COOH) were to be used instead of HCl, I’d except the reaction rate to be slower as acetic acid is a weak acid and it will dissociate in the water unlike the stronger

After getting the results of the 5 different concentrations using the formula c=A/εl (Lambert-Beer Law). So a new graph was created using the concentration plotted versus the time (s). The final rates of the reactions for the five different trial were first trial was 0.5 mM/s , second trial 0.3 mM/s , third trial 0.05 mM/s , fourth trial 0.03 mM/s , and for the fifth was 0.005 mM/s. the final result of the curve line was not that accurate because of the trial 4th that was a little off.

In Part 1, ~0.30g of KHP is dissolved in 25.00 mL of water, and titrated with NaOH. The volume of NaOH added to reach the endpoint is recorded and corrected (Table 1). The number of moles of NaOH reacted in trial 1 and trial 2 are 0.001534 mol and 0.001433 mol (Equation 1). The concentration of NaOH in trial 1 and trial 2 are 0.07331 M and 0.06628 M (Equation2). The average concentration of NaOH is 0.06979 M (Equation 3). The average deviation is 0.003515 M (Equation 4). In Part 2, ~0.13g of Na2CO3 is dissolved in 25.00 mL of water, and titrated with HCl. The volume of HCl added to reach the endpoint is recorded and corrected (Table 2). The number of moles of HCl reacted in trial 1 and trial 2 are 0.001225mol and 0.001253mol (Equation 5). The concentration of HCl in trial 1 and trial 2 are 0.03298M and 0. 0.03477M (Equation 6). The average concentration of HCl is

The rate of a chemical reaction often depends on reactant concentrations, temperature, and the presence of a catalyst. Additionally, the rate law is determined mathematically only from experimental data. The reaction investigated in this experiment is:

The reaction rates for each flow rate can be observed as well in Figure 2 and Figure 3 on the slope equation. The reaction rate of H2CO3 with time for both flow rates was not plotted on the same

For this experiment, titrations on a weak acid, acetic acid, and a buffer were performed. Acetic acid was titrated with NaOH in order to observe the half-equivalence point as well as the equivalence point. Then, the buffer and the buffered acetic acid solution prepared faced additional titration with NaOH and HCl to evaluate the differing buffering effects following the addition of a strong acid and strong base. Finally, the buffer’s buffering capacity was calculated. If the experiment were to be repeated, it would be interesting to observe the buffering effects following a titration between a weak base and a buffer instead with greater concentrations. The change in the concentration following the preparation of buffer with a weak base and its conjugate acid would pose for an interesting experiment to observe an increase in the buffering capacity.

The results of the experiment prove the hypothesis of the group. It was hypothesised that the less diluted the acid is, the faster the ant-acid will dissolve. The results show this, as the ant-acid in the 0.5M acid took 9:02 minutes to dissolve; the ant-acid in the 1.0M acid took 2:50 minutes to dissolve and the ant-acid in the 2.0M acid took 2:09 minutes to dissolve. The trend of this experiment is the downward slope of the time when the acid becomes stronger, however, it is not as mathematically correct as the group had expected. The group had hypothesised that the 1.0M acid would take half the amount of time that the 0.5M acid would take, and the 2.0M would take half the amount of time the 1.0M acid would take. This was not the case, as

Chemical kinetics involving reaction rates and mechanisms is an essential part of our daily life in the modern world. It helps us understand whether particular reactions are favorable and how to save time or prolong time during each reaction. Experiment demonstrated the how concentration, temperature and presence of a catalyst can change the rate of a reaction. 5 runs of dilution and reaction were made to show the effect of concentration on chemical reactions. A certain run from the previous task was twice duplicated to for a “hot and cold” test for reaction rate. The prior run was again duplicated for a test with

In the reaction between HCl and NaOH, the temperature change was 13.474 °C for the first trial and 13.71 °C for the second