Landolt Clock Reaction Lab Report

Investigations into the mechanics of chemical kinetics can reveal invaluable information relating to the rates of reaction. There are numerable applications of reaction rates, knowledge in this area is pivotal for industrial, commercial and research sectors. Thus, allowing them the ability to manipulate a variety of factors of chemical reactions with the use of reaction rates. In the scope of the kinetics of clock reactions, there is a range of information that can be obtained about reaction rates (Shakhashiri, 1992).

“The iodine clock reaction was discovered by the Swiss chemist Hans Heinrich Landolt in 1886.” (Scilearn, 2017). There are many variations of the iodine clock reaction but they all involve two colourless solutions being mixed and after a specific period of time, the mixed solutions turn dark blue. The iodate variation of the clock reaction involves the reaction between iodate ions. These being, IO3-, hydrogen sulfite ions, HSO3-, and hydronium ions, H+. To ensure the blue colour change, starch is needed to be added to the solution, reacting with I2. The delay in colour change occurs because as soon as I2 is formed, it immediately reacts with any HSO3- still present and is converted into colourless I-. After all the HSO3- has been consumed, the concentration of the I2 increases and begins reacting with the starch. This is why the dark blue colour is suddenly present. The reaction of the starch reacting to turn the solution a deep blue colour. The reaction occurs in two steps. The first step generates iodide ion (I in Equation 2, which occurs slowly and is the rate determining step. “A rate determining reaction is the slowest step in a reaction mechanism and is assumed to be equal to the overall reaction rate because the reaction cannot go faster than the slowest step.”(Webcache.googleusercontent.com, 2017 ).

Clock reactions are among the most visually entertaining demonstrations of chemistry. Simply put, a clock reaction is one where two substances are mixed and there is a time delay where there is no visually discernable change in the system. During this delay, one of the chemical species, the clock chemical, has a very low concentration. The end of this induction period is marked by a rapid increase in concentration of the clock chemical (S.J. Preece et al., 1999). This rapid increase in concentration is what triggers effects such as a sudden colour change.

The iodination of O-acetylsalicylamide under the same conditions of this experiment would be slower than Salicylamide, because O-acetylsalicylamide has an amine group, a carboxyl group, and no hydroxyl group, which means that the sodium hypochlorite solution would not be able to deprotonate the Hydroxyl group, thus not allowing a solid precipitate to form, and Hydrochloric acid would not be able to add any Hydrogen atoms due to the Octet rule. The only way that Hydrogen could be added is by breaking double bonds and stabilization (Resonance). And steric hindrance would be another key facet slowing the addition of an Iodine ion to the

Three factors that can affect the rate of reaction are temperature, pH and the salinity.

In next reaction which took place at a temperature 10˚ higher than that had a rate of 39.5 mL/minute. Next was the reaction which took place at 10˚ below room temperature which resulted in a rate of 26.8 mL/minute. In the reaction that 20 mL of distilled water was added to the bleach solution and the temperature was kept constant, the reaction rate dropped to 16.2 mL/minute. Finally the slowest reaction occurred when 60 mL of distilled water was added to the bleach causing a rate of 10.8 mL/minute.

The rate of a chemical reaction often depends on reactant concentrations, temperature, and if there’s presence of a catalyst. The rate of reaction for this experiment can be determined by analyzing the amount of iodine (I2) formed. Two chemical reactions are useful to determining

The aim of the experiment will be to investigate how varying water temperatures influence the time of a chemical reaction, in this case being, a combination of Sodium Thiosulfate and Hydrochloric Acid.

In this Crystal Violet Kinetics Lab, many procedures were performed order to experimentally determine the rate constant(k), for the rate of the reaction of the crystal violet solution and sodium hydroxide. To start, a Beers Law calibration experiment was conduction. Also, two kinetic experiments were ran using both 5ml and 10ml of sedum hydroxide in order to determine the reaction order(k^1), of the sodium hydroxide.

In the future, this experiment could improved or expanded. One area of improvement could be to do this experiment several times and then averaged together or compare the results. This leaves less room for error and could make the data more accurate and precise.

When one looks back at the procedure of the experiment, trials 1, 2, 3, 4, 5, and 6 all use 6 M HCl for their reactions, and if this were to be substituted for 6 M HNO3, then these reactions would have still taken place, except in trial 2 because copper is less reactive than hydrogen in all cases, but their final observations may yield different colored solids and solutions (Beran,

As an enzyme-catalyzed reaction may be the main reason for a reaction to occur faster, many factors can

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

Two catalyst reactants are used in the experiment, thiosulfate and starch, to dictate the time of reactions.

The key aim of this experiment was to determine the rate equation for the acid-catalysed iodination of acetone and to hence consider the insinuations of the mechanism of the rate equation obtained.