The activation energy of a chemical reaction is the amount of energy required to break apart the bonds of the reactants in order to bring them into a transition state, where they form new bonds to form products releasing energy. A catalyst, whether it is biological or not, provides an alternative route for the reactants to take in order to lower the activation energy for the reaction. For a chemical reaction to occur, the particles must have physical contact, the correct orientation, and sufficient kinetic energy to initiate the reaction. Recent research found that hydrogen peroxide is present in the Colorado (USA) groundwater [Xiu, 2015]. Since the consumption of hydrogen peroxide can be very dangerous [Standard Operating Procedure – Hydrogen Peroxide], a method must be found to remove it. If the H2O2 is not removed from the Colorado groundwater, it could pose harmful effects upon ingestion (see Safety Analysis). Hydrogen peroxide naturally decomposes into water and oxygen gas — however this process is very slow. Catalase and manganese dioxide are both catalysts which speed up this process. Manganese dioxide is a heterogeneous catalyst as it is in the solid phase and hydrogen peroxide is aqueous. More importantly, it is a transition metal, which increases the rate of decomposition of hydrogen peroxide by decreasing the activation energy for the reaction. Catalase is an enzyme for the same decomposition reaction and works in the same way as manganese dioxide; however it has different properties. Since it is a biological catalyst temperature changes can affect the catalyst’s effect upon activation energy. This lead to the research question: How does the use of manganese dioxide compare to catalase in its effect on the activation energy for the decomposition of hydrogen peroxide?
To determine the activation energy for each catalyst, the rate constant, k, of the decomposition of hydrogen peroxide is be determined at different temperatures [Bylikin et al, 2014]. To increase the reliability, each independent variable is tested independently and all other variables that could affect the results, i.e: pressure, amount of hydrogen peroxide and catalyst, are controlled throughout the trials.
The dependent value is the rate of reaction for the reaction, which is affected by temperature and the type of catalyst. Therefore multiple temperatures are tested for each type of catalyst. The activation energy is then be calculated mathematically. To measure the rate of decomposition and determine the rate constant, the mass of oxygen lost to the environment is used as an indicator for the rate of reaction. Therefore, an open system is used and the mass is recorded over time to determine the amount of oxygen produced. Conduct this experiment with example masses of catalase and manganese dioxide to determine which amounts yield a sufficient rate of reaction to be noticed on the scale used in the 90 second time frame. All variables for the reaction were denoted in Figure 1.
Figure 1: Variables of the Reaction ▲
Experimental Design & Setup
Hydrogen Peroxide (H2O2) – [10%]
Manganese Dioxide (MnO2) – 96% purity
Catalase – [1 mg/mL] (Source: Liver Extract)
Thermometer (+/- 1°C) (Range: -10 – 110°C)
Stopwatch (+/- 0.001s)
Balance (+/- 0.01 g)
Measuring cylinder – 100 ml (+/- 1 mL)
Pipette – 1 ml (+/- 0.01 mL)
Measure 100 ml of the 10% hydrogen peroxide in a suitable measuring cylinder and add it to a 250 ml erlenmeyer flask.
Weigh 0.100 g of manganese dioxide
Set up a camera to record the stopwatch and and mass of the system
Add the magnesium dioxide to the erlenmeyer flask with the hydrogen peroxide.
Record the initial temperature using an analog thermometer and the initial mass of the system.
Let the camera record the mass of the system and the stopwatch to allow more precision in measurements
Record the mass of the system in 5 second intervals
Repeat steps 1-4 at a minimum of 10 temperatures between 268 and 331 K.
Repeat steps 1-6 with catalase instead of MnO2
Method is based on: [Order of Reaction Experiments]
Hydrogen peroxide is a known irritant and hazardous substance. It is an oxidiser and is corrosive, and can cause fires when in contact with organic materials. Even though hydrogen peroxide itself is not flammable, when it decomposes it produces oxygen and can therefore increase combustion. Hydrogen peroxide should always be stored in dark and cool containers to slow the natural decomposition and be kept away from sources of contamination such as dust, which can rapidly increase the rate of decomposition. Should any H2O2 spill it must be cleaned with excess amounts of water. When in contact with skin or eyes it must be washed out thoroughly using water to rinse for a minimum of 15 minutes, and additional medical help must be sought to ensure no permanent damage. To decrease risk, always wear safety goggles, a lab coat, and gloves while using hydrogen peroxide and minimise the amount of exposed skin to decrease the probability of contact with the chemical [ORCBS – Hydrogen Peroxide]. Catalase can cause “allergy or asthma symptoms or breathing difficulties if inhaled. Should any be inhaled, remove yourself to fresh air.” In case of skin / eye contact; wash with water for at least 20 minutes and seek medical attention if any symptoms occur. [Safety Data Sheet Catalase]. Manganese Dioxide is “hazardous in case of skin contact (irritant), of eye contact (irritant), of ingestion, of inhalation.” [Sciencelab.com]
This raw data displays the mass of oxygen that was lost in grams per 5 second unit interval at which the mass was recorded.
Figure 2: Raw Data (Catalase) ▲
Figure 3: Raw Data (Manganese Dioxide) ▲
For the catalase, there was clear fizzing and air bubble formation which clearly indicated that gas was being produced. Furthermore there was a continuous change from the initial transparent nature of the hydrogen peroxide into light yellow, which most likely originated from the catalase which is yellow in colour. Moreover, the flask became warmer.
Similarly to the catalase, there was fizzing and air bubble formation. However, the reaction was much more rapid. Grey smoke left the Erlenmeyer flask and the warming of the flask indicated a heavily exothermic reaction since the flask became too hot to touch after around 20 seconds. For lower temperatures the temperature changes was less than when the original solution was at a higher temperature. Furthermore there was a colour change from transparent to grey-black (originated from the MnO2 which is black).
Since this is a first order decomposition reaction, meaning the initial concentration of hydrogen peroxide determines the rate of reaction, 2H2O2 —> 2H2O + O2, the rate equation for the reaction is: rate = k[H2O2]. Since the initial concentration of hydrogen peroxide does not change with temperature, the only thing affecting the rate is the rate constant, k; which changes with temperature. In order to calculate k, the rate and initial concentration can be used in a simple re-arrangement of the rate equation to solve for k. Here k = rate / [H2O2].
Since using 10% hydrogen peroxide solution and the density of hydrogen peroxide is 1.135 g/cm3 [Wikipedia – Hydrogen Peroxide] it was assumed that there are 11.35 grams of hydrogen peroxide in 100 ml of solution; since weighing of the hydrogen peroxide wasn’t possible as it was in solution. The molar mass of hydrogen peroxide is 34.02 g mol-1 [IB, 2014]. Therefore 11.35 / (34.02 x (100 / 1000)) = 3.336272781, which is a molarity of 3.34 mol dm-3.
In order to determine the concentration of hydrogen peroxide over the course of the 5 seconds, the number of moles of oxygen produced were subtracted from the initial concentration with regard to the stoichiometric relationship. To find activation energy, the initial rate was determined by finding the gradient between the concentration of hydrogen peroxide at 0 and 5 seconds and using the formula: m = (y2 – y1) / (x2 – x1). The first 5 seconds were used because there was the largest change in mass and the additional energy caused by the exothermic nature of the reaction has a smaller effect. Since the rate law for this decomposition reaction is: rate = k[H2O2]initial [Bylikin et al,2014], k, the rate constant, at a given temperature; is equal to the initial rate of reaction divided by the initial concentration of hydrogen peroxide. The resulting value of k for each temperature was then converted using the natural log function and plotted against T-1 because the gradient of the resulting function is equal to -Ea/R, which was derived from the Arrhenius equation [Bylikin et al, 2014].
Sample Calculation: Manganese Dioxide at 268 K
Mass at 0 seconds = 0.00 g (+/- 0.01g)
Mass at 5 seconds = 0.15 g (+/- 0.01g)
Oxygen in mol at 5 seconds = 0.15 / 32.00 (molar mass of O2) = 0.469 x 10-2
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