Academy of Art University Chemistry Worksheet

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Sn2020

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Academy of Art University

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I have HW in chemistry class, they are 5 Qs and I need the full answers to give them to my teacher, Just watch the videos and answers the Qs

QUESTION 1

  1. Enthalpy of decomposition of H2O2. Please show work in a scratch paper. You must submit your scratch paper to receive credits.Initial temperature (from plot): °CFinal temperature (from plot): °C∆T: °CHeat change for the reaction mixture (qsolution): JHeat of the reaction: KMoles of H2O2 reacted: molEnthalpy of decomposition (experiment): kJ/mol



QUESTION 2

  1. How does your calculated value of the enthalpy change of formation of H2O2 (aq) compare to the literature value of –191.2 kJ/mol? Please calculate the % error.
    % error = [a]


QUESTION 3

  1. Record your observations of color during the experiment in the table below:
    Observation of color
    Fe(NO3)3 catalyst
    H2O2 solution before catalyst was added
    H2O2 solution after catalyst was added
    H2O2 solution at the end of the experiment




QUESTION 4

  1. What is the purpose of the catalyst?


QUESTION 5

  1. What is the evidence that the catalyst is not consumed during the overall reaction?

Unformatted Attachment Preview

CHEM 118 Summer 2020 Experiment 5 THE ENTHALPY OF DECOMPOSITION OF HYDROGEN PEROXIDE INTRODUCTION Thermochemistry investigates the relationship between chemical and physical phenomena and energy changes involving heat. Practical applications of thermochemistry include the development of alternative fuel sources, such as fuel cells, hybrid gas-electric cars or gasoline supplemented with ethanol. Thermochemistry is also important because the forces holding molecules or ionic compounds together are related to the heat evolved or absorbed in a chemical reaction or a physical change. Therefore, chemists are interested in the thermochemistry of numerous processes, whether it is the change in physical state of a substance, the solubility of a salt, or the decomposition of a material, as you will do in today’s lab. Figure 1. Schematic representation of a calorimeter. The amount of heat evolved or absorbed in a chemical or physical process can be studied using a calorimeter. In the simplified schematic of a calorimeter shown below (Fig. 1), the "system" is placed in a well-insulated vessel surrounded by water which acts as the “surroundings”. A thermometer is used to measure the heat transferred to the system from the surroundings, or from the system to the surroundings. Ideally, only the water would be the "surroundings" and the vessel would not allow heat to pass. In reality, the vessel does allow heat to pass from the water to the rest of the universe, and this is typically accounted for by determining the heat change for the vessel (we will use a Styrofoam cup to make our measurements). E N T H A L P Y O F D E C O M P O S I T I O N O F H 2 O 2 All chemical reactions involve an exchange of heat energy; therefore, it seems logical that you could follow a reaction by measuring the enthalpy change (ΔH). But it is often not possible to directly measure the heat energy change of the reactants and products (the system). We can, however, measure the heat change that occurs in the surroundings by monitoring temperature changes. If we conduct a reaction between two substances in aqueous solution, then the enthalpy of the reaction can be indirectly calculated with the following equation: q (joules) = m (g) • DT (K) • Cp (J/g•K) The term q represents the heat energy that is gained or lost. Cp is the specific heat of water, m is the mass of water, and ΔT is the temperature change of the reaction mixture. The specific heat and mass of water are used because water will either gain or lose heat energy in a reaction that occurs in aqueous solution. Furthermore, according to a principle known as Hess’s law, the enthalpy changes of a series of reactions can be combined to calculate the enthalpy change of a reaction that is the sum of the components of the series. This experiment involves the decomposition of aqueous hydrogen peroxide (3%) using iron III nitrate as a catalyst. The reaction is shown below: 2 H2O2 (aq) → 2 H2O (l) + O2 (g) You will be experimentally measure the enthalpy change for this decomposition. A solution of hydrogen peroxide will be added to a calorimeter and its temperature will be monitored for a short time. You will then add a small amount of iron(III) nitrate to catalyze the reaction. This will produce a measureable temperature change. It also produces a distinct color change that can be followed during the course of the reaction. The iron (III) nitrate solution itself has a pale yellow color due to the presence of Fe(OH)(H2O)52+. When added to the aqueous hydrogen peroxide, a dark amber color is produced. As the reaction proceeds and the hydrogen peroxide is decomposed, the dark amber color of the solution will revert back to a pale yellow color. The temperature change that occurs during the course of the reaction will be plotted over time. Using this plot, the mass of the solution, its specific heat, the calorimeter constant (this can be determined experimentally also), and the temperature change, you will be able to determine the enthalpy for the decomposition of hydrogen peroxide. EXPERIMENTAL D E T E R M I N A T I O N PROCEDURE O F T H E C O N C E N T R A T I O N O F H 2 O 2 Before you can run the calorimetry experiment, you will need to titrate the hydrogen peroxide solution in order to determine its concentration. The solution will be titrated with an acidified KMnO4 solution according to the following procedures: Pipette 2.00 mL of the hydrogen peroxide solution into a 250 mL Erlenmeyer flask. Add 10 mL of 4.0 M H2SO4 and enough water to bring the level of the solution to about 75 mL. Fill a burette with the standardized KMnO4 solution that is approximately 0.02 M. BE SURE to record the exact molarity of this standardized solution. Then titrate the hydrogen peroxide solution to the point where the pink color of the added permanganate ceases to turn colorless. D E T E R M I N A T I O N O F T H E H Y D R O G E N P E R O X I D E E N T H A L P Y O F D E C O M P O S I T I O N O F While you are doing this part of the experiment, pay attention to and record colors of solutions and changes in color during the course of the reaction. Record these observations on your report form. Add 50.0 mL of the hydrogen peroxide solution to the Styrofoam cup using a graduated cylinder. Place a magnetic stir bar in the cup and turn the stir plate on at a low setting. Secure and place a thermometer such that it will NOT be struck by the stir bar during the experiment and that it can be easily read. Measure the initial temperature of the hydrogen peroxide solution. This will represent the zero minute reading. Using a graduated cylinder, add 10.0 mL of 0.50 M Fe(NO3)3 to the hydrogen peroxide solution. Remember to record the color of the reaction mixture immediately after the addition of catalyst. Measure and record the temperature. Continue measuring and recording the temperature, at one-minute intervals, until the 15 minute mark is reached. DATA PROCESSING Prepare a plot of temperature versus time (be sure to attach this plot to your data report form). • The last data points should be fairly linear. Extrapolate this last data points, again to the point of mixing. The temperature at this point represents the final temperature. • Determine ∆T. Calculate qsolution, which represents the heat change for the reaction mixture. Remember you will need the mass of water, specific heat of water, and ∆T. We assume that the heat change for the calorimeter (qcalorimeter) is negligible. Use the following relationship to calculate this value, where Csolution = 4.179 J/°C: qsolution = msolutionCsolution∆Tsolution Now you can determine the heat change of the reaction: qreaction = –qsolution Calculate the moles of hydrogen peroxide that reacted for this particular experiment. Recall that n = molarity ´ volume. Now you can calculate ∆H for this reaction. This is the heat change per mole of hydrogen peroxide reacted. (Be sure to indicate the sign of ∆H). • Also determine the literature value for the enthalpy of decomposition of H2O2 (aq) from the enthalpies of formation of H2O2 (aq), H2O (l), and O2 (g) found in your textbook. • Calculate the % error in the enthalpy of decomposition of hydrogen peroxide for this experiment. Finally, calculate the enthalpy of formation of hydrogen peroxide from YOUR experimental value for the enthalpy of decomposition of hydrogen peroxide AND the literature value for the enthalpy of formation of H2O (l). Use Hess’s law to do this.
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Explanation & Answer

Attached.

1. Concentration of H2O2 = 0.875 M
Initial temperature = 21.5 ℃
Final temperature = 37 ℃
So, ∆T = 37 − 21.5 = 15.5℃
Specific heat of solution = 4.179 J/g℃
Density of the solution = 1g/ml
Volume of the solution = 60 ml
Heat change of solution:
g
J
∆q = 1 × 60ml × 4.179
× 15.5℃ = 3886.47J
ml
g℃
Hea...


Anonymous
Awesome! Perfect study aid.

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