Calculating DHdis Using Hess’ Law

the system that are always at the same temperature as the reactants or products because of the experimental design, for example, the inside wall of the calorimeter, the stirrer, the thermometer, and the solvent. Since DH is a function of state and independent of the path taken:DHdec = DHA + DHB (4)The first step is carried out adiabatically so by definition DHA = 0 and thereforeDHdec = DHB (5)If we assume that Cp is temperature independent, then DH for the system is given byDHB = DHdec = Cp(T1 - T0) = CpDT (6)DHB can be calculated asDHB = Cp(Sys)DT + Cp(H2O)DT + Cp(O2)DT (7)The contributions to the change in heat from the products of the reaction (H2O and O2) are small compared with the contribution due to the calorimeter.

The last two terms can be ignored andDHB = Cp(Sys)DT (8)By measuring the heat capacity of the whole calorimeter Cp(Sys), the temperature change due to the decomposition of H2O2, and the number of moles of H2O2 decomposed, the molar heat of decomposition of H2O2 can be determined. The heat capacity of the calorimeter filled with water is determined by measuring the heat evolved by the reaction of H2SO4 with water DHsol. H2SO4 is added to the calorimeter and the change in temperature is measured. The number of moles of H2SO4 is determined by titration.

Values for DHsol (H2SO4) per 100 ml depend on the final molarity and can be found in tables.(9)The decomposition of H2O2 in water is a slow reaction. In fact, it does not proceed to any appreciable extent without the addition of a catalyst. Therefore the H2O2 can be added to the calorimeter

and the concentration determined by titrating aliquots taken from the calorimeter before the reaction is initiated with a catalyst.

A solution of permanganate MnO4- is used as the titrant and the following oxidation-reduction reaction occurs:5 H2O2 + 2 MnO4- + 6 H+ ? 5 O2 + 8 H2O + 2 Mn+2 (10)MnO4- is intensely colored whereas the reduction product Mn+2 is practically colorless. The equivalence point in these titrations with MnO4- is a faint pink color. The pink color of permanganate that is just barely detectable to the eye corresponds to about 10-4 M.MnO2 is added to the calorimeter to catalyze the decomposition of H2O2 and the change in temperature DT is measured.

The molar heat of decomposition DHdec is determined using DT, Cp(Sys), and the number of moles of H2O2. Using DHdec and the heats of reaction for processes given in Table 1, the heat of dissociation DHdis can be determined. You need to arrange the equations in Table 1 such that they sum to giveH2O2(g) ? 2 OH(g) DHdiss (11)In summary, if we make only one measurement, that of DHdec and use Hess' Law, the dissociation energy of HO-OH can be determined. To obtain DHdec, the heat capacity of the calorimeter Cp(S), the amount of H2O2 decomposed, and the temperature change due to the reaction must be determined.III.

PROCEDUREA. Determination of the Heat Capacity of the Calorimeter1. Before starting, ensure that the calorimeter, stirrer, and thermistor are clean and dry. Transfer 235 mL of distilled water into the calorimeter.

Clamp the thermistor on a ring stand and make sure that the spin bar will not strike the thermistor. Calibrate the LabWorks thermistor (review

Expt 2). Caution: do not place stirrer on top of the ring stand; this will ruin the magnet.2.

Switch on the stirrer, adjust to a moderate speed, and follow the temperature until a steady temperature reading is obtained. Because the temperature changes are so small in this experiment, all temperature readings need to be made carefully. Record temperature vs. time using the procedure for the LabWorks thermistor from Experiment 2.

3. Once a steady temperature has been reached, add approximately 15 mL of concentrated sulfuric acid using a small graduated cylinder. Take care that the acid is added directly to the water. Continue taking temperature readings at intervals until the temperature remains steady or falls regularly for at least 10 minutes.4. Review Appendix B from Experiment 1 and Harris 2-6 on pipetting techniques.

Open the calorimeter and determine the concentration of the dilute sulfuric acid by titrating two 5 mL aliquots (use a pipet !) of the acid with standard 0.5 M NaOH using phenolphthalein as an indicator.5. The temperature rise due to the addition of the sulfuric acid may be determined by plotting the temperature, T, against the time and extrapolating back to the time when the acid was added.

Because of the time required to completely mix the solution some of the first temperature readings may be higher or lower than should be the case and should be disregarded when extrapolating the graph. Only the linear portion of the graph corresponding to a steadily decreasing temperature should be extrapolated back to zero time. The difference between the temperature of the water before the addition of the acid and the extrapolated temperature at t=0 will

be the observed temperature change. (See Fig. 1, Expt. 2)6.

The quantity of heat evolved by the reaction of H2SO4 with water may be determined from the measured concentration of acid and the following data on the heat of solution of H2SO4, DHsol(H2SO4), given in terms of Joules per 100 mL of H2O or, in our case, total solution volume.DHsol(H2SO4)J/100 mL H2OFinal Molarity of Solution-8030.109-1,0100.139-1,9800.276-3,2800.

459Plot the data and, by extrapolation, find the heat corresponding to the molarity of the acid in the calorimeter. From this value and your experimental temperature change, calculate Cp(S) per 100 mL of solution.7. Repeat this procedure and obtain an average of the effective heat capacities.B.

Determination of Heat of Decomposition1. Again, be certain that the styrofoam cup, stirrer, and thermistor are clean and dry before starting. Transfer 200 mL of distilled water into the calorimeter and add 50 mL of 6% hydrogen peroxide.2. Remove two 5 mL aliquots (use a pipet !) and save for subsequent standardization against the 0.04 M potassium permanganate (KMnO4) solution following the procedure in (6) below.

3. Position the cover, stirrer, and thermistor carefully and set the stirrer for a moderate speed. Measure the temperature until a steady reading is obtained.4.

Add approximately 6 g of powdered MnO2 to catalyze the decomposition of the peroxide. Continue taking temperature readings until the temperature remains steady or is falling regularly.5. Obtain the temperature rise by taking the difference between the maximum temperature and the temperature before addition of the MnO2.

(Extrapolation takes 30 minutes or more and is not generally worthwhile in this part.)6. To determine the concentration of H2O2 in the initial solution, dilute the 5 mL

aliquots with 100 mL of distilled water, add 10 mL of diluted (3M) sulfuric acid and titrate with 0.04 M permanganate.

7. Repeat if necessary and if time allows.IV. CALCULATIONS1.

Plot the data for DHsol(H2SO4) versus H2SO4 molarity. Indicate your value of H2SO4 molarity and find the heat evolved per 100 mL of solution. From this value and your experimental temperature change, calculate Cp(Sys) per 100 mL of solution.2. Calculate the number of moles of H2O2 in the calorimeter and the resulting value for DHdec in units of kJ/mol of H2O2.

3. Finally, calculate values for the O-O bond energy in hydrogen peroxide. Estimate uncertainties in all values. Compare your final thermochemical values to those reported in the literature. Check Oxtoby and Nachtrieb, CRC, JANAF Tables, or NBS Tables of thermodynamic values. Usually, only heats of formation are tabulated.

A literature value for D(HO-OH) can be calculated from these.-Investigate the enthalpy values for the reaction between calcium carbonate and calcium oxide with Hydrogen Chloride.Ruth Spink12 AquinasChemistry courseworkAimI will carry out this practical to investigate the enthalpy values for the reaction between calcium carbonate and calcium oxide with Hydrogen Chloride. I will collect the results and analyse them to reach a conclusion about the reactions between the reactants.TheoryI will measure the temperature changes when calcium oxide and calcium carbonate react with hydrochloric acid solution.

I can use a Hess' Law cycle to calculate the enthalpy changes using the equation.Enthalpy change = mass of liquid x temperature change x specific heat capacityThis equation will help me to calculate the enthalpy changes, and which reaction was more exothermic, and why.I drew a Hess' Law cycle to illustrate the enthalpy

values.CaCO� (s) CaO (s) + CO� (g)H�H� H�CaCl� (aq)Fair test� We used scales correct to 2 decimal places.� We measured the calcium oxide and calcium carbonate straight into the beaker to avoid leaving particles on the paper I would otherwise have measured it on.

� We measured the hydrochloric acid using a measuring cylinder including the meniscus to make it accurate, and poured out as much as possible. We couldn't wash it out and this would have affected the concentration of the acid and the rate of reaction.� We left as small time gaps as possible to get as accurate measurements as possible.Safety� Used pestle and mortar and spatula to handle chemicals.� Wore lab coat and goggles.� Wore gloves to handle chemicals.

� Handled equipment safely.Equipment listMeasuring cylinderScalesBeakerSpatulaPestle ; MortarThermometerMethod1. Weigh out a beaker containing between 2.4g and 2.6g of calcium carbonate.

Record the results.2. Using the measuring cylinder provided place 50cm� of 2 mol dm� hydrochloric acid in a 250cm� beaker.3. Measure the temperature of the acid using the thermometer provided and record this value. The thermometer should not be left unsupported or it may break.

4. Add the calcium carbonate to the acid. Take the temperature, then again when the reaction is complete. Record this value.5. Weigh the beaker afterwards.

6. Repeat steps 1 to 5 using an accurately weighed mass of calcium oxide, in the range 1.3g to 1.5g, instead of calcium carbonate.Results1st time2nd timeAverageMass of CaCO� ; beaker102.

60g104.37g103.49gMass of empty beaker100.13g101.95g101.

04gMass of CaCO�2.47g2.42g2.45gTemp.

of acid initially18�c19�c18.5�cTemp. after mixing19�c20�c19.5�cTemp. change during reaction3�c1�c2�cMass of CaO and beaker102.

66g103.99g103.33gMass of empty beaker101.23g102.

59g101.91gMass of CaO1.42g1.41g1.

41gTemp. of acid initially18�c18.5�c18.5�cTemp.

after mixing19�c20�c20�cTemp. change during reaction8�c9.5�c8.75�cI will use the average calculations in my

analysis and conclusion as they represent a fairer result.Calcium carbonate + hydrochloric acidEnthalpy change = Mass of liquid x specific heat capacity x temperature changeEnthalpy change = 50cm� x 4.2 x 2�cEnthalpy change = 420 joulesH� = - 420 j molCalcium oxide ; hydrochloric acidEnthalpy change = Mass of liquid x specific heat capacity x temperature changeEnthalpy change = 50cm� x 4.

2 x 8.75�cEnthalpy change = 1837.5 joulesH� = - 1837.5 j molEvaluationThe main problem was that my experiment lost a lot of heat into the atmosphere and into the beaker glass.

This meant that our calculations would have been very inaccurate as they didn't account for the energy lost as heat and through other methods.As our reactants weren't totally pure, we cannot say whether the masses recorded were correct. For example, when I measured out the calcium carbonate and the calcium oxide, not all of it was pure and I tried to remove some rocks, so I only crushed the reactants I needed.I had no problems with the equipment.There was also a problem with the hydrochloric acid as it may have been contaminated and it might not all have reacted the way it should have.ConclusionFrom my results I can tell that both reactions were exothermic, but that the reaction with calcium oxide was much more exothermic than the reaction with calcium carbonate.

They were different because the bonds that were made in the calcium oxide reaction required less energy to be made than in the calcium carbonate reaction. My Hess' Law cycle can be labelled correctly. I have calculated one enthalpy change, but I would need more data to complete my cycle. I also

need to find out how inaccurate my results are.