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CALORIMETER EXPERIMENTS & CALCULATION of ENERGY TRANSFER

(c) doc bDoc Brown's Chemistry KS4 science GCSE/IGCSE/O level/A Level Chemistry Revision Notes

PART D Exothermic and Endothermic Energy Changes – Chemical Energetics  Methods of determining energy transfers and calculation of energy changes from calorimetric data.

Experimental methods for obtaining vales for energy transfer changes in chemical reactions are described and how to do the calculations based on calorimeter experiment results. Calculation of energy transferred from experimental data is explained. A simple calorimeter is described and how to obtain energy transfer measurements. Revision notes for GCSE/IGCSE/O Level/basic stuff for GCE Advanced Level AS students. These revision notes on calorimeter experiments, procedures and calculations of energy transfers in chemical reactions should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.


Sub–index for ENERGY CHANGES: 1. Heat changes in chemical/physical changes – exothermic and endothermic  *  2. Reversible reactions and energy changes  *  3. Activation energy and reaction profiles  *  4. Catalysts and activation energy  *  5. Introduction to bond energy/enthalpy calculations  *  6. Calorimeter methods of determining energy changes and examples of experiments  *  7. Energy transfer calculations from calorimeter results (this page)

See also Advanced A Level Energetics–Thermochemistry – Enthalpies of Reaction, Formation & Combustion

and enthalpy calculations from calorimetry data for Advanced A Levels students

6. The experimental determination of energy changes using simple calorimeters

The basic principles of calorimetry

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Using this apparatus you can observe changes in heat energy accompanying the following changes -
(i) salts dissolving in water, (ii) neutralisation reactions, (iii) displacement reactions, (iv) precipitation reactions
and when these reactions take place in solution, the
temperature changes can be measured to calculate and compare the relative heat changes.

This method 7.1 is can be used for any non–combustion reaction that will happen spontaneously at room temperature involving liquids or solid reacting with a liquid. The reactants are weighed in if solid and a known volume of any liquid (usually water or aqueous solution). The mixture could be a salt and water (heat change on dissolving) or an acid and an alkali solution (heat change of neutralisation). It doesn't matter whether the change is exothermic (heat released or given out, temperature increases) or endothermic (heat absorbed or taken in, temperature decreases). See calculations below.

For extra insulation you can place the polystyrene container in a larger beaker containing a thick layer of cotton wool wrapped around.

You measure the temperature of the reactants at the start and the final maximum/minimum temperature when the reaction is done. Subtracting one from the other gives the temperature change.

The main problem is heat loss from the calorimeter, though a poorly conducting cup and a sealing lid to stop convection in air, both help reduce the heat loss and quite accurate results can be obtained. You can also stand the plastic cup in a beaker of cotton wool to provide even greater thermal insulation.

From the temperature rise or temperature fall and the heat capacity of the water you can calculate how much heat was released or how much heat energy was absorbed by a specific quantities of chemicals reacting (or even the energy transfer when a salt dissolves in water, which can be exothermic or endothermic).

Energy transferred in J = mass of solution in grams X specific heat capacity of water (4.2J/goC) X temperature change in oC

You can then calculate the amount of energy released/absorbed per gram or per mole.

You can use this method for adding a solid to water or a solution or mixing two solutions of reactants. In either case you must carefully measure the initial temperature of all the reactants.

e.g. (i) An acid plus an alkali gives and exothermic neutralisation reaction, temperature rises,

(ii) adding the salt ammonium nitrate to water gives a temperature fall, an endothermic dissolving change

(iii) adding zinc to copper sulfate produces an exothermic displacement reaction shown by the rise in temperature.

(c) doc b This method 7.2 is specifically for determining the heat energy released (given out) for burning fuels. The burner is weighed before and after combustion to get the mass of liquid fuel burned. The thermometer records the temperature rise of the known mass of water (1g = 1cm3).

The heat from the fuel combustion heats up the water. From the heat capacity of the water and the temperature rise you can calculate how much heat was released by a specific mass of fuel.

You measure the temperature of the reactants at the start and the final maximum/minimum temperature when the reaction is done. Subtracting one from the other gives the temperature change.

Energy transferred in J = mass of water in grams X specific heat capacity of water (4.2J/goC) X temperature change in oC

You can then calculate the amount of energy released per gram or per mole.

You can use this system to compare the heat output from burning various fuels. The bigger the temperature rise, the more heat energy is released. See calculations below for expressing calorific values.

BUT you must conduct the experiments under 'fair test' conditions.

apart from repeating experiments (to eliminate anomalous results), you must use the same burner & wick (if possible), same volume (mass) of water, same calorimeter, burn for the same length of time, same insulation set-up

This is a very inaccurate method because of huge losses of heat e.g. radiation from the flame and calorimeter, conduction through the copper calorimeter, convection from the flame gases passing by the calorimeter etc. BUT, at least using the same burner and set–up, you can do a reasonable comparison of the heat output of different fuels. You can simple hydrocarbons like hexane, alcohols like ethanol and even vegetable oils and it is possible to do a crude calibration of the calorimeter using a fuel of known energy output on complete combustion.

You can investigate the temperature rise produced in a known mass of water by the combustion of the alcohols, methanol, ethanol, propanol, butanol using this simple calorimeter system.

See GCSE/IGCSE/O Level notes on chemistry of alcohols

BOMB CALORIMETER Advanced Level students need to know about the bomb calorimeter for determining enthalpies of combustion as well as the methods described above.

top7. Calculations from the experimental calorimeter results

  • PLEASE NOTE that section 7. is for higher GCSE students and an introduction for advanced level students of how to do energy change (enthalpy change) calculations from experimental data.

  • The calculation method described below applies to both experimental methods 6.1 and 6.2 described above.

  • You need to know the following:

    • the mass of material reacting in the calorimeter (or their concentrations and volume),

    • the mass of water in the calorimeter,

    • the temperature change (always a rise for method 6.2 combustion),

    • the specific heat capacity of water, (shorthand is SHCwater), and this is 4.2J/goC (for advanced 4.2J g–1 K–1),

      • this means it means the addition of 4.2 J of heat energy to raise the temperature of 1g of water by 1oC.

  • Example 7.1 typical of calorimeter method 7.1

    • Measuring the energy transfer when a salt dissolves in water

    • 5g of ammonium nitrate (NH4NO3) was dissolved in 50cm3 of water (50g) and the temperature fell from 22oC to 14oC.

    • Temperature change = 22 – 14 = 8oC (endothermic, temperature fall, heat energy absorbed)

    • Heat absorbed by the water = mass of water x SHCwater x temperature

      • = 50 x 4.2 x 8 = 1680 J (for 5g)

      • heat energy absorbed on dissolving = 1680 / 5 = 336 J/g of NH4NO3 

    • this energy change can be also expressed on a molar basis.

      • Relative atomic masses Ar: N = 14, H = 1, O = 16

      • Mr(NH4NO3) = 14 + (1 x 4) + 14 + (3 x 16) = 80, so 1 mole = 80g

      • Heat absorbed by dissolving 1 mole of NH4NO3 = 80 x 336 = 26880 J/mole

      • At A level this will be expressed as enthalpy of solution = ΔHsolution = +26.88 kJ/mol

      • The data book value is +26 kJmol–1

  • Example 7.2 typical of calorimeter method 7.2

    • A typical fuel combustion reaction

    • 100 cm3 of water (100g) was measured into the calorimeter.

    • The spirit burner contained the fuel ethanol C2H5OH ('alcohol') and weighed 18.62g at the start.

    • The initial temperature of the water is taken.

    • After burning some time, the flame is extinguished, the water stirred gently and the final water temperature is taken to get the temperature rise.

    • The burner and fuel are then reweighed to see how much fuel had been burned.

    • After burning it weighed 17.14g and the temperature of the water rose from 18 to 89oC.

    • The temperature rise = 89 – 18 = 71oC (exothermic, heat energy given out).

    • Mass of fuel burned = 18.62–17.14 = 1.48g.

    • Heat absorbed by the water = mass of water x SHCwater x temperature

      • = 100 x 4.2 x 71 = 29820 J (for 1.48g)

      • heat energy released per g = energy supplied in J / mass of fuel burned in g

      • heat energy released on burning = 29820 / 1.48 = 20149 J/g of C2H5OH

    • this energy change can be also expressed on a molar basis.

      • Relative atomic masses Ar: C = 12, H = 1, O = 16

      • Mr(C2H5OH) = (2 x 12) + (1 x 5) + 16 + 16 = 46, so 1 mole = 46g

      • Heat released (given out) by 1 mole of C2H5OH = 46 x 20149 = 926854 J/mole or 927 kJ/mol (3 sf)

      • At AS level this will be expressed as the ...

      • Enthalpy of combustion of ethanol = ΔHcombustion (ethanol) = –927 kJmol–1

      • This means 926.9 kJ of heat energy is released on burning 46g of ethanol ('alcohol').

      • The data book value for the heat of combustion of ethanol is –1367 kJmol–1, showing lots of heat loss in the experiment!

      • It is possible to get more accurate values by calibrating the calorimeter with a substance whose energy release on combustion is known.

    • Need GCSE level acid-alkali example but see example 5. on an A Level enthalpy calculation page

  • In some exothermic changes, no heat is not released e.g. in batteries and fuel cells, where the energy is released as electrical energy.

  • See also Advanced Level Chemistry Notes on Experimental methods for determining enthalpy changes and treatment of results

  • See other pages related to fuels


Sub–index for ENERGY CHANGES: 1. Heat changes in chemical/physical changes – exothermic and endothermic  *  2. Reversible reactions and energy changes  *  3. Activation energy and reaction profiles  *  4. Catalysts and activation energy  *  5. Introduction to bond energy/enthalpy calculations  *  6. Calorimeter methods of determining energy changes  *  7. Energy transfer calculations from calorimeter results

See also Advanced A Level Energetics–Thermochemistry – Enthalpies of Reaction, Formation & Combustion

and enthalpy calculations from calorimetry data for Advanced A Levels students


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