Calorimetry (Edexcel AS Chemistry)

Revision Note

Test yourself
Richard

Author

Richard

Last updated

Calculating Energy Transferred, Q

Measuring enthalpy changes

  • Calorimetry is the measurement enthalpy changes in chemical reactions
  • A simple calorimeter can be made from a polystyrene drinking cup, a vacuum flask or metal canChemical Energetics Calorimeter, downloadable AS & A Level Chemistry revision notes

A polystyrene cup can act as a calorimeter to find enthalpy changes in a chemical reaction

  • The energy needed to increase the temperature of 1 g of a substance by 1 oC is called the specific heat capacity (c ) of the liquid
  • The specific heat capacity of water is 4.18 J g-1 K-1
  • The energy transferred as heat can be calculated by:Chemical Energetics Equation for Calculating Energy Transferred in Calorimeter, downloadable AS & A Level Chemistry revision notes

Equation for calculating energy transferred in a calorimeter

Worked example

Specific heat capacity calculations

In a calorimetry experiment 2.50 g of methane is burnt in excess oxygen.

30% of the energy released during the combustion is absorbed by 500 g of water, the temperature of which rises from 25 °C to 68 °C.

The specific heat capacity of water is 4.18 J g-1 K−1

What is the total energy released per gram of methane burnt?

   Answer

   Step 1:

    • q = m x c x ΔT
    • m (of water) = 500 g
    • c (of water) = 4.18 J g-1 oC-1
    • ΔT (of water) = 68 oC - 25 oC = 43 oC

   Step 2: 

    • q = 500 x 4.18 x 43 = 89 870 J

   Step 3: 

    • This is only 30% of the total energy released by methane
    • Total energy x 0.3 = 89 870 J
    • Total energy = 299 567 J

   Step 4:  

    • This is released by 2.50 g of methane
    • Energy released by 1.00 g of methane = 299 567 ÷ 2.50 = 120 000 J g-1 (to 3 s.f.) or 120 kJ g-1

Calculating Enthalpy Changes

  • Aqueous solutions of acid, alkalis and salts are assumed to be largely water so you can just use the m and c values of water when calculating the energy transferred.
  • To calculate any changes in enthalpy per mole of a reactant or product the following relationship can be used:

ΔHq over n space or space fraction numerator m space cross times space c space cross times straight capital delta T over denominator n end fraction

  • When there is a rise in temperature, the value for ΔH becomes negative suggesting that the reaction is exothermic
    • This means that your value should be negative for an exothermic reaction, e.g. combustion
  • When the temperature falls, the value for ΔH becomes positive suggesting that the reaction is endothermic
    • This means that your value should be positive for an endothermic reaction, e.g. combustion

Worked example

1.50 g of an organic liquid (Mr = 58.0) underwent complete combustion. The heat formed raised the temperature of 100 g of water from 20 oC to 75 oC.

Calculate the enthalpy of combustion for the organic liquid

Answer

Step 1: Calculate the energy released by the organic liquid

    • mcΔT
    • = 100 x 4.18 x (75 - 20)
    • Q = 22990 J
    • Q = 22.99 kJ

Step 2: Calculate the number of moles of the organic liquid

    • Number of moles = equals fraction numerator m a s s over denominator m o l a r space m a s s end fraction equals fraction numerator 1.50 over denominator 58.0 end fraction equals0.0259 moles (to 3s.f.)

Step 3: Calculate the enthalpy change of combustion

    • ΔcHθ equals Q over n equals fraction numerator 22.99 over denominator 0.0259 end fraction equals-887 kJ mol-1 (to 3s.f.)
      • Remember, combustion is an exothermic process and will, therefore, be a negative enthalpy change value

You've read 0 of your 10 free revision notes

Unlock more, it's free!

Join the 100,000+ Students that ❤️ Save My Exams

the (exam) results speak for themselves:

Did this page help you?

Richard

Author: Richard

Expertise: Chemistry

Richard has taught Chemistry for over 15 years as well as working as a science tutor, examiner, content creator and author. He wasn’t the greatest at exams and only discovered how to revise in his final year at university. That knowledge made him want to help students learn how to revise, challenge them to think about what they actually know and hopefully succeed; so here he is, happily, at SME.