THERMOCHEMISTRY ENERGETICS/ENTHALPY

INTRODUCTION All reactions require energy to break bonds in the reactants, and all reactions give off energy when new bonds form to make products The difference in energy required to break the bonds and to make the new bonds can tell you whether the reaction is endothermic or exothermic.

PRIOR KNOWLEDGE FROM GCSE Before you start remind yourself of what you already should know: Atoms bond together to fill their outer shells of electrons Atoms may gain or lose electrons to get full outer shells of electrons Chemical reactions may give out heat – exothermic reactions. Chemical reactions may take in heat energy from their surroundings – endothermic reactions

4.1 Endothermic and exothermic reactions Thermochemistry is the study of heat changes during chemical reactions. When a chemical reaction takes place bonds must be broken and new bonds must be formed. Breaking bonds uses energy from the surroundings; forming bonds gives off heat energy to the surroundings. During a reaction there is always an overall change in energy. The energy may be in different forms – light, electrical or, most usually, heat.

Define the terms endothermic and exothermic. Endothermic reactions are reactions that take in heat from the surroundings by the end of the reaction. Exothermic reactions are reactions that have given out heat by the end of the reaction. It is always the case that a reaction that is endothermic in one direction is exothermic in the reverse direction.

ENERGY PROFILE DIAGRAM OF EXOTHERMIC REACTIONS The diagram below shows an energy profile diagram for an exothermic reaction In an exothermic reaction the products have less energy than the reactants, so energy is given out Also in an exothermic reaction more bonds are made than are broken, so overall energy is given out Activation energy (Ea) – is the minimal amount of energy needed to start the reaction (i.e. minimum amount of energy to break bonds in the reactants).

ENERGY PROFILE DIAGRAM OF ENDOTHERMIC REACTIONS The diagram below shows an energy profile diagram for an endothermic reaction In an endothermic reaction the products have more energy than the reactants, so energy is taken in Also in an endothermic reaction more bonds are broken than are made, so overall energy is taken in

EXAMPLES OF EXO AND ENDOTHERMIC REACTIONS DEMO 1 – THERMAL DECOMPOSITION OF LIMESTONE DEMO 2 – HEATING HYDRATED COPPER II SULPHATE CRYSTALS

QUANTITIES Energy is always measured in kJmol-1 in Chemistry In general we always talk about 1 mole of a fuel burning So when writing balanced equations such as: CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) DH = -890kJmol-1 This means that when ONE mole of methane reacts with 2 MOLES of oxygen then 890kJ of energy is given out OR C2H6(g) + 3½O2(g)  2CO2(g) + 3H2O(l) DH = -86kJmol-1 (NOT 2C2H6(g) + 7O2(g)  4CO2(g) + 6H2O(l)) So in this case the value of -86kJmol-1 is the energy given off by 1 MOLE of ethane NOT for two moles of ethane

REMEMBER FROM GCSE THAT If DH is negative in value then the reaction is exothermic If DH is positive in value then the reaction is endothermic So, C(s) + O2(g)  CO2(g) DH = -394kJmol-1 (This means that 394kJ of energy is given out for 1 mole of carbon burnt NH4NO3(s) + (aq)  NH4NO3(aq) DH = +26kJmol-1 (This means that 26kJ of energy is absorbed by 1 mole of NH4NO3 solid)

So why do we bother to work out the energy given out for different fuels: ANSWER: It allow us to compare the efficiency of different fuels.

WORKING OUT ENERGY OF COMBUSTION OF FUELS IN kJg-1 Chemists generally refer to the energy given out when a fuel burns in kJmol-1 because this compares the same number of molecules of each fuel. For use as fuels it is sometimes better to convert the units from kJmol-1 to kJg-1 (OR the energy density) of a fuel NOTICE: That petrol stores significantly more energy than either ethanol or methanol. This is important when considering vehicles fuelled by either of these alcohols.

Also at first sight we might think the hydrogen energy density is amazing, but that’s not the only thing to consider. All the other three fuels are liquids whereas hydrogen is a gas. Although hydrogen stores more energy per gram, a gram of hydrogen takes a lot of space because of the low density of gases. How to store the hydrogen efficiently is a challenge for designers.

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