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- 1. 3 Enthalpy
- 2. <ul><li>Units SI unit = joule </li></ul><ul><li>1KJ = 1000J = 239.0 cal </li></ul><ul><li>1st law of Thermodynamics </li></ul><ul><li>The total energy of the universe is constant </li></ul><ul><li>i.e energy cannot be created or destroyed but can be changed from one form to another. </li></ul>
- 3. <ul><li>Heat (or thermal energy) q is the energy transferred between a system and its surroundings as a result of a difference in temperature only. </li></ul><ul><li>All other forms of energy transfer (mechanical, electrical etc) involve some type of work, w, the energy transferred when an object is moved by a force. </li></ul><ul><li>State functions A state function is a property dependant on the current state of the system (e.g. its composition, volume, temp etc) </li></ul><ul><li>It is independent of the path the system took to reach that state. </li></ul>
- 4. <ul><li>Energy of a system is a state function </li></ul><ul><li>So E is a constant for any given change but q and w may vary </li></ul><ul><li>(q and w are not state functions) </li></ul><ul><li>and E = q + w </li></ul>
- 5. <ul><li>Energy can be converted from one form to another </li></ul><ul><li>e.g. from mechanical heat electrical light etc. </li></ul><ul><li>Energy is released when bonds are made </li></ul><ul><li>Energy is used when bonds are broken </li></ul><ul><li>Chemical reactions involve </li></ul><ul><li>Bond breaking </li></ul><ul><li>Bond forming </li></ul>
- 6. <ul><li>Energy changes for an exothermic reaction – one where heat is released to the surroundings </li></ul>
- 7. Energy changes for an endothermic reaction – one where heat is absorbed from the surroundings
- 8. <ul><li>Examples of exothermic reactions </li></ul><ul><li>neutralisation </li></ul><ul><li>burning hydrocarbons </li></ul><ul><li>respiration </li></ul><ul><li>Examples of endothermic reactions </li></ul><ul><li>photosynthesis </li></ul><ul><li>dissolving ammonium nitrate in water </li></ul>
- 9. <ul><li>The chemical energy which a system possesses is its enthalpy. Enthalpy is the change in energy at constant pressure </li></ul><ul><li>symbol H </li></ul><ul><li>And H = H products – H reactants </li></ul><ul><li>if energy is absorbed by a system H is positive </li></ul><ul><li>if energy is released by a system H is negative </li></ul>
- 10. Energy changes for an endothermic reaction A—B + C—D Reactants A—C + B—D Products A B C D Bond breaking Bond forming Overall energy change Products have more energy than reactants E
- 11. <ul><li>Energy changes for an exothermic reaction </li></ul>A—B + C—D A—C + B—D Products Bond forming Bond breaking Overall energy change E A B C D Products have less energy than reactants Reactants
- 12. <ul><li>Energy Diagrams </li></ul>E Activation energy H Reactants Products An exothermic reaction Time
- 13. <ul><li>Reactants </li></ul>Activation energy H Products An endothermic reaction E Reactants Time
- 14. <ul><li>Standard enthalpy of formation is the heat absorbed when 1 mole of a substance is formed from its elements in their standard state. </li></ul><ul><li>The standard state of a substance is 1 mole of the substance in a specified state (solid, liquid, gas) at 1 atmosphere of pressure. The value of an enthalpy change is quoted for standard conditions: gases at 1 atm, solutions at unit concentration and substances in their normal states at a specified temperature. (usually 273K or 0 0 C) </li></ul><ul><li>All elements in their standard state are assigned an enthalpy of formation of 0 </li></ul>
- 15. <ul><li>Standard enthalpy of reaction is the heat absorbed in a reaction at constant pressure between the number of moles of reactant shown in the equation for the reaction. </li></ul><ul><li>Standard enthalpy of combustion is the heat absorbed when 1 mole of a substance is completely burned in oxygen at constant pressure. </li></ul><ul><li>Standard enthalpy of solution is the heat absorbed when 1 mole of a substance is dissolved at constant pressure in a stated amount of solvent. This may be 100g or 100ml or an ‘infinite’ amount, i.e. a volume so large that on further dilution there is no further heat change. </li></ul>
- 16. H r Ө depends only on the difference between the standard enthalpy of the reactants and the standard enthalpy of the products, not on the route by which the reaction occurs. This is Hess’s Law – If a reaction can proceed by more than one route the overall enthalpy is the same regardless of which route is followed.
- 17. <ul><li>Find the enthalpy change for the reaction </li></ul><ul><li>CH 2 = CH 2(g + HCl (g) C 2 H 5 Cl (g) </li></ul><ul><li>Given the following data </li></ul><ul><li> H f Ө = standard enthalpy change of formation </li></ul><ul><li> H f Ө CH 2 CH 2 = +52.3 KJ mol -1 </li></ul><ul><li> H f Ө HCl = - 92.3 KJ mol -1 </li></ul><ul><li> H f Ө C 2 H 5 Cl = -105 KJ mol -1 </li></ul><ul><li> H = H f Ө products - H f Ө reactants </li></ul><ul><ul><li> H r Ө = -105 – (+52.3) + (-92.3) </li></ul></ul><ul><li>= -65KJ mol-1 </li></ul>
- 18. <ul><li>Note that the formula H = H f Ө products - H f Ө reactants </li></ul><ul><li>Applies when the data given is the enthalpy of formation. </li></ul><ul><li>If the data given is the enthalpy of combustion the following formula is used H = H c Ө reactants - H c Ө products </li></ul><ul><li>It doesn’t matter what the reaction is called . It may be called a combustion reaction but if the data given is the enthalpy of formation use the formula H = H f Ө products - H f Ө reactants ! </li></ul>
- 19. <ul><li>Calculate the standard enthalpy of reaction for the following </li></ul><ul><li>2C + 2H 2(g) + O 2 CH 3 CO 2 H (l) </li></ul><ul><li>Given the following enthalpies of combustion </li></ul><ul><li>C -394 KJ/mol </li></ul><ul><li>H 2(g) -286 KJ/mol </li></ul><ul><li>CH 3 CO 2 H (l) -876 KJ/mol </li></ul><ul><li> r = (2x -394) + (2x-286) – (-876) KJ/mol = -484 KJ/mol </li></ul>

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