1. Topics 5 and 15

2. Hess’s Law Calorimetry Enthalpy Enthalpy of Formation Bond Energy

3. Assessment Statements 5.1.1 Define the terms exothermic, endothermic and Standard Enthalpy Change (∆H). 5.1.2 Understand that combustion and neutralisation are exothermic processes. 5.1.3 Apply the relationship between temperature change, enthalpy change and the classification of a reaction as endothermic or exothermic. 5.1.4 Deduce, from an enthalpy level diagram, the relative stabilities of reactants and products and the sign of the enthalpy change for the reaction.

4. Energetics Relationship between chemical reactions and heat What causes chemical reactions to occur The concepts of Enthalpy, Entropy, and Free Energy

5. Two Trends in Nature Order  Disorder  High energy  Low energy 

6. System v. Surroundings System – The actual chemical reaction taking place Surroundings – The rest of the universe outside of the system

7. Heat Flows Heat always flows from hot to cold Heat flows (not cold) Cold is an absence of heat

8. Assessment Statemen t 5.1.1 Define the terms exothermic, endothermic and Standard Enthalpy Change (∆H). 5.1.2 Understand that combustion and neutralisation are exothermic processes.

9. Exothermic Reactions Examples include: Burning reactions including the combustion of fuels. Detonation of explosives. Reaction of acids with metals. Thermit reaction Magnesium reacting with acid Exothermic reactions increase in temperature.

10. Say whether these processes are exothermic. 1. Charcoal burning 2. A candle burning. 3. A kettle boiling 4. Ice melting 5. A firework exploding yes no in You have to put heat in for boiling and melting. out You get heat out from all the other processes Activity

11. Magnesium + Hydrochloric acid Gets hot 25 o C 45 o C magnesium Hydrochloric acid Heat energy given out Exothermic Reactions

12. If heat is given out this energy must have come from chemical energy in the starting materials (reactants). Reactants convert chemical energy to heat energy. The temperature rises. 25 o C 45 o C Exothermic Reactions

13. 45 o C Almost immediately the hot reaction products start to lose heat to the surroundings and eventually they return to room temperature. 25 o C Chemical energy becomes heat energy. The reaction mixture gets hotter. Eventually this heat is lost to the surroundings. It follows that reaction products have less chemical energy than the reactants had to start with. Exothermic Reactions

14. Energy / kJ) Progress of reaction (time) Energy Level Diagram for an Exothermic Reaction reactants Reactants have more chemical energy. Some of this is lost as heat which spreads out into the room. products Products now have less chemical energy than reactants.

15. Energy / kJ Progress of reaction reactants products  H=negative Energy Level Diagram for an Exothermic Reaction 2.  H is how much energy is given out  H is negative because the products have less energy than the reactants. Energy Level Diagram for an Exothermic Reaction

16. Exothermic reactions give out energy. There is a temperature rise and  H is negative. Exothermic Reaction - Definition products Energy / kJ) Progress of reaction reactants  H is negative

17. Activity

18. Exothermic reactions

19. Endothermic Reactions Endothermic chemical reactions are relatively rare. A few reactions that give off gases are highly endothermic - get very cold. Dissolving salts in water is another process that is often endothermic. Endothermic reactions cause a decrease in temperature.

20. Cools Heat energy taken in as the mixture returns back to room temp. Starts 25°C Cools to 5°C Ammonium nitrate Water Endothermic reactions cause a decrease in temperature. Returns to 25°C Endothermic Reactions

21. Extra energy is needed in order for endothermic reactions to occur. This comes from the thermal energy of the reaction mixture which consequently gets colder. Reactants convert heat energy into chemical energy as they change into products. The temperature drops. 25 o C 5 o C Endothermic Reactions

22. 25 o C The cold reaction products start to gain heat from the surroundings and eventually return to room temperature. 5 o C The reactants gain energy. 25 o C This comes from the substances used in the reaction and the reaction gets cold. Eventually heat is absorbed from the surroundings and the mixture returns to room temperature. Overall the chemicals have gained energy. Endothermic Reactions

23. products Energy / kJ) Progress of reaction reactants  H=+ Energy Level Diagram for an Endothermic Process This is positive because the products have more energy than the reactants. This is how much energy is taken in

24. Endothermic reactions take in energy. There is a temperature drop and  H is positive. Endothermic Reaction Definition  H=+ products Energy / kJ Progress of reaction reactants

26. Are these endothermic or exothermic? 1. A red glow spread throughout the mixture and the temperature rose. 2. The mixture bubbled vigorously but the temperature dropped 15 0 C. 3. Hydrazine and hydrogen peroxide react so explosively and powerfully that they are used to power rockets into space. 4. The decaying grass in the compost maker was considerably above the outside temperature. exo endo exo Activity

27. Sketch the two energy diagrams and label exothermic and endothermic as appropriate.  H=+ products Energy / kJ Progress of reaction reactants products Energy / kJ) Progress of reaction reactants  H=- Activity

28. Endothermic reactions

29. Energy Changes in Endothermic and Exothermic Processes In an endothermic reaction there is more energy required to break bonds than is released when bonds are formed. The opposite is true in an exothermic reaction. 29

30. Exothermic and endothermic reactions

31. Exothermic reaction one that gives off heat – transfers thermal energy from the system to the surroundings. 2H 2 (g) + O 2 (g) 2H 2 O (l) + energy H 2 O (g) H 2 O (l) + energy 6.2 Reactants had more energy than the products

32. Endothermic reaction one in which heat has to be supplied to the system from the surroundings. energy + 2HgO (s) 2Hg (l) + O 2 (g) energy + H 2 O (s) H 2 O (l) Products end with more potential energy the reactants

33. Types of enthalpy change

34. Standard enthalpies: examples The standard enthalpy of formation of methane can be represented by: C (s, graphite) + 2H 2(g)  CH 4(g) By definition, the standard enthalpy of formation of an element, in its standard state, must be zero. ∆H f ө = -74.9 kJ mol -1 The standard enthalpy of combustion of methane can be represented by: CH 4(g) + 2O 2(g)  CO 2(g) + 2H 2 O (l) ∆H c ө = -890 kJ mol -1

35. Enthalpy change summary

36. Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.  H = H (products) – H (reactants)  H = heat given off or absorbed during a reaction at constant pressure H products < H reactants  H < 0 H products > H reactants  H > 0 6.4

37. Thermochemical Equations H 2 O (s) H 2 O (l)  H = 6.01 kJ Is  H negative or positive? System absorbs heat Endothermic  H > 0 6.01 kJ are absorbed for every 1 mole of ice that melts at 0 0 C and 1 atm. 6.4

38. Thermochemical Equations CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O (l)  H = -890.4 kJ Is  H negative or positive? System gives off heat Exothermic  H < 0 890.4 kJ are released for every 1 mole of methane that is combusted at 25 0 C and 1 atm. 6.4

39. Heat v. Temperature Heat – form of energy measured in J (SI) Temperature – average kinetic energy of molecules in a system – Measured in K (SI) 0 o C= 273 K