TOPIC 5 ENERGETICS/THERMOCHEMISTRY

1. TOPIC 5ENERGETICS/THERMOCHEMISTRY 5.1 MEASURING ENERGY CHANGES

2. ESSENTIAL IDEA The enthalpy changes from chemical reactions can be calculated from their effect on the temperature of their surroundings. NATURE OF SCIENCE (2.6) Fundamental principal – conservation of energy is a fundamental principle of science. NATURE OF SCIENCE (3.1) Making careful observations – measurable energy transfers between systems and surroundings.

3. INTERNATIONAL-MINDEDNESS The SI unit of temperature is the Kelvin (K), but the Celsius scale (◦C), which has the same incremental scaling, is commonly used in most countries. The exception if the USA which continues to use the Fahrenheit scale (◦F) for all non-scientific communication.

4. THEORY OF KNOWLEDGE What criteria do we use in judging discrepancies between experimental and theoretical values? Which ways of knowing do we use when assessing experimental limitations and theoretical assumptions?

5. UNDERSTANDING/KEY IDEA 5.1.A Heat is a form of energy.

6. ENERGY • All chemical reactions are accompanied by energy changes. • Energy is a measure of the ability to do work, that is to move an object against an opposing force. • Examples: heat, light, sound, electricity and chemical energy which is the energy released or absorbed during chemical reactions.

7. HEAT • Heat is a form of energy which is transferred as a result of a difference in temperature and produces an increase in disorder of the behavior of particles. • Heat increases the average kinetic energy of the molecules in a disordered fashion.

8. Needed definitions • Enthalpy • Heat content of a substance (at constant pressure) • Enthalpy is also the internal energy stored in the reactants. • The absolute value for the enthalpy of reactants and products cannot be known, but the what can be measured is the difference between the two.

9. Needed definitions • System – area of interest (Example: beaker and its contents) • Surroundings – everything else in the universe

10. An “open” system can exchange matter and energy with the surroundings. • A “closed” system can only exchange energy, not matter, with the surroundings.

11. The joule (J) is the SI unit for energy and work.

12. UNDERSTANDING/KEY IDEA 5.1.B Temperature is a measure of the average kinetic energy of the particles.

13. Temperature • Def – measure of the average kinetic energy of the particles • Temperature increase depends upon: • Mass of the object • Amount of heat added • Nature of the substance • Different substances need different amounts of heat to increase the temp of a unit mass by 1K or 1ºC.

14. UNDERSTANDING/KEY IDEA 5.1.C Total energy is conserved in chemical reactions.

15. UNDERSTANDING/KEY IDEA 5.1.D Chemical reactions that involve transfer of heat between the system and the surroundings are described as endothermic or exothermic.

16. EXOTHERMIC REACTIONS A reaction which results in a transfer of energy from the system to the surroundings. • Heat is given off or produced. • Products have less energy or heat content than the reactants. • ΔH is negative. • The bonds in the products are stronger than the bonds in the reactants.

17. Exothermic reaction

18. Exothermic reactionEnthalpy diagram reactants enthalpy H ΔH = negative products extent of reaction

19. ENDOTHERMIC REACTIONS • A reaction which results in a transfer of energy from the surroundings to the system. • Heat is absorbed. • Reactants have less energy than the products. • ΔH is positive. • The bonds in the reactants are stronger than those in the products.

20. Endothermic reaction

21. Endothermic reactionEnthalpy diagram products enthalpy H ΔH = positive reactants extent of reaction

22. UNDERSTANDING/KEY IDEA 5.1.E The enthalpy change (ΔH) for chemical reactions is indicated in kJ mol-1

23. UNDERSTANDING/KEY IDEA 5.1.F ΔH values are usually expressed under standard conditions, given by ΔH◦, including the standard states.

24. GUIDANCE Standard state refers to the normal, most pure stable state of a substance measured at 100 kPa. Temperature is not part of the definition of standard state, but 298K is commonly given as the temperature of interest.

25. GUIDANCE Be familiar with the enthalpy change of combustion (ΔHc ◦) and the enthalpy change of formation (ΔHf ◦).

26. APPLICATION/SKILLS Be able to calculate the heat change when the temperature of a pure substance is changed using q = mcΔT.

27. GUIDANCE The specific heat capacity of water is provided in the data booklet in section 2.

28. APPLICATION/SKILLS Be able to evaluate calorimetry experiments for the enthalpy of a reaction.

29. GUIDANCE Be familiar with reactions occurring in aqueous solutions and combustion reactions.

30. GUIDANCE Students can assume the specific heat and density of aqueous solutions are equal to those of water, but be aware of this limitation.

31. GUIDANCE Heat losses to the environment and the heat capacity of the calorimeter in experiments should be considered, but the use of a bomb calorimeter is not required.

32. THERMOCHEMICALEQUATION EXAMPLES Combustion of Methane CH4(g) + 2O2(g) CO2(g) + 2H2O(l)ΔH = -890 kJ mol-1 Photosynthesis 6CO2(g) + 6H2O(l) C6H12O6(aq) + 6O2(g)ΔH = +2802.5 kJ mol-1 Thermite Reaction 2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe(s)ΔH = -841 kJ mol-1

33. You must give the “state” symbols such as (s), (g), (l), (aq) in thermochemical equations because energy changes depend upon the state of the reactants and the products.

34. Standard enthalpy change of reaction ΔHº • This is the heat or enthalpy change of a reaction when carried out at standard conditions. • Temperature = 298K or 25ºC • Pressure = 101.3 kPa (1atm) • Solution concentration = 1mol dm-3 (1M) • All substances are in their standard states (how they are found in nature)

35. The actual amount of heat absorbed or produced in a chemical reaction depends upon several factors: • Nature of the reactants and products • The amount or concentration of the reactants. (The greater the amount that reacts, the greater the heat change.) • The states of the reactants and products. Changing states involves an enthalpy change so this will affect the total heat change. • The temperature of the reaction.

36. All combustion reactions are exothermic processes. • All neutralization reactions are exothermic processes.

37. Apply the relationship between temperature change, enthalpy change and be able to classify the reaction as endothermic or exothermic. • This will be determined by calorimetry.

38. Specific Heat Capacity • The following relationship allows the heat change in a material to be calculated from the temperature change. q = mcΔT heat = mass x specific heat x ΔT

40. Specific heat capacity (c) is the heat needed to increase the temperature of a unit mass (usually 1g) by 1K or 1ºC. • Specific heat of water c = 4.18 J K-1g-1

41. Be able to deduce from an enthalpy level diagram, the relative stabilities of reactants and products and the sign of the enthalpy change for the reaction.

42. There is a natural direction for change. • The direction of change is in the direction of lower stored energy. • We generally expect a reaction to occur if ΔH is negative (exothermic), but some endothermic reactions do occur. These occur when the entropy (S) of the system is large.

43. Enthalpy Diagram for Hydrogen Peroxide (H2O2) H2(g) + O2(g) ΔH1 H2O2(l) ΔH3 ΔH2 Enthalpy H2O(l) + 1/2 O2(g) In this diagram, hydrogen peroxide in the middle is stablecompared to H2 and O2, but unstable compared to the decomposition of water and oxygen on the bottom line.

44. The heat produced when one mole of a substance is burned in excess oxygen is called the enthalpy of combustion. • ΔHc is exothermic and always negative.

45. Calculate the heat energy change when the temperature of a pure substance is changed.

46. How to calculate heat changes from temp changes. • When heat released by an exothermic reaction is absorbed by water, the temperature of the water increases. • The heat produced by the reaction can be calculated if it is assumed that all the heat is absorbed by the water. heat change of rxn = -heat change of water = - mH2O x cH2O x ΔTH2O As the water has gained the heat produced by the reaction, the heat change of the reaction is negative when the temperature of the water increases.

47. Example Problem 1 • Calculate the enthalpy of combustion of ethanol from the following data. Assume all heat from the reaction is absorbed by the water. Compare your value with the IB Data booklet value and suggest reasons for any differences.

48. Mass of water in Cu Calorimeter = 200.00g • Temperature increase in water = 13.00ºC • Mass of ethanol burned = 0.45g heat change of rxn = - mH2O x cH2O x ΔTH2O = -(200.00g)(4.18Jg-1ºC-1)(13.00ºC) = -10868 J The heat calculated above is the heat gained by the water which is also the heat lost by the combustion of the ethanol. Since ΔHc is per mole of substance combusted, you have to find moles of ethanol burned and divide the heat it lost by the moles. To find moles, divide mass by molar mass: 0.45g / 46.08 g mol-1 = .0098mol ΔHc = -10868J / .0098mol = -1112883 J/mol = -1112.883 kJ/mol Using sig figs, the answer is limited by the mass = -1100 kJ/mol

49. The IB Data booklet value is -1367 kJ/mol. • What are some reasons for the difference in values? • Not all the heat produced by the combustion is transferred to the water. • Some is needed to heat the Cu calorimeter and some has passed to the surroundings. • The combustion of ethanol is unlikely to be complete due to the limited oxygen available.

50. Heat of Combustion Set up