Unit 1 – Thermochemistry

1. General Outcomes: • Determine and interpret energy changes in chemical reactions • Explain and communicate energy changes in chemical reactions Unit 1 – Thermochemistry

2. Key Concepts • Enthalpy of formation • Enthalpy of reaction • ΔH notation • Hess’ Law • Molar enthalpy • Energy diagrams • Activation Energy • Catalysts • Calorimetry • Fuels and Energy Efficiency

3. Energy • the potential to do work. • examples of energy changes: - chemical energy in food to motion - solar energy to power - geothermal energy to heat - nuclear energy to power, heat, motion

4. Thermodynamics: study of energy changes in systems • Thermochemistry–study of energy changes involved in physical and chemical changes.

5. Laws of Thermodynamics First Law - Energy can be neither created nor destroyed, it simply changes forms. Esystem= -Esurroundings Second Law – In any energy conversion, energy is always lost in the form of heat.

6. How do we know energy changes exist? • We can observe the effect it has on our surroundings. • Energy systems either release energy to the surroundings or absorb energy from the surroundings. • Universe = System + Surroundings

8. Three types of systems: • Open systems – matter and energy move freely in / out. • Closed systems – only energy moves in / out. • Isolated system – neither matter nor energy moves in / out.

10. Q=mcΔTYouTube - Chemistry Music Video: EmCee Delta Tee (The Calorimetry Song) Calculating Thermal Energy

11. Q =quantity of heat flowing in / out of that substance, measured as energy in units of joules (J) or kilojoules (kJ) • m-mass of substance undergoing the change measured in units of grams (g) or kilograms (kg) • c=heat capacity -heat required to change the temp of 1 g of substance by 1oC

12. Example Calculations (Q=mcΔT) • Many hot water heaters use the combustion of natural gas to heat the water in the tank. When 150 L of water at 10oC is heated to 65oC, how much energy flows into the water?

13. A solid substance has a mass of 250.0 g. It is cooled by 25.0oC, and loses 4.937 kJ of heat. What is the specific heat capacity and the identity of the substance?

14. If a copper pot weighing 0.500 kg is heated with 2.500 kJ of energy, what is the expected temperature change?

15. Energy From the Sun • Energy from the sun is stored in the bonds of chemical substances such as fossil fuels and hydrocarbons.

16. Solar energy is captured by plants during photosynthesis. • Living tissues from plants and animals become buried. Instead of decaying, these compounds form hydrocarbons or fossil fuels. • The energy from fossil fuels is released during combustion reactions and converted into heat.

17. Photosynthesis is endothermic: 6CO2(g) + 6H2O(l) + energy  C6H12O6(s) + 6O2(g) • Cellular Respiration is exothermic: C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l) + energy • Combustion is exothermic: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) + energy

19. Energy Types • There are two types of energy that matter may possess: 1. Kinetic energy – the energy of motion 2. Potential energy – the energy of position

20. Kinetic Energy (Ek) • faster the particle is moving, the greater the kinetic energy. • Can be found in three types: • Translational • Rotational • Vibrational

21. Ekmeasured indirectly by measuring the temperature of a system. • is proportional to the amount of heat a system holds.

22. Potential Energy (Ep) • Stored energy found in the bonds and inter-molecular forces of molecules • Changes in potential energy include phase changes, chemical reactions and nuclear reactions

23. Chemical Changes • Are changes in potential energy • energy is required to break bonds and the energy is released when new bonds are formed. • chemical reactions:micro/macro

24. Enthalpy and Enthalpy Change • enthalpy describes the total amount of energy a system has. • Enthalpy Change (ΔH = kJ) refers to how much the total energy of a system increases or decreases during a chemical reaction. • Molar enthalpy (ΔH = kJ/mol) describes the amount of enthalpy change per mole of substance.

25. Two types of reaction enthalpies • Endothermic reaction – less energy is released in forming bonds than was required to break bonds. • (+ΔH) • Surrounding temperature decreases • Exothermic reaction – more energy is released in making bonds than is required to break them. • (-ΔH) • Surrounding temperature increases

27. Calculating enthalpy changesH = n H H = change in enthalpy (kJ) H = molar enthalpy kJ/mol n = moles n=m/M

28. Example Calculations ΔH = nH • Hydrogen and oxygen gas react to produce 0.534 g of water vapor. If the molar enthalpy of reaction is 241.8 kJ/mol for water, what is the enthalpy change?

29. The molar enthalpy of combustion for ethanol is -725.9 kJ/mol. What mass of methanol must be burned to generate 2.34 x 104 kJ of energy?

30. What is the molar enthalpy of combustion if a 10.0 g sample of pure acetic acid is burned in oxygen and produces 144.7 kJ of energy?

31. Expressing Enthalpy Change 1. Molar enthalpy of a specific reaction relative to one species in the reaction. 2H2(g) + O2(g) 2H2O(g)H: -241.8 kJ/mol H2O 2. Enthalpy change for a balanced reaction equation. 2H2(g) + O2(g) 2H2O(g) ∆H: -483.6 kJ

32. 3. Including the energy value as part of the balanced equation. 2H2(g)+ O2(g)  2H2O(g)+ 483.6 kJ 2C(g) + 2H2(g) + 52.4 kJ  C2H4(g) 4. Drawing a potential energy diagram for the reaction.

34. Calorimetry • used to determine energy changes during physical or chemical processes • calorimeter : ideally an isolated system Energy in Chemical Reactions - Learning Activit - Flash Player Installation

35. All components of the calorimeter are included in the total energy change. • Mathematically, calorimetry may be described as: Heat lost = heat gained (* remember heat lost is expressed as a negative value)

36. Simple Calorimeter • used for reactions in aqueous solution • can calculate the heat change in the water and use to determine enthalpy change • P 353 lists assumptions made when using a simple calorimeter

38. Example Questions • 50.00mL of 0.300mol/L CuSO4(aq) is mixed with 50.00mL of NaOH(aq). The initial temperature of both solutions is 21.40oC and the highest temperature recorded after mixing is 24.60oC. • Determine the enthalpy change for the system • Determine the molar enthalpy change for CuSO4 • Write the thermochemical equation.

39. Barium chloride and sodium sulfate react to produce a precipitate of barium sulfate. When 0.100 mol of precipitate is formed the temperature of 3.00 kg of water is raised from 20.123 to 20.316oC. Calculate the molar enthalpy of reaction for the production of barium sulphate.

40. Bomb Calorimeter • used for combustion reactions • total energy change of the calorimeter includes changes in the container itself, the water, the thermometer and the stirrer.

41. bomb calorimeters may report a heat capacity for the entire calorimeter system. (C measured in J/oC) • Q = CT

42. Example Calculations • A bomb calorimeter has a heat capacity of 40.00kJ/oC. Complete combustion of 1.00g of hydrogen causes the temperature of the calorimeter to increase by 3.54oC. What is the molar enthalpy of combustion from this evidence?

43. When 1.00 g of propane is burned, about 2.36 J of heat is given off. What mass of water at 50.0oC can be heated to 80oC when 4.00 mol of propane is burned?

44. A strip of magnesium metal having a mass of 1.22 g is placed in 100 mL of 1.00 mol/L HCl(aq) in a metal can with a heat capacity of 562 J/oC. After the reaction, the temperature increased from 23.0oC to 45.5oC. Assuming the solution in the calorimeter has the same heat capacity as water, determine the heat of reaction per mole of magnesium.