Topic 5: Energetics

1. Topic 5: Energetics 5.1 Exothermic and endothermic reactions 5.2 Calculation of enthalpy changes

2. Thermochemistry (Energetics) • The study of energy involved during chemical reactions • Heat: • the energy of motion of molecules • All matter has moving particles at stp • Temperature: • transfer of heat to a substance because of faster molecular movement (as long as there is no phase change)

3. A temperature change is explained as a change in kinetic energy (movement) • Temperature depends on the quantity of heat (q) flowing out or in of the substance. • Energy flowing in the system = Endothermic • Has a positive value • Energy entering (feels cool) • Energy flowing out of the system = Exothermic • Has a negative value • Energy exiting (feels warm)

4. Heat (q) q=mc ∆t • q=heat • m=mass • ∆t=change in temperature (tf-ti) • c=specific heat capacity (J (g oC)-1) Specific heat capacity is the quantity of heat required to raise the temperature of a unit mass of a substance by one degree Celsius.

5. Law of conservation of energy • ∆E universe = O • The total energy of the universe is constant, it is not created or destroyed, however it can be transferred from one substance to another. • ∆E universe = ∆E system + ∆E surroundings

6. First Law of thermodynamics • Any change in energy of a system is equivalent by an opposite change in energy of the surroundings. • ∆E system = - ∆E surroundings • According to this law, any energy released or absorbed by a system will have a transfer of heat, q. • So, q system = - q surroundings

7. Sample Problem • 15 g of ice was added to 60.0 g of water. The Ti of water was 26.5 oC, the final temperature of the mixture was 9.7 oC. How much heat was lost by the water? • q=mc ∆t q=(60.0 g) (4.18 J/g oC) (9.7-26.5 oC) q= - 4213.44 J [-4.2 kJ]

8. Watch this flash video about heat flow http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/enm1s3_4.swf

9. Enthalpy (∆H) • Total kinetic and potential energy of a system under constant pressure. • The internal energy of a reactant or product cannot be measured, but their change in enthalpy (heat of reaction) can. ∆ H = Hproducts – Hreactants • A change in enthalpy occurs during phase changes, chemical reactions and nuclear reactions. ∆ H system = q surroundings

10. Endothermic Reactions Method 1: enthalpy level diagram

11. Endothermic Reactions Method 2: enthalpy term outside of the equation 2 HgO (s) 2 Hg (l) + O2(g)ΔH=181.67 kJ Method 3: enthalpy term within the equation 2 HgO (s) + 181.67 kJ 2 Hg (l) + O2(g)

12. Exothermic ReactionsMethod 1: Enthalpy level diagram

13. Exothermic Reactions Method 2: 4Al(s)+ 3O2(g) 2 Al2O3(g)ΔH=-1675.7 kJ Method 3: 4 Al(s) + 3O2(g) 2 Al2O3(g) +1675.7 kJ • Neutralization and combustion reactions

14. Calorimeters (qwater = -qsystem) • Used to measure the amount of energy involved in a chemical reaction. • To be treated like • Isolate/closed system • Specific mass of water used. • Energy flows to or from the water in the cups • Measure the temperature change related to the water.

15. Problem • An 25.6 g of an unknown metal with an initial temperature of 300 oC, is placed in 150.0 g of water with an initial temperature of 35.0 oC. If the water’s temperature stabilizes at 55.0 oC, calculate the specific heat capacity of this metal.

16. Bomb Calorimeter

17. Heat Capacity • Related to bomb calorimeters • Unit is (J/oC) because its always with a set mass, so it is redundant to repeat the term over.

18. 15.1 Standard enthalpy changes of reaction Higher level 15.1.1 Define and apply the terms standard state, standard enthalpy change of formation, and stand enthalpy change of combustion 15.1.2 Determine the enthalpy change of a reaction using standard enthalpy changes of formation and combustion.

19. Standard Molar Enthalpy of Formation • Standard implies the states of the particle at 1 atm and at 0oC. • Quantity of energy released (-) or absorbed (+) when one mole of a compound is formed directly from its elements at standard temperature and pressure. • We use a table to find them. • Unit for ΔHof: kJ/mol • Watch your states!

21. Practice: • What is the standard molar enthalpy of formation for the following reaction? • 2 Na (s) + Cl2(g) 2 NaCl (s) + 814 kJ By definition, the standard molar enthalpy of formation is for ONE mole of product formed. ∆Hfo = -407 kJ/mol for NaCl

22. Standard Molar Enthalpy of Combustion (∆Hcombo • Energy changes involved with combustion reactions of one mole of a substance. • Remember that these reactions are only measured once cooled to 25oC • Combustion is a reaction with oxygen as a reactant (burning) • Will need a table of values to use.

23. Combustion reaction with alkanes: • Always form water and carbon dioxide. • Ex: CH4 + 2O2 CO2 + 2H2O • Remember your alkanes: CnH2n+2 • Meth: C=1, Eth: C=2, Prop: C=3, But: C=4, Pent: C=5, Hex: C=6, Hept: C=7, Oct: C=8, Non: C=9 and Dec: C=10.

24. Standard heats of reactions (∆Hrxno) • Measured from all products and reactants at their standard states. (if a solution, concentration = 1M) • All elements at standard state ΔHof = 0 • Most compounds have a negative ΔHof • Use balanced equation, where n= number of moles

25. Calculating enthalpy changes • Amount of a substance reacting matters, so can use q= nΔH. • Remember n=amount of moles. • If you are given a mass (g) and molar mass (g/mol), then you can solve for n by dividing mass by molar mass. (review from Topic 1 stoichiometry section)

26. Practice: • Calculate the ΔHorxn for: • 4NH3 (g) + 5O2 (g)  4NO (g) + 6H2O (g) • CO (g) + H2O (g)  CO2 (g) + H2 (g) • Calculate the ΔHocomb for: • 2CH3OH(l) + 3O2(g)  2 CO2(g) + 4H2O(l) • 2C2H6 (g) + 7 O2 (g)  4CO2 (g) + 6 H2O (l)

27. Practice: 1/8 S8 (s) + H2 (g)  H2S (g) ∆Hrxno= -20.2 kJ • Is this an endo or exothermic reaction? • What is the ∆Hrxno for the reverse reaction? • What is the ∆Hwhen 2.6 mol of S8 reacts? • What is the ∆H when 25.0 g of S8 reacts?

28. Don’t Forget to… • Read your text book • Look over the questions assigned • Use your course companion • Use your study guide