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2. Definitions: recall physics 2  Energy (E):  The ability to do work; measured in Joules (J)  Work:  Amount of energy applied or transferred over a distance  Potential Energy (E p ):  Energy of object due to its position or composition; stored energy  Kinetic Energy (E k ):  Energy of object due to its motion; movement energy

3. Definitions: recall physics 3  Thermal Energy (H):  Total quantity of kinetic & potential energy within a substance, E p + E k = H  Heat (q):  The transfer of thermal energy from a warm body to cooler body  Temperature:  A measure of the average kinetic energy of particles in a substance/object

4. Definitions: new ones 4  Chemical System:  The group of reactants & products being studied  Surroundings:  The environment in which the chemical reaction takes place; substance that is not part of the reaction  Open System  Matter & Energy able to leave system  Closed System  Energy able to leave system  Isolated System  Neither matter nor energy able to leave system

5. System and Surroundings can exchange energy and matter with surroundings universe = system + surroundings the system is the sample being observed the surroundings is everything else interactions between a system and its surroundings involve exchange of energy and matter can exchange energy, not matter, with surroundings cannot exchange matter or energy with surroundings Three types of systems based on this type of exchange are:

6. Calorimetry Measuring energy changes in a reaction  We measure the ΔT of the surroundings and calculate the thermal energy, q, lost or gained by the surroundings -q surroundings = +q system q = mcΔT 6

7. ΔT = T final – T initial Measurable properties of a system include: volume, mass, pressure, temperature, and specific heat capacity. Calculating the Amount of Heat Entering and Leaving a System A calculated property of a system includes heat (Q) that enters or leaves an object. When heat enters a system, ΔT is positive and so is Q.

9. Specific Heat Capacity,(c); 9

10. A 1.0 g sample of copper is heated from 25.0°C to 31.0°C. How much heat did the sample absorb? L EARNING C HECK

11. The First Law of Thermodynamics: Energy is Conserved The first law of thermodynamics states that: Energy can be converted from one form to another but cannot be created or destroyed. Since any change in energy of the universe must be zero, ΔE universe = ΔE system + ΔE surroundings = 0 ΔE system = –ΔE surroundings if a system gains energy, that energy comes from the surroundings if a system loses energy, that energy enters the surroundings

12. A 400 g iron rod is heated in a bunsen flame and then plunged into an insulated beaker containing 1.00 L of water (1000 g). The original temperature of the water was 20ºC. After the iron rod and the water have reached the same temperature, it is 32.8ºC. (specific heat capacity of iron is 0.45 J/gºC, and water is 4.18J/gºC) What was the original temperature of the iron rod? U SING THE F IRST L AW OF T HERMODYNAMICS

13. Enthalpy One way chemists express thermochemical changes is by a variable called enthalpy, H. The change in enthalpy, ΔH, of a system can be measured. It depends only on the initial and final states of the system, and is represented by ΔH = ΔE + Δ(PV) For reactions of solids and liquids in solutions, Δ(PV) = 0 If heat enters a system ΔH is positive the process is endothermic If heat leaves a system ΔH is negative the process is exothermic

14. The Second Law of Thermodynamics When in thermal contact, energy from hot particles will transfer to cold particles until the energy is equally distributed and thermal equilibrium is reached. The second law of thermodynamics states that: When two objects are in thermal contact, heat is transferred from the object at a higher temperature to the object at the lower temperature until both objects are the same temperature (in thermal equilibrium)

15. Definitions: new ones 15  Enthalpy (H): another term for thermal energy  Enthalpy Change (ΔH): the change in thermal energy from reactants to products; this equals q under constant pressure: ΔH system = q system ΔH system = H products – H reactants Where q surroundings is the measurable quantity  Endothermic:  intake of energy, -q surroundings & +q system & + ΔH  Exothermic :  release of energy, +q surroundings & -q system & - ΔH

16. Comparing Categories of Enthalpy Changes: Enthalpy of Solution The orange arrow shows the overall ΔH. Three processes occur when a substance dissolves, each with a ΔH value. 1.bonds between solute molecules or ions break 2.bonds between solvent molecules break 3.bonds between solvent molecules and solute molecules or ions form Sum of the enthalpy changes: enthalpy of solution, ΔH solution

17. The ΔH for one phase change is the negative of the ΔH for the opposite phase change. Heat must be added to or removed from a substance in order for the phase of the substance to change. The ΔH for each phase change has a particular symbol. For example, ΔH melt is called the enthalpy of melting.

18. Thermochemical Equations and Calorimetry Chemical reactions involve initial breaking of chemical bonds (endothermic) then formation of new bonds (exothermic) ΔH r is the difference between the total energy required to break bonds and the total energy released when bonds form.

19. Thermochemical Equations Thermochemical equations include the enthalpy change. The enthalpy term can also be written beside the equation. Exothermic reaction: CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O(g) + 890.8 kJ Endothermic reaction: N 2 (g) + 2O 2 (g) + 66.4 kJ → 2NO 2 (g) CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O(g) ΔH r = –890.8 kJ N 2 (g) + 2O 2 (g) → 2NO 2 (g) ΔH r = +66.4 kJ

20. Enthalpy Diagrams Enthalpy diagrams clearly show the relative enthalpies of reactants and products. For exothermic reactions, reactants have a larger enthalpy than products and are drawn above the products. For endothermic reactions, the products have a larger enthalpy and are drawn above the reactants.

21. Molar Enthalpy Change ΔH n : 21  Molar Enthalpy Change (ΔH n ): enthalpy change associated with any change (physical, chemical or nuclear) in one mole of a substance; J/mol or kJ/mol ΔH rxn = nΔH n

22. Common Molar Enthalpies 22

23. Example:  Ethanol is used to disinfect skin before receiving a flu shot. How does this feel on your skin? The enthalpy of vaporization (ΔH vap )of ethanol is 38.6 kJ/mol. What is the enthalpy change (ΔH rxn ) of this reaction if 1.0g of ethanol is rubbed on your skin? 23

24. Determine ΔH comb for 15.0 g of propane. L EARNING C HECK The molar mass of propane: 44.1 g/mol Therefore, 15.0 g is 0.340 mol ΔH comb = ΔH° comb ΔH comb = (0.340 mol) (–2219.2 kJ/mol) ΔH comb = 755 kJ/mol The symbol for a standard enthalpy change is ΔH°, read as "delta H standard" or, perhaps more commonly, as "delta H nought“. The standard enthalpy change of a reaction is the enthalpy change which occurs when equation quantities of materials react under standard conditions, and with everything in its standard state (SATP (25° C and 100 kPa))

25. Expressing Molar Enthalpy Changes: 1. As part of chemical equation: 2. Associated ΔH value: 3. Potential Energy Diagram: 25

26. 3) Potential Energy Diagram: 26