1. Thermodynamics Principles of Chemical Reactivity

2. Basic Principles Thermodynamics: The science of heat and work Energy: the capacity to do work -chemical, mechanical, thermal, electrical, radiant, sound, nuclear -affects matter by raising its temperature, eventually causing a state change -All physical changes and chemical changes involve energy Potential Energy: energy that results from an object’s position -gravitational, chemical, electrostatic Kinetic Energy: energy of motion

3. Basic Principles Law of Energy Conservation: Energy can neither be created nor destroyed -a.k.a. The first law of thermodynamics -The total energy of the universe is constant Temperature vs. Heat: –Temperature is the measure of an object’s heat energy –Heat ≠ temperature

4. The Measurement of Heat Thermal Energy depends on temperature and the amount (mass or volume) of the object -More thermal energy a substances has the greater the motion its atoms/molecules have -Total thermal energy of an object is the sum of the individual energies of all atoms/molecules/ions that make up that object SI unit: Joule (J) 1 calorie = 4.184 J English unit = BTU

5. Converting Calories to Joules Convert 60.1 cal to joules

6. System: object or collection of objects being studied –In lab, the system is the chemicals inside the beaker Surroundings: everything outside of the system that can exchange energy with the system –The surroundings are outside the beaker Universe: system plus surroundings Exothermic: heat transferred from the system to the surroundings Endothermic: heat transferred from the surroundings to the system Basic Principles

7. Specific Heat Capacity (C) amount of heat required to raise the temperature of 1 gram of a substance by 1 degree Celsius SI Units: Specific heat capacity =J/g. °C Specific heat of water = 4.184 J/g. °C

8. Heat Transfer Heat transfer equation used to calculate amounts of heat (q) in a substance J g J/g·°C °C q 1 + q 2 + q 3 … = 0 or q system + q surroundings = 0

9. Heat Transfer Calculate the amount of heat to raise the temperature of 400 g of water from 10.0 o C to 100 o C

10. Heat Transfer Calculate the amount of heat energy (in joules) needed to raise the temperature of 12.50 g of water from 45.0°C to 79.0°C

11. Heat Transfer Specific heat of gold is 0.13 Therefore the metal cannot be pure gold. A 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold? J g. °C

12. Changes of State occurs when enough energy is put into a substance to over come molecular interactions Solid-liquid: molecules in a solid when heated move about vigorously enough to break solid-solid molecular interactions to become a liquid Liquid-gas: molecules in a liquid when heated move about more vigorously enough to break liquid-liquid molecular interactions to become a gas Note: This happens in reverse by removing heat energy

13. Energy and Changes of State Heat of fusion: heat needed to convert a substance from a solid to a liquid (at its melting/freezing point) 333 J/g for water Heat of vaporization: heat needed to convert a substance from a liquid to a gas (at its boiling/condensation point) 2256 J/g for water Example: Calculate the amount of heat involved to convert 100.0 g of ice at -50.0°C to steam at 200.0°C.

14. The First Law of Thermodynamics This law can be stated as, “The combined amount of energy in the universe is constant” Also called-The Law of Conservation of Energy: –Energy is neither created nor destroyed in chemical reactions and physical changes.

15. Changes in Internal Energy (  E)  E is negative when energy is released by a system -Energy can be written as a product of the process

16. Changes in Internal Energy (  E)  E is positive when energy is absorbed by a system undergoing a chemical or physical change –Energy can be written as a reactant of the process

17. Enthalpy Changes for Chemical Reactions Exothermic reactions: release energy in the form of heat to the surroundings (  H < 0) -heat is transferred from a system to the surroundings Endothermic reactions: gain energy in the form of heat from the surroundings (  H > 0) -heat is transferred from the surroundings to the system For example, the combustion of propane: Combustion of butane:

18. Enthalpy Changes for Chemical Reactions Exothermic reactions generate specific amounts of heat –Because the potential energies of the products are lower than the potential energies of the reactants Endothermic reactions consume specific amounts of heat –Potential energies of the reactants are lower than the products  H for the reverse reaction is equal, but has the opposite sign to the forward reaction

19. Thermochemical Equations balanced chemical reaction with the  H value for the reaction  H < 0 designates an exothermic reaction: heat is a product, the container feels hot  H > 0 designates an endothermic reaction: heat is a reactant, the container feels cold

20. Calorimetry An experimental technique that measures the heat transfer during a chemical or physical process Constant pressure calorimetry: A styrofoam coffee-cup calorimeter is used to measure the amount of heat produced (or absorbed) in a reaction –This is one method to measure q P (called  H) for reactions in solution q reaction + q solution = 0 Note: Assuming no heat transfer to the surroundings

21. Calorimetry If an exothermic reaction is performed in a calorimeter, the heat evolved by the reaction is determined from the temperature rise of the solution –This requires a two part calculation When we add 25.00 mL of 0.500 M NaOH at 23.000 o C to 25.00 mL of 0.600 M CH 3 COOH already in the calorimeter at the same temperature, the resulting temperature is observed to be 25.947 o C. Determine heat of reaction and then calculate the change in enthalpy (as KJ/mol) for the production of NaCH 3 COO. CH 3 COOH (aq) + NaOH (aq)  NaCH 3 COO (aq) + H 2 O (l)