1. Energy Relationships in Chemical Reactions

2. The nature of Energy and Types of Energy Energy – The capacity to do work Chemists define work as directed energy change resulting from a process Radiant Energy – energy produced from the sun Thermal Energy – the energy associated with the random motion of atoms and molecules Increase in thermal energy will increase the temperature but temperature does not define the amount of thermal energy For example a cup of coffee at 70 degree verus a bathtub of water at 40 degrees. The bathtub will have more thermal energy due to the number of molecules

3. Types of Energy Chemical Energy – the stored energy within the structural units of the substance Kinetic Energy – energy produce by a moving object Potential Energy – the amount on energy an object possibly has due to the composition or position

4. Conservation of Energy It is deducted that nothing can be created or destroyed by the law of conservation of mass Therefore energy can be transfer into a different type of energy THE TOTAL QUANTITY OF ENERGY IN THE UNIVERSE IS ASSUMED CONSTANT

5. Energy Changes in Chemical Reactions Almost all chemical reactions absorb or produce (release) energy Heat – the transfer of thermal energy between two bodies at different temperatures Welcome to Thermochemistry

6. Thermochemistry To analyze the energy changes associated with chemical reactions we must define the system in relationship to the rest of the universe Systems – open, closed or isolated

7. Systems Open – can exchange mass and energy with its surroundings (usually in the form of Heat) Closed – transfer of energy but not mass Isolated – does not allow transfer of mass or energy

9. Introduction to Thermodynamics The study of interconversion of heat and other kinds of energy Comparing the State of the System V = V f - V i

10. First Law of Thermodynamics Energy can be converted from one form to another but cannot be created or destroyed The change of internal energy in denoted by U U = U f - U i

11. Example of 1 st Law Consider the reaction S (s) + O 2(g) SO 2(g) U = U (products) – U (reactants) = energy content of 1 moles of SO 2(g) – energy content of [1mole S (s) + 1mole O 2(g) ] The energy of the products is less than that of the reactants and U is negative Therefore the reaction gives off heat

12. Interpreting the release of heat The release of heat is a transfer of energy to the surrounding universe Based on the 1 st law the total energy of the universe is not changed Therefore the total sum of the energy changes must be zero U sys + U surr = 0

13. Common Chemical Equation for 1 st Law U = q + w U of the system is the sum of heat exchange q between the system and the work done on the system w. Positive q – endothermic (heat absorbed by the system) Negative q - exothermic (heat released into the surroundings) Positive w – work done by the system on the surroundings Negative w – work done on the system by the surroundings

14. Work equation W = -P V Example A certain gas expands in volume from 2L to 6L at constant temperature. Calculate the work done by the gas if it expands against a pressure of 1.2 atm

15. Converting to Joules In the equation for work the units will be L*atm To convert to Joules Given in L * atm 101.3J 1L * atm

16. Work problem A gas expands from 264ml to 971ml at a constant temperature. Calculate the work done in joules by the gas if it expands against a pressure of 4atm

17. Heat (q) Energy can be gained by adding it directly to the system Bunsen burner Stirring

18. Heat Problem The work done when a gas is compressed in a cylinder is 384J. During this process there is a 152J heat transfer to the surroundings. Calculate the energy change for this system U = q + w Hint: must figure out signs of q and w

19. Enthalpy of Chemical Reactions Enthalpy of a reaction or energy change of a reaction H, is the amount of energy or heat absorbed in a reaction H is calculated in a pressure constant system H = H(products) – H(reactants)

20. H 2 O (l) H 2 O (s) Heat absorbed by the system from the surroundings H = 6.01 kJ/mol Heat given off by the system from the surroundings H = -890.4 kJ/mol CH 4 + 2O 2 CO 2 + 2H 2 O

21. Thermochemical Relationship between U and H U = H - P V Consider the reaction between carbon monoxide and oxygen (write out on your paper) U = H - n(RT)

22. Calorimetry Measurement of heat changes Specific Heat (s) – the amount of heat required to raise the temperature of one gram of the substance by one degree Celcius Heat Capacity (C) – the amount of heat required to raise the temperature of a given quantity of the substance by one degree Celcius

23. C=ms What is the heat capacity of 60g of water is the specific heat is 4.184J/g*C

24. Calculating the amount of heat lost or gained (q) q = ms T or q = C T Reminder – Pos (q) = Endothermic Neg (q) = Exothermic

25. Constant Volume Calorimetry q sym = q cal + q rxn = O q cal = C cal T Pg.192

26. Example: It is known that 1g of benzoic acid releases 26.42kJ of heat. If the temp rises 4.673 then what is the heat capacity of the calorimeter

27. Pressure Constant Calorimetry Much simpler device to determine heat changes for non- combustion reactions Can use two Styrofoam cups to contain heat for measurement