1. Unit 13 Thermochemistry

2. Energy u The ability to do work or cause a change u Often measured in joules (J) u Law of Conservation of Energy – energy is neither created nor destroyed It is often transferred & can change form

3. Heat (q) u Heat – energy transfer between objects at different temperatures Goes from warmer object to cooler object and not vice-versa u Temperature (T) – measure of how hot or cold OR measure of average kinetic energy Higher temp = faster movement

4. Heat vs. Temperature u Heat is an EXTENSIVE PROPERTY Physical property that depends on the amount of sample present Match vs campfire u Temperature is an INTENSIVE PROPERTY Physical property that does not depend on the amount of sample

5. Enthalpy (H) u Total energy content in a sample under a constant pressure u The change in enthalpy (  H) is often used interchangeably with heat (q)

6. 6 Units of Energy u Energy can also be measured in calories u calorie is amount of heat to change 1 g of water by 1  C u Food C alories are kilocalories 1 C alorie = 1 kilocalorie = 1000 calories u 1 c al = 4.184 J u 1 C al = 4184 J

7. Heat in phase changes u Remember, the temperature does not change during a phase change

8. Latent heat u Heat absorbed or released by a substance during a process that does not involve a change in temperature u Solid  Liquid = Latent Heat of fusion (L f ) u Liquid  Gas = Latent Heat of Vaporization (L v ) u Varies for substances based on IMF & mass

10. Calculating latent heat u Q=m L f OR Q=m L v u How much energy does it take to melt 16 g of ice at it’s melting point? The L f of of ice is 80 cal/g u Q = m L f u Q = 16g * 80 cal/g u Q = 1280 calories OR 1.280 C alories

11. Molar heat of fusion ( H fus ) u Heat absorbed by 1 mole of a solid as it melts to a liquid at a constant temperature u Molar heat of solidification ( H solid ) u Heat lost when 1 mole of a liquid becomes a solid u ( H fus ) = ( H solid ) but opposite sign

13. Using molar heat of fusion u How many grams of ice at 0 deg C melts if 2.25 kJ heat added? The molar heat of fusion of water is 6.01 kJ/mol. 2.25 kJ mol kJ 1 6.01 g mol 18 1

14. Molar heat of vaporization u Amount of heat to vaporize 1 mole of a substance u Equal but opposite to the molar heat of condensation

15. UNDERSTANDING CHECK u The heat of vaporization of water is 40.79 kJ/mol. How much heat is required to turn 14 g of water into vapor at the boiling point?

16. Exothermic Reaction u Change in matter where E is released, often as heat u Surroundings get hotter u q or  H of the system is negative System Surroundings Energy

17. 17 Exothermic Reactions u Joules of energy created are written with the products of the balanced reaction u On a graph, the products are lower in energy than the reactants u Makes sense because energy is RELEASED during the reaction

18. 18 C + O 2  CO 2 Energy ReactantsProducts  C + O 2 CO 2 -395kJ + 395 kJ

19. Endothermic u Change in matter in which energy is absorbed, often as heat u Surroundings get cooler u q or  H of the system is positive System Surroundings Energy

20. 20 Endothermic u Joules of energy absorbed are written with the reactants of the balanced reaction u On a graph, the products are higher in energy than the reactants u Makes sense because energy is ABSORBED during the reaction

21. 21 CaCO 3  CaO + CO 2 Energy ReactantsProducts  CaCO 3 CaO + CO 2 +176 kJ CaCO 3 + 176 kJ  CaO + CO 2

23. 23 Heat capacity (c) u How much heat it takes to heat an object by 1  C u Affected by identity of the substance (sand vs. water)& amt present (ocean vs. pan) u Molar heat capacity (C) – Heat needed to increase temperature of 1 mole by 1 degree Celsius Affected by identity of substance q = n C ▲T

24. Heat capacity u Specific heat capacity (c) is the amount needed to heat 1 g of a substance by 1  C u Affected by identity of the substance u The higher the specific heat, the more energy it takes to change its temperature u Formula q = m c  T u For water, c = 1 cal/gºC or 4.186 J/gºC

25. u How much heat is needed to change the temperature of 12 g of silver with a specific heat of 0.057 cal/g  C from 25  C to 83  C? u Q = m c  T u m = 12 g u c = 0.057 cal/g  C u  T = 58  C (no need to change to Kelvin)

26. u If you put 6500 J of heat into a 15 g piece of Al at 25  C, what will the final temperature be? ( c = 0.90 J/g  C ) u Q = m c  T u m = 15 g u c = 0.90 J/g  C u Q = 6500 J

27. Applying specific heat u Equal masses of copper and water are heated until the temperature is increased by 10 degrees. Why did the copper’s temperature increase more quickly? u Water has a very HIGH heat capacity and therefore changes temp SLOWLY u Walking on water or sand on a hot day?

28. 1 st law of thermodynamics u Law of Conservation of Energy – The change in the internal energy of a system is equal to the amount of energy added by heating minus the amount lost by the work done by the system u Eating a cheeseburger gives you the amount of energy equal to what’s in the burger minus the energy used to chew & digest

29. Calorimetry u Measures heat using a device called a calorimeter u Measure temp before and after reaction u Since the change in enthalpy (  H) can be used interchangeably with heat (q), we can use the equation q = m c  T u Water c = 1 calorie/gram * degree Celsius c = 4.186 Joule/gram * degree Celsius

30. Bomb calorimeter

31. Coffee Cup Calorimeter

32. Example u A chemical reaction is carried out in a coffee cup calorimeter. There are 75.8 g of water in the cup, and the temperature rises from 16.8 ºC to 34.3 ºC. How much heat was released?