1. Thermochemistry Some Like It Hot!!!!!

2. The Flow of Energy ► Thermochemistry – concerned with heat changes that occur during chemical reactions ► Energy - capacity for doing work or supplying heat

3. Is there a difference between heat and temperature? ► Temperature – Average kinetic energy of the particles in matter ► Heat - represented by “q”, is energy that transfers from one object to another, because of a temperature difference between them.  only changes can be detected!  flows from warmer  cooler object

4. Exothermic and Endothermic Processes ► Essentially all chemical reactions and changes in physical state involve either:  release of heat, or absorption of heat ► In studying heat changes, think of defining these two parts:  the system - the part of the universe on which you focus your attention  the surroundings - includes everything else in the universe ► Law of Conservation of Energy – Energy is neither created nor destroyed

5. Exothermic vs Endothermic ► Endothermic – Heat flowing into a system from its surroundings  Surroundings get cooler but the system gets warmer  q has a positive value ► Exothermic - Heat flowing out of a system into it’s surroundings  System gets colder but the surroundings get warmer  q has a negative value

6. ► Endothermic ► Exothermic System Heat System increases in temperature, Surroundings decrease in temperature System decreases in temperature, surroundings increase in temperature

7. Endothermic and Exothermic Reaction Review ► Every reaction has an energy change associated with it ► Exothermic reactions release energy, usually in the form of heat. ► Endothermic reactions absorb energy ► Energy is stored in bonds between atoms

8. Measuring Heat Flow ► Measured with 2 common units  Calorie – quantity of heat needed to raise the temperature of 1 g of pure water 1 o C.  Joule - SI unit of heat and energy ► How are the units similar?  1 J = 0.2390 calories  1 Calorie = 4.184 J

9. Heat Capacity and Specific Heat ► Heat Capacity – the amount of heat needed to increase the temperature of an object exactly 1°C  Dependant on 2 things: ► Mass of the object ► Chemical Composition ► Specific Heat – The amount of heat it takes to raise the temperature of 1 g of the substance 1°C  Written as J/(g x °C)

10. The higher the specific heat the more energy needed to change the temperature Note the tremendous difference in Specific Heat. Water’s value is VERY HIGH.

11. Calculating the Quantity of Heat Energy ► To calculate, use the formula: q = mass (in grams) x  T x C ► heat is abbreviated as “q” ►  T = change in temperature   T = Final Temp. – Initial Temp ► C = Specific Heat  Units are either: J/(g o C) or cal/(g o C)

13. Measuring and Expressing Enthalpy Changes ► Enthalpy (H) - The heat content of a system at constant pressure ► Calorimetry - the measurement of the heat into or out of a system for chemical and physical processes.  Based on the fact that the heat released = the heat absorbed ► The device used to measure the absorption or release of heat in chemical or physical processes is called a “Calorimeter”

14. ► For systems at constant pressure, the “heat content” is the same as a property called Enthalpy (H) of the system A foam cup calorimeter – here, two cups are nestled together for better insulation

15. Calorimetry ► Changes in enthalpy =  H ► q =  H These terms will be used interchangeably in this textbook ► Thus, q =  H = m x C x  T 4  H is negative for an exothermic reaction 4  H is positive for an endothermic reaction

16. Thermochemical Equations ► In a chemical equation, the heat change for the reaction can be written as either a reactant or a product

17. C + O 2 → CO 2 Energy ReactantsProducts  C + O 2 CO 2 395kJ given off + 395 kJ

18. Exothermic ► The products are lower in energy than the reactants ► Thus, energy is released. ► ΔH = -395 kJ  The negative sign does not mean negative energy, but instead that energy is lost.

19. CaCO 3 → CaO + CO 2 Energy ReactantsProducts  CaCO 3 CaO + CO 2 176 kJ absorbed CaCO 3 + 176 kJ → CaO + CO 2

20. Endothermic ► The products are higher in energy than the reactants ► Thus, energy is absorbed. ► ΔH = +176 kJ  The positive sign means energy is absorbed

21. Heat in Change of State ► Heat of Fusion  What happens if you place an ice cube on a table in warm room? ► Heat from the room travels to the cooler ice cube causing the ice to melt ► The gain of heat causes a change in state instead of a change in temperature ► The heat required to change the water from a solid to a liquid is called the molar heat of fusion

22. Heat of Fusion Molar Heat of Fusion (  H fus. ) = the heat absorbed by one mole of a substance in melting from a solid to a liquid q = mol x  H fus. (no temperature change) Values given in Table 17.3, page 522 Example: How much heat is required to melt 5 moles of ice? q = mol x  H fus. q = 5 moles H 2 O (s) x 6 kJ/mole q = 30 kJ

23. Another Example How many grams of ice at 0°C will melt if 2.25 kJ of heat are added? q = mol x  H fus. 2.25 kJ x (1 mol ice/6.01 kJ) x (18 g Ice/1 mol ice) 6.7 g of Ice

24. Heat in Change of State ► Molar Heat of Vaporization  Similar to Molar Heat of Fusion, however the state of matter change is from a liquid to a vapor  The amount of heat necessary to vaporize one mole of a given liquid

25. Heat of Vaporization Molar Heat of Vaporization (  H vap. ) = the amount of heat necessary to vaporize one mole of a given liquid. q = mol x  H vap. (no temperature change) ► Table 17.3, page 522 ► How much heat is transferred when 5 moles of mercury (Hg) is vaporized? q = mol x  H vap. q = 5 moles X 59.1 kJ /mol q = 295.5 kJ

26. Another Example ► Example: How much heat (in kJ) is absorbed when 24.8 g H 2 O(l) at 100°C and 101.3 kPa is converted to steam at 100°C? 24.8 g H 2 O x (1 mol H 2 O/18 g H 2 O) x (40.7 kJ/1mol H 2 O) ΔH = 56.1 kJ

27. The solid temperature is rising from -20 to 0 o C (use q = mol x ΔT x C) The solid is melting at 0 o C; no temperature change (use q = mol x ΔH fus. ) The liquid temperature is rising from 0 to 100 o C (use q = mol x ΔT x C) The liquid is boiling at 100 o C; no temperature change (use q = mol x ΔH vap. ) The gas temperature is rising from 100 to 120 o C (use q = mol x ΔT x C) The Heat Curve for Water, going from -20 to 120 o C, similar to the picture on page 523 120

28. How to determine ΔH Fus. and ΔH Vap ► This is for water only:  ΔH Fus. and ΔH Vap. are in kJ/mol ΔH Fus. = (.334 kJ/g) x ( 18g H 2 0/mole) = 6.01 kJ/mol  334 J/g can be found in your reference tables for the heat of fusion of water. You must convert it to.334 kJ because ΔH Fus. is in kJ/mol. ΔH Vap = (2.26 kJ/g) x (18 g H 2 0/mol) = 40.7 kJ/mol  2260 J/g can be found in your reference tables for the heat of vaporization of water. You must convert it to 2.26 kJ/g because ΔH Fus. is in kJ/mol.

29. Entropy ► Defined as a measure of the disorder of a system  Systems tend to go from a state of order (low entropy)to a state of maximum disorder (high entropy) Increasing Entropy Solids  Liquid  Gas Ionic Compounds  Ions in solution

30. - Page 570 Entropy of the gas is greater than the solid or liquid Entropy is increased when a substance is divided into parts Entropy increases when there are more product molecules than reactant molecules Entropy increases when temperature increases