chemistry and energy

How are they related? Energy Encountered Daily What is Energy?  Defined as the ability to work or create heat  Many types of energy  Thermal  Light  Gravitational  Kinetic  Potential Light Energy Review  How is light energy produced?  Electrons release light energy when they fall from a high energy level to a lower energy  We’re now going to talk about energy released or used in a chemical reaction Heat energy Thermochemistry  The study of heat used or released in a chemical reaction  Let’s investigate heat as it compares to temperature using the Heat vs Temperature Handout Specific Heat Calculations  q = mCΔT q = heat (J or cal or Cal) 4.184 cal = Joule 1000 cal = Cal (dietary calorie) m = mass (g) C = specific heat (J/g oC or cal/g oC) ΔT = change in temperature (o C or K) = Tf - Ti Specific Heat  Specific heat of water = cal /g o C or = 4.184 J / goC  Specific heat of most metals = < J / goC  Do metals heat slowly or quickly compared to water?  Do metals stay warm longer or shorter than water? Practice Problem  How much energy is required to heat 120.0 g of water from 2.0 oC to 24.0oC? q = mCΔT m= 120.0 g C = 4.184 J/goC ΔT= (24.0 – 2.0)oC = 22.0oC q = (120.0g)(4.184 J/goC)(22.0oC) = Practice Problem  How much heat (in kJ) is given off when 85.0 g of lead cools from 200.0oC to 10.0 oC? (Specific heat of lead = 0.129 J/g oC) q = mCΔT m = 85.0 g C = 0.129 J/g oC ΔT = (10.0 – 200.0)oC = - 190.0oC q = (85.0 g)(0.129 J/g oC)(- 190.0oC) = - How Do Chemical Reactions Create Heat energy?  Consider the combustion of gasoline (octane) C8H18 +25 O2  16 CO2 +18 H2O  Potential Energy: Stored energy  Potential energy is stored in the bonds of the reactant s and the products  When bonds are broken, the energy is available  When produce bonds form, some energy is used in these bonds  The excess energy is released as heat How does ice melt? Molar Heat of Fusion  Heat absorbed by one mole of a substance during melting  Constant temperature  ∆Hfus  H2O(s) → H2O(l) ∆H = 6.01 kJ/mol Molar Heat of Solidification  Heat lost when mole of a liquid solidifies  Temperature is constant  ∆Hsolid  ∆Hfus = -∆Hsolid  H2O(l) → H2O(s) ∆H = -6.01 kJ/mol Molar Heat of Vaporization  Heat needed to vaporize mole of a liquid  ∆Hvap  H2O(l) → H2O(g) ∆Hvap = 40.7 kJ/mol Molar Heat of Condensation  Heat released when mole of vapor condenses  ∆Hcond  H2O(g) → H2O(s)  ∆Hvap = -∆Hcond ∆Hcond = -40.7 kJ/mol Phase Change Diagram for Water Phase Change Diagram The House that Heats Itself  http://www.sciencefriday.com/videos/watch/ 10007 Calorimetry  Method used to determine the heat involved in a physical or chemical change  Relies on the law of conservation of energy Calorimeter Simple Calorimeter Calorimetry Math  Heat gained by the water = q  Heat lost by the system = -q mC∆T = q ∆T = Tf –Ti , m = mass, C = specific heat q gained by water = q lost by system  q water = - q system  mC∆T = -mC∆T (mass H2O)(spec heat H2O)(∆T H2O) = - (mass sys)(spec heat sys)(∆T sys) Standard Heat of Reaction  Heat change for the equation as it is written ∆H = ∆Hf(products) - ∆Hf(reactants) Standard Heats of Formation (∆Hf)  Change in enthalpy when mole of the compound is formed from its elements in their standard states at 25oC and 101.3 kPa Hess’s Law  A way to calculate the heat of a reaction that may be too slow or too fast to collect data from  Add together several reactions that will result in the desired reaction Add the ΔH for these reactions in the same way  ∑∆Htotal = ∑∆Hproducts - ∑∆Hreactants [...]... Conservation of Energy  Energy is not lost or gained in a chemical reaction  In a chemical reaction potential energy is transferred to kinetic energy Thermochemical Equations  An equation that includes the heat change  Example: write the thermochemical equation for this reaction  CaO(s) + H2O(l) →Ca(OH)2(s)∆H = -65.2 kJ CaO(s) + H2O(l) →Ca(OH)2(s) + 65.2 kJ Stoichiometry and Thermochemistry Tin...Kinetic Energy  Directly related to temperature Is Heat Used or Released?  Endothermic reactions used heat from the surroundings  Sweating  Refrigeration  Exothermic heat releases heat to the surroundings  Hot hands  Combustion  Exercise Endothermic Reactions  Decrease in kinetic energy  decrease in      temperature  heat will transfer... heat sys)(∆T sys) Standard Heat of Reaction  Heat change for the equation as it is written ∆H = ∆Hf(products) - ∆Hf(reactants) Standard Heats of Formation (∆Hf)  Change in enthalpy when 1 mole of the compound is formed from its elements in their standard states at 25oC and 101.3 kPa Hess’s Law  A way to calculate the heat of a reaction that may be too slow or too fast to collect data from  Add together... change  Relies on the law of conservation of energy Calorimeter Simple Calorimeter Calorimetry Math  Heat gained by the water = q  Heat lost by the system = -q mC∆T = q ∆T = Tf –Ti , m = mass, C = specific heat q gained by water = q lost by system  q water = - q system  mC∆T = -mC∆T (mass H2O)(spec heat H2O)(∆T H2O) = - (mass sys)(spec heat sys)(∆T sys) Standard Heat of Reaction  Heat change for... Exothermic Reactions  Increase in kinetic energy  increase in      temperature of system heat released to the environment resulting in a hotter environment Releases heat to its surroundings The system loses heat Negative value for q ∆H = q =