Transcript MODULE 4
MODULE 5 Energy and Thermodynamics Thermodynamics & Energy • Thermodynamics - The science of heat and work • Work - A force acting upon an object to cause a displacement • Energy - The capacity to do work and transfer heat Kinetic Energy • Kinetic Energy – KE = 1/2 mv2 – The energy of a moving object, depending on it’s mass and velocity – “Energy of motion” Potential Energy • Potential Energy – PE = mgh – The energy that results in an object’s position – “Stored energy” PE, KE and Work Internal Energy • Internal energy () - The sum of the potential energies and kinetic energies of the particles within a thermodynamic system ETOTAL = PE + KE + First Law of Thermodynamics • Law of Conservation of Energy – The total energy of the universe is constant – Heat, work, and other energy transfers in an event equal the total energy content both before and after the event has occurred – A battery stores chemical potential energy Measurement • 1 cal = 4.184J • 1000cal = dietary calorie = 1kcal • 1 calorie = amount of energy required to raise 1 gram of H2O, 1°C. Temperature & Heat • Heat is not the same as temperature • The more thermal energy, the more kinetic energy, the more motion the atoms and molecules will have • The total thermal energy of an object is the sum of all the individual energies • Thermal energy depends on the amount of substance as well as the temperature • Temperature changes are measured with a thermometer by heat transfer Heat Transfer • Occurs when 2 objects of different temperatures are brought into contact • Heat is transferred from the hotter object to the colder one • Transfer will continue until the 2 objects are at the same temperature, we call the system at thermal equilibrium • The amount of heat lost by the hotter one = the amount of heat gained by the colder one Heat Transfer • Exothermic – process where heat is transferred from a system to it’s surroundings • Endothermic – process where heat is transferred to the system from it’s surroundings Specific Heat The quantity of heat transferred depends on: – The amount of material – The overall change in temperature – The identity of the material transferring the energy Specific Heat Capacity • Specific heat is the quantity of heat required to raise the temperature of 1 gram of a substance by one kelvin • (See Table 6.1 on page 210) q = C x m x T Heat Calculations Calculate the heat absorbed by 15.0g of water required to raise the temperature from 20°C to 50°C. Where q=C·m·∆T Let q = heat = unknown C = Heat Capacity for H2O = 4.184J/gK m = mass = 15.0g ∆T = Tf – Ti = 50°C-20°C = 323K-293K = 30K q = (4.184J/gK)(15.0g)(30K) = 1.88x103J=1.88kJ Exothermic • • • • Heat is given off Q<0 Enthalpy (H) is negative (-) Energy of the products is less than the energy of the reactants • The balanced equation is written: 4Fe(s) + 3O2(g) 2Fe2O3(s): H = -1648kJ Endothermic • • • • Heat is absorbed Q>0 Enthalpy (H) is positive (+) Energy of the products is more than the energy of the reactants • The balanced equation is written: 2Fe2O3(s) 4Fe(s) + 3O2(g) : H = +1648kJ Standard Enthalpies • See Table 20 on page A-31 Hess’s Law • Hess’s Law states: If a reaction is the sum of 2 or more other reactions, the H for the overall reaction is the sum of all the H values of those individual reactions. • See CD-ROM Screen 6.17 • Calculate H for the lab data. Enthalpy Change for a Rxn. • The enthalpy change for a reaction can be calculated by the sum of the products H values minus the sum of the reactants H values. • See Screen CD-ROM 6.18 • See Table 20, Appendix L starting on page A.31 • Hrxn = [Hf (products)] -[Hf (reactants)]