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 –KE = 1/2 mv 2 –The energy of a moving object, depending on it’s mass and velocity –“Energy of motion”

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 E TOTAL = 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 H 2 O, 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 H 2 O = 4.184J/gK m = mass = 15.0g ∆T = T f – T i = 50°C-20°C = 323K-293K = 30K q = (4.184J/gK)(15.0g)(30K) = 1.88x10 3 J=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) + 3O 2(g)  2Fe 2 O 3(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: 2Fe 2 O 3(s)  4Fe (s) + 3O 2(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  H rxn =  [  H  f (products)] -  [  H  f (reactants)]