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Topic 6: Part I Thermochemistry. Thermochemistry : The Study of heat change in chemical reactions. Energy – the capacity to do work or produce heat Types of Energy: Radiant – solar energy, comes from the sun Thermal – energy associated with the random motion of atoms and molecules

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Topic 6: Part IThermochemistryThermochemistry: The Study of heat change in chemical reactionsEnergy – the capacity to do work or produce heatTypes of Energy: Radiant – solar energy, comes from the sun Thermal – energy associated with the random motion of atoms and molecules Chemical – energy stored within the structural units of chemical substances Potential – energy available by virtue of an object’s position Every chemical reaction obeys two fundamental Laws: (1) The Law of Conservation of Mass, and (2) The Law of Conservation of Energy The nature of energyEnergy is the capacity to do work or produce heat. Law of Conservation of Energy: energy can be converted from one form to another, but not created or destroyed.Potential – energy of position or composition Kinetic – energy of motion Temperature reflects the avg KE of the molecules. Heat involves the transfer of energy between two objects. Energy is ……Change in energy (∆E) is calculated using heat (q) and work (w). It is a state function – meaning it is independent of the pathway, or how you get from point A to B. Work is a force acting over a distance. Heat is energy transferred between objects because of temperature difference. Chemical Energy in Reactions ( keeping track )The Universe is divided into two parts: the system (sometimes a reaction) and the surroundings. The system is the part we are concerned with. The surroundings are the rest. Exothermic reactions release energy to the surroundings. Endothermic reactions absorb energy from the surroundings. HeatPotential energyHeatPotential energyDirectionEvery energy measurement has three parts. A unit ( Joule, J= kg∙m2 / s2 ). A number indicating how many A sign to tell direction. negative - exothermic positive - endothermic SurroundingsSystemEnergyDE <0SurroundingsSystemEnergyDE >0We apply the same rules of energy change for heat and work.Heat given off is negative. Heat absorbed is positive. Work done by system on surroundings is negative. Work done on system by surroundings is positive. Thermodynamics- The study of energy and the changes it undergoes. First Law of ThermodynamicsThe energy of the universe is constant. Law of conservation of energy. q = heat w = work DE = q + w Take the system’s point of view to decide signs. Enthalpy abbreviated H H = E + PV (that’s the definition) E = internal energy of a system P = pressure of the system V = volume of the system at constant pressure. DH = DE + PDV Enthalpy Con’tthe heat, q, at constant pressure can be calculated from DE = q + w (w = - PDV) Substitute for work (w) and solve for heat (q) q = DE + P DV = DH Therefore q = DH at constant pressure We use DH as a measure of the change in enthalpy or heat in joules of a system. Enthalpy con’tHeat is a measure of change in enthalpy of a system at constant P. Heat of a reaction = change in enthalpy For a chemical reaction: ∆H = Hproducts – Hreactants + ∆H = heat absorbed by system (endothermic) - ∆H = heat lost by system (exothermic) Example6.4 When one mole of methane (CH4) is burned at constant pressure, 890 kJ of energy is released as heat. Calculate ∆H for a process in which a 5.8g sample of methane is burned at constant pressure. CalorimetryMeasuring heat. Use a calorimeter. Two kinds Constant pressure calorimeter (called a coffee cup calorimeter) heat capacity for a material, C, is the measure of the energy needed to raise the temp. of an object 1oC. C= heat absorbed/ DT = DH/ DT Calorimetryspecific heat capacity, s, is the heat capacity given per gram of substance (C/g). molar heat capacity = C/moles the heat of reaction, q = s x m x DT S = specific heat capacity m = mass of substance DT = change in temperature CalorimetryA coffee cup calorimeter measures DH. An insulated cup, full of water. The specific heat of water is 4.18J/ºC∙g Heat of reaction= DH = s x m x DT ExamplesIf 50ml of 1.0M HCl at 25oC is reacted with 50ml of 1.0M NaOH at 25oC. The final temperatrue of this exothermic reaction is 31.9oC. Calculate the change in heat Calculate ∆H in kJ/mol ExampleWhen 1.00 L of 1.00 M Ba(NO3)2 solution at 25oC is mixed with 1.00L of 1.00M Na2SO4 solution at 25oC in a calorimeter, the white solid BaSO4 forms and the temperature of the mixture increases to 28.1oC. Assuming the calorimeter absorbs only a negligible quantity of heat, that the specific heat capacity of the solution is 4.18 J/oC ∙ g and that the density of the final solution is 1.0g/mL, calculate the enthalpy change per mole of BaSO4 formed. CalorimetryConstant volume calorimeter is called a bomb calorimeter. Material is put in a container with pure oxygen. Wires are used to start the combustion. The container is put into a container of water. The heat capacity of the calorimeter is known and tested. Since DV = 0, PDV = 0, DE = q Bomb Calorimeterthermometer stirrer full of water ignition wire Steel bomb sample Propertiesintensive properties not related to the amount of substance. density, specific heat, temperature. Extensive property - does depend on the amount of stuff. Heat capacity, mass, heat from a reaction. Hess’s LawEnthalpy is a state function meaning it is independent of the path. We can add equations to come up with the desired final reaction and add the associated DH values to get the final ∆H. Two rules using Hess’s Law: If the reaction is reversed the sign of DH is changed If the reaction is multiplied by a constant, so is DH for that reaction. Nitrogen + Oxygen → Nitrogen DioxideOverall Reaction:N2 + 2O2 → 2NO2∆H1 = 68kJReaction steps added: N2 + O2 → 2NO ∆H2 = 180kJ2NO + O2 → 2NO2 ∆H3 = -112kJ____________________________________________________________________________________________________________________________________________________N2 + 2O2 → 2NO2∆H2 + ∆H3 = 68kJO2NO2-112 kJH (kJ)180 kJNO268 kJN22O2Characteristics of Enthalpy ChangesIf a reaction is reversed the sign of ∆H is also reversed. The magnitude of ∆H is directly proportional to the quantities of reactants and products. If coefficients are multiplied by an integer, then ∆H is multiplied by the same integer Examples: 6.7 and 6.8 Standard EnthalpyThe enthalpy change for a reaction at standard conditions (25ºC, 1 atm , 1 M solutions) Symbol DHº When using Hess’s Law, work by adding the equations up to make it look like the answer. The other parts will cancel out. Standard Enthalpies (Heat) of FormationHess’s Law is much more useful if you know lots of reactions. Made a table of standard heats of formation. The amount of heat needed to form 1 mole of a compound from its elements in their standard states. Standard states are 1 atm, 1M and 25ºC For an element it is 0 There is a table in Appendix 4 (pg A21) Standard Enthalpies (Heat) of FormationNeed to be able to write the equations. What is the equation for the formation of NO2 ? ½N2 (g) + O2 (g) ® NO2 (g) Have to make one mole to meet the definition. Write the equation for the formation of methanol CH3OH. Since we can manipulate the equationsWe can use heats of formation to figure out the heat of reaction. Examples 6.9 and 6.10Using the standard enthalpies of formation listed in the appendix, calculate the standard enthalpy change for the overall reaction that occurs when ammonia is burned in air to form nitrogen dioxide and water. 4NH3(g) + 7O2(g) → 4NO2(g) + 6H2O (l) Using enthalpies of formation, calculate the standard change in enthalpy for the thermite reaction: 2Al (s) + Fe2O3(s) → Al2O3(s) + 2Fe (s)

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