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Thermochemistry

Thermochemistry. Powerpoint #2. Enthalpy. The chemical and physical changes that occur around us essentially occur at constant atmospheric pressure.

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Thermochemistry

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  1. Thermochemistry Powerpoint #2

  2. Enthalpy • The chemical and physical changes that occur around us essentially occur at constant atmospheric pressure. • Most commonly, the only kind of work produced by chemical and physical changes open to the atmosphere is the mechanical work associated with change in the volume of the system.

  3. Pressure – Volume Work • The work involved in the expansion or compression of gases if called P-V work. • When pressure is constant, as in our example, the sign and magnitude of the pressure-volume work is given by • w = -P DV • The negative sign is necessary to conform with the sign conventions given earlier.

  4. W = - PΔV • When the volume expands, ΔV is a positive quantity and w is a negative quantity. • Energy leaves the system as work, indicating that work is done by the system on the surroundings. • When the volume contracts, ΔV is a negative quantity and w is a positive quantity. • Energy enters the system as work, indicating that work is done on the system by the surroundings

  5. Enthalpy, denoted H • From a Greek word meaning “to warm” • Accounts for heat flow in processes occurring at constant pressure when no forms of work are performed other than P-V work. • H = E + PV (its definition) • Is a state function because internal energy, pressure, and volume are all state functions.

  6. Enthalpy • When a change occurs at constant pressure, the change in enthalpy is given by the following relationship: • ΔH = DE + PDV • the heat at constant pressure qp can be calculated from • ΔE = qp+ w = qp- PΔV • qp= ΔE + P ΔV = ΔH

  7. The change in enthalpy equals the heat gained or lost at constant pressure. • Because qp is something we can either measure or readily calculate and because so many physical and chemical changes occur at constant pressure, enthalpy is a more useful function than internal energy. • For most reactions the difference in ΔH and ΔE are so small because P ΔV is small.

  8. Enthalpies of Reaction • Because ΔH = ΔH final – ΔH initial, the enthalpy change for a chemical reaction given by the enthalpy of the products minus the enthalpy of the reactants. • ΔH = ΔHproducts – ΔHreactants • Enthalpy of reaction–the enthalpy change that accompanies a reaction • Sometimes called –Heat of reaction • ΔHrxn

  9. Guidelines when using thermochemical equations: • 1. Enthalpy is an extensive property –ΔH is directly proportional to the amount of reactant consumed in the process. • 2. The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to the ΔH for the reverse reaction. • 3. The enthalpy change for a reaction depends on the state of the reactants and products. Important to note the states in your equations.

  10. Calorimetery • Measurement of heat flow. • Use a calorimeter. • The temperature change experienced by an object when it absorbs a certain amount of heat. • It is a unique value for all substances. • Heat Capacity – the amount of heat required to raise its temperature by 1 K. •  • The greater the heat capacity, the greater the heat required to produce a given increase in temperature.

  11. The heat capacity of one mole of a substance. • C – has units of J/gxK or J/gx◦C • Specific heat capacity (specific heat) –the heat capacity of one gram of a substance • q • C = ------ or q = C x m x ΔT • m x ΔT

  12. Bomb Calorimetry • Constant 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 ΔV = 0, PΔV = 0, ΔE = q

  13. Bomb Calorimeter • thermometer • stirrer • full of water • ignition wire • Steel bomb • sample

  14. Properties • Intensive 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.

  15. Hess’s Law • Enthalpy is a state function. • It is independent of the path. • We can add equations to come up with the desired final product, and add the ΔH • Two rules • If the reaction is reversed the sign of ΔH is changed • If the reaction is multiplied or divided, so is ΔH

  16. Remember… • H is a state function, so for a particular set of reactants and products, ΔH is the same whether the reaction takes place in one step or in a series of steps.

  17. Enthalpies of Formation • The enthalpy change associated with the formation of a compound from its constituent elements. • Labeled ΔHf • The enthalpy change for a reaction at standard conditions (25ºC, 1 atm , 1 M solutions). These are standard state conditions. • Symbol ΔHfº • When using Hess’s Law, work by adding the equations up to make it look like the answer. • The other parts will cancel out.

  18. Standard Enthalpies of Formation • Hess’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 for 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

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