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Thermochemistry

Systems

  • In thermodynamics we must define the part of the universe where the energy changes occur

  • System: part of the universe that we are studying

  • Surroundings: everything outside of the system

  • Open system: matter & energy exchange with surroundings

  • Closed system: exchanges energy as work and/or heat (but not matter)

  • Isolated system: neither matter nor energy exchanged

First Law of Thermodynamics

  • First law of thermodynamics: energy is conserved

  • Internal energy, U, of a system: the sum of ALL potential and kinetic energies of the components of the system

  • A precise, numerical value cannot be determined

  • We consider only the change in internal energy

    • delta U = Ufinal – Uinitial

  • delta U = q + w

State Functions

  • Depend only on the present state of the system and NOT on the PATH by which the system arrived at that state

  • The internal energy of the system is the same regardless of which PATH was used

Enthalpy

  • “PV-work” is work involved in expansion or compression of gases

  • At constant pressure: w = -P(delta V)

  • If delta V is + → expansion of gas; system does work on surroundings and w is negative

  • If delta V is - → compression of gas; work is done on the system and w is positive

  • Like internal energy E, both P and V are state functions

  • We can combine these state functions to define “ENTHALPY”, H

    • H = U + PV

  • At constant pressure: delta H = delta E + P(delta V)

    • delta H = delta U + P(delta V) = (qpw) –w = qp

    • delta H = qp

      • This equation is important because qp is easily measured

  • If delta H is positive, the reaction is endothermic

  • If delta H is negative, the reaction is exothermic

  • Reversing the direction of a reaction changes the sign of delta H

  • If the coefficients of a chemical reaction are multiplied by some factor, the enthalpy change must also be multiplied by that factor

    • (delta H is an extensive property).

  • Enthalpy change depends on the physical states of reactants and products and these states must be specified.

Constant Pressure Calorimetry

  • The reaction is carried out in aqueous solution

  • System = reactants and products Surroundings includes the water

  • Heat released or absorbed by the reaction changes the temperature of the solution

  • delta H can be determined experimentally by measuring heat flow for a reaction at constant pressure

Heat Capacity and Specific Heat

  • Heat capacity: the amount of heat required to raise the temperature of an object by one degree C or K

  • Specific heat: the heat capacity of one gram of a substance; Cs

Enthalpies of Formation

  • Enthalpy of formation: the enthalpy change for the reaction in which one mole of a substance is made from its constituent elements in their elemental (most stable) forms.

  • Standard enthalpies of formation: measured under standard conditions

Thermochemistry

Systems

  • In thermodynamics we must define the part of the universe where the energy changes occur

  • System: part of the universe that we are studying

  • Surroundings: everything outside of the system

  • Open system: matter & energy exchange with surroundings

  • Closed system: exchanges energy as work and/or heat (but not matter)

  • Isolated system: neither matter nor energy exchanged

First Law of Thermodynamics

  • First law of thermodynamics: energy is conserved

  • Internal energy, U, of a system: the sum of ALL potential and kinetic energies of the components of the system

  • A precise, numerical value cannot be determined

  • We consider only the change in internal energy

    • delta U = Ufinal – Uinitial

  • delta U = q + w

State Functions

  • Depend only on the present state of the system and NOT on the PATH by which the system arrived at that state

  • The internal energy of the system is the same regardless of which PATH was used

Enthalpy

  • “PV-work” is work involved in expansion or compression of gases

  • At constant pressure: w = -P(delta V)

  • If delta V is + → expansion of gas; system does work on surroundings and w is negative

  • If delta V is - → compression of gas; work is done on the system and w is positive

  • Like internal energy E, both P and V are state functions

  • We can combine these state functions to define “ENTHALPY”, H

    • H = U + PV

  • At constant pressure: delta H = delta E + P(delta V)

    • delta H = delta U + P(delta V) = (qpw) –w = qp

    • delta H = qp

      • This equation is important because qp is easily measured

  • If delta H is positive, the reaction is endothermic

  • If delta H is negative, the reaction is exothermic

  • Reversing the direction of a reaction changes the sign of delta H

  • If the coefficients of a chemical reaction are multiplied by some factor, the enthalpy change must also be multiplied by that factor

    • (delta H is an extensive property).

  • Enthalpy change depends on the physical states of reactants and products and these states must be specified.

Constant Pressure Calorimetry

  • The reaction is carried out in aqueous solution

  • System = reactants and products Surroundings includes the water

  • Heat released or absorbed by the reaction changes the temperature of the solution

  • delta H can be determined experimentally by measuring heat flow for a reaction at constant pressure

Heat Capacity and Specific Heat

  • Heat capacity: the amount of heat required to raise the temperature of an object by one degree C or K

  • Specific heat: the heat capacity of one gram of a substance; Cs

Enthalpies of Formation

  • Enthalpy of formation: the enthalpy change for the reaction in which one mole of a substance is made from its constituent elements in their elemental (most stable) forms.

  • Standard enthalpies of formation: measured under standard conditions

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