Chem 1B Ch. 9 - Thermochemistry

conservation of energy

energy lost by the system = energy gained by its surroundings (ΔE system = -ΔE surroundings)

chemical energy

• when ΔE_system is negative the system loses energy and heat/work to the surroundings

• when ΔE_system is positive the system gains energy as heat/work from the surroundings

work & energy

• work performed by or on the system is energy transfer that results in macroscopic changes in the system

• w = -PΔV

• ΔE_system = q + w

quantifying heat

• (heat) q = m x c x ΔT

convention

• if energy flows into the system, it has a positive sign because the system’s energy is increasing

• if energy leaves the system, it has a negative sign because the system’s energy is decreasing

constant-volume process

• ΔV = 0

• change in internal energy is equal to the heat exchanged

constant-pressure process

• ΔH = ΔE + PΔV

• under constant pressure, the heat exchanged is defined as the enthalpy (H)

• enthalpy = most useful state function for constant pressure processes

enthalpy

• q_p = heat flow at constant pressure

• ΔE = q_p + w

• w = -PΔV

• therefore: ΔE = q_p - PΔV = ΔH - PΔV

• ΔH = ΔE + PΔV

enthalpy change of a reaction (1)

• if sign of ΔH is positive: reaction is endothermic (surroundings will cool in the reaction)

• if the sing of ΔH is negative: reaction is exothermic (the surroundings will warm up)

enthalpy change of a reaction (2)

• ΔH = ∑ (ΔH bonds broken) + ∑ (ΔH bonds formed)

• ΔH is always positive for breaking bonds and negative for forming them

• endothermic reactions = positive ΔH → strong bonds break & weak bonds form

• exothermic reactions = negative ΔH → weak bonds break & strong bonds form

conditions of constant volume and constant pressure

• under constant pressure conditions some of the energy released is to do work on the surroundings by expanding against it → less energy is manifested as heat

• under constant volume conditions all of the energy released is evolved as heat

• ex. if the same mass of a fuel is combusted under the 2 conditions, the reaction that produces a smaller heat (kJ) corresponds to constant pressure & a larger heat corresponds to constant volume

specific heat capacity

• the quantity of heat needed to raise the temperature of 1g of a substance by 1°C

• when substances are equal in mass…

  • objects with a low specific heat capacity need to absorb less heat energy to increase in temperature

  • objects with a high specific heat capacity need to absorb more heat energy to increase in temperature

If there are equal masses of each, which substance must absorb the greatest amount of heat to increase in temperature?

the substance with the highest heat capacity must absorb the greatest amount of heat to increase in temperature

bomb calorimetry

q_cal = -q_rxn

• All of the heat lost by the reaction will be gained by the calorimeter

• This means the reaction and the calorimeter have equal magnitudes of q with opposite signs

• Because heat is lost by the reaction, it has a negative sign.

coffee cup calorimetry

endothermic reactions are defined by…

• a positive ΔH

• absorption of heat (reaction is cooler than the surroundings)

exothermic reactions are defined by…

• a negative ΔH

• release of heat (reaction is hotter than surroundings)

What is the ΔH∘reaction in units of kJ/mol for the chemical reaction:

Ag⁺(aq) + I^-(aq) → AgI(s)

Given:

• ΔH∘fAg⁺(aq)=105.8 kJ/mol

• ΔH∘fI^-(aq) =−56.78 kJ/mol

• ΔH∘fAgI(s) =−61.8 kJ/mol

-110.82 kJ/mol = ∑ products (AgI(s)) - ∑ reactants (Ag+(aq) & I-(aq))