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If a chemical change is carried out at constant pressure and the only work done is caused by expansion or contraction, q for the change is called the enthalpy change with the symbol Δ H , or for reactions occurring under standard state conditions. The value of Δ H for a reaction in one direction is equal in magnitude, but opposite in sign, to Δ H for the reaction in the opposite direction, and Δ H is directly proportional to the quantity of reactants and products. Examples of enthalpy changes include enthalpy of combustion, enthalpy of fusion, enthalpy of vaporization, and standard enthalpy of formation. The standard enthalpy of formation, is the enthalpy change accompanying the formation of 1 mole of a substance from the elements in their most stable states at 1 bar (standard state). Many of the processes are carried out at 298.15 K. If the enthalpies of formation are available for the reactants and products of a reaction, the enthalpy change can be calculated using Hess’s law: If a process can be written as the sum of several stepwise processes, the enthalpy change of the total process equals the sum of the enthalpy changes of the various steps.
Explain how the heat measured in
[link] differs from the enthalpy change for the exothermic reaction described by the following equation:
The enthalpy change of the indicated reaction is for exactly 1 mol HCL and 1 mol NaOH; the heat in the example is produced by 0.0500 mol HCl and 0.0500 mol NaOH.
Using the data in the check your learning section of [link] , calculate Δ H in kJ/mol of AgNO 3 ( aq ) for the reaction:
Calculate the enthalpy of solution (Δ H for the dissolution) per mole of NH 4 NO 3 under the conditions described in [link] .
25 kJ mol −1
Calculate Δ
H for the reaction described by the equation. (
Hint : use the value for the approximate amount of heat absorbed by the reaction that you calculated in a previous exercise.)
Calculate the enthalpy of solution (Δ H for the dissolution) per mole of CaCl 2 .
81 kJ mol −1
Although the gas used in an oxyacetylene torch ( [link] ) is essentially pure acetylene, the heat produced by combustion of one mole of acetylene in such a torch is likely not equal to the enthalpy of combustion of acetylene listed in [link] . Considering the conditions for which the tabulated data are reported, suggest an explanation.
How much heat is produced by burning 4.00 moles of acetylene under standard state conditions?
5204.4 kJ
How much heat is produced by combustion of 125 g of methanol under standard state conditions?
How many moles of isooctane must be burned to produce 100 kJ of heat under standard state conditions?
1.83 10 −2 mol
What mass of carbon monoxide must be burned to produce 175 kJ of heat under standard state conditions?
When 2.50 g of methane burns in oxygen, 125 kJ of heat is produced. What is the enthalpy of combustion per mole of methane under these conditions?
802 kJ mol −1
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