Chapter 15: Sec. 4 Hess’s Law and Standard Enthalpies of Formation NOTES

Hess’s Law
o Hess’s Law:
 States that if you can add two or more thermochemical equations to produce a
final equation for a reaction, then the sum of the enthalpy changes for the
individual reactions is the enthalpy change for the final reaction.
o Steps for using Hess’s Law:
 Step 1: Chemical equations must include the desired substances and have known
enthalpy changes.
 Example:

Step 2: Equations should be balanced individually as well as with the other
equations. If equations aren’t balanced with each other multiply each coefficient
and also the enthalpy change by the number needed to balance all equations.
 Example:

Step 3: In the desired equation, if the desired reactant is a product or the desired
product is a reactant, reverse the chemical equation and reverse the sign of the
enthalpy change.
 Example:

Step 4: Add all the equations together as well as the changes in enthalpy. Cancel
any terms that are common to both sides of the chemical equation.
 Example:
o Hess’s Law Example:
 Use the two equations below to determine ΔH for the decomposition of hydrogen
peroxide (2H2O2(l) → 2H2O(l) + O2(g)).
 2H2(g) + O2(g) → 2H2O(l) ΔH = -572 kJ
 H2(g) + O2(g) → H2O2(l) ΔH = -188 kJ

Standard Enthalpy (Heat) of Formation
o Standard state = 1 atm and 298 K (25°C)
o Standard enthalpy (heat) of formation (ΔHf°):
 The change in enthalpy that accompanies the formation of one more of the
compound in its standard state from its elements in their standard states.
o Steps for using standard enthalpies of formation:
 Step 1: Find the chemical equations that have the desired products and reactants
and also the standard enthalpies of formation for those equations.
 Example:

Step 2: Rewrite each equation to show the reactants and products desired.
Reverse the sign of the enthalpies of formation if the chemical equation was
reversed.
 Example:

Step 3: Multiply any chemical equation and its standard enthalpy of formation by
a number that will balance all the equations.
 Example:

Step 4: Add all the equations together as well as their standard enthalpies of
formation. Cancel out any terms that are the same on both sides of the chemical
equation.
 Example:
o The Summation Equation:
 ΔH°rxn =
 ∑=
 ΔH°f (products) = standard enthalpies of formation of all the products
 ΔH°f (reactants) = standard enthalpies of formation of all the reactants
o Summation Equation Example:
 Use the standard enthalpies of formation to calculate ΔH°rxn for the combustion of
methane (CH4(g) + 2O2(g) → CO2(g) + 2H2O(l))
 ΔH°f (CO2) = -394 kJ
 ΔH°f (H2O) = -286 kJ
 ΔH°f (CH4) = -75 kJ
 ΔH°f (O2) = 0.0 kJ