Name: _______________________
Date: _______________________
Learning Target:
I can use thermochemical equations to calculate energy
changes that occur in chemical reactions and classify
reactions as exothermic or endothermic.
Criteria for Success:
I can use Hess’s Law to calculate the enthalpy (heat) of
reaction.
I can use standard enthalpies (heats) of formation to
calculate the enthalpy (heat) of reaction.
Performances of Understanding:
-I will listen attentively while taking written notes in
Cornell format.
Hess’s Law
A. Sometimes it is impossible or impractical to measure the ΔH of a reaction by using a calorimeter. In these situations,
ΔH can be calculated using ________ Law.
1. ________ Law states that if you can ______ 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.
2. When manipulating reactions to achieve the final reaction keep in mind the following:
a. A compound or element must appear on the reactant side of one reaction and the product side of a
second in order to __________ out.
b. If a reaction is ________________, the sign of ΔH is also ______________.
c. If the coefficients in a balanced reaction are ____________ by an integer, the value of ΔH is
_________ by the same integer.
Standard Enthalpy (Heat) of Formation & Summation Equation
B. It would be an insane task to record the ΔH values for all known chemical reactions. Instead, scientists record what is
known as standard enthalpy (heat) of ______________, ΔH°f for a compound.
1. The standard enthalpy (heat) of ____________, ΔH°f, is defined as the change in enthalpy that accompanies
the formation of _____ mole of the compound in its standard state from its _________ their standard states (at
1 atm and 298K).
Examples
C(s) + 2H2(g) → CH4(g)
ΔH°f = -74.9kJ
3
S(s) + 2 O2(g) → SO3(g)
ΔH°f = -396kJ
a. For comparison purposes, note that you will see _______________ coefficients to ensure only
_______ mole of compound is formed.
b. The ΔH°f for __________ is always _______ kJ as they are used as the reference for the ΔH°f of
compounds.
C. The step-wise process outlined in Hess’s Law can be summed up using the following equation:
ΔHrxn = ƩΔH°f products - ƩΔH°f reactants
Guided Practice
1. Find the enthalpy change for the formation of lead(IV) chloride by the reaction of lead(II) chloride with chlorine.
PbCl2(s) + Cl2(g)  PbCl4(l)
∆H = ?
Use the following thermochemical equations:
Pb(s) + 2Cl2(g)  PbCl4(l)
∆H = –329.2 kJ
Pb(s) + Cl2(g)  PbCl2(s)
∆H = –359.4 kJ
2. Use standard heats of formation ( H f ) to calculate the change in enthalpy for the following reaction:
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)

Independent Practice
3. Find the enthalpy change for the formation of pentane, C5H12, by the reaction of carbon with hydrogen.
5C(s) + 6H2(g)  C5H12(g)
Use the following thermochemical equations:
C(s) + O2(g)  CO2(g)
∆H = -393.5kJ/mol
H2(g) + ½ O2(g)  H2O(l)
∆H = -285.8kJ/mol
C5H12(g) + 8O2(g)  5CO2(g) + 6H2O(l)
∆H = -3535.6kJ/mol
4. Calculate the heat of formation for sulfur dioxide, SO2, from its elements sulfur and oxygen. Use the balanced
chemical equation and the following information.
S(s) + 3 O2(g)  SO3(g)
∆H = -395.2kJ/mol
2SO2(g) + O2(g)  2SO3(g)
∆H = -198.2kJ/mol
2

5-16. Use standard heats of formation ( H f ) to calculate the change in enthalpy for the following reactions. Then
indicate whether the reaction is endothermic (A) or exothermic (B):
5-6. C(s) + O2(g)  CO2(g)
5. ( H f )

6. endothermic (A) or exothermic (B)
7-8.
 CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
7. ( H f )
8. endothermic (A) or exothermic (B)
9-10.
 CaCO3(s)  CaO(s) + CO2(g)
11. ( H f )
12. endothermic (A) or exothermic (B)
13-14.
 H2O(g)  H2O(l)
13. ( H f )
14. endothermic (A) or exothermic (B)
15-16.
 3CO(g) + Fe2O3(s)  2Fe(s) + 3CO2(g)
15. ( H f )
16. endothermic (A) or exothermic (B)
