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In Class Exercise #4
Energy and Enthalpy
When chemical reactions occur, energy commonly is released (an exothermic
reaction) or consumed (an endothermic reaction). Energy is also consumed in
many physical processes; melting ice requires energy and thus is an endothermic
process. The amount of energy consumed or produced is designated as the
enthalpy change (ΔH) for the chemical reaction or physical process. When
energy is released the sign of ΔH is negative and the reaction is exothermic;
consumption of energy gives a positive sign of ΔH and the reaction is endothermic.
Chemical reactions involve rearrangements of the bonds that hold atoms close
together. Moving atoms far apart breaks their bonds, requires energy, and is an
endothermic process. Conversely, moving atoms together to form a bond is
exothermic because it releases energy. In chemical reaction bonds of the reactants
are broken and new bonds are formed in the products. Whether the reaction is
overall endothermic or exothermic depends on the energy needed to make the
bonding rearrangements.
1. Consider the following reaction in which H2 is separated into two H atoms.
H2  H + H or, alternatively,
H2 2H
Is this reaction endothermic or exothermic?
Is the sign of ΔH positive or negative?
Add energy (or heat) to the reactant or product side of the equations above, as
appropriate.
Why is the energy released when a bond is formed exactly equal to the energy
required to break a bond?
2
Examine the following table of information
Reaction
H2  2H
O2  2O
N2  2N
F2  2F
ΔH (kJ/mole of reactant)
+453
+498
+945
+158
Why are all of the ΔH values positive?
Order these diatomic molecules according to increasing strength of bonds between
the two atoms.
Energy Diagrams
Enthalpy changes associated with chemical and physical processes can be
graphically depicted. The vertical scale is enthalpy; if a process gives off energy it
is depicted as going to decreasing enthalpy. The difference in enthalpy between
the upper and lower levels is the amount of energy released.
2H  H2
Enthalpy (H)
2 moles H
ΔH=-453 kJ
1 mole H2
On the same scale depict the conversion O2  2O. Which way does the arrow go
for this reaction?
3
The formation of one mole of methane from one mole carbon atoms and four
moles of hydrogen atoms releases 1662 kJ of heat.
C + 4H  CH4 ΔH=-1662 kJ/mol
Methane has four C-H bonds. What is the average strength of a C-H bond? Is this
stronger or weaker than the H-H bond?
How much energy is released when two moles of CH4 are generated from H and C
atoms?
Examine the enthalpy diagram shown below for taking one mole of C atoms and
two moles of H2 to methane.
C + 2H2  CH4
4 moles H + 1 mole C
Enthalpy (H)
ΔH=+906 kJ
ΔH=-1662 kJ
2 mole H2 + 1 mole C
One mole CH4
4
Based on this diagram, which arrow represents the enthalpy change for 2 mole H2
+ 1 mole C  one mole CH4? Circle the arrow.
Is this reaction endothermic or exothermic? Why?
What is the value of ΔH for the reaction?
Let’s assume that the reaction of C + 4F  CF4 has the same enthalpy (ΔH=-1664
kJ/mole CF4) as the reaction C + 4H  CH4. Draw an enthalpy diagram for the
reaction of 1 mole C + 2 mole F2  CF4. Compute the enthalpy of the reaction
Because we assumed that C + 4F  CF4 and C + 4H  CH4 have the same
enthalpy, the main difference in the energies of the reactions
2 mole H2 + 1 mole C  one mole CH4
2 mole F2 + 1 mole C  one mole CF4
can be traced to the difference in bond enthalpies for F2 and H2. Do your results
support the statement that “reactions that release heat become more exothermic as
the reactant bonds become weaker”.
5
The energy consumed in breaking all of the bonds of one mole of graphite to make
one mole of C atoms is 717kJ.
C(graphite)  C
ΔH=+717kJ
Construct an energy diagram for the reaction
C(graphite) + 2H2  CH4
and compute the enthalpy of reaction.
6