SCH 4U
Name: ____________________
Unit: Thermochemistry and Rates
Date: ____________________
To determine the enthalpy change involved in a particular reaction, we can use
(1) Calorimetry
Although we have not yet used a calorimeter, we will next week. Figure 1, page 309, shows a coffee cup calorimeter like
the one we will use in class to measure temperature change in a system of a given mass. With that data, we can use
q=mcΔT to calculate ΔHr.
(2) Hess’s Law
(3) Tables of Standard Enthalpies of Formation (with or without Hess’s Law)
(4) Tables of Bond Energies
Standard Enthalpies of Formation (pg. 331 – 340)

Chemical equations written as formation equations are always written for one mole of a particular product, which
may be in any state or form. The reactants must be elements in their standard states.
Example:
kJ
2 C(s) + 3 H2(g) + ½ O2(g)  C2H5OH(l) ∆𝐻𝑓𝑜 = – 235.2
(See Table C6, page 799 for the ΔH value.)
mol

The standard enthalpy of formation of an element in standard state is equal to zero. So the enthalpy of formation for
H2(g), Mg(s), etc. will all be zero.
Enthalpies of formation, ΔHf, are available in tables. Your text has one such table, C6, on page 799. (Note that it is
expected that you know that the enthalpy of formation for all elements is zero. These values are never in a table of
enthalpies of formation.)
Standard enthalpies of formation for a series of steps can be used in Hess’s Law to calculate the overall enthalpy
change for a given target reaction. (See the example at the top of page 333.)
The enthalpy change for any given equation can also be calculated from heats of formation in tables using the
following equation:



𝑜
𝑜
ΔH=∑ 𝑛∆𝐻𝑓(𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠)
− ∑ 𝑛 ∆𝐻𝑓(𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠)
Practice
1) Do #1 page 332.
2) Do #2 – 5 page 335.
Using Bond Energy to Calculate Enthalpy Change
Generally speaking, energy is required to break bonds and released when bonds are formed. We can use this information and
tables of bond energies to calculate the overall change in enthalpy for a particular reaction. Two things to keep in mind are:

Tables of bond energies provide average values. Even in methane, where all the C – H bonds look to be identical,
breaking one of these four bonds will change the environment for the other 3 bonds and potentially change the bond
energy associated with them. The use of average bond energy values helps us to be able to do the math but
sometimes introduces a source of error.

Tables of bond energies are compiled for substances in the gaseous state. ONLY. Thus, other manipulations have to
occur if one is faced with substances in the liquid, solid or aqueous states.

Overall:
ΔH=  𝑛∆𝐻°𝑟𝑒𝑞𝑢𝑖𝑟𝑒𝑑 𝑡𝑜 𝑏𝑟𝑒𝑎𝑘 𝑡ℎ𝑒 𝑏𝑜𝑛𝑑𝑠 𝑖𝑛 𝑡ℎ𝑒 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 − ∑ 𝑛∆𝐻°𝑟𝑒𝑙𝑒𝑎𝑠𝑒𝑑 𝑏𝑦 𝑏𝑜𝑛𝑑 𝑓𝑜𝑟𝑚𝑎𝑡𝑖𝑜𝑛 𝑖𝑛 𝑡ℎ𝑒 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
Practice
How much energy is absorbed or released (identify which) when each of the following reactions takes place? For efficiency in
counting up bond types it is strongly suggested that you draw Lewis structures for all substances in each reaction.
1) CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
2) 2 H2 (g) + O2 (g)  2 H2O (g)
3) H2 (g) + Cl2 (g)  2 HCl (g)
4) CH4 (g) + Cl2 (g)  CH3Cl (g) + HCl (g)
The relevant bond energies are:
The answers are:
C—H 414 kJ/mol
C = O 730 kJ/mol
H—H 435 kJ/mol
H—Cl 431 kJ/mol
1. - 656 kJ released
3. - 184 kJ released
O = O 502 kJ/mol
O—H 464 kJ/mol
Cl—Cl 243 kJ/mol
C—Cl 331 kJ/mol
2. - 484 kJ released
4. - 105 kJ released