# Enthalpy and thermochemical equations

ENTHALPY, HESS’ LAW, AND
THERMOCHEMICAL
EQUATIONS
Quick Review of Concepts


We have been introduced to heat producing
(exothermic) reactions and heat using (endothermic)
reactions
Heat is a measure of the transfer of energy from a
system to the surroundings and from the surroundings
to a system
Enthalpy

The heat given off or absorbed during a
reaction (change in heat of a system) is called
the change in enthalpy (ΔH) when the pressure
of the system in kept constant
DH = H (products) – H (reactants)
More on Enthalpy
Hproducts &lt; Hreactants
DH &lt; 0
Hproducts &gt; Hreactants
DH &gt; 0
6.3
Calorimetry


We measure the transfer of heat (at a constant
pressure) of a chemical reaction by a technique
called calorimetry
In calorimetry:
 The
heat released by the system is equal to the heat
absorbed by its surroundings
 The heat absorbed by the system is equal to the heat
released by its surroundings

The total heat of the system and the surroundings
remains constant
 First
Law of Thermodynamics
Calorimetry


We use an insulated device called a calorimeter to
measure this heat transfer
A typical device is a “coffee cup calorimeter”
 Reaction
is open to the atmosphere
 Therefore,
q
constant pressure
= ∆H at constant pressure
Calorimetry

To measure ΔH for a reaction using a
coffee-cup calorimeter:
1. Dissolve the reacting chemicals in known
volumes of water
2. Measure the initial temperatures of the
solution
3. Mix the solutions
4. Measure the final temperature of the
mixed solution
Calorimetry





The heat generated by the reactants is absorbed by
the water
We know the mass of the water, mwater
We know the change in temperature, ∆Twater
We also know that water has a specific heat of
Cwater = 4.18 J/&deg;C-g.
We can calculate the heat of reaction by:
qsys = ∆H = −qsurr
qrxn = ∆H = −qwater
qrxn = ∆H = -(mwater &times; Cwater &times; ∆Twater)
Practice Problem
•
When 25.0 mL of water containing 0.025 mol of HCl at 25.0&deg;C is
added to 25.0 mL of water containing 0.025 mol of NaOH at 25.0&deg;C in
a coffee cup calorimeter, a reaction occurs. Calculate ∆H (in kJ) of one
mole of water during this reaction if the highest temperature observed is
32.0&deg;C.
•
Note:
•
•
Assume the densities of the solution are 1.00 g/mL
Assume all heat was absorbed by water
Summary




Heat is a measure of the transfer of energy from a system to
the surroundings and from the surroundings to a system
The change in heat of a system is called the change in
enthalpy (ΔH) when the pressure of the system in kept
constant
We measure the transfer of heat (at a constant pressure) of
a chemical reaction by a technique called calorimetry
We use an insulated device called a calorimeter to measure
this heat transfer
Thermochemical Equations
Is DH negative or positive?
System absorbs heat
Endothermic
DH &gt; 0
H2O (s)
H2O (l)
DH = + 6.01 kJ
6.3
Thermochemical Equations
Is DH negative or positive?
System gives off heat
Exothermic
DH &lt; 0
CH4 (g) + 2O2 (g)
CO2 (g) + 2H2O (l)
DH = -890.4 kJ
Thermochemical Stoichiometry

Enthalpy is STOICHIOMETRIC!

Recall that the stoichiometric coefficients always refer
to the number of moles of a substance
H2O (s)
H2O (l)
DH = 6.01 kJ
You would read the above equation as:
“6.01 kJ are absorbed for every 1 mole of ice that melts at
00C and 1 atm”

CH4 (g) + 2O2 (g)
CO2 (g) + 2H2O (l)
DH = -890.4 kJ
“890.4 kJ are released for every 1 mole of methane that is
combusted at 250C and 1 atm”
Practice Problem
How much heat is evolved when 266 g of white
phosphorus (P4) burn in air?
P4 (s) + 5O2 (g)
266 g P4
x
1 mol P4
123.9 g P4
P4O10 (s)
DH = -3013 kJ
x 3013 kJ = 6470 kJ
1 mol P4
6.3
Let’s Try Some More!
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Page 165
Manipulating ∆H
•
If you reverse a reaction, the sign of DH changes
H2O (l)
•
H2O (s)
DH = -6.01 kJ
If you multiply both sides of the equation by a factor n,
then DH must change by the same factor n
2H2O (s)
2H2O (l)
DH = 2 x 6.01 = 12.0 kJ