Chemistry 111 Lab: Thermochemistry
Page I-3
THERMOCHEMISTRY
Heats of Reaction
The Enthalpy of Formation of Magnesium Oxide
T
ransfer of heat energy often occurs during chemical reactions. A reaction
may transfer heat to its surroundings—it is exothermic. Alternatively,
heat may be transferred from the surroundings to the reaction system—it is
endothermic. In this experiment, you will explore some exothermic reactions
with two objectives in mind:
• To discover the relationship between the quantity of material undergoing
reaction and the quantity of heat evolved.
• To determine the molar enthalpy of formation of a compound, MgO, using
calorimetry and Hess’s Law.
MOLAR ENTHALPY OF FORMATION
The heat absorbed when a mole of compound in its standard state is formed
at a given temperature from the appropriate elements, also in their standard
states, is the molar enthalpy of formation, ΔH˚f, of the compound at that temperature. The molar enthalpy of formation is positive if heat is absorbed when
the compound is formed under these conditions (an endothermic reaction)
and negative if heat is released (an exothermic reaction). In this experiment,
the molar enthalpy of formation of solid magnesium oxide, MgO(s), will be
determined by calorimetry.
The molar enthalpy of formation of MgO(s) is simply the enthalpy of the
reaction
(1) Mg(s) + 1/2 O2(g, 1 atm) → MgO(s)
∆H1
Unfortunately, the enthalpy change for this reaction cannot be measured
directly in a simple experiment. Therefore, we take advantage of the following
sequence of reactions, whose sum is equivalent to reaction 1.
(2) Mg(s) + 2 H+(aq) → Mg2+(aq) + H2(g, 1 atm)
(3) Mg2+(aq) + H2O(l) → MgO(s) + 2 H+(aq)
∆H2
(4) 1/2 O2(g, 1 atm) + H2(g, 1 atm) → H2O(l)
∆H4
∆H3
∆H˚f [MgO(s)] = ∆H1 = ∆H2 + ∆H3 + ∆H4
Thus, to determine the enthalpy of formation of MgO(s) (reaction 1) we shall
have to determine the enthalpies for reactions 2, 3, and 4. This is really quite
straightforward, since ΔH2 can be measured directly in a simple experiment
(Part 1 of the Experiment). Reaction 3 normally proceeds in the direction
opposite to the way it is written above, so it is convenient to measure the
enthalpy of the reaction
∆H5
The enthalpy of reaction 5 is the negative of that for reaction 3 (ΔH5 = –ΔH3),
and it is measured in Part 2 of the experiment. Finally, ΔH4 is the enthalpy of
Revised: June 2003
oxide. MgO has a solid state
structure that consists of a
lattice of oxide ions with Mg2+
ions in the holes in the lat-
Applying Hess's law to these equations, we see that
(5) MgO(s) + 2 H+(aq) → Mg2+(aq) + H2O(l)
Structure of magnesium
tice. See the Models folder
on the General ChemistryNow
CD-ROM.
Page I-4
Consult your textbook for
the value of the enthalpy
of formation of liquid
water.
Chemistry 111 Lab: Thermochemistry
formation of liquid water and is obtained from tables of standard enthalpies.
Consult your textbook for this value.
An excellent way to prepare for this laboratory is to visit the Dartmouth
College web site for their Chemistry Labs. The figure at the bottom of the page
shows you a screen from a “tool” that allows you to piece together the data
needed to calculate the molar enthalpy of formation of magnesium oxide.
CALORIMETRY
The enthalpy change, ΔH, for a chemical reaction can be measured conveniently at a constant pressure under conditions where virtually no heat is
exchanged with the surroundings. For a reaction that takes place in aqueous
solution, the measurement is made by observing the increase or decrease in
temperature that accompanies the reaction when it is carried out in a special
type of vessel (called a calorimeter) that keeps the exchange of heat with the
surroundings to a negligible amount. (In this experiment your calorimeter
consists of two styrofoam cups, nested one inside the other. Styrofoam is an
excellent insulator.) Thus, there is no heat transferred to or from the surroundings when the reaction takes place in this way.
However, the desired enthalpy change applies to the reaction when it is
carried out in such a way that the products end up at the same temperature
that the reactants had to begin with. Thus, it is necessary to find out how
much heat would be absorbed or released if the solution containing the reaction products were returned to the original temperature. This amount of heat
is the only heat transferred in the entire process and, therefore, is the same as
the enthalpy of the reaction. And how is this quantity of heat obtained? It is
See the Dartmouth College
Chemistry Department web
site at
a) http://www.dartmouth.
edu/~chemlab/techniques/
calorimeter.html
b) http://www.dartmouth.
edu/~chemlab/chem3-5/
calor1/overview
/start.html
Figure This is a screen shot of an interactive module that allows you to experiment with the chemistry of this experiment. Please use it to examine the reactions you need to use to determine the
enthalpy of formation of MgO. You can find it at
http://www.dartmouth.edu/~chemlab/info/resources/deltah/deltah.html
Revised: June 2005
Chemistry 111 Lab: Thermochemistry
Page I-5
calculated from the relationship
q=
–(mass of solution)•(specific heat of solution)•(∆T)
where ΔT is the observed “change in temperature.” That is,
∆T = final temperature – original temperature
The specific heat of the solution is the heat required to raise the temperature
of 1.00 gram of the solution 1.00 degree Celsius. It is a positive quantity.
Finally, be sure to notice the negative sign in the equation above. It is there
for the following reason:
If the temperature of the solution is observed to increase (i.e, ∆T
is positive), the reaction is exothermic, and so q (= ∆H) must be
negative. On the other hand, if the temperature is observed to
decrease in the reaction (i.e., ∆T is negative), then the reaction
is endothermic, and q (= ∆H) is positive.
Finally, the heat calculated, as above, corresponds to the specific number of
grams of the reactant that you actually use (in this experiment, either Mg or
MgO). In order to convert this measured value of the heat into molar enthalpy,
it is only necessary to divide the measured heat by the number of moles of the
reactant.
ΔH (per mol) = q (per mol) =
qmeasured
quantity of reactant (mol)
EXPERIMENTAL PROCEDURE
PART I The Relation Between the Quantity of Material Reacting and
the Heat Transferred
In this portion of the experiment we want to explore the relationship between
the quantity of magnesium metal reacting with hydrochloric acid and the heat
evolved by the reaction.
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
(a)
Set up your coffee-cup calorimeter and support it in a ring stand or
beaker as illustrated by your instructor and in the Figure.
(b)
Weigh out, to the nearest 0.001 g, three different portions of magnesium metal, say about 0.2 g, about 0.4 g, and about 0.5 g. Record the
masses in the data table on the report form.
(c)
Place one of the magnesium samples in a clean, dry coffee-cup calorimeter.
(d)
Using a graduated cylinder, measure out as accurately as possible 100.
mL of 1.0 M HCl. Measure the temperature of the HCl solution and
record this initial temperature on the report form. Replace the thermometer in the calorimeter setup.
(e)
Add the HCl solution to the calorimeter and swirl gently but steadily. At
30 second intervals record the temperature (to the nearest 0.5 ˚C) until
the temperature has held constant or decreased for three consecutive
readings.
(f)
Repeat the steps above with another sample of Mg. Do all three sam-
Revised: June 2005
Although it is possible for
an individual student to
perform this experiment
satisfactorily, it is easier if
students work in pairs.
Page I-6
Chemistry 111 Lab: Thermochemistry
(a) Calorimeter consists of thermometer, two styrofoam cups, paper cover, and beaker.
(b) Two cups together, supported in a beaker. Cover
in place with thermometer
through the hole in the cover.
(c) Diagram of the experimental
setup for the thermochemistry
experiment.
Figure Equipment and setup for the thermochemistry experiment.
ples, making sure the calorimeter is clean and dry each time.
(g)
Revised: June 2005
In each case, determine ΔT, the change in temperature between the
temperature of the HCl solution before adding it to the magnesium
and the maximum temperature of the reacting system. If you were to
plot the temperature of the reacting system versus time, you would see
something like that illustrated here.
Chemistry 111 Lab: Thermochemistry
PART II: Heat of Reaction of Magnesium Oxide with Hydrochloric
Acid
(a)
Set up the calorimeter as in PART I. Make sure it is clean and dry.
(b)
Weigh out 0.7 g of MgO to the nearest 0.001 g and place the powder
in the calorimeter.
(c)
Repeat steps (d) and (e) as in PART I above with another sample of
magnesium oxide.
Revised: June 2005
Page I-7
Note: One problem here
is that the MgO tends to
stick on the paper. The
best way to avoid an error
is to weigh the MgO on a
small piece of paper on the
balance in such a way that
you know the net weight of
the MgO. Place both the
paper and the MgO powder
in the calorimeter.
Page I-8
Revised: June 2003
Chemistry 111 Lab: Thermochemistry