Enthalpy Changes by Calorimetry
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Experiment 22. Enthalpy Changes by
Calorimetry
Objectives
The aims of the experiment are: (i) to determine the enthalpy change which
accompanies the melting of a solid, and (ii) to determine the enthalpy
change for the formation of a chemical compound by using calorimetric data
and applying Hess' Law.
Introduction
The heat evolved or absorbed when a process occurs at constant
pressure is equal to the change in enthalpy. Since H is defined
by the equation:
H=U+pV
then DH = DU + pDV at constant pressure; where DH represents the
change in enthalpy.
Reactions which occur in unsealed containers in the laboratory, occur
essentially at constant pressure (= atmospheric pressure). Chemical
processes which occur in plants and animals also occur at constant pressure.
This is why enthalpy is such an important thermochemical parameter
for physical and chemical processes. It can be related directly to the
heat evolved or absorbed when the processes occur under "natural" conditions.
When processes occur in a pressure-tight, sealed container, such
as a bomb calorimeter, the heat evolved or absorbed is equal to
the change in internal energy, DU, since the process occurs at
constant volume.
Enthalpy is a state function, and so if one wants to define uniquely
the enthalpy change in a physical or chemical process, one needs to
define only the initial and final states of the system when the process
occurs. For a physical process such as the melting of ice, once the
pure substance is identified and the pressure is specified,
the enthalpy change is uniquely defined.
The value which is now most often quoted for the enthalpy change
in this process, is the molar enthalpy of melting (or "latent
heat" of melting) when the process occurs at a pressure of 1 bar.
(1 bar = 10x5 Pa)
For chemical reactions, one can define a "standard" enthalpy of
reaction by specifying "standard" initial and final states of the
reacting system. The standard enthaly of formation of a
chemical compound, DHf, is the heat evolved or absorbed when the
compound is formed in its standard state from its constituent
elements in their standard states.
The standard state of a substance is defined as the stable form
of that substance at a pressure of 1 bar and a specified temperature.
The standard molar enthalpies of formation of elements are zero at
all temperatures - by definition.
The standard molar enthalpy of fonnation of a compound is
therefore a uniquely defined quantity, DHf(T), and values given in
thermodynamic tables are usually at 298.15 K. These quantities are
useful because they can be used to obtain enthalpy of any
reactions in which the individual compounds are involved.
The heat, Q required to change the temperature of a substance from
Ti to Tf is given by:
Q = mC(Tf - Ti)
where m is the mass whose temperature is changed from Ti to Tf and
C is the heat capacity of the substance. When m is in kg, C is in
J K-1 kg-1, and T is in C or K, the heat is in Joule.
Note that the heat capacity, C, quoted here, bears no indication
of conditions, that is, whether it is Cp or Cv. This is because
only solids and liquids are usually involved in calorimetry at
this level, and Cp and Cv are very nearly the same value for
matter in these "condensed" phases.
I Enthalpy of Melting of Ice
Theoretical Considerations
If ice is mixed with warm water in a calorimeter, the ice will
melt and the water so formed will be raised to a final
temperature, Tf up from 0 C. If one can assume that no heat
enters or leaves the calorimeter, then:
Heat lost by the warm water = Heat gained by the ice and cold
water produced by melting.
mwCw(Ti-Tf) = mice(DHmelt) + miceCw(Tf- 0)
where mw and mice = mass of warm water and ice respectively.
Tf and Ti = final and initial temperatures of the water
Cw = heat capacity of water (= 4.18 JK-1g-1).
DHmelt = enthalpy of melting of ice per unit mass.
Procedure
Heat approximately 200 cm3 of water to about 50C and measure
carefully 100 cm3 of this warm water into the calorimeter. Start
your clock/watch and record the temperature of the water at 1
minute intervals, stirring constantly, until the temerature is
about 40 - 45 C. Temperatures should be read to the nearest
0.1 C, or better if the thermometer allows, and the lid should be
on the calorimeter while you are doing this.
Take 4 ice cubes, shake off the excess moisture, and wipe dry as
quickly as you can, with paper towels. Add them to the calorimeter
exactly on one of the minutes after you have read the water temperature,
taking care not to splash water out of the calorimeter. Continue
stirring and now read and record the temperature every 30 seconds.
Continue reading the temperature for another ten minutes in order that
you have a continuous record of temperature of the water in the
calorimeter for at least fifteen minutes.
Measure carefully, using a measuring cylinder, the final volume of
water in the calorimeter. The difference between the initial and
final volumes, is of course the volume of water obtained as a
result of the melting of added ice. Its mass is the mass of ice
added. (Assume density = 1.0 gcm-3).
Display your data graphically, and obtain Ti and Tf from your
graph following the instructions given by the lab supervisor. Use
the equation developed above to obtain DHmelt of ice in kjmol-1.
II. Enthalpy Change in the Formation of Chemical Compound
Theoretical Considerations
From our definition, the enthalpy of formation of MgO(s) is
the heat produced (or absorbed) when one mole of magnesium solid
reacts with a half mole of oxygen gas, the reactants and products
being in their standard states.
It would be difficult to carry out this process in the laboratory
particularly because a gaseous reagent is involved, but the
difficulty can be avoided by selecting more convenient reactions
for investigation, and combining the results using Hess' Law.
Consider the following reactions:
(a) Mg(s) + 2H+(aq) ---> Mg++(aq) + H2(g) : DHl
(b) MgO(s) + 2H+(aq) ---> Mg++(aq) + H2O(l) : DH2
(c) H2(g) + 1/2 02(g) ---> H2O(l) : DH3
Combination of these equations (a - b + c) results in
Mg(s) + 1/2 02(g) ----> MgO(s) : DHf (MgO)
The enthalpy of formation of magnesium oxide can be obtained from
experimental observation of reactions (a) and (b) and by using
data for the DHf of water from the literature:
DHf(298) of water = -285.8 kJmol-1
Procedure
(a) Determination of DH1.
Make sure your calorimeter is clean and dry. Weigh it empty and
again with about a 10 cm length of clean magnesium ribbon. The
mass should be taken to at least +/- 0.001 g.
Measure out 50 cm3 (to +/- 0.5cm3) of 1 M HCl (density (HCl) = 1.018 gcm-3)
into a measuring cylinder and record its temperature at four one
minute intervals. On the fifth minute pour the HCl solution into the
calorimeter and put the lid on. Insert the thermometer and stirrer quickly
through the lid and continue to take the temperature at 30-second
intervals for about seven minutes after mixing, stirring the mixture constantly.
Display your data graphically and follow the instructions given to
find the initial and final temperatures. Calculate the heat
evolved, using the temperature rise determined above.
Q = M.HCl C.HCl(Tf - Ti) (C.HCl = 4.00JK-1g-1)
Convert this to heat evolved when a mole of magnesium reacts.
This is DH1 (kJ mol-1). Remember this is heat evolved so DH1 is
negative according to the normal convention.
(b) Determination of DH2
Make sure your calo