Experiment 22. Determination of 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 ΔH = ΔU + pΔV at constant pressure;
where ΔH 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, ΔU, 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, ΔHf, 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 formation of a compound is
therefore a uniquely defined quantity, ΔHf(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.
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) ---> Mg2+(aq) + H2(g) : ΔHl
(b) MgO(s) + 2H+(aq) ---> Mg2+(aq) + H2O(l) : ΔH2
(c) H2(g) + 1/2 02(g) ---> H2O(l) : ΔH3
Combination of these equations (a - b + c) results in
Mg(s) + 1/2 02(g) ----> MgO(s) : ΔHf (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 ΔHf of water from the literature:
ΔHf(298) of water = -285.8 kJmol-1
Procedure
(a) Determination of ΔH1.
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.
Graphically display your data 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 ΔH1 (kJ mol-1). Remember this is heat evolved so ΔH1 is
negative according to the normal convention.
Combine the three values ΔHl, ΔH2 and ΔH3, (paying attention to
the signs), to obtain ΔHf(MgO), and determine the uncertainty
(error) in this final value.
Calculate ΔUf of magnesium oxide, assuming that ΔHf is the value
you have found in this experiment.
Comment on the accuracy and precision of the procedure as
pointedly as you can.
Return to Chemistry, UWI-Mona,
Home Page
Copyright © 2006 by Robert John Lancashire,
all rights reserved.
Created and maintained by Prof. Robert J.
Lancashire,
The Department of Chemistry, University of the West Indies,
Mona Campus, Kingston 7, Jamaica.
Created Apr 1996. Links checked and/or last
modified 5th March 2006.
URL
http://wwwchem.uwimona.edu.jm/lab_manuals/c10expt22.html