E10 ENTHALPY OF REACTION
Dissolution, acid/base reactions
Forms of energy
Energy can be stored within any substance in a number of different ways, including as translational, rotational, and
vibrational energy associated with the particles of that substance. The temperature of any substance will reflect the level
of stored energy within the molecules of that substance. Given that no chemical reaction takes place, any increase in the
temperature of that substance corresponds to an increase in the stored energy. The temperature change observed in a
substance that absorbs (stores) a certain amount of energy will be determined by the heat capacity of that substance,
which is defined as the amount of energy required to raise the temperature by 1 K.
Enthalpy
Enthalpy, H, is a property of matter related to its total energy. Enthalpy is an extensive property, meaning that its value
depends upon the amount of the substance being examined.
Besides being an extensive property, enthalpy is also a function of state. That is, its value changes when the
environmental conditions (temperature and pressure) change. To compare the enthalpy of one substance with another,
we need to choose standard values for the following: amount and pressure. If these are set at 1.00 mole and 101 kPa, the
value of the enthalpy is called the standard enthalpy of the substance and is designated as H°.
Systems and changes in enthalpy
A "system" may consist of a given amount of a single substance (for example, a beaker containing some water) or it
may be composed of more than one substance (for example, 0.10 mole of potassium nitrate in 200 mL of water). The
state of a system is defined by its temperature and pressure. A particular combination of temperature and pressure may
be designated as State 1, and it associated enthalpy as H1. If a change occurs in the temperature, the pressure, or both,
the system is in a new state which may be designated as State 2 with an associated enthalpy H2. The change from State
1 to State 2 is thus accompanied by a change in enthalpy from H1 to H2. This enthalpy change is designated as ΔH,
where ΔH = H2 - H1, the difference between the final and initial states. Depending on whether the system gained or lost
energy during the change from State 1 to State 2, ΔH may be positive, zero, or negative.
A change where ΔH is positive is called endothermic and when ΔH is negative is called exothermic.
It is not possible to measure absolute values of H, but the change in enthalpy, ΔH, that occurs when changing states is
measurable. This arises because we can measure changes in temperature and we can relate temperature changes to
changes in enthalpy. For a particular amount of a substance then
ΔH = (molar heat capacity) x (moles of substance) x (ΔT)
or
ΔH = (specific heat capacity) x (grams of substance) x (ΔT)
The molar heat capacity is the heat capacity of 1 mole of a substance whilst the specific heat capacity is the heat
capacity of 1 gram of a substance.
Hess' law
If the change from State l to State 2 (accompanied by an enthalpy change ΔH) is carried out in a series of small steps
(eg. Assuming pressure of the system remains constant, if a single temperature increase of 30 K is replaced by the
following series of steps - raise temperature 15 K, raise temperature 10 K, raise temperature 5 K), the enthalpy changes
accompanying these steps are each smaller than ΔH. Measurements show that the sum of these small enthalpy changes
is always equal to ΔH (Hess's Law). Or, in other words, regardless of how many intermediate states were passed
through to get from State 1 to State 2, the overall enthalpy change remains the difference in enthalpy between States 1
and 2.
Isothermal systems
Isothermal systems are systems where any changes that occur do not change the temperature (T remains constant).
Since energy must be conserved, for an isothermal system, an increase in enthalpy (ΔH > 0) must be accompanied by
absorption of an equal quantity of heat energy from the surroundings of the system (so that T will remain constant).
Because the surroundings are unlimited in extent, they can provide an unlimited supply of heat energy without suffering
any observable fall in temperature. On the other hand, when ΔH < 0, heat energy flows to the surroundings. A
thermostat is a device for providing a constant temperature environment in which to study isothermal changes.
Isolated systems
An isolated system is a system where there can be no exchange of heat energy with the surroundings. In such a system,
any chemical reaction that results in a change in the stored energy would be reflected by a change in temperature of the
system. For example, the energy change of a system which experiences a decrease in enthalpy (exothermic reaction)
results in an increase in the temperature of the system. Conversely, an endothermic change is accompanied by an
increase in the energy of the system and a drop in temperature.
The temperature change accompanying a isolated process, combined with the heat capacity of the system, can provide
an exact measure of the enthalpy change occurring. Changes taking place in a isolated system are called adiabatic. A
calorimeter is a device for measuring temperature changes in an isolated system.
The Calorimeter
For this practical exercise, a simple calorimeter will be used to estimate the enthalpy changes resulting from several
chemical reactions under isolated conditions. Our calorimeter (which you will construct) is a simple device consisting
of a 250 mL glass beaker, with a thermometer and a glass stirring rod. The volume of solution used is 200 mL, which is
in contact with most of the beaker and stirrer, but only a small part of the thermometer. The observations consist of
temperature readings before and after some change in the system, these readings being used to deduce enthalpy
changes. Because some heat is lost to, or gained from, the calorimeter, allowance must be made for the heat capacity of
the calorimeter. Both the enthalpy of the reaction and the allowance for the calorimeter are calculated in exactly the
same way.
More details are given in the example section below.
Units
Throughout these experiments the quantities of heat are conveniently expressed in calories (mainly because the specific
heat capacity of water is defined as 1 cal K-1g-1). For general and theoretical purposes, the final results of each
experiment should be converted to the SI unit of energy, the joule, where l calorie = 4.18 joule.
EXAMPLE 1
Experimental procedure
This exercise involves measurements of enthalpy changes, namely heats-of-reaction. In particular, the heats-of-reaction
are measured for dissolution processes, acid/base reactions, and combinations of these two. In order to simplify the
procedure, a basic calorimeter is used to isolate the system from the surroundings, thus allowing the experimental
observations to consist of temperature readings before and after some change in the system. That is, the changes in the
temperature readings are used as a measure of the enthalpy changes that occur during these reactions.
The Calorimeter
The simple calorimeter used in these experiments consists of a 250 mL glass beaker, a glass stirring rod, and a
"standard" volume of solution (for these experiments, a total of 200 mL of solution was used). A 400 mL glass beaker,
some paper, and a thermometer complete the equipment required.
The calorimeter is constructed by insulating the 250 mL beaker from the
"surroundings" by wrapping it with the paper so that its entire outer surface is
covered. The 400 mL beaker is used to hold this paper against the smaller beaker
by slipping the paper wrapped 250 mL beaker inside the 400 mL beaker.
Throughout the duration of each experiment, the solution will be in contact with most of the inner beaker and stirrer, but
only a small part of the thermometer. Thus, the actual components of the calorimeter are really only the 250 mL beaker
and the glass stirring rod. Consequently, the mass of the calorimeter is taken to be the sum of the masses of these two
items.
Adiabatic or isothermal?
Each experiment is carried out in an insulated glass beaker surrounded by air. Thermal equilibration between this
system and the surroundings is slow, so for short time intervals the system may be regarded as thermally isolated and
thus adiabatic. Therefore, it is possible to measure the enthalpy change accompanying a reaction by obtaining the
difference in the temperature of the contents of the beaker before and after the change in the system.
Sources of error
Likely sources of experimental error in this experiment include improper mixing, the placement of the thermometer
bulb onto slowly dissolving solids, incorrect reading of the thermometer, not enough thermometer readings, and
spillages. The most probable sources of error are the first two. In each case layering of the solution occurs so that the
temperature of the region around the thermometer bulb is not representative of the entire calorimeter. Consequently,
good mixing of the solutions (with a stirring rod and not the thermometer) is important.
Calculating the total heat capacity of the calorimeter
For a typical calorimeter, the heat capacity can be calculated as follows:
Specific heat capacity of glass = 0.84 J K -1 g -1
Mass of beaker + stirrer = 92 g (example only)
Total heat capacity of calorimeter = 92 x 0.84 = 77 J K-1
In a typical experiment the chemical change in the solution (200 mL) was accompanied by a +6.0 K change in
temperature. The corresponding enthalpy change is calculated as follows:
Total heat change = heat change in calorimeter + heat change in solution
Heat change in calorimeter = calorimeter heat capacity x temperature change = 77 x 6.0 = 462 J
Given the density of dilute water solutions as 1.0 g mL-1 and the specific heat of dilute water solutions as 4.18 J K -1 g-1,
then
Heat change in solution = mass of solution x specific heat of x temperature change = 200 x 4.18 x 6.0 = 5016 J.
Total heat change = 462 + 5016 = 5478 J
that is, approximately = 5.5 x 10 3 J = 5.5 kJ. If the above chemical change had been carried out isothermally, the
accompanying enthalpy change of the system would have been -5.5 kJ.