THERMOCHEMISTRYHeats of Reaction

The Enthalpy of Formation of Magnesium Oxide

Transfer 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 FORMATIONThe 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 tem-perature. 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) ∆H2

(3) Mg2+(aq) + H2O(l) → MgO(s) + 2 H+(aq) ∆H3

(4) 1/2 O2(g, 1 atm) + H2(g, 1 atm) → H2O(l) ∆H4

Applying Hess's law to these equations, we see that

∆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

(5) MgO(s) + 2 H+(aq) → Mg2+(aq) + H2O(l) ∆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

Chemistry 111 Lab: Thermochemistry Page I-3

Structure of magnesium

oxide. MgO has a solid state

structure that consists of a

lattice of oxide ions with Mg2+

ions in the holes in the lat-

tice. See the Models folder

on the General ChemistryNow

CD-ROM.

Revised: June 2003

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.

CALORIMETRYThe enthalpy change, ΔH, for a chemical reaction can be measured conve-niently 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 surround-ings 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 reac-tion 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

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Figure This is a screen shot of an interactive module that allows you to experiment with the chem-

istry 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

Consult your textbook for the value of the enthalpy of formation of liquid water.

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

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) = (per mol) = quantity of reactant (mol)

measuredqq

EXPERIMENTAL PROCEDUREPART I The Relation Between the Quantity of Material Reacting and

the Heat TransferredIn 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 magne-sium 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 calorim-eter.

(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 ther-mometer 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-

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Although it is possible for an individual student to perform this experiment satisfactorily, it is easier if students work in pairs.

ples, making sure the calorimeter is clean and dry each time.

(g) 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.

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Figure Equipment and setup for the thermochemistry experiment.

(a) Calorimeter consists of thermometer, two sty-rofoam cups, paper cover, and beaker.

(b) Two cups together, sup-ported in a beaker. Cover in place with thermometer through the hole in the cover.

(c) Diagram of the experimental setup for the thermochemistry experiment.

Revised: June 2005

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.

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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.

Revised: June 2005

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