NAME______SECTION______

THE DETERMINATION OF THE ENTHALPY OF FORMATION OF MgO USING HESS’S LAW

HOWARD R. LEO, Ph.D.

Hess’s Law is used to determine the enthalpy of a reaction when it is either very difficult or impossible to measure it directly. According to Hess’s Law, if two or more reactions are added together the ΔH of the resulting reaction is the sum of the ΔH’s of the individual reactions. For example, it would be impossible to measure directly the ΔHf° of CO(g), since there is no way to insure that the combustion of carbon would produce pure CO(g). However carbon can be made to combust completely to CO2(g). In addition CO(g) can be isolate and then combusted to CO2(g). These two reactions can be add together, along with their ΔH’s to find the ΔHf° of CO(g).

C(s) + O2(g) → CO2(g)ΔH1 = ΔHf° of CO2(g)

CO2(g) → CO(g) + ½O2(g)ΔH2 = –ΔH°comb of CO(g)

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C(s) + ½O2(g) → CO(g)ΔHf° = ΔH1 + ΔH2

In this experiment you will determine the ΔHf° of MgO(s):

Mg(s) + ½O2(g) → MgO(s)

To do so you will measure the ΔH° of the following reactions:

Mg(s) + 2H+(aq) + 2Cl‾(aq) → Mg+2(aq) + 2Cl‾(aq) + H2(g)

and

MgO(s) ) + 2H+(aq) + 2Cl‾(aq) → Mg+2(aq) + 2Cl‾(aq) + H2O(l)

Using the ΔH°’s of these two reactions and the standard enthalpy of formation of H2O(l) (ΔHf° = –285.9 kJ/mol) you will be able to calculate the ΔHf° of MgO(s) from Hess’s Law.

Experimental Procedure 1) Wear your safety glasses!

2) If you get any acid on yourself

immediately wash it off thoroughly.

3) Work in the HOOD with the reactions.

4) When adding Mg or MgO to the

acid, keep your face away from the cup.

......

A. Reacting Mg With HCl

1. In the HOOD, set up your styrofoam cup calorimeter and thermometer as shown in the diagram below.

2. Tare a piece of weighing paper on the top-loading balance (i.e. zero in the balance with the paper on it). Add about 0.12 grams of magnesium to the weighing paper on the balance. Record the mass of the Mg to a hundreth of a gram on line 1A) of the first data page, page 5.

3. Measure out about 60 mL of 6.00M HCl in a dry graduate cylinder. Record the amount used on line 2A of the first data page. Do not rinse the cylinder at this time. It will be used again in Part B.

4. Add the HCl to the calorimeter, and begin recording its temperature at one minute intervals, up to 4.0 minutes. Be sure to gently swirl the cup during the entire 20 minutes of the experiment. Record the temperatures to the nearest tenth of a degree on the second data page, page 6

5. At the 5.0 minute mark, carefully and slowly (about 10 seconds) add the Mg to the HCl. DO NOT put your face near the cup while adding the Mg.

6. Take a temperature reading at 6.0 minutes, and continue to read the temperature at one minute intervals up to the 20 minute mark. Record the temperatures to the nearest tenth of a degree on the second data page.

7. Rinse and dry both the cup and thermometer. They will be used for the next part of the experiment.

B. Reacting MgO With HCl

Follow the exact same procedure as in Part A. Only, this time use about 1.3 grams of MgO. Be careful, DO NOT INHALE the MgO. It is a very finely divided solid, and is easily airborne. As before, you must record the mass of the MgO (line 1B), the volume of the HCl (line 2B), and your temperature data in column 2 of the second data page.

Calculations

The quantity of heat absorbed by the solution, Q, from the reactions is calculated from the change in temperature (ΔT) of the solution, its mass (m), and its specific heat (S.H.)

Q = S.H. x m x ΔT(1)

Qreaction = –Qsolution(2)

The ΔH° of the reaction is determined by dividing the Q of the reaction by the number of moles of the limiting reagent. In Part A the limiting reagent is the Mg. in Part B it is the MgO.

ΔH° = Qreaction/n(3)

The mass of the solution is the mass of the HCl used plus the mass of the other reactant. (In the case of reacting Mg with HCl, H2(g) was lost, but the mass of the H2(g) is small enough to ignore.)

msolu = mHCl + mreactant(4)

The mass of the HCl can be determined from its volume and its density (1.09g/mL)

D = m/V(5)

The specific heat, S.H., of both solutions is estimated to be 3.80J/g °C.

1. Determine the ΔT, for both reactions. To do this the temperature data must be plotted on a graph. Since there are two reactions, you must make two separate graphs. Put time on the x-axis, and temperature on the y-axis. The initial temperature (Ti) is determined from the acid readings. Extrapolate the best fit line to the 5.0 minute mark (See abbreviated example below.) Record Ti on line 3 of the first data page.

To determine the final temperature (Tf), extrapolate the mixture’s temperature backwards from 20 minutes to 5.0 minutes. Be sure to draw the best fit line from the data points. Record Tf on line 4 of the first data page.

Calculate the ΔT of the solution, line 5.

T

e

m

p

(°C)

time (min)

2. Determine the mass of the HCl used for each reaction with equation(5), line 6.

3. Determine the mass of the solution resulting from each reaction using equation(4), line 7.

3. Calculate the quantity of heat, Q, absorbed by the HCl + Mg solution using equation (1). Use 3.80J/g°C for the value of the S.H. of the solution. Report this on line 8A.

4. Calculate the quantity of heat, Q, absorbed by the HCl + MgO solution using equation (1). Use 3.80J/g°C for the value of the S.H. of the solution. Report this on line 8B.

5. Calculate the moles of Mg and MgO used. Record the number of moles on lines 9A and 9B respectively.

6. Calculate the ΔH° of the Mg + HCl reaction using equation (3). Remember Qreact =

–Qsolution. Record the ΔH° on line 10A in kilojoules per mole.

7. Calculate the ΔH° of the MgO + HCl reaction using equation (3). Record the ΔH° on line 10B.

8. Calculate the the ΔHf° of MgO(s). In the space provided on page 5, write the three reactions in the appropriate directions, so that they add up to: Mg(s) + ½O2(g) → MgO(s). Also write next to each equation its ΔH° with the appropriate algebraic sign. Add the equations and their ΔH°’s to find the enthalpy of formation of MgO(s).

First Data Page

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Mg(s) + HCl(aq) Reaction

1A) Mass of Mg (grams) ______

2A) Volume of HCl (mL) ______

3A) Initial Temperature (°C) ______

4A) Final Temperature (°C) ______

5A) Change in Temp. (°C) ______

6A) Mass of HCl (grams) ______

7A) Mass of Solution (grams) ______

8A) Q of the Solution (J) ______

9A) Moles of Mg used ______

10A) ΔH° of reaction (kJ/mol) ______

CALCULATION OF THE ΔHf° of MgO(s):

Show Your Work.

MgO(s) + HCl(aq) Reaction

1B) Mass of MgO (grams) ______

2B) Volume of HCl (mL) ______

3B) Initial Temperature (°C) ______

4B) Final Temperature (°C) ______

5B) Change in Temp. (°C) ______

6B) Mass of HCl (grams) ______

7B) Mass of Solution (grams) ______

8B) Q of the Solution (J) ______

9B) Moles of MgO used ______

10B) ΔH° of reaction (kJ/mol) ______

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Second Data Page

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Mg(s) + HCl(aq) Reaction

Time (Minutes)Temperature (°C)

0.0______

1.0______

2.0______

3.0______

4.0______

5.0 MIX

6.0______

7.0______

8.0______

9.0______

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16.0______

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18.0______

19.0______

20.0______

MgO(s) + HCl(aq) Reaction

Time (Minutes)Temperature (°C)

0.0______

1.0______

2.0______

3.0______

4.0______

5.0 MIX

6.0______

7.0______

8.0______

9.0______

10.0______

11.0______

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18.0______

19.0______

20.0______

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QUESTIONS

1. Define precisely the heat of formation of a compound.

2. In this experiment, if a student had measured an incorrectly high ΔT for the Mg + HCl reaction, would the ΔHf° of MgO(s) be too exothermic, not exothermic enough, or unaffected? Explain your answer.

3. In this experiment, if a student had measured an incorrectly high ΔT for the MgO + HCl reaction, would the ΔHf° of MgO(s) be too exothermic, not exothermic enough, or unaffected? Explain your answer.

4. Prove that Mg is the limiting reagent if 2.00 grams are added to 60.0 mL of 6.00M HCl (Remember to balance the equation.)

5. If 2.00 grams of MgO are added to 60.0 mL of 1.00M HCl, which substance is the limiting reagent? (Remember to balance the equation.)

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