CHEM 151
ENTHALPY OF FORMATION OF MgO
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FALL 2008
Name __________________________
Stamp Here
Prelab attached (p 8-9)
Partner _________________________
Lecture instructor _________________
Date ___________________________
INTRODUCTION
Chemical reactions either produce heat as they proceed (exothermic) or require heat to proceed
(endothermic). Very sophisticated and expensive pieces of equipment, called calorimeters, can
accurately measure the amount of heat energy associated with a reaction by isolating the reaction in
a well insulated container and measuring either the rise or fall in temperature. This temperature
change, when multiplied by the heat capacity of the system, provides a measure of the energy change
associated with converting the reactants to products.
An expensive calorimeter can measure temperature changes of less than 1°C, a precision required for
some reactions. However, many reactions produce sufficient amounts of heat to be measured using
a much simpler system. In the experiment today, you will use a very simple and inexpensive
calorimeter to indirectly measure the heat of formation of MgO.
Discussion
By using various thermochemical measurements and applying Hess's Law, one can derive basic
information about a chemical reaction in more than one way. It is possible, for instance, to directly
measure the heat of formation (∆Hf) of MgO simply by burning magnesium in oxygen and
measuring the heat evolved in the reaction.
Mg (s) +
1
O 2 (g) ! MgO (s) + heat
2
(1)
This simple reaction, however, requires a rather elaborate calorimeter to perform the measurements.
We therefore need to find a simpler method that will provide the same information.
Hess's Law states that for any process that can be defined as the sum of several stepwise processes,
the value of ∆H for the whole process must be equal to the sum of the ∆H values for the individual
steps. Therefore, if you dissolve Mg in acid and then dissolve MgO in the same acid, the difference
between these two heats of reaction will be the heat of formation of MgO from its elements,
providing you account for all side reactions that occur. The only side reaction involves the
formation of water. If you can find (try your textbook!) the ∆Hf for the reaction
H2 (g) + ½ O2 (g)
→ H2O (l)
ΔHC
then Hess’s Law can be used to determine the heat of formation (∆Hf) for MgO, reaction (1).
You will experimentally determine the heat of reaction for the following reactions A and B.
a)
Mg (s) + 2 HCl (aq) →
b)
MgO (s) + 2 HCl (aq) → MgCl2 (aq) + H2O (l)
∆HB
c)
H2 (g)
∆HC (text appendix)
#7 ∆Hf of MgO
+ ½ O2 (g)
MgCl2 (aq) + H2 (g)
→ H2O (l)
Rev F08AEM Fall 2008
∆HA
Page 1 of 9
In this experiment you will determine experimentally the heats of reaction for Mg with HCl and
MgO with HCl. Then, by appropriately manipulating the data gathered, you will determine the heat
of formation of MgO.
The calorimeter you will be using is a thermally insulated container (two Styrofoam cups) with a
cardboard cover in which a reaction can be conducted, accompanied by a thermometer that can
measure the temperature change of the system (see Figure 1).
Figure 1: A Styrofoam Cup Calorimeter
You will determine if the reactions run in lab are exothermic or endothermic. Since both reactions
are run in dilute water solutions of HCl, the heat produced by the reaction with the HCl (the reaction,
or the “system”) will be transferred to the water (the “surroundings”). By knowing the heat capacity
of water, (assuming that the heat capacity of water and that of dilute HCl are the same, which is
reasonable) the heat of reaction can be calculated. Some heat will also be transferred to the
calorimeter whose heat capacity is unknown. However, using the double cup technique, this heat
transfer to the cups is assumed to be minimal and therefore, should not affect the final determination
of your determination of ΔH.
We know that when a reaction occurs, if the heat is lost by the system (the reaction) then the heat
will be gained by the surroundings (in this case, the water). Conversely, if the heat is gained by the
system (the reaction), then the heat will be lost by the surroundings (the water). Therefore, the q
values will be numerically equal but they will be opposite in sign.
heat system(rxn) = -heat surroundings(H2O)
m1C1ΔT1 = -(m2C2ΔT2)
m = mass; C = specific heat; ∆T = temperature change = (Tf – Ti)
when a reaction is run under a constant pressure, then q = ΔH
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 2 of 9
EXPERIMENTAL PROCEDURE
All chemicals and supplies can be obtained from the reagents bench. In all parts, try to keep the
total volume of solution at 100 mL.
C H 2O = 4.184
J
.
goC
1 calorie = 4.184 Joules
Determination of ΔH values:
In order to calculate the heats of reaction for Mg and MgO with HCl, we must measure the heat
transferred to the water. We know:
Energy released or absorbed = m H O iC H O i!T (3)
2
2
This formula would tell us how much heat was produced by some amount of Mg or MgO reacting
with HCl. Specific heats are typically given in J/g•°C; we, however, are interested in knowing the
heat of reaction in terms of kJ/mole Mg or MgO. We will reconcile units at the end of our
calculations. Assume that the specific heat of your solution is the same as that for water (it is
very close).
A. Procedure for Determining Heat of Reaction of Mg with HCl
First rinse and dry your calorimeter setup (two cups, nestled together – referred to as CAL), and
weigh them together to the nearest 0.01g; record this mass in the table on the next page. You must
use these two cups for both parts of the experiment. Then place about 100 mL of 2M HCl into the
inner cup. Weigh, the CAL and HCl solution, to the nearest 0.01 g and record this mass in the table.
Subtract CAL weight to determine mass of the HCl used and record this in the table. Cover the CAL
and insert the thermometer. While this is coming to thermal equilibrium, obtain about 0.50 g of Mg
turnings and weigh them accurately to the nearest 0.01 g and record this mass in the table.
Use the Fume Hood for this reaction!!
Now read the temperature of your CAL + HCl. Record this initial temperature of the HCl solution to
the tenths place and record in the table on the next page. Drop in the weighed magnesium, replace
the cover and read the temperature at half minute intervals (stirring constantly) until three
successive readings show no longer show a change in temperature. All of the Mg should react.
From this reaction, the temperature either went up or it went down. In the table, record the highest
(or lowest) temperature reached – record this Tfinal. Place the spent reaction mixture in the waste
container in the hood.
NOTE: record all masses to the hundredths place and all temperature readings to the tenths place in
the table on the next page.
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 3 of 9
DATA TABLE: ALL MASSES SHOULD BE TO THE HUNDREDTHS PLACE, ALL
TEMPERATURE READINGS SHOULD BE TO THE TENTHS PLACE!!
Mass of CAL
Mass of CAL + HCl solution
Mass of HCl solution
Mass of Mg
Initial Temp of HCl solution
Final Temp of HCl + Mg
∆T
Calculations
Because the HCl is such a dilute solution: assume that the mass of HClsolution = mass of H2O
Energy = q = m H O iC H O i!T
2
•
2
Calculate the amount of heat (q) in Joules. Show your work and put units on your answer.
q= _____________________
•
Remember that the energy calculated, q, is also for the mass of Mg that you used. Use
conversions to calculate the value in kJ/mol Mg. Show your work and units in your answer.
q = ______________________
•
Is the reaction exothermic or endothermic?
•
Therefore, what is ∆H for the reaction?
Mg (s) + 2 HCl (aq) → MgCl2 (aq) + H2 (g)
______________________________
∆HA = _______________________
Remember, use the correct sign for ∆H, that is consistent with your answer to the previous question!
Instructor initials for value of ∆HA:
(note: you must get initials before moving on)
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 4 of 9
B. Procedure for Determining Heat of Reaction of MgO with HCl
You will be following the same procedure for this experiment, EXCEPT you will use MgO. Place
100 mL of 2M HCl (measured from the same stock used in the above experiment) into the clean, dry
CAL. Weigh as before to determine mass of HCl and proceed with MgO. Calculate and weigh out
accurately an amount of MgO that is equimolar (equal in moles) to the amount of magnesium
used above. Show this calculation in your work and weigh the MgO on weighing paper to ±0.01 g.
Record the temperature of the 2M HCl in the CAL. Add the MgO to the acid (stirring constantly)
and record the highest/lowest temperature reached. This could take up to 20 minutes as the
temperature changes more slowly in this reaction than in the reaction of Mg with HCl. When you
are finished, place the spent reaction materials in the waste container in the hood.
•
Calculation for determining amount of MgO needed:
Mass MgO = _____________ g
Instructor initials for MgO
amount
__________________________
DATA TABLE: REMEMBER TO RECORD YOUR MASSES TO THE HUNDREDTHS
PLACE AND YOUR TEMPERATURE TO THE TENTHS PLACE!
Mass of CAL
Mass of CAL + HCl solution
Mass of HCl solution
Mass of MgO
Initial Temp of HCl solution
Final Temp of HCl + MgO
∆T
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 5 of 9
Calculations
Again, because the HCl is a dilute solution, assume that the mass of HClsolution = mass of H2O
Energy = q = m H O iC H O i!T
2
•
2
Calculate the amount of heat (q) in Joules. Show your work and put units on your answer.
q= _____________________
•
Remember that the energy calculated, q, is also for the mass of MgO that you used. Use
conversions to calculate the value in kJ/mol MgO. Show your work and put units on your
answer.
q = ______________________
•
Is the reaction exothermic or endothermic?
•
Therefore, what is ∆H for the reaction?
MgO (s) + 2 HCl (aq) → MgCl2 (aq) + H2O (l)
______________________________
∆HB = _____________________
Remember, use the correct sign for ∆H, that is consistent with your answer to the previous question!
Instructor initials for value of ∆HB:
(note: you must get initials before moving on)
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 6 of 9
HESS’ LAW CALCULATION
Using the information determined above and Hess’s Law, the heat of formation (∆Hf) for MgO can
be obtained.
• Show how you can calculate the heat of formation of MgO using the equations A, B, and C
from the first page. Hint: set-up reactions so that the sum is the formation equation for MgO.
Perform the calculation to determine ∆Hf MgO. PAY ATTENTION TO SIG FIGS!
A.
Mg (s) + 2 HCl (aq) →
B.
MgO (s) + 2 HCl (aq) → MgCl2 (aq) + H2O (l)
∆HB = ___________________
C.
H2 (g)
∆HC = ___________________
+ ½ O2 (g)
MgCl2 (aq) + H2 (g)
→ H2O (l)
∆HA = ___________________
Rearrange the equations above to determine ΔH for: Mg (s) + ½ O2 (g)  MgO (s): show your
work!
Your value for ∆Hf MgO (pay attention to sig figs!) = ____________________
•
Look up the theoretical (textbook) value for the heat of formation of MgO and calculate your
percent difference.
Textbook value for ∆H = ______________ kJ/mol
% difference = ____________________
•
List at least two sources of error (other than human error) that could lead to a % difference
between your calculated and the actual value. Also explain how they would lead to a higher
or lower value for ∆Hf MgO.
a.
b.
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 7 of 9
Stamp:
Prelab Exercises
1. Define exothermic and endothermic:
Exothermic:
Endothermic:
2. When we look at the ΔH for a reaction, how can we identify (simply!) if the reaction is exo or
endo thermic?
3. The amount of heat, q, obtained in a reaction of 0.10 mole HCl(aq) with excess NaOH(aq) is 96
calories.
a. Calculate the heat in terms of the number of Joules (note: 1 calorie = 4.184 Joules)
q = ________________
b. Calculate q in kJ/mol HCl.
q = ________________
c. If the reaction gives off the heat, is the reaction endothermic or exothermic?
_______________________
d. Therefore, what is ∆H for the reaction in kJ/mol?
_____________________kJ/mol
Remember, use the correct sign for ∆H, that is consistent with your answer to the previous question!
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 8 of 9
4. On page 2 of the lab, there is an explanation about how qrxn will be determined. Read more about
this and answer the following questions about the reactions you will be performing in lab: Circle
the appropriate answer:
a.
If the temperature of the water increases, did the water absorb heat or release heat?
Absorbs heat
Releases heat
b. Therefore, will qH2O be a +q or –q?
+q
-q
c. If the temperature of the water increases, then did the reaction absorb heat of release heat?
Absorbs heat
Releases heat
d. How are qH2O and qrxn related to one another?
-
They are exactly the same
-
They are numerically the same with opposite charges
-
They have nothing to do with one another
e. Therefore, will the qrxn be +q or –q?
+q
-q
5. Given the following thermochemical equations, using Hess’ law, calculate ∆H for the
decomposition of one mole of acetylene (C2H2 gas) into its elements in their stable/natural state
at 25˚C and 1 atm pressure.(hint: write the balanced chemical reaction for the decomposition of
1 mole of acetylene into its elements!)
C2H2 (g) + 5/2 O2 (g) → 2 CO2 (g) + H2O (l)
∆H1 = -1299.5 kJ
C (s) + O2 (g) → CO2 (g)
∆H2 = -393.5 kJ
H2 (g) + ½ O2 (g) → H2O (l)
∆H3 = -285.8 kJ
Balanced equation: ____________________________________________________________
∆Hrxn = ___________________
#7 ∆Hf of MgO
Rev F08AEM Fall 2008
Page 9 of 9