Download Thermochemistry Thermodynamics is the study of energy and its

Thermochemistry
Thermodynamics is the study of energy and its transformations.
Thermochemistry is the branch of thermodynamics that is concerned with
the study of energy changes that accompany chemical reactions. According to
the First Law of Thermodynamics, energy can neither be created nor destroyed,
but it can be converted from one form to another. In a chemical reaction, the
energy stored in chemical bonds may be converted into other forms of energy
such as heat and light. In a chemical reaction, bonds in the reactant molecules
are broken and new bonds in the product molecules are formed. Bond breakage
requires an input of energy, whereas bond formation results in a release of
energy. If the energy released in bond formation exceeds the energy absorbed
in bond breakage, then the reaction occurs with a net release of energy. In
many reactions this energy appears in the surroundings as heat, symbolized by
q. For example, the combustion of natural gas produces heat energy. The
reaction is:
CH4 (g) + 2O2 (g) →CO2(g) + 2H2O(l)
(1)
A chemical reaction can also occur with the absorption of heat from the
surroundings. In this case, the energy required for bond breakage exceeds the
energy released in bond formation. The heat released or absorbed by a
chemical reaction at constant pressure is the enthalpy, ΔH. Where
ΔH = Enthalpy of Products - Enthalpy of Reactants (2)
The enthalpy of a substance is a measure of its stored chemical energy. By
convention, if the reaction releases heat ΔH is negative, and the reaction is said
to be exothermic. On the other hand, if the reaction occurs with the absorption of
heat it is said to be endothermic, ΔH is positive.
The enthalpy change for a chemical reaction depends on the temperature
and pressure of the reactants and products. The standard enthalpy of
reaction, ΔHo, is the enthalpy change for a reaction in which the reactants and
products are all in their standard states. The standard state of a substance is
taken to be the pure substance at 1 atm pressure, and 25oC.
For example, the combustion of methane is an exothermic process. For
the combustion of 1 mole of methane, ΔHo = -890 kJ/mole. This can be
summarized in terms of a thermochemical equation:
CH4 (g) + 2O2 (g) → CO2(g) + 2H2O(l) ΔHo = -890 kJ/mole (3)
This equation shows both the amount of material reacting, and the enthalpy
change. Each molecular formula represents 1 mole of substance.
The standard enthalpy of formation of a substance, ΔfHo is the enthalpy change
for the formation of 1 mole of the substance in its standard state from the
elements in their standard states. By convention, the enthalpies of formation of
the pure elements are zero. Equation (2), can be rewritten in terms enthalpies of
formation:
ΔH = Σ ΔfHo Products - Σ ΔfHoReactants (4)
In the case of reaction (3), we have:
ΔH = {2ΔfHo(H2O) + ΔfHo(CO2)} – {2ΔfHo(O2) + ΔfHo(CH4)} (5)
Numerical values of ΔfHo are available in tables of thermodynamic data.
Substitution into equation (5) yields:
ΔH = 2(-285.83 kJ/mol) + 1(-393.51 kJ/mol) -2(0 kJ/mol) -1(-74.81 kJ/mol)
= -890.31 kJ/mol
(6)
Hess's Law: According to Hess's Law, the enthalpy change of a chemical
reaction is independent of whether the reaction is carried out in one step, or in a
series of steps. The enthalpy change depends only on the initial and final states
of the reaction. For example, reaction (3) could be carried out in a two-step
process:
CH4(g) + O2(g) → CO2(g) + 2H2O(g) ΔHo1 =-802.29 kJ/mol
2[H2O(g) → H2O(l)]
ΔHo2= 2[-44.01 kJ/mol]
___________________________
____________________________
(7)
(8)
CH4 (g) + 2O2 (g) → CO2(g) + 2H2O(l) ΔHoRXN = ΔHo1 + ΔHo2 = -890.31 kJ/mol (3)
In this process 1 mol of methane is burned to produce 1 mol of gaseous CO2
and 2 mols of gaseous water. The gaseous water is then condensed to the
liquid. Note that ΔHo2 is equal to twice the enthalpy of the condensing of water.
The overall change is the same as in reaction (3). According to Hess's Law, ΔH
for the overall process is the sum of the ΔH's of the individual steps.
Purpose: In this experiment, you will measure the enthalpy change of each of
the three reactions below. From your data, you will analyze the results in terms of
Hess's Law.
HCl(aq) + NaOH(aq) → H2O(l) +NaCl(aq) (9)
HCl(aq) +NaOH(s) → H2O(l) + NaCl(aq) (10)
NaOH(s) → NaOH(aq)
(11)
Reactions (9) and (10) are acid-base neutralization reactions. They involve the
reaction of an acid,HCl, and a strong base, NaOH, to produce water and a
salt,NaCl. Reaction (11) represents the dissolution of solid NaOH in water.
Method
The measurement of the enthalpy change of a reaction involves measuring the
heat of the reaction at constant pressure. This is accomplished by using a
device known as a calorimeter. There are many different kinds of calorimeters.
The one used in this experiment is a "Coffee Cup Calorimeter” consisting of two
nested coffee cups, a lid and a thermometer (or thermistor if available). See
below.
Figure 2
In making a measurement, a known amount of acid is poured into the
calorimeter and the temperature measured. Also, a know amount of base is
measured out and the temperature measured. The average of the two
temperatures is the initial temperature, Ti . The measured amount of base
is added to the acid in the calorimeter, and, since the reaction is exothermic, the
temperature of the calorimeter and its contents rises. The final temperature after
the reactants are mixed, Tf, is then measured. The temperature change of the
calorimeter and its contents, in oC is:
ΔT =Tf -Ti (12)
The hear absorbed by the calorimeter and its contents is CCAL*ΔT + Cp*m*ΔT,
where
CCAL = the heat capacity of the calorimeter, in J/oC
Cp = the specific heat of the calorimeter contents, in J/(g-oC )
m=
the mass of the calorimeter contents in g
It is assumed that all of the heat produced by the reaction, qRXN is absorbed by
the calorimeter and its contents:
qRXN = -[CCAL*ΔT + Cp*m*ΔT]
(13)
From equation (12) we see that the determination of qRXN requires, in addition to
ΔT, the values of m, Cp and CCAL. The value of m will be calculated from mass
data, and that of Cp will be given. The value of CCAL will be determined
experimentally as described in the procedure below.
PROCEDURE
Part 1. The heat capacity of the calorimeter
Prepare about 200 mL of chilled water in a beaker by adding ice to
it. Weigh your empty dry calorimeter. Add 50.0 mL of chilled water into
the calorimeter and record the mass of the calorimeter and the water (the goal
here is to know the mass of water added. Make sure that no ice is transferred to
the graduated cylinder. Turn on the stirrer, and adjust for a slow stirring rate.
Place the lid on the calorimeter. Record the temperature vs. time readings for 2
minutes.
While collecting the temperature data for the cold water, place 50.0 mL of
water that has been heated to 50-55oC into a graduated cylinder. Using a
thermometer, measure and record the temperature of the hot water, and
immediately add it to the cold water in the calorimeter. The temperature in the
calorimeter will rise, reach a maximum, and the start to decrease. Collect data
as time vs. temperature for 3 minutes so that a graph of the data will look similar
to the one in Figure 2, below.
Repeat this measurement once and average your results.
Part 2. Calculation of the heat capacity of the calorimeter
When the hot water is added to the cold water in the calorimeter, the hot water
loses heat to the cold water and the calorimeter. The temperature of the hot
water decreases from some initial temperature TH to a final temperature TM , and
its temperature change is
ΔT H = TH - TM
The temperature of the cold water and the calorimeter increases from and initial
temperature, TC ,to a final temperature, TM , and its temperature change is
ΔT C = TM - TC
The temperature TM is the temperature of the water in the calorimeter at the time
of mixing, and can be obtained graphically as shown in Figure 2.
Figure 3
The heat lost by the hot water is
; qH = Cp*mH*ΔT H
The heat gained by the cold water is
qc = Cp*mC*ΔT C
and the heat gained by the calorimeter is
qCAL = CCAL*ΔT C
where Cp is the specific heat of water (4.184 J/(K-g)), and mC and mH are the
masses in grams of the cold and hot water respectively. The heat lost by the hot
water is equal to the heat gained by the cold water and the calorimeter. That is,
qH = qC + qCAL
Example: In an experiment to calibrate her coffee-cup calorimeter, a student
obtained the data in Figure 2. In addition, she found that the initial temperature
of the hot water, TH, was 51.0oC. What is the heat capacity of the calorimeter?
Solution: From the graph in Figure 2, TM = 27.0 oC, and TC = 5.0 oC. Therefore,
heat lost by hot water =
4.184 J/(oC-g)*50.0 g*(51.0 - 27.0)oC = 5.02*103 J
heat gained by cold water = 4.184 J/(oC-g)*50.0 g*(27.0 - 5.0)oC = 4.60*103J
heat gained by calorimeter = CCAL*ΔT C = 5.02*103 J - 4.60*103 J = 418. J
CCAL =
418. J/(27.0 - 5.0)oC =
19.0 J/oC
Part 3. The heat of neutralization of HCl(aq) and NaOH(aq)
Place 50.0 mL of 1.00 M HCl(aq) and the stir bar in the calorimeter, and 50.0
mL of 1.00 M NaOH(aq) in a graduated cylinder. Measure the temperature of the
HCl and the NaOH with the thermometer. The two solutions should be at the
same temperature (within + 0.2oC) If the temperatures are not the same adjust
the temperature of the NaOH by running warm or cold tap water over the
cylinder. Place the calorimeter on the stir plate. Start the stirrer and add the
NaOH to the HCl solution. The temperature will quickly rise and then level off.
Continue collecting data for about 1 min after the initial temperature rise.
Evaluate ΔT in equation (12) from a graph of temperature versus time as shown
in Figure 3. Calculate the heat of the reaction, qRXN, using equation (13). The
density of the 0.5 M NaCl(aq) solution produced is 1.02 g/mL and its specific
heat is 4.02 J/(g-oC). Finally calculate ΔH. Here, ΔH, is the enthalpy change of
reaction (9), and represents the heat produced when 1 mol of HCl reacts with 1
mol of NaOH. It is the heat per mol of HCl or NaOH reacted. Hence,
ΔH = qRXN/(moles of HCl reacted)
Repeat this measurement once.
Figure 4
Part 4. The Heat of Solution of NaOH(s).
Carefully weigh about 2.00 g (to + 0.01g) of NaOH(s). Since NaOH(s) is very
hygroscopic, weigh it by difference in a stoppered 50 mL Erlenmeyer flask. This
will require about 19 pellets. Measure 50.0 mL of deionized water into your
calorimeter.
Turn on the stirrer. Proceed as before and measure the temperature for about 1
min. Carefully add the NaOH(s) to the calorimeter by sliding the lid aside, and
quickly replacing it once the entire sample of solid has been added. Continue
taking readings for about 4 minutes after mixing. Measure ΔT at the time of
addition by extrapolation, as shown in Figure 4.
Figure 5
Calculate the heat of solution per mole of NaOH(s) dissolved. The final
concentration of NaOH is 1 M. The mass of solution is about 52 g and its
specific heat is 3.93 J/(g-oC).
Perform this measurement in duplicate.
Part 5. The Heat of Reaction of HCl(aq) and NaOH(s).
Again weigh out about 2.00g of NaOH(s), using the procedure of Part 4.
Carefully Measure 55 mL of 1 M HCl in a 100 mL graduated cylinder, and dilute it
to 100 mL with deionized water. This ensures that an excess amount of HCl is
present and that all of the NaOH will react. Transfer the entire 100 mL of HCl(aq)
to the calorimeter. Replace the lid and thermometer. Turn on the stirrer. Again
measure the temperature for about 1 min. Carefully add the NaOH(s) to the
calorimeter, and continue taking readings for 3-4 min. Prepare a graph of
temperature versus time, and, as in Part 4, determine ΔT at the time of mixing by
extrapolation.
Assume that the density and specific heat of the NaCl solution is the same as
that in part 2. In this case, the NaOH(s) is the limiting reactant, and ΔH is given
by
ΔH = qRXN/(moles of NaOH reacted)
Repeat this measurement once and average your results.
CALCULATIONS and GRAPHS
Use Excel to prepare graphs of temperature versus time for Parts 1 through 4.
Include these in your notebook. Clearly show your calculations of:
1.
The average heat capacity of the calorimeter
2.
The average enthalpy change for reaction (9), ΔH9.
3.
The average enthalpy change for reaction (10), ΔH10.
4.
The average enthalpy change for reaction (11),ΔH11.
CONCLUSION
1.
Report the average enthalpy change for each of the reactions 9, 10
and11.
2.
Explain how your results can be used to check the validity of Hess's
Law.
3.
Are your results consistent with Hess's Law? Please explain.
Given the enthalpies of formation at 25oC, in Table 1, below, calculate the
enthalpy change for reaction 9.
Table 1
Compound ΔfHo, kJ/mol, at 25 oC
HCl(aq)
-167.16
NaOH(aq)
-470.11
NaCl(aq)
-407.27
H2O(l)
-285.83
4.
5. Compare your measured value of ΔH9 with the value calculated in
question 3. What is the percent difference between your experimental
result, and the calculated value?
6. Can you think of any reasons, other than experimental error, why your
experimental value for ΔH9 might be different from the calculated value?