Lab # 4
Experimental Application of Hess's Law
Experiment 1: Combustion of Magnesium Ribbon
Diana Bakos
Stein van der Plas
Liam Beck
SCH4U 01
November, 2014
Introduction
Hess's law is one of the most important concepts in chemistry. It is the direct
consequence of the first law of thermodynamics, which states that energy is conserved. Hess's
law states that enthalpy does not depend on the pathway taken to convert reactants into products.
This is because enthalpy is a state function, which only depends only on the initial and final state
of the reaction. Furthermore, if a reaction were to proceed in a series of steps the enthalpy
change for the overall reaction would equal the sum of all the enthalpies for the individual steps
(Clark). Hence, it is possible to arrive at a target chemical equation by adding up two or more
separate equations. The sum of all the separate enthalpies of each reaction would equal that of
the target equation. Therefore, Hess's law can be used to determine the overall energy required
for a chemical reaction, without even carrying out said experiment. This comes in very useful
when chemists want to find the enthalpy values for reactions that may be difficult to perform
directly or too dangerous perform directly (ChemTeam). In the following experiment,
applications of Hess's law were used to determine the change of enthalpy for combustion of
magnesium oxide (Mg(s) + ½O2(g) →MgO(s)) by adding three separate equations. A coffee cup
calorimeter was used to calculate the enthalpies of two separate reactions. The two reactions
conducted were that of Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) and MgO(s) + 2 HCl(aq) →
MgCl2(aq) + H2O. The enthalpy of the last reaction needed H2(g) + ½O2(g) → H2O(l) was
researched.
Objective
Please refer to page 1of 2
Materials
Please refer to page 1of 2
Procedure
Please refer to page 1of 2
Observations: Temperature Measurements for the Enthalpy of Combustion of Magnesium ribbon
Time
(seconds)
0s
30
60
90
120
150
180
210
240
270
300
330
360
390
420
450
480
510
540
570
600
630
660
690
720
750
780
810
840
870
900
Experiment 1: Equation 3
Mg(s) + 2 HCl(aq) -----> MgCl2(aq) + H2(g)
Experiment 1: Equation 4
MgO(s) + 2 HCl(aq) ----> MgCl2(aq) + H2O
20.8 oC
25.6 oC
30.5 oC
33.3 oC
38.8 oC
38.4 oC
38.3 oC
38.1 oC
38.0 oC
37.9 oC
37.7 oC
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
N/A
21.2 oC
22.1 oC
22.5 oC
22.7 oC
22.8 oC
22.9 oC
23.1 oC
23.1 oC
23.2 oC
23.3 oC
23.4 oC
23.5 oC
23.5 oC
23.5 oC
23.6 oC
23.7 oC
23.7 oC
23.7 oC
23.8 oC
23.8 oC
23.8 oC
23.9 oC
23.9 oC
24.0 oC
24.1 oC
24.1 oC
24.1 oC
24.1 oC
24.2 oC
24.2 oC
24.2 oC
Data Analysis- Experiment 1: Part 1:
Equation 3: Mg(s) + 2 HCl(aq) -----> MgCl2(aq) + H2(g)
Data Collected
Volume of 1.0 M HCl solution used
Mass of magnesium
Initial temperature of 1.0 M HCl
solution before mixing
Final temperature of solution after Mg
added
Calculations
ΔT
Energy produced in reaction
100mL
0.14g
20.8 oC
38 oC
=Tfinal -Tinitial
=39 oC -20.8 oC
=18.2 oC
Q= mc ΔT
= (100g)(4.184J/g oC) (18.2 oC)
=7615J
1𝑚𝑜𝑙 𝑀𝑔
Moles of magnesium oxide used
100gHCl×
ΔH
1𝑚𝑜𝑙 𝐻2
0.14gMg× 24.3𝑔𝑀𝑔 × 1 𝑚𝑜𝑙 𝑀𝑔 = 0.0058molMg
=
=
qrxn
n
7615J
0.0058mol
= -1313kJ/mol-1
1𝑚𝑜𝑙 𝐻𝐶𝑙
36.46𝑔𝐻𝐶𝑙
×
1𝑚𝑜𝑙 𝐻2
2 𝑚𝑜𝑙 𝐻𝐶𝑙
= 1.37molHCl
Temperature of Mg(s) + 2 HCl(aq)
40
38
Temperature (deg C)
36
34
32
30
Temperature
28
26
24
22
20
30
60
90
120
150
180
210
240
270
300
Time (s)
Data Analysis- Experiment 1: Part 2:
Equation 4: MgO(s) + 2 HCl(aq) ----> MgCl2(aq) + H2O
Data Collected
Volume of 1.0 M HCl solution used
Mass of magnesium oxide used
Initial temperature of 1.0 M HCl
solution before mixing
Final temperature of solution after
MgO added
Calculations
ΔT
Energy produced in reaction
100mL
0.66g
21.2 oC
24.2 oC
=Tfinal -Tinitial
=24.2 oC -21.2 oC
=3.1 oC
Q= mc ΔT
= (100g)(4.184J/g oC) (3.1 oC)
=1297J
Moles of magnesium oxide used
0.66gMgO×
1𝑚𝑜𝑙 𝑀𝑔𝑂
40.30𝑔𝑀𝑔
1𝑚𝑜𝑙 𝐻𝐶𝑙
1𝑚𝑜𝑙 𝐻2𝑂
× 1 𝑚𝑜𝑙 𝑀𝑔𝑂 = 0.016molMgO
1𝑚𝑜𝑙 𝐻2𝑂
100gHCl× 36.46𝑔𝐻𝐶𝑙 × 2 𝑚𝑜𝑙 𝐻𝐶𝑙 = 1.37molHCl
ΔH
=
=
qrxn
n
1297J
0.016mol
= -81.06kJ/mol-1
Temperature of MgO(s) + 2 HCl(aq)
25
24.5
23.5
23
Temperature
22.5
22
21.5
21
30
60
90
120
150
180
210
240
270
300
330
360
390
420
450
480
510
540
570
600
630
660
690
720
750
780
810
840
870
900
Temperature (deg C)
24
Time (s)
eqn (2) H2(g) + ½O2(g) → H2O(l) ΔH2= -572 kJ/mol-1
eqn(3) Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) ΔH3= +1313kJ/mol-1
eqn (4) MgCl2(aq) + H2O → MgO(s) + 2 HCl(aq) ΔH4= -81.06kJ/mol-1
______________________________________________________________________________
eqn (1) Mg(s) + ½O2(g) →MgO(s)
ΔH1 = ΔH2 + ΔH3 + ΔH4
= (-572
kJ/mol-1) + (+1313kJ/mol-1) + (-81.06kJ/mol-1)
= 660 kJ/mol-1
% 𝑒𝑟𝑟𝑜𝑟 =
(−660kJ · mol − 1) − (−601.8 kJ · mol − 1)
(−601.8 kJ · mol − 1)
% error = 9.67%
Discussion
Through the following experiment the molar enthalpy of the combustion of magnesium
ribbon was calculated to be -660 kJ/mol-1. The following result was calculated by using a coffee
cup calorimeter to determine the enthalpies of two separate equations; the molar enthalpy of the
reaction Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) was +1313kJ/mol-1and the molar molar
enthalpy of the reaction MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O was +81.06kJ/mol-1. The
enthalpy of the last reaction needed was researched. According to Chemwiki H2(g) + ½O2(g) →
H2O(l) has the molar enthalpy of -572 kJ/mol-1. In order to calculate the enthalpy change of the
equation for the combustion of magnesium the separate enthalpies were manipulated using
Hess's law. It was observed that the reactions were all exothermic because the final temperature
was greater than the initial temperature. In addition, this provides evidence as to why the value of
q was positive. However, the following experiment was not accurate because the value
calculated holds a 9.67% error compared to the theoretical value. The error could have been a
result of heat being lost to the surroundings instead of being absorbed into the system and the
assumptions made in the calculations. Heat was lost to the surrounding because the coffee-cup
calorimeter was a closed system. To maximize the efficiency of the calorimeter one could wrap
the calorimeter with an insulator in the form of a cotton mitten because cotton is an excellent
insulator. Another, source of error was the assumptions needed to be made to complete the lab
calculations. The experiment called for one to assume that the solution had the same specific heat
capacity as water. This error cannot truly be improved, but one could conduct the experiment
several times to calculate several heats of combustion for magnesium and then arrive at an
average. This would probably help with the accuracy of the experiment.
Source of error
The errors that occurred in this lab could have been a result of heat being lost to the
surroundings instead of being absorbed into the system and the assumptions made in the
calculations. Heat was lost to the surrounding because the coffee-cup calorimeter was a closed
system. Consequently, the temperature of the air surrounding the system increased. Due to the
fact that warm air is less dense than cold air, cold air forcibly replaced the warm air. This process
turned into a continues circle because it is well known that heat is always transferred from an
object at higher temperature to the object at lower temperature until thermal equilibrium is
reached. To maximize the efficiency of the calorimeter one could wrap the calorimeter with an
insulator in the form of cotton mittens because cotton is an excellent insulator. The mitten would
decelerate the effect of heat flow by conduction of convection because the thousands of tiny air
spaces between the fibers of the cotton would slow down the rate of transmission of energy.
Another, source of error was the assumptions needed to be made to complete the lab
calculations. The experiment called for one to assume that the solution had the same specific heat
capacity as water. Chemists usually assume that the specific heat capacity of a dilute solution is
very close to the specific heat capacity of pure water. Hence, the specific heat capacity of HCl
may not be 4.184 J K-1 g-1, but is somewhat precise to what the true value may be. This error
cannot truly be improved, but one could conduct the experiment several times to calculate
several heats of combustion for magnesium and then arrive at an average. This would probably
help with the accuracy of the experiment.
Conclusion
In conclusion, the molar enthalpy of the combustion of magnesium was calculated to be
-660 kJ/mol-1. The following result was calculated by determining the molar enthalpy of the
reaction Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) which was +1313kJ/mol-1and the molar
enthalpy of the reaction MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O which was +81.06kJ/mol-1.
The enthalpy of the last reaction needed was researched. According to Chemwiki H2(g) + ½O2(g)
→ H2O(l) has the molar enthalpy of -572 kJ/mol-1. A series of steps were followed to do the
calculation, which applied Hess's law. This law states that enthalpy is a state function and that it
is possible to arrive at a target equation (Mg(s) + ½O2(g) →MgO(s)) by adding up two or more
separate equations because the sum of all the enthalpies for the individual steps would equal the
overall reaction. Overall, the reaction only had a 9.67% error, which suggests that the experiment
was only slightly inaccurate.
Work Cited
Unknown, Author. "ChemTeam: Hess' Law - Using Two Equations and Their Enthalpies."
ChemTeam: Hess' Law - Using Two Equations and Their Enthalpies. N.p., n.d. Web. 01 Dec. 2014.
<http://www.chemteam.info/Thermochem/HessLawIntro1.html>.
Unknown, Author. "Hess's Law." - Chemwiki. N.p., n.d. Web. 01 Dec. 2014.
<http://chemwiki.ucdavis.edu/Physical_Chemistry/Thermodynamics/Thermodynamic_Cycles/Hess%27s
_Law>.
Unknown, Author. "Magnesium Oxide." Wikipedia. Wikimedia Foundation, 30 Nov. 2014. Web.
01 Dec. 2014. <http://en.wikipedia.org/wiki/Magnesium_oxide>.
Clark, Jim. "Hess's Law and Enthalpy Change Calculations." Hess's Law and Enthalpy Change
Calculations. N.p., May 2013. Web. 30 Nov. 2014.
<http://www.chemguide.co.uk/physical/energetics/sums.html>.