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Laboratory 5 Hess s Law
Concepts In Biochemistry And Microbiology (SHGB6115 )
Universiti Malaya
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CHEM 106Preview text
Laboratory 5- Hess’s Law
Hannah Overman, Chem 161 NCSS Preformed February 11, 2014 LP : Jina Schun
Introduction
By using calorimetric based experiments, energy released (exothermic reaction) or energy absorbed (endothermic reaction) can be measured in a thermodynamic situation using Hess’s Law. When looking at energy transfer in the chemical reaction for instance, between magnesium and oxygen, it is important to distinguish between what is undergoing change, the system and the surrounding environment. When energy is released or absorbed from a reaction to its surroundings, it is in the form of heat represented by q, measured in joules in the equation q=mcΔT. The m in the equation represents the mass of the surroundings in grams that interacts. In this experiment, mL of the HCl solution is used to represent grams. The variable c represents the heat capacity of the surroundings in Joule/Gram ° Celsius), in this experiment, 4 is the recorded heat capacity and ΔT represents the change in temperature of the surroundings, determined by each trial.
Based on the results in our experiment, it is clear that Hess’s law is an acceptable method to calculate enthalpies of reactions. Despite the sources of error, Hess’s Law still proved to provide a result, which was close to the accepted value.
To measure the enthalpy change for the combustion of magnesium oxide, we used a coffee cup calorimeter to calculate the enthalpies of two separate reactions. The two reactions we conducted were:
Mg(s) + 2HCl(aq) → MgCl2(g) +H 2 and 2HCl (l) + MgO(aq) → MgCl2(s) + 2H 2 O(aq)
By calculating the enthalpies of the two prior equations, Hess’s Law allows you to combine those enthalpies, along with the known enthalpy of the formation of water, 4 J/g C, to figure the enthalpy of a combustion reaction that takes place between Mg and O 2. The enthalpy change would be negative (exothermic) based our previous knowledge of combustion as a reaction that releases energy.
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) ΔH1o=experimentally determined MgO(s) + 2 HCl(aq) → MgCl2(aq) + H2O(l) ΔHo2=experimentally determined H2(g) + 1⁄2 O2(g) → H2O(l) ΔH3o= -285 kJ
The sum of these three reactions produces:
Mg (s) + ½ O 2 (g) → MgO (s)
The desired enthalpy of formation of magnesium oxide is equal to:
ΔHf (MgO) = ΔH 1 + ΔH 2 + ΔH
Materials: Chemicals
2 Styrofoam cups
distilled water
thermistor ring stand
utility clamp
calorimeter
magnetic stirrer and stir bar cover for foam cup 100 ml- graduated cylinder
weighing paper
electronic scale
0 g Mg strips
1 MgO
200 ml 1 M HCl
Procedure:
We will have to pick which equipment to use to measure the temperature from various equipment depending on what’s available.
Obtain 2 nested Styrofoam cups that will function as your calorimeter for this experiment. Carefully measure out exactly 25 ml of HCl (density = 1. g/ml) and pour it into your calorimeter.
Measure and record the mass of a magnesium strip (~0 g), using an analytical balance.
Slide the temperature probe into the small notch cutout on the cover and place the probe into the HCl. Stir the HCl with the probe to maintain a uniform temperature throughout the solution. Wait until the temperature stabilizes.
Roll the magnesium ribbon into a loose ball. Collect temperature readings. After a ‘initial temperature’ is determined, slide the cover aside and drop the ball of magnesium into the calorimeter. Slide the cover back into place. Continue stirring and reading temperature until enough data points are taken until you have the ‘final temperature’.
Repeat the procedure 3 more times.
Perform Steps 2-5 now using ~0 g of magnesium oxide in place of the magnesium strips.
Repeat procedure 7 three more times.
Data /Calculations
Part 1: Reacting a magnesium ribbon with 25 mL of hydrochloric acid in a coffee cup calorimeter.
Magnesium Ribbon weight (grams)
Amount of 1M HCl added (mL)
Initial Temperature in Celsius
Final Temperatur e in Celsius
Tfinal-Tinitial
Trial 1 0 25 mL 21°C 20 °C -0°C
Trial 2 0 25 mL 21 °C 43 °C 22°C
Trial 2- Final Temperature 26 C
Calculation of data Part 1, using Trial 2: ΔH for the Mg/HCl reaction in kJ/mole of Mg
to find the enthalpy of Mg(s) + 2HCl+(aq) → MgCl2(aq) + H 2 > q = mcΔT Δ T = TFINAL - TINITIAL ΔT = 43 °C -21 °C = 22 °C 0 G MG
0 X (
1
24/mol
) = 0 MOL MG
= (25G HCL) (4 J/G °C ) (22 °C) + (10) (22°C) = 2555 J > 2. 55558 KJ
= (-2 KJ/0 MOL MG) =
ΔH = -qmix = qrxn
ΔH = -481 KJ/MOL
Part 2, using Trial 1: to find the enthalpy of MgO(s) + 2HCl(aq) → H 2 O (l) + MgCl(aq)
ΔH for the MgO/HCl reaction in kJ/mole of MgO.
q = mcΔT
Δ T = Tfinal - Tinitial ΔT = 27 °C -20 °C = 6 °C
0 g MgO/ (1 mol /40 MgO) = 0 mol MgO
q = (25 g HCl) ( 4 J/g·°C) ( 6 °C) +(10)(6 °C)
q = 767 or .76782 kJ
0 kJ/ 0 mol =
ΔH = -qmix = qrxn
ΔH= -122/mol
Mg (s) +2 HCl (aq) → MgCl2(aq)+H2 (g) MgO (s)+2 HCl (aq) →MgCl2(aq)+ H2O 2 H2(g)+O2 (g) → 2 H2O (l)
the result is the thermochemical equation above for the combustion of magnesium. AH for this reaction: 2 Mg (s)+4 HCl (aq) → 2 MgCl2(aq)+ 2 H2(g) 2 MgCl2(aq)+2 H2O (l) → 2 MgO (s) + 4 HCl (aq) 2 H2 (g) + O2 (g) → 2 H2O (l)
2 Mg (s) + O2 (g) → 2 MgO (s)
to find the enthalpy of Mg + 1/2O 2 → MgO using these equations and their enthalpies:
ΔH = Enthalpy: -481 kJ/mol 2 Mg (s)+4 HCl (aq) → 2 MgCl2(aq)+ 2 H2(g)
ΔH = Enthalpy: +122 kJ/mol 2 MgCl2(aq)+2 H2O (l) → 2 MgO (s) + 4 HCl (aq)
ΔH = Enthalpy: -285 kJ/mol 2 H2 (g) + O2 (g) → 2 H2O (l)
Enthalpy: -644 kJ/mol Mg + 1/2O 2 → MgO
value, based on our average calculations. The experiment demonstrated evidence that calculated values were close to the accepted values. Based on our resources of this lab, and the lack of containers to prevent heat escape, the percent error is relatively minimal.
Possible sources of experimental error could have occurred from improper lab technique. First off, when measuring the magnesium oxide powder into the weighing boat, the scale fluctuating, possibly giving an incorrect weight of the powder. This could have affected the enthalpy calculated. Also, when moving the magnesium oxide powder from the weighing boat to the Styrofoam cup, a small amount of powder could have slipped onto the lab station table or remained on the paper. This could have decreased the mass of the amount of magnesium oxide for the reaction, affecting the final enthalpy change. Another source of error that affected the accuracy of the value was that the magnesium strip had a coating of magnesium oxide. When pure magnesium is exposed to the air, it reacts oxygen to form magnesium oxide, potentially yielding inaccurate results. Finally, our initial trial resulted in an “out of control” result. This could have been due to the size of the Mg strip, the solution or contamination of the containers or solutions. We in turn switched out calorimeter reader, which resulted in another trial and inconsistent data.
Conclusion
Based on the results in our experiment, it is clear that Hess’s law is an acceptable method to calculate enthalpies of reactions. Despite the many sources of error that potentially may have made the results inaccurate, Hess’s Law still proved to provide a result, which was only 1% off from the accepted value and 1% from our average of trials combined.
Our enthalpy value was negative, and thus the reaction was exothermic. Hess’s law could be used to calculate the heat of combustion of magnesium. The limiting reagent was determined to be Mg and MgO respectfully and heat transferred to the water for Mg was -2 kJ and - kJ for MgO. Our heat of combustion result was relative to the standard heat of combustion rate listed -644 to std -633.
(Chemistry: The Molecular Nature of Matter, 6th Edition, Jesperson, Neil. Brady, James. Hyslop, Alison. January 2011. )
Laboratory 5 Hess s Law
Course: Concepts In Biochemistry And Microbiology (SHGB6115 )
University: Universiti Malaya
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