The quantity of heat liberated by the decomposition of hydrogen peroxide is measured experimentally by allowing the reaction to take place in a thermally insulated vessel called calorimeter. The heat liberated in the decomposition will cause an increase in the temperature of the solution and of the calorimeter. If the calorimeter was perfect, not heat would be radiated to the laboratory.
How can temperature change inside a calorimeter be used to calculate ΔH for a reaction? When a calorimeter and contents absorb a given amount of heat, the temperature rise that results depends on the calorimeter’s heat capacity. Heat capacity(C) is the amount of heat required to raise the temperature of an object or substance a given amount. The exact amount of heat absorbed is equal to the heat capacity times the temperature rise: q = C×ΔT.
2. Theory
2.1 The Thermodynamic Standard State
The value of the enthalpy change ΔH reported for a reaction is the amount of heat liberated or absorbed when reactants are converted to products at the same temperature and in the molar amounts represented by coefficients in the balanced equation. In the decomposition of hydrogen peroxide, the reaction of 1 mol of hydrogen peroxide decomposed to give 1 mol of water and half mol of oxygen gas releases approximately 94.64kJ. The actual amount of heat released in a specific reaction, however, depends on the actual amounts of reactants. Thus, reaction of 0.05 mol of peroxide releases 0.05x94.64= 4. 732kJ.
Note that the physical states of reactants and products must be specified as solid(s), liquid (l), gaseous (g), or aqueous (aq) when enthalpy changes are reported. The enthalpy change for the reaction of peroxide is ΔH= -94.64 if water is produced as liquid but ΔH=-54.026kJ if water is produced as gas.
H₂O₂(aq) →H₂O(l) + O₂(g) ΔH=-94.64
H₂O₂(aq) →H₂O(g) + O₂(g) ΔH=-54.026kJ
In addition to specifying the physical state of reactants and products when reporting an enthalpy change, it is also necessary to specify the pressure and temperature. To ensure that all measurements are reported in the same way so that different reactions can be compared, a set of conditions called thermodynamic standard state has been defined- Most stable form of a substance at 1 atm pressure and at a specified temperature, usually 25⁰C; 1 M concentration for all substances in solution.
Measurements made under these conditions are indicated by addition of the superscript ⁰ to the symbol of the quantity reported. Thus, an enthalpy change measured under standard conditions is called a standard enthalpy of reaction and is indicated by the symbol ΔH⁰. The reaction of peroxide can be written
H₂O₂(aq) →H₂O(l) + O₂(g) ΔH⁰=-94.64
2.2 Enthalpies of Physical and Chemical Change
Physical Change:
Almost every change in a system involves either a gain or loss of enthalpy. The change can be either physical, such as the melting of a solid to a liquid, or chemical, such as decomposition of hydrogen peroxide.
Imagine what would happen if you started with a block of ice at a low temperature, say -10⁰C, and slowly increased its enthalpy by adding heat. The initial input of heat would cause the temperature of ice to rise until it reached 0⁰C. Additional heat would then cause the ice to melt without raising its temperature as the added energy was expended in overcoming the forces that hold H₂O molecules rigidly together in ice crystal. The amount of heat necessary to melt a substance is called the enthalpy of fusion, or heat of fusion, and has a value of 6.01kJ/mol for H₂O.
Once the ice has melted, further input of heat raises the water’s temperature until it reaches 100⁰C, and additional heat then causes the water to boil. Once again, energy is necessary to overcome the forces holding the molecules together in the liquid (van der Waals and hydrogen bonds), so the temperature does not rise again until all the liquid has been converted into vapor. The amount of heat required for evaporation is called the enthalpy of vaporization, or heat of vaporization, and has a value of 40.7kJ/mol at 100⁰C for H₂O.
Another kind of physical change in addition to melting and boiling is sublimation-the direct conversion of solid to a vapor without going through a liquid state. Solid CO₂ (dry ice), for example, changes directly from solid to vapor at atmospheric pressure without first melting to liquid. Since enthalpy is a state function, the enthalpy change on going from solid to vapor must be constant regardless of the path taken. Thus, a substance’s enthalpy of sublimation, or heat of sublimation, equals the sum of heat of fusion and heat of vaporization.
Chemical Change:
Enthalpy change is often called a heat of reaction because it is measure of the heat flow into or out of a system at constant pressure. If the products have more enthalpy that reactants, then heat has flowed into the system from surrounding and ΔH has a positive sign. Such reactions are said to be endothermic (endo means “within” so heat flows in). The reaction of 1 mol of barium hydroxide octahydrate with ammonium chloride, for example absorbs 80.3kJ from the surroundings (ΔH⁰=+80.3kJ). The surroundings, having lost heat, become so cold that water freezes around the outside of the container.
If the products have less enthalpy than reactants, then heat has flowed from the system to the surroundings and ΔH has a negative sign. Such reactions are said to be exothermic (exo means “out”, so heat flows out). The so called thermite reaction of aluminum with iron(III) oxide, for example, releases so much heat and the surroundings get so hot (ΔH⁰=-852kJ) that is used in construction work as building railroads to weld iron.
As noted previously, the value of ΔH⁰ given for an equation assumes that the equation is balanced for the number of moles of reactants and products, that all substances are in their standard states, and the physical state of each substance is as specified. The actual amount of heat released in a specific reaction depends on the amount of reactants.
It also emphasized that ΔH⁰ refer to the reaction going in the direction written. For the reverse reaction, the sign of ΔH⁰ must be changed. Because of the reversibility of state functions, the enthalpy change for any reaction is equal in magnitude but opposite sign to that for the reverse reaction.
2.3 Calorimetry and Heat Capacity
The amount of heat transferred during a reaction can be measure with a device called calorimeter, shown schematically in Figure 1. At its simplest, a calorimeter is just an insulated vessel with a stirrer, a thermometer, and a loose-fitting lid to keep the contents at atmospheric pressure. The reaction is carried out inside the vessel, and heat evolved or absorbed is calculated from the temperature change. Because the pressure inside the calorimeter is constant (atmospheric pressure), the temperature measurement makes it possible to calculate the enthalpy change ΔH during a reaction.
Figure 1: Simple calorimeter (McMurry & Fay 2003).
A somewhat more complicated device called a bomb calorimeter is used to measure the heat released during a combustion reaction, or burning flammable substance. (More generally, a combustion reaction or burning of a flammable flame.) The sample is placed in an insulated, water-filled container (Fig 2). The reactants are ignited electrically, and the evolved heat is calculated from the temperature change of the surrounding water. Since the reaction takes place at constant volume but not constant pressure, the measurement provides a value of ΔE rather than ΔH.
Figure 2: Bomb calorimeter (McMurry & Fay 2003).
How can temperature change inside a calorimeter be used to calculate ΔH for a reaction? When a calorimeter and contents absorb a given amount of heat, the temperature rise that results depends on the calorimeter’s heat capacity. Heat capacity(C) is the amount of heat required to raise the temperature of an object or substance a given amount, a relationship that can be expressed by the equation
Where q is the quantity of heat transferred and ΔT is the temperature change (ΔT= T final – T initial). The greater the heat capacity, the greater the amount of heat needed to produce a given temperature change. A bathtub full of water, for instance, has greater heat capacity than a coffee cup full, and it therefore takes far more heat to warm the tubful than the cupful. The exact amount of heat absorbed is equal to the heat capacity times the temperature rise:
q = C×ΔT
Heat capacity is an extensive property, so its value depends on both size of an object and its composition. To compare different substances, it’s useful to define a quantity called specific heat, the amount of heat necessary to raise the temperature of 1g of a substance by 1⁰C. The amount of heat necessary to raise the temperature of a given object, then, is the specific heat times the mass of the object times the rise in temperature:
q = (Specific heat) × (Mass of substance) x ΔT
Now let us relate back to the first assignment in the experiment. We measured the heat lost by warm water and heat gain by cool water by the expression above. The heat lost should be equal to the heat gain according to the first law of thermodynamics which states that in any process in an isolated system, the total energy remains the same. However, when calculation has been established, we notice that the heat gain and heat lost are not the same; hence, heat must be absorbed by the calorimeter. With this relationship, we can deduce the heat capacity of the calorimeter. This heat capacity will be useful to determine the heat change for calorimeter in second assignment.
2.4 Standard Heats of Formation
With this we will be able to explain where do the literature value of -94.6kJ/mol for the decomposition of hydrogen peroxide shown in the lab manual come from. There are many chemical reactions-several hundred million are known-that it’s impossible to measure ΔH⁰ for all of them. A better way is needed.
The most efficient way to manage with the smallest number of experimental measurements is to use what are called standard heats of formation, symbolized.
Standard heat of formation is defined as the enthalpy change for the formation of 1 mol of a substance in its standard state from its constituent elements in their standard states.
Note several points about this definition. First, the “reaction” to form a substance from its constituent elements can be (and often is) hypothetical. We can’t combine carbon and hydrogen in the laboratory to make methane, for instance, yet the heat of formation for methane is -74.8kJ/mol, which corresponds to the standard enthalpy change for the hypothetical reaction.
Second, each substance in the reaction must be in its most stable, standard form at 1atm pressure and the specified temperature (usually 25⁰C). Carbon, for example is most stable as solid graphite rather than diamond under these conditions, and hydrogen is most stable as gaseous H₂ rather than H atoms.
The most stable form of any element in the standard state like O₂ has = 0kJ. (That is, the enthalpy change for formation of an element from itself is zero.)
The standard enthalpy change for any chemical reaction is found by subtracting the sum of heats of formation of all reactants from the sum of heats of formation of all products, with each heat of formation multiplied by the coefficient of the substance in the balanced equation.
Let us determine the literature value of the decomposition of hydrogen peroxide with the equation above.
3. Procedure
Assignment 1: Heat capacity of calorimeter
Foremost, a calorimeter was constructed with Styrofoam cup, cover and a thermometer. Next 30mL of tap water has been placed in the calorimeter cup and replaced cover and thermometer.
The system has reached its equilibrium after a wait of 5 to 10 min and the temperature has been recorded to the nearest 0.5oC. Subsequently, 30mL of water was placed in a 250mL beaker and heated using a hot plate until the temperature is approximately 20oC above room temperature while ensuring that water does not boil as appreciable water was not lost that may lead to an erroneous result. After heating, the water was left to stand for a minute or two and quickly recorded its temperature to the nearest 0.5oC before it was poured completely into the calorimeter. The cover was replaced with the thermometer and the cup was swirled gently once.
Without further delay, the temperature was watched over for the next 3 minutes and recorded for every interval of 15 seconds. Ultimately, the temperature as the function time was plotted and ΔT together with heat capacity of the calorimeter were then determined from the graph.
Assignment 2: Enthalpy of decomposition of hydrogen peroxide solution
First of all, the calorimeter and thermometer were dried. Next, 50mL of 1.0M H2O2 has been carefully measured and added to the calorimeter. Immediately, the cover was replaced on the calorimeter and the thermometer. Without further delay, the solution was swirled once slowly and recorded the temperature every minute for 4 minutes. Upon reaching the 5 minute mark, the cover along with the thermometer were removed before adding 10mL of 0.50M Fe(NO3)3 to the solution and the cover and thermometer were quickly replaced after the addition. Followed by that, the temperature was measured at the 6 minute mark and every succeeding minute up to a total of about 20 minutes. Later on, a temperature-versus time curve was plotted and the value of ΔT was then determined. Enthalpy of decomposition of hydrogen peroxide was then calculated. The initial temperature has been obtained by extrapolating the 5 points prior to adding the catalyst to the point of mixing. The final temperature has also been attained by extrapolating the linear portion of the graph to the point of mixing.
4. Results and Calculation
Assignment 1: Heat Capacity of Calorimeter
Temperature of cold tap water, T1 = 25.5oC
Temperature of warm water, T2 = 41oC
Temperature of cold and warm tap water after mixing:
Maximum temperature after mixing from table above, Tmax = 33.0oC
Assignment 2: Enthalpy of Decomposition of Hydrogen Peroxide Solution
Record of temperature every minute:
Initial temperature of H2O2 solution, Ti = 24.5oC
Final maximum temperature of H2O2 solution (from graph), Tf = 43.0oC
To calculate heat capacity of calorimeter, c (J/OC):
Heat lost to calorimeter = heat lost by warm tap water – heat gained by cold tap water
Heat lost by warm tap water = mass of water x specific heat of water x (T2 – Tmax)
= 30g x 4.187 J/oC g x 8
= 1004.88 J
Heat gained by cold tap water = mass of water x specific heat of water x (Tmax – T1)
= 30g x 4.187 J/oC g x 7.5
= 942.075 J
Therefore, Heat lost by calorimeter = 1004.88 – 942.075
= 62.805 J
Where heat lost to calorimeter = (Tmax – T1) x c
Thus heat capacity of calorimeter, c = 8.374 J/OC
To calculate qsol and qcalorimeter:
Heat change for solution, qsol = -specific heat of solution x mass of solution x (Tf - Ti) = -4.187 J/OC x 60g x (43.0 – 24.5) =-4647.57J
Heat change for calorimeter, qcalorimeter = heat capacity of calorimeter x (Tf - Ti) =154.919J
To calculate the enthalpy of decomposition of H2O2 (literature value = -94.6 kJ/mol)
ΔH=q/n
Where
q = qtotal = qsol + qcalorimeter
n = number of moles of H2O2 reacted
= (1.0 M)(0.05 L)
= 0.05 moles
Therefore, ΔH = (-4647.57 + 0.154919)/0.05 ≈ -89.9 kJ/mol (about 5.02% error from the literature value)
5. Discussion
The enthalpy of decomposition of Hydrogen peroxide obtained from the experiment has a value of -89.9kJ/mol. However, the literature value is about -94.6kJ/mol, whereby we an error of 5.02% from the literature value. Overall the experiment is a success but we may identify some errors and precautions to take:
During heating of the 30mL of water in the 250mL beaker, water may boil and evaporate and causes a significant drop in the volume of water in the beaker. Thus, we will not be able to accurately calculate the heat capacity of the calorimeter. It is advisable for us to place a filter funnel covering the beaker to prevent water loss while heating.
Moreover, the temperature change is rather small and subtle (there have be instances whereby temperature changes by 0.1 or 0.2 degree Celsius in a minute). However, we’ve only recorded our data to nearest 0.5 degree Celsius and this explains why when we extrapolate some of the points of the linear portion of the graph it is a perfect straight line instead it should have a slope and we would have gotten a very different value for ΔT. Therefore, I would recommend the use of magnifying glass while reading the temperature and enabling us to record the temperature to the precision of 0.2 degree Celsius.
Adding to that, in this experiment we are very concern with the volume of water added (accuracy) to determine the heat capacity of calorimeter. However, we have chosen to use graduated (measuring) cylinders to take measurement. Since the area of the surface of the liquid is much greater than graduated flask, the accuracy is not very high. Graduated cylinders cannot therefore be employed for work demanding even a moderate degree of accuracy. They are, however, useful only rough measurements are required. Therefore, in my humble opinion I would recommend the use of burette to measure out the volume of water needed to greater accuracy.
Furthermore, the simple calorimeter we have used in the experiment poorly insulates the system from the surrounding and significant heat will be lost from the system to the surrounding. Therefore, we could stack two Styrofoam cups on top of one another in order to trap air and create a layer of insulation shown in figure 3 or even better if we can use bomb calorimeter.
Figure 3: Suggested setup for simple calorimeter.
Human errors such as slowness and parallax errors affect the reading. For instance, when we are told to cover the calorimeter immediately, there will still be delay of fraction of a second or more which will result in loss of heat and temperature will differ from the ideal. While taking reading from the thermometer or measuring out the volume of reactants need with measuring cylinder parallax errors might occurs and we have to careful while reading the meniscus level.
All types of volumetric glassware have a cylindrical shape in the measuring region which causes the surface of most liquids whose volumes are to be measured to be curve downwards. Take reading from the bottom of the curved surface called meniscus with your eyes at the same level.
To read the position of the meniscus, the eye must be at the same level as the meniscus, in order to avoid errors due to parallax.
6. Conclusion
The objectives of this experiment such as the heat of decomposition of hydrogen peroxide, the heat capacity of calorimeter and enthalpy of decomposition of hydrogen peroxide were measured and calculated. The heat capacity of the calorimeter was 8.374J/oC and the enthalpy of decomposition of hydrogen peroxide was -89.9kJ/mol. Therefore, we have successfully met the literature value with an error of approximately 5.02%. Therefore, I have stated a few suggestions that may be put to use in if we can repeat the experiment to achieve a more accurate and precise value of enthalpy.
8. References
Internet:
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Wisegeek, 2008. What is calorimeter? [online]. Available from: [Accessed 7 January 2010].
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MicroLab, 2010. Enthalpy and entropy of zinc with copper sulfate [online]. Available from: [Accessed 6 January 2010].
Books:
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McMurry Fay, 1998. Chemistry. 4 ed. USA: Prentice Hall.
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G.H. Jeffery and J. Bassett, 1989. Vogel’s Textbook of Quantitative Chemical Analysis. 5 ed. UK: Longman Scientific & Technical.
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Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten (2009), Chemistry the Central Science, 11th Edition.
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