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In this experiment, the primary objective is to determine the heat of combustion of glucose using a constant-volume bomb calorimeter. Subsequently, this value will be utilized to calculate both the enthalpy of combustion and the standard enthalpy of formation of glucose.
The initial step involves using benzoic pellets as a standard for determining the calorimeter's heat capacity. Additionally, benzoic acid is employed along with glucose pellets as a spiking material to facilitate the ignition of glucose, which is challenging to ignite directly.
The results obtained indicate percentage errors of 0.0985% for the heat of combustion and 0.6897% for the heat of formation of glucose. These errors, although small, could potentially arise from various factors discussed in the subsequent Discussion section.
In this section, we will outline the objectives of the experiment in our own words, avoiding direct copying of theories provided in handouts. We will clearly define what we aim to measure and describe how we intend to obtain the desired quantity from experimental measurements.
The primary goal of this experiment is to determine the heat of combustion of glucose.
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Subsequently, we will utilize this value to calculate both the enthalpy of combustion and the standard enthalpy of formation of glucose. These calculations will be performed using a constant-volume bomb calorimeter.
It is important to note that the process ideally should be isothermal, but due to practical limitations, the temperature will not remain constant throughout. Thus, the isothermal process is carried out in two steps.
Firstly, we conduct an adiabatic process in which we obtain our products (A (T₀) + B (T₀) + S (T₀) → C (T₁) + D (T₁) + S (T₁)).
Subsequently, we bring the products back to their initial temperature by either adding or removing heat from the system (C (T₁) + D (T₁) + S (T₁) → C (T₀) + D (T₀) + S (T₀)).
It is important to account for the formation of nitrogen dioxide (NO₂) gas in air, which will react with oxygen to produce nitric acid (HNO₃).
To quantify this reaction, we employ sodium carbonate for titration to determine the volume of the resulting acid.
During calculations, we will determine the heat capacity of the calorimeter by employing a standard benzoic acid pellet. Subsequently, we will calculate the heat of combustion of glucose and, based on this value, derive the enthalpy of combustion and the enthalpy of formation of glucose.
The experimental procedure outlined in the Lab Manual was meticulously followed, with one exception: sodium carbonate was used for titration instead of sodium hydroxide. The apparatus employed for this experiment was a constant-volume bomb calorimeter, and titration was conducted to determine the quantity of nitric acid formed.
To prepare the mixed pellets, the following proportions of glucose and benzoic acid were used:
Table 1: Given values to be used
| Quantity | Value |
|---|---|
| ∆H combustion of glucose | -2808 kJ/mol |
| ∆H0f CO2 | -393.5 kJ/mol |
| ∆H0f H2O | -285.83 kJ/mol |
| ∆U wire combustion | -1400 cal/g |
| ∆U nitric acid | -14.1 cal/g |
| ∆U Benzoic acid | -6318 cal/g |
| Normality of Na2CO3 | 0.07163N |
Table 2: Data for benzoic acid (standardization pellets)
| Standard Pellet | T1 | T2 |
|---|---|---|
| Mass of pellet (benzoic acid) (±0.0001g) | 1.0033 | 0.9880 |
| Mass of the wire (±0.0001g) | 0.0173 | 0.0168 |
| Mass of the wire remaining (±0.0001g) | 0.0037 | 0.0031 |
| Mass of burned wire (±0.0001g) | 0.0136 | 0.0137 |
| Initial temperature T0 (±0.0001oC) | 19.4489 | 21.1457 |
| Rise in temperature ∆T (±0.0001oC) | 2.6128 | 2.6292 |
| Volume of Na2CO3 (±0.02mL) | 0.20 | 0.10 |
Table 3: Data for mixed benzoic acid and glucose pellets
| Mixed Pellet | T3 | T4 |
|---|---|---|
| Mass of glucose ±0.0001g | 0.7904 | 0.7906 |
| Mass of benzoic acid ±0.0001g | 0.2101 | 0.2101 |
| Mass of pellet ±0.0001g | 0.9963 | 0.9931 |
| Mass of the wire (±0.0001g) | 0.0157 | 0.0156 |
| Mass of the remaining wire (±0.0001g) | 0.0023 | 0.0011 |
| Mass of burned wire (±0.0001g) | 0.0133 | 0.0145 |
| Initial temperature T0 (±0.0001oC) | 23.2998 | 24.5531 |
| Rise in temperature ∆T (±0.0001oC) | 1.7725 | 1.7711 |
| Volume of Na2CO3 (±0.02mL) | 1.60 | 0.60 |
Heat capacity (Cv) of the calorimeter from standardization calculation:
Standardization 1:
Standardization 2:
Cv (average) = (2433.44 + 2381.147) / 2 = 2407.2935 cal/°C
Heat of combustion of glucose (ΔU) calculation:
Mixed pellet 1:
Mixed pellet 2:
ΔUglucose (average) = (-3716.1304 -3728.9334) / 2 = -3722.5319 cal/g
Enthalpy of combustion of glucose (ΔH) calculation:
Combustion of Glucose:
C6H12O6 (s) + 6O2 (g) → 6CO2 (g) + 6H2O (l)
ΔHcombustion of glucose = -2810.7665 kJ/mol
Heat of formation of glucose (ΔHf) calculation:
Hess’s law: ΔH∘r= ΔH∘fproducts−ΔH∘freactants [3]
Percentage Error = [(exp. - theoretical) / theoretical] * 100
Sample Calculation: % Error ΔHcombustion = [(-2810.7665 + 2808) / -2808] * 100 = 0.0712%
In reality, measuring ∆H directly in an isothermal process is not feasible due to temperature fluctuations caused by heat transfer. Therefore, an adiabatic process is employed, which consists of two steps:
For the isothermal process, we have:
ΔUisothermal = ΔUcal + ΔUimg
ΔUcal is zero because it's adiabatic and occurs at a constant volume. Therefore:
ΔUisothermal = ΔUT1→To = Cv (To - T1) = -Cv ΔT
where Cv = Cv (H2O) + Cv (CO2) + Cv (s)
It's worth noting that Cv (H2O) and Cv (CO2) are negligible in comparison to Cv (s) due to their much smaller masses.
In this experiment, we made several assumptions:
Table 4: Experimental and Real ΔHcombustion & ΔHf0 of Glucose Comparison
| Experimental Result | Real Value | Percentage Error | |
|---|---|---|---|
| ΔHcombustion (kJ/mol) | -2810.7665 | -2808 | 0.0985% |
| ΔHf0 (kJ/mol) | -1265.2135 | -1274 [4] | 0.6897% |
We considered that some air might still be present in the bomb, leading to the possibility of nitrogen gas reacting to form HNO3. To account for this, we used sodium bicarbonate for titration to determine the heat of combustion.
We also weighed the wire before and after combustion to include its heat of combustion in our calculations.
Benzoic acid served as a standard due to its ability to facilitate the combustion of glucose, which is not easily ignited.
The percentage errors were found to be 0.0985% and 0.6897%, both of which are considered small and acceptable.
Possible sources of errors include systematic errors from the machine or the digital balance, slight losses of glucose and benzoic acids during the preparation of mixed pellets, incorrect volume readings on the burette during titration, and adding excess sodium carbonate beyond the equivalence point.
Cellular respiration is a process that releases energy in the form of ATP by breaking down food. It has two types: aerobic and anaerobic respiration, with both using glucose as a reactant.
Table 5: Comparison between Aerobic and Anaerobic Respiration
| Aerobic Respiration | Anaerobic Respiration | |
|---|---|---|
| Occurrence | Takes place more in animals and plants | Takes place more in microorganisms |
| Oxygen | Takes place in the presence of Oxygen | Takes place in the absence of Oxygen |
| Place | Only takes place inside the cell | Takes place anywhere |
| Adenosine Triphosphate | Gives 36 ATPs per glucose molecule. | Gives 2 ATPs per glucose molecule. In animals: lactic acid. In Bacteria: methane and hydrogen sulfide. |
| End Products | End products: CO2 & H2O | End products: CH3CH2OH & CO2. In animals: lactic acid. In Bacteria: methane and hydrogen sulfide. |
| Oxidation of Substrate | Substrate is oxidized completely into carbon dioxide and water | Substrate is oxidized incompletely |
| Chemical Reaction | C6H12O6 + 6O2 → 6CO2 + 6H2O + 2,900 kJ/mol | C6H12O6 → 2C2H5OH + 2CO2 + 118 kJ/mol |
Some molecular geometries require molecules to bend or stretch their bonds from their normal state, resulting in strain energy (S). Cyclopropane is a well-known molecule with strain energy, as its carbon-carbon bond angles are constrained to be 60º.
Strain Energy Calculation:
S = ∆HII - ∆HI
Therefore, bomb calorimetry is used to determine the enthalpies of these two compounds, and their difference yields the strain energy of cyclopropane (S).
Benzene exhibits two equivalent structures, known as Kekule Structures:
C6H6
However, these are resonance structures where pi electrons are delocalized among all the carbon atoms. This equal sharing of electrons between all carbons contributes to benzene's enhanced stability.
To determine the resonance energy of benzene, we can compare the enthalpies of the real benzene molecule, which exhibits resonance, with one of the Kekule benzene structures (cyclohexatriene). The difference between these two enthalpies represents the remaining energy, known as the resonance energy.
These two molecules can be TTCC and Cyclohexane:
| TTCC | Cyclohexane | Benzene |
|---|
By employing bomb calorimetry, we can determine the enthalpies of TTCC and cyclohexane and then subtract their values. This calculation will yield the imaginary enthalpy of the Kekule benzene structure.
Subsequently, using bomb calorimetry again, we would combust benzene and calculate its enthalpy.
Finally, by subtracting the enthalpy of benzene from the enthalpy of the calculated Kekule structure, we can determine the resonance energy of benzene:
R = ∆Hbenzene - ∆Hkekule benzene
In this experiment, our primary objective was to calculate the enthalpy of combustion and the standard enthalpy of formation of glucose using a bomb calorimeter.
The experimental enthalpy of combustion was found to be -2810.7665 kJ/mol with a percentage error of 0.0985%. The enthalpy of formation was determined to be -1265.2135 kJ/mol with a percentage error of 0.6897%. These errors, although present, are relatively small and may be attributed to various sources of error discussed earlier.
It is important to note that Kekule structures do not exist in reality. Therefore, we need to explore other molecules with similar structures.
While there are no exact matches, we can consider using two molecules with differing structures (bonds and atoms) that, when combined, would theoretically yield the benzene structure. This approach assumes that these two molecules have no other source of energy, such as strain energy.
Determination Of The Heat Of Combustion And The Standard Enthalpy Of Formation Of Glucose. (2024, Jan 11). Retrieved from https://studymoose.com/document/determination-of-the-heat-of-combustion-and-the-standard-enthalpy-of-formation-of-glucose
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