how much energy would be released in the following reaction
CH4 + 2O2--> CO2 + 2H2O
To calculate the energy released in a chemical reaction, we need to use the concept of enthalpy change (∆H). The enthalpy change (∆H) represents the heat energy released or absorbed during a reaction.
The enthalpy change (∆H) for the reaction you provided, which is the combustion of methane (CH4) to form carbon dioxide (CO2) and water (H2O), can be determined by subtracting the sum of the enthalpies of the reactants from the sum of the enthalpies of the products.
The balanced equation for the reaction is:
CH4 + 2O2 --> CO2 + 2H2O
The standard enthalpy of formation (∆H˚f) for CH4 is -74.8 kJ/mol.
The standard enthalpy of formation (∆H˚f) for CO2 is -393.5 kJ/mol.
The standard enthalpy of formation (∆H˚f) for H2O is -285.8 kJ/mol.
Using these values, we can calculate the energy released as follows:
Reactants: CH4 + 2O2
∆H = 1*(-74.8 kJ/mol) + 2*(0 kJ/mol) [The standard enthalpy of formation of O2 is 0 kJ/mol since it is in its elemental form.]
Products: CO2 + 2H2O
∆H = 1*(-393.5 kJ/mol) + 2*(-285.8 kJ/mol)
Now, subtracting the sum of the enthalpies of the reactants from the sum of the enthalpies of the products:
∆H = [1*(-393.5 kJ/mol) + 2*(-285.8 kJ/mol)] - [1*(-74.8 kJ/mol) + 2*(0 kJ/mol)]
Simplifying the equation:
∆H = -820.4 kJ/mol - (-74.8 kJ/mol)
∆H = -745.6 kJ/mol
Therefore, in the given reaction, approximately 745.6 kJ of energy would be released.
To determine the amount of energy released in a chemical reaction, we need to calculate the change in enthalpy (ΔH).
Here's how you can calculate the energy released in the given reaction:
1. Identify the balanced equation:
CH4 + 2O2 → CO2 + 2H2O
2. Determine the stoichiometric coefficients:
In the balanced equation, the stoichiometric coefficient for CH4 is 1, and for O2 it is 2.
3. Look up the standard enthalpy of formation values:
Find the standard enthalpy of formation (ΔHf) values for all the reactants and products in the reaction. These values represent the heat energy released or absorbed when one mole of a substance is formed from its constituent elements in their standard states.
Using these values, we have:
ΔHf(CH4) = -74.8 kJ/mol
ΔHf(O2) = 0 kJ/mol
ΔHf(CO2) = -393.5 kJ/mol
ΔHf(H2O) = -285.8 kJ/mol
4. Calculate the change in enthalpy (ΔH):
ΔH = Σ(n * ΔHf(products)) - Σ(n * ΔHf(reactants))
Here, "n" represents the stoichiometric coefficient of each substance in the reaction.
ΔH = (1 * ΔHf(CO2) + 2 * ΔHf(H2O)) - (1 * ΔHf(CH4) + 2 * ΔHf(O2))
Plugging in the values, we can calculate:
ΔH = (1 * -393.5 kJ/mol + 2 * -285.8 kJ/mol) - (1 * -74.8 kJ/mol + 2 * 0 kJ/mol)
ΔH = -1181.1 kJ/mol + 149.6 kJ/mol
ΔH = -1031.5 kJ/mol
Therefore, the reaction releases approximately 1031.5 kilojoules (kJ) of energy per mole of CH4 reacted.
delta Hrxn = (n*sum delta Hf products) - (n*sum delta Hf reactants).
You can find a table of delta Hf values in your text or notes.