estimate Delta H for the following reaction
C2H4 (g)+3O2(g)→2 CO2(g)+2 H2O(g)
So we have (4) C-H bonds and (1) C=C bond or (H = 4*414 + 615= 2271)
For oxygen,
(3) O=O bonds or (H = 3*502=1506)
2CO2 is (2) O=C=O
(H=730*4=2920)
2H2O is (2) H-O-H
or (H=4*464=1856)
Delta H =(2271+1506)-(2920+1586)=-729
But the actual answer is -1614.
Can you please tell me where am I wrong?
To estimate the enthalpy change (ΔH) for a chemical reaction, you can use the bond enthalpy approach. The bond enthalpy is the amount of energy required to break one mole of a specific bond in a gaseous molecule.
It seems like you have correctly determined the bond enthalpies for the C-H and C=C bonds in ethene (C2H4) and the O=O bond in oxygen (O2). However, there are a couple of errors in your calculations.
First, you correctly calculated the sum of the bond enthalpies for the reactants: (4 C-H bonds) + (1 C=C bond) + (3 O=O bonds) = 2271 kJ. This represents the energy required to break these bonds.
Next, you determined the bond enthalpies for the products: (2 O=C=O bonds) + (2 H-O-H bonds) = 2920 kJ. This represents the energy released when these bonds are formed.
To calculate the ΔH for the reaction, you need to subtract the energy released from the energy required: ΔH = (energy required) - (energy released). Therefore, ΔH = 2271 kJ - 2920 kJ = -649 kJ (rounded to the nearest whole number). This means the reaction is exothermic, releasing energy.
However, I believe the discrepancy in your calculated value and the actual answer (-1614 kJ) might be due to the fact that the values you used for the bond enthalpies are not accurate. Bond enthalpies can vary based on factors such as molecule environment and bond characteristics.
It is more accurate to refer to experimentally determined bond enthalpies, which can be found in reliable databases or textbooks. Alternatively, you can calculate ΔH using other methods such as Hess's Law or using standard enthalpy of formation values.
I hope this clarifies where the discrepancy lies and how you can approach estimating ΔH for a chemical reaction.
You made a mistake in calculating the enthalpy change for the products. The correct calculations are:
For CO2:
2 O=C=O bonds or (H = 2*799=1598)
For H2O:
2 H-O-H bonds or (H = 2*464=928)
Therefore, the enthalpy change for the products is:
Delta H = (2271 + 1506) - (1598 + 928) = 3251 - 2526 = -275 kJ/mol
So, the correct estimate for Delta H is -275 kJ/mol, not -729 or -1614 kJ/mol.