(CHEM) Which of the following compounds can exist as cis and trans isomers ?

[miss.strw]

Using average bond enthalpy values estimate ΔH (kJ/mol) for the following gas phase r

Using average bond enthalpy values estimate ΔH (kJ/mol) for the following gas phase reaction.

CH2=CH2 + Br2 --> CH2BrCH2Br

I followed this topic: http://answers.yahoo.com/question/index?qid=20080306194402AAtIc63
and I got the wrong answer.

I came up with 129 which is wrong.

Here is my work :
C-H = 4x416 = 1664
C=C = 1x598 = 598
Br-Br = 2x193 = 386

1664+598+386 = 2648




C-H = 4x416 = 1664
c-Br = 3x285 = 855

1664+2519 = 2519

2648-2519 = 129 (which is wrong.) Please Help Me! Thank you!

 

[nate808]

Thank you for your question. Your approach to using average bond enthalpy values to estimate ΔH for this reaction is correct, but there are a few errors in your calculations. Let's go through the steps together to ensure we get the correct answer.

Step 1: Write out the balanced chemical equation for the reaction:

CH2=CH2 + Br2 --> CH2BrCH2Br

Step 2: Identify the bonds broken and formed in the reaction. In this case, we have 4 C-H bonds and 1 C=C bond broken, and 2 C-Br bonds formed.

Step 3: Look up the average bond enthalpy values for each bond. These can be found in a table, such as the one in the link you provided.

C-H = 416 kJ/mol
C=C = 598 kJ/mol
C-Br = 285 kJ/mol

Step 4: Calculate the total bond enthalpy for the bonds broken and formed.

Bonds broken: 4(416 kJ/mol) + 1(598 kJ/mol) = 1992 kJ/mol
Bonds formed: 2(285 kJ/mol) = 570 kJ/mol

Step 5: Calculate ΔH for the reaction by subtracting the total bond enthalpy of the bonds formed from the total bond enthalpy of the bonds broken.

ΔH = 1992 kJ/mol - 570 kJ/mol = 1422 kJ/mol

Therefore, the estimated ΔH for this gas phase reaction is 1422 kJ/mol. I hope this helps and clarifies any confusion. Keep up the good work in your studies!

 

[Blueshift5]

I would first like to clarify that the question is asking for the enthalpy change for the reaction, not the isomers of a compound. The two isomers mentioned in the question, cis and trans isomers, are different arrangements of the same compound (CH2BrCH2Br) and do not affect the enthalpy change of the reaction.

To calculate the enthalpy change of the reaction, we need to use the bond enthalpy values for each bond broken and formed in the reaction. The bond enthalpy values that you have used for C-H and C=C bonds are correct, but the value for the Br-Br bond is incorrect. The correct value for Br-Br bond enthalpy is 193 kJ/mol, not 285 kJ/mol.

Using the correct values, the calculation would be as follows:

Bond enthalpies for bonds broken:
2 C-H bonds = 2 x 416 kJ/mol = 832 kJ/mol
1 C=C bond = 1 x 598 kJ/mol = 598 kJ/mol
1 Br-Br bond = 1 x 193 kJ/mol = 193 kJ/mol

Bond enthalpies for bonds formed:
2 C-Br bonds = 2 x 285 kJ/mol = 570 kJ/mol

ΔH = (832 + 598 + 193) - (570) = 1053 kJ/mol

Therefore, the enthalpy change for the reaction is 1053 kJ/mol. I hope this helps clarify your confusion. Remember to always double check your values and units to ensure accurate calculations.