Posted by Rossina on Wednesday, November 1, 2006 at 6:20pm.
Use the following data (in kJ/mol) to estimate the standard enthalpy of formation (in kJ/mol) for hypothetical compound CsF2 in the solid state.
Lattice energy of CsF2 -2347
First ionization energy of Cs 375.7
Second ionization energy of Cs 2422
Electron affinity of F -328
Bond energy of F2 158
Enthalpy of sublimation of Cs 76.1
How am I supposed to put all of that information together? Can someone please start me...
deltaH = lattice energy+first IP+second IP+ (2xEA) + bond energy F2 + deltaH sub = ??
Write each equation above, then add them together to make sure that the signs are right and that all of them add together to give you the equation you want.
Cs(s) ==> Cs(g) 76.1 kJ/mol
Cs(g) ==> Cs^+(g) + e. +375.7 kJ/mol
Cs^+(g) ==> Cs^+2(g) + e 2422 kJ/mol
F(g) + e ==> F^-(g) -328 kJ/mol and you will need 2 mols of this which multiplies the -328 also.
What you want is for the final equation, after everything cancels, is
Cs(s) + F2(g) ==> CsF2(s)
There is a good site on the Internet. Go to www.google.com and look for Born-Haber cycle.
Post your work if you run into trouble. Check my work carefully. It is easy to add too many + signs and () in the wrong place.
I think I still haven't grasped some concepts to fully understand the problem. Right now I am reading the book... so when I come back to the problem, I will ask again if I need any more help from what you already provided me with.
Thanks for the help and site :)
OK. Please be specific about what you don't understand if you repost. Here are two web sites that help a great deal. Good luck. http://chemistry.bd.psu.edu/jircitano/BH.html
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