ek 1001 chemistry #496 endothermic reaction and energies of bonds

[HalfPast6]

I can't seem to get my head wrapped around the solution to this problem.

EK 1001 Chemistry #496

"496. When solid KCl is dissolved in water, the energies of the bonds formed are lower than the energies of the bonds broken. Why does the reaction proceed anyway?

A. Undissolved KCL compensate for the portion that dissolves.
B. The reaction does not take place under standard conditions.
C. The electronegativity of the water is increased by the interaction of K and Cl ions.
D. The increased disorder due to mixing results in an increased entropy of the system.

Solution:
D. If the bonds formed have lower energy than the bonds broken, the reaction is endothermic. Therefore, using dG=dH-T*dS, the entropy of the system must increase if the reaction is to be spontaneous."

First of all, I answered D because I had ruled out the other answers and knew that entropy increases when a salt is dissolved in water. However, it's the first sentence in the solution that really gets me.

Initially, I thought that it was talking about potential energy. From my understanding, doesn't a lower potential energy of the bonds (more stable) formed make a reaction exothermic? Or is this line of reasoning wrong?

Or is EK talking about bond energy, in which case the solution would make more sense?

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[Gengar]

If you look in EK chemistry page 70, it explains this pretty well.

Solution formation is broken down into 3 steps.

Step 1: breaking of solute-solute bonds (endothermic since bond breaking is always endothermic)
Step 2: breaking of solvent-solvent bonds (endothermic)
Step 3: formation of solvent-solute bonds (exothermic, bond formation is always exothermic)

Thus, deltaH of the solution = deltaH1 + deltaH2 + deltaH3

If you form stronger bonds than those broken, then total deltaH will be negative as forming stronger bonds will make deltaH3 release more heat than the energy taken for deltaH1 and deltaH2.

The question says the energies of the bonds formed are at lower energies, so total deltaH will be positive here, because the energy released by the bonds formed is less than the energy it takes to break the solvent-solvent and solute-solute bonds.

Since deltaH is positive, for the reaction to be spontaneous, deltaS has to be positive as well (which makes sense because of the increased disorder of breaking apart KCl bonds).

Hope that helps.

 

[MedPR]

I can't seem to get my head wrapped around the solution to this problem.

EK 1001 Chemistry #496

"496. When solid KCl is dissolved in water, the energies of the bonds formed are lower than the energies of the bonds broken. Why does the reaction proceed anyway?

A. Undissolved KCL compensate for the portion that dissolves.
B. The reaction does not take place under standard conditions.
C. The electronegativity of the water is increased by the interaction of K and Cl ions.
D. The increased disorder due to mixing results in an increased entropy of the system.

Solution:
D. If the bonds formed have lower energy than the bonds broken, the reaction is endothermic. Therefore, using dG=dH-T*dS, the entropy of the system must increase if the reaction is to be spontaneous."

First of all, I answered D because I had ruled out the other answers and knew that entropy increases when a salt is dissolved in water. However, it's the first sentence in the solution that really gets me.

Initially, I thought that it was talking about potential energy. From my understanding, doesn't a lower potential energy of the bonds (more stable) formed make a reaction exothermic? Or is this line of reasoning wrong?

Or is EK talking about bond energy, in which case the solution would make more sense?

EK is talking about bond energy. I'm not sure what you are talking about though.


You need to input energy to break bonds. If your products have lower energy than your reactants, then the BDE of your products is less than the BDE of your reactants. If you input 500kJ of energy to break the reactant bonds, and come out with only 200kJ of energy in the product bonds, you have an endothermic reaction. This is unfavorable unless the TdeltaS value is large enough to overcome the positive deltaH.