6 - LectureNotes3:Solubility,Solutions,andmixing...

Info iconThis preview shows pages 1–4. Sign up to view the full content.

View Full Document Right Arrow Icon
Lecture Notes 3: Solubility, Solutions, and mixing Up to now we have been dealing with pure compounds. Now we are going to begin to look at mixtures. The next two sets of notes will address mixtures. This first one is one mixing, solutions, and solubility. The next will be on the effect of mixing on phase transitions . A few terms to define before we begin Solution : A homogeneous mixture of two (or more) substances Solvent : The majority substance Solute : The minority substance For example, if you put sugar into water you make a sugar water solution. The water is the solvent. The sugar is the solute. This is to be contrasted with putting sand into water. Then you have solid sand sitting at the bottom of your water- a heterogeneous system. Let’s answer a very important question about mixing: what happens to the entropy? First ,I’ll just tell you that when you mix two things up, the entropy increases. (there are rare cases when this is not true, but this is only because of strong interactions between the molecules). What is an example of this? The mixing of two ideal gases. This process is spontaneous (it will actually happen). As such, we know that the change in free energy is negative (remember that a negative change in free energy ( Δ G) for a process is the requirement for spontaneity). Things that actually occur lower the free energy. Why does the free energy decrease? Is it due to the enthalpy (energy) or the entropy or both?
Background image of page 1

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full Document Right Arrow Icon
Since these are ideal gases we know it cannot be the enthalpy since there are no intermolecular forces (remember the ideal gas approximation is that there are no intermolecular forces at all). So we know that Δ H = 0 . Therefore since Δ G < 0, and Δ G = Δ H –T Δ S, we know that Δ S >0. That is, mixing increases the entropy! What about systems in which there are intermolecular forces? These are the ones we are generally interested in. Then Δ H is heavily dependent on the identity of the solute and the solvent. For example, for NaCl dissolving in water Δ H > 0 while N 2 gas dissolving in water Δ H < 0 To look at Δ H we have to think about what is happening when we make a solution. The differences in energy result from differences (or changes) in the intermolecular forces (IMF). If you can’t remember your IMF review them! Let’s look at how the IMF are changed during the formation of the solution. There are two key parts to the changes in the forces. The largest changes are for the solute. Before mixing, the solute molecules only interact with other solute molecules. In the solution,the solute molecules only interact with solvent molecules. Thus the changes result from the loss of the solute-solute interaction and the gain of the solute- solvent interactions (it should be noted that along with these changes there is a small loss of solvent-solvent interactions).
Background image of page 2
The process described above can be expressed mathematically as the following: Δ H solution = Δ H lattice energy + Δ H solvation The first term is the change in enthalpy on forming the solution. The other two terms are the change broken down
Background image of page 3

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full Document Right Arrow Icon
Image of page 4
This is the end of the preview. Sign up to access the rest of the document.

{[ snackBarMessage ]}

Page1 / 9

6 - LectureNotes3:Solubility,Solutions,andmixing...

This preview shows document pages 1 - 4. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online