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UT CH 302 - Lecture notes
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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 (ΔG) 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?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 ΔH = 0 . Therefore since ΔG < 0, and ΔG = ΔH –TΔS, we know that ΔS >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 ΔH is heavily dependent on the identity of the solute and the solvent. For example, for NaCl dissolving in water ΔH > 0 while N2 gas dissolving in water ΔH < 0 To look at ΔH 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).The process described above can be expressed mathematically as the following: € ΔHsolution= ΔHlattice −energy+ΔHsolvation The first term is the change in enthalpy on forming the solution. The other two terms are the change broken down into the two changes in details above: the loss of solute-solute interaction and the gain in solute-solvent interactions. The first term is noted here as the Lattice Energy. This is what we typically associate with the energy of forming an ionic solid from separated ions. This is the energy requir ed to pull the solute apart. The second term is the enthalpy of solvation. This is the energy released when the solute interacts with the solvent. VERY IMPORTANT NOTE. It is critical that you keep track of the sign of these two enthalpy changes. Often the lattice-energy is negative number. However, here you need to use it as a positive. It is the energy input to overcome the solute/solute interactions. Thus you need it as a positive number. Energy in. Positive change. Likewise, the enthalpy of solvation should be negative. It is energy out. Negative change. Also, this term often assumes that the solvent is water. In these cases it is noted as the enthalpy of hydration. If you see the enthalpy of hydration, it is referring to the enthalpy of solvation (where the solvent is specified to be water). This is a great example of why understanding the actual process is so important. People choose different notation. People choose different sign conventions. You have to know what they mean. Breaking up the solute/solute interactions will always cost energy. Creating the solute/solvent interactions will always release energy. Now the question remains--which is bigger?IN NEARLY ALL CASES the “lattice-energy” term is larger than the solvation term. This means that generally the enthalpy of solution is positive. So how do we figure out what might dissolve and what won’t? Again we can look at the free energy ΔG =ΔH –TΔS We have established that generally for forming solutions ΔHsolution > 0 and ΔSsolution > 0. Thus, the entropy contribution helps to lower the free energy, but the enthalpy is raising it. If the change in free energy is going to be negative we need |TΔSsolution| > ΔHsolution There are two ways we can think about this. The first is to raise the temperature. This will lead to the entropy term dominating the process. The second is to make sure the enthalpy term is very small. How can we do this? We need to make sure that the changes in the intermolecular forces in going from the unmixed solute and solvent to the solution are minimal. With this we happen upon the general idea of understanding what substances will dissolve (mix). You may have heard the phrase “Like dissolves like”This means that similar things dissolve in similar things. What do we mean by similar? Similar intermolecular forces. Similar intermolecular forces between solute and solvent means there isn’t much change in energy upon making the solution. This is what we mean by minimizing ΔHsolution. Polar things dissolve polar (ionic) things. Non-polar things dissolve non-polar things. When is similar not similar enough? We can get a good idea looking at ionic compounds and water. Water is polar. It isn’t ionic. So already we know the IMF are not exactly the same. However, dipole-ion forces are substantial, even though they are likely not as large as ion-ion forces. When will this difference be too much? When the ion-ion forces are really large. If you remember from


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UT CH 302 - Lecture notes

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