It is proper under Hess' Law of Heat Summation, in chemical thermodynamics, to imagine a simplified pathway for a chemical process to understand the changes in thermodynamic state functions. Let us reprise our imaginary thermochemical path for the solution process.
The imaginary path from pure solute and solvent to solution can be conceptualized by increasing electric potential energy along lines of intermolecular force as the molecules of solvent and solute are separated. This is followed by decreasing potential energy as the molecules come together to form the solution.
In addition to net internal energy change (which equals the enthalpy change in transformations where the total volume changes little or not at all, as in solution processes generally) the equilibrium will also be determined by the effects of changes in the state of disorder in the system, or entropy. For nonpolar solvents dissolving nonpolar solutes, entropy generally increases, while for electrolytes dissociating in water, entropy decreases. Remember that! For salts dissolving in water, entropy decreases. This is because of the order imposed on the water through the formation of hydration shells.
Factors determining the interaction of a molecule or ion and a solvent include hydrogen bonding, the creation of a cavity in the solvent, the electrostatic and dispersion forces between solvent and solute molecules or ions, and an increase or decrease of solvent-solvent structure on introducing the solute.
Look at the topics in the Concept Map list and make sure you can follow this core reasoning, which is one of the foundations of a conceptual understanding of chemistry. First, think about the view at the intermolecular level of the separated solvent and solute versus the solution in terms of electric potential energy. What happens to internal energy as you go from one to the other? For most situations in which the solute is soluble in the solvent, the answer tends to be 'not much'. In other words 'like dissolves like'. If the internal energy doesn't change much, the solution process won't be too endothermic, although dissolving even a soluble solute is usually a little endothermic. (Do you notice how we've moved from the conceptual discussion of internal energy to enthalpy? Now we are using concepts of Thermochemistry). Finally, though, remember that if we are asking if a process is spontaneous, what we are really discussing is Chemical Thermodynamics. We are asking about the position of equilibrium.
The imaginary path from pure solute and solvent to solution can be conceptualized by increasing electric potential energy along lines of intermolecular force as the molecules of solvent and solute are separated. This is followed by decreasing potential energy as the molecules come together to form the solution.
In addition to net internal energy change (which equals the enthalpy change in transformations where the total volume changes little or not at all, as in solution processes generally) the equilibrium will also be determined by the effects of changes in the state of disorder in the system, or entropy. For nonpolar solvents dissolving nonpolar solutes, entropy generally increases, while for electrolytes dissociating in water, entropy decreases. Remember that! For salts dissolving in water, entropy decreases. This is because of the order imposed on the water through the formation of hydration shells.
Factors determining the interaction of a molecule or ion and a solvent include hydrogen bonding, the creation of a cavity in the solvent, the electrostatic and dispersion forces between solvent and solute molecules or ions, and an increase or decrease of solvent-solvent structure on introducing the solute.
Look at the topics in the Concept Map list and make sure you can follow this core reasoning, which is one of the foundations of a conceptual understanding of chemistry. First, think about the view at the intermolecular level of the separated solvent and solute versus the solution in terms of electric potential energy. What happens to internal energy as you go from one to the other? For most situations in which the solute is soluble in the solvent, the answer tends to be 'not much'. In other words 'like dissolves like'. If the internal energy doesn't change much, the solution process won't be too endothermic, although dissolving even a soluble solute is usually a little endothermic. (Do you notice how we've moved from the conceptual discussion of internal energy to enthalpy? Now we are using concepts of Thermochemistry). Finally, though, remember that if we are asking if a process is spontaneous, what we are really discussing is Chemical Thermodynamics. We are asking about the position of equilibrium.
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