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Electrolytes, Non-Electrolytes and Phase Transition
Name: Angela
Status: other
Grade: other
Country: Canada
Date: Winter 2009-2010
Question:
Why do solutions of electrolytes have a greater effect on
the boiling point and freezing temperature of solvents than
solutions of non electrolytes?
Replies:
The effect of solutes on the freezing point and boiling point of a solvent
depends upon the number of "particles" released by the solute into the
solvent. This is true in dilute solution whether the solute is an
electrolyte or a non-electrolyte. Soluble non-electrolytes typically produce
a single molecular species.
Soluble electrolytes typically produce at least two molecular species. So
the effect of an electrolyte is approximately double (or more) than of a
non-electrolyte. In a text book you will find this discussed under the
heading of "colligative properties", which means that the effect just
depends upon the number of particles of solute per mol. This is a
simplification because when you get down into the nitty gritty things get
more complicated, but for your purposes the above suffices.
Vince Calder
Because each mole of an electrolyte contains at least two moles of
ions, which dissociate in solution. Non-electrolytes do not dissociate.
This sounded like a homework question. I hope you inform your
teacher how you obtained your answer.
Richard Barrans
Department of Physics
Angela,
Electrolytes contribute more than one particle per formula unit when they
dissolve. e.g. NaCl dissolves to give two particles per formula unit (Na+
and Cl- ions). Solutions of a non-electrolyte will only give one particle
per formula unit. e.g. glucose C6H12O6 dissolves to give one glucose molecule
per formula unit, it does not dissociate. This means that solutions of the
same concentration of each solute will give different numbers of solute
particles in solution (in fact the NaCl will have almost twice as many).
Freezing point depression and boiling point elevation are colligative properties:
they depend only on the number of solute particles present as a ratio of the
number of solvent particles. With a 1m solution of NaCl you get about 2 moles
of particles per kg of solvent, where as a 1m solution of C6H12O6 gives you
1 mole of particles per kg of solvent.
This "effective number of particles per formula unit" (1 for C6H12O6 and about
2 for NaCl) is given as the van't Hoff factor. It is actually experimentally
found and doesn't always agree exactly with what we expect. For NaCl it is
only 2 for dilute solutions; we think the ions pair up at higher concentrations
and reduce the "effective number of particles per formula unit" below the ideal
situation.
Interestingly, it does not seem to affect the boiling point elevation and freezing
point depression very much if you change what the solute particles are, only the
number of particles present seems to matter. This leads us to think that this
phenomenon is not based on enthalpy, i.e. the attractions of the solute particles
to the solvent particles are not important as the solute is assumed to remain in
solution while the pure solvent freezes or boils.
So how do we explain this phenomenon? It is usually done by saying that the mixing
of the solute with the solvent increases the entropy of the solution making the
freezing process involve a larger negative change in entropy, and the boiling
process involve a smaller positive change in entropy. This requires a lower
temperature for the freezing to occur and a higher temperature for the boiling to
occur as a higher temperature will exacerbate any entropy effects and a lower
temperature will lessen any entropy effects. The role of entropy change in phase
changes can be explained from the Gibbs equation: DG = DH -TDS, where any spontaneous
change will only occur if DG is negative.
Incidentally, this topic is one of those "red flags" in high school: it is VERY easy
to have this taught in a misleading way, involving ions "holding onto" water molecules,
preventing boiling or freezing, which, although easy to grasp, is not a helpful model
and is not backed up by experimental observation. The entropy argument is a bit
involved, though!
Best wishes,
Tom Collins
Angela,
In most textbooks, the freezing point depression or boiling point
elevation equation is written as:
dT = km; where dT is the change in freezing or boiling point, k is
the constant and m is the molality.
A more complete equation is:
dT = ikm; where i is the van't Hoff factor which gives how much
ionization is happening in solution.
In dilute solutions, typically, the i can be estimated directly from
how many ions can be seen in the molecular formula.
For example: NaCl = 2; MgCl2 = 3; Fe(NO3)3 = 4.
From this you can figure out why electrolytes - on a per mole
basis, or of the same molality - can cause higher changes in
boiling or freezing points.
Greg (Roberto Gregorius)
Canisius College
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