Three interesting questions to ponder - answers

[Zellmer, Robert]

I sent these out a couple of weeks ago.  We only discussed #1 in
class, although due to the cancellation of classes I'm not sure I
covered it in both classes.  Plus, I only discussed #1.  I actually
did discuss #3 in class.  I've included explanations below.

Dr. Zellmer

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Here are three interesting questions you should be able to answer
at this point based on things which have been discussed in
chapters 11 and 13.

1)

Last night I removed a bottle of pop (Coke, Pepsi, etc.) from my
freezer.  The liquid showed no signs of ice formation.  I swirled
the bottle a little to equilibrate the temperature.  As soon as
I opened the bottle the pop turned to slush (i.e. ice formed).

Why?  Give a good reasonable explanation.  There is one from
chapter 11 and one from chapter 13.  It is not a supercooled
liquid.

If you want to try this for yourself, here's what to do.
Take a BOTTLE of pop (you need to be able to see inside).
Put it in the freezer.  Check on it periodically.  Let it
freeze a little.  Take it out and let it warm up until the
last piece of solid melts (swirl gently).  As soon as this
occurs open it and see what happens.

I did address this one in lecture on Tue.

This could occur for two reasons.  The first is because of the
way the m.p. line leans for H2O and it's solutions, it leans to
the left.  That means as pressure inc. at a constant temp the
solid actually melts.  Beverages are bottled at about 3-5 atm of
pressure.  This could possibly put it in the liquid region, depending
on the temp.  When it's opened the pressure drops to 1 atm
and it's possible it drops below the m.p. line and freezes.

A more probable reason deals with colligative prop.  It's bottled
at high pressures of CO2.  The CO2 is nonpolar but pretty soluble
in water, particularly because it reacts with water to produce
carbonic acid, H2CO3.   Since this is a soln its f.p. is lower than
that of pure water.  You put it in the freezer and it won't freeze
until it's below the f.p. of pure water.  When you take it out and
open it CO2 comes out and the conc. of solute decreases which
raises the f.p.  However, the soln has not had a chance to warm
up.  If the f.p. rises enough it could go above the temp of the soln
in the bottle and it will freeze.

2)

What's one reason to use tap water rather than distilled
water in a steam iron?

Tap water has a higher boiling point and can thus get hotter
in the iron w/o boiling.  The steam coming out will be at a
higher temp. making it easier to get rid of wrinkles.

3)

When it gets really cold out and there's a bad snowstorm
they use CaCl2 (or MgCl2) instead of NaCl on the roads.  Why?

Look at "i", the van't Hoff factor, assuming ideal behavior
(no ion-pairing).

CaCl2    i = 3
NaCl    i = 2

So, for the same conc. the CaCl2 would have a greater conc.
of particles and a larger freezing point depression (lower
freezing pt).  It's more expensive than NaCl so it's not normally
used unless it's really cold with lots of snow.  It can also cause
more damage to concrete, particularly new concrete.

Another reason to use CaCl2 is because when it dissolves the process
releases heat.  A slurry of CaCl2 (a saturated soln with undissolved
CaCl2) is made and dropped on the ice.  This heat can help melt
some of the existing ice and as it melts some of the undissolved
CaCl2 dissolves.  The resulting solution now has a lower f.p.

In any case, below certain temps neither will work.  Then they
dump sand for traction (but it doesn't lower the f.p. because
sand isn't soluble in water, otherwise we wouldn't have beaches).


Dr. Zellmer
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