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<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif">I sent this out sometime last week and I addressed #1 and #3 in<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif">class. I've included explanations for all the questions below.<br>
<br>
Dr. Zellmer <o:p></o:p></span></p>
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*******************************************************<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif">Here are three interesting questions you should be able to answer<br>
at this point based on things which have been discussed in<br>
chapters 11 and 13.<br>
<br>
1)<br>
<br>
Last night I removed a bottle of pop (Coke, Pepsi, etc.) from my<br>
freezer. The liquid showed no signs of ice formation. I swirled<br>
the bottle a little to equilibrate the temperature. As soon as<br>
I opened the bottle the pop turned to slush (i.e. ice formed).<br>
<br>
Why? Give a good reasonable explanation. There is one from<br>
chapter 11 and one from chapter 13. It is not a supercooled<br>
liquid.<br>
<br>
If you want to try this for yourself, here's what to do.<br>
Take a BOTTLE of pop (you need to be able to see inside).<br>
Put it in the freezer. Check on it periodically. Let it<br>
freeze a little. Take it out and let it warm up until the<br>
last piece of solid melts (swirl gently). As soon as this<br>
occurs open it and see what happens.<br>
<br>
<span style="color:red">This could occur for two reasons. The first is because of the<br>
way the m.p. line leans for H2O and it's solutions, it leans to<br>
the left. That means as pressure inc. at a constant temp the<br>
solid actually melts. Beverages are bottled at about 5 atm of<br>
pressure for the CO2. This could possibly put it in the liquid<o:p></o:p></span></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">region, depending on the temp. When it's opened the pressure<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">drops to 1 atm total pressure (way less than this for the partial<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">pressure of CO2) and it's possible it drops below the m.p. line<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">and freezes.<br>
<br>
A more probable reason deals with colligative prop. It's bottled<br>
at high pressures of CO2. The CO2 is nonpolar but pretty soluble<br>
in water, particularly because it reacts with water to produce<br>
carbonic acid, H2CO3. Since this is a soln its f.p. is lower than<br>
that of pure water. You put it in the freezer and it won't freeze<br>
until it's below the f.p. of pure water. When you take it out before<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">it freezes and open it CO2 comes out and the conc. of solute<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">decreases which raises the f.p. However, the soln has not had a<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">chance to warm up. If the f.p. rises enough it could go above the<o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-size:12.0pt;font-family:"Arial",sans-serif;color:red">temp of the soln in the bottle and it will freeze.</span><span style="font-size:12.0pt;font-family:"Arial",sans-serif"><br>
<br>
2)<br>
<br>
What's one reason to use tap water rather than distilled<br>
water in a steam iron?<br>
<br>
<span style="color:red">Tap water has a higher boiling point and can thus get hotter<br>
in the iron w/o boiling. The steam coming out will be at a<br>
higher temp. making it easier to get rid of wrinkles.<br>
</span><br>
3)<br>
<br>
When it gets really cold out and there's a bad snowstorm<br>
they use CaCl2 instead of NaCl on the roads. Why?<br>
<br>
<span style="color:red">Look at "i", the van't Hoff factor, assuming ideal behavior<br>
(no ion-pairing).<br>
<br>
CaCl2 i = 3<br>
NaCl i = 2<br>
</span><br>
<span style="color:red">So for the same conc. the CaCl2 would have a greater conc.<br>
of particles and a larger freezing point depression (lower<br>
freezing pt). It's more expensive than NaCl so it's not normally<br>
used unless it's really cold. It can also cause more damage to<br>
concrete, particularly new concrete.<br>
<br>
Another reason to use CaCl2 is because when it dissolves the process<br>
releases heat. A slurry of CaCl2 (a saturated soln with undissolved<br>
CaCl2) is made and dropped on the ice. This heat can help melt<br>
some of the existing ice and as it melts some of the undissolved<br>
CaCl2 dissolves. The resulting solution now has a lower f.p.<br>
<br>
In any case, below certain temps neither will work. Then they<br>
dump sand for traction (but it doesn't lower the f.p. because<br>
sand isn't soluble in water, otherwise we wouldn't have beaches).</span> <br>
<br>
<br>
Dr. Zellmer <o:p></o:p></span></p>
<p class="MsoNormal"><span style="font-family:"Arial",sans-serif"><o:p> </o:p></span></p>
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