Re: procedure to find heat of vaporization of water

*Excellent* question;  the subject is not usually discussed until the 
junior or senior year of undergrad in a physical chem course, and then 
again in the first year of grad school in physical chemistry.  Also, my 
apologies for taking so long to answer: a combination of computer problems 
and work requirements just made it impossible for me.  The procedure for 
measuring the Heat of Vaporization is easy in theory, but a pain in real 
life;  perhaps you can persuade your teacher to let you get out of doing 
the experiment.  I hope that the following discussion helps: 

1. Just so we're both talking the same thing, the vapor pressure is defined 
as "the pressure exerted when a solid or liquid is in equilibrium with its 
own vapor."  The "Heat of Vaporization" is defined as the "heat required to 
change a liquid to a vapor at identical temperature".

2. The Heat of Vaporization of water (and many other substances) is found 
in the "Handbook of Chemistry and Physics" [HCP]; your library should have 
it or any local college library will have it;  politely encourage your 
teacher to have a look.  This is how most people in the real world get this 
information.

3. To get the heat of vaporization, you need to measure the vapor pressure 
of the liquid at at least two, but preferably about a dozen for water, at 
different temperatures.  (The better you can control the temperature, the 
better your heat of vaporization results will be.)  Then you graph your 
results and you should get a straight line.  Take two widely separated 
points on this line and plug the values into an equation I will present 
below, and behold! the vapor pressure will pop out.  In other words, it's 
extremely simple to state the procedure, but a real pain to do it 
correctly.

4.  There are many ways of calculating or measuring vapor pressure, 
ranging from the very simple to the extremely complex.  A simple but 
pretty  accurate measurement technique is found in Sienko & Plane, 
"College Chemistry", page 63.  What you do--and please both refer to 
the text *AND DO THE WORK USING RUBBER GLOVES UNDER A FUME HOOD*--is to 
take a long tube filled with mercury 1 cm in diameter (to simplify the math 
involved) least 40 cm long and invert it into a basin of mercury, so that 
you've created a vacuum at the top.  Note this is the same as a mercury 
barometer. Carefully note the height of the mercury column in mm--to 
tenth's of mm if you can.  Then take a U-shaped bent glass dropper and 
inject a just a small amount of water into the column, so that the liquid 
floats to the top.  All you need is for a small droplet of water to be 
showing at the top.  The mercury level will drop due to the vapor pressure 
of the liquid;  just measure the distance that the column of mercury drops 
and that's the vapor pressure of the liquid at that temperature.  AGAIN, do 
this under a fume hood and wearing rubber gloves to protect yourself from 
the mercury vapor--nasty stuff.  For more detailed (and accurate) vapor 
measurement techniques and explanation, try W. J. Moore, Physical 
Chemistry, page 106; available in your library or any college library.  
Important point for accuracy:  do all the measurements on the same 
afternoon, so barometric pressure changes will not affect your results.  Do 
the experiment first at room temperature. Your teacher may complain that 
the vapor pressure of the mercury screws things up, but it's negligible at 
these temperatures (Moore, Table 4.1, page 107).

5. Now for the really hard part of the experiment:  you need to change the 
temperature of the mercury column from about +5 to +90 C, so that the water 
 vapor pressure changes as well.  The high temps can be done by wrapping 
the column with a heating tape connected to a rheostat, and using some sort 
of thermometer rig to measure the temperature.  ONE LAST TIME:  DO THIS 
UNDER A FUME HOOD WEARING GLOVES, and be careful about electricity and 
burning yourself.  The cold temperatures can probably be done by using a 
fan blowing over a mixture of salt and ice;  this will be tricky, but it 
can be done:  just be sure to measure the mercury column temperature 
carefully.

6.  Now we have the vapor pressures of water at various temperatures.  A 
single formula is all that is necessary with the above data.  The equation 
for calculating the heat of vaporization is called the Clapeyron-Clausius 
equation found in Moore, pages 14,104--105 and working from that we get the 
final equation:

  (ln(P2/P1)*8.3143)/((1/T2)-(1/T1)) = Heat_of_Vaporization! in Joule/mole
   
  P1 is the vapor pressure measured at T1; P2 is the v.p. at T2
  All temperatures are expressed in K;  K=C+273.15
Please note that the heat of vaporization is a function of temperature.   

So, IF you know (or can measure as we did) the vapor pressure of water at a 
number of temperatures you can easily calculate the Heat of Vaporization. 
Same hold true for any other liquid.  And if you use a solid like 
napthalene (mothballs) you can get the same thing, only then it's called 
the Heat of Sublimation.

6. I stated that the calculation was easy, but the experimental procedure 
was no fun.  But thank you for asking this *very* good question;  it took 
me back 25 years and was a lot of fun to figure out the answer!