Name____________________________________________Date_________________Per______ST#______
BOIL WATER LAB
1. Hypothesis: ______________________________________________________________
____________________________________________________________________________________
2. How many cal/gram does it take to vaporize liquid water at any temperature? _____________________
3. How many dietary Calories is your answer to the prior question? _______________________
4. How many cal/g does it take to raise the temp of 1 mL of water by 1oC while between 0 and 100oC?____
5. Show a calculation for raising the 300mL of water used in the lab from room temp to boiling temp (20oC
to 100oC)
6. Did bubbles form before the temp reached 100oC? _______________ What does that mean? _________
____________________________________________________________________________________
7. What was the salt supposed to do to the boiling temp? _________________________ Note: This has to
do with the solute ( Salt or water? Circle one ) essentially “getting in the way” of the water as it tried
to vaporize out of the solvent ( Salt or water? Circle one )
8. What is in the bubbles in the boiling water? _________________________________________________
9. What happened to the temp after all of the water in the beaker reached a vigorous boil (rolling boil)?
____________________________________________________________________________________
10. What are 3 reasons for your prior answer?
a. _________________________________________________________________________________
b. _________________________________________________________________________________
c. _________________________________________________________________________________
11. State whether or not your hypothesis was supported. Use lab data to back up your point (short conclusion)
____________________________________________________________________________________
____________________________________________________________________________________
____________________________________________________________________________________
12. Create a diagram that shows the relationship between vaporization, evaporation, and boiling.
ICE LAB
13. Hypothesis: ______________________________________________________________
____________________________________________________________________________________
14. How many cal/gram does it take to freeze liquid water at zero degrees C? _____________________
15. How many dietary Calories is your answer to the prior question? _______________________
16. How many cal/g does it take to lower the temp of one mL of water by 1oC anywhere between 0 and
100oC? ________
17. Show a calculation for changing the 220 mL of water used in the lab from room temp to freezing temp
(20oC to 0oC)
18. What happened to the level of the water in the graduated cylinder when the ice melted?
VolI ______________ VolF ________________ Change in vol _________________
19. What do you call the curve of the top of the water in the gradated cylinder?______________________
20. Why is that curve at the top of the water? __________________________________________________
____________________________________________________________________________________
21. What is the density of liquid water?___________________ Density of ice? ___________________
22. How much mass does 20g of salt and 80 mL of water have?__________g. How much space does the salt
water occupy when it’s mixed up (what’s the volume)? _____________. How do you account for the
difference? __________________________________________________________________________
23. Explain your answer to question 18.______________________________________________________
___________________________________________________________________________________
___________________________________________________________________________________
24. Draw an explanation picture of the lab, too.
http://www.chem.wisc.edu/deptfiles/genchem/sstutorial/Text11/Tx117/tx117.html
Physical Properties of Solutions
The physical properties of a solution are different from those of the pure solvent. Many
differences in physical properties are predictable if the solute in the pure state is nonvolatile that is, if it has a very low vapor pressure. Sugar, sodium chloride, and potassium nitrate are
examples of nonvolatile solutes. Colligative properties are those physical properties of solutions
of nonvolatile solutes that depend only on the number of particles present in a given amount of
solution, not on the nature of those particles. We will consider four colligative properties: vapor
pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure.
A. Vapor Pressure Lowering
At any given temperature, the vapor pressure of a solution containing a nonvolatile solute is less
than that of the pure solvent. This effect is called vapor pressure lowering. The solid line in
Figure 11.6 is a plot of the vapor pressure of pure water versus temperature. The break in the
curve at 0°C is the intersection of the curve of the vapor pressure of the solid with the curve of
the vapor pressure of the liquid. The dashed line in Figure 11.6 is a plot of the vapor pressure of
an aqueous solution of sugar verses temperature. Notice that the vapor pressure of the solution is
always less than that of the pure solvent. What causes this difference?
FIGURE 11.6 The vapor pressure of pure water is shown as a
solid line; the vapor pressure of an aqueous solution is shown
as a dashed line. Note the differences between the solution
and the pure substance in melting point and boiling point.
The surface of a pure solvent (Figure 11.7a) is populated only by solvent molecules. Some of
these molecules are escaping from the surface, and others are returning to the liquid state. The
surface of a solution is populated by two kinds of molecules; some are solvent molecules,
FIGURE 11.7 Vapor pressure lowering: (a) the vapor pressure of
a pure liquid; (b) the vapor pressure of a solution. In (b), the
number of solvent molecules on the surface of the liquid has been
decreased by the presence of the solute molecules. Fewer solvent
molecules can vaporize, and the vapor pressure is lower.
others are solute molecules. Only the solvent molecules are volatile. They alone can escape to
build up the vapor pressure of the solution. There are fewer solvent molecules on the surface of
the solution than on the surface of the pure liquid. Fewer will vaporize and, as a consequence,
the vapor pressure of the solution will be less than that of the pure liquid at the same temperature
(see Figure 11.7b).
B. Boiling Point Elevation
The boiling point of a substance is the temperature at which the vapor pressure of the substance
equals atmospheric pressure. A solution containing a nonvolatile solute, having a lower vapor
pressure than the pure solvent, must be at a higher temperature before its vapor pressure equals
atmospheric pressure and it boils. Thus, the boiling point of a solution containing a nonvolatile
solute is higher than that of the pure solvent (see Figure 11.6) This effect is called boiling point
elevation.
C. Freezing Point Depression
Recall that freezing and melting point are two terms that describe the same temperature, the
temperature at which the vapor pressure of the solid equals the vapor pressure of the liquid and
at which the solid and the liquid are in equilibrium. Remember, too, that vapor pressure
decreases as the temperature decreases. The vapor pressure of a solution is lower than that of the
solvent, so the vapor pressure of a solution will equal that of the solid at a lower temperature
than in the case of the pure solvent. Thus, the freezing point will be lower for a solution than for
the pure solvent (see Figure 11.6). This effect is called freezing point depression. Remember
that, just as it is the solvent that vaporizes when a solution boils, it is the solvent, not the
solution, that becomes solid when a solution freezes. When a salt solution freezes, the ice is pure
water (solid); the remaining solution contains all the salt.
Application of this principle leads us to add antifreeze (a nonvolatile solute) to the water in the
radiators in our cars. We thus lower the freezing point of the solvent (water), and the solution
remains a liquid even at subfreezing temperatures.
Boil Water at Room
Temperature
How To Boil Water at Room Temperature Without Heating It
(http://chemistry.about.com/od/foodchemistryfaqs/f/Does-Adding-Salt-Lower-TheBoiling-Point-Of-Water.htm)
If you pull back on a syringe containing water, you create a vacuum which boils the water at room temperature. Chris
Stein, Getty Images
Anne Marie Helmenstine, Ph.D.
Chemistry Expert
SHARE
PIN
You can boil water at room temperature without heating it. Here's an easy way to see this for yourself.
Materials

water

syringe
You can get a syringe at any pharmacy or lab. You don't need the needle, so it's safe project, even for
kids.
How To Boil Water Without Heating It
1. Use the plunger to pull up a bit of water into the syringe. Don't fill it -- you need air
space in order for this to work. You just need enough water that you can observe
it.
1. Next, you need to seal the bottom of the syringe so that it won't be able to suck up
more air or water. You can put your fingertip over the opening, seal it with a cap (if
one came with the syringe), or press a piece of plastic against the hole.
2. Now you'll boil the water. All you need to do is pull back as quickly as you can on
the syringe plunger. It may take a couple of tries to perfect the technique, so you
can keep the syringe still enough to watch the water. See it boil?
How It Works
The boiling point of water or any other liquid depends on vapor pressure. As you lower the pressure,
the boiling point of the water drops.
You can see this if you compare the boiling point of water at sea level with the boiling point of water on
a mountain. The water on the mountain boils at a lower temperature, which is why you see high-altitude
instructions on baking recipes!
When you pull back on the plunger, you increase the amount of volume inside the syringe. However,
the contents of the syringe can't change because you have sealed it.
The air inside the tube acts as gases do and the molecules spread out to fill the whole space. The
atmospheric pressure inside the syringe drops, creating a partial vacuum. The vapor pressure of the
water becomes high enough compared to the atmospheric pressure that the water molecules can
easily pass from the liquid phase into the vapor phase. This is boiling.
Compare it with the normal boiling point of water. Pretty cool. Any time you lower the pressure around
a liquid, you lower its boiling point. If you increase the pressure, you raise the boiling point. The
relationship is not linear, so you would need to consult a phase diagram to predict how great the effect
of a pressure change would be.
Does water boil faster if you put salt in the
water? (http://www.swri.org/10light/water.htm)
Yes and no. If you look at how fast water boils when you add a small amount of salt to
it, such as when cooking your noodles, the change is insignificant between pure water
and the salted water. However, if you take two identical pots and add one gallon of pure
water to one pot and one gallon of 20 percent salt water to the other and heat the two
pots on identical stoves, the pot containing the salt water will come to a boil first.
Surprised?
To truly answer the question, one must look at what it takes to boil a container of water.
The time it takes a bucket of liquid to boil is controlled by essentially three things. The
first is how much heat or energy you put into the bucket. The second is how fast the
temperature rises in response to the heat input (the liquid's heat capacity), and the third
is the boiling point of the liquid. Assuming that we can control our stoves and add the
same amount of energy to each pot, this variable becomes insignificant.
The boiling point of water does rise if you add salt to it, but only by about 2°C (4°F) to
102°C (216°F). Remember, water boils at 100°C (212°F). This is an insignificant
change for adding such a large amount of salt. For you science nerds out there, the
boiling point increase is calculated using the "ebullioscopic" constant of water. This
leads us to the important variable, how fast water or salt water heats up, or the
solution's heat capacity.
The heat capacity of water is very high. What this means is that it takes a lot of energy
to raise the temperature of water 1°C; in fact, the calorie is defined as the amount of
energy that it takes to heat one gram of water to 1°C. Not to digress, but the high heat
capacity of water is good, especially if you live on a planet where two-thirds of the
surface is covered by water - it helps regulate the global temperature.
Now back to the question. If you look at the heat capacity of salt water, you will find that
it is less than pure water. In other words, it takes less energy to raise the temperature of
the salt water 1°C than pure water. This means that the salt water heats up faster and
eventually gets to its boiling point first.
Why does salt water have a lower heat capacity? If you look at 100 grams of pure
water, it contains 100 grams of water, but 100 grams of 20 percent salt water only
contains 80 grams of water. The other 20 grams is the dissolved salt. The heat capacity
of dissolved salt is almost zero when compared to the high heat capacity of water. This
means that the heat capacity of a 20-percent salt solution is 80 percent that of pure
water. Twenty percent salt water will heat up almost 25 percent faster than pure water
and will win the speed race to the boiling point.
Please note that this will not hold true if you take two identical pots containing one
gallon of water each and add the salt to one pot because then the volume of liquid in
the salted pot will be greater than the one gallon starting point.
This month's Whizard is Mike Dammann, manager of the Inorganics Section in the
Chemistry and Chemical Engineering Division.
http://van.physics.illinois.edu/qa/listing.php?id=16388
Baloney with salt water and boiling temps.
http://www.physlink.com/Education/AskExperts/ae643.cfm
Our first Baloney example (an answer with five major errors)comes from a generally good site which happened to slip pretty
badly on a few questions. The site administrator was informed of the difficulties with the answer some weeks ago. More
baloney will follow. We invite you to send in examples. The Web is a target-rich environment. "Question What does salt do to
water that raises the boiling temperature? Is it a chemical reaction?"
- Mike W. (age 61)
A:
From the posted answer:
1. "...it is a conservation of momentum. "
Conservation of momentum has nothing to do with this question. Since there are external forces (e.g. with a pot holding the
water) momentum wouldn't have to be conserved. At any rate, in this case there is no change in momentum as the water
boils, since the molecules in either the liquid or the gas are randomly heading all directions, giving zero average momentum.
2. "...for something to boil, enough energy must be absorbed by it to cause vibrations large enough to enhance the kinetic
energy of each molecule to the point where they break away..." "...the energy goes into not only exciting each water
molecule to a higher kinetic energy but also each salt molecule..."
The whole issue of the solute soaking up energy is irrelevant, since the question is not how much energy is required to boil
the water but rather what the boiling temperature is. If the solute did soak up some energy, that would temporarily lower the
temperature, not change the boiling point.
3. "...more massive salt molecules themselves need a larger contribution of energy..." Although that whole issue is irrelevant,
it's also handled wrong. More massive particles do not need more energy to be at a given temperature. In fact, the
equipartition theorem says that they have the same kinetic energies as lighter particles at a given temperature.
4. The freezing point will be lowered by "...even better, calcium chloride, an even heavier compound..." Not that it's relevant,
but it's not even true that dissolving salts in water consistently soaks up energy (enthalpy, to be precise). Some salts release
enthalpy (exothermic), some soak it up (endothermic). Oddly enough, the particular salt mentioned (CaCl2) is one of the
exothermic ones, releasing heat as it dissolves.
5. "...even better, calcium chloride, an even heavier compound..." It's not close to being true that more massive ions raise
the boiling point more than less massive ones. In fact, there's a pretty good rule of thumb that the boiling point effect
depends only on the number of solute particles, not their detailed properties, and certainly not particularly on the mass. The
rule becomes exact for dilute solutes.
The real physical chemistry involves a fundamental thermodynamic argument that any stable solute will raise the boiling
point, regardless of microscopic details. At the microscopic level, the reason that most solutes have about the same effect
per solute particle is that the main effect usually has to do with how the entropy of the solute goes down as the solvent
evaporates, not the detailed and variable ways in which the energies change. A search of our site will turn up some info on
this issue.
e.g.
http://van.physics.illinois.edu/qa/listing.php?id=1447
http://van.physics.illinois.edu/qa/listing.php?id=1464
Mike W.
\
Follow-Up #1: Why does salt water have a higher boiling point than distilled
water?
https://van.physics.illinois.edu/qa/listing.php?id=1447
Q:
Why does salt water have a higher boiling point than distiller water? Can you please explain in terms of the positive
and negative charges of the particles and the rubbing off of electrons please? Thank you. My teacher taught it to us
but I missed the lesson.
- Ash (age 14)
Australia
A:
Ash- I've marked your question as a follow-up to a similar question, which has a version of the answer.
Your teacher may have given an answer "in terms of the positive and negative charges of the particles and the rubbing off of
electrons", but that sounds fundamentally wrong. Any solute in water raises the boiling point, so long as the solute stays in
the liquid water. It's true that part of why salt dissolves well in water is that it falls apart into charged particles, but some
uncharged molecules also dissolve in water and also raise the boiling point.
Another way of putting the answer is to say that if the solute stays in the liquid, there is less room for it to find various
different states as the liquid boils away. When not so many states are available, we say that the "entropy" is reduced. The
basic rule that tells us what will happen is that nature always heads toward an increase in entropy. So having solutes in there
goes against boiling, which doesn't then occur until the temperature is higher. At the higher temperature it turns out that
boiling still increases net entropy, thanks to the water molecules getting more space to run around.
By the way, something about this topic seems to bring out goofy answers. You might enjoy this:
http://van.physics.illinois.edu/qa/listing.php?id=16388
Mike W.
Most recent answer: 02/23/2013
https://van.physics.illinois.edu/qa/listing.php?id=1447
Q:
Why does adding salt make the boiling temperature of water rise???
- soeun lee
Auckland Girls Grammar, Auckland, New Zealand
A:
Souen- That's a good question. It turns out to be easy to give an answer to someone who's studied a little Statistical
Mechanics, but I'll try to give an answer that doesn't assume that sort of background.
Salt (or other solutes, like sugar) can easily dissolve in liquid water. However, taking the solute out of the water and putting it
in the gas phase (air) requires a lot of energy. At temperatures around the water boiling point, these solutes stay in the
liquid.
Now the total pressure in the liquid and the air at the boundary are the same- otherwise one would push the other into a
smaller space. Part of the pressure in the liquid comes from the solutes, not the water. So the pressure due to the water
alone is reduced compared to that of pure water at the same temperature. The vapor pressure, meaning the pressure of
water vapor that would stay in equilibrium with the liquid, is reduced by the same amount because of the solutes. (I've
simplified and approximated a little here, since the pressure doesn't quite break up into separate parts due to the salt and the
water.)
Water boils when the vapor pressure of the water gets to be as big as the pressure of the atmosphere. At that point, vapor
bubbles in the water can grow. You have to heat the liquid with solutes up more to get the vapor pressure in it to equal the
atmospheric pressure, so it has a higher boiling point.
A very similar argument explains why solutes also lower the freezing point. Since the solutes are almost completely excluded
from the solid (like from the gas) they stabilize the liquid. A search of this site will turn up some answers about freezing salt
water.
Mike W. (and Tom J.)