water vapour in experiment

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EXPERIMENT 11: VAPOUR PRESSURE

Introduction.

Some molecules of a liquid can escape from the liquid surface (evaporate) because of their kinetic energy.  If these molecules, now in the vapour phase, are collected in a closed container, they will exert a pressure that is known as the vapour pressure, P * , of the liquid, as indicated in Figure 1.

Figure 1 : The microscopic process of evaporation and condensation at the liquid surface, by HellTchi , licensed under CC BY-SA 3.0

As the temperature increases so does the kinetic energy of molecules.  Then more molecules escape the liquid phase (evaporate) and the number of molecules above the liquid increases, yielding a higher vapour pressure.  Consequently, the vapour pressure of a liquid depends heavily on temperature.  The higher the temperature, the higher the vapour pressure.  The relationship is not linear, though, but follows the expression

where T is the absolute temperature, and A , B are constants for any given liquid.

When the temperature is such that the vapour pressure of a liquid reaches the outside pressure (which is often atmospheric pressure), the liquid will start boiling.  When boiling happens, if additional energy is supplied, the temperature and pressure will remain constant until all the liquid evaporates, unless the system is in a closed container.  In a closed container, the generated vapours will increase the pressure and the liquid will stop boiling (because the container pressure will be higher than the vapour pressure), unless its temperature also increases.

To measure the vapour pressure of a liquid we exploit this principle that the liquid boils when its temperature is such that the vapour pressure is equal to the system pressure.  We set the pressure of the system and slowly heat-up the  liquid until it starts boiling.  If the process happens slow-enough, as it should be, it is difficult to visually assess when boiling happens (there will be no vigorous bubbling).  Therefore, we rely on temperature to assess boiling: boiling happens when the temperature of the liquid stays constant while being heated because all the added energy is used for phase change.

If we use this technique to calculate the vapour pressure at a few temperature levels, we can then employ equation (1) to calculate the vapour pressure at any temperature.  The easiest way to do this is by plotting log 10 P versus 1/T which yields a straight line with intercept A and slope B.

The apparatus used in the E030 lab is a simple one consisting of parts that can be found in any chemistry lab.  The main parts of the apparatus are shown in Figure 2.  Essentially, a small flask is placed under low (vacuum) pressure using the lab’s vacuum line.  Methanol is added into this low-pressure space and is heated through a hot-water bath.  The methanol’s temperature is constantly monitored to assess when the methanol boils under the system’s pressure, which, as mentioned above, provides the boiling temperature at the system’s pressure or the vapour pressure at this temperature.  This process is continuously repeated at increasing pressures to provide a set of vapour pressures at different temperatures.

Figure 2: Vapour pressure apparatus in E030

The purpose of the experiment is:

• To understand what vapour pressure is and how it changes with temperature.
• To understand the relationship between vapour pressure at a given temperature and boiling temperature at a given pressure.
• To determine the vapour pressure of a pure liquid at various temperatures.

Before proceeding, check your understanding by performing the following drag-and-drop task.

You are given a set of temperatures and vapour pressures for a substance.  The data sets are not ordered; so, you must match the temperature with the vapour pressure.  Recall that the relationship is monotonic: higher temperature leads to higher vapour pressure.

• Ensure that the stopcock from the funnel is closed.  Place about 10 mL of methanol into the funnel of the vapour pressure apparatus. Set the water bath around the flask and start its stirring but do not heat yet.
• Connect the apparatus to the vacuum tap and turn on the tap.  Evacuate the entire apparatus until the pressure inside is about 120 mmHg.  *NOTE : Make sure there are no leaks in the apparatus by observing the “vacuum pressure”; it should be constant.
• Allow a small amount of liquid (enough so that there is visible liquid in the flask) from the funnel to run onto the cotton fibre around the thermometer bulb. Heat the water bath until the temperature is 5 to 10°C above the thermometer reading in the flask.  Heat the water slowly, as you do not want to overshoot the bath temperature.
• Start recording (1) the temperature of the flask thermometer, (2) the temperature of the bath, and (3) the system pressure every 30 seconds. (three readings every 30 seconds for the duration of the experiment).
• When the flask thermometer reading remains constant, the liquid on the cotton should be in equilibrium with its vapours at the pressure in the apparatus. This happens because all the energy transfer from the hotter bath to the methanol in the flask is used for evaporating the methanol (i.e. methanol boils at the apparatus pressure), which occurs at constant temperature.  Take a special note of this flask temperature and the corresponding pressure.  This constitutes a pair of vapour pressure at that temperature (or boiling point at that pressure).
• Increase the pressure into the apparatus by about 50 mmHg. This can happen by tightening the clamp on the hose leading to the vacuum tap.  Increasing the pressure will increase the boiling temperature of methanol.  Then, the temperature of the flask will start increasing as the methanol gets heated by the hotter bath (ensure that the hot bath is always 5-10°C hotter than the flask).   This temperature increase will stop when the boiling point of methanol at the new pressure is reached.  Take a special note of this flask temperature and the corresponding pressure .  This constitutes a new pair of vapour pressure at that temperature (or boiling point at that pressure).
• Repeat the above procedure several times until atmospheric pressure of approximately 760 mmHg is reached.
• When the experiment is finished, allow air into the apparatus until the pressure inside and outside are equalized. Disassemble the system and remove the cotton fibre.  Clean and dry the flask and put cold water back into the water bath.
• Repeat the experiment.

First we will look at the raw data collected during the experiment

• Provide two Tables, one for each run, with your recordings of time, flask temperature, hot bath temperature, and system pressure.
• Plot the data of these two Tables.  One graph for each run.  Place time on the x-axis.  The graph must have three sets of data/lines.  The two temperatures must be on one y-axis and pressure on another axis.
• Plot on one graph the vapour pressure and temperature pairs for each run.  The Table must have two sets of data, one for each run, with temperature on the x-axis.
• Get a literature value for the “Normal Boiling Point” (NBP) of methanol (clearly state your source) and add it to the graph created in step 4.
• What is the relationship between vapour pressure and temperature?
• Are there differences between the two runs?  What are any causes of such differences?
• How well do your experimental measurements compare to the NBP of methanol? Why are there differences, if any?

We want to calculate the parameters A and B in order to be able to predict the vapour pressure of methanol and any temperature. For this task you should use the data collected during both runs .  Combine the vapour pressure – temperature data from both runs into one Table.  If your plot created during step 4 above clearly indicates that the data from one of the two runs is flawed, use data from only one run but clearly state that you are doing this.

• Tabulate: temperature, vapour pressure (P * ), absolute temperature, log(vapour pressure), inverse of absolute pressure.
• Create a plot of log(P * ) versus inverse of absolute temperature.
• Generate a linear trendline through the data.  Get Excel (or any graphing software you are using) to show the trendline equation on the plot.  Clearly state the values of A and B of equation (2) above.
• State the relationship between log(P * ) and 1/T.
• Vapour pressure of methanol at 50°C, and at 70°C
• The boiling temperature of methanol at 0.5 atm, 1 atm, and 1,2 atm
• Comment on how the boiling point at 1 atm compares with the literature NBP (normal boiling point) of methanol.
• Calculate the vapour pressure of methanol at 50°C and at 70°C using Antoine’s equation (if covered in class) and compare with your predictions during step 6 above. https://ecampusontario.pressbooks.pub/app/uploads/sites/562/2020/05/Video_2.mp4

PROCTECH 2CE3 Lab Manual Copyright © by Kostas Apostolou. All Rights Reserved.

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April 25, 2013

Steamy Science: Demonstrating Condensation

A fun physics demonstration from Education.com

By Education.com

Key concepts Physics Liquids Gasses Pressure

Introduction Ever wonder where those little drops of water on the outside of your cold can of soda pop or bottle of water come from? That’s condensation! Cold surfaces can cause water vapor in the air to cool down, condense and form tiny beads of liquid. The molecules in these miniscule droplets of water are grouped far more closely together than when they were in their gas phase, and exert less pressure—a fact that has some pretty cool physical implications.

Perhaps you have seen the classic science demonstration where a hard-boiled egg is “sucked” into a bottle using a match. The effect is definitely cool, but understanding how it works is tough. Air molecules are spaced differently and exert different levels of pressure depending on how hot or cold they are. This is a fun experiment where the physics are more observable, the effect more dramatic and the pyrotechnics totally unnecessary.

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Background Molecules, which make up everything around us—including air—are in a constant state of motion. The hotter water molecules become, the faster they move, turning from water (their liquid phase) to steam (their gas phase). When liquid water turns to gas, not only do the molecules move much faster, they also are spaced much farther apart. They spread out so much that they generate pressure by pushing on each other and everything else they come into contact with. What happens when we take the heat source away from that steam? The molecules form liquid water again. This is called condensation.

The air in our atmosphere is also a gas that exerts a fairly strong pressure of its own. This experiment will illustrate what can happen when the changing pressure of condensing steam goes up against the pressure of air, which remains relatively constant.

Materials • One large, thick plastic bottle with a wide neck (an empty, 64-ounce fruit juice bottle will work or a three-gallon water-dispenser jug is great). Use caution with thinner plastic containers—hot water can cause them to melt; and avoid glass—boiling water can cause glass to break. • Small, empty water balloons (Keep more than one handy, in case of breakage.) • Water • Stove • Oven mitt • Pot or teakettle for boiling water (Use caution and adult help when dealing with hot water.)

Procedure • Set a kettle or pot of water to boil on the stove. • While you’re waiting for your water to boil, fill your balloon full of water using a faucet or a hose. Don’t overinflate the balloon! It should be too large to slip through the neck of the bottle via gravity alone but not so large that it would burst were it to get pushed through. • Once your water reaches a rapid boil, very carefully pour it into your bottle to about a quarter of the way full. • Place the filled water balloon in the neck of the bottle. • Stand back and watch as the balloon gets sucked into the bottle. Knowing what we know now about water and steam pressure, why do you think this happens? • Extra: Try sketching a diagram that includes illustrations of what the air and water molecules look like during each phase of the experiment. Read “Observations and Results” below for some hints. • Extra: Suction is a misleading concept. Condensing steam doesn’t have attractive power of its own, like a magnet does. It doesn’t actually pull or suck the balloon into the bottle. When the steam molecules stop pushing out of the bottle, and stop pushing on the balloon, something else outside the bottle becomes strong enough to push the balloon into the bottle—and it’s not gravity. What might it be? • Extra: What happens if the balloon is too big? Why? Observations and Results When the water was heated, its molecules began to move rapidly, turning some into its gas phase: steam. When in a gas phase, water molecules are spaced much farther apart and take up more space. The pressures inside and outside the bottle reach a state of equilibrium, meaning that they are the same. Why? With the neck of the bottle unobstructed, the expanding steam can move from inside the bottle out into the surrounding air.

Here’s when everything changes: When the steam in the bottle starts cooling down and we place the balloon in the bottle’s neck. Without heat, the water molecules inside the bottle start condensing—that is, they start turning from steam back into liquid water. When matter turns from its gas phase back into its liquid phase, the molecules take up much less space and exert far less pressure. In fact, the condensing steam creates a partial vacuum—a region of much lower pressure than that of the surrounding atmosphere—inside the bottle. Remember, unlike the condensing steam the air outside the bottle doesn’t change, and still exerts a pressure of its own. We call the resulting difference between these two areas a pressure gradient. The pressures aren’t able to equalize easily because the balloon blocks the gases from flowing from one area into another. So what happens? The gas on the outside (air) pushes harder than gas on the inside (the condensing steam), so the balloon gets pushed—and pulled—into the bottle.

Another way to describe what happened is to use the word “suction,” because the water balloon was sucked through the neck and into the bottle. But suction can be a misleading concept! What we’re really talking about when we talk about “suction” is a liquid or gas force that pushes on something in the absence of an equal force pushing back. You can crunch an empty water bottle simply by sucking the air out of it. The outside air pressure is what causes the bottle to collapse, because you’ve removed the air inside that was pushing back!

More to explore Condensation Balloon Trick  , from ScienceFix.com Crunch a Can , from Education.com Balloon in a Bottle: An Air Pressure Experiment , from Education.com Balloon Air Pressure Magic , from Education.com

This activity brought to you in partnership with  Education.com

Experiment 24. Enthalpy of vaporisation of water from valour pressure measurement.

Introduction.

In this experiment, a sample of air is trapped over water in an inverted measuring cylinder in a beaker. When the temperature of the apparatus is changed the number of moles of water vapour in the gas phase will vary according to the Clausius-Clapeyron equation, while that of air will remain constant. The number of moles of air in the mixture can be found by reducing the temperature of the whole apparatus to about 5 C. At that temperature it can be assumed that the vapour pressure of water is so small that the volume of gas measured corresponds only to the air present. The enthalpy of vaporisation can then be calculated from a plot of ln p(H 2 O) (the vapour pressure) versus 1/T.

Apparatus Required:

Calculations, discussion questions.

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