evaporating experiment report

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Science Fair Project Ideas for Kids, Middle & High School Students ⋅

Simple Evaporation Experiments

Viewing Evaporation Experiment for Kids

Evaporation happens when liquids turn into vapors. You can often see water evaporate on a hot day. In addition, there are fun and simple evaporation experiments you can do at home to learn more about the process. Read on and try the following evaporation science experiments.

Experiment With Covered and Uncovered Jars

Fill two identical mason jars with water. Leaving one of the jars uncovered, cover the other one with an improvised aluminum foil lid. Make the lid as secure as possible. Then, take the jars outside and place them both in an equally sunny spot. Draw a picture of the jars, noting the current water levels. Return to the experiment every day for the next week to observe and draw the current state of the water jars. You will observe that the water in the uncovered jar “disappears” more every day, while the water in the covered jar evaporates at a much slower rate because the evaporation process gets blocked by the aluminum foil.

Experiment With Sun and Shade

After filling up two identical bowls with water, take them outside and locate a spot where direct sunlight and shadow stand side by side. Place one water bowl in the direct sunlight, and the other beside it in the shade. Observe both bowls and use pencil and paper to illustrate current water levels in each bowl. Return to the experiment every hour for the rest of the day, continuing to make observations and illustrations of the water levels. You will see that the water in the bowl placed in direct sunlight evaporates much more quickly than the shaded water due to the higher levels of heat, which increase molecular activity in the water, thus expediting evaporation.

Experiment With Wet Cloth

Wet two identical pieces of cloth and wring the excess water out. Place one of the pieces of cloth in an airtight plastic bag. Place the other piece of cloth in an open tray. Position both items near a window with plenty of sunlight. Make predictions regarding which item will dry up first: the cloth in the sealed bag, or the one exposed to the air. Leave the items by the window overnight. When you return to the experiment the next day, you’ll see that the exposed cloth dried up, while the one sealed inside the bag remains moist. This is because the water molecules in the sealed cloth can’t escape into the air like the ones in the exposed cloth.

Experiment With Salt Water

Add a decent amount of salt to a large glass of water. Then, pour the salty water onto a sheet of black construction paper placed inside a baking tray. If necessary, weigh down the paper with rocks or waterproof paper weights. Place the tray outside in a beam of direct sunlight. Predict what will happen to the water and salt. In a few hours, return to the tray to discover the outcome of the experiment. You will see that the water is gone, and that the salt remains on the black paper. The water disappeared due to the process of evaporation, but the salt stayed because it would require a lot more energy than provided by the sunlight to fully evaporate.

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About the Author

Bill Reynolds holds a Bachelor's degree in Communications from Rowan University. He has written hundreds of articles for print and online media, drawing inspiration from a wide range of professional experiences. As part of the UCLA Extension Writer's Program, he has been nominated for the James Kirkwood Prize for Creative Writing.

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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

$\begin{equation*} \log_{10} P^\star = A + \frac{B}{T} \end{equation*}$

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?