photosynthesis experiment soda lime

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Experiment to show that carbon dioxide is needed for photosynthesis

In my school test, the question is as follows:

• Will the experiment work satisfactorily?
• What alterations can you make to obtain expected result?

What I wrote:

Yes, it will work satisfactorily. Lime water absorbs carbon dioxide and hence there is no carbon dioxide for the leaf and hence photosynthesis does not take place and no starch is prepared. Hence on iodine test, the presence of starch is negative and thus proves that carbon dioxide is required for photosynthesis.

We can put half of the leaf inside the flask and half of it outside. So, the inside part will not show presence of starch and the outside part will show presence of starch.

[Note]: It was marked incorrect by my biology teacher. It got it reviewed by another biology teacher and HE SAID WHAT I WROTE IS CORRECT and the teacher who marked it wrong gave the opinion as follows:

It will not work satisfactorily because lime water does not absorb carbon dioxide. [which to me, seems incorrect as lime water ABSORBS carbon dioxide]

The alteration is that you have to use Potassium Hydroxide instead of lime water. [which to me, seems incorrect too since both lime water and potassium hydroxide absorbs carbon dioxide and then there is no need to use potassium hydroxide when lime water works!]

Now I am confused what is correct and what should I write in my final examination?? Please help. Any help will be appreciated, Thank you.

• photosynthesis

• 1 $\begingroup$ Not an expert answer but from experience, I’ve seen limewater absorb $\ce{CO_2}$ mostly when $\ce{CO_2}$ is bubbled through the solution. However, in respirometers, KOH is used to absorb any neighboring $\ce{CO_2}$ in the chamber. Analogically thinking, it seems to me that limewater won’t be very efficient in removing $\ce{CO_2}$ from the flask in your experiment. Thus, you are correct in stating limewater absorbs $\ce{CO_2}$, but it might not work “satisfactorily,” making the teacher’s answer correct. For q2, I feel you should combine both your teacher’s and your answers. A control is needed. $\endgroup$ –  lightweaver Commented Apr 2, 2016 at 5:23

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• Revision notes >
• GCSE Biology Revision Notes >
• AQA GCSE Biology Revision Notes

Photosynthesis: Limiting Factors Affecting the Rate of Photosynthesis - (GCSE Biology)

Limiting factors affecting the rate of photosynthesis, single factors affecting rate of photosynthesis.

• Limiting factors affect the rate of a reaction. A limiting factor is a condition, that when in shortage, slows down the rate of a reaction. Light intensity, carbon dioxide concentration and temperature are limiting factors of photosynthesis. They all affect the rate of the photosynthetic reaction, but in different ways.

Photosynthesis Equation

The equation for photosynthesis is shown below in words. Factors affecting photosynthesis will affect the rate of reaction, will affect the amount of glucose and oxygen produced.

• Carbon dioxide + water + (light energy) → glucose + oxygen gas

Carbon Dioxide

• The higher the CO₂ concentration, the higher the rate until a certain point. As you increase the concentration of carbon dioxide (the reactant), the reaction is driven forwards.
• At high CO₂ concentration, rate levels off. until you reach a point at which the enzymes required are saturated. At this point, the carbon dioxide is no longer a limiting factor.

• You can prove the need for carbon dioxide in photosynthesis by placing a plant in a sealed bell jar with some soda lime. The soda lime will absorb any carbon dioxide present in the jar. After a while you can test the leaves of the plant for the presence of starch: 1. Dip the leaf in boiling water. This stops any further reactions in the leaf. 2. Place the leaf in a tube with ethanol. 3. Place the tube in an electric water bath and heat till it’s boiling. This removes any chlorophyll and turns the leaf a white colour. 4. Rinse the leaf with cold water and add drops of iodine solution. 5. Observe your results. The leaf won’t turn blue-black which means there is no starch present hence, no photosynthesis took place.

Light Intensity

• The higher the light intensity, the higher the rate until a certain point. If you increase the intensity of the light, more light can be trapped by the chloroplasts to provide energy to drive the photosynthetic reaction.
• At high light intensity, rate levels off. Again, this is up until a certain point, when the maximum amount of light has already been trapped. When you reach this point, light is no longer a limiting factor.

• You can prove the need for light by testing leaves of a plant left in the dark for 48 hours. You can test the leaves for starch in the same way as described above and will find that they won’t turn blue-black as no starch is present. The plant would have used up its starch stores without being able to make more due to the lack of light.

Temperature

• The higher the temperature, the higher the rate until a certain point. As the temperature increases, the enzymes gain more kinetic energy and so can catalyse the reaction at a greater rate. At the optimum temperature, the rate of reaction is the highest.
• At high temperatures, the rate falls as enzymes denature. When the temperature is greater than the optimum, the enzymes denature (change shape), so the reaction can no longer take place. This means the reaction is no longer feasible and so will not occur. This temperature is often around 45ºC.

Chlorophyll

• Chlorophyll can act as a limiting factor for photosynthesis. As the amount of chlorophyll increases, more light can be trapped by the chloroplasts in order to drive the reaction. However, once there is a certain amount of chlorophyll, it will no longer be a limiting factor and the values of other factors will limit the rate of the reaction. Infection and disease can reduce the amount of chlorophyll in a plant.
• You can prove the need for chlorophyll by testing variegated leaves. Variegated leaves are green and white – only the green areas contain chlorophyll so only they photosynthesise and produce starch. You can test the leaves for starch in the same way as described earlier. You will find that the areas of the leaf that were initially green (with chlorophyll) will turn black due to the presence of starch. The areas that were initially white (without chlorophyll) will remain orange.

Multiple Factors Affecting Rate of Photosynthesis

• Realistically, more than one factor is involved at a period in time. In the wild, plants are not kept under experimental conditions and so will be affected by different temperatures, carbon dioxide concentrations and light intensities.

Graphs can be used to map the effect of different limiting factors on the rate of photosynthesis at the same time:

Light Intensity and Temperature

• In the above graph, CO2 concentration is constant, so the effect of light intensity and temperature on photosynthesis is shown.
• As you increase light intensity, rate increases until we reach a plateau, after which light is no longer the limiting factor.
• As you increase temperature (shift from 20ºC to 30ºC), the rate increases. Also, the light intensity effect plateau’s at a higher level, so light intensity is a limiting factor for longer.

Light Intensity and CO2 concentration

• In the above graph, temperature is constant, so the effect of light intensity and carbon dioxide concentration on photosynthesis is shown.
• As you ice increase CO2 (shift from 0.05% to 0.5%), the rate increases. Also, the light intensity effect plateau’s at a higher level, so light intensity is a limiting factor for longer.

Canadian Pondweed Experiment

Setting up the experiment.

This test can specifically measure the effect of light intensity on photosynthesis. The rate of the plant’s oxygen production is proportional to the rate of photosynthesis .

• Place the plant in a boiling tube. Place the plant underwater in a boiling tube.
• Connect a gas syringe. Place the boiling tube on a clamp and add a capillary tube and a gas syringe. There should be an air bubble in the capillary tube.
• Use a ruler to measure movement of the air bubble. You can then use a ruler to see how much the air bubble moves. The amount the bubble moves is proportional to the rate of photosynthesis.
• Change the variables. You can now vary the light intensity, temperature and CO₂ concentration and observe the effect on bubble displacement (and hence rate of photosynthesis).

Investigating the Effect of Light Intensity

• Vary light intensity. Place a lamp at differing distances from the plant and observe the oxygen production through the movement of the bubble. This will give you an opportunity to see the effect of light intensity on the rate of photosynthesis.
• Control the other variables, and repeat for reliability. Control the temperature of the room and the time taken for the experiment. Make sure to repeat the experiment three times at each distance, then take a mean value in order to attain more reliable results.

Investigating the Effect of Temperature

• Vary the temperature. Place the boiling tube in different temperature water baths and observe the oxygen production through the movement of the bubble. This will give you an opportunity to see the effect of light intensity on the rate of photosynthesis
• Control the other variables, and repeat for reliability. Control the time taken for the experiment and the distance at which the lamp is placed. Make sure to repeat the experiment three times at each distance, then take a mean value in order to attain more reliable results.

Investigating the Effect of CO ₂ Concentration

• Vary the CO₂ concentration. Dissolve different amounts of sodium hydrogen-carbonate in the boiling tube. This releases carbon dioxide in water. You can dissolve different amounts into different boiling tubes to see the effect of differing carbon dioxide concentration on the rate of photosynthesis
• Control the other variables, and repeat for reliability. Control the temperature, distance of the lamp and the time taken for the experiment. Make sure to repeat the experiment three times at each carbon dioxide concentration, then take a mean value in order to attain more reliable results.

Photosynthesis takes place in plants to give them energy. Plants use sunlight, water and carbon dioxide to make food (energy in the form of sugars) and oxygen.

Photosynthesis takes place in chloroplasts, which contain chlorophyll. Chlorophyll is a pigment that gives plant leaves their green colour. This pigment absorbs the sun’s energy to use in the process of photosynthesis.

Chloroplast organelles are located inside the plant cell. Inside the chloroplast is a thylakoid membrane, housing a pigment called chlorophyll and this pigment absorbs the sun’s energy.

Photosynthesis is an endothermic reaction. This is because the sun’s energy is absorbed by the plant’s cells.

Plants absorb carbon dioxide and water from the air and soil. Inside the plant cell, the water is turned into oxygen. The carbon dioxide is transformed into energy (sugars in the form of glucose).

The rate of photosynthesis can be affected by various limiting factors such as light intensity, carbon dioxide concentration, temperature and water availability.

The rate of photosynthesis increases with an increase in light intensity until it reaches an optimal level. Beyond this level, an increase in light intensity will not increase the rate of photosynthesis.

The rate of photosynthesis increases with an increase in carbon dioxide concentration until it reaches an optimal level. Beyond this level, an increase in carbon dioxide concentration will not increase the rate of photosynthesis.

The rate of photosynthesis increases with an increase in temperature until it reaches an optimal level. Beyond this level, an increase in temperature will decrease the rate of photosynthesis.

Water is essential for photosynthesis, as it provides the hydrogen ions needed for the process. A shortage of water can limit the rate of photosynthesis.

Limiting factors are important in photosynthesis as they can determine the maximum rate of photosynthesis that can be achieved under a set of environmental conditions. Understanding these factors can help to improve plant growth and crop yield.

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Practical Biology

A collection of experiments that demonstrate biological concepts and processes.

Observing earthworm locomotion

Practical Work for Learning

Published experiments

How do plants and animals change the environment around them, class practical.

In this experiment hydrogencarbonate indicator shows the concentration of carbon dioxide in the environment of aquatic animals and plants. Over a 24 hour period, explore which combinations of plants and animals make a stable ecosystem in light or dark conditions.

Lesson organisation

This activity involves handling aquatic plants and animals collected in advance from a suitable pond. Students may therefore have some contact with pond water which can carry infectious diseases, so take hygiene precautions.

If necessary, reduce the apparatus required (and the time involved) by students working in larger groups, with each group making up one set of experimental tubes.

The experiment needs to run for 12 hours, so set up one day to read the results the next. Switch the lesson focus to interpretation and evaluation of method rather than practical skills development if you are short of time.

The indicator used to show levels of carbon dioxide is hydrogencarbonate indicator and the investigation uses soda lime which should only be handled by qualified staff.

Apparatus and Chemicals

For each group of students:.

Flat-bottomed specimen tubes (7 cm x 2.5 cm), 8 (or boiling tubes) ( Note 3 )

Bungs (to fit tubes), 8

Means of labelling tubes

Racks to support the tubes

Light source

Lightproof box or cupboard

Water plants, such as Elodea or filamentous alga, 8–10 lengths per working group ( Notes 4 and 5 )

Water animals such as water snails or Gammarus ( Notes 4 and 5 )

For the class – set up by technician/ teacher:

Hydrogencarbonate indicator solution (Recipe card 34) diluted 1 part in 10, equilibrated with atmospheric carbon dioxide, 100 cm 3 per group ( Note 1 )

Samples of hydrogencarbonate indicator showing the range of colours produced by atmospheres at different concentrations of carbon dioxide ( Note 2 )

Health & Safety and Technical notes

To collect Elodea or filamentous alga and pond animals you will have to work with pond water. Take hygiene precautions to minimise risk of infections from pond water ( Note 5 ).

Soda lime is corrosive. (Refer to CLEAPSS Hazcard 91.) It should be handled only by technician/ teacher wearing gloves and eye protection. Be aware of first aid response if any gets into the eyes.

Read our standard health & safety guidance

1 The indicator used to show levels of carbon dioxide is hydrogencarbonate indicator. Refer to CLEAPSS Recipe card 34 for details. The risks of indicators are unknown, so it should be made up by qualified staff using eye protection and in a fume cupboard. The resulting solution is Low hazard . You must use freshly prepared indicator solution.

2 Set up three tubes to show the range of colours produced by this indicator with carbon dioxide.

Tube 1: breathe into this through tubing or a straw before inserting the stopper. Exhaled air contains increased levels of carbon dioxide.

Tube 2: equilibrate this with normal air before stoppering. A fish tank bubbler might be useful. This will show normal levels of carbon dioxide.

Tube 3: place a muslin bag of soda lime granules in this tube, held in place with the stopper. You will need to do this some hours in advance. Soda lime absorbs carbon dioxide, so this shows the effect of no carbon dioxide in the air. Soda lime is Corrosive – refer to CLEAPSS Hazcard 91 – and should only be handled by qualified staff. Be aware of the first aid and medical response required should soda lime get into the eyes.

3 Wash all glassware very carefully and rinse with 0.01 M potassium hydrogencarbonate before use, as traces of acidic or alkaline chemicals will affect the indicator.

4 Rinse the plants and animals with 0.01M potassium hydrogencarbonate to prevent a colour change due to acids on their leaves or exoskeletons.

5 Pond water could contain disease-causing microbes, so take hygiene precautions and ensure students cover cuts or abrasions with waterproof plasters or gloves. Students will need to wash their hands with appropriate cleansers before leaving the teaching room. Refer to CLEAPSS Supplementary Risk Assessment SRA 09 09/06 for more details.

6 At some times of the year (when background light levels are lower) plants are less active. At all times of the year, the continuous light you provide needs to be intense enough to promote photosynthesis. Fluorescent striplights or halogen lamps are recommended, even in the summer months.

Ethical issues

Take care not to damage natural environments when harvesting from a pond.

Put used plant material and aquatic animals into fresh pond water for a few hours before returning to the pond, and return only the living material. This will limit the amount of indicator that is introduced to the pond.

SAFETY: Wear eye protection and gloves to make up indicator in a fume cupboard. (Refer to CLEAPSS Recipe card 34)

Wear eye protection and gloves to handle soda lime. (Refer to CLEAPSS Hazcard 91) Take hygiene precautions when handling pond organisms and pond water. (Refer to CLEAPSS Supplementary Risk Assessment SRA 09 09/06 for more details.)

Preparation

a Collect plant material such as Elodea or filamentous alga and animal material such as water snails or Gammarus from a pond.

b Rinse with 0.01 M potassium hydrogencarbonate ( Note 4 ).

c Cut the plant material into 5 cm lengths.

d Make up a fresh indicator solution ( Note 1 ).

e Wash the tubes thoroughly and rinse with 0.01 M potassium hydrogencarbonate ( Note 3 ).

f Set up three sample tubes to show range of indicator colours ( Note 2 ).

Investigation

a Take eight flat-bottomed specimen tubes in a rack (or boiling tubes).

b Label the tubes A, B, C, D, E, F, G and H.

c Fill each tube to about 2 cm from the top with hydrogencarbonate indicator solution. This will detect changes in carbon dioxide concentration. Try not to breathe directly over the open tubes.

d Add only aquatic animals to tubes A and E, and only aquatic plants to tubes B and F. Add some aquatic plants and animals together in tubes C and G. Put only indicator solution in tubes D and H. The table below shows these combinations.

Detailed Description of the Experiment (written for students)

Introduction, materials and methods, questions for further thought and discussion, references and links.

• Tools for Assessment of Student Learning Outcomes (written for faculty)

Every good gardener knows that the key to healthy plants is a fertile soil. Plants get water and nutrients from soil and it is the inherent characteristics of the soil in combination with environmental factors that determine soil fertility. Soils are complex and dynamic ecosystems with communities of organisms. Like all ecosystems they have a food web that may include bacteria, fungi, algae, protists, insects, worms, plant roots and burrowing animals. Soils also carry out essential ecosystem functions like water storage and filtration and, perhaps most importantly, decomposition.

Decomposition in soils is a key ecosystem function that in part determines the productivity and health of the plants growing there.  Decomposers feed on dead organic matter and in the process break it down into its simplest components: carbon dioxide, water and nutrients ( organic matter consists of material or molecules produced by living organisms). The process of decomposition releases large quantities of essential nutrients to the soil solution, thereby making them available to plant roots.  In northern hardwood forests, for example, about 85% of a tree’s nitrogen comes from decomposition (Bormann and Likens 1979).  Thus, if decomposition of a forest is impaired by drought, acid rain or some other stress, the vegetation may experience nutrient deficiencies.

Students looking for snails at the end of the experiment. Photograph by T.W. Stewart.

Decomposition is also important because it is part of the global carbon cycle. The carbon cycle is the cyclical movement of carbon atoms from the atmosphere to the biosphere/lithosphere and back to the atmosphere (Figure 1). In the atmosphere, carbon is in the form of carbon dioxide gas. Through the process of photosynthesis, some of that carbon is converted into organic carbon which makes up organic matter or biomass. Plants and animals perform cellular respiration and convert a small percentage of that organic carbon back to CO 2 .

A larger portion of that organic carbon in plants is transferred to the soil  when plants shed their leaves or when they die. Decomposers then begin their work of breaking down the organic matter. Some of the organic carbon in the organic matter is converted into CO 2 which is released into the soil pore spaces leading to relatively high concentrations of CO 2 compared to the atmosphere. This difference in concentration causes CO 2 to diffuse from the soil to the atmosphere. This movement or flux of CO 2 is known as CO 2 emission (Figure 1).

Decomposition is not the only source of CO 2 in soil. In a forest or grassland ecosystem, plant roots are abundant in the soil and root cells perform cellular respiration , metabolizing carbohydrates that are sent down from the leaves. This CO 2 is released to the soil and can be responsible for anywhere between 0 and 60% of a soil’s CO 2 emission. Note that CO 2 emission is the movement of CO 2 from soil to the atmosphere, whereas decomposition and root respiration are processes that produce CO 2 in the soil (Figure 2).

Release of CO 2 from soils has global implications because it occurs in ecosystems worldwide and its magnitude is such that it contributes significantly to the greenhouse effect . The greenhouse effect is a natural property of our atmosphere in which greenhouse gases prevent the transfer of heat from the earth’s surface to outer space, thereby warming the atmosphere. Since the industrial revolution human activity (e.g., fossil fuel combustion and deforestation) has led to global increases in the concentrations of greenhouse gases (such as CO 2 ) in our atmosphere. This rapid increase will likely lead to a cascade of environmental impacts such as global warming, sea level rise, alteration of precipitation patterns, and increased storm severity (IPCC 2007).

A great deal of research money and effort has been invested in studies of soil CO 2 emission in recent years because of the potential impacts of this process on the greenhouse effect (Schlesinger and Andrews 2000). The amount of organic carbon stored in soils worldwide is about 1600 gigatons (Gt) compared to 750 Gt in the atmosphere mostly in the form of CO 2 (Rustad et al. 2000).  Thus, if soil respiration increased slightly so that just 10% of the soil carbon pool was converted to CO 2 , atmospheric CO 2 concentrations in the atmosphere could increase by one-fifth!

Several environmental factors control the rates of decomposition and root respiration and therefore, the rate of CO 2 emission from soils. Since decomposition is an enzyme-mediated biological process carried out by bacteria and fungi, it is very sensitive to temperature. In most soils, the decomposition rate peaks at about 25°C and declines as temperature varies from this maximum.  Soil moisture also affects the activity of microorganisms.  Very dry or very wet (flooded) conditions tend to reduce decomposition rates (Hanson et al. 1993). A history of acid deposition can also lower the pH of soils thereby inhibiting decomposers.

Respiration rates will also depend on how fast CO 2 molecules can diffuse to the soil surface. Diffusion will be affected by soil moisture (how much of the pore space is filled with water) and soil texture (the size distribution of soil particles). Thus, it is likely that soil temperature, moisture, pH, density and texture will all influence soil respiration rates. In this exercise, you will investigate the effects of these (and perhaps other) environmental factors on CO 2 emission (Figure 2).

One of the most common methods for measuring soil respiration, the soda-lime method , is remarkably easy and does not require expensive equipment. As a result scientists all over the world have employed it (Grogan 1998). Soda lime is a variable mixture of sodium hydroxide (NaOH) and calcium hydroxide (Ca(OH) 2 ) in granular form. It’s commonly used in laboratories as a desiccant because it readily absorbs water vapor from the air.  Because of its alkaline properties soda lime also removes carbon dioxide very efficiently from the atmosphere according to these reactions: 

            2NaOH (s) +  CO 2 (g)        «        Na 2 CO 3 (s)  +   H 2 O (ads)    [1]

            Ca(OH) 2 (s) +  CO 2 (g)       «        CaCO 3 (s)  +   H 2 O (ads)     [2]

Note that for every molecule of CO 2 adsorbed, a molecule of water is created. These water molecules remain temporarily adsorbed (ads) to the soda lime but can be evaporated off at boiling temperatures.

Figure 3. Schematic diagram of the soil respiration chambers. CO 2 diffuses from the soil into the chamber air space. It is then absorbed by the soda lime.

The soda lime method involves placing a pre-weighed, open dish of soda lime on the ground and covering it with a chamber of known diameter (Figure 3).  As the soil CO 2 diffuses into the chamber it is quickly absorbed by the soda lime (along with water vapor).  After 24 hours, the chamber is removed and the soda lime is dried at 105°C to evaporate the water and then weighed.  The increase in mass of the soda lime is attributable to CO 2 (Edwards 1982, as modified by Grogan 1998).

Study Site(s):    With your Instructor, choose appropriate study sites that are relatively flat and are not extremely stony. You need to be able to place an 18 cm (7.1 in) diameter chamber on the ground where there are no living plants and no large stones. Depending on your experimental question you may want two contrasting sites like conifer site vs. hardwood site, north slope vs. south slope, or dry vs. wet.

Overview of Data Collection and Analysis Methods:   1 to 2 Days Before Lab Session 1:

• Label glass jars (40- to 100-mL glass jar with screw top) with a piece of tape and permanent marker. Add approximately 8 grams of soda lime to each jar. Place the jars with soda lime in an oven at 105°C for at least 24 hours to evaporate the water from the granules. You will need 8 - 10 jars per group plus one extra that the whole class can use for the blank.

Lab Session 1:

• Remove jars from the oven (use gloves or tongs!) and place in a desiccator to cool for 2-5 minutes. Remove jars from desiccator one at a time, weigh to the nearest milligram (0.001 g) or tenth-milligram (0.0001g) and cover immediately. Record the mass as the initial mass in Table 1 .
• Take the jars, chambers, thermometers and sampling equipment and go out to your field site. Take a few minutes to note the variations in microclimate and microtopography within the forest.
• In small groups design your experiment. You will be comparing the rate of soil CO 2 emission of two sites with different microclimates and/or soil characteristics. As a group, decide on the sites or the microclimates you would like to compare. Here are some suggestions but you are encouraged to think of your own: Conifer site vs. hardwood site Sun vs. shade Ridgetop vs. valley bottom With leaf layer vs. without leaf layer (i.e., the layer of dead leaves on the soil surface is removed)
• As a group write out your Experimental Design according to the handout, Experimental Design Requirements .  Show it to your Instructor for approval before proceeding. As homework type up your answers to the questions on the handout.
• Place a chamber upside down on a relatively flat area of the soil. The entire rim of the chamber must be inserted at least 1 cm into the soil so as to minimize gas exchange with the atmosphere. So, carefully remove twigs and small rocks that lie under the rim without disturbing the leaves and soil surface under the chamber. Remove any green plants by pinching or cutting them at soil level. It is essential that the soil be disturbed as little as possible!
• Slowly and carefully push down while rotating the chamber back and forth to force the edges about 1 – 2  cm into the soil surface. If there are subsurface roots or rocks in the way, you may need to move to another location. The key here is to get a good seal all along the edge of the chamber so there are no gaps.
• Obtain a jar containing soda lime. Remove the cap and place the jar under the chamber so that it rests on the soil surface. Make sure it is not likely to tip over.
• Replace the chamber and place a weight on it (like a fist-sized rock or a thick branch) to maintain pressure and keep it from blowing away or tipping over.
• Record the number of the soda lime jar and the number and location of the chamber. Repeat these steps for each of the chambers at each site.
• At one of the sites used by the class, place an opened jar of soda lime in an upright chamber and seal the chamber with a lid. This will serve as a blank to document the amount of CO 2 absorbed from the air in the chamber and during the opening and closing of the jars. Only one blank is needed for all of the groups.
• Let all chambers incubate for 24 ( + 4) hours. If the ambient daytime air temperature is below 16ºC, then incubate the chambers for 48 ( + 4) hours.
• Before leaving the site quantify the differences in environmental factors between your two sampling sites. You may measure any or all of the following. Your Instructor may have additional parameters for you to measure. Click here for instructions on measuring these variables. Soil temperature Soil moisture Soil pH

1 or 2 Days After Lab Session 1:

• Return to the field site after the designated time has elapsed. Remove the chambers and cap the soda lime jars. Return all materials to the lab. Uncover the soda lime jars and place them in the drying oven at 105ºC.

Lab Session 2:

• Remove the dry soda lime from the oven and place in a desiccator to cool for 5 minutes. Remove jars one at a time from the desiccator, weigh to the nearest milligram (0.001 g) or tenth-milligram (0.0001 g). Record this as the final mass (which includes the mass of the jar) in Table 1 .
• Calculate the mg of soil CO 2 absorbed by the soda lime in each chamber: Change in Mass of Blank (g) = M b = (Final Mass of Blank – Initial Mass of Blank) Soil CO 2 Absorbed (g) = Final Mass – Initial Mass – M b
• Calculate the CO 2 Emission Rate (E) for each chamber: Ac = Area of ground covered by chamber (m 2 ) E (g CO 2 m -2 d -1 ) = (Soil CO 2 Absorbed * 1.69) / A c / Days of Incubation [ The 1.69 in the equation above is used to correct for the water molecule that is lost when a molecule of CO 2 is adsorbed. ] Click here for a data sheet in EXCEL format.
• Perform a Student’s t-Test on the CO 2 Emission Rates to test for significant differences between the two experimental treatments. Click here for a step-by-step procedure .
• With help from the Instructor summarize your environmental variables and create a table in the proper format to present these data. Homework: Write a lab report using the proper format. Click here for report guidelines . Your Instructor will assign a due date for the first draft of the report and for the final draft of the report.
• How did your two sampling sites differ in terms of temperature, moisture, pH or other characteristics? Could these differences explain the differences you observed in CO 2 emission rate?
• The soil under your chambers probably contained plant roots. How might these plant roots have affected your CO 2 emission rates? Explain. Design an experiment using these chambers that would allow you to determine what proportion of the CO 2 emitted came from roots and what proportion came from decomposition.  
• Explain how decomposition in soils is linked to the greenhouse effect.
• If just 5% of the world’s soil organic carbon pool was decomposed , how many tons of carbon would be released?
• Calculate the average CO 2 emission rate  and standard deviations for each sample location (or perform a statistical test). Put these values in a table. Then write two to three paragraphs describing and interpreting the results of your experiment.
• The temperature and moisture data you collected represent point-in-time measurements. Do you think the temperatures and soil moisture values are representative of the microclimate during the entire incubation period? What would be a more accurate way to quantify the microclimate during the incubation period?
• Are there other environmental or site factors that you did not measure that could explain the differing rates of CO 2 emission between your sampling locations? Explain how they would affect the emission rate.
• CO 2 Emission varies with geographic location and with season. Conduct a literature search for soil CO 2 emission values from around the world. Try to find some from your area. Some key words that will aid you in your search are: soil respiration, soil CO 2 , soil carbon, carbon emissions, CO 2 emissions, soda lime, carbon cycle.  What range of values can you find? Where are the values the highest? Where are they the lowest? How does your area compare? [ Note: make sure when you compare values from different studies that you convert all the values to the same units! ]
• Because decomposition is a temperature-dependent process, it is expected to be affected by global warming. Write down one or two predictions about how decomposition in soil will change and how those changes will affect plants. Then conduct a literature search to find out what the experts are predicting. Were your predictions correct? If not, why not? What other predictions have the experts made? Some search phrases that will aid you in your search are: soil CO 2 , CO 2 emissions, soil respiration, global warming soil carbon, tundra soils, global warming positive feedback, soil respiration temperature, decomposition temperature.
• Bormann, F.H. and G.E. Likens. 1979. Pattern and Process in a Forested Ecosystem . Springer-Verlag, New York, NY.
• Buchman, N. 2000. Biotic and abiotic factors controlling soil respiration rates in Picea abies stands. Soil Biology and Biochemistry 32 :1625-1635.
• Edwards, N.T. 1982. The use of soda-lime for measuring respiration rates in terrestrial systems. Pedobiologia 23 :321-330.
• Grogan, P. 1998. CO 2 flux measurement using soda lime: correction for water formed during CO 2 adsorption. Ecology 79 :1467-1468.
• Hanson, P.J., S.D. Wullschleger, S.A. Bohlman, and D.E. Todd. 1993. Seasonal and topographic patterns of forest floor CO 2 efflux from an upland oak forest. Tree Physiology 13 :1-15.
• Hogberg, P., A. Nordgren, N. Buchmann, A.F. Taylor, A. Ekblad, M.N. Hogberg, G. Nyberg, M. Ottosson-Lofvenius, D.J. Read. 2001. Large scale forest girdling shows that current photosynthesis drives soil respiration. Nature 411 :789-92.
• IPCC, 2007: Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change [Solomon, S., D. Qin, M. Manning, Z. Chen, M. Marquis, K.B. Averyt, M. Tignor and H.L. Miller (eds.)]. Cambridge University Press, Cambridge, United Kingdom and New York, NY, USA, 996 pp.
• Raich, J.W., and W.H. Schlesinger. 1992. The global carbon dioxide flux in soil respiration and its relationship to vegetation and climate. Tellus 44B :81-99.
• Rustad, L.E., T.G. Huntington, R.D. Boone. 2000. Controls on soil respiration: Implications for climate change. Biogeochemistry 48 :1-6.
• Schlesinger, W.H., and J.A. Andrews. 2000. Soil respiration and the global carbon cycle. Biogeochemistry . 48 :7-20.
• Simmons, J.A., I.J. Fernandez, R.D. Briggs and M.D. Delaney. 1996. Forest floor carbon pools and fluxes along a regional climate gradient in Maine, USA. Forest Ecology and Management 84 :81-95.
• Toland, D.E., and D.R. Zak. 1994. Seasonal patterns of soil respiration in intact and clear-cut northern hardwood forests. Canadian Journal of Forest Research 24 :1711-1716.

Useful Web Sites

• Caprette, D.R. 2005. Student’s t-Test (for independent samples). ( http://www.ruf.rice.edu/~bioslabs/tools/stats/ttest.html , accessed 4 December 2007)
• Davidson, E. 2006. Soil respiration, temperature and moisture. ( http://harvardforest.fas.harvard.edu/data/p00/hf006/hf006.html , accessed 12 December 2007)
• Dolphin, W.R. 1997. Writing Lab Reports and Scientific Papers. ( http://www.mhhe.com/biosci/genbio/maderinquiry/writing.html , accessed 13 December 2007)
• Environmental Literacy Council. 2006. Soil and the Carbon Cycle. ( http://www.enviroliteracy.org/article.php/700.html , accessed 4 December 2007)
• Harmon, M. 2003. LIDET: Long-term intersite decomposition experiment team. ( http://www.fsl.orst.edu/lter/research/intersite/lidet.htm , accessed 13 December 2007)
• Rice, C.W. 2005 (posting date). What is the carbon cycle? ( http://soilcarboncenter.k-state.edu/carbcycle.html , accessed 4 December 2007)
• Roche, J. 2007. How to write a lab report. ( http://inpp.ohiou.edu/~roche/371_fall07/how_to_write_a_lab_report.pdf , accessed 10 December 2007)
• Trochim, W.M.K. 2006. The t-test. ( http://www.socialresearchmethods.net/kb/stat_t.php , accessed 13 December 2007) United States Geological Survey. 2006. Assessing carbon stocks in soil. ( http://edcintl.cr.usgs.gov/carbon_cycle/carbonstocks.html , accessed 11 December 2007)
• Woods Hole Research Center. 2005. Soil Respiration. ( http://www.whrc.org/new_england/Howland_Forest/soil_respiration.htm , accessed 5 December 2007)

Tools for Assessment of Student Learning Outcomes

Assessment You will be assessed on two aspects of this project -  the experimental design and the written lab report. The experimental design will be used to assess your ability to use the scientific method appropriately to answer a question. The lab report will be used to test your comprehension of the principles behind soil respiration and your ability to communicate in writing in proper scientific format.

Experimental Design Guidelines Download Experiment Design Guidelines document (DOC)

Lab Report Guidelines See Lab Report Guidelines document (DOC)

Rubrics Download Experimental Design Rubric (DOC) Download Lab Report Prime Trait Assessment ( EXCEL file (XLS) and WORD file (DOC))

Sample Exam Questions Q. The process that converts atmospheric CO 2 into organic C in plants is_________.

A.   Photosynthesis

Q.   If global warming were to lead to warmer soil temperatures and therefore faster decomposition worldwide, what would you expect to happen to the levels of CO 2 in the atmosphere (all else being equal)? Explain.

A.   Faster decomposition would lead to greater CO 2 emission rates which would lead to an increase in atmospheric CO 2 concentration.

Q.   Acid deposition tends to inhibit soil microbial populations and lead to slower decomposition. What effect, if any, will this have on the vegetation? Explain.

A.   Plants obtain most of their nutrients from the decomposition process. If decomposition is slowed, plants may become nutrient deficient or their growth will be slowed .

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The carbon dioxide problem

By Derek Cheung 2006-03-01T00:00:00+00:00

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Measuring carbon dioxide from plant debris provides an opportunity for an inquiry-based experiment aimed at 14-15 year olds. Similar experiments are done by soil scientists and ecologists in their efforts to understand the global carbon cycle

Source: Petmal/iStock

Plant debris could help students to understand the global carbon cycle

Chemistry educators worldwide advocate inquiry-based practical work in the secondary school chemistry curriculum. However, good contextualised scientific inquiries for chemistry students are scarce.1 Here I describe an inquiry-based experiment done by Secondary 4 chemistry students (aged 14-15 years) in Hong Kong. This investigation can provide students with an excellent opportunity to understand part of the global carbon cycle.2,3

The carbon dioxide problem 

The purpose of this investigation is not for students to arrive at a single, 'correct' answer but to develop inquiry skills. Box 1 illustrates the problem. Students are seldom aware that bacteria and fungi feed on the dead remains of plants and, via respiration, release CO 2 to the atmosphere. Leaf litter decomposition is a key process in the global carbon cycle. Each year, the mass of CO 2 released by decomposing leaves is approximately equal to that released by animal and plant respiration. 

In a trial of the experiment, working with teachers, we divided a class into groups of three-four students based on their abilities. After about 30 minutes, we collected their plans for evaluation. The next day, we invited the students to present their plans to the rest of the class. This is done to turn assessment activities into a teaching-learning activity. Twenty minutes were allocated for each presentation. Each group needed to answer questions raised by other students. After the presentations, students did experiments according to their plans. The problem has been designed to be an open-ended investigation so students should be allowed to try out their own plans. 

Common procedural errors and limitations

We found that our students had the following common procedural errors or limitations in their experimental designs: 

• some students did not recognise that CO 2 is present in the air even without the decomposing leaves. They did not realise, therefore, that they had to determine how much more CO 2 is produced as a result of the leaf decomposition, and a control set-up was needed;
• many students wanted to use lime water to absorb the CO 2 released by dead leaves and determine the mass of CO 2 by weighing a bottle of lime water before and after a measurement period. This method does not work because lime water does not absorb CO 2 efficiently;
• some students used sodium hydroxide solution to absorb the CO 2 released by dead leaves and then did a back titration with hydrochloric acid to calculate any excess sodium hydroxide. But sodium hydroxide reacts with CO 2 to form sodium carbonate. Therefore, the resulting solution would be a mixture of carbonate ions and hydroxide ions. During the back titration, hydrochloric acid would react with both carbonate ions and hydroxide ions. (The composition of the hydroxide-carbonate mixture could be found by the double indicators method. 4 Alternatively, barium chloride solution may be added to the mixture to precipitate the carbonate ions before the back titration is done. 5 Unfortunately, these two methods are beyond Secondary 4 students.)  

Sample procedure

Here we present a sample procedure for teacher information. It should not be distributed to students as a cook-book style experiment.

Materials (per group of three-four students):

• Safety goggles; 
• Two airtight containers (at least 28cm × 21cm × 6cm); 
• 10 dry dead leaves (same type of leaf);
• Soda lime (1.0-2.5mm granular size,  ca 40g, hazard warning label = corrosive); 
• Deionised or distilled water; 
• Two glass Petri dishes (at least 9cm diameter); 
• Measuring cylinder (10cm 3 );
• Spatula; 
• Access to balance (±0.001g); 
• Oven, and oven gloves. 

Pre-lab (by technician)

• Put a thin layer of granulated soda lime in the bottom plate of a glass Petri dish. Find out the mass of soda lime used. This is approximately half of the mass of soda lime to be used by one group of students. Estimate the total amount of soda lime needed by the whole class. 
• To inactivate soda lime, put the estimated total amount of soda lime in a beaker and dry it in an oven at 105°C for 24 hours. Place the soda lime in a desiccator immediately upon removal from the oven. 

Experimental details for students

(Put on your safety goggles and wear lab coat and gloves.) Obtain two airtight containers. Label one container as 'experimental' and the other as 'control'. Obtain eight to 10 leaves (use the same type of leaf and if the leaves are large in size, reduce the number of leaf). Record the total mass of leaves. Put the leaves into the experimental container. 

Fill the top plates of the two glass Petri dishes half-full with distilled water. Place one plate of water into the experimental container and the other plate into the control. Using a marker, label the bottom plates of the two Petri dishes as 'experimental' and 'control'. Weigh each plate separately and record its exact mass. Using a spatula, spread out a thin layer of the oven-dried soda lime granules in the plates. ( CAUTION : Soda lime is corrosive. Do not touch it with your bare hands.) Weigh each plate separately and record the exact mass of soda lime used. Place the plates containing soda lime into the experimental and control containers. Using a measuring cylinder, carefully add 5.0cm 3 distilled water to the soda lime in the experimental and control containers to activate the soda lime. Immediately seal the containers and place them in the location specified by your teacher for one week.

After one week, open the experimental and control containers and remove the two soda lime dishes. Dry the soda lime at 105°C in an oven for 24 hours. Find out the mass of CO 2 absorbed by re-weighing the two soda lime dishes. 

Notes for teachers

Risk assessments should be done in advance by the teacher. Teachers should also note that: 

• solid sodium hydroxide alone is not satisfactory to absorb CO 2 in this experiment because it absorbs water vapour from the air and puddles of concentrated (and corrosive) sodium hydroxide solution will be formed. A lot of heat will also be released when sodium hydroxide solution is formed;
• absorption of CO 2 using solid calcium hydroxide alone is not efficient because CO 2 must be dissolved in water before it can react with calcium hydroxide. Soda lime is more efficient because sodium hydroxide is added to calcium hydroxide to absorb moisture. The moisture in soda lime granules is not visible when the water content is less than 20 per cent. Because CO 2 is chemically bound but the moisture is not, soda lime can be dried and weighed before and after a measurement period to determine the amount of CO 2 absorbed. Since soda lime is inefficient for CO 2 absorption unless moisture is available, this sample procedure allows water to evaporate from a dish to increase humidity. You should remind your students that soda lime is corrosive and so they should not touch it with their bare hands. I recommend the use of granular (1.0-2.5mm) soda lime. You may use other granular sizes, but it should be small enough for a large surface-to-volume ratio and large enough to prevent losses of fine particles during drying and handling.  

Sample calculations and results 

Although soda lime is a variable mixture of sodium hydroxide and calcium hydroxide, you do not need to know the exact percentages of these chemicals to calculate the mass of CO 2  absorbed. You may guide your students to apply the following chemical concepts to solve this problem.  

Absorption of CO 2 occurs by: 

Ca(OH) 2 + CO 2   → CaCO 3 + H 2 O  ( i ) 

2NaOH + CO 2   → Na 2 CO 3 + H 2 O  ( ii ) 

For every mole of CO 2 that is reacted with Ca(OH) 2 in the soda lime, one mole of H 2 O is formed that is subsequently evaporated during oven-drying. Thus, the increase in mass of dried soda lime measured before and after the experiment is not equal to the mass of CO 2 absorbed.  

From equation ( i ), if one mole (44g) of CO 2  has been absorbed, the increase in mass = molar mass of CaCO 3- molar mass of Ca(OH) 2 = 100 - 74 = 26g. The mass of CO 2 absorbed by Ca(OH) 2 is proportional to the increase in mass after the experiment. Thus, 

Mass of CO 2 absorbed/44g = Increase in mass/26g

Mass of CO 2 absorbed = increase in mass × 44g/26g 

= increase in mass × 1.69

Similarly, from equation ( ii ), if one mole (44g) of CO 2 has been absorbed by two moles of NaOH in the soda lime, the increase in mass 

= molar mass of Na 2 CO 3

- the mass of two moles of NaOH 

= 106 - (2 × 40)

= 26g.  

Like Ca(OH) 2 , the mass of CO 2 absorbed by NaOH = increase in mass × 1.69. Thus, the relative amounts of Ca(OH) 2 and NaOH in a sample of soda lime is not important in this experiment. Some chemical suppliers also add potassium hydroxide to soda lime, but the mass of CO 2 absorbed by potassium hydroxide can also be obtained by multiplying the increase in mass by 1.69.  

To find the amount of CO 2 released by decomposing leaves, students need to subtract the mass of CO 2 absorbed in the control container from the mass of CO 2 absorbed in the experimental container. 

Mass of dry dead leaves used = 3.073g  Name of plant =  Liquidambar formosana Duration = seven days   See Table 1 for the results to this example. Therefore, the mass of CO 2 released per gram of leaf per day = (0.059 - 0.008)g/3.073g × 7 days 

And finally

The mass of CO 2 released depends upon factors such as the type of plant, the amount of bacteria and fungi, and the temperature. With  Bauhinia purpurea leaves, for example, I found that 0.012g of CO 2 was released per gram of leaf per day. In countries with cooler climates than Hong Kong, a treatment period longer than seven days may be needed. Students could do the following additional activities as extensions: 

• compare the amounts of carbon dioxide produced by different types of plant; 
• write a paper on the role of plants in global climate change; 
• debate whether society should act now to halt global warming. 

Overall we have found that such practical activities set in a real-life context are of prime importance to chemistry students because they help to structure the learning process, and give purpose to learning chemistry.

Dr Derek Cheung is associate professor in the department of curriculum and instruction at The Chinese University of Hong Kong, Shatin, Hong Kong (e-mail: [email protected]).

Acknowledgements

I would like to thank the Quality Education Fund for financial support of this project (grant code 2003/0750). 

Box 1 

Plant debris decomposes in the soil, releasing carbon dioxide. This is because bacteria and fungi break down dead leaves by the following chemical reaction: 

C x (H 2 O) y +  x O 2 →  x CO 2 +  y H 2 O + energy 

This reaction is one of the components of the global carbon cycle. Different types of plant material decompose at different rates. Green, leafy matter decomposes easily, but woody, stem debris takes longer. Low temperatures, dry conditions and flooding will also slow down decomposition. Unfortunately, human activities - especially burning of fossil fuels - have led to changes in the natural carbon cycle. The cycle is now out of balance. The amount of carbon dioxide in the atmosphere has increased from 280ppm to 370ppm over the past 140 years, resulting in global warming via the greenhouse effect. The slight warming of the Earth's surface may cause bacteria and fungi to decompose dead leaves more rapidly, releasing even more carbon dioxide into the atmosphere. This investigation will provide you with an opportunity to understand part of the carbon cycle dynamics. 

Suppose that you work as a chemist in an environmental protection company. Your challenge is to plan and do an investigation to determine the number of grams of carbon dioxide released by one gram of dead leaves per day. Submit the plan as group work by __(enter date)__. Your group will be presenting on __(enter date)___ in front of the company representatives. You will have 10 minutes to present your plan, followed by 10 minutes in which you will be expected to respond to queries.  

Your presentation needs to answer the following questions: 

• how will you measure the rate of formation of CO 2 for a particular type of leaf ( eg , oak)? 
• what variables will you need to keep constant in this investigation?
• will the proposed procedure be feasible and safe?

After reviewing your experimental design, the teacher will discuss any safety precautions that are specific to your design. Obtain teacher approval before beginning any lab work.  

Submit a lab report in writing by __(enter date)__. Make sure you include the following information in your report: 

• the purpose of your investigation; 
• the actual method used for investigation; 
• data and results; 
• conclusion.
• V. L. Lechtanski,  Inquiry-based experiments in chemistry . Washington, DC: ACS, 2000. 
• P. Buell and J. Girard,  Chemistry fundamentals: an environmental perspective . Sudbury, MA: Jones and Bartlett, 2003. 
• J. W. Raich,  Forest Ecol. and Manage ., 1998,  107 , 309. 
• D. A. Skoog, D. M. West, F. J. Holler and S. R. Crouch,  Analytical chemistry: an introduction . Fort Worth: Saunders College, 2000. 
• D. J. Wink, S. F. Gislason and J. E. Kuehn,  Working with chemistry: a laboratory inquiry program . New York: W. H. Freeman, 2005. 
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Investigation and Possibilities of Reuse of Carbon Dioxide Absorbent Used in Anesthesiology

Absorbents used in closed and semi-closed circuit environments play a key role in preventing carbon dioxide poisoning. Here we present an analysis of one of the most common carbon dioxide absorbents—soda lime. In the first step, we analyzed the composition of fresh and used samples. For this purpose, volumetric and photometric analyses were introduced. Thermal properties and decomposition patterns were also studied using thermogravimetric and X-ray powder diffraction (PXRD) analyses. We also investigated the kinetics of carbon dioxide absorption under conditions imitating a closed-circuit environment.

1. Introduction

Soda lime is one of the most popular carbon dioxide absorbents used in order to maintain a safe level of this gas. Its composition has slightly changed over time; however, calcium hydroxide is still the main component. Often an indicator signalizing its consumption is added, as well as small amounts of sodium (or potassium) hydroxide since NaOH is more reactive than Ca(OH) 2 . Moreover, the hygroscopic properties of NaOH reduce interphase mass transfer barriers and speed up the CO 2 sorption process. Soda lime is most commonly used in environments characterized by reduced or no connection with fresh air, like anesthetic and diving apparatus or spacecraft. Such environments are commonly called “closed” and “semi-closed” circuit environments. It is an important issue since humans are there to the most degree, exposed to increased levels of CO 2 . Some of the most important examples of such systems, where soda lime and similar absorbents are used are anesthetic and diving apparatus, submarines, spacecraft and mine refuge chambers [ 1 , 2 ]. Breathing in such environments results in increased carbon dioxide concentrations. This may cause occurrence of several symptoms or even lead to death. In order to prevent it, absorbents based on alkali hydroxides are used to capture carbon dioxide, with soda lime being the most common one. Despite relatively low costs of production and simple operating principle, soda lime also has numerous flaws and limitations. It may undergo reaction with some of the gases used in general anesthesia, especially with sevoflurane, desflurane, isoflurane and enflurane to form a number of degradation products [ 3 , 4 , 5 ]. One of the gases, fluoromethyl 2,2-difluoro-1-(trifluoromethyl)vinyl ether, often abbreviated in the literature as Compound A , has been proven nephrotoxic to rats [ 6 , 7 , 8 ]. It may also contain viruses derived from the exhaled air, including rhinovirus, respiratory syncytial virus, parainfluenza virus, adenovirus, coronavirus, human metapneumovirus and influenza virus [ 9 , 10 , 11 , 12 ]. Under certain conditions, soda lime may also support the formation of carbon monoxide, one of the most toxic gases. Some of the factors that increase the possibility of carbon monoxide formation are the use of volatile anesthetic agents in question, their concentration and flow rate, dryness, type of used absorbent and the temperature in which absorption takes place [ 13 , 14 , 15 , 16 ]. Another complication is the need to control the absorbent exhaustion level since an excess of CO 2 can cause hypercapnia. In order to signalize soda lime consumption, indicators like ethyl violet or ethyl orange are added. It does not, however, provide the necessary level of safety since sometimes these indicators may return to their previous color despite absorbent exhaustion [ 17 ]. Thus, capnometers are used in order to control the level of carbon dioxide in a patient’s organism. These devices are, however, relatively expensive and are not used in less developed areas. Soda lime with additions of indicators has also been withdrawn from U.S. Navy Fleet since it was suspected of releasing harmful compounds [ 18 ]. During the absorption process, extensive amounts of heat are produced, especially when baralyme (modification of soda lime in which calcium hydroxide is replaced with barium hydroxide) is used. Indicator color change monitoring and extensive heat emission additionally complicate soda lime use in diving apparatus. Furthermore, soda lime dust inhalation was observed, which may contribute to the occurrence of airway diseases in divers [ 19 ]. There are also no reports of reliable recycling methods of exhausted soda lime, which is most commonly considered medical waste. Taking into consideration all the mentioned problems and limitations, soda lime requires rigorous and careful handling. It is also the basis for seeking new alternative absorbents that would be more reliable and versatile. Some of them are carbonaceous materials, solid and liquid organic amines, mixtures of metals peroxides, hyperoxides or superoxides with water, membranes and zeolites [ 20 ]. There is no doubt that carbon dioxide plays one of the most significant roles taking into consideration both biological and environmental issues.

In this paper, we present a critical evaluation of soda lime performance as carbon dioxide absorbent. We have also investigated composition (volumetric analysis, photometric analysis) and thermal properties (thermogravimetric analysis, X-ray powder diffraction analysis) of two soda lime commercial samples, as well as proposed its possible recycling method.

2. Experimental

2.1. materials and analysis.

All chemicals used during analysis were purchased from Avantor Performance Materials, Gliwice, Poland S.A.: pure calcium hydroxide, pure calcium carbonate, hydrochloric acid (1 mol·L −1 ), sodium hydroxide (4 mol·L −1 ), EDTA (0.05 mol·L −1 ), phenolphthalein, methyl orange, Patton and Reeder’s indicator. Analyzed soda lime samples came from the company producing carbon dioxide absorbents. The first sample was a fresh, unused sample, while the second one was used and considered exhausted prior to the research. The samples were marked as follows:

• SL (F)—fresh soda lime sample;
• SL (U)—used soda lime sample.

Chemical composition and thermal decomposition curves of samples were investigated using volumetric, photometric, thermogravimetric analysis. In order to better understand thermal decomposition pathways and to exclude the theoretical presence of other products of decomposition, a PXRD analysis was performed for sinters of samples SL (F) and SL (U) prepared at 950 °C. The sinters were obtained by heating the samples to the temperature defined from the thermal curves.

In the second part, we conducted an experiment under conditions imitating carbon dioxide absorption in closed circuit anesthetic apparatus, which allowed us to draw conclusions about the kinetics of carbon dioxide absorption and soda lime performance as a carbon dioxide absorbent. Chemical composition and thermal destruction ways after absorption were investigated in the same way as in the first part of our study. These samples were marked in the following way:

• SL (5 min)—fresh soda lime sample after 5 min of carbon dioxide absorption;
• SL (15 min)—fresh soda lime sample after 15 min of carbon dioxide absorption.
• SL (30 min)—fresh soda lime sample after 30 min of carbon dioxide absorption.

The experimental setup is schematically shown in Figure 1 . A compressor was used to flow atmospheric air through a water bubbler and then a packed bed of sorbent. The carbon dioxide concentration in the inlet air (C o ) was around 4% (average concentration of carbon dioxide in exhaled air). The experiment was conducted at room temperature (23–25 °C), and relative humidity was maintained at around 55% during the experiment.

Experimental setup for carbon dioxide absorption under conditions imitating closed-circuit environment.

2.2. Methods and Instruments

A bigger batch of each sample was ground in a mortar. After homogenization, around 1.0 grams of each sample was stirred in 1 L of distilled water on a magnetic stirrer for 24 h. Suspensions prepared this way were then investigated using volumetric analysis. In order to ensure repeatability of results, no less than three portions per each sample were collected and titrated. Powders resulting from grinding in a mortar were investigated using thermogravimetric analysis. We have also performed thermal decompositions of two main soda lime components—calcium hydroxide Ca(OH) 2 and calcium carbonate CaCO 3 . Volumetric analysis was performed using automatic burettes at room temperature (23–25 °C). P alkalinity and M alkalinity determinations were performed using 2 mol·L −1 hydrochloric acid solution in the presence of phenolphthalein and methyl orange, respectively. Calcium ion concentration determinations were performed using 0.05 mol·L −1 EDTA solution in the presence of Patton and Reeder’s indicator. Photometric analysis was performed using BWB-XP flame photometer (BWB Technologies, Newbury, England). The content of sodium in samples was measured at an analytical spectral line 589 nm with the limit of detection 0.02 ppm. Thermal behavior and decomposition patterns of samples were investigated using IRIS 209 (Netzsch, Selb, Germany) in the temperature range 25–980 °C at a heating rate of 4°·min −1 in flowing dynamic nitrogen atmosphere (v = 30 mL·min −1 ) using platinum crucibles; as reference material, platinum crucibles were used. PXRD analysis was performed using a X’Pert Pro MPD diffractometer (PANalytical, Malvern, England) in the Bragg–Brentano reflection geometry using CuK α radiation in the 2θ range 5–90° with a step of 0.0167° and exposure per step of 50 s.

3. Results and Discussion

The performed analyses allowed us to determine the composition of the investigated samples SL (F), SL (U), SL (5 min), SL (15 min), SL (30 min). Volumetric and thermogravimetric analyses allowed us to calculate the percentage of the contents of calcium hydroxide and calcium carbonate for each sample. The contents of water were derived from thermal decomposition, as dehydration is the first process of thermolysis. Photometric analysis allowed us to establish the content of sodium hydroxide in investigated samples. Table 1 presents the collected data.

Composition of investigated samples.

Determination of each component’s content was done separately using different analytical techniques, which may have caused propagation of error. This is the reason why, in some samples, the summed contents of components may exceed 100%. The biggest error occurred in SL (15 min) sample (contents sum up to 102.32%); however, it was still within the error tolerance.

3.1. Thermal Decomposition of Samples: SL (F), SL (U), Samples After Absorption: SL (5 min), SL (15 min) and SL (30 min)

Figure 2 presents the thermal decomposition of calcium hydroxide and calcium carbonate. Thermolysis began at 280 °C for calcium hydroxide (DTA peak at 430 °C) and at 560 °C for calcium carbonate (DTA peak at 740 °C). Mass losses and decomposition curves were consistent with reactions that took place during the process. For calcium hydroxide, it was the release of one molecule of water, and for calcium carbonate, it was the release of a molecule of carbon dioxide. The final product of decomposition in both cases was pure calcium oxide.

TG-DTA curves of decomposition of calcium hydroxide and calcium carbonate in nitrogen.

Figure 3 presents TG-DTA curves of the investigated fresh soda lime sample SL (F) and the used sample SL (U) that came from a hospital and were considered exhausted. It is clear that the first step of decomposition was the dehydration of samples. For SL (F), the sample mass loss related to this process was 0.89% at a temperature range 25–275 °C, while for the SL (U) sample, it was 1.93% at a temperature range 25–300 °C. In the second step, one of the samples’ components decomposed—calcium hydroxide. For the SL (F) sample, it took place above 275 °C (DTA peak at 405 °C), while for the SL (U) sample—at a temperature range of 300–400 °C (DTA peak at 390 °C). It was clearly visible that above 400 °C, for the SL (U) sample, decomposition of calcium carbonate took place (DTA peak at 675 °C).

TG-DTA curves of decomposition of fresh soda lime sample (SL (F)) and used soda lime sample (SL (U)) in nitrogen.

Figure 4 presents the TG curves of investigated soda lime samples after 5, 15 and 30 min of carbon dioxide absorption, as well as the TG curve of the SL (F) sample (after 0 min of absorption). The decomposition path was analogous in all cases. The first step (up to 300 °C) was associated with dehydration. It is clear that along with the increasing time of CO 2 absorption, the content of water increased. We could also observe how the second step of decomposition, associated with decomposition of calcium hydroxide, shortened, while the last step, associated with decomposition of calcium carbonate, increased along with time. These curves also show that the process of absorption slowed down with time.

TG curves of decomposition of SL (F) sample and samples after carbon dioxide absorption: SL (5 min), SL (15 min) and SL (30 min) in nitrogen.

Soda lime is a mixture of different chemicals, and thus its thermal decomposition is a multistage process. In all cases, the first step is dehydration. Later, decomposition of calcium hydroxide and calcium carbonate takes place. In order to thoroughly investigate the thermal properties of such absorbents, we decided to study the composition of two sinters prepared at the end of the decomposition process (950 °C). Both the fresh sample’s (SL (F)) and the used sample’s (SL (U)) sinters were prepared. Their X-ray powder diffraction patterns are shown in Figure 5 . These patterns correlate very well with calcium oxide, proving it is a final product of the thermal decomposition.

X-ray powder diffraction patterns of analyzed SL (F) and SL (U) sinters prepared at 950 °C.

3.2. Chemical Kinetics of Carbon Dioxide Absorption by SL (U) Sample

Three samples of SL (F) were exposed to CO 2 absorption for 5 min, 15 min and 30 min, and then analyzed in the same way as SL (F) and SL (U) samples. The absorption of carbon dioxide and water is a multistage process. The first stage involves the formation of carbonic acid from CO 2 and water. Then, NaOH (or KOH) added in small amounts acts as an activator to speed up the process through the formation of sodium (or potassium) carbonate. It can also be concluded that the absorption and the hydration of CO 2 and the formation of CO 3 2− are rapid steps, and the dissolution of Ca(OH) 2 is the slowest step of the carbonation process [ 21 , 22 ]. Calcium hydroxide reacts with the carbonates within minutes to form an insoluble precipitate of calcium carbonate as well as results in a regeneration of NaOH [ 23 ]. Some carbon dioxide may also react directly with Ca(OH) 2 to form calcium carbonates, but this reaction is much slower. In addition, calcium bicarbonate may be formed on the surface of the sorbent particles. The higher solubility of bicarbonate enhances CO 2 diffusion through the bulk of the particle [ 24 ]. Soda lime is exhausted when all hydroxides become carbonates.

Figure 6 shows the experimental results of CO 2 sorption on soda lime as a relationship between conversion rate α and time t . It clearly indicates that the conversion of the sorbent was incomplete and would be difficult to reach under the typical working conditions. According to the results shown in the figure, about 20–30 min from the beginning of the experiment, the carbonation rate slowed down noticeably. We can observe that the curve is composed of two sections. The initial upslope of the curve depicts the fastest rate of carbonation; its initial rate was 3.3 min −1 . After 30 min, the reaction slowed down and reached a rate of 0.17 min −1 . It was a result of significant limitations of CO 2 transport from the surface to the bulk of the sorbent particles, and differentiation between kinetics-controlled and diffusion-controlled ranges occurs.

Relationship between the fractional conversion of soda lime and time of carbon dioxide absorption.

The reaction rate of a solid-state process, d α d t , can be related to the process temperature, T , and to the fraction reacted, α , by means of the following general equation [ 25 ]:

where k is a constant rate.

The kinetic curve of CO 2 absorption of soda lime can be described by the pseudo-first or pseudo-second order kinetic equation [ 26 , 27 ]. In the first section of the kinetic curve, the carbonation is controlled by the surface reaction, whereas in the second section, a heterogeneous system is controlled mainly by diffusion [ 28 ]. Assuming a driving force of CO 2 removal to be proportional to the difference between its concentrations in sorbent at any time prior to equilibrium and its concentration at equilibrium, we can use the equation:

The fittings of the experimental data to the linear form of the two kinetic models, i.e., pseudo-first order and pseudo-second order, are shown in Figure 7 .

Linearized equation of the pseudo-first (right axis) and pseudo-second (left axis) order kinetics models.

The values of the correlation coefficient for linear forms of both kinetic equations are significantly different. The pseudo-second order model describes the kinetic data better than pseudo-first order model when the process is diffusion-limited [ 29 ]. The obtained values of the correlation coefficients were, therefore, 0.999 and 0.712, respectively. Thus, pseudo-first order model does not cover both stages of the CO 2 sorption, i.e., the chemical reaction and the diffusion process. However, the carbonation rate constant determined using first order reaction was greater than for the second order reaction and amounted to 0.011 min −1 and 0.0022 min −1 , respectively. Experimental data have shown that the carbonation process ends before all lime is converted into a calcium carbonate [ 23 ]. On the other hand, the first, fast absorption stage is completed within one hour, and the experimental and calculated values of fractional conversion ( Figure 6 ) were in good agreement with values calculated for both kinetic models: 65.5% and 67.1%, respectively [ 27 ].

One of the possible ways to express soda lime exhaustion rate is a relationship between calcium carbonate content and calcium hydroxide content α   C a C O 3 α   C a O H 2 in the bed and time t . This relationship is presented in Figure 8 .

Relationship between calcium carbonate amount and calcium hydroxide amount ratio and time.

We used this relationship to determine what time of absorption corresponds with the chemical composition of the SL (U) sample. For example, a bed exhaustion rate of 1.247 could be obtained after 21.7 min.

4. Conclusions

Using various analytical techniques, we determined the chemical composition of several soda lime samples: fresh sample, exhausted sample after use in hospital and three samples after carbon dioxide absorption under conditions imitating semi-closed circuit apparatus. Thermogravimetric and XRD analyses comprehensively described the thermal properties and decomposition ways of the investigated samples. This product decomposed in a stable manner, releasing water and carbon dioxide. It is possible to recycle and reuse soda lime in different forms; however, the calcination process would require relatively high temperatures. On the other hand, high temperatures would ensure the biological neutrality of recycled soda lime. Calcium oxide itself could be reused in many different areas, e.g., in absorption and desiccation, in the construction industry or in the manufacturing of chemicals.

Soda lime is a fairly efficient carbon dioxide absorbent that has been used for a long time. It has, however, some limitations and drawbacks that require further investigations, as it is a product used in environments where dependability is a factor of great importance. One of the issues that should be addressed is possible interactions between absorbent and anesthetic gases, which can lead to the release of harmful compounds. Another limitation is the speed of carbon dioxide absorption, which is the highest at the beginning of the process and slows down relatively fast. On the surface of the soda lime granules, water forms a film less than three molecular layers thick, and the reaction rate is reduced [ 21 ]. As the carbonation proceeds, the product particles precipitate on the surface of Ca(OH) 2 and cover it a thicker, porous deposit layer, which inhibits the exchange of reacting species between the surface of calcium hydroxide and bulk solution. Therefore, the diffusion rate of reacting species is an important factor affecting the final stage of carbonation. The carbonation of Ca(OH) 2 was observed to stop before one hour. However, carbonation may go on by diffusion through the covering layer, but its rate is too slow to be detected in the range of carbonation time used [ 28 ]. Thorough research on soda lime properties is an important step in the evaluation of its performance as an agent responsible for preventing carbon dioxide poisoning. Lack of data concerning the kinetics of this process causes this problem to be still very interesting and important in both anesthetic and medical science. Such data would provide a reliable tool to compare different types of absorbents and thus would allow proper absorbent choice taking into consideration all other aspects of the environment or apparatus. Our research, only to some extent, covers the main problems of soda lime use. Additional further studies must be performed in order to ensure the required level of safety and efficiency and to determine the best recycling method.

Author Contributions

Conceptualization, A.C., P.A., P.S. and B.R.; methodology, A.C., P.A., P.S., B.R.; validation, A.C., P.A. and P.S.; formal analysis, P.A, B.R.; investigation, A.C., P.A., P.S., B.R.; resources, A.C., P.S.; data curation, P.A., B.R.; writing—original draft preparation, P.A., B.R.; writing—review and editing, A.C., P.A., P.S., B.R.; visualization, P.A., B.R.; supervision, A.C.; project administration, A.C., B.R; All authors have read and agreed to the published version of the manuscript.

This research received no external funding.

Conflicts of Interest

The authors declare no conflict of interest.

Publisher’s Note: MDPI stays neutral with regard to jurisdictional claims in published maps and institutional affiliations.

Purpose of the Soda & Lime Experiment

By alejandro leopardi / in hobbies.

Soda and lime, sometimes simply referred to as soda lime, is used in numerous scientific experiments in schools, colleges and in the field. Although soda and lime is used for various different experiments, the purpose of their use tends to remain the same for each of those experiments. Understanding the purpose of soda and lime in science experiments is an important foundation for future scientific endeavours.

Basic Purpose of Soda Lime

Soda lime is used for its unique chemical make-up which is a combination of soda ash (Na2CO3) and slaked lime (Ca(OH)2). Although soda lime can be used for various purposes within a given experiment, it is often used to remove carbon dioxide from a particular organism, and sometimes from the surrounding air. The process of removing carbon dioxide (CO2) is either conducted to produce a certain effect, or to remove CO2 so that the experiment can continue.

Softening Water with Soda Lime

Hard water (containing Ca2+ and Mg2+) can be found in many homes and businesses and is identifiable by mineral deposits it leaves behind. Examples of hard water can be soap scum or water residue left on previously washed dishes. In order to remove Ca2+ and Mg2+ from hard water, soda lime is used in the process. What happens when soda lime is added to hard water is that the magnesium (Mg2+) in the water turns into (Mg(OH)2) and the calcium (Ca2+) in the water turns into (CaCO3) after mixing with the soda lime. The Mg(OH)2 and CaCO3 are solids that can be removed from the water, changing it from hard to soft water.

• Hard water (containing Ca2+ and Mg2+) can be found in many homes and businesses and is identifiable by mineral deposits it leaves behind.
• The Mg(OH)2 and CaCO3 are solids that can be removed from the water, changing it from hard to soft water.

Soil Respiration

The forest floor is essentially the top layer, or surface, of soil in the forest. Decomposition, mostly carried out by bacteria and fungi, is necessary in order for plants and trees to receive the proper amount of nutrients. The bacteria and fungi convert carbon (C) into carbon dioxide (CO2) as part of the decomposition process, which is called respiration. To test the rate at which decomposition occurs in a given section of forest, the soda lime experiment is used. A dish of soda lime is placed on the forest floor and then covered by a chamber. The soda lime is pre-weighed as a control in the experiment. The soda lime then absorbs all of the carbon dioxide from the ground, as well as water vapour. Incubation times may vary, but eventually the soda lime is removed and left to dry at 40.6 degrees C. The weight of the dried soda lime is taken and scientists can determine the decomposition time from that.

• The forest floor is essentially the top layer, or surface, of soil in the forest.
• Incubation times may vary, but eventually the soda lime is removed and left to dry at 40.6 degrees C. The weight of the dried soda lime is taken and scientists can determine the decomposition time from that.

Oxygen and Carbon Dioxide of a Small Animal

Respiration in all animals is a combination of processes that include the intake/consumption of oxygen and release of carbon dioxide, among other factors. Soda lime is used in experiments in which the amount of oxygen a small animal consumes is to be recorded. In order to remove carbon dioxide from the equation and final analysis, soda lime is utilised for its ability to absorb all carbon dioxide in a given area. This makes it easier for scientists or researchers to determine only the amount oxygen consumed without the complication of carbon dioxide. At the same time, researchers can determine how much carbon dioxide is released by the animal.

• Respiration in all animals is a combination of processes that include the intake/consumption of oxygen and release of carbon dioxide, among other factors.
• In order to remove carbon dioxide from the equation and final analysis, soda lime is utilised for its ability to absorb all carbon dioxide in a given area.

Syllabus Edition

First teaching 2020

Last exams 2024

Investigating RQs ( CIE A Level Biology )

Revision note.

Biology Lead

Investigating RQs

• Respirometers are used to measure and investigate the rate of oxygen consumption during respiration in organisms
• They can also be used to calculate respiratory quotients
• The experiments usually involve organisms such as seeds or invertebrates

The typical set-up of a respirometer

Equation for calculating change in gas volume

• The volume of oxygen consumed (cm 3 min -1 ) can be worked out using the diameter of the capillary tube r (cm) and the distance moved by the manometer fluid h (cm) in a minute using the formula:

Using a respirometer to determine the Respiratory Quotient

• Always read from the side of the U-tube manometer closest to the respiring organisms
• Reset the apparatus: allow air to re-enter the tubes via the screw cap and reset the manometer fluid using the syringe
• Run the experiment again: remove the soda-lime from both tubes and use the manometer reading to calculate the change in gas volume in a given time,  y cm 3 min -1

Calculations

• x  tells us the volume of oxygen consumed by respiration within a given time
• y  may be a positive or negative value depending on the direction that the manometer fluid moves (up = negative value, down = positive value)
• remembering to read the scale on the side of the U-tube manometer closest to the respiring organisms
• The two measurements  x and  y can be used to calculate the RQ

RQ Equation for Respirometer experiment

Worked example: Calculating RQ from a respirometer experiment

x = 2.9 cm 3 min -1

y = -0.8 cm 3 min -1

( x + y ) / x = RQ

(2.9 - 0.8) / 2.9 = 0.724

When equal volumes of oxygen are consumed and carbon dioxide produced (as seen with glucose) the manometer fluid will not move and  y will be 0, making the RQ 1.

• E.g. temperature – using a series of water baths
• When an RQ value changes it means the substrate being respired has changed
• This is because the RQ of glucose is 1 and the RQ of lipids is 0.7
• Under normal cell conditions the order substrates are used in respiration: carbohydrates, lipids then proteins
• An RQ value of more than 1 suggests excessive carbohydrate/calorie intake
• An RQ value of less than 0.7 suggests underfeeding

There are several ways you can manage variables and increase the reliability of results in respirometer experiments:

• Use a controlled water bath to keep the temperature constant
• Have a control tube with an equal volume of inert material to the volume of the organisms to compensate for changes in atmospheric pressure
• Repeat the experiment multiple times and use an average

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Author: Lára

Lára graduated from Oxford University in Biological Sciences and has now been a science tutor working in the UK for several years. Lára has a particular interest in the area of infectious disease and epidemiology, and enjoys creating original educational materials that develop confidence and facilitate learning.