lime cycle experiment

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Lesson Video: Limestone Chemistry

In this video, we will learn how to describe the uses of limestone and the related reactions, by using the lime cycle.

Video Transcript

In this video, we will learn how to describe the uses of limestone and the related reactions by using the lime cycle.

Limestone is an abundant, naturally occurring rock made up of primarily calcium carbonate. Often we use the terms limestone and calcium carbonate interchangeably. Although there are other oxides and minerals found in limestone, their presence is small enough to be able to refer to calcium carbonate as limestone without much issue. Hundreds of millions of tons of limestone are mined around the world each year because it has a variety of uses. The majority of limestone is used in the construction industry. Limestone is a key ingredient in many useful materials, like brick, cement, and glass.

Limestone is also added to foods like bread to provide calcium, a necessary mineral for your body. It’s used in toothpaste as an abrasive to gently scrub your teeth. It can provide color and texture to cosmetics. Also, since it comes in a variety of colors and tends to chip instead of fracture, it makes an ideal material for sculptures. For the most part, these uses rely on the physical properties of calcium carbonate. As we’ll see later on in the video, there are also a variety of uses that rely on the chemical properties of limestone.

Limestone can be processed to form other substances with similar qualities. One such process is the breakdown of calcium carbonate into carbon dioxide and calcium oxide. This reaction takes a lot of heat to take place, so we may see the word heat written above the arrow in the equation. Another way to signify that heat is added to the reaction is a 𝛥 symbol above the arrow instead of the word heat. This reaction is an example of thermal decomposition or, in other words, breaking down a compound using heat. Since the reactant, calcium carbonate, breaks down into multiple products, it’s a decomposition reaction. Since we use heat to do it, it’s a thermal decomposition reaction. When we thermally decompose calcium carbonate, frequently referred to as limestone, one of the products is carbon dioxide. This product is usually released as waste.

The other product, calcium oxide, is also known as quicklime. Quicklime has a variety of uses. Quicklime is used to refine steel, to make paper, to make fiberglass, and its water-absorbing nature makes it an excellent drying agent. We can take the processing of limestone another step farther. If we take quicklime and hydrate it or add water to it, we end up with calcium hydroxide, also known as slaked lime. Our starting compound, quicklime, is extremely dry. It can absorb a lot of water. Slaked lime is a dry powder as well, but it won’t absorb much water at all. Adding water to slaked lime forms a wet slurry. For this reason, it’s used in wet materials like mortar, plaster, and cement.

If we mix quicklime with an excess of water or take slaked lime and add even more water, we’ll get limewater, which is the name we give the aqueous form of calcium hydroxide. We might want calcium to be present in water, for example, to treat drinking water or to preserve the mineral balance in an aquarium. If we used slaked lime, we’d have to wait for the solid to dissolve in the water. In its aqueous form of limewater, the calcium hydroxide is already dissolved, so it can more easily spread to the boundaries of the container. And as we will see in a moment, limewater can also be used as a test for CO2.

In our next equation, if we take calcium hydroxide, either slaked lime or limewater, and carbonate it or add carbon dioxide gas, we will create water and limestone. We can use this reaction to our advantage during a clever process that lets us test for the presence of carbon dioxide gas. We may want to know if carbon dioxide is the product of a certain reaction, in which case we can place the reactants in a test tube. We can place a stopper connected to tubing in the top of the test tube. That way, the gas product that is released can only go through the tube. The other end of the tube is placed in a solution of saturated limewater, saturated meaning there’s just enough water to dissolve the calcium hydroxide.

When the gas bubbles through the limewater, we may see it turn milky white. In this case, the test for CO2 is simple. If the limewater turns cloudy white in the presence of a gas, that’s a sign that the insoluble calcium carbonate has formed. Based on this third equation, we can see that limewater forms limestone when mixed with carbon dioxide. So the presence of the milky-white precipitate is a sign that the gas is indeed carbon dioxide. If the limewater does not turn white, then the gas is not carbon dioxide. You may have noticed that we ended right where we started. We heated limestone to get quicklime. We hydrated quicklime to get calcium hydroxide. And we carbonated calcium hydroxide to get limestone once again.

We can visualize this family of reactions another way. This simple representation is called the lime cycle. Using the same three reactions as before, if we heat limestone, we get quicklime. If we hydrate quicklime, we get limewater or slaked lime. And if we carbonate limewater or slaked lime, we get limestone once again. We can also draw the lime cycle to include the waste at each step. When we heat limestone, CO2 is released. When we hydrate quicklime, heat escapes. And when we carbonate limewater or slaked lime, water is also produced. An umbrella term for all of these calcium-containing compounds is lime. Since they have similar chemical properties, they share many uses. Which specific variant we select may depend on whether we want a wet, dry, or unprocessed substance.

Let’s take a look at some more uses of lime. One significant use of the chemical properties of lime is in agriculture. One common problem on farms is soil that is too acidic. When something is too acidic, we can neutralize it with a neutralization reaction, also known as an acid–base reaction. In a neutralization reaction, an acid and a base combine to form water and a salt. In the case of acidic soil, we can add it to the basic limestone to neutralize it. The products of this reaction are neutral, so adding limestone to the acidic soil raises the pH of the soil. Limestone allows farmers to boost the productivity of their soil. While other lime compounds have been used for this purpose in the past, today we primarily use limestone as it’s less harmful to the plant and animal life surrounding the farm.

Another use of lime is in a process called flue gas desulfurization, relevant to the pollution produced by factories. One problem that factories face is that the burning of fossil fuels produces sulfur dioxide gas. SO2 can react with the water in the atmosphere, contributing to acid rain. Acid rain is harmful to the plants, animals, soil, and buildings in the area. One solution to this problem is to install a flue gas desulfurization system. This system is a series of processing units that filter, stir, heat, mix, or otherwise process the gas in order to minimize its harmful effects. There are a variety of chemical reactions that take place inside the flue gas desulfurization system.

One example for a system that uses limestone is this reaction here, where sulfur dioxide combines with limestone to form carbon dioxide and calcium sulfite. This reaction takes the harmful sulfur dioxide gas and makes new products out of it. Calcium sulfite can be used to make gypsum, a key ingredient in making plaster and drywall. The CO2 is released as waste. So while the system isn’t perfect, it is preferable to polluting with the more dangerous sulfur dioxide. This reaction uses limestone in order to capture the sulfur, but we could use another lime product in its place.

For example, another reaction that could occur in a flue gas desulfurization system would be the combination of sulfur dioxide and slaked lime to produce water and calcium sulfite. For reactions like this, we say that the lime captures the sulfur. Instead of being released as waste, it’s captured in a product, calcium sulfite, that we can use for other purposes. Overall, flue gas desulfurization systems use lime to capture sulfur, reducing sulfur dioxide pollution from factories. Limestone is obtained by mining it from a quarry or a large open pit where the stone can be removed from the earth. However, mining limestone is not without controversy. Let’s review the pros and cons of this process.

Mining limestone provides necessary materials for the variety of industries we’ve mentioned already. Without mining limestone, it would be difficult or impossible for these industries to find substances that could do the job instead of limestone. Next, many limestone mines are in rural areas where work is harder to come by. A limestone mine offers steady employment for a variety of people in the area. As an extension of that, it also helps the economy. On the local level, shop owners and service providers in the area have a larger, better paid customer base. On a national level, limestone can be exported to other countries to make money.

There are also some downsides to mining limestone. Limestone is loosened from the earth by blasting it with dynamite. It is then pulled up with heavy equipment and carried away by trucks. These explosions and heavy machinery create a lot of noise pollution. The noise is unpleasant to the surrounding townsfolk and harmful to their well-being. All of this machinery releases exhaust fumes that pollute the air. More noticeably, the blasting of limestone creates airborne dust particles. These particles can carry through the air and settle in a wide range far beyond the mine itself. This is harmful to the respiratory health of the residents as well as the local wildlife.

Next, limestone mines are often found in remote areas, often near regional or national parks. Creating a limestone mine essentially takes a hill or mountain and turns it into a road covered rock pit. This transformation disrupts the visual beauty of the countryside. Lastly, digging into the ground to pull up mine stone can destabilize the ground and water in the area. Water sources upstream from the mine can flow into the empty space created by the mine. Water that flows through the mine can take on additional sediment and pollutants, lowering the water quality in the area. Also, when limestone is removed and water flows underground, it can create sinkholes, where the land collapses due to a lack of support. Overall, mining limestone is a useful and profitable yet disruptive and harmful process.

Now that we’ve learned about limestone, let’s do a practice problem to review.

In old brick limekilns, such as the one in the picture, large quantities of limestone were continually added and heated to very high temperatures. What was the main waste gas that escaped through the chimney? What solid limestone derivative was collected from holes in the bottom of the kiln?

This question is asking about a reaction involving limestone, also known as calcium carbonate. While the limestone pulled from the earth does have other minerals in it, it is composed primarily of calcium carbonate. For that reason, in chemistry, we use the terms limestone and calcium carbonate interchangeably. Another clue about the nature of this chemical reaction is that the limestone is being heated. When a chemical reaction requires heat to proceed, we may write the reaction arrow with the 𝛥 symbol above it to signify that heat is being added. We might also see the word heat in place of the 𝛥.

The two parts of this question are asking about the substance that escaped or the substance that was collected at the end of the reaction. In other words, this question is asking, what are the products of this chemical reaction? We may be familiar with this reaction because it’s part of the lime cycle. The lime cycle consists of products related to limestone and the reactions that create them. The reaction that’s part of the lime cycle that involves heating limestone creates carbon dioxide and calcium oxide as products. We call a reaction like this thermal decomposition because we’re using heat to break down a compound into multiple products.

But the question remains, which product corresponds to which part of the question? Well, the waste gas is carbon dioxide. When carbon dioxide is produced, it is almost always as a gas and not as a solid like the other part of the question would suggest. So our answer to the first part of the question is carbon dioxide.

Next, which product is a solid limestone derivative? That’s the other product, calcium oxide. As we move around the lime cycle, we obtain a variety of limestone derivatives. The thing that they have in common is that they all contain calcium. So the solid limestone derivative that the second part of the question is asking about is calcium oxide, which we can write in as our answer. Another name for calcium oxide is quicklime. It’s not necessary to include the name quicklime in our answer. But it’s important to recognize that calcium oxide and quicklime are different names for the same substance.

Because it can be easily created by heating naturally occurring minerals, the use of quicklime as a construction material dates back quite far, at least 6000 years. Quicklime was used in mortars and plasters to create the pyramids of Egypt, the Roman aqueducts, and the Great Wall of China, among other historical structures. So in brick limekilns, such as the one in the picture, what was the main waste gas that escaped through the chimney? That’s carbon dioxide. And what solid limestone derivative was collected from holes in the bottom of the kiln? That’s calcium oxide, also known as quicklime.

Now that we’ve learned about limestone, let’s review the key points of the video. Limestone is an important ingredient in construction materials, food, cosmetics, and more. Limestone, or calcium carbonate, can be processed into quicklime, calcium oxide; slaked lime, solid calcium hydroxide; and lime water, which is aqueous calcium hydroxide. These materials and the processes that turn one into another are referred to as the lime cycle. Limestone can also be used to neutralize acidic soil in agriculture. Lime is used in flue gas desulfurization, a process that removes sulfur dioxide from the waste gas of factories. Lastly, limestone must be mined from the earth. This process is controversial. While it helps the economy with profits and jobs, it creates visual pollution, noise pollution, and air pollution that can be harmful.

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The Cycle of Reactions involving Limestone and Products Made From it

• Limestone is a sedimentary rock composed mainly of calcium carbonate (CaCO3).
• When limestone is heated, it undergoes a chemical reaction called thermal decomposition. This process turns calcium carbonate into calcium oxide and carbon dioxide (CaCO3 -> CaO + CO2).
• The common name for calcium oxide is ‘quicklime’ or ‘burnt lime.’ It has a strong, alkaline reaction with water and generates significant heat.
• Adding water to quicklime (calcium oxide) creates slaked lime, also known as calcium hydroxide (CaO + H2O -> Ca(OH)2).
• The process of adding water to calcium oxide is called slaking, which is highly exothermic, releasing a large amount of heat.
• Calcium hydroxide is slightly soluble in water and is often referred to as limewater in its aqueous state.
• When carbon dioxide is passed through limewater, a reaction occurs resulting in the formation of calcium carbonate (Ca(OH)2 + CO2 -> CaCO3 + H2O).
• This reaction implies the process can be cycled, going from limestone (calcium carbonate) back to limestone via quicklime (calcium oxide) and slaked lime (calcium hydroxide).
• Limestone and its products have many industrial uses, including in the manufacture of cement, glass, iron, steel, and concrete.
• Limestone is also used to neutralise acidity in soils and in lakes, to purify sugar, and in the making of toothpaste.
• Quicklime is used in the steel industry to remove impurities and in the manufacture of bricks.
• Slaked lime is used in the food industry, such as for pickling and pH adjustment. It’s also used in wastewater treatment.

Remember, each of these substances - limestone, quicklime, and slaked lime - has distinct physical and chemical properties. Also worth noting is the environmental impact of limestone quarrying. While it provides essential raw materials, it can damage landscapes, create noise and dust pollution, and affect local wildlife. The industry must balance the demand for limestone with the environmental impact of quarrying.

Try to take time to understand this cycle thoroughly, as knowing how each reaction works will help you predict other similar chemical reactions.

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Lime and its Production

The Lime Cycle The lime cycle shows the stages from quarrying the limestone through to the production of mortars and plasters for our buildings and how it slowly, through the re-absorption of Carbon Dioxide, reverts to its original chemical form (Calcium Carbonate) in the wall.

Lime burning

Limestone (Calcium Carbonate – CaCO3) is burnt in a kiln giving off Carbon Dioxide (CO2) gas and forming Calcium Oxide (CaO) which is commonly known as Quicklime or Lumplime. 

It needs to be burnt at 900°C to ensure a good material is produced. The temperature at which it is burnt will affect its reactivity in all other stages of the limecycle – slaking and carbonating. The resulting lime is at its most volatile and dangerous at this stage.

Lime slaking The Burnt Lime or Quicklime is then combined with water (slaked) as quickly as possible. From the moment it is burnt the material starts to degrade by ‘air-slaking’. Combining Quicklime (CaO) and water (H20) produces Calcium Hydroxide (Ca(OH)2 - slaked lime and heat. There are three main ways of slaking the Quicklime:

• in an excess of water to produce a putty;
• in a shortfall of water to produce a powder - hydrated or bag lime;
•  in damp sand to produce a hot mix. 

Lime carbonation Lime sets by absorbing water soluble Carbon Dioxide from the air. This process is called carbonation. The ‘set’ or carbonation must occur slowly – the slower the set the better (it is not a case of just drying), therefore direct heaters or dehumidifiers do not help and may cause failures – it is therefore vitally important that the conditions are right to enable the water-borne Carbon Dioxide (CO2) to be absorbed. In our Lime Handbook, specific conditions are described to control the carbonation process. Failure to properly control carbonation will lead to problems and potentially failure on site. It is, therefore, an extremely important process to come to understand.

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Joint Earth Science Education Initiative - limestone weathering

This activity is designed for students aged 11-14. It can be used to reinforce work on the reactions of carbonates with acids as well as the chemical weathering of rock.

The chemical attack on limestone by rain that is naturally acidic (containing dissolved carbon dioxide) and ‘acid rain’ (rain that is more acidic because of dissolved pollutants such as sulfur dioxide and nitrogen oxides).

If you teach primary science, see the headings below to find out how to use this resource:

Skill development

Children will develop their working scientifically skills by:

• Drawing conclusions and raising further questions that could be investigated, based on their data and observations.
• Using appropriate scientific language and ideas to explain, evaluate and communicate their methods and findings.

Learning outcomes

Children will:

• Compare and group together different kinds of rocks on the basis of their appearance and simple physical properties.

Concepts supported

Children will learn:

• That rocks have different properties, often as a result of the type of rock they are and how they are formed.
• That rocks are natural materials whereas bricks are man-made (also referred to as manufactured).

Suggested activity use

This activity can be used as a whole-class investigation into the properties of rocks, in particular limestone. This leads on to children being able to make suggestions about the properties and possible uses of rocks, based on their findings from this experiment. The activity can lead onto subsequent investigations into testing other properties of rocks, such as hardness, permeability and reaction with acids.

Practical considerations

Equipment for the activity will need to be sourced prior to the lesson, including universal indicator and limestone. Also, if you are progressing on to the properties of different types of rocks, a selection of other rocks will be needed.

In order for children to understand what is happening, prior knowledge of the colours of universal indicator in the presence of acidic, alkaline and neutral solutions is needed. Also, children require knowledge of gases, in particular carbon dioxide and oxygen, as well as understanding what we breathe in and breathe out.

If not carefully managed, you may introduce or reinforce the misconception that carbon dioxide is an ‘acidic gas’, whereas in fact it produces an acidic solution when dissolved in water.

As with all experiments, a thorough risk assessment should be carried out along with other health and safety considerations. It must be stressed to children that they need to blow out through the straws and not suck in.

You should note that the national curriculum links at the beginning of the document are now out of date.

Limestone weathering

• 11-14 years
• Practical experiments
• Acids and bases
• Reactions and synthesis

Specification

• a metal carbonate + an acid → a salt + water + carbon dioxide
• (f) the acid/carbonate reaction as a test for acidic substances and CO₃²⁻ ions
• (g) the environmental effects and consequences of the emission of carbon dioxide and sulfur dioxide into the atmosphere through the combustion of fossil fuels
• 2.5.8 explain that cracking involves the breakdown of larger saturated hydrocarbons (alkanes) into smaller more useful ones, some of which are unsaturated (alkenes); and
• Hardness in water. Causes of temporary and permenant hardness.
• Acid rain and its effects on the environment.
• Carbonates react with dilute acids to form carbon dioxide gas. Carbon dioxide can be identified with limewater.
• The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.
• Students should be able to: describe the use of universal indicator or a wide range indicator to measure the approximate pH of a solution.
• Recall that acids react with some metals and with carbonates, and write equations predicting products from given reactants.
• Evaluate the evidence for additional anthropogenic causes of climate change, including the correlation between change in atmospheric carbon dioxide concentration and the consumption of fossil fuels, and describe the uncertainties in the evidence base.
• 3.11 Explain the general reactions of aqueous solutions of acids with: metals, metal oxides, metal hydroxides, metal carbonates to produce salts
• 8.25 Evaluate the evidence for human activity causing climate change, considering: the correlation between the change in atmospheric carbon dioxide concentration, the consumption of fossil fuels and temperature change; the uncertainties caused by the…
• 9.5C Describe tests to identify the following ions in solids or solutions as appropriate: carbonate ion, CO₃²⁻, using dilute acid and identifying the carbon dioxide evolved
• C1.3.2 evaluate the evidence for additional anthropogenic causes of climate change, including the correlation change in atmospheric carbon dioxide concentration and the consumption of fossil fuels, and describe the uncertainties in the evidence base
• C5.2.4 describe tests to identify aqueous cations and aqueous anions and identify species from test results including: tests and expected results for metal ions in solution by precipitation reactions using dilute sodium hydroxide (calcium, copper, iron(I…
• C6.1.1 recall that acids react with some metals and with carbonates and write equations predicting products from given reactants
• C6.1.4 recall that relative acidity and alkalinity are measured by pH including the use of universal indicator and pH meters
• C3.3f recall that carbonates and some metals react with acids and write balanced equations predicting products from given reactants

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Limestone Cycle

Subject: Chemistry

Age range: 14-16

Resource type: Lesson (complete)

Last updated

2 October 2014

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Excellent visuals on the powerpoint. Makes the chemical reactions clear.

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Troubleshooting the lime cycle at Kraft pulp mills

White liquor is used in the Kraft cooking process to liberate the fibers in the wood. A vital step in the white liquor preparation is causticization, where green liquor is converted to white liquor, which requires good quality lime. The lime is produced in the lime cycle, which is sensitive to different disturbances.

In order to ensure the adequate production of white liquor for the cooking process, the mill relies on a smooth operation of the lime cycle: the causticizing and the lime reburning stages. However, these vital pulp mill units can experience numerous problems, the most common being low causticizing degree, poor filtration and dewatering properties of lime mud as well as dead load and ring formation in the lime kiln. In many cases, these problems are connected to accumulation of non-process elements, e.g. Si, P, Al or Mg. 

Laboratory experiments at RISE

Investigation of the reasons behind, and the solutions to, a lime cycle problem usually starts with a thorough analysis of the involved streams. Elemental composition, including non-process elements, of green and white liquor, lime mud, lime, lime kiln ESP dust and makeup is measured. Other important quantities, like free CaO of lime, EA and TTA of liquors etc. are also evaluated. If necessary, the chemical analysis of the solids is accompanied by other methods, e.g. electron microscope imaging. The results are compared with data from RISE’s extensive database of mill measurements.

The causticization experiments are conducted at carefully controlled temperature. Causticization degree is calculated after ABC-titration of filtered liquor. The composition of liquor and lime can be modified in order to simulate different conditions at the mill, e.g. introduction of a new process solution and its effect on the operation of the lime cycle. Next step involves calcination experiments in RISE’s laboratory oven. The oven can be heated to 1500°C and the atmosphere can be adjusted, from inert (N2) to custom-supplied gas mixtures. This type of experiment is a very good approximation of a regular lime kiln operation and provides valuable knowledge about the condition of your lime mud and the achievable levels of free CaO.

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iLia Fresco (Anossov) | artist

The medium of fresco – carbon based art for carbon based life | fresco painter, sculptor, teacher, lime cycle – fresco plaster.

Lime Cycle is the process by which limestone or marble (pure form of lime stone) is converted to quicklime by heating, then it is slaked (hydrated), and reverted back to limestone (marble) by carbonation.

Lime Cycle – fresco plaster, lime plaster

Lime Plaster is composed of slaked lime (calcium hydroxide) and an aggregate such as sand, mixed with water. In Western culture scholars are dating lime plasters as far back as to 4th century BC. It was widely used in Ancient Rome where it came from Greece. Greek origins of Roman Plaster are mainly responsible for its mixture proportions of 1 part lime 1 part fine aggregate for the final color coat (referred to as marmorino). During the Renaissance the mixture proportion of the final color coat was changed to more painting friendly proportion of 5 parts lime to 8 parts aggregate. Technically, the proportion itself did not change, rather the final, marmorino coat of polished plaster was omitted and painting began to be done on finely troweled undercoat. I believe that at this point, to distinguish the difference in plaster finish, the “new final” painting plaster coat began to be referred to as “Intonaco” assuming its’ name from the word “plaster” – “Intonaco” in Italian. Hence – “painting on wet plaster.”

Complete Lime Cycle: CaCO3 + heat → CaO + CO2 → CaO + H2O → Ca(OH)2 → Ca(OH)2 + CO2 → CaCO3 + H2O

How to get started with fresco painting?

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