Calthemite is a secondary deposit, derived from concrete, lime, mortar or other calcareous material outside the cave environment. Calthemites grow on or under man-made structures and mimic the shapes and forms of cave speleothems, such as stalactites, stalagmites.

Degrading concrete has been the focus of many studies and the most obvious sign is calcium-rich leachate seeping from a concrete structure. Calthemite stalactites can form on concrete structures and "artificial caves" lined with concrete (e.g. mines and tunnels) significantly faster than those in actual caves. This is because the majority of calthemites are created by chemical reactions which are different from normal "speleothem" chemistry. Calthemites are usually the result of hyperalkaline solution (pH 9–14) seeping through a calcareous man-made structure until it comes into contact with the atmosphere on the underside of the structure, where carbon dioxide (CO₂) from the surrounding air facilitates the reactions to deposit calcium carbonate as a secondary deposit. CO₂ is the reactant (diffuses into solution) as opposed to speleothem chemistry where CO₂ is the product (degassed from solution). Calthemites are generally composed of calcium carbonate (CaCO₃) which is predominantly coloured white, but may be coloured red, orange or yellow due to iron oxide (from rusting reinforcing) being transported by the leachate and deposited along with the CaCO₃. Copper oxide from copper pipes may cause calthemites to be coloured green or blue.

The way stalactites form on concrete is due to different chemistry than those that form naturally in limestone caves and is the result of the presence of calcium oxide (CaO) in cement. Concrete is made from aggregate, sand and cement. When water is added to the mix, the calcium oxide in the cement reacts with water to form calcium hydroxide (Ca(OH)₂), which under the right conditions can further dissociate to form calcium (Ca²⁺) and hydroxide (OH⁻) ions. The chemical reactions are reversible and several may occur simultaneously at a specific location within a concrete structure, influenced by leachate solution pH.

Calcium hydroxide will readily react with any free CO₂ to form calcium carbonate (CaCO₃). This reaction occurs in newly poured concrete as it sets, to precipitate CaCO₃ within the mix, until all available CO₂ in the mix has been used up. Additional CO₂ from the atmosphere will continue to react, typically penetrating just a few millimetres from the concrete surface. Because the atmospheric CO₂ cannot penetrate very far into the concrete, there remains free Ca(OH)₂ within the set (hard) concrete structure. Any external water source (e.g. rain or seepage) which can penetrate the micro cracks and air voids in set concrete will readily carry the free Ca(OH)₂ in solution to the underside of the structure. When the Ca(OH)₂ solution comes in contact with the atmosphere, CO₂ diffuses into the solution drops and over time the reaction deposits calcium carbonate to create straw shaped stalactites similar to those in caves.

The presence of soluble potassium and sodium hydroxides in new concrete complicate things as it supports a higher solution alkalinity of about pH 13.2 – 13.4, thus the predominant carbon species is CO₃²⁻ and the leachate becomes saturated with Ca²⁺ and a slightly different chemical reaction occurs.  As the soluble potassium and sodium hydroxides are leached out of the concrete along the seepage path, the solution pH will fall to pH ≤12.5. Below about pH 10.3, the more dominant chemical reaction will return the concrete setting type reaction. The leachate solution pH influences which dominant carbonate species (ions) are present, so at any one time there may be one or more different chemical reactions occurring within a concrete structure.

In very old lime, mortar or concrete structures, possibly tens or hundreds of years old, the calcium hydroxide (Ca(OH)₂) may have been leached from all the solution seepage paths and the pH could fall below pH 9. This could allow a similar process to that which creates speleothems in limestone caves to occur. Hence, CO₂ rich groundwater or rainwater would form carbonic acid (H₂CO₃) (≈pH 7.5 – 8.5) and leach Ca²⁺ from the structure as the solution seeps through the old cracks.

Leachate solutions creating calthemites can typically attain a pH between 10–14, which is considered a strong alkaline solution with the potential to cause chemical burns to eyes and skin – dependent on concentration and contact duration.

The growth rates of calthemite stalactite straws, stalagmites and flowstone etc., is very much dependent on the supply rate and continuity of the saturated leachate solution to the location of CaCO₃ deposition. The concentration of atmospheric CO₂ in contact with the leachate, also has a large influence on how quickly the CaCO₃ can precipitate from the leachate. Evaporation of the leachate solution and ambient atmospheric temperature appears to have very minimal influence on the CaCO₃ deposition rate. Calthemite straw stalactites precipitated (deposited) from hyperalkaline leachate have the potential to grow up to ≈200 times faster than normal cave speleothems precipitated from near neutral pH solution.

If the drip rate is quicker than one drop per minute, most of the CaCO₃ will be carried to the ground, still in solution. The leachate solution then has a chance to absorb CO₂ from the atmosphere (or degas CO₂ depending on reaction) and deposit the CaCO₃ on the ground as a stalagmite.

In most locations within manmade concrete structures, calthemite stalagmites only grow to a maximum of a few centimetres high, and look like low rounded lumps. This is because of the limited supply of CaCO₃ from the leachate seepage path through the concrete and the amount which reaches the ground. Their location may also inhibit their growth due to abrasion from vehicle tires and pedestrian traffic.

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