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A Brief Review of Natural Water’s Influence on Scale Formation in Heat Exchangers

| By Brad Buecke

Understanding the water chemistry is a first step in preventing heat-exchanger fouling

At the two power plants and chemical plant in which this author formerly worked, the fresh-water-makeup supplies (two lakes and an underground aquifer adjacent to a river) all had a consistent pH at or slightly above 8.0. This pH range is common for many surface supplies, and typically comes from bicarbonate alkalinity (HCO3) that naturally dissolves in these water bodies. But from where does this alkalinity arise? Per the excellent discussion of this subject in Ref. 1, we briefly explore this issue and also examine what can happen in heat exchangers without proper chemistry control.

 

Geology is one key

One of the most common mineral deposits near the earth’s surface is limestone, whose principal component is calcium carbonate (CaCO3). This versatile mineral serves as a raw material for numerous important industrial and infrastructure applications including concrete, water-treatment chemicals, scrubbing reagent, and simply for gravel roads.

Many surface waters are in contact with limestone formations, and groundwaters often percolate through limestone and settle in aquifers that are bounded by mineral. Calcium carbonate has a strong crystal lattice, and thus CaCO3 is only slightly soluble in water.

 

CaCO3   Ca2+ + CO32– (1)

 

Ksp (25°C) = [Ca 2+][CO32– ] = 4.6 × 10–9

 

Per this solubility product, the molar concentrations of calcium and carbonate in neutral water would be 6.8 × 10–5, which is indeed very slight. But an additional factor must be considered. CO32– is a fairly strong base and will hydrolyze water to some extent.

 

CO32– + H2O   HCO3 + OH (2)

 

Combining Equations (1) and (2) shows the overall reaction of CaCO3 in neutral water.

 

CaCO3 + H2O   Ca2+ + HCO3 + OH (3)

 

Calculations indicate that CO32– hydrolysis of water increases the limestone solubility from 6.8 × 10–5 M at 25°C to 9.9 × 10–5 M. An important point to keep in mind is that these reactions produce hydroxide alkalinity (OH ), even if only in slight concentrations.

Now, let’s look at the other end of the spectrum and a primary reaction that makes natural waters non-neutral and which greatly influences chemistry.

 

Atmospheric influences

Natural waters absorb carbon dioxide from the atmosphere. While it is often argued that much of the CO2 exists as hydrated molecules, the following equations sufficiently represent the chemistry.

 

CO2 + H2O H2CO3 (4)

 

H2CO3 H+ + HCO3 (5)

 

The lowest pH in natural surface waters that can be achieved by these reactions (excluding acid rain issues) is around 5.6, but the solution is still acidic, which is very important. Consider again, Equation (3). When the acidity generated by CO 2 absorption interacts with the alkalinity generated by the fractional CaCO3 dissolution, the hydrogen and hydroxyl ions combine to form water, and per Le Chatelier’s Principle, Reactions (3) and (5) are both driven to the right. This synergistic effect can produce water with a HCO3 concentration of 1 × 10–3 M (equivalent to about 60 parts per million as the species), and “a pH of about 8.3” [1]. The relationship of the carbonate species is clearly illustrated in Figure 1.

FIGURE 1. The relationship between carbonate species in water is shown here. For waters that pass along or through limestone deposits and that absorb CO2 from the atmosphere, the reactions are driven towards the maximum HCO3
– alkalinity

This same acid-base synergy is what makes high-purity limestone (high CaCO3 content) quite reactive and economical as a scrubbing agent, when ground to very fine particles, in wet fluegas desulfurization systems. Aqueous sulfur dioxide (SO2) is a stronger acid than CO2, and analytical data have shown nearly complete CaCO3 reactivity in well-designed scrubbers [2].

 

Carbonate chemistry in reverse

Now let’s briefly examine some of the above chemistry in reverse, and show how it can influence equipment operation and performance at facilities in the chemical process industries (CPI).

From the time humans began heating water for cooking and sanitary purposes, they have undoubtedly observed mineral deposition in heated vessels. These issues became acute following the invention and expanding use of steam engines during the Industrial Revolution of the 18th and 19th centuries. The primary culprit is calcium carbonate.

 

Ca2+ + 2HCO3 + heat

CaCO3 + CO2 + H2O (6)

FIGURE 2. This graph shows the inverse solubility of two of the most common minerals in natural waters

A key aspect of this chemistry is the inverse solubility of CaCO3, as illustrated in Figures 2 and 3.

FIGURE 3. Shown here is CaCO3 scale in an extracted and bisected heat exchanger tube

Accordingly, in many systems that have cooling towers for primary heat exchange, the cycling up in concentration of dissolved elements and compounds in the cooling water, combined with the heat increase in various types and styles of heat exchangers, can lead to serious scale formation. This article highlights CaCO3, which is usually the first deposit to form in untreated or poorly treated cooling water, but other well-known and problematic deposits include, but are not limited to, those shown in Table 1.

And now, as industrial plant personnel, either by choice or mandate, switch from fresh-water supplies to alternatives, such as municipal wastewater-treatment plant effluent, proper scale (and corrosion) control methods have become even more critical. Additional details are available in Ref. 3.

 

References

1. Baird, C., “Environmental Chemistry,” Second Edition, W.H. Freeman and Company, New York, N.Y., 1999.

2. Buecker, B., Wet Limestone FGD Solids Analysis by Thermogravimetry, proceedings of the 24th Annual Electric Utility Chemistry Workshop, May 11–13, 2004, Champaign, Illinois.

3. Buecker, B., (Tech. Ed.), “Water Essentials Handbook”; 2023. ChemTreat, Inc., Glen Allen, Va., available as a downloadable eBook at www.chemtreat.com.

Acknowledgment

All figures courtesy of ChemTreat, Inc.

 

Author

Brad Buecker is president of Buecker & Associates, (Lawrence, Kan.; [email protected]) specializing in technical writing and consulting. Most recently, he served as senior technical publicist with ChemTreat, Inc. Buecker has over four decades of experience in or affiliated with the power and water-treatment industries, much of it in steam generation chemistry, water treatment, air quality control, and results engineering positions with City Water, Light & Power (Springfield, Illinois) and Kansas City Power & Light Company’s (now Evergy) La Cygne, Kansas station. He also spent two years as acting water/wastewater supervisor at a chemical plant, and eleven years with two engineering firms, Burns & McDonnell and Kiewit. Buecker has a B.S. in chemistry from Iowa State University with additional course work in fluid mechanics, energy and materials balances, and advanced inorganic chemistry. He has authored or co-authored over 250 articles for various technical trade magazines, and has written three books on power plant chemistry and air pollution control. He is a member of the ACS, AIChE, AIST, ASME, AWT and the Electric Utility Chemistry Workshop planning committee.