The Weight of Carbonate
The words echoed between my ears as the facility manager continued shouting into his end of the phone.
“THE ENTIRE COOLING TOWER IS SCALED.”
I had been working in water treatment for just over a year at this point, and had never been yelled at like that. I’d seen a handful of problems come up, but hadn’t dealt with any serious issues. So as I started up my truck and headed towards the site, I prayed that they were exaggerating.
The tower was a moderately sized evaporative condenser. Water was pumped from the basin to the top of the unit, sprayed over the condenser coils, and then distributed across tightly packed sheets of plastic fill before cascading back to the basin.
The plastic fill was key to the entire process. Hundreds of angular plastic sheets divided the falling water into thin films and droplets, exposing as much surface area as possible to evaporation into the flowing air. Altogether, the tower contained roughly 60,000 square feet of wetted surface area, approximately the size of a football field.
When I arrived on site the facility manager was waiting next to the cooling tower. “I want to know what the hell happened,” he said as he began spreading apart sections of fill to expose the inside. Nearly every square inch was coated in scale. Some sections were so heavily scaled that adjacent sheets of fill had fused together.
He had not been exaggerating.
I was speechless and immediately began searching for an explanation for what went wrong.
The blowdown solenoid was working correctly. The chemical feed pumps were primed and plugged in. The tower conductivity matched my meter exactly, and was operating below the blowdown setpoint. Everything I had learned to associate with scale looked completely fine.
There was only one thing that stood out.
The controller displayed a pH reading of 7.9, but my meter measured the pH at 8.8. I pulled the pH sensor to inspect the glass electrode and found it coated in scale, which wasn’t surprising based on the condition of the tower.
The pH sensor controlled acid feed to the tower, and I noted that the sulfuric acid level in the tank hadn’t moved since my previous service visit. I checked the pH of the makeup water and found that it was only 7.8.
“I think that something increased the pH of the cooling tower,” I told the facility manager. “The tower’s pH is 8.8 and that’s definitely high enough for calcium carbonate to scale… maybe a leak?”
I was correct, but I was looking in the wrong direction.
Carbonate Does Not Arrive. It Evolves.
Most dissolved ions retain a recognizable identity.
Calcium enters water when minerals dissolve, and it remains calcium. Sodium may come from different salts, but the result is always sodium. Chloride does not spontaneously turn into some other ion because the water becomes warmer or the pH changes.
Inorganic carbon is different.
It enters water from two primary directions:
From rocks: Bicarbonate and carbonate (HCO₃⁻ and CO₃²⁻) are produced as water dissolves limestone and other carbonate-bearing minerals.
From the air: Carbon dioxide (CO₂) dissolves into water from the atmosphere.
However carbon arrives in water, it doesn’t often stay that way long. The moment carbon dissolves, it becomes part of a reversible chemical system, called the carbonate equilibrium:
CO₂(g) ⇌ CO₂(aq) ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺ ⇌ CO₃²⁻ + 2H⁺
There is no true starting point for a chemical equilibrium. It is a chain of reversible reactions that proceeds in either direction, dictated by the contributions of bicarbonate (from rocks) and carbon dioxide (from air). We will choose to begin at the left side of the equation, where carbon dioxide dissolves from the air.
Carbonation: The First Move
The first move begins through physical equilibrium. Any amount of carbon dioxide in the air surrounding water will push some into solution, following Henry’s Law from Chapter 3:
CO₂(g) ⇌ CO₂(aq)
Once dissolved, CO2 does not just float around. It reacts.
CO₂ is a nonpolar molecule, but its individual bonds are highly polar. This leaves the carbon atom electron-deficient and vulnerable to attack by a lone pair on water’s oxygen atom. The result is a chemical reaction, and the first step in the chemical equilibrium, that produces carbonic acid:
CO₂(aq) + H2O ⇌ H₂CO₃
Only a very small fraction of dissolved CO₂ hydrates into true carbonic acid, but the two forms exchange rapidly enough that they are commonly treated together as CO₂/H₂CO₃.
Carbonic acid can donate hydrogen to form bicarbonate:
H₂CO₃ ⇌ H⁺ + HCO₃⁻
Bicarbonate can donate hydrogen again to form carbonate:
HCO₃⁻ ⇌ H⁺ + CO₃²⁻
Each step toward bicarbonate or carbonate releases hydrogen.
Each step back toward carbon dioxide consumes it.
That is why dissolving carbon dioxide generally lowers pH, while removing carbon dioxide can allow pH to rise. But carbon does not get to choose where it rests. Its position is determined by the water surrounding it.
The pH Lever
Several conditions influence the carbonate equilibrium, including concentration, temperature, pressure, and gas exchange.
The most useful lever for understanding which species will dominate is pH.
Looking at the equilibrium equation, this should not be much of a surprise. It’s the only other ion that participates directly in the equilibrium, and it appears in both dissociation constants.
H₂CO₃ ⇌ H⁺ + HCO₃⁻
HCO₃⁻ ⇌ H⁺ + CO₃²⁻
This makes hydrogen (H⁺), and hydroxide (OH⁻) by extension, fundamental determinants of the equilibrium position.
At low pH, hydrogen ions (H⁺) are abundant. Carbonate accepts hydrogen and becomes bicarbonate. Bicarbonate accepts additional hydrogen and becomes carbonic acid and dissolved carbon dioxide. The equilibrium shifts toward the left.
Below approximately pH 6.3, dissolved carbon dioxide and carbonic acid predominate.
At intermediate pH, bicarbonate dominates. There is enough hydrogen to suppress most carbonate, but not enough to force all inorganic carbon into dissolved carbon dioxide and carbonic acid. Between approximately pH 6.3 and 10.3, bicarbonate is the largest carbonate species. That is an enormous range, and it includes most natural and industrial waters.
At high pH, hydrogen becomes scarce. Bicarbonate increasingly releases hydrogen and shifts toward carbonate. Above approximately pH 10.3 carbonate dominates.
Long before that, though, carbonate matters: near pH 8.3, carbonate is still only about one percent of the dissolved inorganic carbon, but calcium carbonate is so sparingly soluble that even that small fraction becomes operationally important once calcium and alkalinity are concentrated.
The same carbon atoms remain in the water. Only their arrangement changes.
The Rise of Carbonate
Back at the scaled cooling tower, the difference between pH 7.9 and 8.8 looked small on the controller. But chemically, it was a big leap in the wrong direction.
The relationship between bicarbonate and carbonate is governed by the second dissociation equilibrium:
HCO₃⁻ ⇌ H⁺ + CO₃²⁻
Because pH is logarithmic, the carbonate-to-bicarbonate ratio changes approximately tenfold for every one-unit change in pH, so a difference of 0.9 units represents a change of almost eightfold.
At an actual pH of 8.8, the ratio was nearly eight times greater than the controller's 7.9 reading suggested. The probe was not wrong by "less than one point." It was misrepresenting the position of the carbonate equilibrium by almost an order of magnitude.
But that still did not explain why the tower pH had risen. The makeup entered at pH 7.8, no caustic had been added, and no alkaline leak had been found. I had spent the entire investigation looking for something that entered the water.
The answer was that something had left.
How the Cooling Tower Broke the Balance
Cooling towers intensify a process that also occurs in rivers, waterfalls, and aeration systems: gas exchange. Every pass across the fill creates the ideal conditions for CO₂ removal.
The tower had raised its own pH by allowing carbon dioxide to escape.
Warm return water reduces gas solubility and encourages dissolved carbon dioxide to leave. Massive airflow continually replaces the air near the water surface, maintaining a strong driving force for gas transfer. Enormous surface area divides the water into thin films and droplets, giving carbon dioxide countless opportunities to escape.
As dissolved carbon dioxide leaves, the carbonate equilibrium shifts to replace it:
CO₂(g) ← CO₂(aq) ← H₂CO₃ ← HCO₃⁻ + H⁺
Carbonic acid converts back into dissolved carbon dioxide. Bicarbonate consumes hydrogen to replace the carbonic acid. The hydrogen concentration decreases. The pH rises.
Nothing was dosed.
Nothing failed chemically.
The system simply obeyed equilibrium.
The Failure Fed Itself Twice
The chemical response was natural. The control failure was not.
The acid feed was designed to hold the line. As carbon dioxide left and pH increased, the controller should have activated the sulfuric pump. The added acid would have replenished hydrogen, pushed the equilibrium back toward bicarbonate and carbonic acid, and reduced the amount of carbonate available to form scale.
But there was a problem with the probe.
Scale and debris had accumulated around the electrode. Its response time slowed and the reading began to sag. The controller saw pH 7.9, within the upper bound of the programmed control range, and it saw no need to call for acid.
As the real pH climbed, more carbonate formed, more scale coated the glass, and the reading fell further behind the truth. The failure fed itself once.
The problem was further exacerbated by the operation of the cooling tower itself. Cycles of concentration made the consequences much worse.
Evaporation removed nearly pure water while calcium, alkalinity, and other dissolved minerals accumulated in the basin. Carbon dioxide stripping continued to push the equilibrium toward higher pH, while increasing concentrations of calcium and alkalinity moved the water closer to calcium carbonate saturation.
The tower concentrated the ingredients while changing the conditions that kept them dissolved. The system fed itself twice.
Once carbonate found calcium, it took less than a week to lay down over 2,000 pounds of scale.
I had plenty of time to think through the chemistry while vacuuming and shoveling most of it from the basin and fill. It was among the most menial labor I had ever done, but it was also one of the most effective chemistry lessons I have ever received.
pH didn’t cause the scale. It gave carbonate permission to build it.
And cooling towers aren’t the only system that pushes carbon in the wrong direction.
Where Alkalinity Goes in Boilers
Cooling towers create carbonate trouble by stripping CO₂ from recirculating water. Boilers create a similar problem in reverse.
Boilers aren’t exposed to atmospheric exchange the way cooling towers are. Instead, carbonate chemistry enters with the makeup water, primarily as bicarbonate alkalinity.
Once inside the boiler, a new lever takes over: Heat.
As water is heated, bicarbonate becomes thermally unstable and converts into carbonate, water, and carbon dioxide:
2HCO₃⁻ ⇌ CO₃²⁻ + H₂O + CO₂↑
The resulting carbon dioxide cannot remain dissolved under boiler conditions. It volatilizes with steam and is carried into the header.
Steam generation continuously removes the carbon dioxide from the boiler water and carries it into the distribution system.
The carbonate remains behind.
That creates two separate risks.
Scaling Potential
The newly formed carbonate creates an intense scaling risk with any calcium present.
Boiler water is hot and concentrated. CaCO₃ becomes less soluble as temperature rises, and heat-transfer surfaces provide ideal locations for precipitation. Even a minor hardness excursion can produce deposits where heat transfer matters most.
This is why boiler feedwater must be effectively softened.
At higher operating pressures and temperatures, further conversion of carbonate alkalinity to hydroxide and carbon dioxide becomes increasingly important:
CO₃²⁻ + H₂O ⇌ 2OH⁻ + CO₂↑
The hydroxide remains in the boiler water.
The carbon dioxide leaves with the steam, once again.
Condensate Corrosion
The released carbon dioxide travels through the steam distribution system.
As steam cools and condenses, the carbon dioxide dissolves into the newly formed condensate.
CO₂(g) ⇌ CO₂(aq)
It then reacts with water to reform carbonic acid:
CO₂(aq) ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺
Carbonic acid releases hydrogen and lowers the condensate pH.
Condensate contains very little alkalinity. There is almost no chemical reserve available to absorb the disturbance. A relatively small amount of carbon dioxide can therefore create a sharply acidic and corrosive environment. Carbon steel begins dissolving. Iron enters the condensate. Grooves and channels form along the bottoms of horizontal return lines where acidic condensate flows.
The carbonate equilibrium has completed a full circuit.
Bicarbonate entered with the feedwater. Heat converted it into carbon dioxide. Steam carried the carbon dioxide out of the boiler. Condensate absorbed it. Carbonic acid formed.
The equilibrium successfully transformed a buffering species into a corrosion mechanism, only hundreds of feet downstream.
How We Interrupt the Cycle
Carbonate chemistry cannot be ignored.
But it can be managed.
Cooling Towers
In cooling towers, acid adds hydrogen back into the water and pushes the equilibrium away from carbonate:
CO32⁻ + H+ ⇌ HCO₃⁻
HCO₃⁻ + H+ ⇌ H2CO₃
Lowering pH does not remove calcium or alkalinity.
It changes the form in which inorganic carbon primarily exists and reduces the carbonate available for calcium carbonate precipitation.
The strategy works only when the control system measures pH accurately and the acid feed remains operational.
Boilers
In boilers, there are two different strategies for dealing with equilibrium.
Remove Alkalinity Upstream
Reverse osmosis and dealkalization can remove bicarbonate before it enters the boiler.
Less bicarbonate entering the boiler means less carbonate formed in the boiler water, less alkalinity-derived carbon dioxide leaving with steam, and lower condensate-treatment demand.
This strategy removes much of the precursor before the reaction begins.
Neutralize Acidity Downstream
Neutralizing amines travel with the steam and enter the condensate system.
When the steam condenses, the amine raises condensate pH and neutralizes acidity before carbonic acid can aggressively attack the metal. This does not prevent carbon dioxide from forming. It manages the consequence after the carbon dioxide has entered the steam.
One strategy removes the precursor.
The other neutralizes the product.
Both are attempts to control where the carbonate equilibrium is allowed to land.
The Other Side of Equilibrium
Moving carbon through dissolved CO₂, carbonic acid, bicarbonate, carbonate produces very different consequences under different conditions. For most of the industrial systems we treat, it’s not good news.
As we saw in cooling towers, sustained carbon dioxide loss drives pH upward and promote scale. In boilers, heat converts bicarbonate into an even greater scaling concern, while simultaneously delivering carbonic acid to its unbuffered condensate system.
But somehow, this same equilibrium provides incredible stability to natural water.
The reaction is the same. The conditions decide the consequence.
The Buffered Response
In rivers, lakes, and groundwater, the carbonate system acts as a chemical shock absorber.
It evolves as air, water, soil, and rock interact.
Carbon dioxide enters water from the atmosphere and from biological activity in soil. Some of it reacts to form carbonic acid, lightly depressing pH. The added acidity allows the water to dissolve carbonate-bearing minerals beneath the surface.
CaCO3 + CO2 + H2O ⇌ Ca2⁺ + 2HCO₃⁻
Carbon dioxide came from the air and soil.
Carbonate came from rock.
Their encounter produced bicarbonate.
Within the pH range occupied by many natural waters, bicarbonate becomes the dominant form of dissolved inorganic carbon. Its position in the middle of the carbonate equilibrium gives it an extraordinary advantage: it can respond in either direction.
When acid enters, bicarbonate accepts hydrogen and shifts toward carbonic acid and dissolved carbon dioxide:
HCO₃⁻ + H⁺ ⇌ H2CO3 ⇌ CO2 + H2O
When base enters, bicarbonate can release hydrogen and shift toward carbonate:
HCO₃⁻ ⇌ H⁺ + CO32⁻
Bicarbonate does not prevent pH from changing.
Its role is to moderate disturbances.
That may sound like a modest task, but it is difficult to overstate its importance. Bicarbonate is one of the most consequential ions on the planet. It carries carbon through rivers, groundwater, oceans, and living systems. It controls acid demand, moderates pH, participates in mineral weathering, and helps determine whether water remains chemically stable.
Without bicarbonate, many natural waters would be dangerously fragile. Small additions of acid or base would produce much larger swings in pH, disrupting the chemical conditions required by aquatic organisms and altering the solubility and toxicity of substances throughout the water. Bicarbonate stands in the middle of the equilibrium, able to negotiate from either direction.
Managing Carbonate
Carbonate chemistry is often ignored because it is easy to mistake it for a side reaction.
It is not.
Carbonate is the dominant buffer system in natural waters, the hidden driver behind pH stability, acid demand, cooling tower scaling, and condensate corrosion. It is not misbehaving. It is doing exactly what equilibrium demands.
The real question in water treatment is never whether carbonate chemistry will act.
It is where it will act, and how much leverage it will have when it does.
Fortunately, the equilibrium that governs it behaves predictably. If you understand carbonate chemistry, you are no longer reacting to scale, corrosion, or acid demand.
Your job as a water treatment consultant is to decide deliberately where that equilibrium is allowed to land. I learned that lesson shoveling two thousand pounds of scale out of a cooling tower. The physical weight of an equilibrium I had failed to control.
Engineering Notes: Carbonated Water
“If you are reading this straight through, you can skip this section and lose nothing essential to the story. These notes are for the operators, engineers, and technicians who need to do the math.”
Why This Matters in the Field
Carbonate chemistry is not a mystery, it just takes time to internalize. These calculations let you:
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Predict pH drift in open systems: CO₂ stripping drives pH upward even if nothing is added.
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Estimate scale risk: carbonate speciation controls CaCO₃ precipitation potential.
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Quantify boiler CO₂ generation: alkalinity in makeup becomes carbonic acid in condensate.
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Budget chemical demand: neutralizing amines and acid feed are often proportional to alkalinity-derived CO₂.
Core Tools & Constants
| Constant / Conversion | Value |
|---|---|
| CaCO₃ molecular weight | 100 g/mol |
| CaCO₃ equivalent weight | 50 g/eq |
| CO₂ molecular weight | 44 g/mol |
| 1 meq/L alkalinity | = 50 mg/L as CaCO₃ |
| CO₂ from alkalinity (boiler rule) | 0.79 ppm CO₂ per 1 ppm alk as CaCO₃ |
| Total (M) alkalinity titration endpoint | pH ≈ 4.3 |
| CO₂/HCO₃– crossover | pH ≈ 6.3 |
| Phenolphthalein (P) alkalinity titration endpoint | pH ≈ 8.3 |
| HCO₃–/CO₃²– crossover | pH ≈ 10.3 |
Carbonate System Recap
The carbonate equilibria:
CO₂(g) ⇌ CO₂(aq) ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺ ⇌ CO₃²⁻ + 2H⁺
Although carbonic acid is often shown explicitly, in practice it is convenient to treat dissolved CO₂ and H₂CO₃ as a combined species. This is done because they exist in rapid equilibrium with one another, and analyses generally measure their combined concentration.
Practical translation:
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Low pH: CO₂/H₂CO₃ dominate (corrosion zone)
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Neutral range: HCO₃⁻ dominates (buffering zone)
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High pH: CO₃²⁻ grows rapidly (scale zone)
Concentrations Expressed “mg/L as Calcium Carbonate”
In water treatment, alkalinity is rarely reported as the actual mass of bicarbonate, carbonate, or hydroxide present. Instead, it is usually expressed as mg/L as CaCO₃. That convention exists because these species vary in molecular weight and, more importantly, they vary in the amount of acid they can neutralize.
A milligram of bicarbonate does not neutralize the same amount of acid as a milligram of carbonate. A milligram of hydroxide is different again. If we compare them only by mass, we are comparing apples to oranges. What matters in alkalinity chemistry is not just how much material is present, but how much neutralizing capacity it carries.
That is why water treatment uses equivalent weight.
Molecular weight tells you how much a mole weighs, which determines how much of a substance is present. Equivalent weight tells you how much of a substance is needed to do a fixed amount of chemistry. In terms of alkalinity, that usually means: how much of this substance is needed to neutralize one mole of hydrogen ion (H+)?
For example:
-
OH⁻ neutralizes 1 H⁺
-
HCO₃⁻ neutralizes 1 H⁺
-
CO₃²⁻ neutralizes 2 H⁺
So even though these species have different masses, we can convert them to a common basis by asking how much acid each one can neutralize.
Calcium carbonate became the standard reference because its numbers are convenient. One mole of CaCO₃ weighs 100 g, and one mole can neutralize 2 moles of H⁺. That means its equivalent weight is:
100 ÷ 2 = 50 g/eq
This makes the conversion useful:
50 mg/L as CaCO₃ = 1 meq/L
Once concentrations are expressed as mg/L as CaCO₃, different chemicals can be placed on the same footing. It converts them into a common chemical currency based on neutralizing capacity.
Converting Alkalinity to milliequivalents
With alkalinity expressed as mg/L as CaCO3, regardless of which species are present, it can be converted directly to milliequivalents per liter:
(mg/L as CaCO₃) ÷ 50 = meq/L
This works because all of those concentrations have already been converted to the same neutralizing basis per unit volume.
If alkalinity is 100 mg/L as CaCO₃:
meq/L = 100 ÷ 50 = 2 meq/L
Which can be expressed in equivalents per liter:
eq/L = 2 ÷ 1,000 = 0.002 eq/L
This idea extends beyond alkalinity. A milliequivalent per liter expresses concentration in terms of reactive charge. One meq/L of positive charge will always balance one meq/L of negative charge.
That is why hardness ions (like calcium and magnesium) are often reported as mg/L as CaCO3. This places both hardness and alkalinity on the same equivalent basis, where they can be compared directly in softening calculations, scaling indices, and charge-balance relationships. Without that common basis, every comparison would require a separate molecular-weight conversion.
A mole counts molecules. An equivalent counts reactive charge. In water treatment, meq/L is often the more useful unit because it tells you how much chemistry the water can do.
P- and M-alkalinity Speciation
M-alkalinity is measured to about pH 4.3 and represents the total acid-neutralizing capacity contributed by hydroxide, carbonate, and bicarbonate.
P-alkalinity is measured to about pH 8.3 and represents the alkalinity neutralized before the phenolphthalein endpoint. It includes all hydroxide alkalinity and half of the carbonate alkalinity, but it does not include bicarbonate alkalinity.
That half-carbonate behavior is the key.
During the P-alkalinity titration, hydroxide is neutralized and carbonate is converted to bicarbonate:
OH⁻ + H⁺ → H₂O
CO₃²⁻ + H⁺ → HCO₃⁻
But the titration stops at pH 8.3, before bicarbonate is neutralized to carbonic acid. The M-alkalinity titration continues to about pH 4.3, where bicarbonate is converted to carbonic acid/CO₂:
HCO₃⁻ + H⁺ → H₂CO₃ ⇌ CO₂ + H₂O
This is why comparing P to M allows us to infer which alkalinity species are present.
Relationships (as CaCO₃):
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If P = 0 → all alkalinity is bicarbonate
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If P < ½M → mix of bicarbonate + carbonate
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If P = ½M → all alkalinity is carbonate
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If P > ½M → mix of carbonate + hydroxide
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If P = M → all alkalinity is hydroxide
| Condition | Hydroxide (OH−) | Carbonate (CO32−) | Bicarbonate (HCO3−) |
|---|---|---|---|
| P = 0 | 0 | 0 | M (All Bicarbonate) |
| P < ½ M | 0 | 2P | M−2P |
| P = ½ M | 0 | 2P (All Carbonate) | 0 |
| P > ½ M | 2P−M | 2(M−P) | 0 |
| P = M | M (All OH) | 0 | 0 |
Doing multiple titrations in the field can feel like busywork. But when you understand the role of carbonate equilibrium, it becomes chemical reconnaissance.
Estimating pH from OH-Alkalinity
If you know the Hydroxide Alkalinity (calculated from the table above), you can estimate high-range pH without a meter.
Example: OH-Alkalinity: 300 ppm as CaCO3.
Step 1: Convert to Molarity
Alkalinity is reported as Calcium Carbonate Equivalents:
Equivalent Weight of CaCO3 = 50,000 mg/eq
Therefore:
300 mg/L ÷ 50,000 mg/eq = 0.006 eq/L
Because hydroxide is monovalent, it carries one equivalent per mole. This means that eq/L and mol/L are numerically identical here:
[OH-] = 0.006 mol/L
Step 2: Calculate pOH
pOH = -log10(0.006) = 2.22
Step 3: Calculate pH
pH = 14 – 2.22 = 11.78
Note: Some handbooks calculate this by dividing by 100 (Molecular weight) instead of 50 (Equivalent weight). That is incorrect for alkalinity conversions and leads to a pH error of ~0.3 units.
CO₂ Generation in Boilers (Rule-of-Thumb)
When bicarbonate decomposes inside a boiler, it produces CO₂ that travels with steam and re-dissolves in condensate.
Thermal Decomposition (driven by heat):
2HCO₃⁻ ⇌ CO₃²⁻ + H₂O + CO₂↑
Further Breakdown (at high pressure):
CO₃²⁻ + H₂O ⇌ 2OH⁻ + CO₂↑
Field Estimation:
For many low- to medium-pressure boilers, a useful rule of thumb is:
1 ppm M-Alkalinity as CaCO₃ (feedwater) → 0.79 ppm CO₂ (in steam/condensate)
Note: This is a field estimate, not a universal stoichiometric law. Actual CO₂ generation depends on boiler pressure, alkalinity form, cycles, steam purity, and operating conditions.
Interpretation:
100 ppm alkalinity as CaCO3 in feedwater → ~79 ppm CO₂ in steam/condensate
This can be managed in two places: upstream by removing alkalinity before the boiler, or downstream by neutralizing the carbonic acid formed in the condensate.
For downstream management, amine demand is estimated:
Target: pH 8.3 – 8.8
Dosage: A stoichiometric starting estimate is approximately 2–3 ppm active neutralizing amine per ppm CO₂, depending on the amine. Actual product dosage must account for blend composition, distribution ratio, condensate return, recycle, losses, and the target condensate pH.