Where acidic groundwater comes from
While hard water leaves minerals behind, acidic groundwater does the exact opposite: it aggressively strips materials away. It is highly corrosive and can silently dissolve a home's plumbing infrastructure from the inside out.
Acidic groundwater is primarily a result of local geology and surface biology failing to buffer the natural acidity of rain.
When rain falls, it absorbs atmospheric carbon dioxide (CO₂) to form a dilute carbonic acid (H₂CO₃). This gives natural, unpolluted rainwater a slightly acidic pH of around 5.6. As this water percolates down through the topsoil, it encounters decaying organic plant matter, which releases even more CO₂ and natural organic acids into the water, driving the pH down further.
What happens next depends entirely on the rock beneath the soil:
- Neutralization (hard water): If the water flows through sedimentary rocks like limestone or dolomite, the calcium and magnesium in the rock neutralize the water, raising its pH while making it hard in the process.
- Continued acidity (corrosive water): If the water flows through inert, non-carbonate rocks like granite, sandstone, or quartz, there are no neutralizing minerals available. The groundwater reaches the aquifer retaining its low pH — typically between 5.0 and 6.5.
The chemistry of acidic corrosion
Acidity is a measure of the concentration of free hydrogen ions (H⁺) in a liquid. Acidic water is highly reactive because it is seeking equilibrium — for electrons to balance itself out.
When this water enters a home's plumbing system, the abundant H⁺ ions begin to aggressively attack and oxidize the metal pipes, stripping away electrons and causing the solid metal to dissolve into the water as metallic ions. For example, copper metal (Cu⁰) is oxidized into aqueous copper ions (Cu²⁺).
Effects on household plumbing
- Blue-green stains: The hallmark symptom of acidic water in homes with copper plumbing. As acidic water dissolves the interior of copper pipes, it carries copper ions to fixtures. When the water evaporates in sinks, tubs, and toilets, it leaves behind vivid, hard-to-remove blue or green stains.
- Pinhole leaks: Acidic water does not dissolve pipes uniformly. It attacks specific weak points, gradually thinning the metal until it breaches. The result is microscopic pinhole leaks that can cause severe hidden water damage inside walls or ceilings.
- Heavy metal leaching: The most significant health risk. Acidic water dissolves not just copper, but also the lead solder used to join older pipes together, and the brass used in modern fixtures. This leaches toxic heavy metals directly into the drinking water supply.
- Metallic taste: As iron, copper, or zinc is stripped from pipes and suspended in the water, it often gives drinking water a sharp, bitter, or distinctly metallic taste.
- Fixture and appliance degradation: The corrosive water slowly eats away at the metallic heating elements in water heaters, the internal valves in washing machines, and the chrome plating on faucets, leading to premature appliance failure.
pH adjustment choices
There are a few practical ways to adjust the pH of acidic well waters and it depends on what else may be in your water. Consider the following example of a typical water you may find in the area's wells: moderately high hardness and alkalinity of 150 ppm. The example describes the effects and outcome when raising the pH from 6 to 7.5*.
*I use this example of a minimum for upper pH adjustment since in practice a final pH of 7.5–8.0 appears to be the most effective protection range for copper pipes. I've often seen corrosion below a pH of 7.5, particularly when the TDS is either very low or very high.
Because alkalinity is 150 mg/L while the pH is thoroughly acidic at 6.0, the water contains a very high concentration of dissolved carbon dioxide (carbonic acid). To raise the pH to exactly 7.5, the acid load needs to be neutralized. Here is how sodium hydroxide (NaOH) and calcite (CaCO₃) compare in their effectiveness at reaching your target and their impact on total dissolved solids (TDS).
1. Calcite (CaCO₃)
Effectiveness at reaching pH 7.5: highly ineffective. Calcite is a weak, self-limiting base. When acidic water passes through calcite, it dissolves according to this equilibrium reaction:
According to Le Chatelier's principle, a reaction is slowed if the products are already present in high concentrations. Because the well water already has 150 mg/L of calcium hardness and 150 mg/L of alkalinity, the water will resist dissolving more calcite. Even with infinite contact time, calcite will stall out around pH 7.0 to 7.2. It will not reach 7.5. Typically, a portion of the treatment tank volume will contain magnesium oxide (MgO) — Corosex brand name — to buffer the pH higher.
Influence on final TDS & hardness: massive increase. Any calcite that does dissolve directly adds calcium hardness and bicarbonate to your water. TDS will increase significantly, and hardness could easily push past 200–250 mg/L, making the water highly scale-forming and likely requiring a water softener downstream.
2. Sodium hydroxide (NaOH)
Effectiveness at reaching pH 7.5: highly effective. NaOH is a strong liquid base injected via a chemical metering pump. Because it does not rely on natural dissolution of a mineral, the pump can be tuned to dose the correct amount to dial in the desired pH — 7.5 in this example.
Influence on final TDS & hardness: moderate TDS increase, zero hardness increase. Injecting NaOH adds sodium to the water as it neutralizes the acid. TDS increase will be moderate. However, NaOH contains no calcium or magnesium, meaning your hardness will remain exactly at 150 mg/L.
Conclusion: For a well with an already high starting hardness and alkalinity, sodium hydroxide is the best of the two for neutralizing your water. Using calcite will fail to reach the target pH while simultaneously making your hard water much harder. However, when TDS is very low, the dissolution of calcite/MgO adds needed TDS that may make the water less corrosive.
Tank storage can raise the pH
When this well water is pumped from its pressurized underground environment to an open storage tank at sea level, it transitions from a closed system (isolated from the atmosphere, relatively higher pressure) to an open system (in equilibrium with atmospheric gases).
As the water equilibrates with the atmosphere at 25 °C, it undergoes massive off-gassing of carbon dioxide. Consequently, the pH rises significantly, stabilizing at approximately 8.5.
Sidebar: closed vs open system chemistry
1. The initial state (underground well water from ~120 ft at 15 °C). In the aquifer, the water is under hydrostatic pressure (from being 40 meters deep) and is likely insulated from the atmosphere. Biological or geological processes often pump CO₂ into groundwater, acting as a weak acid and driving the pH down to 6.0.
2. The transition: open-system equilibrium. When the water is pumped into the open storage tank at sea level and warms to ~25 °C, two critical things happen:
- Degassing: The excess CO₂ escapes into the atmosphere until the dissolved concentration matches the atmospheric partial pressure.
- Alkalinity is conserved: As long as no solid minerals precipitate, the total alkalinity remains strictly constant, even as CO₂ off-gasses.
In this open system at 25 °C, the dissolved carbonic acid concentration is fixed by Henry's Law and the atmospheric CO₂ concentration (assuming a modern atmospheric level of roughly 415 ppm). This should theoretically demonstrate a shift in the pH to 8.5. Under these conditions a well water with moderate alkalinity of, say, 100 mg/L will almost certainly become supersaturated with calcium carbonate as the pH hits 8.5, and the tank can develop scaling. However, in practice, drinking water storage tanks are usually sized to turn over within a few days and can never reach "equilibrium," or that final pH of ~8.5, so scaling is not typical.
Water in your home doing any of this?
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