Introduction: water earns its “taste” by travelling underground
Up to #10, the story was about water level — a matter of physics. From here we step into the water’s contents: its chemistry.
A spring is soft, a hot spring is salty, another smells metallic. Why do waters that are all just “water” differ so much? The answer is simple: as water travels underground, it slowly dissolves the rocks and gases it meets along the way. Rain, which starts out containing almost nothing, takes up carbon dioxide in the soil, dissolves minerals, and sometimes mixes with ancient seawater. That history is written into the water’s “taste” — its chemistry.
This article reviews the major ions that set groundwater chemistry, organizes the four mechanisms that create it, then follows the most universal of them — calcite dissolution — by calculating it with the geochemical code PHREEQC. Finally we compare river, groundwater, seawater and hot spring on a Piper diagram.
The major ions — an alphabet
The composition of natural water can be described almost completely by eight major constituents: four cations (Na⁺, K⁺, Ca²⁺, Mg²⁺) and four anions (Cl⁻, SO₄²⁻, HCO₃⁻, NO₃⁻). These eight are the alphabet of the language of water chemistry.
Each ion carries a story about where the water has been:
| Type | Ion | Main origin | What high values indicate |
|---|---|---|---|
| Cation | Na⁺ | halite, deep thermal water, seawater | Na-Cl thermal water, seawater mixing |
| Cation | K⁺ | weathering of feldspar, volcanic thermal water | deep high-temperature thermal water |
| Cation | Ca²⁺ | limestone, gypsum, carbonate rocks | Ca-HCO₃ (bicarbonate) water, hard water |
| Cation | Mg²⁺ | dolomite, mafic (Fe–Mg) minerals | young, immature groundwater |
| Anion | Cl⁻ | sea salt, deep thermal water, halite | deep Na-Cl thermal water, seawater |
| Anion | SO₄²⁻ | gypsum, sulfide oxidation, volcanic gas | sulfate springs, mine drainage |
| Anion | HCO₃⁻ | dissolution of CO₂ and carbonate rocks | bicarbonate springs |
| Anion | NO₃⁻ | fertilizer, sewage, forest soil | anthropogenic (human) contamination |
Cl⁻ deserves special note: because almost no rock contains it, chloride in groundwater is a near-conservative tracer of sea salt, deep thermal water, or halite. The shape these eight ions make on a hexadiagram (Stiff diagram) is a water’s chemical “face”.
The metallic smell noted at the outset, however, is not among these eight. Reducing groundwater can carry dissolved iron (Fe²⁺); drawn up and exposed to air, it oxidizes into reddish-brown iron hydroxide — the source of that metallic tang, and of the rust-colored stains on taps and laundry. Such trace constituents are governed less by the balance of major ions than by oxidation–reduction (redox) — a theme we return to later in the water-quality thread of this series.
Why chemistry varies — four mechanisms of water-quality formation
The reactions that create groundwater chemistry come down to four.
- Dissolution (weathering) — CO₂-bearing water dissolves minerals. The most basic reaction: calcite dissolution (→ Ca-HCO₃ type), feldspar weathering (→ Na, K), and so on.
- Cation exchange — on clay-mineral surfaces, Ca²⁺ in the water swaps with Na⁺ held by the mineral. Important when a seawater-invaded aquifer freshens, producing Na-HCO₃ waters.
- Evaporative concentration — in arid settings or long-residence waters, evaporation concentrates ions until salts precipitate.
- Mixing — fresh water blends with deep thermal water, or with seawater. Ubiquitous in freshwater lenses (#7) and in coastal / geothermal areas.
Of these, dissolution is the road almost every groundwater travels first. Below we follow its archetype — rainwater dissolving calcite (CaCO₃) into a Ca-HCO₃ groundwater — calculating with PHREEQC.
Calculating water-quality formation with PHREEQC — calcite dissolution
Setup
Rain, just before it infiltrates, is almost pure water. But the soil is rich in CO₂ — tens of times atmospheric — from root and microbial respiration. The rain takes up this CO₂ to become a weak carbonic acid, and begins to dissolve limestone (calcite). The reaction is a single line:
\[\mathrm{CaCO_3 + CO_2 + H_2O \;\rightleftharpoons\; Ca^{2+} + 2HCO_3^-}\]
Assuming an open system in which soil CO₂ is continuously supplied, we fix the partial pressure at \(p\mathrm{CO_2}=10^{-2}\,\mathrm{atm}\) and let calcite dissolve to saturation. The PHREEQC input is short:
SOLUTION 1 Rainwater
temp 25
units mmol/kgw
pH 5.6 # rain that has taken up CO2
-water 1 # kg
EQUILIBRIUM_PHASES 1
CO2(g) -2.0 # fix soil CO2 (open system, pCO2 = 10^-2 atm)
Calcite 0.0 10 # equilibrate with calcite (ample supply)
SELECTED_OUTPUT
-pH true
-totals Ca C(4)
-si Calcite
ENDEQUILIBRIUM_PHASES fixes the CO₂ pressure while equilibrating with calcite — that is the open-system, calcite-saturation calculation.
Result: evolution toward a Ca-HCO₃ water
The result is Figure 1. The horizontal axis is the amount of calcite dissolved (= Ca²⁺ taken into solution).
Three things stand out.
- Ca²⁺ : HCO₃⁻ = 1 : 2, exactly the stoichiometry of the reaction above.
- pH rises, as the acid (CO₂) is consumed and alkalinity (HCO₃⁻) accumulates.
- SI approaches 0 and dissolution stops. The water reaches equilibrium with calcite — a classic Ca-HCO₃ groundwater (pH ≈ 7.3, Ca ≈ 1.4 mmol/L ≈ 56 mg/L).
The Ca concentration at saturation scales roughly as \(p\mathrm{CO_2}^{1/3}\). In settings with high soil CO₂ (\(p\mathrm{CO_2}=10^{-1.5}\) to \(10^{-1}\) atm), Ca reaches 2–4 mmol/L. PHREEQC from scratch #5 defines a carbonate groundwater with “Ca ≈ 3–4 mmol/kg, calcite SI ≈ 0” under higher CO₂ — continuous with the calculation here. The full PHREEQC workflow is taught from the ground up in the PHREEQC series.
This calculation uses a dilute approximation (activity coefficients = 1). For high-ionic-strength waters such as seawater, PHREEQC applies Debye–Hückel activity corrections automatically; for quantitative work, let PHREEQC handle it.
Checking against real data — from the literature
Fluoride enrichment in Kumamoto: dissolution sets the chemistry
A telling example of dissolution controlling chemistry is the groundwater of western Kumamoto. Hossain et al. (2016) analyzed the geochemical processes by which fluoride (F⁻) becomes locally enriched there. Dissolution of volcanic-rock minerals, combined with residence time and pH, can push fluoride above drinking-water limits. Which mineral dissolves under which conditions decides the chemistry — sometimes the water-quality problem.
Classifying Beppu hot springs: a multivariate view
The archetype of mixing-controlled chemistry is a hot spring. Yang (2021) measured the major constituents of many Beppu spring sources and used multivariate analysis (cluster and principal-component analysis) to classify the spring types and their vertical distribution. Deep Na-Cl thermal water and shallow Ca-HCO₃ groundwater mix in varying proportions to give each source its character; one source was estimated at 12% seawater mixing from its hexadiagram.
Not just reading the major ions but classifying them with multivariate statistics — treating water-quality data with the same mindset as the time-series methods from #8 onward.
The Piper diagram — a whole-picture view
Finally, we view four contrasting waters — river, groundwater, hot spring and seawater — on a Piper (trilinear) diagram (Figure 2). The lower-left triangle is the cations, the lower-right the anions, each plotting relative proportions (equivalent %). The upper diamond is obtained by projecting the points from the two triangles until they intersect, and it tells the water type at a glance.
- The diamond is divided into four water-type fields (I: Na-Cl, II: Ca-Mg-SO₄, III: Ca-HCO₃, IV: Na-HCO₃).
- Groundwater and river fall in field III (Ca-HCO₃) — precisely the product of the calcite dissolution calculated above.
- Hot spring and seawater fall in field I (Na-Cl), chemistry governed by deep thermal water and seawater mixing.
On a single figure, the differences in formation mechanism appear directly as differences in position. That is the power of the Piper diagram.
Summary
- Groundwater chemistry can be understood as the history of four mechanisms: dissolution, cation exchange, evaporative concentration, and mixing.
- The most basic — calcite dissolution — can be tracked quantitatively in PHREEQC. Rain takes up CO₂, dissolves calcite in a Ca : HCO₃ = 1 : 2 ratio, evolves toward a Ca-HCO₃ type, and stops at saturation.
- The Piper diagram shows water type and formation process at a glance; the diamond point is the intersection of the projections from the two ternary plots.
Taking the Kumamoto fluoride enrichment of Hossain et al. (2016) touched on above, the next article reads that paper closely and reproduces its water-quality formation in PHREEQC — testing how far real groundwater chemistry can be explained with public data and geochemical modeling.
References
- Appelo, C.A.J. & Postma, D. (2005) Geochemistry, Groundwater and Pollution, 2nd ed. Balkema.
- Parkhurst, D.L. & Appelo, C.A.J. (2013) Description of input and examples for PHREEQC version 3. U.S. Geological Survey Techniques and Methods, book 6, chap. A43.
- Hossain, S., Hosono, T., Yang, H., Shimada, J. (2016) Geochemical Processes Controlling Fluoride Enrichment in Groundwater at the Western Part of Kumamoto Area, Japan. Water, Air, & Soil Pollution, 227(10), 385.
- Yang, H. (2021) Classification and vertical distribution of Beppu hot-spring water quality using hydrochemical methods and multivariate analysis. Journal of Groundwater Hydrology, 63(3), 137–149. (in Japanese)
- Yang, H., Mishima, T., Katazakai, S., Kagabu, M. (2023) Analytical approach using a chemical equilibrium formula and geochemical modeling for alkalinity measurements of small natural water samples. Applied Geochemistry, 148, 105535.