St. Johns River Water Management District St. Johns River Water Management District St. Johns River Water Management District St. Johns River Water Management District St. Johns River Water Management District St. Johns River Water Management District
St. Johns River Water Management District -

Water supply

Groundwater quality in the St. Johns River Water Management District

Water quality variables

Alkalinity, total, mg/L as CaCO3 (EPA Storet 410) — Alkalinity (as CaCO3) includes bicarbonate (HCO3), carbonate (CO32−), and other anions that can be neutral­ized by a strong acid. Alkalinity is produced by the reaction of carbon dioxide and water in the atmosphere or the soil zone, the dissolution of calcite and dolomite, and the oxidation of organic materials by microorganisms. Bicarbonate is the dominant anion in natural waters and is the dominant constituent of alkalinity. Alkalinity is an indicator of the capacity of solutes in water to react with and neutralize acid.

Calcium, total, mg/L as Ca (EPA Storet 916) — The calcium ion (Ca2+) is dissolved from most soils and rocks and especially from limestone, dolomite and gypsum. Calcium is a dominant cation resulting from the chemical weathering of the mineral calcite (CaCO3) in limestone and the mineral dolomite (CaMg [CO3]) in dolostone. Calcium is also present in large concentrations in some brines, seawater and connate water. The concentration of dissolved calcium is a result of the water’s contact with aquifer materials, residence time and flow path, with highest concentration in alkaline waters that are in equilibrium with aquifer materials. Calcium causes water to be hard and contributes to the scale-forming properties of water.

Chloride, total, mg/L as Cl (EPA Storet 940) — The sources of chloride (Cl) in Florida’s aquifer systems are the lateral intrusion of recent seawater, residual seawater in an aquifer and marine aerosols. Chloride is a relatively inert element that does not readily enter into mineral reactions, so the occurrence of chloride is useful for identifying laterally intruded seawater and mixing of fresh groundwater with residual seawater. Concentrations above 300 mg/L in combination with sodium give a salty taste to water. The drinking water standard for chloride is 250 mg/L.

Fluoride, total, mg/L as F (EPA Storet 951) — Fluoride (F) in Florida’s aquifer systems is derived mostly from the weathering of carbonate-fluorapatite in the Hawthorn Group, indicating that the water has come into contact with the Hawthorn Group at some time in the past, or from the mixing of seawater with groundwater. The drinking water standard for fluoride, which addresses toxicity, is 2.0 mg/L. In small amounts of less than 1.0 mg/L, fluoride is considered beneficial for preventing cavities.

Hardness, total, mg/L as CaCO3 (EPA Storet 900) — Hardness is caused primarily by the presence of calcium and/or magnesium ions in the water, and commonly is reported as an equivalent concentration of calcium carbonate (CaCO3). Natural sources of hardness principally are limestones that are dissolved by percolating rainwater made acidic by dissolved carbon dioxide. Water of hardness up to 60 mg/L is considered soft; 61 to 120 mg/L, moderately hard; 121 to 200 mg/L, hard; and more than 200, very hard. Hardness is often expressed as grains per gallon (gpg). Multiply gpg by 17.1 to get hardness in mg/L.

Magnesium, total, mg/L as Mg (EPA Storet 927) — Magnesium (Mg2+) is dissolved from practically all soils and rocks, especially magnesium-rich clays in the Haw­thorn Group and dolomite in the Floridan aquifer. Magnesium is present in large concentrations in some brines, seawater and connate water. It causes water to be hard and contributes to the scale-forming properties of water.

Nitrate + Nitrite, total, mg/L as N (EPA Storet 630) — Nitrogen compounds occur in groundwater as a result of specific land uses, the leaching of organic soils and from precipitation. Sources of nitrogen from human activities include agricultural fertilizers, animal wastes and human wastes. Nitrogen is transformed between organic nitrogen (TKN), ammonium (NH4+), nitrite (NO2), nitrate (NO3) and other nitrogen compounds depending on oxidation/reduction conditions, microbial activity and plant utilization. Nitrate is the stable form of nitrogen in oxidizing environments. Nitrite is unstable in the presence of oxygen and is present in small concentrations in most waters. Once nitrate enters the aquifer and is isolated from the environments in which denitrification and plant fixation occur, nitrate behaves more-or-less conservatively and can move long distances in aquifers. The drinking water standard for nitrate is 10 mg/L as nitrogen, which is intended to protect human health and is not based on the protection of ecological systems. Nitrate + nitrite concentrations above 0.2 mg/L are considered as above‑natural background conditions. Elevated nutrient concentrations may lead to increases in algae growth, which decreases water clarity and changes the aesthetic qualities and ecology of springs.

Orthophosphate, total, mg/L as P (EPA Storet 70507)/​Phosphorus, total, mg/L as P (EPA Storet 665) — Sources of phosphate (orthophosphate, PO43−) in groundwater include the dissolution of phosphate-rich sediments, leaching from organic sediments and plant materials, agricultural fertilizers, human waste effluent (such as from septic tank systems), industrial effluent and other waste disposal practices. The orthophosphate ion is soluble in acidic waters, such as in siliciclastic horizons of the surficial aquifer system, but insoluble in alkaline aquifers, such as in the Floridan aquifer. In carbonate-rich aquifers, orthophosphate is removed by precipitation of carbonate-hydroxylapatite and alkaline waters seldom have detectable phosphate as a result.

pH, field (EPA Storet 400) — The variable pH reflects the potential for acid-base reactions in water. The pH of aqui­fer water is a result of past chemical reactions, and it is also a measure of the potential for reactions, if chemical equilibrium between the water and surrounding rock has not been established. The pH of a water sample is a measure of the activity of hydrogen ions and is expressed in logarithmic units, with pH representing the negative base-10 log of the hydrogen ion activity in moles per liter. The pH of pure water at 25°C is 7.00, or neutral. Values less than 7.00 are acidic and values greater than 7.00 are basic, or alkaline. Dissolved gases such as CO2, SO2, and NO2 increase the number of hydrogen ions and cause waters to be acid. Carbonates, bicarbonates, hydroxides, phosphates, silicates, and borates decrease the number of hydrogen ions and cause waters to be basic. The hydrogen ion (H+) is generally the cause of acidity, and bicarbonate (HCO3) is generally the source of alkalinity. Water with pH less than 6.5 is likely to be corrosive and have high iron and phosphate levels. Water with pH greater than 8.5 is likely to be corrosive and may result in turbidity from the precipitation of carbonate minerals, or high pH values may be the result of well drilling fluids remaining in the well after construction. Precipitation has a pH of about 5.5. As rainfall infiltrates into the ground, it reacts with carbon dioxide in the soil, resulting in low pH values typically measured in surficial aquifer wells. As water moves downward and interacts with aquifer minerals, acidity is consumed and alkalinity is produced, resulting in higher pH levels. For example, carbonate minerals such as calcite in the Floridan aquifer are highly reactive, raising pH levels.

Potassium, total, mg/L as K (EPA Storet 937) — The principal sources of the potassium ion (K+) are the mixing of fresh groundwater with seawater in coastal transition zones and with connate water, and from the weathering of clays and feldspars. Weathering of potassium feldspars and clays is not considered a dominant process in Florida due to the scarcity of these minerals in aquifer sediments and slow weathering reaction rates. Potassium is rarely present in concentrations over a few milligrams per liter because potassium-rich sediments are scarce in the aquifer system and because potassium is immobilized as a nutrient by plants and sorbed onto clays.

Specific conductance, µmhos/cm at 25°C, field (EPA Storet 94) — Specific conductance is a measure of the ability of water to conduct electric currents and is depend­ent on the concentrations and types of ions in the water and on temperature. Specific conductance has a strong positive correlation with total dissolved solids (TDS) and with chloride concentrations. It is expressed in micromhos per centimeter at 25°C.

Sodium, total, mg/L as Na (EPA Storet 929) — The primary sources of sodium (Na+) in Florida’s aquifer systems are the mixing of fresh groundwater with sea­water along the coast and in coastal transition zones, upconing of seawater from deeper zones, the weathering of sodium-rich minerals such as clays or feldspars, and marine aerosols. Sodium-rich minerals are found in siliciclastic horizons of the surficial and intermediate aquifer system and minor amounts occur in the Floridan aquifer. The drinking water standard for sodium is 160 mg/L.

Sulfate, total, mg/L as SO4 (EPA Storet 945) — Sources of sulfate (SO42−) in groundwater include the dissolution of gypsum and anhydrite, the weathering of pyrite and iron sulfides, residual formation water, the lateral intrusion of seawater, precipitation that contains sulfur oxides and marine aerosols. Sulfate is usually present in mine waters and some industrial waters. Sulfate in water containing calcium forms a hard scale in steam boilers; in large concentrations, sulfate in combination with other ions gives a bitter taste to water. The occurrence of sulfate depends upon the reduction/oxidation potential of the water. In reducing conditions, sulfate reduction produces hydrogen sulfide, which is the cause of the rotten egg odor in some wells. In oxidizing conditions, sulfides may be oxidized to sulfates. The drinking water standard for sulfate is 250 mg/L.

Temperature, °C, field (EPA Storet 10) — The temperature of groundwater is controlled by climatic conditions, cultural activities, heat flow from the earth’s interior, and chemical reactions in the aquifer system. In shallow aquifers, water temperature is usually controlled by climatic conditions, as opposed to other possible causes. Water that has recently entered the aquifer system normally reflects atmospheric temperature at the time of recharge. In deeper aquifer systems, temperature can be affected by recharge from shallow environments, earth heat flow and chemical reactions. Temperature measurements are recorded in the field.

Total dissolved solids, mg/L — Dissolved solids in groundwater are the result of mineral dissolution reac­tions. TDS is a general measure of the total mass of ions dissolved in water and is usually determined from the weight of the dry residue remaining after evaporation of the volatile portion of an aliquot of the sample. Total dissolved solids values are widely used in evaluating and comparing water quality. Water that has recently entered an aquifer in recharge areas has had less time to dissolve minerals than has water that has traveled for some greater distance through the aquifer. TDS concentrations are higher for water from an aquifer composed of reactive materials, such as limestone, than for an aquifer composed of quartz sand materials, which are relatively inert. TDS has a guidance criteria in groundwater of 500 mg/L. Water with a high TDS concentration may have an unpleasant taste, may contribute to the development of kidney stones, and may result in scale or precipitates in hot water heaters.

Additional data

Units: µmhos/cm = micromhos per centimeter
mg/L = milligrams per liter


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