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Carbonate hardness
Carbonate hardness
Carbonate hardness
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Carbonate hardness, is a measure of the water hardness caused by the presence of carbonate (CO2−
3) and bicarbonate (HCO−
3) anions. Carbonate hardness is usually expressed either in degrees KH (°dKH) (from the German "Karbonathärte"), or in parts per million calcium carbonate ( ppm CaCO
3 or grams CaCO
3 per litre|mg/L). One dKH is equal to 17.848 mg/L (ppm) CaCO
3, e.g. one dKH corresponds to the carbonate and bicarbonate ions found in a solution of approximately 17.848 milligrams of calcium carbonate(CaCO
3) per litre of water (17.848 ppm). Both measurements (mg/L or KH) are usually expressed as mg/L CaCO
3 – meaning the concentration of carbonate expressed as if calcium carbonate were the sole source of carbonate ions.

An aqueous solution containing 120 mg NaHCO3 (baking soda) per litre of water will contain 1.4285 mmol/l of bicarbonate, since the molar mass of baking soda is 84.007 g/mol. This is equivalent in carbonate hardness to a solution containing 0.71423 mmol/L of (calcium) carbonate, or 71.485 mg/L of calcium carbonate (molar mass 100.09 g/mol). Since one degree KH = 17.848 mg/L CaCO3, this solution has a KH of 4.0052 degrees.

Carbonate hardness should not be confused with a similar measure Carbonate Alkalinity which is expressed in either [milli[equivalent]s] per litre (meq/L) or ppm. Carbonate hardness expressed in ppm does not necessarily equal carbonate alkalinity expressed in ppm.

whereas

However, for water with a pH below 8.5, the CO2−3 will be less than 1% of the HCO−3 so carbonate alkalinity will equal carbonate hardness to within an error of less than 1%.

In a solution where only CO2 affects the pH, carbonate hardness can be used to calculate the concentration of dissolved CO2 in the solution with the formula

[CO2] = 3 × KH × 107 − pH,

where KH is degrees of carbonate hardness and [CO2] is given in ppm by weight.[citation needed]

The term carbonate hardness is also sometimes used as a synonym for temporary hardness, in which case it refers to that portion of hard water that can be removed by processes such as boiling or lime softening, and then separation of water from the resulting precipitate.[1]

See also

[edit]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Carbonate hardness

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from Grokipedia
Carbonate hardness, also referred to as temporary hardness or carbonate in some contexts, is the portion of attributable to the presence of calcium (Ca²⁺) and magnesium (Mg²⁺) ions combined with (HCO₃⁻) and (CO₃²⁻) anions, typically measured and expressed in milligrams per liter (mg/L) as (CaCO₃) equivalents. This type of arises primarily from the dissolution of minerals such as () and dolomite in geological formations, where (CO₂) in or facilitates the formation of soluble bicarbonates. Unlike noncarbonate , which involves sulfates, chlorides, or nitrates of calcium and magnesium, carbonate can be reduced or eliminated by boiling , as heat decomposes bicarbonates into insoluble carbonates that precipitate out, leaving softer . In water quality assessment, carbonate hardness contributes to the total hardness of , which is the sum of carbonate and noncarbonate components, and is often equivalent to the bicarbonate fraction of in regions dominated by geology. It is measured through methods, such as acid to determine the concentration of these ions, or by direct analysis of calcium, magnesium, bicarbonate, and levels using or spectrometry, with results standardized to CaCO₃ for comparability. While total hardness levels are classified by the U.S. Geological Survey as soft (0–60 mg/L), moderately hard (61–120 mg/L), hard (121–180 mg/L), or very hard (>180 mg/L) as CaCO₃, carbonate hardness specifically influences buffering capacity, helping to stabilize by neutralizing acids in natural and treated systems. The significance of carbonate hardness extends to practical applications in , , and ; high levels can lead to scale formation in pipes, boilers, and appliances, reducing efficiency and increasing maintenance costs, while also interfering with lathering by forming insoluble precipitates. In aquatic ecosystems, it supports biological processes like shell formation in and egg calcification but helps mitigate to aquatic organisms by reducing ; in , it poses no direct risks but affects aesthetic and operational , often prompting softening treatments like or lime precipitation. Management strategies focus on monitoring and adjustment to balance hardness with other parameters, ensuring suitability for industrial, agricultural, and municipal uses.

Fundamentals

Definition

Carbonate hardness refers to the component of hardness resulting from the dissolved s (HCO₃⁻) and s (CO₃²⁻) of calcium and magnesium, specifically the salts Ca(HCO₃)₂, Mg(HCO₃)₂, CaCO₃, and MgCO₃. These ions contribute to hardness by forming complexes that interfere with lathering and other water uses, and the concentration is conventionally expressed as equivalents of (CaCO₃) for standardization in water chemistry assessments. This measure reflects the chemical equivalence to the bicarbonate and carbonate fractions of , particularly in waters that have interacted with or similar geological formations. Historically termed "temporary hardness," carbonate hardness can be reduced or eliminated through , a process that decomposes the soluble bicarbonates into insoluble carbonates and releases gas. For instance, the of calcium bicarbonate follows the reaction: \ceCa(HCO3)2−>[Δ]CaCO3↓+CO2↑+H2O\ce{Ca(HCO3)2 ->[\Delta] CaCO3 \downarrow + CO2 \uparrow + H2O} A similar reaction occurs with magnesium bicarbonate, leading to precipitation of MgCO₃. This distinguishes it from permanent, or non-carbonate, , which arises from calcium and magnesium combined with anions such as sulfates (SO₄²⁻), chlorides (Cl⁻), and nitrates (NO₃⁻)—for example, in salts like CaSO₄, MgCl₂, or Ca(NO₃)₂—that remain dissolved even after boiling. Total represents the combined contribution of both carbonate and non-carbonate components.

Relation to Total Hardness

Total (TH) in is defined as the sum of carbonate (CH) and non-carbonate (NCH), expressed in milligrams per liter (mg/L) as (CaCO₃) equivalents. This relationship arises because TH measures the total concentration of divalent cations like calcium (Ca²⁺) and magnesium (Mg²⁺), while CH specifically accounts for those associated with (HCO₃⁻) and (CO₃²⁻) anions, and NCH covers associations with other anions such as (SO₄²⁻) or (Cl⁻). The equation can be written as: TH=CH+NCHTH = CH + NCH where all terms are in mg/L as CaCO₃. Water classification is typically based on TH levels, providing a framework for assessing potential impacts on , appliances, and usage. According to the U.S. Geological Survey, water is classified as soft (0–60 mg/L TH), moderately hard (61–120 mg/L TH), hard (121–180 mg/L TH), or very hard (>180 mg/L TH). In natural waters, CH often constitutes the majority of TH, particularly in regions with or formations, due to the prevalence of bicarbonate-rich dissolution processes. A high proportion of CH relative to TH indicates water dominated by temporary hardness, which is prone to forming scale deposits—primarily —upon heating, as seen in boilers or hot water systems. However, this type of hardness is easier to mitigate compared to NCH, as it can be reduced through simple boiling (which precipitates bicarbonates) or chemical precipitation methods like , avoiding the need for full ion-exchange treatment in many cases. For instance, in , the dominance of CH frequently leads to higher overall , enhancing the water's buffering capacity against pH changes.

Measurement

Analytical Methods

The primary laboratory technique for quantifying carbonate hardness involves determining total alkalinity via acid-base titration, as carbonate hardness is typically equivalent to the bicarbonate and carbonate alkalinity in waters dominated by these ions. A water sample is titrated with a standardized acid, typically 0.02 N sulfuric acid (H₂SO₄) or hydrochloric acid (HCl), using a pH meter to reach the endpoint at pH 4.5, corresponding to the conversion of HCO₃⁻ + H⁺ → H₂CO₃. The volume of acid required is recorded, and the alkalinity (which approximates CH) is calculated in mg/L as CaCO₃ based on stoichiometry, where 1 mL of 0.02 N acid equates to 1 mg/L CaCO₃. This follows standard protocols such as APHA Method 2320. This approach differs from the method for total hardness, which uses EDTA at 10. Carbonate hardness is derived by subtracting non-carbonate hardness (NCH) from total hardness (TH), where NCH is determined as the excess of TH over total (both expressed as CaCO₃ equivalents). In practice, if TH exceeds total alkalinity, CH equals the total alkalinity value; otherwise, CH equals TH. The titration may include an initial endpoint at 8.3 (CO₃²⁻ + H⁺ → HCO₃⁻) to measure alkalinity for waters with significant content. Instrumental approaches include (IC), which separates and quantifies HCO₃⁻ and CO₃²⁻ ions directly via conductivity detection after anion exchange, enabling CH calculation by converting measured (in meq/L) to mg/L CaCO₃ equivalents (multiplying by 50). This method offers high specificity for low-level analysis in varied matrices. For field applications, colorimetric kits use acid reagents with indicators like , detecting the pH 4.5 endpoint through color change for rapid CH estimation without lab equipment. Method accuracy depends on minimizing interferences from non-carbonate contributors, such as borates, phosphates, or organic acids, which can inflate volumes. Sample preservation is essential, involving collection in CO₂-impermeable bottles filled to the brim, cooling to , and analysis within 24 hours to prevent atmospheric CO₂ absorption or loss that alters ionic equilibria.

Units and Standards

Carbonate hardness is most commonly expressed in milligrams per liter (mg/L) as (CaCO₃) equivalent, a standardized measure that normalizes the concentrations of (HCO₃⁻) and (CO₃²⁻) ions contributing to temporary . This unit reflects the equivalent amount of that would produce the same , where 1 mg/L corresponds to 0.02 milliequivalents per liter (meq/L) of , facilitating comparisons across samples. Alternative units for carbonate hardness include regional degree-based systems, which originated in for practical assessments. The German degree (°dH) defines 1 °dH as equivalent to 17.8 mg/L CaCO₃, while the French degree (°fH) sets 1 °fH at 10 mg/L CaCO₃; the English degree (°e, also known as Clark degrees) equates to approximately 14.3 mg/L CaCO₃ or parts per million (ppm). These units allow for straightforward conversions, such as 1 °dH ≈ 1.78 °fH or 1.25 °e, and are still used in specific industries like and despite the global preference for mg/L CaCO₃. To calculate the hardness contribution from measured calcium and magnesium concentrations (for total hardness; carbonate hardness is then the portion balanced by carbonate/bicarbonate anions, typically min(TH, alkalinity as CaCO₃)), the following formula is applied: Hardness (mg/L as CaCO₃)=[Ca2+ (mg/L)20+Mg2+ (mg/L)12.15]×50\text{Hardness (mg/L as CaCO₃)} = \left[ \frac{\text{Ca}^{2+} \text{ (mg/L)}}{20} + \frac{\text{Mg}^{2+} \text{ (mg/L)}}{12.15} \right] \times 50 This expression accounts for the molar contributions of divalent cations, converting them to CaCO₃ equivalents based on equivalent weights (Ca²⁺: 20 g/eq, Mg²⁺: 12.15 g/eq, CaCO₃: 50 g/eq). Regulatory and industry standards for hardness primarily address total hardness (TH), with limited specific limits for carbonate hardness (CH) due to its role in buffering rather than direct health impacts. The (WHO) guidelines recommend TH below 500 mg/L as CaCO₃ to avoid aesthetic issues like scaling, but provide no dedicated CH limit, emphasizing that values up to this threshold pose no health risks. While the Drinking Water Directive (98/83/EC, revised 2020/2184) lacks a mandatory TH cap or specific CH limit, many member states apply non-binding national guidelines of around 250 mg/L CaCO₃ for TH to mitigate infrastructure effects. In aquarium applications, standards recommend CH levels of 3–6 °dH (approximately 53–107 mg/L CaCO₃) for most freshwater setups to maintain stability and support invertebrate health.

Chemical Properties

Ionic Composition

Carbonate hardness in aqueous solutions arises primarily from the presence of (HCO₃⁻) and (CO₃²⁻) anions paired with divalent cations, forming the key ionic contributors to temporary hardness. The carbonate system in water is governed by the following stepwise equilibria: CO2+H2O⇌H2CO3⇌H++HCO3−⇌2H++CO32−\mathrm{CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons H^+ + HCO_3^- \rightleftharpoons 2H^+ + CO_3^{2-}} These reactions establish the speciation of , with dissociation constants pK₁ ≈ 6.3 and pK₂ ≈ 10.3 at 25°C. In typical natural waters with ranging from 6 to 9, the dominant species is the bicarbonate ion (HCO₃⁻), which constitutes the majority of the and pairs with cations to form soluble bicarbonates. At higher values above approximately 10.3, the ion (CO₃²⁻) becomes predominant, shifting the equilibrium toward more basic conditions and potentially leading to if cation concentrations exceed limits. This pH-dependent directly influences , as higher correlates with increased concentrations of these anions, enhancing the overall level. The primary cations associated with carbonate hardness are calcium (Ca²⁺) and magnesium (Mg²⁺), forming species such as Ca(HCO₃)₂ and Mg(HCO₃)₂ in acidic to neutral conditions, or and at higher . These pairings are limited by the low of the corresponding solids; for instance, the solubility product constant (K_{sp}) for () is 3.36 × 10^{-9} at 25°C, restricting the product [Ca²⁺][CO₃²⁻] in saturated solutions. Magnesium carbonate has higher (K_{sp} ≈ 6.8 × 10^{-6} at 25°C), allowing greater contribution from Mg in carbonate-rich environments, though Ca typically dominates due to geological abundance. In natural aquatic systems, carbonate hardness originates largely from the dissolution of (CaCO₃) in the presence of , as described by the reaction: CaCO3+CO2+H2O→Ca2++2HCO3−\mathrm{CaCO_3 + CO_2 + H_2O \rightarrow Ca^{2+} + 2HCO_3^-} This , common in regions, increases both Ca²⁺ and HCO₃⁻ concentrations, elevating carbonate hardness while maintaining charge balance. Speciation of the system, and thus carbonate hardness, can be predicted from total (DIC) using equilibrium equations that account for and . DIC is defined as the sum [CO₂(aq)] + [HCO₃⁻] + [CO₃²⁻], and carbonate hardness is approximated by twice the carbonate alkalinity (primarily from HCO₃⁻ and CO₃²⁻) multiplied by the equivalent concentrations of Ca²⁺ and Mg²⁺, assuming minimal other contributors. Such calculations often employ software like PHREEQC for precise modeling in complex waters.

Thermal Decomposition

Carbonate hardness, primarily due to dissolved calcium and magnesium bicarbonates, undergoes thermal decomposition when water is heated, converting the soluble bicarbonates into insoluble carbonates and releasing carbon dioxide gas. This process is responsible for the "temporary" nature of carbonate hardness, as it can be largely eliminated through boiling. The key reaction for calcium bicarbonate is: \ceCa(HCO3)2(aq)−>[heat]CaCO3(s)v+CO2(g)+H2O(l)\ce{Ca(HCO3)2 (aq) ->[heat] CaCO3 (s) v + CO2 (g) ^ + H2O (l)} A similar decomposition occurs for magnesium bicarbonate: \ceMg(HCO3)2(aq)−>[heat]MgCO3(s)v+CO2(g)+H2O(l)\ce{Mg(HCO3)2 (aq) ->[heat] MgCO3 (s) v + CO2 (g) ^ + H2O (l)} These reactions lead to the precipitation of scale, predominantly , in heated systems. Significant decomposition begins above 60°C, with the process becoming complete at the boiling point of 100°C, as the solubility of the resulting carbonates decreases with rising temperature. The reaction is driven by the reduction in partial pressure of CO₂ as the gas escapes from the boiling water, shifting the bicarbonate equilibrium toward carbonate formation according to Le Chatelier's principle. This shift favors precipitation, particularly for calcium, since calcium carbonate has very low solubility (approximately 0.013 g/L at 25°C, decreasing further with heat). The primary residual effect is the formation of , consisting mainly of CaCO₃ deposits, which accumulate in boilers, kettles, and , insulating heat transfer surfaces and reducing operational efficiency by up to 20-30% in severe cases. Magnesium ions are less prone to because magnesium carbonate is more soluble (approximately 0.22 g/L at ) and often remains partially dissolved or forms less adherent scales. Overall, can remove up to 90% of carbonate hardness, depending on the initial concentration and boiling duration, though complete removal may require or of the precipitate.

Impacts and Applications

Effects on Infrastructure

Carbonate hardness, primarily from (CaCO₃), leads to scale formation in systems, heaters, and household appliances such as kettles. When with elevated carbonate levels is heated, dissolved minerals precipitate as hard deposits on interior surfaces, narrowing pipe diameters and restricting flow. This buildup can reduce flow rates substantially over time, with even thin layers (e.g., 1/8 inch) causing significant impediments in hot water lines. Additionally, scale acts as an insulator on heating elements, increasing ; for instance, a 1/8-inch layer of CaCO₃ scale can raise use by 25-30% in heaters due to decreased efficiency. While high carbonate hardness (>200 mg/L as CaCO₃) predominantly causes scaling issues, moderate levels (50-150 mg/L) can offer corrosion inhibition by forming a thin protective carbonate film on metal surfaces like pipes. This film, composed mainly of , acts as a barrier against further oxidation and pitting, reducing corrosion rates in distribution systems. However, at higher concentrations, the protective benefit is outweighed by excessive scaling, which can instead promote localized corrosion under deposits or accelerate wear in valves and fittings. The economic repercussions of carbonate hardness-induced scaling are substantial through maintenance, reduced equipment efficiency, and premature replacements. In the United States, hard water scaling shortens the lifespan of residential water heaters from an expected 10-15 years to as little as 5-7 years, increasing replacement frequency and energy bills. These impacts extend to commercial infrastructure, where scale-related downtime and cleaning add to operational expenses across industries reliant on hot water systems. In regions dependent on limestone aquifers, where groundwater often exceeds 100 mg/L in carbonate hardness due to natural dissolution of carbonate rocks, scaling necessitates frequent descaling of appliances like dishwashers and coffee machines. For example, in areas such as parts of the in , high CaCO₃ levels result in visible deposits that clog spray arms and heating coils, requiring manual or chemical cleaning every few months to maintain performance. This ongoing maintenance highlights the infrastructure challenges in karst terrains with elevated hardness.

Role in Water Buffering and Ecosystems

Carbonate hardness (CH) is approximately equivalent to total alkalinity (TA) in most natural waters, as both are primarily derived from the dissolution of carbonate minerals like , providing similar concentrations when expressed as mg/L CaCO₃. This equivalence arises because the (HCO₃⁻) and (CO₃²⁻) ions from CH dominate TA in typical freshwater and slightly alkaline systems. The HCO₃⁻/CO₃²⁻ buffering system inherent to CH resists changes from CO₂ fluctuations by absorbing or releasing protons, thereby stabilizing pH in the range of 7 to 8.5, which is common in buffered aquatic environments. The buffering capacity of this system can be approximated by the buffer intensity β: β≈2.303([\ceHCO3−]+2[\ceCO32−]+[\ceOH−]−[\ceH+])\beta \approx 2.303 \left( [\ce{HCO3-}] + 2[\ce{CO3^2-}] + [\ce{OH-}] - [\ce{H+}] \right) This expression highlights how higher concentrations of HCO₃⁻ and CO₃²⁻ enhance resistance to acidification, with the factor of 2 for CO₃²⁻ reflecting its contribution to proton acceptance and the 2.303 arising from the base-10 logarithm in definition. In practice, this capacity prevents rapid drops from dissolved CO₂, which forms and could otherwise lower below stable levels. In aquatic ecosystems and controlled settings like aquariums, an optimal CH of 50-150 mg/L as CaCO₃ maintains stability against pH swings that are lethal to , such as those caused by CO₂ buildup from respiration or , which can drop pH below 6 and induce stress or mortality. This range also supports essential biological processes, including calcification, where sufficient ions facilitate the formation of shells by providing the necessary ionic environment for in species like oysters and clams. Environmentally, waters with low CH, such as rainwater (typically <5 mg/L), lack this buffering and acidify quickly upon exposure to atmospheric CO₂ or pollutants, resulting in pH values around 5.6. In contrast, river waters with CH exceeding 20 mg/L exhibit greater resistance to acidification from anthropogenic pollution, such as acid mine drainage or atmospheric deposition, thereby protecting aquatic biodiversity. This threshold ensures minimal disruption to pH-dependent ecological processes in flowing systems.

Treatment Methods

Boiling and Precipitation

One common low-tech method for reducing carbonate hardness in household settings is domestic , where water is heated to 100°C for 10-30 minutes to drive off dissolved , promoting the of (CaCO₃) from ions via the reaction Ca(HCO₃)₂ → CaCO₃ + H₂O + CO₂. This process can achieve a significant reduction in carbonate hardness, with studies showing up to 40% removal after brief , though longer durations and settling allow for higher rates, making it effective for small-scale applications such as tea kettles to prevent scale buildup. However, it primarily targets temporary (carbonate) hardness and leaves non-carbonate hardness unaffected. For larger-scale treatment, employs chemical to remove carbonate by adding hydrated lime (Ca(OH)₂) to , raising the above 10 to facilitate the formation of insoluble (CaCO₃) and (Mg(OH)₂) precipitates. The cold lime process operates at ambient temperatures and reduces calcium to 35-50 ppm as CaCO₃, while the hot lime process, conducted at 108-116°C, achieves greater efficiency by lowering total to approximately 8 ppm as CaCO₃ and magnesium to 2-5 ppm. After and , recarbonation with CO₂ stabilizes the treated by lowering the to 8.0-9.0 and converting excess carbonates back to bicarbonates, preventing further scaling. Like , this method selectively removes carbonate but does not address non-carbonate components. Despite their effectiveness, both and have notable limitations. is energy-intensive for volumes beyond household use and requires for complete precipitate removal. generates substantial sludge—up to 2 pounds per pound of lime added—which necessitates disposal or recycling in large systems, adding operational costs. Additionally, the cold lime process is incomplete for waters with high magnesium content, limiting magnesium reduction to around 70 ppm as CaCO₃, whereas the hot process mitigates this but demands even more energy for heating.

Ion Exchange and Filtration

Ion exchange is a widely used method for reducing carbonate hardness by targeting the calcium (Ca²⁺) and magnesium (Mg²⁺) ions that contribute to it, as these cations form bicarbonates responsible for temporary hardness. In this process, water passes through a bed of cation-exchange resin, typically in the sodium (Na⁺) form, where the resin selectively exchanges its Na⁺ ions for the hardness-causing Ca²⁺ and Mg²⁺ ions due to the resin's higher affinity for divalent cations. This ion swap can remove nearly all hardness (typically over 99%), effectively lowering carbonate hardness to near zero in treated water. The resin becomes saturated over time and requires periodic regeneration using a concentrated brine solution (sodium chloride), which displaces the captured Ca²⁺ and Mg²⁺ ions, restoring the resin to its Na⁺ form. This technique is commonly employed in household water softeners, providing consistent softening for domestic use, though it increases sodium content in the output water, which may be a concern for individuals on low-sodium diets. Reverse osmosis (RO) represents another engineered approach to carbonate hardness removal, utilizing semi-permeable membranes to reject divalent ions like Ca²⁺ and Mg²⁺ under high pressure. The process forces water through the membrane, which allows water molecules to pass while blocking over 98% of dissolved salts, including those contributing to hardness, resulting in permeate water with significantly reduced carbonate hardness. RO systems typically achieve recovery rates of 50-80%, producing a concentrated brine stream that constitutes 20-50% of the feed volume, necessitating proper disposal to avoid environmental impacts. This method is scalable for both residential and industrial applications, offering comprehensive purification beyond just hardness removal, such as the elimination of other contaminants. Electrodeionization (EDI), particularly for production, employs an across ion-exchange resins and membranes to continuously remove ionized , including (HCO₃⁻) ions that pair with Ca²⁺ and Mg²⁺ to form carbonate hardness. In EDI modules, bipolar membranes facilitate the and removal of HCO₃⁻ by splitting into H⁺ and OH⁻ ions, which then regenerate the resins in without chemical additives, achieving demineralization efficiencies exceeding 99% for target ions. This chemical-free process is often integrated downstream of RO for polishing, making it suitable for high-purity needs in pharmaceuticals and electronics manufacturing. In municipal water treatment, ion exchange and filtration methods like RO and EDI are applied to achieve residual hardness levels below 60 mg/L as CaCO₃, corresponding to soft water classifications that minimize scaling in distribution systems and appliances. These techniques provide permanent softening by removing hardness ions outright, unlike temporary methods, but drawbacks include the addition of sodium via ion exchange and the generation of brine waste from RO and EDI, which require management strategies. Emerging approaches, such as magnetic water treatment and adsorption using bentonite, offer potential low-waste alternatives as of 2025. Overall, they enable efficient, large-scale treatment while maintaining water quality standards.

References

  1. https://edis.ifas.ufl.edu/publication/SS540
  2. https://pubs.usgs.gov/wri/wri024094/pdf/mainbodyofreport-3.pdf
  3. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_%28Inorganic_Chemistry%29/Descriptive_Chemistry/Main_Group_Reactions/Hard_Water
  4. https://www.usgs.gov/faqs/do-you-have-information-about-water-hardness-united-states
  5. https://pubs.usgs.gov/of/1982/0119/plate-1.pdf
  6. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Main_Group_Reactions/Hard_Water
  7. https://water.mecc.edu/courses/Env211/lesson12_print.htm
  8. https://water.mecc.edu/courses/ENV148/lesson20b.htm
  9. https://www.sciencedirect.com/topics/engineering/carbonate-hardness
  10. https://kgs.uky.edu/kgsweb/datasearching/water/analyteDescr.asp
  11. https://documents.thermofisher.com/TFS-Assets/LPD/Product-Information/Alkalinity-ST-ALKABASICS-EN.pdf
  12. https://www.hach.com/asset-get.download-en.jsa?id=7639983916
  13. https://www.standardmethods.org/
  14. https://www.umass.edu/mwwp/pdf/epa310_1singleendpointalk.pdf
  15. https://www.hach.com/asset-get.download-en.jsa?id=60542000201
  16. https://www.waterboards.ca.gov/water_issues/programs/swamp/docs/cwt/guidance/3211.pdf
  17. https://www.lenntech.com/ro/water-hardness.htm
  18. https://aqua-chem.com/water-chemistry-caco3-equivalents/
  19. https://www.aqion.de/site/water-hardness
  20. https://www.lenntech.com/calculators/hardness/hardness.htm
  21. https://floridadep.gov/sites/default/files/dw-abcd-formula_1.pdf
  22. https://cdn.who.int/media/docs/default-source/wash-documents/wash-chemicals/hardness-bd.pdf?sfvrsn=a13853a9_4
  23. https://fns.uniba.sk/fileadmin/prif/envi/kge/LifeWater/PDF/Regulation_for_Ca_and_Mg__EU.pdf
  24. https://www.aquariumcoop.com/blogs/aquarium/ph-gh-kh
  25. https://www.pmel.noaa.gov/co2/files/dickson_thecarbondioxidesysteminseawater_equilibriumchemistryandmeasurementspp17-40.pdf
  26. https://chem.libretexts.org/Bookshelves/General_Chemistry/Map:_General_Chemistry_(Petrucci_et_al.)/18:_Solubility_and_Complex-Ion_Equilibria/18.01:_Solubility_Product_Constant_Ksp
  27. https://www.engineeringtoolbox.com/solubility-product-equilibrium-constant-ionic-solution-salt-Ksp-d_1952.html
  28. https://www.margaretriverattractions.com/wp-content/uploads/2019/05/CAVE-CHEMISTRY-FOR-SENIOR-STUDENTS.pdf
  29. https://web.njit.edu/~kebbekus/week_2_class_notes.htm
  30. https://www.elac.edu/sites/elac.edu/files/2022-08/Chem_101_Lab_Manual.pdf
  31. https://melscience.com/US-en/chemistry/experiments/water-hardness/
  32. https://geoweb.uwyo.edu/geol2100/CarbonateChemistry.pdf
  33. https://water.mecc.edu/courses/ENV115/lesson1_3b.htm
  34. https://www.kgs.ku.edu/General/Geology/Sherman/05_gw4.html
  35. https://water.usgs.gov/owq/hardness-alkalinity.html
  36. https://www.pnnl.gov/main/publications/external/technical_reports/pnnl-22921.pdf
  37. https://nepis.epa.gov/Exe/ZyPURL.cgi?Dockey=9100CILD.TXT
  38. https://nepis.epa.gov/Exe/ZyPURL.cgi?Dockey=10003FIW.TXT
  39. https://wqa.org/learn-about-water/perceptible-issues/scale-deposits/
  40. https://pubs.usgs.gov/sir/2018/5060/sir20185060.pdf
  41. https://doi.org/10.3390/w13010042
  42. https://extension.msstate.edu/publications/total-alkalinity
  43. https://www.waterontheweb.org/under/streamecology/09_pH-draft.html
  44. https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Acids_and_Bases/Buffers/Buffer_Capacity
  45. https://www.fdacs.gov/Consumer-Resources/Recreation-and-Leisure/Aquarium-Fish/Aquarium-Water-Quality-Total-Alkalinity-and-Hardness
  46. https://shellfish.ifas.ufl.edu/wp-content/uploads/pH-and-Alkalinity-Fact-Sheet-for-Hatcheries-Final-Draft.pdf
  47. https://www.fondriest.com/environmental-measurements/parameters/water-quality/ph/
  48. https://www.intechopen.com/chapters/69568
  49. https://aslopubs.onlinelibrary.wiley.com/doi/10.1002/lol2.70047
  50. https://www.watertechnologies.com/handbook/chapter-07-precipitation-softening
  51. https://www.mdpi.com/2073-4441/17/9/1330
  52. https://water.mecc.edu/courses/ENV115/lesson14b.htm
  53. https://extension.psu.edu/water-softening/
  54. https://www.epa.gov/sdwa/overview-drinking-water-treatment-technologies
  55. https://www.dupont.com/content/dam/water/amer/us/en/water/public/documents/en/RO-NF-FilmTec-Scale-Control-Softening-Manual-Exc-45-D01548-en.pdf
  56. https://extensionpubs.unl.edu/publication/g1491/2014/html/view
  57. https://www.sciencedirect.com/science/article/pii/S1944398624203921
  58. https://axeonsupply.com/blogs/news/reverse-osmosis-vs-deionization
  59. https://www.dupont.com/water/technologies/electrodeionization-edi.html
  60. https://www.watertechnologies.com/products/e-cell-electrodeionization-edi-solutions
  61. https://www.researchgate.net/publication/394546280_Magnetic_water_treatment_for_calcium_carbonate_scale_prevention_a_review
  62. https://iwaponline.com/wpt/article/20/5/1071/108095/New-insights-into-the-management-of-hard-water
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