**3.5 Larson-Skold index (LI)**

In contrast to the previously descried indices, the Larson–Skold Index (LI) describes the corrosivity of water towards iron or mild steel due to the presence of chloride and sulfate ions [18]. No consideration is given to calcium concentration or pH. This index looks at the relative ratio of chloride and sulfate ions to alkalinity in the water. The reactive anions on one-hand have a strong acidic effect in the anodic pits generated in the exposed corroding metal. The alkalinity due to combined bicarbonate and carbonate ions counter this effect by creating a stabilized and buffered environment that reduces the acidic tendency of the water. The Larson-Skold Index is calculated as follows:

$$\text{LI} = \frac{[\text{Cl}^-] + [\text{SO}\_4^{2-}]}{[\text{HCO}\_3^-] + [\text{CO}\_3^{2-}]}$$

where the concentrations of each of the species involved is expressed in milliequivalents per liter (meq/L).

The water is then classifies as:


## **4. Processes used for remineralization of desalinated water**

The most common techniques that are currently employed worldwide for water remineralization and stabilization of desalinated or naturally soft water can be divided into three categories: (1) direct dosing of two or more chemical solutions, (2) lime dosing systems, and (3) calcite contactors. Each technique comes with its own benefits and disadvantages, from both a water quality perspective as well as process considerations. Of these three processes, lime dosing systems and calcite contactors see the most widespread application, especially for large desalination plants. This is due to the major drawbacks and costs associated with chemical dosing, as demonstrated below.

#### **4.1 Direct dosing of two or more chemical solutions**

One of the simplest and most effective ways of controlling the quality of the final water leaving a treatment facility is through the direct dosing of chemical solutions, either prepared offsite, or involving a simple slurry make down on site when supplied as a solid material. Whilst any combination of chemicals is possible, the most common combination is calcium chloride (CaCl2) for hardness addition in conjunction with sodium bi-carbonate (NaHCO3) to increase the alkalinity of the final water. Expensive sodium hydroxide may in some cases also be needed to reach the required pH. Direct dosing of two or more chemicals has the advantage over other techniques of being able to precisely regulate the quantities of ions added, by controlling the volumes dosed of known solutions. In addition, full dissociation of these highly pure compounds in water avoids any complications of the generation of waste by-products or residual turbidity resulting from low solubility of products.

Despite being a very simple process the greatest drawback of this process is the cost of the chemicals. For this reason, this approach sees very limited application, used at best on small plants where relatively speaking CAPEX has a much greater impact than OPEX.

As an example, in order to achieve 80 mg/l of hardness and alkalinity within the final water (measured as CaCO3), then the required dosage of CaCl2 and NaHCO3 is equivalent to:

$$\text{CaCl}\_2\left(\frac{\text{mg}}{l}\right) = 80\left(\frac{\text{mg}}{l}\right) \div 100.08\left(\frac{\text{g}}{mol}\right) \times 110.98\left(\frac{\text{g}}{mol}\right) = 89\left(\frac{\text{mg}}{l}\right)$$

$$\text{NaHCO}\_3\left(\frac{\text{mg}}{l}\right) = 80\left(\frac{\text{mg}}{l}\right) \div 100.08\left(\frac{\text{g}}{mol}\right) \times 84.006\left(\frac{\text{g}}{mol}\right) \times 2 = 134\left(\frac{\text{mg}}{l}\right)$$

As two molecular equivalents of NaHCO3 are required to generate one molecular equivalent of alkalinity (measured as CaCO3).

Using the modest values of chemical costs of 300 €/tonne for calcium chloride and 400 €/tonne for sodium bicarbonate (in some locations sodium bicarbonate can be as expensive as 900–950 US\$/tonne), then the cost of consumables of this process alone exceeds the total treatment cost of the other treatment processes when energy costs and amortization of investment costs are taken into account (refer **Table 1** below).

Another major drawback of this process and something that is often overlooked, is the increase of the undesired chloride ions that are introduced to the process, as for every mol of calcium that is added, two mol of chloride is also added. This increases the tendency of the water to be corrosive to iron and steel pipe and equipment, as measured and indicated by the previously mentioned Larson-Skold Index.

#### **4.2 Lime dosing systems**

Lime dosing systems involve the on-site preparation of a saturated lime solution from either a hydrated lime powder (calcium hydroxide) or quicklime (calcium oxide) which is slaked on-site to produce hydrated lime. The saturated solution is produced by feeding a calcium hydroxide slurry into a lime saturator along with make-up water and a flocculant. The saturator allows for the separation of a clear


**Table 1.**

*Chemical costs of calcium chloride and sodium bicarbonate dosing to achieve 80 mg/l of hardness and alkalinity.*

*Remineralization and Stabilization of Desalinated Water DOI: http://dx.doi.org/10.5772/intechopen.99458*

#### **Figure 2.**

solution of calcium hydroxide from the insoluble contents which are settled to the bottom with the aid of the flocculant (**Figure 2**).

The saturated lime solution is then dosed into the final water stream along with carbon dioxide to form bicarbonate alkalinity in the following reaction:

$$\text{Ca(OH)}\_{2\text{ (aq)}} + 2\text{ CO}\_{2\text{ (g)}} \rightarrow \text{Ca(HCO}\_{3}\text{)}\_{2\text{ (aq)}}\tag{2}$$

One of the advantages of lime and reasons that it has been a popular choice for system designers is its worldwide availability as a commercial product and relatively low cost in comparison to other chemicals. It's solubility up to 1700 mg/l at 20°C makes ideal for its preparation within a side stream lending to its smaller footprint relative to calcite contactors.

Lime dosing systems do however have a number of drawbacks and are often the bane of many operators assigned the task to clean and maintain the lime slurry pipework or lime saturators. In comparison to calcium carbonate, lime is more expensive per kilogram of available CaCO3. This can be attributed to the fact that lime is produced by the calcination (burning) and further slaking (hydration) of calcium carbonate which is then dried to produce a powdered hydrated lime. As a result, lime is not only more expensive to produce, but has a much higher carbon footprint. For every kilogram of quicklime that is produced, approximately 700 kcal is required for dissociation and 0.785 kg of carbon dioxide is released [19]. This difference in carbon footprint is not only a concern for plants which strive for environmentally sustainable solutions, but will no doubt further increase the price of lime production as carbon emission taxes are set to play a bigger role in the future.

The operational costs of lime systems are also increased in comparison to calcium carbonate due to the fact that for the same desired quantity of calcium hardness and alkalinity within the final water, twice as much carbon dioxide is required. As can be seen from the chemical reaction equations, remineralization using calcium hydroxide requires two molecular equivalents of carbon dioxide for each mole of calcium hydroxide (refer Eq. (2)). Remineralization using calcium carbonate however, requires only one molecular equivalent of carbon dioxide for

each mole of calcium carbonate (refer Eq. (3)). Carbon dioxide is in most cases the largest operating cost for post treatment systems, so this has an important impact on the overall cost per cubic meter of treated water.

As mentioned previously, hydrated lime contains an insoluble content. This insoluble content is for the most part is unburnt calcium carbonate but can also include silicates and other impurities. These impurities are usually in the order of 5–15% and must be removed and dealt with as a waste product. Needless to say the lower the insoluble content, the purer the product and the higher the price of the product. If the impurities are not effectively removed, they will add to the turbidity in the final water. To improve the efficacy of the clarification process, a flocculant is often dosed to aid in the settling. This waste then needs to be thickened on site and sent away for proper disposal. These factors further add to the operational cost and complexity of lime dosing systems. The operation of a lime clarifier is very sensitive to factors such as temperature, flocculent dosing, and throughput flow rate, meaning they do not handle fluctuations in plant flow very well. Even a perfectly functioning lime clarifier can still be susceptible to turbidity problems. Absorption of carbon dioxide from the atmosphere can lead to the increase in dissolved inorganic carbon within the lime solution resulting in precipitation of calcium carbonate and producing a cloudy solution.

Extending the example for chemical dosing, in order to achieve 80 mg/l of hardness and alkalinity within the final water (measured as CaCO3), then the required dosage of Ca(OH)2 and CO2 is equivalent to:

$$\text{Ca}(\text{OH})\_2\text{(}\frac{\text{mg}}{l}\text{)} = \text{80}\left(\frac{\text{mg}}{l}\right) \div 100.08 \left(\frac{\text{g}}{mol}\right) \times 74.1 \left(\frac{\text{g}}{mol}\right) \div 90\text{\%} = \text{65}\left(\frac{\text{mg}}{l}\right).$$

(assuming a Ca(OH)2 purity of 90%)

$$\text{CO}\_2\left(\frac{\text{mg}}{l}\right) = \text{80}\left(\frac{\text{mg}}{l}\right) \div 100.08\left(\frac{\text{g}}{mol}\right) \times 44.01\left(\frac{\text{g}}{mol}\right) \times 2 = 70\left(\frac{\text{mg}}{l}\right)$$

As two molecular equivalents of CO2 are required to generate one molecular equivalent of alkalinity (measured as CaCO3). Considering also 10% of product is removed as waste from the lime saturator, and then dewatered to a maximum of 30% solids:

$$\text{Waste generated } \left(\frac{\text{mg}}{l}\right) = 65 \left(\frac{\text{mg}}{l}\right) \times 10\text{\%} \div 30\text{\%} = 22 \left(\frac{\text{mg}}{l}\right)^2$$

This generates treatment costs that are a fraction of that required for chemical dosing as presented below in **Table 2** below.

#### **4.3 Calcite contactors**

Remineralization of demineralized water using calcite contactors is achieved by passing a stream of acidified water through a bed of calcite chips. These chips are


**Table 2.**

*Approximate operational costs for lime dosing systems to achieve 80 mg/l of hardness and alkalinity.*

*Remineralization and Stabilization of Desalinated Water DOI: http://dx.doi.org/10.5772/intechopen.99458*

usually limestone or marble, but in some cases dolomite, chalk, precipitated calcium carbonate or even muscle shells is used. The acidified water dissolves the calcium carbonate, increasing the calcium hardness of the water, the carbonate alkalinity and the pH through the following reaction (when carbon dioxide is used as the acidifying agent) (**Figure 3**):

$$\text{CaCO}\_{3\text{ (s)}} + \text{CO}\_{2\text{ (g)}} + \text{H}\_2\text{O} \to \text{Ca(HCO}\_3\text{)}\_{2\text{ (aq)}}\tag{3}$$

As noted earlier, calcium carbonate is one of the most common minerals available, taking up almost 4% of the earth's crust [10]. As a result it is a readily sourced product and processing requirements are minimal as the chemical composition does not need to be altered before use. Consequently, and as alluded to previously, calcium carbonate can be supplied at a lower cost than lime. In addition, it requires half as much carbon dioxide for the same quantity of calcium bi-carbonate (refer Eqs. (2) and (3)). Added to this, calcium carbonate used for water remineralization can be very pure, with an insoluble content as little as 0.1%. This further reduces operating costs, with more available product for use and less produced waste which requires further handling and disposal. Calcium carbonate is also chemically stable and non-corrosive making it easy for manual handling. In comparison, calcium hydroxide is classified as hazardous, and exposure can cause burning and irritation.

The main disadvantage of calcium carbonate is its chemical solubility which is only 13 mg/l in pure water, and requires the addition of an acid to dissolve quantities above this. As the water gets closer to the saturation point, the rate of reaction slows down dramatically. As a result, thermodynamic equilibrium is almost impossible to reach in such a dissolution reactor [20], and requires excessive contact time between the water and the calcite bed. This increases the overall size of the plant required to treat a certain volume of water and the resultant capital investment.

**Figure 3.** *Schematic of typical calcite contactor process.*

In order to increase the rate of reaction and circumvent this problem, the pH of the feed water is often decreased before the calcite contactor, in excess of what would normally be required. This results in a faster dissolution rate so that the required hardness and alkalinity increase occurs in a shorter time period. The water however, does not achieve saturation levels with respect to calcium carbonate, as the pH of the effluent leaving the reactor is much lower than the saturation pH for its calcium carbonate content. This also decreases the efficiency of the process described by Eq. (3), as it requires excess carbon dioxide to be dosed for the increased acidity. This not only increases the chemical costs for the process, but produces effluent leaving the reactor with a quantity of unreacted CO2. The pH of the water must then be adjusted either by the dosage of a strong base (e.g. sodium hydroxide) or the liberation of carbon dioxide to achieve a zero or slight positive LSI (Langelier Saturation Value) value, as required by most treatment facilities.

This pH adjustment step adds to both the investment costs: due to the requirement of additional infrastructure for stripping equipment or an additional chemical dosing system, and the operational costs: due either to additional chemical consumption when a strong base is used, or additional power consumption when the excess CO2 is stripped. In some instances both are required. Sodium hydroxide is a relatively expensive chemical, and even a small dosage can add significantly to the overall cost of the remineralization process. In many cases, it is most cost effective to waste the excess CO2 through stripping rather than convert it to additional alkalinity through the dosage of a strong base. For plants that require a low level of remineralization though, (e.g. 50 mg/l) the quantity of free of CO2 that remains in the water at saturation level is so low that this cannot be achieved through stripping alone and requires some dosage of a strong base for partial or total pH adjustment.

Other issues facing operators of calcite contactors are turbidity spikes in the treated water leaving the contactor. These are generally a result of the introduction of new material to the calcite reactor and the accompanying "fines" supplied with the raw material. To deal with this calcite contactors require frequent backwashes, in particular following the loading of new material. This adds an additional need for a backwash treatment and handling system at plant to deal with this waste stream. The change in bed heights between fills also result in a change in water quality leaving the calcite contactor due to variances in the quantity of product available for reaction. In most cases calcite contactors are loaded manually, increasing the operational costs due to additional operator utilization. Lime systems on the other hand are fed automatically from a lime storage silo, requiring operator input only to receive new deliveries. Calcite contactors also require regular backwash with both air and water. This serves to remove excess fines and insoluble waste material caught on the calcite chips as well as resettling the calcite bed to prevent preferential flow paths of the water through the bed.

Completing the cost comparison example to achieve 80 mg/l of hardness and alkalinity within the final water (measured as CaCO3), then the required dosage of CaCO3 and CO2 is equivalent to:

$$\text{CaCO}\_3\left(\frac{\text{mg}}{l}\right) = 80\left(\frac{\text{mg}}{l}\right) \div 99\% = 65\left(\frac{\text{mg}}{l}\right)$$

Assuming a CaCO3 purity of 99%,

$$1\text{ CO}\_2\left(\frac{\text{mg}}{l}\right) = 80\left(\frac{\text{mg}}{l}\right) \div 100.08\left(\frac{\text{g}}{mol}\right) \times 44.01\left(\frac{\text{g}}{mol}\right) \times 130\text{\%} = 46\left(\frac{\text{mg}}{l}\right)$$

Assuming 30% extra carbon dioxide is required to increase the rate of reaction, and finally an addition of 2.5 ppm of sodium hydroxide is required after degassing to bring the pH to saturation conditions.


*Remineralization and Stabilization of Desalinated Water DOI: http://dx.doi.org/10.5772/intechopen.99458*

#### **Table 3.**

*Approximate operational costs for calcite contactors to achieve 80 mg/l of hardness and alkalinity.*

The approximate treatment costs for calcite contactors are shown below in **Table 3** demonstrating that they are not only less expensive to operate than lime dosing systems, but they offer a more environmentally friendly solution. The big drawback for calcite contactors, which is not demonstrated here are the large physical footprint and investment required. These elements alone can often drive a designers decision towards lime dosing system, particularly in locations such as Singapore, where space is a premium.

### **5. Latest developments and trends**

Historically, the main focus for developments within the field of desalination has been the improvement and optimization of the reverse osmosis and pre-treatment processes. This is due to the fact that these areas comprise the largest fractions of capital investment and are responsible for the largest portion of operating costs. In comparison, only minor inroads have been made to improve and optimize post treatment processes. In spite of this, the value and importance of post treatment processes should not be underestimated, as it is these processes that are ultimately responsible for the final water quality sent to the consumer.

Whilst calcite contactors have many advantages over lime dosing systems, their major drawbacks are centered around their slow reactivity which result in large physical footprint and investment for the dissolution reactors. In order to address the issue of slow dissolution kinetics of calcite chips, new and innovative processes have been developed over the last few years that utilize micronized calcium carbonate. These processes take advantage of the increased surface area and reaction kinetics available from the micronized products to achieve decreased contact times, higher concentrations, improved carbon dioxide efficiency, or a combination of these factors. Micronized calcium carbonate is dissolved in a Membrane Calcite Reactor (MCR) which combines a submerged ultrafiltration membrane immersed in a suspension of micronized calcium carbonate. Carbon dioxide is added to the calcium carbonate suspension, which in turns reacts to form a calcium bi-carbonate solution. The membrane acts as a barrier between the dissolved and undissolved calcium carbonate enabling a perfectly clear solution to be extracted from the reactor that can be dosed into the desalination permeate. The use of micronized calcium carbonate results in fast reaction times, and hence decrease footprint and investment offering an improvement over current processes (**Figure 4**).

Although remineralization processes primarily concern themselves with the replenishment of calcium hardness and alkalinity, more recently attention has been given to the need to replenish magnesium ions, with some countries considering the implementation of legislation for these purposes. Magnesium is arguably the most important mineral for the body, being utilized by every organ, in particular the heart, muscles and kidney. It is the fourth most abundant cation in the body, and the second most in intracellular fluid [21]. Magnesium deficiency has also been

**Figure 4.**

scientifically proven to either trigger or cause the following health problems: heart disease, diabetes, migraines, anxiety, hypertension, depression, fatigue, blood clots, liver disease, kidney disease, osteoporosis, insomnia, fatigue, cystitis, nerve problems and hypoglycemia [22]. Despite these facts and the relative importance of magnesium, up to 75 percent of people do not receive the recommended daily intake of magnesium (based on studies performed in the US – global intakes vary greatly based on local diets) [23].

Magnesium deficiency has also been specifically linked to higher rates of mortality in terms of cardiovascular deaths as well as general mortality. A recent German study sampling over 4000 people showed a strong correlation between low serum magnesium levels (< 0.73 mmol/l) and cardio-vascular deaths at a rate of 3.44 deaths per 1000 person years, in comparison to 1.53 deaths per 1000 person years for those with higher magnesium concentrations. More importantly though, an even stronger correlation was found between serum magnesium levels and allcause deaths. For those with low serum magnesium levels, the mortality rate was 10.95 deaths per 1000 person years compared to 1.45 deaths per 1000 person years at higher serum magnesium concentrations [24]. Researchers from the Bar Ilan University together with the Tel HaShomer Hospital gave more weight to this argument based on their review of death rates in Israel in areas serviced by desalinated water in comparison to those supplied by natural water. In their study they noted a marked difference in the number of deaths from heart disease in the areas that were supplied water from desalination in comparison to those supplied by natural water which had not been previously recognized when comparisons were made before desalination was introduced [25]. The conclusion was drawn by the researcher that this was as a result of decreased magnesium intake in these areas after the introduction of desalination. Only causal links however were established with direct links still to be proven.

Furthermore when reviewing the total number of epidemiological studies from the 1950's until present that had been performed on the link between magnesium in drinking water and cardiovascular mortality, it was determined that there is enough evidence to support a link, especially for concentrations above 5 mg/l [26]. The World Health Organization (WHO) also recommends maintaining a minimum Mg2+ concentration of 10 mg/l in all drinking waters [27]. Despite these recommendations, the replenishment of magnesium salts is rarely performed, if at all. One of the main obstacles is the relative cost of current methods, which significantly increases the total cost to desalinate and stabilize the water. The addition of magnesium to drinking water is commonly achieved through the dosing of chemical solutions such as magnesium chloride or magnesium sulphate. The high solubility of both salts allows for the supply of highly concentrated solutions, or simple solution make-down systems on site using crystalline salts, and the accurate dosing of these

### *Remineralization and Stabilization of Desalinated Water DOI: http://dx.doi.org/10.5772/intechopen.99458*

solutions to produce the magnesium concentration in the final water as desired. Whilst this process is very effective, it is also extremely expensive. This tends to position this subject as a question of luxury rather than necessity. The use of these chemicals in fact renders an additional cost to treat the drinking water by as much as 10 US cents/m<sup>3</sup> [28]. For this reason, more cost-effective replenishment methods are becoming of high interest in order to be prepared to face the coming soon changes in drinking water regulation.

The use of natural minerals for the replenishment of magnesium offers both a low cost and sustainable alternative to chemical dosing. Like natural calcium carbonate, magnesium minerals are found within the earth's crust as a range of sparingly soluble compounds, that naturally replenish themselves through dissolution and precipitation cycles. The fact that these minerals are sparingly soluble, increases their prevalence and concentrations in nature, as they are more likely to precipitate from solution than their highly soluble counterparts such as magnesium chloride and magnesium sulphate, which are only found in limited locations such as the dead sea, and in these cases require further refining. The limited solubility of the minerals means that they alone, struggle to provide the required levels of dissolved magnesium without the addition of an acid, to increase both their rate of dissolution and total concentration. Additionally, because the low solubility and slow reaction kinetics of these minerals, large contact tanks and an expensive installation are often required to achieve the required levels of dissolution. These issues can be effectively countered through the dissolution of powdered products within Membrane Calcite Reactors, similar to that for calcium carbonate.

Alternatively, some companies and research institutions are investigating the "mining" of natural resources from the brines rejected by desalination processes. These are often rich in a number of metals and minerals that are essential for industry and otherwise not scarce in availability [29]. These include magnesium, scandium, vanadium, gallium, boron lithium, indium, molybdenum and rubidium. These processes could offer a cheap source of magnesium at desalination sites where they could be immediately reinjected into the final water to replenish some of what has been extracted.
