Introduction:

Water has unique properties that make it a perfect medium for industrial applications, in particular for heat transfer. However, issues like corrosion and scale formation might be brought on by the solvent’s characteristics. To solve these issues, it is mandatory to have a comprehensive knowledge of water chemistry.

Water is only dissolving many compounds but also dissolves more substances than any other liquid, thus it is considered the universal solvent. Water molecule is polar with partially-positive and negative charges, which gives water the ability to dissolve ions and polar molecules.

Figure 1: Solvent characteristics of water

 

Due to the property of water as a good solvent many impurities can exist in water where anything other than water molecules in water is considered as an impurity.

The amount of impurity in water is measured in milligrams per liter (mg/L), which is equal 0.0001%, in diluted solutions it is common to refer to level of impurities in parts per million (ppm) since a one liter of water weighs 1,000,000 mg, considering that in density of water is ~ 1 g/cm³

1 mg contaminant in 1,000,000 mg of water = 1 mg in 1Kg = 1 mg/L

In water treatment application it is important to define these impurities, then based on the planned use of water we specify criteria for the acceptable levels of each impurity, and develop cost-effective treatment strategies to meet the predetermined quality standards. It is essential to understand that the terms impurity, contamination, and pollution are relative qualities of water.

Types of Impurities:

Impurities in water classified in many ways based on different categories, the most common classification is based on solubility of impurities whether it is soluble or insoluble in water, but still the classification will be segmented into further groups such as types of impurities and whether it is inorganic and organic in nature, generally dissolved impurities in freshwaters are mostly inorganic.

Figure 2: General classification of water impurities

Organic compounds are classified in a separate category, because some organics, even at low levels, can be harmful to human health in drinking water or cause disturbance in industrial systems, whether they are soluble or insoluble.

Biological growth like microorganisms, bacteria, or algae is also categorized separately, owing to the complexity of these living creatures and the generated by-products caused by them.

The following classification of the impurities enables a logical explanation of each set of components and their impact on industrial systems:

• Soluble matter
• Insoluble matter (suspended material)
• Organic contaminants
• Biological contaminants
• Dissolved gases
• Radioactive materials

The below list demonstrates the complexity of water-based systems as well as the difficulties that might arise when employing natural water in industrial, or municipal water systems.

• Dissolved Substances
• Bicarbonate
• Calcium
• Chloride
• Silica
• Sodium
• Sulfate
• Aluminum
• Barium
• Borate
• Bromide
• Copper
• Fluoride
• Iron
• Manganese
• Heavy Metals
• Nitrate
• Phosphate
• Potassium
• Strontium
• Zinc
• Suspended Matte
• Organic Matter
• Biological Organisms
• Bacteria
• Fungi and Algae
• Dissolved Gases

 

Solubility and Solubility Product Constant:

Solubility is defined as the maximum amount of a substance that will dissolve in a given amount of solvent at a specified temperature. Solubility is a characteristic property of a specific solute–solvent combination, and different substances have greatly different solubilities.

Figure 3: Solubility levels of a substance in water

The following are the solubility rules for common ionic solids, these rules are general and have some exceptions, but it gives an overview on salts’ solubility in water:

• The salts of sodium, potassium, and ammonium are highly soluble.
• Most halides (Cl, Br, I, etc.) are soluble, except for fluoride.
• Certain heavy metal cations (Pb and Ag) form insoluble halides.
• Salts containing nitrate ions (NO₃^–) are generally soluble.
• Most sulfate salts are soluble. Important exceptions to this rule include CaSO₄, BaSO₄, PbSO₄, Ag₂SO₄and SrSO₄.
• Most carbonates, hydroxides, and phosphates are only slightly soluble, with the exception of those associated with Na⁺, K⁺, and NH4^+.

The solubility product constant, K_sp is the equilibrium constant for a solid substance dissolving in an aqueous solution at a given temperature. So, the more soluble is a substance, the higher the K_sp value it has.

As an example of a common substance in water, calcium carbonate (CaCO₃) solubility product can be calculated as followed:

K_sp = [Ca⁺²][CO₃^-2]

Where the ions between square brackets express the concentration of each ion in mol/L.

Adding an extra amount of carbonate ions will be reflected in the reduction of calcium ions concentration, at a given temperature, since the solubility product is constant. This concept enables us to predict the possibility of having precipitation of a certain chemical inside an industrial water system.

Water pH:

It is well known in the water treatment field that pH is one of the key parameters considered while studying water characteristics for a system. pH is defined as a quantitative measure of the acidity or basicity of aqueous or other liquid solutions.

It is important to understand how pH affects the treated systems. The goal is to apply these principles while putting a treatment program for a processed water. To understand the pH value, it is mandatory to know the equilibrium reaction of water dissociation into Hydrogen and hydroxyl ions.

H₂O → H⁺ + OH^–

This equilibrium is governed by the dissociation constant in the following equation:

K_w= [H⁺][OH^–] = 10^-14

This means that as the hydrogen ion concentration decreases, the hydroxide ion concentration increases proportionally. At neutral pH, [H⁺] and [OH^–] are equal, which means the molar concentration of both hydrogen and hydroxyl ions is 10^–7 mol/L.

The term pH is measured as the negative logarithm of the hydrogen ion concentration [H⁺]:

pH = -log [H⁺]

Figure 4: pH scale and examples of the pH for different solutions

So, the pH value 7.0, is the neutral point as both concentrations are equal as explained, while pH values below 7.0 indicate an increasing concentration of [H⁺] and thus an increase in acidity, and pH values above 7.0, indicate increasing in the [OH^–] and hence the alkalinity is increasing.

The pH scale is an indicator of the balance between hydrogen and hydroxide ions. It is not a quantitative measure of the level of acid or alkaline substances dissolved in water.

Alkaline Species in Water:

Alkalinity is a measurement of dissolved alkaline substances in water (higher than 7.0 pH) that can neutralize the acid. The three main alkaline species are:

• Bicarbonate (HCO₃⁻)
• Carbonate (CO₃^-2)
• Hydroxides (OH^–)

The relationship between pH, carbon dioxide, and alkalinity describes the general conditions of water, but it is still important to know the amount of individual alkalinity species in water.

Figure 5 illustrates the distribution of these species against pH, at pH below 4.3, only CO₂ is significant, while at pH between 4.3 and 12.3 two inorganic carbon species exist, at pH above 12.3, only CO₃^–2 is significant, while at pH below 8.3, only HCO₃^– and dissolved CO₂ exist in water. Above this pH, HCO₃^– and CO₃^-2 are the predominant species. Of course, above pH 7, the OH^– concentration increases and contributes to alkalinity. The concentration of OH^– must be known to describe the alkalinity of water completely.

OH^– does not have an appreciable effect on alkalinity until pH is above 9.3, where the OH^– concentration is above 1 mg/L as CaCO3.

Figure 5: Alkalinity distribution vs pH

Using these relationships, along with titration measurements of alkalinity described earlier, three types of alkalinities can be defined:

• Total or M alkalinity at the methyl orange endpoint (ph about 4.3)
• P alkalinity at the phenolphthalein endpoint (pH about 8.3)
• Caustic alkalinity, which can be calculated from the other two

Figure 6: Relationship between P and M Alkalinity and the Hydroxide, in Water

Oxidation and Reduction Reactions:

Redox reactions, or what is commonly referred to as oxidation and reduction reactions, are very important in water chemistry, it is part of the fundamental processes in water such as disinfection, corrosion of metallic pipelines, and oxidation of sulfide for odor elimination from wastewater.

Redox reaction can be defined as a chemical reaction in which electrons are transferred between two reactants participating in it.

Figure 7: Redox reaction illustration

In an oxidation reaction, electrons (e^–) are lost from a chemical reactant, like the oxidation of iron metal in corrosion, which releases two electrons:

Fe → Fe⁺² + 2e-

In this oxidation reaction, iron is an electron donor, and a reducing agent, since it supplies the electrons, which can reduce other reactants in water. This reaction is called a half-cell reaction, because the electrons must be accepted by another half-cell reaction. To maintain
electrical neutrality, every oxidation reaction in a system must be accompanied by a reduction reaction. In this reduction reaction, another reactant gains the electrons donated by the reducing agent, as in the half-cell reaction below:

In this case, oxygen is an electron acceptor and is reduced, since it gains the electrons. Oxygen can also be called an oxidizing agent or an oxidant.

These reactions should be familiar since they are the corrosion reactions for iron-containing metals in water. They are called electrochemical reactions since redox reactions involve the transfer of electrons. The power, or potential, of redox reactions, can be determined from the electrochemical series

Figure 8: Table of Standard Electrode Potentials for Common Oxidizing Agents.

In redox reactions, chemicals at the extremes of the series are more powerful oxidants or reductants. Sodium metal is a very strong reducing agent with a negative potential of –2.7 V. Fluorine is a very strong oxidizing agent with a positive potential of 2.87 V. There would be a very strong driving force for a reaction between sodium and fluorine.