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Ions, Acid and Alkali Chemistry
1.
Ions
Dissociation of an ionic substance into its constituent ions is called ionisation. The electrically
charged species formed by dissociation of an ionic substance in water solution are called ions. There
are an equal number of positive and negative charges formed in the water. Thus, the water has no
net electrical charge; the solution is electrically neutral.
Ionic substances are electrically charged when dissolved in water. Non-ionic substances remain
electrically neutral in water solution.
As an example:
- Salt dissolves in water and is ionic, i.e., salt molecules dissociate in water, forming positively
charges sodium ions and negatively charged chloride ions.
- Sugar, on the other hand, is non-ionic and simply exists as molecules of sugar in water solution.
The electric charge or charges on an ion determine the properties of that ion. Electrical charges on
ions are similar to magnets in that like repels like. Positive ions repel each other and combine only
with negative ions.
A more common name for a positively charged ion is the term CATION. The term for a negatively
charged in is ANION.
+
Ions may consist of individual atoms with an electrical charge (e.g. sodium, Na ) or a group of atoms
2“+”
“-“
with a net electrical charge, such as carbonate (CO 3 ). The symbols and are used as superscripts
“+”
“-“
to denote a positive or negative charge, respectively. A number preceding the or symbol
indicates that the ion has a multiple charge, for example, carbonate has a charge of minus two (2 ).
Table 1 lists ions are encountered in water processing.
Table 1 – Common Ions Found in Water
Name
Formula
-
Ammonium
NH4
Bicarbonate
HCO3
Bisulphate
HSO4
-
2+
Calcium
Ca
Carbonate
CO3
2-
-
Chloride
Copper:
Cl
Cuprous
Cupric
Cu
Cu
+
2+
-
Fluoride
F
Hydrogen
H
Hydroxide
(OH)
Hypochlorite
(OCl)
+
-
-
Iodide
Iron:
-
I
2+
Ferrous
Fe
3+
Ferric
Fe
Magnesium
Mg
Manganese: Manganous
Mn
2+
Mn
4+
2+
Manganic
-
Nitrate
(NO3)
Nitrite
(NO2)
-
Permanganate
(MnO4)
Phosphate
(PO4)
Potassium
K
-
3-
+
Silicate
(SiO2)
Sodium
Na
Sulphate
(SO4)
Zinc
Zn
-
+
2-
2+
The ionisation of sodium chloride can be written as an equation in the form of:
+
+
NaCl ------------------- Na + Cl
In a similar way, the reaction of a cation and an anion can be written. For example, the hydrogen
cation combines with the hydroxyl anion to form water; the reaction is symbolized by the following
equation:
+
-
H + OH ---------------- H2O
Note that the chemical formula for water is written as H2O, which is equivalent to HOH.
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1.
Neutralization of Acids and Bases (Alkalis)
Acids and bases react together to cancel or neutralize the chemical properties of each other. Acids
+
are ionic compounds which, when dissolved in water, ionise to release hydrogen ions (H ). Bases, also
called alkalies, are ionic compounds which, when dissolved in water, ionise to release hydroxide ions
(OH ).
The properties of acids and bases are actually properties of hydrogen cations and hydroxide anions.
Combining an acid with a base, results in a neutralization reaction. The hydrogen cation combines
with the hydroxide anion to form a water molecule and the remaining cations and anions are a “salt.”
The neutralization of hydrochloric acid (HCL) with sodium hydroxide (caustic soda) NaOH, forms
water and salt, or sodium chloride (NaCl), and is illustrated in the following chemical equation:
HCL + NaOH ------------ H2O + NaCl
The salt remains in solution.
2.
pH – An Indicator of Acidity or Alkalinity
The term pH, commonly used in water chemistry, is expressed as a number from 0 to 14 and
indicates whether the water is acidic or alkaline. While an absolutely 100% pure water has a pH of 7, a
water with pH 7 is not necessarily 100% pure. That is, a pH of 7 indicates that a water is neutral,
neither acidic nor basic. A pH value of less than 7 denotes to acidic water, while a pH value greater
than 7 indicates alkaline water.
The pH scale is not a simple numeric scale, but is logarithmic. Thus, water at pH 4 is 10 times more
acid than water at pH 5 and 100 times more acid that water at pH 6.
An excess of hydrogen ions makes a solution acidic, while an excess of hydroxyl ions makes a solution
alkaline. The pH of solution is measured using a pH meter, or indicator paper on a scale of 0 to 14,
and is actually defined as – log 10[H+], where [H+] is the hydrogen ion concentration. Solutions with
values from 0-7 are acidic, while those from 7-14 are alkaline. A solution with pH 7 is said to be
neutral, and has an equal amount of hydrogen and hydroxyl ions.
Because it is measured on a negative logarithmic scale, a 1 unit increase in pH corresponds to a 10
fold decrease in hydrogen ion concentration. The pH of a solution indicates the relative strength, but
cannot be used to determine actual concentration of an acid or alkali.
While solutions in certain pH ranges are acidic or alkaline, these values should not be confused with
actual determinations of % alkalinity and % acidity, which are done by chemical titration and are
proportional to concentration.
Natural occurring waters rarely have a pH less than 5.5. Low-mineral-content waters typically have a
pH of 6.0 to 7.0, because the carbon dioxide dissolved in the water is converted to carbonic acid,
more alkaline waters have typical pH values of 7.0 to 9.0.
3.
Alkalinity
Alkalinity is the term applied to the measure of a waters capacity to absorb acids. A small portion of
alkalinity may be cause by silicates, borates, or phosphates, especially in highly mineralised waters.
Most of water’s alkalinity, however, is due to the presence of carbonates.
Alkalinity of water is measured in a laboratory and reported as ppm (or mg/L) of calcium carbonate
(ppm as CaCO3).
Alkalinity is determined by titration of a known weight of sample, with a strong acid whose strength
is known. A solution may have only a slightly alkaline pH of 8, but have high alkalinity or great ability
to neutralise acid.
Acidity is determined by titration of a known weight of sample with a strong alkali.
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Most natural waters contain some alkalinity and have a pH <8.0 Alkalinity is normally due to the
presence of hydroxides, carbonates and bicarbonates, and is therefore used as a measure of the
concentrations of these constituents.
Because of the almost universal presence of the alkaline ions, few natural sources of water contain
measurable acidity or have a pH of 4.4 or less.
Most potable water supplies have a pH in the range of 6.0-8.0.
-
In the pH range 4.4 - 8.2, alkalinity in water is due to the bicarbonate ion (HCO 3) and carbon dioxide,
CO2.
Note: Two water samples with the same pH can have different ratios of bicarbonate to carbon
dioxide, and different alkalinities. The pH alone of water is not a measure of its alkalinity).
Although natural water does not usually have pH >8.4, above a pH of 8.2 to 8.4 the carbonate ion
2(CO3 ) becomes the source of alkalinity.
4. Definitions and Terminology
Alkalinity is often expressed as either “active/free” or “total”.
Alkalinity measured above pH 8.2-8.4 is generally known as “free alkalinity” or (or “P” alkalinity) as it
is free to react with a strong acid.
“Total alkalinity” (or “M” alkalinity) is free alkalinity plus that alkalinity measured from pH 4.3-8.3.
The acid reacts with the source of alkalinity to produce a colour change and an endpoint.
Alkalinity in water is important for several reasons, but the most notable is that highly alkaline waters
have a greater tendency to form bicarbonate and carbonate scales when calcium and magnesium
ions are present. This tendency is greater in the presence of heat.
When used in the production of drinking water, alkaline water can affect flavour.
Water with high alkalinity can also reduce the efficacy of some acidic sanitizers through
neutralization.
As alkalinity limits are often imposed on effluent water, a higher starting concentration can mean
more treatment will be necessary before disposal.
Water for use in any type of food or water plant must be microbiologically fit for human
consumption. This includes water used for rinsing, cleaning and sanitising as well as processing.
Local health authorities are responsible for determining the microbiological standard for water in a
particular area. A certain group of bacteria some of which are found in the intestines and faeces of
warm blooded animals (including man) are used as indicators of the probable microbiological safety
of water. If the number of these organisms (called coliforms) is below the specified limit, (usually <1
per 100 ml of water) it may be assumed that there are no disease causing organisms present.
While potable water is free from pathogenic organisms, it may still contain other types of microbes,
including those that can contribute to food spoilage. In plants where post sanitizer rinses (i.e. a water
rinse after the final cleaning and/ or sanitising step) are used, the final rinse water could be a vehicle
for introducing spoilage organisms onto equipment. In these cases, potable water may need to
undergo additional treatment to ensure suitability for this use.
These non-pathogenic organisms can also be responsible for affecting water flavour and slime
formations in the processing environment.
Micro-organisms that are suspended in water are part of the total solids content. Most can be filtered
out using very fine filters (0.2 micron), or can be killed with chemicals commonly used to treat water,
like chlorine or ozone.
Aside from the health reasons for using microbiologically acceptable water, there are other problems
associated with microbes in water. Some microbes can form corrosive by products in water. A group
of organisms known as "sulphate reducing bacteria" are often times responsible for microbial
corrosion through the formation of sulphides. The best known of these are members of the genus
Desulfovibrio.
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5.
Water Hardness
Water containing calcium, magnesium, and iron salts in dissolved form is known as hard water.
This term originated many years ago as people noticed that when using water from some sources, it
was harder to make suds when doing laundry. More soap was necessary to produce an acceptable
amount of suds, which made cleaning harder, and the term "hard water" was born.
Hardness in water is due primarily to the presence of salts of calcium and magnesium. These two
metals can form salts that range from very soluble to totally insoluble, but heat and reactions with
other chemicals, including cleaning chemicals, can make normally soluble salts, containing these
ions, insoluble.
These insoluble materials interfere with detergency by making less of the active material available
for cleaning, by suppressing foam, falling out of solution and depositing onto surfaces as precipitate
or scale.
Soft water is always better than hard water for cleaning and rinsing purposes.
2+
Water hardness is generally thought of as the sum of calcium (Ca2+) and magnesium (Mg ) ions,
although other metals like aluminium and iron can also be troublesome. These ions have a positive
charge and are called cations, and can combine with many different negative charged ions (anions)
to form salts. Regardless of the ions responsible, it is the convention to express hardness as calcium
carbonate (CaCO3) in mg/L or ppm).
Water harness is generally divided into two types, temporary and permanent, with total hardness
being the sum of the two.
Soluble calcium and magnesium bicarbonates are responsible for temporary hardness in water. Heat
is all that is needed to precipitate these compounds into insoluble carbonates, and this type of
hardness causes most of the scale seen in various types of heat exchange equipment.
Hard water is objectionable for two main reasons:
1. The formation of insoluble soaps; and
2. The formation of a hard, insoluble, tightly adherent scale.
This build-up is most notable in boilers. At high temperatures, much of the mineral matter dissolved
in hard water is precipitated as scale. Boiler explosions have been linked to scale build-up.
These insoluble alkaline materials can be dissolved by treating with acid. Acidic cleaners are used to
remove hard water deposits in Water Plants.
It is possible to have all permanent hardness, or temporary hardness or both in any given sample of
water. Regardless of the type of hardness, detergents must work harder when it is present. Most
cleaners for industrial use now contain materials for keeping hardness from interfering with the
cleaning process, or precipitating out onto surfaces.
It is more difficult for detergent to keep soil suspended in hard water, and various types of scales
often result. Film and scale formation result in both aesthetic and sanitation problems.
Not only are scales unsightly, but they provide an ideal place for bacteria to attach and be protected
from the action of sanitizers. Once scale is present on smooth surfaces, it is easier for additional
layers of scale to deposit. Scales can interfere with heat transfer and reduce the efficiency of heating
equipment leading to excessive power consumption, poor performance, and increased chemical.
Heavy scale reduces the diameter of pipes, impeding flow. Water hardness is typically expressed as
ppm or mg/L of calcium carbonate (ppm of CaCO3). The U.S. Geological Survey (USGS) classification
system of the relative hardness or softness of water is shown in Table IB-5.
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Table 2 - USGS Classification of Hard/Soft Waters as an Example
Classification
ppm (mg/L) as CaCO3
Soft
0-60
Moderately hard
61-20
Hard
121-180
Very Hard
>180
Since water dissolves many substances that it contacts, the "total solids" content of water can be an
important measure of quality.
Solids may affect water in a number of ways, including imparting a bad taste or colour if
concentrations are high enough.
Potable water usually contains less than 500 mg/L of total dissolved solids.
7.
Oxidation-Reduction Reactions
Oxidation / reduction reactions always occur together. They involve the transfer of electrons. The
substance losing electrons, becoming more positively charged is said to have been oxidised while the
substance gaining electrons, often referred to as the oxidising agent is said to be reduced.
Thus the reaction used to clarify iron containing water of oxidising iron 2 to iron 3 with chlorine is an
oxidation / reduction reaction:
2+
2 Fe + Cl2  2 Fe
3+
+ 2 Cl
-
An oxidising agent is a substance that can oxidize other materials e.g., ozone or chlorine. Oxidizers
are commonly used as disinfectants to kill micro-organisms. Air is an oxidizer because of its oxygen
component. Oxidising agents are also used to convert certain ions to more insoluble forms that can
be filtered out, as shown above.
8. Corrosion
Dissolved minerals in water increase the electrical conductivity of water, this can accelerate the
corrosion rate of metal equipment. Chloride and sulphate ions are of particular importance here.
Most potable and raw water has conductivity in the 20-500 mhos range.
Acidity, alkalinity and pH are all important measures of the suitability of water for a particular
application:
+
Alkalinity is defined as the ability of a solution to neutralise acid (hydrogen ions H ).
Acidity is the ability of a solution to neutralise alkali (base) or hydroxyl ions (OH-).
These are measures of chemical concentration.
pH is a measure of the hydrogen ion concentration in solution and stands for "potency of hydrogen".
Chemically neutral water contains equal numbers of hydrogen and hydroxyl ions.
Corrosion refers to the tendency of all metals, when exposed to the elements, to revert back to the
more stable forms found in earth’s crust. The products of corrosion are usually in the form of oxides,
carbonates or sulphides.
Corrosion occurs more quickly at low pH and in a high concentration of oxidising agents such as
ozone, chlorine, or oxygen.
For Further Information:
Australian Beverages Council Ltd
[email protected]
Correct as at 17th October, 2012.
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