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The Environmental Science of Drinking Water
By Patrick J. Sullivan Franklin J. Agardy James J. J. Clark
BUTTERWORTH-HEINEMANN
Copyright © 2005 Elsevier Inc.
All right reserved.
ISBN: 978-0-08-045772-7
Chapter One
The Water We Drink
"The boundaries between water and wastewater are already beginning to fade." The American Water Works Association (2001)
In a landscape dominated and modified by human activity, it should not be surprising that the water we drink contains both chemical and biological pollutants. After all, the world is dependent on the ever expanding chemical wonders of our agricultural/industrial/pharmaceutical-based society that helps support a growing population. An increasing population in turn pollutes water resources with chemicals and biohazards from the production and use of consumer products, agricultural and animal production, and human waste disposal. This ever increasing spiral of population and pollution means that naturally pure sources of drinking water are almost extinct.
With our technological advances and increased population, there is a price we all pay. This price is drinking water that contains a mixture of manufactured chemicals or biohazards that are a potential threat to human health. This threat can be immediate when an individual is exposed to a water-borne disease, or long-term when drinking chemical pollutants. For those living in industrialized nations (e.g., United States, Canada, European Union), the threat from water-borne biohazards has been significantly reduced or eliminated. However, widespread exposure to specific chemicals or chemical mixtures in drinking water still remains. This condition illustrates the necessity to understand the environmental science of our drinking water if this expanding problem is to be managed or reduced.
Environmental science evaluates how humans use the earth's natural resources (in this case, water resources), appraises the repercussions that occur as the result of this use, and evaluates how to mitigate these impacts. Therefore, the objective of this book, by focusing on the environmental science of drinking water, is to provide a basis for understanding the threat posed by manufactured chemicals and biohazards in water resources and the solutions available for minimizing the potential health risk associated with our abuse of this natural product.
Natural Water
Before human activity on this earth, the chemistry of water was initially influenced by the dissolution of minerals from soil, rock, biosynthesis, and biodegradation of organic matter. The chemical compounds that dissolve from minerals, biosynthesis, and biodegradation represent natural or background levels in the water we drink. In some cases natural water can contain elevated concentrations of trace elements (e.g., arsenic, copper, fluorine, lead, zinc) that are known to be detrimental to human health. Because many natural waters (i.e., advertised to contain little or no manufactured pollutants) from around the world are bottled and sold as a very expensive healthy alternative to tap water, it is important to at least understand the basic chemistry of natural water.
The hydrologic cycle is the process that has the greatest influence on the chemistry of natural water. For example, when precipitation falls on the land, it follows one of many paths that constitute what is known as the hydrologic cycle (Figure 1-1). As water runs off the earth's surface, it can infiltrate into the soil, as well as form into bodies of surface water such as ponds, lakes, and rivers. Much of this surface water will evaporate back into the air, which can reform as clouds and subsequently result in precipitation, or it can continue to percolate into the soil. Once water moves into the soil, it can be removed by plants and evaporated back to the atmosphere (transpiration), or it can continue to move downward to form water-saturated zones in soil and fractured bedrock. These saturated zones are called aquifers. In aquifers, groundwater can find its way back onto the land as a spring, a natural artesian well, or by seeping into ponds, lakes, and rivers. Thus, water is naturally cycled from the earth's surface into groundwater and back again. This cycle may be only hours long or may occur over centuries. While groundwater remains in an aquifer (i.e., it is not being lost to the surface through a spring or being pumped from a well), it is essentially a stored water resource that can be tapped by a well. Based on this cycle, water resources are usually classified as being either surface water or groundwater. This is an important difference, as the chemistry of each resource is unique to its origin.
Contact with Soil and Rock
Because water that exists at and below the earth's surface is in contact with soil and rock, some mineral or organic matter will be dissolved into the water. For the most part, the chemical elements that will be dissolved in water are generally preordained by their abundance. For example, the average abundance of the most common chemical elements in the earth's crust (in decreasing order of abundance, see Appendix 1-1) are oxygen (O), silica (Si), aluminum (Al), iron (Fe), calcium (Ca), sodium (Na), magnesium (Mg), and potassium (K). These first eight elements are the building blocks of the most common minerals that make up the earth's crust. The next 20 most common elements are titanium (Ti), hydrogen (H), phosphors (P), manganese (Mn), fluorine (F), barium (Ba), strontium (Sr), sulfur (S), carbon (C), zirconium (Zr), vanadium (V), and chlorine (Cl). All combined, these 28 elements make up 99.93 percent of the earth's crust.
Therefore, it would be anticipated that natural water would contain some or most of these dissolved elements. Based on the natural properties of water (H2O) and the properties of each chemical element and mineral combination, the actual occurrence and concentration of an element in water can vary widely. When minerals dissolve in water, the chemical elements are usually ionized (i.e., they form a charged chemical species called an ion). These ions occur as either cations (a positively charged ion) or as anions (a negatively charged ion). The major cations and anions that occur in both surface water and groundwater are given in Table 1-1. These common ions usually occur in water at levels measured in the part-per-million range.
Another crucial chemical property of water is its relative acidity or alkalinity. This chemical characteristic has a direct influence on the natural water. For example, acid water will generally tend to have more dissolved trace elements at higher concentrations than alkaline water. Therefore, it is important that the relative acidity or alkalinity be measured by determining the pH. The pH of water is the negative logarithm of the hydrogen ion concentration. Although this is the exact definition, it is more important to understand its meaning in everyday use. Therefore, it is necessary to define the common terms associated with this measured value.
Since the 17th century, acids have been described as substances with a sour taste and the ability to dissolve many different substances. Recently, the simplest chemical definition of an acid was proposed by Arrhenius to be a substance containing hydrogen, which, upon its dissolution in water, gives off hydrogen ions (H+) into solution. This is illustrated by Equation 1-1, which represents sulfuric acid dissolved in water.
H2SO4 = 2H+ + SO42- (1-1)
In this case, one molecule of sulfuric acid yields two hydrogen ions and one sulfate ion.
Bases have been described as substances that when dissolved in water feel soapy or slippery (that is because your skin is being dissolved), have a bitter taste, and can neutralize acids. According to Arrhenius, a base is a substance that gives free hydroxide ions (OH-) when dissolved in water. For example, the ionization of calcium hydroxide would be represented by Equation 1-2.
Ca(OH)2 = Ca2+ + 2OH- (1-2)
In this case, one molecule of calcium hydroxide yields one calcium ion and two hydroxide ions.
If an acid and base are mixed in equal proportion, the hydrogen ion and hydroxide ion will combine to form water so that there is no hydrogen or hydroxide ions dissolved in water. When such a reaction occurs the water is not acid or alkaline but neutral. For example, a neutralization reaction between sulfuric acid and calcium hydroxide is represented by Equation 1-3.
2H+ + SO42- + Ca2+ + 2OH- = CaSO4 + 2H2O (1-3)
(Continues...)
Excerpted from The Environmental Science of Drinking Water by Patrick J. Sullivan Franklin J. Agardy James J. J. Clark Copyright © 2005 by Elsevier Inc.. Excerpted by permission of BUTTERWORTH-HEINEMANN. All rights reserved. No part of this excerpt may be reproduced or reprinted without permission in writing from the publisher.
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