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5. Surface Water Quality
Quality of water: Physical, chemical, and biological
attributes that affect the sustainability of water for drinking,
agriculture, recreation, and other uses. In this section, we
will look at the pollution of surface water by suspended
sediments and various municipal and industrial wastes.
Effects of sediments
Suspended load may be considered pollutant when it
exceeds natural concentration.
An increase in turbitity affects the biotic balance. For
example, the increase in turbitity and the clogging of
interstices in stream-bed gravel by silt has a detrimental
effect on the survival of eggs and developing minnow of
Turbitity also reduces the depth to which sunlight penetrates
and thus alters the rate of photosynthesis; e.g. algae growth.
This may increase dissolved oxygen during hours of
sunlight, but the decay of organic material consumes oxygen
during hours of darkness.
Chemical characteristics of natural water
Chemical characteristics : determined by the kinds of
dissolved matters and their concentrations.
Example: “clean” rain water
- free from metals (Na+, Ca2+ etc.) and organic matter.
- significant dissolved gases, in particular O2 and CO2,
which play major roles in weathering processes.
mg/L and ppm (parts per million)
1ppm = 1g/106g = 10-3g/103g = 1mg/L
The mol is a unit to express chemical reactions. 1 mol of
Na and 1 mol of Cl form 1 mol of NaCl. We could have
said 22.99 g of Na and 35.45 g of Cl, but this is not so
Notation: [Cl-] indicates the concentration of Cl in mol/L.
pH = -log10[H+]
H2O ⇔ H+ + OH[H+][OH-] = 10-14
In neutral water, [H+] = [OH-] = 10-7. i.e. pH =
pH of rain water
Dissolved CO2 makes “pure” rain water acidic.
H2O + CO2 ⇔ H+ + HCO3- (bicarbonate)
Rain water is acidic (pH ≅ 5.6) without air pollution. With
the addition of H2SO4 and HNO3 from burning fossil
fuels, it becomes the “acid rain” (pH < 5.6).
Chemical weathering
H+ attacks rock-forming minerals.
H+ + CaCO3(calcite) ⇔ Ca2+ + HCO32H+ + 2KAlSi3O8(feldspar) + 9H2O
⇔ Al2Si2O5(OH)4(kaolinite) + 4H4SiO4 + 2K+
In each case, H+ is neutralized by the mineral. Dissolution of
calcite is fast. Calcite-rich soils offer a large buffering
capacity. Dissolution of feldspar and other silicate minerals
is very slow, not effective as a buffer.
The prairie soil contains a large amount of calcite derived
from limestone.
Upon weathering, soil water gains the elements necessary
for plant growth. Excessively high concentration, however,
is harmful to plants. → soil salinity
Major ions
Most natural water is characterized by:
Ca2+, Mg2+, Na+, K+, Cl-, SO42-, and HCO3-.
Note that pH of natural water is controlled by
H2O + CO2 ⇔ H+ + HCO3Minor species
Metals: Al, Si, Fe, Zn, etc. occurring in various forms
Nutrients: NO3-, PO43-, and dissolved organic matter
At high concentration, some species cause environmental
The maximum amount that can dissolve in water. When
water becomes concentrated during evaporation, the
concentration of some ions may become too high and the
minerals start precipitating.
Different minerals have different solubility.
e.g. NaCl has high, and CaCO3 has low solubility.
In general, the solubility of salts increases in this order,
Ca < Mg < Na < K
CO3 < SO4 < Cl
CaCO3 is the first to precipitate, and KCl is the last.
Effects of pH on metal solubility
Many metal species forms hydroxy complexes.
Al3+ + H2O ⇔ AlOH2+ + H+
AlOH2+ + H2O ⇔ Al(OH)2+ + H+
This process is called hydrolysis. In low-pH water the
abundance of the H+ ion prevents hydrolysis so that
dissolved Al exists primarily as free Al3+ ions. Under
neutral pH, a substantial portion of Al exists as Al(OH)3,
which has a low solubility. Therefore, the total solubility of
Al is the lowest in neutral water and higher in acidic and
basic water.
In this graph, the vertical
axis is essentially the
concentration of each
Drever, J.I, 1988. The geochemistry of
natural waters. Fig. 6-3.
Fe3+ have similar tendency. Precipitation of hydroxides
dictates the behavior of Al and ferric Fe in natural water.
In general, solubility of metals drastically increase in lowpH water.
Cation exchange
The surface of clay minerals have negative charge. Cations
(Ca2+, Na+, etc.) are attached to the clay mineral surface by
electrostatic force.
Any given soil has fixed amount of
clay, and hence fixed number of
- - - - clay mineral
negatively charged “sites”.
All sites have to be occupied by cations. A site can
exchange one cation with another. For example, a site can
release 2Na+ and take Ca2+ from water in exchange. The
negatively charged sites are called “exchange sites” and the
number of exchange sites per unit mass of dry soil is called
cation exchange capacity.
Higher [Na+]/[Ca2+] in water results in higher Na+/Ca2+ on
clay surface. When exchange sites are dominated by Na+,
soils becomes unsuitable for growing crops.
- poor drainage and aeration.
Adsorption of metals and anions
Some metals are adsorbed onto soil particles by the above
mechanism of cation exchange, but others are adsorbed by
more complicated mechanism. Also, negatively charged
ions (anions) such as SO42-, NO3-, and PO43- are adsorbed by
complex mechanisms specific to each ion. In general, PO43has much higher tendency to be adsorbed than NO3-. This
dictates the behavior of the two species in the environment.
Decay of organic materials
CH2O + O2 → CO2 + H2O
Pyrite oxidation and acid generation
FeS2 + 3.75O2 + 3.5H2O → Fe(OH)3 + 4H+ +2SO42In these reactions, electrons are removed from carbon and
sulfur atoms and added to oxygen atom. Oxygen is called
electron acceptor, and carbon and sulfur are called electron
donors. Each reaction release a certain amount of energy,
which may be utilized by bacterial communities.
In well aerated soil O2 is continuously replenished; this is
called aerobic condition. In poorly aerated environments O2
is quickly used up by reactions. This is commonly called
anaerobic condition.
Under anaerobic condition, species other than O2 are used as
electron acceptors;
NO3- → N2
SO42- → H2S
(sulfate reduction)
CO2 → CH4
The condition becomes more strongly anaerobic, or
reducing, as NO3- and SO42- are used up. In general, more
energy is released by reactions involving O2 than NO3-. In
other words bacterial communities consume O2 first, then
NO3-, and then SO42-. Therefore, denitrification does not
occur under aerobic condition.
Salinity of water
Total dissolved solids (TDS), expressed as mg/L.
Sea water : 35,000 mg/L
River water : 100 - 500 mg/L
TDS higher than a few thousands mg/L has adverse effects
on crops. TDS is related to the electrical conductivity (EC).
In a rough sense, TDS (mg/L) ≅ 0.7 EC (µS/cm)
Evaporative enrichment
Suppose 1 L of water in a pan, having TDS of 100 mg/L.
Evaporate 900 mL,
Composition also changes.
Start with (Ca, Mg, Na) - (HCO3, SO4).
As CaCO3 precipitates, water becomes more enriched in
Na, Mg, and SO4.
Origin of salts in the prairies
Glacial tills contain limestone and shale, which are high in
CaCO3 and FeS2, respectively.
FeS2 + 3.75O2 + 3.5H2O → Fe(OH)3 + 4H+ +2SO422H+ + CaCO3 + SO42- → Ca2+ + SO42- + CO2 + H2O
The “unoxidized” dark gray glacial tills have significant
FeS2. The “oxidized” olive brown tills have significant
Fe(OH)3. Soils derived from the oxidized tills have high
[HCO3-] and [SO42-], so as groundwater in the prairies.
Major ions affected by human activity
Return flow from irrigated fields has raised concentration of
major ions due to evaporative enrichment and the
dissolution of minerals and organic matter. If the stream
water is used several times over for irrigation, the salinity of
water in rivers or canals increases downstream. Water may
become unsuitable for irrigation.
Change in salt content of the Sevier River, Utah, as a result of repeated diversion for irrigation.
(D&L, Fig. 20-4)
Urban and industrial use also increases salt concentration.
Treated municipal sewage adds to the receiving water about
35 kg of inorganic salts per person per year. The water used
in industrial plants picks up a large quantities of solutes.
Besides the direct impact of municipal and industrial waste
disposal, the landuse may impact the salt concentration in
rivers by drastically altering the hydrologic cycle.
Trace metals
Heavy metals such as Fe, Zn, Pb, and Cu have low
solubility in neutral-pH water and are often mobilized by
forming soluble complex with organic molecules or by
becoming adsorbed to clay particles. At high concentration,
they can be toxic and can cause major disruption of aquatic
ecosystems. In Canada and US, waste water from industrial
plants are treated to reduce the concentration of heavy
However, groundwater originating from acid mine tailings,
when it discharges into stream, may lower the pH of water
and thereby dissolve metals. Note that fish are not killed by
low pH itself, but rather by the increased concentration of
metals. In particular, moderately high concentration of Al
can cause a serious damage to gills of some species.
The acidity comes from the oxidation of metal sulfides,
which is almost unavoidable in many operating and
decommissioned mines.
The emission of SO2 and NOx from industrialized area
causes acid rain. The catchments surrounding prestine
lakes on the Canadian Shield has very thin soil cover,
which offers little buffering capacity against acidification of
Suggested reading: Schindler, D.W., 1988. Effects of acid rain on freshwater ecosystems. Science,
239: 149-157.
Biodegradable wastes and the oxygen balance
The major consumption of dissolved O2 in streams occurs
through the aerobic chemical and microbial breakdown of
long-chained organic molecules into simpler, stable endproducts.
carbohydrate → CO2 + H2O
proteins → amino acids → ammonia → nitrite → nitrate
→ sulfate and phosphate
If O2 is exhausted, aerobic decomposition ceases and further
breakdown must be accomplished by anaerobic bacteria. A
moderately high dissolved O2 is necessary for the
maintenance of healthy aquatic ecosystems. When a large
amount of industrial or municipal waste enters the stream,
the breakdown of the waste may depletes O2 in the stream.
The pollution strength of a waste is measured by
Biochemical Oxygen Demand (BOD), which is the amount
of O2 consumed by living organisms (mostly bacteria) while
utilizing the organic matter in the waste. The biochemical
consumption of O2 is a complex process that may take
weeks and is dependent on many environmental factors (see
next page). It is usually measured under arbitrary conditions
in the laboratory; 5-day digestion at 20 ºC.
Chemical Oxygen Demand (COD) can be used to obtain a
rough estimate of BOD because it can be measured in a few
D&L, Fig. 20-7
D&L, Table 20-4
Table 20-4 lists 5-day
BOD of various wastes.
For example, 1 L of
untreated domestic
sewage demands up to
400 mg of dissolved O2.
D&L, Table 20-6
Table 20-6 lists the solubility
of O2 in water. At 20 ºC, it
takes 44 L of water to
accommodate 400 mg of O2.
When a waste enters a stream, it becomes completely mixed
with stream water after a short distance. The assimilation of
oxidizable materials begin to consume O2 and increase O2
O2 is replenished by reaeration from the atmosphere, the
rate of which depends mainly on the O2 deficit, the width of
the stream, turbulence, and water temperature.
The balance between O2 consumption and replenishment
leads to a profile of net O2 deficit, which shows a
characteristic dissolved-oxygen sag.
D&L, Fig. 20-9
The profile of O2 deficit is estimated by the Streeter-Phelps
− kt
− rt
− rt
k :
r :
t :
) + Di e
O2 deficit of the stream [mg/L]
initial O2 deficit of the waste mixture [mg/L]
rate constant of O2 consumption [day-1]
initial BOD of the waste mixture [mg/L]
reaeration rate constant [day-1]
time from initial mixing [day]
The word “initial” indicates the time when the waste is
completely mixed with stream water. The initial waste
mixture may have significant O2 deficit. The rate of O2
consumption per unit BOD is dependent on the type of
waste. Many agricultural processes release wastes that can
be rapidly oxidized. In contrast, pulp wastes are assimilated
slowly. Therefore, the rate constant k depends on the type of
materials in the waste. The constant r represents how fast O2
is being replenished. Both k and r are dependent on water
The S-P equation is very simplistic, but it illustrates how the
O2 sag depends on several important factors. Environmental
planners need to be familiar with this concept so that they
can effectively interact with sanitary engineers.
For example, what happens if r is drastically reduced by ice
covering a major portion of the river?
Variation of the dissolved O2 sag. Q is the stream discharge. The lower Q results in the higher
initial BOD of the waste mixture. (D&L, Fig. 20-12)
Strategies for reducing dissolved O2 problems
The BOD in wastes can be reduced by a set of treatment
processes. Primary treatment involves the settling of
suspended mineral and organic matters, which removes 4090 % of the suspended solids and 25-85 % of the 5-day
Secondary treatment involves biological processing. In a
trickling filter, wastewaters are sprayed onto columns of
crushed stone and flow in thin films. Organisms grazing on
the surface of stone break down the dissolved organic
materials. In a activated sludge process, flocs of bacteria
are stirred in the wastewater. These processes remove 5095 % of the BOD.
Tertiary treatment includes such practices as using a strong
oxidizing agent like ozone to remove BOD and the
addition of alum to precipitate phosphates which are not
removed by other treatment processes.
Alternatively, wastes can be processed in stabilization
ponds in which aerobic microbial organisms digest the
organic substances, but the pond is also colonized by
aquatic plants, which release oxygen during
photosynthesis. This method of treatment is gaining a large
popularity in recent years and more and more artificial
wetlands are constructed for this purpose.
More advanced reference:
Linsley et al., 1992. Water-resources engineering. 4th ed., Chapter 19.
Plant nutrients and eutrophication
The breakdown of organic compounds (see 10-16) removes
BOD, but it generates nitrates and phosphates, which are the
stable forms of nitrogen and phosphorous under the aerobic
condition. The concentration of NO3 and PO4 increases
immediately below the outfall of waste and decreases
gradually due to dilution. Note that the majority of PO4
exists in the forms HPO42- and H2PO4- in neutral-pH water.
PO43- + H2O ⇔ HPO42- + OHHPO42- + H2O ⇔ H2PO4- + OH-
Diagram of typical changes in water chemistry (D&L, Fig. 20-17).
N and P are among the essential nutrient elements for plants.
Eutrophication occurs when the excess nutrients causes a
significant increase in the rate of plant growth. In particular
algae can grow prolifically (algal bloom) with the excess
nutrients to cause a number of problems.
Algal blooms are unsightly; they may increase the cost of
water treatment; and upon their death and decay they
impose an O2 demand on the water, which may lead to fish
Plant growth is usually limited by the availability of one or a
few of the essential nutrients, for example P, N, and C. The
nutrient that most frequently limits algal production is P.
Note that N and C are contained in air. Therefore, by the
diffusive exchange, the algae can take those elements from
the air. P is only available from the transport of dissolved
PO4 by rivers or groundwater.
PO4 easily becomes immobilized in the soil by forming
insoluble calcium phosphate minerals. Natural waters from
undisturbed forest has dissolved PO4 concentration of 0.0050.05 mg/L. Usually 0.08-0.1 mg/L is necessary to trigger
blooms. The bulk of P transport in streams occur as organicP and PO4 attached to suspended sediments. These
“immobile” form of P may become available to plants later.
In agricultural lands, P is applied as fertilizers. The plant
uses only about 15 % of applied P, and the rest is fixed in
the soil. Therefore, P accumulates in the soil every year.
Soil erosion can release a large amount of P in streams and
lakes. Animal wastes from feedlots contains large amount of
P. Domestic sewage is an important contributor of P,
because treatment processes remove only 20-30 % of P.
Excess N originates from the same source as P. The most
stable form NO3- is highly mobile, in contrast to immobile
PO4. Agricultural regions, heavily fertilized with N, may
yield runoff containing NO3- exceeding 1 mg/L.
Groundwater beneath fertilized field can have concentration
exceeding 100 mg/L. High concentration (> 45 mg/L) of
NO3- in drinking water can cause methemoglobinemea
(blue-baby syndrome) among infants.
Domestic sewage averages about 20 mg/L of N, because
treatment processes cannot remove the majority of N.
Point source: industrial and domestic sewage, feedlots, etc.
Non-point source: agricultural runoff, etc.
Point sources are easier to control than non-point sources.
For point sources, the nutrients can be removed by passing
the effluent through a treatment pond (artificial wetland)
where aquatic plants and microbial organisms consumes the
For non-point sources, riparian buffer zone can be placed
along the river bank or lake shore, where the lush vegetation
captures suspended soil particles in runoff water and
removes the nutrients. However, a narrow buffer zone may
not be effective in removing the nutrients in groundwater,
which discharges into rivers and lakes.