Chemical Technology • July 2013
s an industrial cooling water treatment option,
ozone has started to gain popularity and could
take over from conventional chemical treatment
methods. In companies such as Air Products, its use is
also seen as a corrosion inhibitor, biocide, dispersant
and pH control measure. However, there is still a poor
understanding of the chemistry and action of ozone as a
water treatment option.
It is required that the ozone chemistry be better un-
derstood. This is so that companies such as Air Products
can minimize ozone use while still obtaining the desired
water quality (but not negatively impacting on other
Ozone is a strong oxidant and a potent disinfecting agent
(Guzel-Seydim et al, 2004). When ozone decomposes
in water, free radicals such as hydrogen peroxy (HO2)
and hydroxyl (OH) are formed. These free radicals have
a large oxidizing capacity and are able to play a role in
disinfection processes (EPA, 1999). Employing ozone as
a primary oxidant before chlorination will usually satisfy
most of the oxidant demands of water being treated,
thus lowering the subsequent demand for chlorine and
minimizing the disinfection by-products of chlorination
(Alsheyab and Munoz, 2007).
Ozone disinfection is generally used at medium to
large sized plants after at least secondary treatment
(EPA, 1999). Although ozone treatment is able to achieve
higher levels of disinfection compared to its competi-
tors, it is often used sparingly. This is because ozone
treatment as a disinfection option tends to have higher
maintenance expenditure and capital costs as compared
to its competitors.
ozone in water
by K G Harding, R Ntimbani, P Mashwama, R Mokale, M Mothapo, N Gina, all of the School of Chemical and
Metallurgical Engineering, University of the Witwatersrand, Johannesburg, South Africa, and D Gina, Air Products
South Africa, Kempton Park. N Gina is also with Air Products South Africa.
There are several advantages and disadvantages of us-
ing ozone as an oxidant or disinfectant as opposed to using
other methods (including chlorination and UV methods).
These are explained by EPA (1999) as tabulated below.
Ozone effectiveness against micro-organisms depends
not only on the amount applied, but also on the residual
ozone in the medium (which is affected by the instability
of ozone and the presence of ozone consuming materi-
als). Because the toxicity of ozone varies depending on
the concentration and the length of exposure, it is impor-
tant to monitor ozone when it is used as a disinfectant.
From studies undertaken for surface disinfection using
ozone, it was shown that moderate doses of ozone, be-
tween 0,5 ppm and 3,5 ppm, both in gaseous form and
as ozonated water, are sufﬁ cient to achieve signiﬁ cant
microbial reductions (Pascual, 2007).
There are several mechanisms that ozone uses when
disinfection takes place (EPA, 1999):
• The direct oxidation of the cell wall with leakage of
cellular constituents outside of the cell;
• The reaction with radical by-products of ozone decom-
• The damage to the constituents of the nucleic acids
(purines and pyrimidines); and
• The breakage of carbon-nitrogen bonds which then
leads to depolymerization.
The half-life of a reaction, t1/2, is the time required
for the concentration of a reactant to reach one-half of
its initial value. The half-life concentration of ozone can
be found by using Equation 1 below. The half-life is a
convenient way to describe how fast a reaction occurs,
especially if it is a ﬁ rst-order process (Brown et al, 2006).
The half-life of a ﬁ rst order reaction can be determined
Ozone is very effective in destroying viruses and bacteria,
even more so than chlorine.
Low dosage may not effectively inactivate some viruses,
spores, and cysts. This would mean higher doses, and
therefore higher costs.
The ozonation process uses only a short contact time.
Ozonation requires more complicated equipment and efﬁ cient
contacting systems since ozonation is a more complex
technology than chlorine or UV disinfection.
After ozonation, there is minimal re-growth of micro-
Ozonation is not economical for wastewater with high levels of
suspended solids, chemical oxygen demand or total organic
Ozone is generated onsite, so there are fewer risks associated
with transportation and handling.
The cost of water treatment can be relatively high in capital
and in power intensiveness when using ozone rather than UV
Table 1: Advantages and disadvantages of using ozone (EPA, 1999)
Chemical Technology • July 2013
uation 2 below
rovided that the rate constant is
sure ozone in water sa
les. These include colorimetric
est kits and spectrophotometry. These methods may not
e able to measure low levels of ozone and are unsuit
able for continuous in-ﬂ owing monitoring and in tha
case, Oxidation-Reduction Potential (ORP) measurement
s used as an alternative
SW Government, 2013).
The objectives of this research were to investigate the
effect of both temperature and pH on the solubility/de
composition of ozone in water as well as determine the
alf-life of ozone.
t the desired conditions
a beaker with ozone from an ozone generator bubbled
hrough it. For both pH and temperature experiments, an
meter was used as
easure of the ozone concentration. Sets of ex
ere performed for pH values ranging from 9-12,3 and
emperatures from 25-4
Sodium hydroxide was used to
re solutions with
values 9, 10, 11 and 12,3. Alkaline
H values were
chosen as it has been observed that at a pH less than
, the rate of ozone decomposition is relatively insensi
tive. A water bath was used to control the temperature at
values of 25, 35, 40 and 45°C.
After the ozone was bubbled in the beaker for 10
minutes, it was assumed a saturated concentration was
obtained. The redox potential was measured and the
results (Eh in millivolts) were recorded every 30 sec-
onds. Experiments were repeated for error analyses and
to check for precision in the results obtained (average
values shown in results below).
Results and discussion
Ozone solubility as a function of pH
As was previously mentioned, Eh shows a direct relation-
ship to ozone concentration and hence it was plotted as
function of time. Figure 2 below shows normalized curves
of Eh (in mV) versus time (in minutes) at different pHs.
Normalized data was used as a comparative measure
since all start values varied. The graphs generally showed
a slight decrease in Eh as time elapsed.
Figure 1: Experimental setup
Figure 2: Normalized Eh results as a function of time for different
Chemical Technology • July 2013
The trends of the plots do not show proper correlation
with the theory. Firstly, values obtained were all negative
and according to Metzger (2009) this could be due to the
interference of high turbidity in the water and the strong
oxidizing power of ozone that can cause ORP readings to
be below what is expected (potentially, even a reducing or
negative value). Suslow (2004) explained that some sys-
tems with moderate turbidity could result in ORP values
far below the expected values and even negative values
which was seen in the data obtained.
One of the reasons for the increase in Eh value as
time elapses could be due to the formation of O3- which
reacts with H2O2 or OH- radicals that are produced in the
ozone decomposition, reforming ozone and thus increas-
ing the ozone concentration with time. The presence of
chemicals in the solution can affect the decomposition
rate of ozone by acting as a promoters and re-generating
O2- from the hydroxyl radical that decomposed and thus
reforming ozone (Erikkson, 2005).
This unexpected behaviour could also be due to the
fact that precise measurement of ozone in water can be
problematic because of potential interferences, unpre-
dictable half-life in water with different pH values and the
time between sample collection and analysis.
Ozone solubility as a function of temperature
The results for normalized Eh (in mV) versus time (in
minutes) at different temperatures is given in Figure 3.
The graph shows a decrease in Eh as time elapses. At
50°C the curve lies below the other temperatures as
drawn (45°C, 40°C, 35°C and 25°C respectively).
The trends of the graphs are not as expected as
they do not correlate well with theory. Again the results
depicted a negative Eh (corresponding to a reducing
rather than oxidising capacity of ozone as explained by
Suslow, 2004). Trends in the ORP readings can be con-
fusing as can be explained by von Gunten (2003): Ozone
decomposes into OH radicals (dOH) which are very strong
oxidizing agents in water and therefore when assessing
ozonation processes, two species; ozone and OH radicals
are always involved.
Since the OH radicals that ozone decomposes into,
are also oxidizing agents, it would seem that the ORP me-
ter measures the increase in OH radicals in the system
as opposed to the decreasing ozone. This is then shown
as an increase in Eh as time elapses. Hydrogen peroxy
(HO2) that may also be produced in ozone decomposition
inﬂ uences the results the same way that OH radicals
do. Another reason for an increase in Eh value with time
could be due to the formation of O3- which reacts with OH
radicals produced in ozone decomposition; this reaction
leads to the reformation of ozone and thus increases the
ozone concentration with time.
Ozone half-life as a function of temperature
The half-life of ozone was estimated by reading the time
to reach 50% of the original ozone concentration. From
Figure 4, the half-life at 50°C is about 17,9 seconds (0,3
minutes).The half-life obtained at 35°C, 40°C and 45°C
was fairly similar at 28,8, 27,6 and 17,6 seconds respec-
tively. At lower temperatures of 25 and 30°C, the half-live
was higher at values of 330 and 220 seconds respectively.
The half-life decreased as the water temperature in-
creased, which is in good agreement with literature. This
meant that the half-life of ozone is longer at high tem-
peratures and it shortens as temperature is increased.
This is since at lower temperatures gas molecules do not
move quickly and therefore collide less frequently, result-
ing in a slow ozone decomposition, while at higher tem-
peratures ozone molecules move faster and collide more
frequently, thus the rapid decomposition and reduction in
ozone concentration in the solution.
From this data, the decay constant can be calculated
for ozone. For these experiments, they were determined
to be: 0,0021, 0,0058, 0,0239, 0,0257, 0,0394 and
0,0387sec-1 for the temperatures from 25 to 50°C at 5
degree increments. As expected, values increase with
Figure 3: Graph of Eh versus time for normalized temperature results
Figure 4: Graph of ozone half-life versus temperature
Chemical Technology • July 2013
As temperature increases, the amount of ozone that
can be absorbed into water decreases, therefore the
solubility of ozone decreases with increasing tempera-
tures. An ORP meter is a good way of indirectly measur-
ing the concentration of ozone in water if all conditions
are kept constant in each run and there is no distur-
bance on the system.
Figure 5: Graph of ozone half-life versus pH.
NOTE: pH 7 half-life excluded from the trend lin
Ozone half-life as a function of pH
Ozone was dissolved in water with pH 1, 3, 5,3, 7, 9
and 11 at a constant temperature of 25°C. The half-
life was then calculated for each and is plotted
in Figure 5.
Low pHs were achieved by addition of sulphuric acid
into deionized water and high pHs were achieved by
addition of sodium hydroxide into deionized water. It
was assumed that any variations in redox potential of
the ozonated solutions after dissolving ozone in water
are due to ozone decomposition.
Compared to Figure 4, half-lives are much shorter
across the range and do not drop off as sharply. At a
pH of 1, the half-life is 90 seconds while at pH 6 the
half-life is 6 seconds. At a pH of 7, an unexpected
value (when compared to other data) of 210 seconds
At pH 7, only hydroxyl and hydrogen ions exist in equi-
librium with each other. Therefore, ozone decomposition
is enhanced by the catalytic effect of other ions such as
sodium and sulphate ions from sodium hydroxide. This
could explain why a longer half-life was obtained at pH 7.
The solubility and decomposition rate of ozone in water
is affected by factors such as pH and temperature. The
pH of the water is important because hydroxide ions
initiate ozone decomposition. At pH values above 7,5,
much of the ozone will decompose into hydroxyl radicals
and will react rapidly with water contaminants. The
concentration of ozone in the solutions decreases as
For the results obtained, Eh values were negative
and they increased as time elapsed. This was due to the
formation of O3- which reacts with HO2 or OH- radicals
that are produced in the ozone decomposition. This
resulted in reformation of ozone and thus increasing the
ozone concentration with time.
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