An investigation into the effect of UV light exposure on the concentration of
hypochlorite (OCl-) ions in pool chlorine
Background Information:
Sodium hypochlorite (NaOCl) is a chemical component of commercial bleach, and is one of the most widely
used disinfecting agents in industry. Due to its strong ability to act as an oxidizing agent, it is capable of
killing and oxidizing the cellular contents of various harmful micro-organisms and unwanted pathogens. For
this reason, sodium hypochlorite is typically found in liquid pool chlorine as a disinfectant. When in contact
with water, it typically reacts in the equation:
NaOCl (aq) + H2O (l) → HOCl (aq) + Na+ + OH- (aq) (1)
to form the weak acid HOCl (hypochlorous acid). This weak acid may then further dissociate to form its
conjugate base OCl- in the reaction:
HOCl (aq) ⇌ H+ + OCl- (aq) (2)
where OCl- are known as hypochlorite ions. Both HOCl- and OCl- are active disinfecting species of pool
water, and are constituents of ‘free chlorine’, which is the measure of chlorine in water that is available to
disinfect. In suitable pool conditions, HOCl- and OCl- typically remain in equilibrium.
However, it is also widely known that free chlorine is prone to photodegradation when exposed to sunlight. In
the presence of ultraviolet light from the sun, both HOCl - and OCl- undergo catalytic decomposition, in these
reactions:
2OCl- (aq) + UV (light) → 2Cl- (aq) + O2 (g)
2HOCl (aq) + UV (light) → 2HCl (aq) + O2 (g)
When these reactions occur, HOCl - and OCl- are decomposed to form the chloride ion and hydrochloric acid,
which are not suitable products for use as disinfectants. This adverse effect of sunlight on the stability of
aqueous free chlorine is widely recognised among scientific research and in swimming pool disinfection. 1
In the context of a swimming pool, this decomposition of free chlorine species under sunlight would result in
a lowered ability of the pool to sanitise. According to data from the NSW Ministry of Health (2010) 2, for
every hour of exposure to UV light, 35% of free chlorine content is destroyed. Particularly for outdoor
swimming pools that are at higher risks of chlorine degradation, this would significantly impact the pool’s
capacity to disinfect. Furthermore, this could result in a build-up of various pathogenic microorganisms and
algae, creating potential health and safety risks for users. Therefore, as this could pose a major issue for pool
disinfection, this led me to my research question: how does increasing the number of hours (0, 0.5, 1, 2, 4)
of exposure to UV light affect the concentration of hypochlorite (OCl -) ions in pool chlorine?
To investigate whether different hours of exposure to UV light degrade the concentration of hypochlorite
(OCl-) ions, liquid pool chlorine containing 13% sodium hypochlorite (NaOCl) will be utilised. The produced
HOCl (1) will then dissociate into OCl- ions (2), which will react in the presence of excess iodide (KI) under
acidified conditions to produce chloride ions and triiodide molecules, in this redox reaction where
hypochlorite oxidises iodide:
(3)
OCl-(aq) + 2H+(aq) + 3I-(aq) → Cl-(aq) + H2O(l) + I3-(aq)
The production of triiodide in (3) is necessary as hypochlorite is not readily titrated with thiosulfate, and the
reaction of OCl- and S2O32- would produce two colourless products (Cl- and S4O62-). The I3- molecules in (3)
1
Nowell, L. and Hoigné, J. (1992). Photolysis of aqueous chlorine at sunlight and ultraviolet wavelengths—I. Degradation
rates. Water Research, 26(5), pp.593-598.
2
Health.nsw.gov.au. (2019). Stabiliser (cyanurate) use in outdoor swimming pools - Fact sheets. [online] Available at:
https://www.health.nsw.gov.au/environment/factsheets/Pages/stabiliser-cyanurate.aspx
1
, are now able to form a triiodide-starch complex with the starch indicator and be titrated against standardised
sodium thiosulfate (Na2S2O3), in this iodometric titration where triiodide is reduced to iodide:
I3-(aq) + 2S2O32-(aq) → S4O62-(aq) + 3I-(aq) (4)
where the end-point of the titration is a colourless solution (both products in the above reaction are colourless)
when all the triiodide is used up. From equation (4), as the number of moles of thiosulfate (S 2O32-) are half that
of I3-, and as the moles of I3- are equal to the moles of OCl - (from equation 3), the concentration of
hypochlorite ions in each sample was stoichiometrically determined from the titres of sodium thiosulfate
obtained. The titre volume was calculated by recording the initial and final burette reading of sodium
thiosulfate for each trial and finding the difference. The volume of this titre is essentially an indicator of how
much OCl- ions remained in the sample, where the bigger the volume of titre, the more hypochlorite remained
(and vice versa), and so it can be used to compare the effects of varying UV light exposure on liquid chlorine
samples. In this experiment, the original sample of liquid chlorine was also diluted in order for all the triiodide
molecules produced to stay in solution, and so it could be titrated with standardised sodium thiosulfate as it
was initially too concentrated.
Research Question: How does increasing the number of hours (0, 0.5, 1, 2, 4) of exposure to UV light
affect the concentration of hypochlorite (OCl -) ions in pool chlorine?
Personal Engagement:
Personally, I am a swimmer and have been swimming approximately 3 times a week for the past 4 years.
More than a few times, I’ve wondered exactly what chemicals I’m swimming in while I train, and exactly how
much bacteria and micro-organisms there were in the pool, as well as how effective chlorine was as a
disinfecting agent in killing such bacteria. Furthermore, especially when swimming in an outdoor pool, I’ve
also been curious about the levels of chlorine in the pool under direct exposure to this sunlight, and whether
this impacts the ability of chlorine to disinfect pool water, as this has various health and safety implications to
swimmers such as myself. Upon consultation with my teacher, I found that the determination of chlorine
concentration in pool chlorine could be done via titration with acid-base and redox chemistry. After doing
some research, I found the real disinfecting species of free chlorine to be hypochlorous acid (HOCl) and
hypochlorite (OCl-) ions. I then decided to focus on hypochlorite ions and investigate exactly how exposure to
sunlight (UV light) would impact its concentration in pool chlorine, and whether it would be degraded by
prolonged exposure to light, resulting in a lowered ability of the pool water to sanitise. This ultimately led to
my research question of how does increasing the number of hours (0, 0.5, 1, 2, 4) of exposure to UV light
affect the concentration of hypochlorite (OCl-) ions in pool chlorine?
Aim:
The aim of this experiment is to determine whether an increased number of hours of exposure to UV light will
affect and degrade the concentration of hypochlorite ions in liquid pool chlorine. This will be determined via
an iodometric titration, with the concentration of hypochlorite ions calculated stoichiometrically from the
titres of sodium thiosulfate obtained.
Variables:
Independent variable:
The independent variable (IV) in this experiment was the duration of the exposure of the samples to the UV
light. UV lamps (each 40W) were used with 5 treatment levels of 0 minutes (a control), 30 minutes, 1 hour, 2
hours and 4 hours. Time was chosen as the IV as concentrations of hypochlorite are said to decrease with
increased duration of exposure to UV light. These intervals of time were chosen as they were fairly evenly
spaced, in order to hopefully observe a steady trend between treatments of the exposure to UV with gradually
degrading hypochlorite concentrations.
Dependent variable:
The dependent variable for this experiment was the concentration of hypochlorite (OCl -) ions in the samples
of liquid pool chlorine after varying hours of exposure to UV light, as measured by the titres of sodium
thiosulfate (Na2S2O3) obtained.
2
hypochlorite (OCl-) ions in pool chlorine
Background Information:
Sodium hypochlorite (NaOCl) is a chemical component of commercial bleach, and is one of the most widely
used disinfecting agents in industry. Due to its strong ability to act as an oxidizing agent, it is capable of
killing and oxidizing the cellular contents of various harmful micro-organisms and unwanted pathogens. For
this reason, sodium hypochlorite is typically found in liquid pool chlorine as a disinfectant. When in contact
with water, it typically reacts in the equation:
NaOCl (aq) + H2O (l) → HOCl (aq) + Na+ + OH- (aq) (1)
to form the weak acid HOCl (hypochlorous acid). This weak acid may then further dissociate to form its
conjugate base OCl- in the reaction:
HOCl (aq) ⇌ H+ + OCl- (aq) (2)
where OCl- are known as hypochlorite ions. Both HOCl- and OCl- are active disinfecting species of pool
water, and are constituents of ‘free chlorine’, which is the measure of chlorine in water that is available to
disinfect. In suitable pool conditions, HOCl- and OCl- typically remain in equilibrium.
However, it is also widely known that free chlorine is prone to photodegradation when exposed to sunlight. In
the presence of ultraviolet light from the sun, both HOCl - and OCl- undergo catalytic decomposition, in these
reactions:
2OCl- (aq) + UV (light) → 2Cl- (aq) + O2 (g)
2HOCl (aq) + UV (light) → 2HCl (aq) + O2 (g)
When these reactions occur, HOCl - and OCl- are decomposed to form the chloride ion and hydrochloric acid,
which are not suitable products for use as disinfectants. This adverse effect of sunlight on the stability of
aqueous free chlorine is widely recognised among scientific research and in swimming pool disinfection. 1
In the context of a swimming pool, this decomposition of free chlorine species under sunlight would result in
a lowered ability of the pool to sanitise. According to data from the NSW Ministry of Health (2010) 2, for
every hour of exposure to UV light, 35% of free chlorine content is destroyed. Particularly for outdoor
swimming pools that are at higher risks of chlorine degradation, this would significantly impact the pool’s
capacity to disinfect. Furthermore, this could result in a build-up of various pathogenic microorganisms and
algae, creating potential health and safety risks for users. Therefore, as this could pose a major issue for pool
disinfection, this led me to my research question: how does increasing the number of hours (0, 0.5, 1, 2, 4)
of exposure to UV light affect the concentration of hypochlorite (OCl -) ions in pool chlorine?
To investigate whether different hours of exposure to UV light degrade the concentration of hypochlorite
(OCl-) ions, liquid pool chlorine containing 13% sodium hypochlorite (NaOCl) will be utilised. The produced
HOCl (1) will then dissociate into OCl- ions (2), which will react in the presence of excess iodide (KI) under
acidified conditions to produce chloride ions and triiodide molecules, in this redox reaction where
hypochlorite oxidises iodide:
(3)
OCl-(aq) + 2H+(aq) + 3I-(aq) → Cl-(aq) + H2O(l) + I3-(aq)
The production of triiodide in (3) is necessary as hypochlorite is not readily titrated with thiosulfate, and the
reaction of OCl- and S2O32- would produce two colourless products (Cl- and S4O62-). The I3- molecules in (3)
1
Nowell, L. and Hoigné, J. (1992). Photolysis of aqueous chlorine at sunlight and ultraviolet wavelengths—I. Degradation
rates. Water Research, 26(5), pp.593-598.
2
Health.nsw.gov.au. (2019). Stabiliser (cyanurate) use in outdoor swimming pools - Fact sheets. [online] Available at:
https://www.health.nsw.gov.au/environment/factsheets/Pages/stabiliser-cyanurate.aspx
1
, are now able to form a triiodide-starch complex with the starch indicator and be titrated against standardised
sodium thiosulfate (Na2S2O3), in this iodometric titration where triiodide is reduced to iodide:
I3-(aq) + 2S2O32-(aq) → S4O62-(aq) + 3I-(aq) (4)
where the end-point of the titration is a colourless solution (both products in the above reaction are colourless)
when all the triiodide is used up. From equation (4), as the number of moles of thiosulfate (S 2O32-) are half that
of I3-, and as the moles of I3- are equal to the moles of OCl - (from equation 3), the concentration of
hypochlorite ions in each sample was stoichiometrically determined from the titres of sodium thiosulfate
obtained. The titre volume was calculated by recording the initial and final burette reading of sodium
thiosulfate for each trial and finding the difference. The volume of this titre is essentially an indicator of how
much OCl- ions remained in the sample, where the bigger the volume of titre, the more hypochlorite remained
(and vice versa), and so it can be used to compare the effects of varying UV light exposure on liquid chlorine
samples. In this experiment, the original sample of liquid chlorine was also diluted in order for all the triiodide
molecules produced to stay in solution, and so it could be titrated with standardised sodium thiosulfate as it
was initially too concentrated.
Research Question: How does increasing the number of hours (0, 0.5, 1, 2, 4) of exposure to UV light
affect the concentration of hypochlorite (OCl -) ions in pool chlorine?
Personal Engagement:
Personally, I am a swimmer and have been swimming approximately 3 times a week for the past 4 years.
More than a few times, I’ve wondered exactly what chemicals I’m swimming in while I train, and exactly how
much bacteria and micro-organisms there were in the pool, as well as how effective chlorine was as a
disinfecting agent in killing such bacteria. Furthermore, especially when swimming in an outdoor pool, I’ve
also been curious about the levels of chlorine in the pool under direct exposure to this sunlight, and whether
this impacts the ability of chlorine to disinfect pool water, as this has various health and safety implications to
swimmers such as myself. Upon consultation with my teacher, I found that the determination of chlorine
concentration in pool chlorine could be done via titration with acid-base and redox chemistry. After doing
some research, I found the real disinfecting species of free chlorine to be hypochlorous acid (HOCl) and
hypochlorite (OCl-) ions. I then decided to focus on hypochlorite ions and investigate exactly how exposure to
sunlight (UV light) would impact its concentration in pool chlorine, and whether it would be degraded by
prolonged exposure to light, resulting in a lowered ability of the pool water to sanitise. This ultimately led to
my research question of how does increasing the number of hours (0, 0.5, 1, 2, 4) of exposure to UV light
affect the concentration of hypochlorite (OCl-) ions in pool chlorine?
Aim:
The aim of this experiment is to determine whether an increased number of hours of exposure to UV light will
affect and degrade the concentration of hypochlorite ions in liquid pool chlorine. This will be determined via
an iodometric titration, with the concentration of hypochlorite ions calculated stoichiometrically from the
titres of sodium thiosulfate obtained.
Variables:
Independent variable:
The independent variable (IV) in this experiment was the duration of the exposure of the samples to the UV
light. UV lamps (each 40W) were used with 5 treatment levels of 0 minutes (a control), 30 minutes, 1 hour, 2
hours and 4 hours. Time was chosen as the IV as concentrations of hypochlorite are said to decrease with
increased duration of exposure to UV light. These intervals of time were chosen as they were fairly evenly
spaced, in order to hopefully observe a steady trend between treatments of the exposure to UV with gradually
degrading hypochlorite concentrations.
Dependent variable:
The dependent variable for this experiment was the concentration of hypochlorite (OCl -) ions in the samples
of liquid pool chlorine after varying hours of exposure to UV light, as measured by the titres of sodium
thiosulfate (Na2S2O3) obtained.
2
