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Research Question: What is the effect of increase in temperature (℃) on the buffer capacity
(β) of borate buffer system, by measuring the volume (cm 3) and thus the amount of HCl
(mol) absorbed before pH change of 1 unit?
1: Introduction
Contact lens solutions are a must have necessity. Due to my bad eyesight, I wear contact lenses every day
and I use multi-purpose solutions (MPS) to clean and to hold my contact lenses in. Because contact lenses
make direct contact with the eyes, it is important that their pH is maintained. Due to the fact that contact
lenses are mainly made of silicone, which is sensitive to pH change, buffer solutions are crucial for the
maintenance of constant pH. Borate/boric acid buffer system is one of the most frequently used buffer
systems in maintaining the pH of multi-purpose contact lens solutions and some ophthalmic drugs as well
(Lehmann, Cavet & Richardson, 2010; Houlsby, Ghajar, & Chavez, 1986). As I was looking at the
additives section of the packaging for my new multi-purpose solution by the brand ReNu, I noticed that
two of the ingredients found were boric acid and sodium borate, which make up the majority of borate-
boric acid buffer system. This led me to investigate more about the borate buffer system as this buffer had
a direct relationship with my day to day life.
As I was learning about buffer systems, I noticed that the systems existed in an equilibrium between weak
acid or base and its conjugate base or acid. From what I had learnt about systems in equilibrium, I began
to wonder if change in temperature would affect the buffer capacity by affecting the equilibrium of the
buffer system. More specifically, I became very curious about whether the change in temperature or heat
would have an effect on the buffer capacity of borate buffer when buffering HCl. Hence, I came up with
the research question, “what is the effect of increase in temperature (℃) on the buffer capacity (β) of
borate buffer system, by measuring the volume (cm3) and thus the amount of HCl (mol) absorbed
before pH change of 1 unit?”, to see if exposure to heat or colder environments would compromise the
buffering properties of my contact lens solutions by shifting the equilibrium present within the borate
buffer system.
2: Background Information
2.1: Equilibrium in Buffers
Buffer solutions are solutions that resist the change of pH upon addition of acid or base. Buffer solutions
are able to resist change due to the fact that a weak conjugate acid-base pair is present in an equilibrium
within the solution. According to the Le Chatelier’s principle of equilibrium, any change in the
concentration of H+ during the neutralization process will cause a shift in the equilibrium to compensate
for the change. pH is a negative logarithmic value of hydrogen ion concentration, therefore change in
concentration of hydrogen ions will lead to a change in pH. Buffering activity is highly dependent on the
equilibrium of weak conjugate acid-base pair in the solution. If hydrogen ion concentration increases, the
equilibrium will shift to favor the acid rather than its conjugate base. Therefore, the conjugate base in the
solution will combine with H+ ions to become an undissociated acid, which will decrease the hydrogen
ion concentration in solution and maintain the pH. If H+ concentration decreases, the equilibrium will shift
to favor the dissociation of the acid in solution. Therefore, the acid will dissociate to create conjugate base
and more H+ ions to compensate for the loss. However, external conditions such as temperature can also
alter the equilibrium. Because buffering activity is dependent on the equilibrium action of the weak
conjugate acid-base pair, anything that may alter the equilibrium present can be assumed to alter the
buffering activity as well.
2.2: Borate Buffer System
Multipurpose solutions (MPS) are used to cleanse and disinfect contact lenses, removing dirt and protein
deposits from them and also to keep the osmolarity of the lenses similar to that of the human eye. Thus,
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MPS are a mixture of different chemical systems, including different buffer systems such as phosphate
and citrate buffers and sodium chloride. The borate buffer system is one of the most common buffer
systems used in MPS due to the disinfectant and anti-fungal agent properties of boric acid
(H3BO3/B(OH)3) and its ability to act as a buffer (Schuerer et al., 2017).
To make a brief account of the nature of borate-boric acid equilibrium relationship, it must be first noted
that boric acid ionization does not follow the typical Arrhenius acid dissociation. Boric acid ionizes by
accepting a hydroxide ion and thus an electron pair, rather than donating an H + ion, making it a Lewis
Acid (Thorsten, 2013). The information below is a summary of information taken from Thorsten (2013).
B(OH)3 + H2O ↔️ B(OH)4- + H+
The tetraborate ion yields a similar product to that of boric acid in ionizing:
B4O72- + 9H2O ↔️ 4B(OH)4- + 2H+
The primary species of borate buffer system are tetraborate and monohydrogen tetraborate ions.
Combining the two equations above, the following equation results:
4B(OH)3 ↔️ B4O72- + 5H2O + 2H+
Both B4O72- and 2H+ are aqueous, making H2B4O7 (aq), a weak acid. This tetraboric acid partially ionizes
into B4O72- and H+ ions following its second ionization.
H2B4O7 ↔️ HB4O7- + H+
HB4O7- ↔️ B4O72- + H+
According to Trejo et al. (2012), there are ten different equilibrium reactions taking place in this system.
Being one of the most complex buffer systems, borate buffer is extremely difficult to characterize
(Thorsten, 2013). In pH above 8, the tetraborate species are predominantly present in the buffer solution.
Because of the complexity of the system, a predetermined buffer recipe (Hardick, 2006) of alkalined boric
acid buffer made by a professional was used to avoid any unnecessary errors. This recipe resulted in 0.1M
borate buffer, with a pH of 9.1
The recipe follows:
Dissolve 6.18g Boric acid (61.8 g/mol) and 1.30g NaOH (40.0g/mol) in 1 dm3 of distilled water.
2.3: Definition Used to Compare Buffer Capacity
The definition of buffer capacity used in this investigation will be the amount in moles of strong acid
absorbed before change in 1 pH unit is observed within the buffer system.
This will be represented by the differential ratio published in Van Slyke’s (1922) publication. Numerical
representation of buffer capacity will make comparison between trials easier.
𝑑𝑛
𝛽=
𝑑𝑝𝐻
Where n is the amount of acid or base added in moles to 1 dm3 of buffer solution to change the pH by 1
unit. Theoretically, in a solution that has no buffer capacity, the amount of H +, thus acid, added to the
solution should have a direct effect on the pH of the solution. However, in a solution with buffering
