Numerous solutions with a pH of less than
Numerous chemical reactions, are affected by acidity or basicity they exist in, thus the ability to maintain a stable pH is crucial (Al-Nidway, 2018). pH is a measure of the concentration of hydrogen ions, which quantitatively measures the acidity or basicity of a solution (Britannica, 2018). pH is measured on a scale that ranges from 0 to 14. At 25ºC, solutions with a pH greater than seven are considered basic, solution with the pH of seven are considered neutral and solutions with a pH of less than seven are considered acidic. Solutions with the pH of seven are considered neutral because, the concentration of OH- equals the concentration of H3O+ in pure water (Britannica, 2018). The equation for calculating pH is, pH = -logH3O +, where log is the base 10 logarithm and H3O+ stands for the concentration of hydronium ion in moles per litre (molarity) (Britannica, 2018).
A buffer is a solution that consists of a weak base and its conjugate acid or a weak acid and its conjugate base (Al-Nidway, 2018). A buffer helps to maintain the pH of solution upon the addition of small amounts of acid or base; the buffer zone is when the addition of base or acid does not affect the pH of the solution drastically (Buffer Effectiveness, n.d.). This process is very important, especially biological systems as they function optimally at a specific pH range (Al-Nidway, 2018). A buffer ‘s ability to resist change in pH is measured by the buffer capacity, which is defined as the amount of acid or base required to change the pH of one litre of a buffer by one unit (Buffer Effectiveness, n.d.). The equation for buffer capacity is:
A buffer is most effective when the pH is equal to the pKa, however, the pH range in which a buffer is effective is equal to the pKa ±1 (Al-Nidway, 2018).
An important biological buffer system that operates in the internal fluid of all cells is the dihydrogen phosphate system (Al-Nidway, 2018). The dihydrogen phosphate system consists of two ions “dihydrogen phosphate ions (H2PO4-) as hydrogen-ion donors (weak acid) and hydrogen phosphate ions (HPO 4 2- ) as hydrogen-ion acceptors (conjugate base)” which are in equilibrium together as demonstrated in the equation:
H2PO4- (aq) H + (aq) + HPO42- (aq) (Al-Nidway, 2018)
If additional hydrogen ions enter the cell, the equilibrium shifts to the left as the ions react with H2PO4- and are consumed (Voet, 2016). However, if additional hydroxide ions enter the system, the equilibrium shifts to the right as the ions react with H2PO4- producing more HPO42- (Voet, 2016).
The acid dissociation constant “Ka” is the equilibrium constant of the dissociation reaction of an acid: it measures the strength of an acid in a solution and is expressed in mol/L (Libretexts, 2017). The equilibrium-constant expression for this equilibrium is:
The value of Ka for the equilibrium used in this experiment is 6.23 x 10-8 at 25ºC (Voet, 2016). The Ka value can be used to measure the position of equilibrium; if the Ka value is large, the formation of the products favored (Voet, 2016). Whereas if the Ka value is small, the undissolved acid is favored (Voet, 2016). The pKa is a value which is related to the Ka, as pKa is the negative logarithm of the acid dissociation constant, Ka and is defined as
pKa = -log10Ka (Libretexts, 2017).
The pKa value is a measure of the strength of an acid or base; if the pKa is small (large Ka) most of the acid has dissociated, and therefore the acid is strong (Libretexts, 2017). However, if the pKa is large (small Ka) a small amount of dissociation has occurred, and therefore the acid is weak (Libretexts, 2017).
The Henderson-Hasselbalch equation:
pH = pKa + log
Given the concentration of the acid- base and the pKa, one can use the Henderson-Hasselbalch equation to approximate the pH of a buffer (Khaira, Khot, 2017). Therefore, when the concentrations of and are equivalent, the “value of the molar concentration of hydrogen ions is equal to the value of the equilibrium constant, and the pH is equal to the pK a , namely 7.21” (Khaira, Khot, 2017).
As Amino acids carry both a basic group (NH3+) and an acidic group(COOH), it naturally act as buffers (Khaira, Khot, 2017). In aqueous solution, amino acids largely exist as dipolar or “zwitter” ions (Khaira, Khot, 2017). In a basic solution the carboxyl group will be deprotonated, whereas in an acidic solution the amino group will be protonated (Al-Nidway, 2018). However, at the isoelectric point, the basic groups become protonated and the acidic groups deprotonate (Al-Nidway, 2018). The isoelectric point (pI), is the pH at which the net electric point is zero, therefore the zwitterion form is dominant and is given by the following equation:
pI = (Al-Nidway, 2018)
Where, the pKas are equivalent to the pH at which the ionizable group is at its best buffering capacity (Al-Nidway, 2018). The titration curve of an amino acid reveals its buffering range and the pKa values of its ionizable groups (Al-Nidway, 2018). Due to its ability to be amphoteric and it’s three ionizable groups; amino, carboxyl chain and imidazole side chain, L- histidine acts a buffer (Khaira, Khot, 2017).