Topic: Environment

Last updated: April 27, 2019

Acids donate hydrogen ions (H+) and bases accept them (Flowers et al, 2018). Buffers work to resist change in a pH solution by binding the hydrogen or hydroxide (OH-) ions (Flowers et al, 2018).

A high buffer capacity solution consists of a weak acid (HA) which does not easily donate an H+ compared to that of a strong acid, as well as its conjugate base (A-) which accepts H+ (Flowers et al, 2018). Both the acid and base need to be in approximately equal concentrations for effective buffering (Flowers et al, 2018). The idea that a buffer has the ability via its characteristics to resist pH change is important.

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Many biological systems either within humans or within the environment rely on buffers to maintain a constant pH level. For example, for optimal cellular respiration to occur, human blood needs to be maintained at pH 7, however stomachs are an acidic environment, at a pH of 3.5 (Levraut et all, 2001). In the bicarbonate buffer respiratory system, carbon dioxide combines with water and forms carbonic acid, which in turn dissociates, forming hydrogen ions and bicarbonate (Levraut et all, 2001). This system is at equilibrium until it is disturbed by an outside factor, in which case the system (i.e. reaction) must shift either left or right to resist the change (Urry et al, 2018). According to Le Châtelier’s Principle, any disturbance of a system at equilibrium is counteracted through a shift to maintain this chemical equilibrium (Flowers et al, 2018).

From the first experiment, it was observed that tap water had a reduced buffering capacity. Its ability to resist pH change to the strong acid (HCI) or strong base (NaOH) was weak, as the addition of a small amount of either resulted in a large linear change in the pH of the solution (see figure 1). The following chemical equations are used to describe what happened: HCI + H2O H3O+ + CI-Water, in this case, acted as the Brønsted-Lowry base and HCI acid, as the Brønsted-Lowry acid. HCI acid is so strong that we got almost 100% ionisation; therefore, the reaction was driven to the right and the products produced were hydronium (H3O+) and chloride (CI-). NaOH + H2O Na+ + OH- is a dissolution reaction. Water dissolved NaOH into Na+ cations and OH- anions. Water has a small acid dissociation (ka). If water has a very small concentration of H+ and OH- floating around, a high number of H2O molecules and a low Ka value, it will start to self-ionise.

An acid is required as well as its conjugate base for a buffer to work. Since water does not have these properties it acts as a very poor buffer. The next experiment tested the bicarbonate/carbonic acid buffer capacity: Na+HCO3- + HCI ? H2CO3 + NaCI. Sodium bicarbonate reacted with the HCI acid to form the products, carbonic acid and sodium chloride (spectator ions in the reaction). Results indicated in both situations that the buffer system resisted pH change when added to either a strong base or strong acid (see Figure 2). The following reaction took place: NaOH- + H2CO3+ ? NaHCO3- + H2O+ The weak carbonic acid reacted with NaOH a strong base to give the products sodium bicarbonate (its conjugate weak base) and water. Hydroxide ions on NaOH pull hydrogen ions off the bicarbonate ions, producing water and sodium bicarbonate ions (CO32-). When a strong base was added, the acid present in the buffer neutralised the hydroxide ions.

To counteract the change in solution of increasing OH- ions the equilibrium shifted to the right, as per Le Châtelier’s Principle. Another important aspect in buffering capacity was the concentration of the buffer components. The results showed that when the buffer concentration was reduced (diluted), so was the buffering capacity (see Figure 3). Some lakes protect themselves from the acid rain with their natural buffering ability (Yu et al, 2016). The results indicated that there was some natural buffering capacity, but this was reduced compared to the bicarbonate/carbonic solution. Lower concentrations of NaOH and HCI acid were used, therefore lake water had a reduced buffering capacity. This was a limited study on lake water buffering capacity.

A direct comparison could not be made due to the different concentrations used. Further studies on the buffering capacities of lake water from differing environments would make for a good experiment. The acid equilibrium reactions that have been discussed so far have focused on a group of monoprotic acids or bases (donate or accept a single H+). Polyprotic acids and bases are those molecules that can donate or accept more than one H+ in a solution. The carbonate ion is an example of a diprotic base since it can accept up to two protons.

Solutions of alkali metal carbonates are quite alkaline. The reaction between a strong acid (HCl) and a weak base (Na2CO3) yielded an acidic solution at the end. The carbonate ion, CO32-, is the conjugate base of the carbonic acid, H2CO3-, a weak acid. The reaction took place as followed: Na2CO3 + 2HCI 2NaCl + CO2 + H2O The first equivalence point was expected to be seen when double the volume of HCl acid was reached, due to the 1:2 ratio in the stoichiometric equation.

However, due to its diprotic nature, it acted as two successive Brønsted-Lowry bases and accepted two protons from the Brønsted-Lowry acid (see Figure 6). The two proton-accepting equilibria were: • Na2CO3 + HCI NaCl + NaHCO3- (equivalence point 1 at pH of about 8.1)• NaHCO3- + HCI NaCl + NaH2CO3- (equivalence point 2 at pH of about 3.8)


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