Acid Base Buffer Systems

ACID / BASE BUFFER SYSTEMS

Abstract
A buffer solution is a solution that contains both an acid and a salt containing the conjugate base / acid in sufficient concentrations so as to maintain a relatively constant pH when either acid or base is added. In this experiment a selection of buffer solutions (Bicarbonate/carbonic acid), Lake water and distilled water were obtained to compare their buffering behaviours when mixed both with an acid and a base. The results showed buffering capacities for all the solutions except distilled water.

Introduction
Many lab experiments require constant pH while acids or bases are added to solutions either by reaction or by the experimenter, buffers are required to do this (Timberlake 2010). Chemists use buffers to balance the pH of a reaction. Biology finds many uses for buffers, which range from controlling blood pH to ensuring that urine does not reach strongly acidic levels (Timberlake 2010).
A buffer is a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. Buffers work by reacting with an acid or base to control the pH. For this experiment, a buffer solution of bicarbonate (weak base), and its (conjugate acid) Carbonic acid was used. The purpose of the experiment was to test the capacity of an undiluted and diluted buffer solution as well as examining the buffering capacity of distilled water using measured concentrations of NaOH and HCL.
These replacements of strong acids and bases for weaker ones give buffers their ability to moderate pH. (Stoker 2013). Part 2 of the experiment, the buffering capacity of lake water was tested. The ions naturally present in rivers are buffering components that allow the pH of the water to remain stable over time. Buffer capacity of river water is very important, usually necessitating narrow pH ranges that are critical to the survival of most organisms. If the buffer capacity of river water is too small or the pH of the water is outside its buffer range, it can be lethal to the river’s ecosystem.

Materials and methods
As per pages 28 – 33 SCI105 Chemistry Practical Manual 201
Results
Part 1: Distilled Water | | Effect of addition of 0.05M NaOH | Effect of addition of 0.05M HCl | mL of 0.05M NaOH added | pH | mL of 0.05M HCl | pH | 0.0 | 6.99 | 0.0 | 6.99 | 1.0 | 11.62 | 1.0 | 2.55 | | | | | Undiluted Buffer Solution | Effect of addition of 0.05M NaOH | Effect of addition of 0.05M HCl | mL of 0.05M NaOH added | pH | mL of 0.05M HCl | pH | 0.0 | 6.78 | 0.0 | 6.90 | 1.0 | 7.03 | 1.0 | 6.67 | 2.0 | 7.47 | 2.0 | 6.45 | 3.0 | 8.43 | 3.0 | 6.25 | 4.0 | 9.32 | 4.0 | 6.01 | 5.0 | 9.68 | 5.0 | 5.72 | 6.0 | 9.90 | 6.0 | 5.17 |

Diluted Buffer Solution | Effect of addition of 0.05M NaOH | Effect of addition of 0.05M HCl | mL of 0.05M NaOH added | pH | mL of 0.05M HCl | pH | 0.0 | 7.07 | 0.0 | 7.35 | 1.0 | 9.03 | 1.0 | 6.62 | 2.0 | 9.91 | 2.0 | 6.10 | 3.0 | 10.30 | 3.0 | 5.23 | 4.0 | 10.63 | 4.0 | 2.64 | 5.0 | 10.97 | 5.0 | 2.30 | 6.0 | 11.37 | 6.0 | 2.12 | | | | | Buffering Capacity of Lake Water | Effect of addition of 0.05M NaOH | Effect of addition of 0.05M HCl | mL of 0.05M NaOH added | pH | mL of 0.05M HCl | pH | 0.0 | 7.38 | 0.0 | 8.14 | 1.0 | 9.72 | 1.0 | 7.03 | 2.0 | 10.17 | 2.0 | 6.43 | 3.0 | 10.44 | 3.0 | 5.62 | 4.0 | 10.64 | 4.0 | 3.68 | 5.0 | 10.81 | 5.0 | 3.24 | 6.0 | 10.95 | 6.0 | 3.05 |

Figure 1: Buffering Capactities of undiluted / diluted buffer solutions and distilled water.

Part 2:

Titration Curve of Carbonate Ion | mL of 0.025M HCl added | pH | 0.0 | 11.32 | 1.0 | 11.02 | 2.0 | 10.79 | 3.0 | 10.61 | 4.0 | 10.44 | 5.0 | 10.30 | 6.0 | 10.12 | 7.0 | 9.94 | 8.0 | 9.73 | 9.0 | 9.43 | 10.0 | 8.72 | 11.0 | 7.60 | 12.0 | 7.18 | 13.0 | 6.94 | 14.0 | 6.74 | 15.0 | 6.57 | 16.0 | 6.37 | 17.0 | 6.22 | 18.0 | 6.03 | 19.0 | 5.83 | 20.0 | 5.52 | 21.0 | 4.82 | 22.0 | 3.15 | 23.0 | 2.77 | 24.0 | 2.58 | 25.0 | 2.35 |

CO32-/HCO3-

Figure 2: Titration curve of carbonate ion

Discussion
The results in figure 1 clearly indicate the different buffering capacities of the different solutions. Distilled H2O showed no buffering effects with the addition of 0.05M NaOH and 0.05 HCL, this is because distilled water has roughly a neutral pH, as well as zero buffering capacity, which could be assumed by the extreme swing of the pH quickly in either direction. The gradual decline and incline in pH of the buffer solution HCO3/H2CO3 showed that it was able to neutralize small amounts of the added NaOH and HCL, thus maintaining a relatively constant pH of the solution. Buffer solutions have a working pH range and dictate how much acid/base can be neutralized to keep pH at a constant, and to the amount by which it will change (Stoker 2012). Furthermore once the buffering capacity of the solution is exceeded it could be assumed that rate of pH change would be more dramatic, because the conjugate acid or base has been depleted through neutralization. Therefore a larger amount of conjugate acid or base will have a greater buffering capacity.
The results also showed the effects of diluting the buffer solution with distilled water. The graph showed less absorbent buffering affects demonstrated by a more dramatic shift of pH. Through diluting a solution it could be assumed to decrease the concentration of the ions in the solution, therefore a lesser ability to neutralize the effects of adding more H30+ and –OH.

The results also provided the equivalence points for the alkaline (carbonate) The Titration curve of Carbonate Ion showed 2 distinct equivalence points. The first proton being released at pH 9, second proton being released at 5 pH. This indicated the acid Neutralizing Capacity / buffering capacity. By gradually adding more HCL to the system, the supply of buffering ions (HCO3) in the sample were exhausted, the endpoint showing a complete conversion of chemical species in solution.

In conclusion the experiment demonstrated the importance of buffering systems and the capability of lakes and river systems to neutralize acid. Without a buffer solution to absorb the changing concentration of H+ ions, water systems would have fluctuating pH levels therefore causing devastating affects on the already fragile marine ecosystem.