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Using Titration to Examine Changes in pH with Hydrochloric Acid and Sodium Hydroxide

By onegreatninja Feb 15, 2014 1017 Words
Using Titration to Examine Changes in pH with Hydrochloric Acid and Sodium Hydroxide

Examine the buffering capacity of H2O, acetic acid, and acetate when titrated with NaOH and HCL. A good buffer will somewhat resist a great change in pH. It is common knowledge that water is not a good buffer, and this was confirmed in this experiment. Acetate demonstrated an excellent buffering capacity because it took a large volume of the acid or base to make a significant change in the pH. Acetic acid, while already having a low pH, was an adequate buffer against NaOH since it took a large volume to raise the pH. Introduction

Examining which solutions make good buffers, a basic understanding of these characteristics is required. Titration is a method of analysis in which a solution of a known concentration is slowly added to a known volume of solution with an unknown concentration until it reaches its equivalence point. Using this method, experimenters can calculate the concentration of the unknown solution.1 A buffer is a weak acid and its base. According to the Brønsted-Lowry acid-base theory, a base accepts protons, while an acid donates protons. In an aqueous solution, the acid donates its proton to form its conjugate base and is a reversible reaction. Most buffers work within a pH range of 6 to 8.2 When equilibrium in this reaction is reached, an ionization constant (Ka) can be used to determine strength of the acid. A weak Ka would be indicative of a weak acid like acetic acid for example. Weak acids do not have a lower tendency to detach a proton, so its pKa is higher. These two values are related as pKa=-log(Ka). When hydrogen or hydroxide ions are added to a buffer, the ratio of these ions are changed, resulting in a pH change. The Henderseon-Hasselbalch equation can be used, when using a weak acid, to relate the pKa, pH, and buffer concentrations. This equation is written as pH=pKa + log([Base])/([Acid]) Materials

The materials for this lab were provided. The 2 N NaOH and 2 N HCL soultions were prepared prior to this experiment. Other materials required were a pH probe, assorted flasks, burettes, H2O, Acetate, and Acetic Acid. Methods

The first burette was filled with 25mL of 2 N NaOH, and the second with 25mL of 2 N HCL. H2O, Acetate, and Acetic Acid was titrated with each of these with the exception of Acetic Acid and HCL since they are both acids. Initial and Final pHs and volumes were recorded until HCL reached a pH of at least 3.5 and NaOH reached a pH of at least 9.5.3 Results

This experiment began with calculating a standard for volume which was mLs/drops and was found to be 0.06mLs per drop. Then, the pH meter was calibrated to 7. The first titration was done with NaOH and H2O. The initial pH was 5.28 and the final was 11.6 after 0.7mLs of NaOH was added as shown in table 1 below. H2O/NaOHpHDropsmLs

Table 1: H2O Titration with NaOH
The second titration was NaOH with Acetate. The initial pH of the Acetate was 4.49 and the final pH was 11.43 after 2mLs of NaOH was added as shown in table 2 below. Acetate/NaOHpHDropsmLs
Table 2: Acetate Titration with NaOH
The next titration was Acetic Acid with NaOH. The initial pH of Acetic Acid was 2.92 and the final was 12.3 after the addition of 4.2mLs of NaOH as shown in table 3 below. Acetic Acid/NaOHpHDropsmLs

Table 3: Acetic Acid Titration with NaOH
The fourth titrations was done with H2O and HCL. The initial pH of the H2O was 5.26 and the final pH was 2.08 after 0.6mLs of H2O was added as shown below in table 4. H2O/HCLpHDropsmLs
Table 4: H2O Titration with HCL
The final titration was HCL with Acetate. The initial pH of Acetate was 4.45 and the final pH was 2.22 after the addition of 1.3mLs of HCL as shown below in table 5. Acetate/HCLpHDropsmLs
Table 5: Acetate Titration with HCL
Figure 1 below shows H2O, Acetate Buffer and Acetic Acid titrated with NaOH.
Figure 1: NaOH Titration

Figure 2 below shows H2O and Acetate Buffer titrated with HCL.
Figure2: HCL Titration
Then, a phosphate buffer was prepared using the calculations that were given in the lecture3 pH = pKa + log([A-]/[HA])
7.0 = 6.91 + log([A-]/[HA])
0.09 = log([A-]/[HA])
10.00.09 = ([A-]/[HA])
1.23 = [A-]/[HA]
If [HA] 1, then [A-] = 1.23
[HA] + [A-] = 1 + 1.23 = 2.23
Fraction of acid = [HA]/([HA][A-])
Fraction of acid = 1/(1+1.23)
Fraction of acid = 1/2.34
Fraction of acid = 0.45
Fraction of base = [A-]/([HA][A-])
Fraction of base = 1.23/(1+1.23)
Fraction of base = 1.23/2.23
Fraction of base = 0.55
Acid form: NaH2PO4 FW = 138 g/mole
Base form: Na2HPO4 FW = 268 g/mole
g(acid) = (138 g/mol)(0.05 mol/L)(0.45)(0.1L)
g(acid) = 0.31 g
g(base) = (268 g/mol)(0.05 mol/L)(0.55)(0.1L)
g(base) = 0.74 g
The pH of the Phosphate Buffer was 6.85

ChemicalRequired Weight (g)Actual Weight (g)
NaH2PO40.31 g0.31 g
Na2HPO40.74 g0.74 g
Table 6: Chemical Weights
Discussion and Conclusion
From this lab we can conclude that H2O is a poor buffer because the pH changed after only a small amount of acid or base was added. It was also demonstrated that Acetate was a good buffer because it took a large amount of acid or base to be added before the pH changed considerably. Acetic Acid required the largest volume to reach the desired pH so we can conclude that it is a good buffer as well. The importance of a buffering system can been seen in biological systems. The body requires very specific pH levels and uses positive and negative feedback loops to maintain homeostasis. Slight changes in pH in the skin for example, could lead to more skin infections, whereas a slight pH change in the blood, could be fatal.

Xavier L. Titration. chemwiki. [cited February 1, 2014]. Available from:
Mills R. 2014. Acids, bases, and buffers lecture notes.
Stephens B. 2014. Acids, Bases and Buffers: How do Buffers Work.

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