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General chemistry – Titration

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Introduction Titration is defined as technique whereby one solution of accurately known concentration also known as the titrant is added slowly into another solution of unknown concentration until a neutralization reaction is reached. There are several types of titration methods in chemistry. Firstly, is the acid-base titration method which was carried out during this experiment. In an acid-base titration, the acid/base with a known concentration and fixed volume will be added into the conical flask with a pipette and this would be the solution that is to be titrated.

This solution is also known as the titrand. The base/acid with a known concentration but no fixed volume would then be added into the burette. This solution is called the titrant [1]. The point where neutralization is reached is usually indicated by an indicator while in some reactions the solutions are self-indicating. Hence, a small amount of acid-base indicator will be added to the titrand to observe the colour changes as the reactions proceeds.

The aim of this laboratory report is to have a better understanding on the importance of selecting suitable indicators for detecting the end points of acid-base titrations. It is also important to select a suitable indicator to ensure that the titration curves plotted are precise and accurate. Therefore, it is critical to choose the right indicator when carrying out an acid-base titration because different indicators changes colour at different pH ranges. For a successful titration, the end point should be close to its equivalence point.

Clancy C et al [2] states that the difference between the end point and the equivalence point is that the ‘end point is the pH at which the indicator changes colour in a titration while the equivalence point is the point in a titration at which the number of moles of the substance added is stoichiometrically equivalent to the number of moles present in the original test solution, or the point at which the acid and base have completely reacted’. The pH changes during an acid-base titration is monitored by using a pH meter. These changes are then plotted on a graph to obtain titration curves for some acid-base reactions.

However, a weak acid-weak base reaction is reversible and thus, is not suitable for titrimetric analysis [1]. Objectives This experiment was carried out to have a better understanding on the importance of choosing a suitable indicator for detecting the end points of acid-base titrations. It was also carried to obtain a titration curve for some acid-base titrations. Materials Chemicals 0. 1 M Hydrochloric acid, 0. 1 M acetic acid, 0. 1 M sodium hydroxide, 0. 1 M ammonium hydroxide, phenolphthalein, screened methyl orange, methyl orange.

Apparatus 100mL beakers, 250mL conical flasks, 50mL burettes, 25mL pipettes, pipette fillers, funnels, pH meters. Procedures Refer to the procedures on the third and fourth page of the General Chemistry Practical Module. Results Part I: Suitability of indicators Strong acid- weak base titration (0. 1M Hydrochloric acid with 0. 1 M Ammonium hydroxide) Indicator: Phenolphthalein Titration Initial burette reading (mL) Final burette reading (mL) Volume of Ammonium hydroxide added (mL) 1 10. 00 40. 00 29. 00 2 10. 00 40. 00 30. 00 3 10. 00 39. 00 30. 00 Average volume of Ammonium hydroxide titre 29. 50mL.


The colour changed from red to greyish green upon reaching the end point. The average volume of NH4OH was 27. 85mL. In part I (B), a strong base was titrated against weak acid. 0. 1M Sodium Hydroxide, NaOH was titrated against 0. 1M Acetic acid, CH3COOH. 2-3 drops of phenolphthalein was added into CH3COOH in the conical flask. The colour changed from colourless to pink for phenolphthalein upon reaching the end point. The average volume of NaOH was 24. 95mL. The experiment was repeated with screened methyl orange as the acid-base indicator. The colour changed from red to greyish green upon reaching the end point.

The average volume of NaOH was 3. 10mL. For the part II, the pH for the titration curve in strong acid-weak base titration, had a starting pH of 1. 04 due to the strong acid HCL. As the weak base, NH4OH was gradually added into the acid, the pH of the solution begins to increase. When the NH4OH was titrated for 28. 5mL, the pH rocketed to 7. 94. The sharp increase in the graph indicates the end point of titration. As NH4OH is continuously added in an excess amount, the graph reaches a plateau at pH 8. 34.

Whereas, the titration curve for weak acid-strong base begins at a pH of 3.00 due to the weak acid, CH3COOH. When NAOH was titrated, the pH starts to rise gradually until it reaches 6. 76. From this point, pH increases sharply to indicate the end point of the reaction. When excess of NAOH was added the pH reached 11. 95. When both titration curves are compared, the main difference is the pH value when the equivalence point occurs. For the strong acid- weak base titration curve, the equivalence point is attained at pH 5. 57. The equivalence point is less than 7 thus acidic. This is due to the formation of Ammonium Chloride, NH4CL.

NH4CL is highly soluble in water and dissociates into NH4+ ions and Cl – ions. NH4+ ions will then undergo hydrolysis in water to produce hydroxonium ions, H3O+ which will result in an acidic condition. Whereas, for the weak acid- strong base titration curve, the equivalence point is achieved at pH 8. 42. The equivalence point is more than 7 thus basic. This is because Sodium Acetate, CH3COOHNa is formed. CH3COOHNa dissociates into CH3COO- ions and Na+ ions. CH3COO- ions will then undergo hydrolysis in water to produce hydroxide ions, OH- which will result in a basic condition.

The most suitable indicator for the strong acid-weak base titration is screened methyl orange rather than phenolphthalein indicator. This is because the pH range of screened methyl orange is 3. 0 – 5. 0 which is close to the equivalence point of the titration which was pH 5. 57. Whereas, phenolphthalein has a pH range of 8. 3-10. 0 which is a more suitable indicator for the weak acid-strong base titration that attained its end point at pH 8. 42, thus are close to each other. A few discrepancies in the findings could have occurred due to parallax error.

This could have been caused due to the differences in eye levels when reading the meniscus. Multiple readings were taken during this experiment that the possibility of the eye level not being directly perpendicular to the meniscus is high. Thus, causing parallax errors. Furthermore, in Part I of the experiment, the precise colour change of the solution when it reached its end point could have been misjudged even with the aid of a white tile This could then affect the volume of titre. Conclusion There are many indicators available to carry out acid-base titrations.

However, it is crucial to select a suitable indicator for a particular titration because each indicator varies in pH ranges. Selecting a wrong indicator can affect drastically affect the accuracy of your readings in the experiment.


[1] Ebbing DD & Gammon SD (2005). General Chemistry, 9th Edition, McGraw Hill.

[2] Clancy C, Farrow K, Finkle T, Francis L, Heimbecker B, Nixon-Ewing B Et al. Salts, Buffers, Titrations, and Solubility. Canada: McGraw-Hill Ryerson, 2011; p. 532.

[3] Acid base indicator [Internet]. Slideshare. net. 2014 [cited 24 October 2016]. Available from: http://www. slideshare. net/mhsn47/acid-base-indicator.

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