Title: ACID BASE TITRATION. Objectives: 1. To determine the concentration of acid using titration. 2. Skills of titration techniques. Apparatus: 1. 250 volumetric flask 2. 10mL measuring cylinder 3. 25mL pipette 4. 50mL burette 5. 250mL beaker 6. 150mL conical flask 7. Retord stand 8. White tile 9. Stopwatch 10. Pipette bulb Chemicals: 1. HCl solution 2. 0. 1M NaOH solution 3. H2SO4 solution 4. Distilled water 5. phenolphthalein Introduction. An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the acid/base with an acid or base of known concentration.
This allows for quantitative analysis of the concentration of an unknown acid or base solution.
It makes use of the neutralization reaction that occurs between acids and bases and the knowledge of how acids and bases will react if their formulas are known. Acid–base titrations can also be used to find percent purity of chemicals. When a weak acid reacts with a weak base, the equivalence point solution will be basic if the base is stronger and acidic if the acid is stronger.
If both are of equal strength, then the equivalence pH will be neutral.
However, weak acids are not often titrated against weak bases because the colour change shown with the indicator is often quick, and therefore very difficult for the observer to see the change of colour. The point at which the indicator changes colour is called the end point. A suitable indicator should be chosen, preferably one that will experience a change in colour (an end point) close to the equivalence point of the reaction. First, the burette should be rinsed with the standard solution, the pipette with the unknown solution, and the conical flask with distilled water.
Secondly, a known volume of the unknown concentration solution should be taken with the pipette and placed into the conical flask, along with a small amount of the indicator chosen. The known solution should then be allowed out of the burette, into the conical flask. At this stage we want a rough estimate of the amount of this solution it took to neutralize the unknown solution. The solution should be let out of the burette until the indicator changes colour and the value on the burette should be recorded.
This is the first (or rough) titre and should be discluded from any calculations. At least three more titrations should be performed, this time more accurately, taking into account roughly where the end point will occur. The initial and final readings on the burette (prior to starting the titration and at the end point, respectively) should be recorded. Subtracting the initial volume from the final volume will yield the amount of titrant used to reach the endpoint. The end point is reached when the indicator just changes color permanently.
This is best achieved by washing a hanging drop from the tip of the burette into the flask right at the end of the titration to achieve a drop that is smaller in volume than what can usually be achieved by just dripping solution off the burette. Acid–base titration is performed with a phenolphthalein indicator, when it is a strong acid – strong base titration, a bromthymol blue indicator in weak acid – weak base reactions, and a methyl orange indicator for strong acid – weak base reactions. If the base is off the scale, i. e. a pH of >13. 5, and the acid has a pH >5. , then an Alizarine yellow indicator may be used. On the other hand, if the acid is off the scale, i. e. a pH of <0. 5, and the base has a pH <8. 5, then a Thymol Blue indicator may be used. Indicator| Color on acidic side| Range of color change| Color on basic side| Methyl Violet| Yellow| 0. 0–1. 6| Violet| Bromophenol Blue| Yellow| 3. 0–4. 6| Blue| Methyl Orange| Red| 3. 1–4. 4| Yellow| Methyl Red| Red| 4. 4–6. 3| Yellow| Litmus| Red| 5. 0–8. 0| Blue| Bromothymol Blue| Yellow| 6. 0–7. 6| Blue| Phenolphthalein| Colourless| 8. 3–10. 0| Pink| Alizarin Yellow| Yellow| 10. 1–12. 0| Red|
Acid-base titrations depend on the neutralization between an acid and a base when mixed in solution. In addition to the sample, an appropriate indicator is added to the titration chamber, reflecting the pH range of the equivalence point. The acid-base indicator indicates the endpoint of the titration by changing colour. The endpoint and the equivalence point are not exactly the same because the equivalence point is determined by the stoichiometry of the reaction while the endpoint is just the colour change from the indicator. Thus, a careful selection of the indicator will reduce the indicator error.
For example, if the equivalence point is at a pH of 8. 4, then the Phenolphthalein indicator would be used instead of Alizarin Yellow because phenolphthalein would reduce the indicator error. Common indicators, their colours, and the pH range in which they change colour are given in the table above.  When more precise results are required, or when the reagents are a weak acid and a weak base, a pH meter or a conductance meter are used. Procedure. 1. 10 mL of HX is put in a volumetric flask and is diluted with distilled water and mixed through in 100mL volumetric flask.
The solution is transferred into a beaker and labelled as HX. 2. The burette is rinsed with 10mL of 0. 1M NaOH aqueous solution and is filled with base. 3. 25mL of HX solution is transferred into a conical flask and 2 drops of phenolphthalein are added. 4. Begin to run the base solution from burette into the flask containing acid solution and phenolphthalein. A pink colour will develop which quickly disappears when solution is swirled. 5. As more bases is added, the pink solution will remain for a longer time, when this occurs slow the addition of base so that only one or ? drop of base at a time. . The flask is swirled on each addition. At the end point, one drop of base solution will turn the entire content of the flask to light pink. At this point, the final volume of the base is recorded to the nearest 0. 01mL. 7. The experiment is repeat 2 times to ensure the reliability of the experiment. 8. All the conical flask is washed and rinsed with the unknown H2SO4. 9. All procedure are repeated using the unknown H2SO4. Result. Using HCl| | 1| 2| 3| average| Volume of acid (mL) (±0. 00)| 25. 0| 25. 0| 25. 0| 25. 0| Final volume of base (mL) (±0. 000)| 9. 70| 19. 0| 29. 70| 19. 70| Initial volume of base (mL) (±0. 000)| 0. 00| 9. 70| 19. 70| 9. 8| Volume of base used (mL) (±0. 000)| 9. 70| 10. 00| 10. 00| 9. 90| Molarity of acid (M)| 0. 04| 0. 04| 0. 04| 0. 04| | | | | | Using H2SO4 | | 1| 2| 3| average| Volume of acid (mL) (±0. 00)| 25. 0| 25. 0| 25. 0| 25. 0| Final volume of base (mL) (±0. 000)| 19. 50| 39. 60| 59. 30| 39. 47| Initial volume of base (mL) (±0. 000)| 0. 00| 19. 50| 39. 60| 19. 70| Volume of base used (mL) (±0. 000)| 19. 50| 20. 1| 19. 70| 19. 77| Molarity of acid (M)| 0. 4| 0. 04| 0. 04| 0. 04| | | | | | Calculation. Discussion & Questions. 1) Why is the volume of base needed to neutralize the acid HCl is different from H2SO4? * It due to the equation which is 1 mole of base NaOH reacted with 1 mole of HCl. For H2SO4: 2 mole of NaOH reacted with 1 mole of H2SO4. 2) Why are HCl and H2SO4 known as strong acid? * Because they reacted 100%. 3) Give one example of a weak acid and explain why it called a weak acid. * Ethanoic acid, when it reacted, it do not dissolves 100%. * It only dissolve 4% and tend to return to its original state. Conclusion. ACID- BASE
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