Titration Acid-Base (Simple)

Titration is a fundamental laboratory technique used to determine the concentration of a substance in a solution. In this article, we will walk you through the steps of performing a simple acid-base titration using sodium hydroxide and hydrochloric acid. We will also discuss important safety precautions to keep in mind, such as wearing appropriate protective gear, handling chemicals with care, and disposing of waste properly. Additionally, we will address environmental concerns related to titration, including the safe disposal of chemical waste and glassware. By following these guidelines, you can perform a titration safely and responsibly while protecting yourself and the environment.

Aim

The purpose of this experiment is to determine the concentration of a solution of sodium hydroxide by titration against a standard solution of sodium hydroxide.

Introduction

Hydrochloric acid is a monoprotic acid in that it produces one mole of hydrogen ions per mole of compound, we can simplify the formula to HA. This simple formula is often used to represent an acid.

Sodium hydroxide reacts with hydrochloric acid according to the equation:

NaOH + HCl –> NaCl + H₂O

To show you when the reaction is complete – the stoichiometric point or equivalence point – you use an indicator called phenolphthalein, which is colourless in acid and pink in alkaline solution. The point at which the addition of one drop (or even less) of alkali changes the solution from colourless to just faintly pink is called the end-point and, in this case, shows that the reaction is just complete.

Requirements

• safety spectacles
• filter funnel, small
• burette, 50 cm³
• 2 beakers, 100 cm³
• sodium hydroxide solution, approx. 0.5 mol dm³ NaOH (CORROSIVE)
• pipette, 25 cm³
• pipette filler
• Hydrochloric acid solution, 0.5 mol dm³ (IRRITANT)
• conical flask, 250 cm³
• phenolphthalein indicator solution
• white tile
• wash-bottle of distilled water

Procedure

1. Using the funnel, rinse the burette with the sodium hydroxide solution and fill it with the same solution. Do not forget to rinse and fill the tip. Record the initial burette reading in the ‘rough’ column of the Results Table.
2. Using a pipette filler, rinse the pipette with some of the Hydrochloric acid solution and carefully transfer 25.0 cm³ of the solution to a clean 250 cm³ conical flask.
3. Add 2-3 drops of the phenolphthalein indicator solution.
4. Run sodium hydroxide solution from the burette swirling, until the solution just turns pink.
5. Refill the burette with the sodium hydroxide solution, and again record the initial burette reading to the nearest 0.05 cm³ (one drop).
6. Using the pipette, transfer 25.0 cm³ of the hydrochloric acid solution to another clean conical flask. Add 2-3 drops of the phenolphthalein indicator solution.
7. Carefully titrate this solution to the end-point, adding the alkali drop by drop when you think the colour is about to change.
8. Repeat steps 5, 6 and 7 at least twice more.
9. Empty the burette and wash it carefully immediately after the titration, especially if it has a ground glass tap.

Accuracy

You should record burette readings to the nearest 0.05 cm³ (approximately one drop). Consecutive titrations should agree to within 0.10 cm³ and, strictly, you should repeat the titration until this is achieved. However, you may have neither the time nor the materials to do this. With practice, your technique will improve so that it is not necessary to do more than four titrations.

Calculate the mean of the two (or preferably three) closest consecutive readings and quote this also to the nearest 0.05 cm³

Note that this does not introduce a fourth significant figure; it merely makes the third figure more reliable.

Pipette Solution		mol dm-3	cm3
Burette Solution		mol dm-3
Indicator

		Rough	1	2	3	(4)
Burette Reading	Initial
	Final
Volume used (titre) cm3
Mean titre cm3

Solution	Molar ratio	Moles (mol)	Volume (dm3)	Concentration (mol dm-3)

Table of Results (PDF)

Questions

1. What effect would each of the errors described below have on the calculated value of the concentration of sodium hydroxide?

(a) The burette is not rinsed with the sodium hydroxide solution.

(c) The tip of the burette is not filled before titration begins.

(d) The conical flask contains some distilled water before the addition of potassium hydrogen phthalate.

2. In using phenolphthalein as an indicator, we prefer to titrate from a colourless to pink solution rather than from pink to colourless. Suggest a reason for this.
3. Why is it advisable to remove sodium hydroxide from the burette as soon as possible after the titration.