This article incorporates, in modified form, material from Illustrated Guide to Home Chemistry Experiments: All Lab, No Lecture.
When an acid reacts with a base, it forms a salt and water. For example, reacting hydrochloric acid with sodium hydroxide produces the salt sodium chloride (common table salt) and water. Reacting nitric acid with potassium hydroxide produces potassium nitrate and water. Reacting acetic acid with aqueous ammonia produces ammonium acetate and water. And so on. Such reactions are referred to as neutralizing the acid with the base (or vice versa).
Unfortunately, the word neutralize is a common source of confusion among beginning chemists, many of whom assume that reacting equivalent amounts of an acid and base should yield a solution that contains only the “neutral” salt that is neither acidic nor basic, and therefore has a pH of 7.0. That’s not how it works.
The pH of a neutralized solution depends on the particular acid and base that are reacted. Reacting equivalents of a strong acid with a strong base in fact does produce a salt solution that has a pH at or near 7.0, as does reacting a weak acid with a weak base. But if the strengths of the acid and base are very different — as for example occurs if you react a strong acid with a weak base or vice versa — the pH of the neutralized solution will not be 7.0. The greater the difference in the strengths of the acid and base, the greater the difference in the pH of the neutralized solution from 7.0.
For example, if you neutralize hydrochloric acid (a very strong acid) with aqueous ammonia (a relatively weak base), the resulting solution of ammonium chloride will have a pH less than 7.0. Conversely, if you neutralize acetic acid (a relatively weak acid) with sodium hydroxide (a very strong base), the resulting neutralized solution of sodium acetate will have a pH greater than 7.0.
In this lab, we’ll determine the pH of aqueous solutions of various salts.
Required Equipment and Supplies
- goggles, gloves, and protective clothing
- beaker, 150 mL
- pH meter or pH test paper
- ammonium acetate, 0.1 M (100 mL)
- ammonium chloride, 0.1 M (100 mL)
- sodium acetate, 0.1 M (100 mL)
- sodium chloride, 0.1 M (100 mL)
All of the specialty lab equipment and chemicals needed for this and other lab sessions are available individually from Maker Shed or other laboratory supplies vendors. Maker Shed also offers customized laboratory kits at special prices, including the Basic Laboratory Equipment Kit, the Laboratory Hardware Kit, the Volumetric Glassware Kit, the Core Chemicals Kit, and the Supplemental Chemicals Kit.
Although none of the salts used in this lab are particularly hazardous, it’s good practice to wear splash goggles, gloves, and protective clothing at all times. If you make up the salt solutions from acid and base solutions as described in the Substitutions and Modifications section, use the normal precautions for working with acids and bases.
- If you have not already done so, put on your splash goggles, gloves, and protective clothing.
- Using what you know about strong and weak acids and bases, predict the approximate pH values for the solutions of the four chemicals and enter your predicted values in Table 11-2 by circling one of the pH numbers. If you are uncertain, circle a range of numbers.
- Pour about 100 mL of 0.1 M ammonium acetate into the beaker.
- Read and follow the directions for your pH meter with respect to calibrating it, rinsing the electrode between measurements and so on. Use the pH meter to measure the pH of the ammonium acetate solution, and record the observed value on Line A of Table 11-2. (If you do not have a pH meter, you can use pH test paper, which is sufficiently accurate for the purpose of this lab session.)
- Repeat steps 3 and 4 for the solutions of ammonium chloride, sodium acetate, and sodium chloride.
|Solute||Predicted pH||Observed pH|
|A. ammonium acetate||0 1 2 3 4 5 6 7 8 9 10 11 12 13 14||___.___pH|
|B. ammonium chloride||0 1 2 3 4 5 6 7 8 9 10 11 12 13 14||
|C. sodium acetate||0 1 2 3 4 5 6 7 8 9 10 11 12 13 14||
|D. sodium chloride||0 1 2 3 4 5 6 7 8 9 10 11 12 13 14||
All of the solutions used in this laboratory can be flushed down the drain with plenty of water.
Q1: What general conclusions can you draw from the pH values you observed for the four solutions?
Q2: If you mixed equal amounts of a 0.1 M ammonium chloride solution and a 0.1 M sodium acetate solution, would you expect the pH of the resulting solution to be strongly acidic, weakly acidic, about neutral, weakly basic, or strongly basic? Why?