046 – pH

Almost every chemical process that takes place in water is sensitive to how acidic or basic that water is. Blood is held within about a tenth of a pH unit of 7.4, soil pH decides which crops will grow, and industrial effluent is regulated on pH before it may be discharged. Measuring pH reliably, and knowing how far to trust a given measurement, is one of the most widely used skills in analytical chemistry.

The quantity itself is defined logarithmically: pH = −log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in moles per litre. Because the scale is logarithmic, one whole unit represents a tenfold change in acidity — a solution at pH 3 carries a hundred times as many hydrogen ions as one at pH 5. Pure water at 25 °C sits at pH 7, the point where [H⁺] and [OH⁻] are both 10⁻⁷ mol/L.

This laboratory uses three methods in increasing order of precision. Litmus paper answers only whether a solution is acidic or basic. A universal indicator produces a colour that can be matched against a chart to within roughly one pH unit. A calibrated pH meter reads to a hundredth of a unit. You will apply all three to three unknown liquids, and then dissolve a salt — ammonium sulfate — and find that a compound containing no hydroxide group and no obvious acidic hydrogen nevertheless gives a distinctly acidic solution. Explaining why is the analytical heart of the exercise.

Educational Goals

Familiarization with the laboratory environment

• Identify the layout, instruments and workflow of the Proteus pH laboratory, and locate the litmus papers, the indicator chart, the well plate, the balance and the pH meter before beginning.

Use of protective equipment

• Wear gloves and eye protection when handling acids, bases and unidentified solutions, and recognise that an unknown liquid is treated as hazardous until it has been characterised.

Choosing an appropriate measurement method

• Distinguish what each of the three methods can establish: litmus gives a yes-or-no answer about acidity, a universal indicator narrows the pH to about one unit, and a pH meter gives a quantitative reading to about 0.01 unit.
• Select the method that matches the question being asked, rather than defaulting to the most precise instrument available.

Correct operation of a pH meter

• Immerse the electrode fully, read the digital display once it has settled, and rinse with distilled water between samples.
• Explain why rinsing matters: carry-over of even a small volume of a strongly acidic or basic solution shifts the next reading, and the effect is largest when moving between extremes of the scale.

Preparing a solution from a weighed solid

• Weigh a solid on the electronic balance, transfer it without loss, dissolve it in a measured volume of distilled water, and stir until the solution is homogeneous.
• Recognise that a concentration is only as good as the poorest of the two measurements it depends on, the mass and the volume.

Quantitative interpretation of the pH scale

• Convert between pH and [H⁺] in both directions, and use the relationship pH + pOH = 14 at 25 °C.

Predicting whether a dissolved salt will be acidic, basic or neutral

• Identify the parent acid and base of a salt, and use their relative strengths to predict the character of the solution it forms — the principle that explains the ammonium sulfate result.

Protocol

PART 1: Determine the acid base character and the pH of a liquid substance

1. Measure 20 mL of solution 1 using the graduated cylinder.
2. Pour 20 mL of solution 1 into a 50 mL beaker.
3. Dip a red litmus paper into it.
4. Then dip a blue litmus paper.
5. Finally dip a pH indicator paper.
6. Compare the colors obtained with those available on the pH chart.
7. Rinse the graduated cylinder with distilled water.
8. Repeat steps 1 to 7 with substances 2 and 3.

PART 2: Determine the pH of a liquid substance using a universal indicator

1. Using the dropper; take solution 1 from the 50 mL beaker.
2. Then empty the dropper into well no. 1 of the well plate.
3. Add one drop of universal pH indicator into the same well.
4. Stir using the glass rod.
5. Compare the color obtained with those available on the pH chart.
6. Repeat steps 1 to 5 with substances 2 and 3.

PART 3: Determine precisely the pH of a liquid substance

1. Immerse the pH meter electrode in the beaker containing solution 1.
2. Read the measurements on the digital display.
3. Rinse the electrode with distilled water.
4. Dry the electrode with absorbent paper.
5. Repeat steps 1 to 4 with substances 2 and 3.

PART 4: Determine the pH of a solid substance

1. Weigh about 1.8 g of ammonium sulfate powder using the electronic balance.
2. Measure 100 mL of distilled water using a graduated cylinder.
3. Pour the weighed ammonium sulfate powder into a 250 mL beaker.
4. Pour the contents of the graduated cylinder into the same beaker.
5. Stir using the glass rod for 5 seconds.
6. Dip a red litmus paper into the 250 mL beaker that contains the ammonium sulfate solution
7. Dip a blue litmus paper into the 250 mL beaker that contains the ammonium sulfate solution
8. Dip a pH indicator paper into the 250 mL beaker that contains the ammonium sulfate solution.
9. Compare the colors obtained with those available on the pH chart.
10. Immerse the pH meter electrode in the ammonium sulfate solution.
11. Read the measurement on the digital display.
12. Rinse the electrode with distilled water.
13. Dry the electrode with absorbent paper.

Anticipated Outcomes

The four solutions. Three unknown liquids are characterised first, and the ammonium sulfate solution is prepared and measured in Part 4.

The definitions behind the numbers. pH is defined as pH = −log₁₀[H⁺]. In any aqueous solution at 25 °C the two ion concentrations are tied together by the ion product of water, [H⁺][OH⁻] = K_w = 1.0 × 10⁻¹⁴, from which pH + pOH = 14. Distilled water has [H⁺] = [OH⁻] = 10⁻⁷ mol/L and therefore pH 7.

Solution 2, a strong base. Sodium hydroxide dissociates completely, so a 0.1 M solution has [OH⁻] = 0.1 mol/L directly. Then pOH = −log₁₀(0.1) = 1, and pH = 14 − 1 = 13. No equilibrium calculation is needed, which is what distinguishes a strong base from a weak one.

Solution 1, a weak acid. Ethanoic acid dissociates only partially, so its concentration alone is not enough — the acid dissociation constant is needed as well. For a weak acid of concentration C that is only slightly dissociated, K_a = x²/(C − x) ≈ x²/C with x = [H⁺], which rearranges to the working form:

pH = ½(pK_a − log₁₀C)

A 5 % v/v solution is about 0.87 M, and ethanoic acid has pK_a = 4.76, so pH = ½(4.76 + 0.06) = 2.41. Note how much less acidic this is than a strong acid at the same concentration: 0.87 M hydrochloric acid would sit near pH 0.06, more than two units lower, because it dissociates completely.

Solution 4, why a salt turns out acidic. Ammonium sulfate contains no hydroxide group and no hydrogen atom that looks obviously acidic, yet its solution is clearly acidic. The reason is that it is the salt of a weak base and a strong acid. On dissolving, it releases ammonium and sulfate ions:

(NH₄)₂SO₄(s) → 2 NH₄⁺(aq) + SO₄²⁻(aq)

The ammonium ion is the conjugate acid of ammonia and donates a proton to water, NH₄⁺ → NH₃ + H⁺, with pK_a = 9.25. The sulfate ion is the conjugate base of a strong acid and is effectively a spectator. Only the ammonium contributes, and it makes the solution acidic.

Each formula unit supplies two ammonium ions, so 1.8 g in 100 mL gives [NH₄⁺] = 2 × (1.8 / 132.14) / 0.100 = 0.272 M. Applying the same weak-acid expression as for solution 1: pH = ½(9.25 − log₁₀0.272) = ½(9.25 + 0.565) = 4.91.

How little of it actually reacts. That pH corresponds to [H⁺] = 1.24 × 10⁻⁵ mol/L, which is only about 45 parts per million of the ammonium present. Fewer than one ammonium ion in twenty thousand gives up its proton — and yet the solution is two full pH units away from neutral. This is the clearest demonstration in the laboratory of how sensitive pH is: an almost undetectable degree of reaction moves the reading a long way, precisely because the scale is logarithmic.

Solubility check. The solubility of ammonium sulfate is about 744 g/L at room temperature. This solution is only 18 g/L, so it dissolves completely and no solid remains to complicate the measurement.

Summary of Assignment by Grade Range

Grade 9–10

• Focus: the pH scale as a measure of acidity, and the difference between a qualitative and a quantitative test.
• Activities: classify the three unknown liquids as acidic, neutral or basic using red and blue litmus; match the universal indicator colour against the chart; read the pH meter and record all three results side by side; observe that the ammonium sulfate solution is acidic even though it is a salt, and describe the observation without yet explaining it.

Grade 11

• Focus: quantitative treatment of the scale, and the distinction between strong and weak electrolytes.
• Activities: convert each measured pH into a hydrogen ion concentration and back; use pH + pOH = 14 to obtain the pH of 0.1 M sodium hydroxide from first principles and compare it with the meter; explain why the same concentration of a strong and a weak acid give very different pH values; compare the precision of the three methods and state which is adequate for which purpose.

Grade 12 / College Level

• Focus: weak-acid equilibrium, salt hydrolysis, and the limits of a simplifying assumption.
• Activities: derive pH = ½(pK_a − log C) from K_a = x²/(C − x) and state the approximation it relies on; predict the pH of the ammonium sulfate solution before measuring it, including the factor of two from the two ammonium ions per formula unit; calculate the fraction of ammonium that dissociates and comment on what the result implies about the sensitivity of a logarithmic scale; discuss whether treating sulfate as a spectator is justified, and estimate the direction and size of the error introduced by that assumption.

Laboratory essentials

Instruments

• Beakers (50 mL & 250 mL)
• Well plate
• Droppers
• Electronic scale
• Glass rod
• Graduated cylinders (25 mL & 100 mL)
• Spatula
• pH meter
• pH colour chart
• Red litmus paper
• Blue litmus paper
• pH indicator paper
• Paper towel
• Gloves and eye protection

Products

• Ammonium sulfate (powder)
• Universal pH indicator (liquid)
• Distilled water
• Unknown liquid #1 — ethanoic acid 5 % v/v
• Unknown liquid #2 — sodium hydroxide 0.1 M
• Unknown liquid #3 — distilled water