Not All Salt Water Is Neutral

Eight Tubes, One Indicator

Each of the eight test tubes holds 14 mL of a salt solution at 0.1 mol/L, with the same universal indicator in every one. They stand in a clear acrylic rack in front of an LED light box, and the light comes through each tube from behind.

From left to right: alum, ammonium chloride, magnesium chloride, sodium chloride (table salt), sodium hydrogencarbonate (baking soda), sodium acetate, sodium hydrogenphosphate and sodium carbonate (washing soda). The colours run from red, orange, yellow and green to blue and violet.

As this page computes them, their pH values are 3.7, 5.1, 6.2, 7.0, 8.3, 8.9, 9.7 and 11.7. All are salts; only the table salt is neutral, some are acidic and some basic.

14 mL each at 0.1 mol/L · 25 °C, water free of carbon dioxide · pH 3.7 to 11.7 · colours after CDH’s pH 4⁠–⁠11 chart

Why Salt Water Stays Green

Water trades protons with itself: one molecule hands an H⁺ to another, making H₃O⁺ and OH⁻, and they soon swap back. At 25 °C the product of their concentrations is Kw = 1.0 × 10⁻¹⁴, so pure water has 1.0 × 10⁻⁷ mol/L of each and a pH of 7.00. That is neutral.

Table salt dissolves into Na⁺ and Cl⁻, each held in a ring of water molecules. Cl⁻ is what hydrochloric acid, a strong acid, leaves behind, and it hardly takes a proton back from water; Na⁺ hardly gives one up. So salt water at 0.1 mol/L stays at pH 7.00 and the indicator stays green.

Left open, it changes: carbon dioxide from the air dissolves into a very weak acid. As this page computes it, at today’s 427 ppm of carbon dioxide both pure water and salt water settle near pH 5.6, the figure the US EPA gives for normal rain. Classroom demonstrations often use water that has been boiled, and this page’s numbers are for water free of carbon dioxide.

H₂O + H₂O ⇌ H₃O⁺ + OH⁻ · Kw = 1.0 × 10⁻¹⁴ at 25 °C · neutral at pH 7.00 · open to air, about pH 5.6

Acetate Takes a Proton

Acetic acid is a weak acid: it lets its proton go reluctantly. The other way round, acetate, CH₃COO⁻, what is left when it does, will take a proton. Now and then it takes an H⁺ from a water molecule, becomes acetic acid again and leaves an OH⁻.

CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻. The extra OH⁻ makes the solution basic: sodium acetate at 0.1 mol/L has a pH of 8.88 as this page computes it, and the indicator turns blue. Na⁺ only looks on.

Titrate acetic acid with sodium hydroxide and the flask holds a sodium acetate solution at the equivalence point, which is why its pH there is not 7; see Acetic Acid Instead in Acid–Base Titration.

CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻ · sodium acetate at 0.1 mol/L, pH 8.88

Ammonium Gives One Away

Ammonia is a weak base, and ammonium, NH₄⁺, is ammonia with one proton taken on. Now and then ammonium hands that proton to a water molecule, becomes ammonia again and leaves an H₃O⁺.

NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺. So the solution is acidic: ammonium chloride at 0.1 mol/L has a pH of 5.12 as this page computes it, and the indicator turns orange. This time Cl⁻ looks on.

When both ions trade protons with water, the stronger wins. In ammonium acetate, ammonium gives protons and acetate takes them, almost equally strongly, and the solution sits at pH 7.00; yet it is not neutral the way salt water is: about 6 acetates in 1,⁠000 have become acetic acid.

NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺ · ammonium chloride at 0.1 mol/L, pH 5.12 · ammonium acetate, pH 7.00

One in Thirteen Thousand

A weak acid’s Ka times its partner’s Kb is exactly Kw: Ka × Kb = Kw. Acetic acid’s Ka is 1.75 × 10⁻⁵, so acetate’s Kb is 5.7 × 10⁻¹⁰. The weaker the acid, the readier its partner is to take a proton.

With a Kb that small, only about 0.0075% of the acetate in sodium acetate at 0.1 mol/L has taken a proton, as this page computes it: about one in 13,⁠000. The salt has not “split”; a very few ions have traded one proton with water.

That is enough. The OH⁻ left behind comes to about 7.6 × 10⁻⁶ mol/L, more than seventy times the amount in pure water, and the pH rises from 7 to 8.88. In ammonium chloride, too, only about 0.0075% of the ammonium has given up its proton. Ions that take protons more strongly react in far larger shares, as carbonate in washing soda does.

Ka × Kb = Kw · acetic acid Ka 1.75 × 10⁻⁵ → acetate Kb 5.7 × 10⁻¹⁰ · 0.0075%, about one in 13,⁠000

Small Ions, Big Charges

The alum tube is red because of aluminium. In water Al³⁺ is held by six water molecules, as [Al(H₂O)₆]³⁺. The ion is small and highly charged; it pulls on the O–H bonds of those water molecules, and one of them readily gives up an H⁺.

[Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅OH]²⁺ + H₃O⁺. In this page’s simple model, alum at 0.1 mol/L has a pH of 3.7; counting the ion pairs aluminium forms with sulfate as well, about 4.1⁠–⁠4.3. Either way the indicator shows red.

More charge and less size usually means more acid. The CRC Handbook gives pKa values of 14.8 for Na⁺, 11.4 for Mg²⁺ and 5.0 for Al³⁺. So magnesium chloride is only faintly acidic, pH 6.2 at 0.1 mol/L, and the indicator leans yellow. Fe³⁺ is more acidic still, pKa about 2.2, and part of the yellow-⁠brown of iron(III) chloride comes from the ion that has lost an H⁺; see Two Ions of Iron.

[Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅OH]²⁺ + H₃O⁺ · pKa: Na⁺ 14.8, Mg²⁺ 11.4, Al³⁺ 5.0 (CRC)

Baking Soda, Washing Soda

Sodium hydrogencarbonate (baking soda) and sodium carbonate (washing soda) are both salts of carbonic acid, yet at 0.1 mol/L one has a pH of 8.34 and the other 11.65. Carbonate, CO₃²⁻, takes protons with a Kb of 2.1 × 10⁻⁴; hydrogencarbonate’s is only 2.3 × 10⁻⁸, about ten thousand times less. In washing soda about 4.5% of the carbonate has taken a proton.

Hydrogencarbonate can both take a proton and give one, and the two cancel, so its pH settles half-⁠way between the two pKa values: (6.35 + 10.33) ÷ 2 = 8.34. Diluting it hardly moves the pH: from 1 mol/L down to 0.001 mol/L it only falls from 8.34 to 8.30, as this page computes it.

Sodium acetate and ammonium chloride behave otherwise: each tenfold dilution moves their pH 0.50 towards 7, and washing soda falls by about 0.5 too. For how a balance shifts when it is diluted, see Pale in Ice, Dark in Heat.

CO₃²⁻ Kb 2.1 × 10⁻⁴ · HCO₃⁻ Kb 2.3 × 10⁻⁸ (OpenStax) · baking soda pH ≈ (pKa₁ + pKa₂) ÷ 2 = 8.34

The Colour Is Only a Guide

Universal indicator is a mix of several indicators, each changing colour over about pKa ± 1, and together they run from red to violet. This page’s colours follow CDH Fine Chemical’s pH 4⁠–⁠11 chart: 4 red, 5 orange red, 6 yellow, 7 light green, 7.5 green, 8.5 greenish blue, 9 blue, 10 violet, 11 deeper reddish violet. Other makers’ charts differ by about a pH unit, and baking soda at 8.3 and sodium acetate at 8.9 differ by only a shade.

To read decimals takes a pH meter: a glass electrode in the solution, reading to 0.01. The beaker holds sodium carbonate at 0.1 mol/L and the meter shows 11.39. It measures the activity of H⁺, which in a solution full of ions comes out a little below the 11.65 computed from concentrations, here corrected with the Davies equation; a handbook gives 11.37 measured for a 1% solution.

CDH chart: 4 red · 6 yellow · 7 light green · 9 blue · 10 violet · a pH meter reads to 0.01 · sodium carbonate: 11.65 by concentration, 11.39 by activity

Laundry, Heartburn, Fields

Washing soda makes a basic solution because carbonate takes protons from water; its makers sell it as a cleaner that cuts grease. The solid stings the eyes.

Baking soda is one of the active ingredients in antacids: the US rules for over-⁠the-⁠counter antacids list sodium bicarbonate. It takes the H⁺ of stomach acid and gives off carbon dioxide, HCO₃⁻ + H⁺ → H₂O + CO₂.

Ammonium fertilisers slowly make soil more acidic, but not mainly because ammonium gives up its own proton: in soil at pH 5⁠–⁠7 only 0.006⁠–⁠0.56% of the ammonium is ammonia, as this page computes it. Soil bacteria oxidise ammonium to nitrate and release two H⁺ for every nitrogen, NH₄⁺ + 2O₂ → NO₃⁻ + 2H⁺ + H₂O; Cornell University’s course calls this often the largest single source of acidity in fertilised fields.

washing soda: carbonate takes H⁺ · baking soda: HCO₃⁻ + H⁺ → H₂O + CO₂ · ammonium fertiliser: nitrification releases H⁺

It Began with Syrup of Violets

In 1664 Robert Boyle wrote in Experiments and Considerations Touching Colours: put a little syrup of violets on white paper, add two or three drops of spirit of salt or vinegar, and the syrup turns red at once; use a solution of potashes instead, and the blue is “turn’d in a moment into a perfect Green”.

He used it as a test: a syrup turned red mostly argues an acid; turned green, it argues a salt “repugnant to that of the Tribe of Acids”. He set down what he saw and did not say why a salt’s solution should turn the syrup green.

In 1909 Søren Sørensen, at the Carlsberg Laboratory, introduced pH; in 1923 Brønsted and Lowry described acids as things that give up protons and bases as things that take them. That is the account this page gives of every tube.

1664 Boyle: syrup of violets · 1909 Sørensen: pH · 1923 Brønsted, Lowry: giving and taking protons

The Bench, Piece by Piece

Eight test tubes, an acrylic rack, an LED light box, a pH meter, a glass electrode, a beaker, an electrode stand, four jars of solids and a bottle of universal indicator. The tubes are 16 × 160 mm borosilicate glass; the colours read clearly only when the light comes from behind.

tubes · rack · light box · pH meter · electrode · beaker · stand · solids · indicator