The pH Landscape

Map the full acid-base spectrum using natural indicators and household chemicals

Difficulty: Easy | Time: 45–60 min | Visual Impact: Very High

Historical Context

Natural pH indicators have been used for centuries. Robert Boyle described the color changes of plant extracts in acids and bases in the 1660s. Anthocyanins — the pigments responsible for the red, purple, and blue colors in many plants — were among the first indicators studied systematically.

The connection between color and acidity was crucial to understanding acids and bases long before pH was defined. In 1909, Søren Sørensen introduced the pH scale while working at the Carlsberg brewery in Copenhagen, where he needed a consistent measure of acidity to standardize yeast fermentation. The pH scale transformed qualitative color observations into a precise logarithmic quantity: each pH unit represents a tenfold change in hydrogen ion concentration.

Red cabbage remains one of the most informative natural indicators, spanning nearly the entire pH range with eight distinct color zones. In this experiment you will map those zones with real chemicals, make a second indicator from turmeric, and observe how buffering slows the color change — a phenomenon critical to blood chemistry.

Materials

For the indicator: - Red cabbage — ¼ head, chopped (~200 g) - Water — 500 mL - Pot for boiling, strainer

For the pH scale: - Citric acid solution — 1 tsp in 100 mL water - Vinegar (white or malt) - Black coffee (freshly brewed) - Water (distilled or tap) - Baking soda solution — 1 tsp in 100 mL water - Sodium carbonate solution — 1 tsp in 100 mL water - Calcium hydroxide (limewater) — pinch in 100 mL water, strained - 8 clear cups or glasses, labelled

For turmeric indicator (optional): - Ground turmeric powder — 1 tsp - Rubbing alcohol or warm water — 50 mL

For buffering demo (optional): - Baking soda — 1 tsp in 100 mL water - Lemon juice or vinegar, added drop by drop

Procedure

Part 1 — Make the indicator

  1. Boil the chopped red cabbage in 500 mL of water for 10 minutes, stirring occasionally. The water will turn deep purple.
  2. Strain out the cabbage. You should have about 400 mL of deep-purple indicator solution. Let it cool to room temperature.
  3. Line up 8 labelled cups. Add 40–50 mL of indicator to each.

Part 2 — Build the pH scale

Add a small volume (10–20 mL) of each test solution to its cup and stir. Note the color change:

Cup Substance Expected pH Expected color
1 Citric acid solution ~2 Crimson red
2 Vinegar ~3 Pink-red
3 Black coffee ~5 Violet-pink
4 Water ~7 Purple (no change)
5 Baking soda solution ~8 Blue-violet
6 Borax solution (if available) ~9 Blue
7 Sodium carbonate solution ~11 Blue-green/teal
8 Limewater (calcium hydroxide) ~12 Green-yellow

Arrange the cups in a row from most acidic to most basic — you now have a visual pH scale.

Part 3 — Turmeric as a second indicator

  1. Dissolve 1 tsp of ground turmeric in 50 mL of warm water (or rubbing alcohol, which gives a brighter result). Let it sit 5 minutes, then strain if needed.
  2. Add a small amount to separate cups of each test solution. Note: turmeric only changes at one threshold — it is yellow in acid and neutral, and turns red-brown in base above about pH 8.
  3. Compare the two indicators side-by-side. Red cabbage gives 8 distinct color zones; turmeric gives only 2. This illustrates an important point: different indicators are useful over different pH ranges, and for different precision needs.

Part 4 — The buffering demonstration

  1. Make a fresh cup of baking soda solution (pH ~8, blue-violet with cabbage indicator).
  2. Add vinegar or lemon juice one drop at a time, stirring after each drop, and count the drops needed to cause a visible color shift toward purple.
  3. Repeat with plain water (start at purple/neutral) and add the same acid, one drop at a time.
  4. Compare: with baking soda, many more drops are needed before the color shifts — the solution resists the change. This is buffering.

The pH Color Chart

Red cabbage anthocyanin changes through the spectrum as pH rises:

pH range Anthocyanin form Color
1–2 Flavylium cation (AH⁺) Crimson / bright red
3–4 Quinoidal base losing one proton Pink-red
5–6 Neutral form Violet-pink
7 Neutral Purple
8–9 Mono-anionic Blue-violet / blue
10–11 Di-anionic Teal / blue-green
12–13 Tri-anionic Green
14 Further deprotonation Yellow

The Science

Why anthocyanins change color with pH

Anthocyanins belong to a class of pigments called flavonoids. The core of the molecule, the flavylium cation, carries a positive charge at low pH. As the pH rises, the molecule progressively loses protons (H⁺ ions) from its phenolic –OH groups, changing the distribution of electrons across the ring system.

Each deprotonation shifts the wavelength of maximum absorption — the color of light the molecule most strongly absorbs. At pH 1 the molecule absorbs green light, so we see red. At pH 7 it absorbs orange light, so we see purple. At pH 11 it absorbs red light, so we see teal. At pH 14 it absorbs blue, so we see yellow.

This is a beautiful example of how molecular structure directly determines color: every time a proton is removed, the electron cloud rearranges, and the whole molecule changes the wavelength it “chooses” to absorb.

\[\ce{AH+ (red) <=> A (purple) + H+}\]

\[\ce{A (purple) <=> A- (blue) + H+}\]

\[\ce{A- (blue) <=> A^{2-} (green) + H+}\]

Each equilibrium is governed by the pH of the solution. Add acid (more H⁺) and the equilibrium shifts left, toward the red forms. Add base (remove H⁺) and it shifts right, toward blue and green.

Why turmeric has only two colors

Turmeric’s active pigment, curcumin, has a simpler molecular structure than anthocyanins. It undergoes only one significant protonation change, around pH 7.4–8.5. Below this threshold it is yellow; above it, the enolate form is red-brown. There are no intermediate color zones — curcumin is essentially a pH “switch” rather than a graduated scale.

Buffering

The baking soda demonstration shows buffering: the ability of a solution to resist pH changes when small amounts of acid or base are added.

Baking soda (NaHCO₃) dissolves to give bicarbonate ions (HCO₃⁻). When acid is added, the bicarbonate absorbs it:

\[\ce{HCO3^-(aq) + H+(aq) -> H2CO3(aq) -> H2O(l) + CO2(g)}\]

The H⁺ ions are “mopped up” before they can significantly change the pH. Only when all the bicarbonate is consumed does the pH drop sharply.

Human blood is buffered at pH 7.4 by the same bicarbonate system, with phosphate and protein buffers as backup. Without this buffering, a glass of orange juice would be fatal — the acid would crash blood pH and stop enzyme function.

Explore Further

Map your household: Test as many household liquids as you can — shampoo, conditioner, dish soap, antacid, sparkling water, milk, egg white, tomato juice. Where do they fall on your color scale?

Concentration vs. pH: Make a series of citric acid solutions at 10×, 5×, 2×, and 1× your standard concentration. Does doubling the acid double the color shift? (It should not — pH is logarithmic, so a tenfold dilution shifts pH by only 1 unit.)

Indicator paper: Soak strips of coffee filter paper in your cabbage indicator and let them dry. You now have reusable pH indicator paper. Compare their accuracy to your cups — do they show the same colors?

The turmeric stain test: Leave a drop of turmeric indicator on white paper and add a drop of baking soda solution. The paper turns bright red. Add a drop of acid and it returns to yellow. This is why turmeric stains turn reddish when treated with alkaline soap — and why rinsing with a little vinegar removes them.

Chemicals used in this experiment: