The pH Landscape
Summary
Red-cabbage juice is an indicator1 that spans nearly the whole pH2 scale in eight colours, and you’ll build a visual pH ladder with household chemicals, compare it with turmeric, and watch a buffer3 resist change. You’ll be able to explain why losing a proton changes the colour, why pH is logarithmic, and why blood needs buffering.
History
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.
Hazards & preparation
PPE: safety glasses. This is a low-hazard, mostly-kitchen experiment.
- The strongest chemicals here — sodium carbonate (washing soda) and limewater — are mild alkalis; rinse splashes off skin and keep them out of eyes.
- Don’t taste any of the test solutions once indicator is added.
Disposal: everything here is dilute and benign — flush down the drain with water. See the Safety page.
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
- Borax solution (optional) — 1 tsp in 100 mL water, for a pH ~9 point
- 8 clear cups or glasses, labelled
For turmeric indicator (optional): ground turmeric — 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, drop by drop
Procedure
Part 1 — Make the indicator
- Boil the chopped cabbage in 500 mL water for 10 minutes — the water turns deep purple.
- Strain out the cabbage (~400 mL of indicator). Cool to room temperature.
- Line up 8 labelled cups and add 40–50 mL of indicator to each.
Part 2 — Build the pH scale
Add 10–20 mL of each test solution to its cup and stir. Note the colour:
| 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 from most acidic to most basic — a visual pH scale.
Part 3 — Turmeric as a second indicator
- Dissolve 1 tsp turmeric in 50 mL warm water (or rubbing alcohol for a brighter result); let sit 5 minutes and strain.
- Add a little to separate cups of each test solution. Turmeric is yellow in acid and neutral, and turns red-brown in base above ~pH 8 — only one threshold.
- Compare: red cabbage gives 8 zones, turmeric only 2 — different indicators suit different ranges and precisions.
Part 4 — The buffering demonstration
- Make a fresh cup of baking-soda solution (pH ~8, blue-violet with cabbage).
- Add vinegar or lemon juice one drop at a time, stirring, counting drops until the colour visibly shifts toward purple.
- Repeat with plain water (starting neutral) and the same acid, drop by drop.
- Compare: baking soda needs many more drops before shifting — it resists change. That’s buffering.
What you should see
Poured into the cabbage indicator, the cups fan out into a full rainbow: crimson in citric acid, through pinks and a neutral purple in water, to blue, teal, and a final yellow-green in limewater — a spectrum set purely by acidity. In the buffer test, the plain water flips colour after a drop or two, while the baking-soda cup stubbornly holds its colour for many more drops.
| Symptom | Likely cause | Fix |
|---|---|---|
| Colours look muddy/weak | Indicator too dilute or old | Boil a fresher, stronger batch; use more per cup |
| No colour difference between cups | Test solutions too dilute | Use stronger acid/base solutions |
| Turmeric doesn’t turn red | Solution not basic enough | Use a stronger base (sodium carbonate, limewater) |
| Buffer shifts as fast as water | Baking soda too dilute | Use a stronger baking-soda solution |
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 dangerous — the acid would crash blood pH and stop enzyme function.
Chemicals used in this experiment:
- Red Cabbage Juice — the anthocyanin indicator
- Citric Acid — strongly acidic test solution (pH ~2)
- Baking Soda — mildly basic solution and buffering demonstration
- Sodium Carbonate — strongly basic test solution (pH ~11)
- Calcium Hydroxide — very basic limewater (pH ~12)
Questions to Explore
Why does removing a proton change the colour so dramatically? How can losing something as small as an H⁺ change which wavelength the whole molecule absorbs?
Hint / answer
Removing a proton frees up electrons and reshapes the molecule’s cloud of alternating double bonds (its conjugation). Because colour is set by the energy gap those electrons absorb, shifting that gap shifts the absorbed wavelength — so a tiny proton has an outsized effect.
Why is pH logarithmic? pH 3 vs 4 is a tenfold H⁺ difference. Why use logs, and what would a linear scale look like?
Hint / answer
H⁺ concentrations range over fourteen powers of ten, so a log scale packs that into a tidy 0–14 where each step is ×10. A linear scale would need numbers from 1 down to 0.00000000000001 — unwieldy and impossible to plot sensibly.
What is buffering and why does blood need it? Blood holds pH 7.4. What happens to enzymes if it drifted to 6, and why is even a small shift dangerous?
Hint / answer
Buffers soak up added acid or base so pH barely moves. Enzymes are exquisitely shaped and only work in a narrow pH window; drop blood to 6 and they misfold and fail. Even 7.0 or 7.7 disrupts oxygen transport and nerve function enough to be a medical emergency — hence the tight bicarbonate buffering.
Why does turmeric give two colours when cabbage gives eight? What determines how many pH-sensitive equilibria a molecule has?
Hint / answer
The number of colours tracks the number of acidic groups that lose protons over the pH range. Anthocyanins have several –OH groups deprotonating at different pHs (many colours); curcumin has essentially one such switch, so just yellow ↔︎ red-brown.
How does an indicator “know” when to change? The dye is itself a weak acid at equilibrium with the solution. What must hold for it to report pH without changing it?
Hint / answer
The indicator’s own acid/base forms equilibrate to the solution’s pH, so its colour reflects that pH. This works only if you add so little indicator that it doesn’t itself shift the pH — a drop reads the solution rather than rewriting it.
Going further
- Map your house. Test shampoo, dish soap, antacid, sparkling water, milk, and tomato juice against your colour scale — where does each fall?
- Make test strips. Soak coffee-filter strips in the cabbage indicator and dry them for reusable pH paper.
- Full circle on the Acid–Base track: revisit the neutralization and gas experiments now that you can read pH by colour.