Acids, Bases and Indicators in Action
In 1909, Danish chemist Søren Sørensen invented the pH scale to save his 10,000-litre brewery fermentation vats from going sour.
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Red cabbage juice turns pink in lemon juice and green in baking soda solution. What do you think this colour-changing ability tells a scientist about those substances? Have you ever seen a natural substance change colour like this?
Before digital pH meters existed, chemists still needed to measure how acidic or basic a substance was. How do you think they did this? What limitations might a natural colour-based indicator have compared to a digital meter?
● Know
- That natural indicators such as red cabbage and turmeric change colour with pH
- How to safely test the pH of common substances
- How to record and present investigation results in a table
● Understand
- The advantages and limitations of natural vs synthetic indicators
- That reliable results require systematic testing and repeated observations
- How to link indicator colour changes to pH values
● Can do
- Conduct a practical investigation to test pH using indicators
- Compare natural and synthetic indicators for accuracy and ease of use
- Record results in a clear table and draw conclusions from data
Before electronic pH meters, chemists relied on indicators - natural or synthetic dyes that change colour depending on the acidity or basicity of the solution. The chemistry behind this is elegant and still relevant today.
Indicators are themselves weak acids or bases. In solution, they exist in equilibrium between two molecular forms: one that predominates in acid (HIn) and one that predominates in base (In-). These two forms have different colours because their electron structures differ, causing them to absorb different wavelengths of light.
When H+ concentration is high (low pH), the equilibrium shifts toward the acid form (HIn). When H+ concentration is low (high pH), the equilibrium shifts toward the base form (In-). The eye sees the colour of whichever form predominates.
Different indicators have different transition pH ranges. Litmus changes at pH 5-8. Phenolphthalein is colourless below pH 8.2 and pink above. Methyl orange is red below pH 3.1 and yellow above pH 4.4. By choosing the right indicator, chemists can detect specific pH ranges.
Phenolphthalein is a dramatic indicator used in many school titrations. In a flask containing hydrochloric acid, it is completely colourless. As sodium hydroxide is added drop by drop, the solution remains colourless until the very last drop that neutralises the final H+ ion. At that moment, the pH jumps from acidic to basic, and the solution suddenly turns bright pink. This sharp colour change makes phenolphthalein ideal for detecting the equivalence point in strong acid-strong base titrations.
Australian natural indicators: Indigenous Australians used natural indicators long before European contact. Red cabbage contains anthocyanins that change from red (acidic) to green (neutral) to yellow (basic). Many native Australian plants contain similar pigment molecules. While not as precise as modern indicators, these natural dyes allowed traditional cultures to assess water quality and prepare materials for cultural practices.
Indicators change colour because the acid destroys the dye. This is false. Indicators change colour through a reversible equilibrium between two molecular forms. If you add acid to turn an indicator red, then add base to turn it blue, then add acid again, it will turn red again. The indicator molecule is not consumed or destroyed - it cycles between forms depending on pH.
You have a colourless solution and add universal indicator. It turns orange. Predict: is the solution acidic, neutral, or basic? Then explain what is happening at the molecular level.
Orange indicates pH around 4-5, so the solution is weakly acidic. The indicator molecules change structure depending on H+ concentration, and the acidic form absorbs light differently from the basic form.
Use these terms in your explanation: indicator · H+ concentration · pH · molecular structure
Titration is a quantitative laboratory technique used to determine the unknown concentration of a solution by reacting it with a solution of known concentration. It is one of the most important analytical methods in chemistry.
The setup consists of: a burette (a graduated tube with a stopcock) containing the titrant (solution of known concentration); a conical flask containing the analyte (solution of unknown concentration) with a few drops of indicator; and a white tile underneath to make colour changes easier to see.
The titrant is added slowly to the analyte while swirling the flask. Near the endpoint, the titrant is added drop by drop. The endpoint is reached when the indicator permanently changes colour. If the indicator is chosen correctly, the endpoint coincides with the equivalence point where stoichiometrically equivalent amounts of acid and base have reacted.
The concentration of the analyte is calculated using the titration formula: C1V1 = C2V2 (for 1:1 reactions), where C is concentration and V is volume.
A student titrates 25.0 mL of unknown hydrochloric acid against 0.100 mol/L sodium hydroxide. The indicator changes colour after adding 18.5 mL of NaOH. The reaction is 1:1 (HCl + NaOH -> NaCl + H2O). Using C(acid) * V(acid) = C(base) * V(base): C(acid) * 25.0 = 0.100 * 18.5. Therefore C(acid) = 1.85 / 25.0 = 0.074 mol/L. The acid concentration is 0.074 mol/L. This precise quantitative result is why titration remains indispensable in analytical chemistry.
Australian analytical chemistry: The National Measurement Institute (NMI) uses titration and other analytical techniques to certify reference materials used by Australian laboratories. Wineries use titration to measure acidity in grapes and wine. Dairy factories titrate to determine fat content and acidity in milk. These applications ensure product quality and regulatory compliance across Australian food and beverage industries.
The endpoint and equivalence point are the same thing. They are close but not identical. The equivalence point is the theoretical point where stoichiometrically equivalent amounts have reacted. The endpoint is the experimental point where the indicator changes colour. A well-chosen indicator makes these points coincide closely, but they are conceptually distinct. For weak acid-strong base titrations, the equivalence point is above pH 7, so phenolphthalein (which changes at pH 8.2) is appropriate. Using methyl orange (which changes at pH 4) would give a very inaccurate result.
- Titration
- Burette
- Endpoint
- Equivalence point
- The point where moles of acid equal moles of base
- A technique to determine concentration by reacting with a solution of known concentration
- A graduated tube that delivers precise volumes of liquid
- The point where indicator changes colour
Good scientists record their results systematically so that others can understand and trust their findings. A well-organised results table is essential.
Tips for recording results
- Use a ruler to draw neat tables with clear headings.
- Include units in column headings where appropriate (e.g., "pH" not just "reading").
- Record what you actually see, not what you expect to see. If the colour is "pinkish-orange," write that rather than just "red."
- Be consistentdescribe colours using the same words each time, or better still, compare to a colour chart.
- Note any anomaliesunexpected results should be recorded, not ignored.
Drawing conclusions
After collecting your data, look for patterns:
- Which substances were acidic? Which were alkaline? Which were neutral?
- Did all three indicators agree on whether each substance was acidic or basic?
- Which indicator gave the most detailed information about pH?
- Did the pH meter readings match the indicator colours?
Wrong: "Natural indicators are less accurate than synthetic ones, so they are useless." No, natural indicators are perfectly valid for many purposes. Indigenous Australians have used natural indicators for thousands of years. The key is understanding the limitations of each tool.
Right: Natural indicators are genuinely useful, they have been used reliably for thousands of years. They have limitations (e.g., limited pH range or colour sensitivity), but so do synthetic indicators. Knowing the limitations of your tool is good science.
Wrong: "If an indicator does not change colour, the substance must be neutral." Not necessarily, some indicators only change in certain pH ranges. Turmeric does not change in acid, so a yellow result with turmeric could mean acid or neutral. You need multiple indicators to be sure.
Right: No colour change only means the pH is outside that particular indicator's response range. To determine whether a substance is truly neutral, you need to test with multiple indicators or use a pH meter.
Wrong: "You only need to test each substance once." No, repeating measurements and using multiple indicators improves reliability and helps identify mistakes or anomalies.
Right: Repeating tests and using multiple indicators improves reliability. A single result could be an error or anomaly, repeated consistent results give you confidence that your conclusion is correct.
Natural Indicators in Indigenous Knowledge
Aboriginal and Torres Strait Islander Peoples have long understood that certain plants can reveal properties of water and soil. Some plants that grow only in specific soil types act as indirect indicators of pH and mineral content. For example, certain species of eucalyptus prefer acidic soils, while others grow better in neutral or slightly alkaline conditions. Observing which plants grow where can give information about soil chemistry.
In Australian agriculture, soil pH testing is essential for healthy crops. Many Australian soils are naturally acidic, and farmers use lime (calcium carbonate, a base) to raise the pH. Regular pH testing with indicators or meters helps farmers decide when and how much lime to apply.
✍ Copy Into Your Books
▾Natural Indicators
- Red cabbage: pink (acid), purple (neutral), green/yellow (base)
- Turmeric: yellow (acid/neutral), reddish-brown (base)
- Made from plant pigments; vary in strength
Synthetic Indicators
- Litmus: red (acid), blue (base)
- Universal indicator: red to purple across pH 0–14
- More consistent and reliable than natural indicators
Recording Results
- Use clear tables with headings and units
- Record what you actually observe
- Repeat measurements for reliability
Indicator Comparison
Data Detective
At the start of this lesson, the hook reminded you that red cabbage juice changes through 11 different shades depending on how acidic or alkaline a liquid is, and that early chemists squeezed berries and flower petals to test their chemicals.
Now that you've used and compared different indicators, how would you explain to a friend why red cabbage juice works as a pH indicator? What limitations does it have compared to a digital pH meter, and how did your initial ideas compare with what you discovered?
Q1. 1. Describe the difference between qualitative and quantitative data. Give one example of each from the pH testing investigation. 4 MARKS
Q2. 2. A student wants to test whether the water from their rainwater tank is acidic. They have red cabbage indicator, universal indicator paper, and a pH meter. Design a systematic method they could use, including which tool(s) they should choose and why. 4 MARKS
Q3. 3. In the practical investigation, a student finds that red cabbage indicator turns slightly different colours for the same substance on two different days. Explain TWO factors that could cause this variation, and suggest how the student could improve the reliability of their results. 4 MARKS
Revisit Your Thinking
Go back to your Think First answer. Has your understanding changed?
- Would you now choose different equipment to test your five mystery liquids?
- What have you learned about recording results that you did not know before?
Model answers (click to reveal)
Answers
▾MCQ 1
BRed cabbage is a natural indicator. Universal indicator, litmus paper and pH meters are all synthetic or manufactured tools.
MCQ 2
DTurmeric only changes colour in alkaline conditions. If it stays yellow, the solution could be acidic or neutral, but it is definitely not alkaline.
MCQ 3
ASynthetic indicators are manufactured with controlled chemical composition, giving more consistent and reproducible results. Natural indicators can vary depending on how they are grown and prepared.
MCQ 4
CPink with red cabbage and orange with universal indicator both indicate an acidic pH. pH 4 is the only acidic option listed.
MCQ 5
BSubstance Y shows blue with universal indicator (alkaline) but pH 7.0 (neutral) on the meter. This is inconsistent and should be retested. Substance X (green + pH 7.0) and Substance Z (red + pH 2.5) are both consistent.
Short Answer 1
Model answer: Qualitative data describes qualities or characteristics that cannot be measured with numbers. For example, recording that universal indicator turned "red" is qualitative data. Quantitative data involves numerical measurements. For example, recording that the pH meter read 3.5 is quantitative data. Both types are important: qualitative data describes what happened, while quantitative data allows precise comparison.
Short Answer 2
Model answer: The student should use the pH meter as their primary tool because it gives a precise numerical reading, which is needed to know if the water is acidic (pH < 7) and by how much. They should also use universal indicator paper as a quick check to confirm the meter reading is reasonable. The method should be: (1) collect a sample of rainwater, (2) calibrate the pH meter, (3) rinse the probe and dip it into the sample, (4) record the reading, (5) dip universal indicator paper into the sample and compare to the colour chart, (6) record both results in a table, (7) repeat the measurement twice for reliability.
Short Answer 3
Model answer: Factor 1: The concentration of the red cabbage indicator may have differed between preparations, more concentrated indicator can produce deeper colours. Factor 2: Lighting conditions when observing colours can affect perception, natural light versus artificial light may make colours appear different. To improve reliability, the student should: (1) prepare the indicator using a standardised method each time, (2) observe colours under consistent lighting, (3) use a colour chart for comparison rather than relying on memory, and (4) repeat tests with the same substance multiple times and compare results.