pH & Indicators
Loading simulation…
flagWhat you'll discover
- arrow_forwardRank common household and lab solutions on the pH scale
- arrow_forwardChoose the right pH tool for the job and state each one's precision
- arrow_forwardExplain the logarithmic nature of the pH scale
- arrow_forwardDemonstrate buffer action and explain it with the H₂CO₃/HCO₃⁻ pair
- arrow_forwardInterpret universal indicator colours across the full 0-14 range
Four tools, four levels of precision
Litmus is the coarsest tool: red below about pH 5, blue above about 8 — it answers only "acid or base?" and says nothing in between. pH paper is litmus's educated cousin, impregnated with a mixture of dyes: match the strip's colour to the chart and you read pH to roughly the nearest whole number. Universal indicator is the same dye mixture in solution, painting the whole range — red (strong acid) through orange, yellow, green (neutral), blue, to violet (strong alkali).
The digital pH meter is the precision instrument: a glass electrode develops a voltage proportional to H⁺ concentration, displayed to 0.01 pH units. But precision needs care — the meter must first be calibrated against standard buffer solutions (pH 4.0, 7.0, 9.2), the bulb kept wet, and readings taken at a known temperature. An uncalibrated meter is confidently, precisely wrong.
The scale is logarithmic
pH = −log₁₀[H⁺], so each pH unit is a factor of ten in hydrogen-ion concentration. Lemon juice at pH 2.4 is not "a bit" more acidic than milk at 6.6 — it has over ten thousand times the H⁺ concentration. This is why mixing equal volumes of pH 1 acid and pH 13 alkali gives a violent reaction while mixing pH 6 and pH 8 solutions barely warms.
Landmarks worth memorising: dilute HCl (0.1 M) pH 1, lemon juice 2.4, vinegar 2.9, milk 6.6, pure water 7.0, blood 7.4, baking soda solution 8.4, soap 9-10, lime water 11, 0.1 M NaOH 13. Note that "neutral" is 7.0 only at 25 °C — pure water's pH drifts slightly with temperature because the ionisation of water itself is temperature-dependent.
Buffers: solutions that fight back
Add a drop of strong acid to pure water and the pH crashes by several units. Add the same drop to blood and the pH barely twitches. Blood is a buffer — a mixture (here carbonic acid H₂CO₃ and hydrogencarbonate HCO₃⁻) containing both a weak acid and its conjugate base, each standing guard. Added H⁺ is absorbed by the base: H⁺ + HCO₃⁻ → H₂CO₃. Added OH⁻ is absorbed by the acid: OH⁻ + H₂CO₃ → HCO₃⁻ + H₂O. Either way, free H⁺ hardly changes.
Your blood must hold pH 7.35-7.45; outside roughly 6.8-7.8 enzymes denature and cells die, so the buffer is life-critical — backed up by the lungs (exhaling CO₂) and kidneys. Buffers fill labs and industry too: calibration standards for pH meters, fermentation tanks, and the acetic acid/acetate buffer you will meet in volumetric analysis.