Wait, What? Two solutions can have the same concentration and very different conductivity.
That is because conductivity is not a direct concentration meter. It depends on how many mobile ions are present, what charges they carry, how easily they move, how completely a substance ionises, the temperature of the solution, and the geometry of the measuring cell.
The practical job is therefore not “dip a probe in and read a number.” It is to understand what the electrical signal actually represents and when conductivity can—or cannot—be used as evidence about ionic composition or concentration.
Why ionic solutions conduct
In a metallic conductor, electrons carry charge. In an electrolyte solution, positive and negative ions move through the liquid under an electric field and carry current.
A solid ionic compound may conduct poorly because ions are locked in a lattice. Dissolve or melt it and those ions become mobile. The Royal Society of Chemistry uses microscale conductivity tests to connect conductivity with the presence of mobile charged particles. See the RSC microscale conductivity practical.
Conductance is not conductivity
The instrument may directly measure electrical resistance R or conductance G:
G = 1/R
But conductance depends on electrode geometry. A wider electrode area or smaller electrode separation can increase conductance even for the same solution.
Conductivity κ attempts to remove this geometry dependence using a cell constant K:
κ = KG
The cell constant depends on probe geometry and is usually determined by calibration with a standard solution.
Why the probe must be immersed consistently
If electrode depth changes, the effective area of contact with solution can change. If the electrodes touch the beaker wall or are placed near sediment, the local electrical path can also differ.
Use a marked immersion depth or probe holder where possible. A digital display does not remove geometry from the measurement.
Temperature changes conductivity
As temperature rises, ion mobility in aqueous solutions usually increases and conductivity often rises. Therefore comparing samples at different temperatures can produce false concentration trends.
Commercial conductivity meters often use automatic temperature compensation, but that compensation depends on a model. For careful school work, measure or control temperature and know whether the displayed value has already been corrected.
More concentration usually means more conductivity—but not linearly forever
For a strong electrolyte at modest concentration, adding more dissolved electrolyte generally adds more charge carriers and increases conductivity. But ion-ion interactions grow at higher concentration, so simple linear proportionality can fail.
If you want to use conductivity as a concentration measurement, build a calibration curve over the concentration range of interest rather than assuming κ ∝ c universally.
Strong and weak electrolytes create different curves
A strong electrolyte such as sodium chloride is highly dissociated in dilute aqueous solution. A weak acid such as ethanoic acid is only partially ionised. Two 0.10 mol dm⁻³ solutions can therefore have very different numbers and types of mobile ions.
This is why “same molarity = same conductivity” is wrong. Conductivity depends on the ionic species actually present, not simply on the formula concentration written on the bottle.
Dilution can make molar conductivity rise even while conductivity falls
Bulk conductivity κ generally decreases as a solution is diluted because there are fewer ions per unit volume. But molar conductivity Λm is defined approximately as:
Λm = κ/c
For weak electrolytes, Λm can increase strongly on dilution because ionisation increases. This is a useful JC-level reminder that the quantity being graphed matters.
AC measurement reduces electrode chemistry
If direct current is applied for too long, electrolysis can occur at the electrodes, changing local ion concentrations and producing gas or reaction products. Conductivity instruments therefore commonly use alternating signals to reduce polarisation and net electrochemical change.
A simple school LED tester is excellent for qualitative comparisons, but a calibrated conductivity meter is needed for stronger quantitative work.
Rinsing prevents carryover
If a probe is moved from concentrated salt solution into distilled water without rinsing, droplets of the first solution contaminate the second. The error can be enormous because low-conductivity samples are especially sensitive to small ionic contamination.
Rinse with deionised or distilled water between samples and use the handling procedure appropriate to the probe. Do not wipe electrode surfaces in ways that damage coatings or alter geometry.
Calibration turns signal into a measurement
Use a conductivity standard with known value to calibrate or verify the cell response. One standard checks one point; multiple standards can test linearity over a wider range.
If a meter displays 1.000 mS cm⁻¹ on every solution but fails its standard, the apparent precision is meaningless. Calibration tests the measurement chain, not just the screen.
Quantitative window: a calibration curve
Suppose sodium chloride standards at 20 °C give:
- 0.010 mol dm⁻³ → 1.2 mS cm⁻¹
- 0.020 mol dm⁻³ → 2.3 mS cm⁻¹
- 0.030 mol dm⁻³ → 3.4 mS cm⁻¹
- 0.040 mol dm⁻³ → 4.4 mS cm⁻¹
An unknown reading of 2.9 mS cm⁻¹ can be interpolated within the calibration range to estimate concentration. But this calibration is valid for sodium chloride under similar temperature and instrument conditions. You cannot safely use it to determine magnesium sulfate concentration because ion charges and mobilities differ.
Conductivity and purity
Distilled water generally has much lower conductivity than a salt solution because it contains far fewer ions. This makes conductivity useful for comparing relative ionic purity of water samples.
But low conductivity does not prove complete chemical purity. Non-ionic contaminants such as some organic molecules may contribute little conductivity. The method detects ionic contamination much more strongly than many neutral substances.
Observation versus inference
Observation: “Sample A gave 5.6 mS cm⁻¹ and Sample B gave 0.08 mS cm⁻¹ at the same measured temperature using the same calibrated probe.”
Inference: “Sample A supports much greater ionic charge transport than Sample B under these conditions.”
Overclaim: “Sample A contains exactly seventy times more dissolved particles.” Conductivity is not a universal particle counter.
Failure modes that cap standards
- Different sample temperatures: mobility changes create false concentration differences.
- Probe not calibrated: displayed conductivity may share a systematic scale error.
- Carryover between samples: low-conductivity solutions become contaminated.
- Different immersion depth: effective cell geometry changes.
- Assuming conductivity equals concentration: ion identity and dissociation are ignored.
- Using DC for long periods: electrode polarisation/electrolysis can distort readings.
- Calibration outside the sample range: extrapolation becomes unreliable.
Unfamiliar transfer: titration monitored by conductivity
During some acid-base or precipitation titrations, conductivity changes because ions are consumed, replaced or produced. Plotting conductivity against titrant volume can reveal an endpoint without an indicator.
The same measurement logic transfers, but now conductivity is tracking a changing mixture of several ions rather than one concentration. Interpretation requires the reaction stoichiometry and relative ionic mobilities.
Secondary → JC → deeper Chemistry
Secondary: distinguish conducting and non-conducting substances, connect conductivity to mobile ions, and compare solutions fairly.
JC: calibrate conductivity, control temperature, interpret strong/weak electrolytes, distinguish κ from Λm, and use calibration curves within valid ranges.
Deeper Chemistry: conductometry extends to ionic mobility, Kohlrausch-type laws, activity, water-quality monitoring, conductometric titrations and electrochemical impedance methods.
Checkpoint 1: equal molarity
0.10 mol dm⁻³ NaCl conducts much better than 0.10 mol dm⁻³ ethanoic acid. Does this mean the NaCl solution contains more formula units per litre?
Checkpoint 2: warm sample
An unknown is measured at 30 °C using a calibration curve collected at 20 °C. What problem arises?
Answer key and WHY reasoning
Checkpoint 1: no. Both have the same analytical concentration, but NaCl is strongly dissociated while ethanoic acid is weakly ionised. The number and mobility of charged particles differ.
Checkpoint 2: conductivity is temperature-dependent, so the unknown cannot be compared directly with the 20 °C calibration without temperature control or validated compensation.
How to study this practical
For every conductivity result, ask five questions: which ions? how many? what charges? what temperature? what probe geometry/calibration? If you cannot answer those, do not convert the number into a concentration claim.
Evidence boundaries
Conductivity measures bulk electrical transport by mobile ions under the instrument conditions. It does not uniquely identify ions, total dissolved solids, chemical purity or concentration unless supported by calibration and composition assumptions.
Authoritative next steps
- Royal Society of Chemistry: Using a microscale conductivity meter
- Royal Society of Chemistry: conductivity as a check on distillate purity
- NIST Physical Measurement Laboratory
- SEAB A-Level syllabus directory
Teaching Guide
Give students equal-molar solutions of a strong electrolyte, weak electrolyte and non-electrolyte. Make them predict conductivity before measuring. Then warm one sample deliberately. The aim is to break the shortcut “more concentration = more conductivity” and replace it with charge-carrier reasoning.