Conductivity in chemistry is a measure of how well a material allows electric charge to flow through it. It is the inverse of resistivity: the lower a material’s resistance, the higher its conductivity. In chemistry specifically, the focus is usually on solutions, where dissolved ions carry the charge, rather than on metals, where electrons do the work.
How Conductivity Works in Solutions
When you dissolve a substance like table salt in water, it splits into positively charged sodium ions and negatively charged chloride ions. Apply a voltage across that solution, and those ions start moving: positive ions drift toward the negative electrode, negative ions drift toward the positive one. That movement of charged particles is electric current, and the solution’s ability to support it is its conductivity.
This is fundamentally different from what happens in a copper wire. In metals, a “sea” of free electrons flows through a fixed lattice of atoms. In a solution, the charge carriers are ions, which are much larger and slower than electrons. That’s why even a highly conductive salt solution conducts far less effectively than a metal. The two types are sometimes called electronic (metallic) conduction and electrolytic (ionic) conduction.
Units and How It’s Measured
The SI unit of conductivity is siemens per meter (S/m). For the dilute solutions chemists typically work with, microsiemens per centimeter (µS/cm) or millisiemens per centimeter (mS/cm) are more practical. Pure copper at 25 °C, for reference, has a conductivity of about 5.8 × 10⁷ S/m. Pure water sits around 0.05 µS/cm, while seawater lands near 50,000 µS/cm.
To measure conductivity in a lab, you dip a conductivity cell (two electrodes at a fixed distance apart) into the solution. The meter reads the electrical resistance between the electrodes and converts it to conductance using a simple relationship: conductance equals 1 divided by resistance. That raw number then gets multiplied by something called the cell constant, which accounts for the electrode geometry (specifically, the distance between the electrodes divided by their area). The result is the specific conductivity of the solution, comparable across different instruments and cell designs.
Molar Conductivity and Concentration
Specific conductivity tells you how well a particular solution conducts, but it doesn’t tell you how efficiently each dissolved unit of a substance contributes to that conduction. That’s what molar conductivity captures: the conducting power of all the ions produced by one mole of electrolyte in a given volume of solution.
Molar conductivity changes with concentration, and it changes differently depending on whether the electrolyte is strong or weak.
- Strong electrolytes (like sodium chloride or potassium nitrate) dissociate completely in water. Their molar conductivity is already high at moderate concentrations and increases slowly as you dilute further. The increase comes from ions interacting less with each other at lower concentrations, so they move more freely. If you plot molar conductivity against the square root of concentration, you get a straight line. The y-intercept of that line gives you the limiting molar conductivity: the theoretical maximum when the solution is infinitely dilute.
- Weak electrolytes (like acetic acid) only partially dissociate. At higher concentrations, most molecules remain intact and contribute nothing to conduction. As you dilute the solution, more molecules split into ions, so molar conductivity rises steeply. The plot is curved, not a straight line, and the limiting molar conductivity can’t be found by simply extending the graph.
Kohlrausch’s Law
For weak electrolytes, you need another approach to find limiting molar conductivity. That’s where Kohlrausch’s law comes in. Friedrich Kohlrausch discovered that at infinite dilution, each type of ion contributes independently to the total conductivity, regardless of what other ion it’s paired with. The limiting molar conductivity of any electrolyte equals the sum of the individual contributions of its cations and anions, each multiplied by how many of that ion the formula produces.
This is useful because you can look up the individual ion conductivities (which are measured from strong electrolytes, where the straight-line extrapolation works) and add them together to get the limiting value for a weak electrolyte you can’t measure directly.
Why Hydrogen Ions Conduct So Well
If you compare the conductivities of different ions in water, hydrogen ions and hydroxide ions stand out. They conduct far better than you’d expect from their size alone. The reason is a special transport process sometimes called the Grotthuss mechanism. Instead of a single hydrogen ion physically traveling from one electrode to the other, it hops to a neighboring water molecule, which then passes a different proton to the next water molecule, and so on down the chain. The net effect is that positive charge moves through the solution faster than any individual proton does, like a bucket brigade passing water rather than one person running back and forth. This chain-like relay makes hydrogen ions the fastest-conducting species in aqueous solution.
Temperature Effects
Temperature influences conductivity in opposite directions depending on the type of conductor. In metals, heating increases the vibration of atoms in the lattice, which scatters electrons and reduces conductivity. In electrolyte solutions, heating does the opposite: it lowers the viscosity of the solvent and gives ions more kinetic energy, so they move faster and conductivity goes up. For ionic conductors, this temperature dependence often follows an exponential pattern where conductivity rises as the energy barrier for ion movement becomes easier to overcome at higher temperatures.
Practical Applications
Water Quality Testing
One of the most common real-world uses of conductivity is checking water quality. Because dissolved salts, minerals, and contaminants all produce ions, a conductivity reading gives a quick estimate of total dissolved solids (TDS). Field meters convert conductivity to TDS using a conversion factor that ranges from about 0.50 at low conductivity (100 µS/cm, yielding roughly 50 ppm TDS) to around 0.80 at high conductivity (7,000 µS/cm, yielding about 5,600 ppm TDS). The factor isn’t constant because different ion mixtures contribute differently to both conductivity and mass.
Drinking water typically falls between 50 and 1,500 µS/cm. Distilled or deionized water used in laboratories has conductivity near zero, which is exactly the point: if conductivity is low, very little is dissolved in it. Industrial processes, pharmaceutical manufacturing, and aquarium keeping all rely on conductivity measurements to monitor water purity in real time.
Chemical Analysis
In the lab, conductivity measurements help track reactions that produce or consume ions. Conductometric titrations, for instance, monitor how conductivity changes as you add a reagent drop by drop. The endpoint of the reaction shows up as a sharp change in the conductivity curve, which can be more precise than a color-change indicator for certain reactions. Conductivity is also used to assess the purity of chemicals, the concentration of salt solutions, and the degree of dissociation of weak acids and bases.