Electrical Conduction and Electrolytes

Conduction of Electricity: Charges, Media, and Mechanisms

Electrical conduction is the directed movement of electric charge. In chemistry, we are especially interested in conduction in liquids and solutions and in how this is different from conduction in metals.

In all cases, an electric current $I$ is defined as the amount of charge $Q$ passing through a cross-section per unit time $t$:
$$
I = \frac{Q}{t}
$$

What differs between materials is:

• What the charge carriers are
• How freely they can move
• What has to happen chemically so they can move

These differences are central for understanding electrolytes.

Types of Electrical Conduction

Electronic conduction (in metals and some solids)

• Charge carriers: Mainly electrons.
• Typical media: Metals, graphite, some semiconductors.
• Mechanism: Electrons move through a lattice of positive metal ions; the ions themselves remain fixed.
• Chemical change: None required for conduction. Passing a current through a metal wire does not change its composition.

This is the type of conduction familiar from simple electrical circuits.

Ionic conduction (in electrolytes)

• Charge carriers: Ions (cations and anions).
• Typical media:
• Molten (fused) ionic compounds (e.g. molten NaCl)
• Aqueous solutions of acids, bases, and salts (e.g. HCl(aq), NaOH(aq), NaCl(aq))
• Mechanism:
• Cations move toward the negative electrode (cathode).
• Anions move toward the positive electrode (anode).
• Chemical change: Often accompanied by chemical reactions at the electrodes; conduction is intimately linked to electrochemical processes.

In ionic conduction, matter is transported together with charge, in contrast to metals where mostly electrons move.

What Is an Electrolyte?

An electrolyte is a substance that:

• Produces mobile ions when molten or dissolved in a suitable solvent (typically water).
• Therefore allows the resulting liquid or solution to conduct electric current.

Common examples:

• Acids: HCl, HNO$_3$, H$_2$SO$_4$
• Bases: NaOH, KOH, Ca(OH)$_2$
• Salts: NaCl, KNO$_3$, CuSO$_4$

An electrolyte can be:

• Molten (fused): Pure ionic compound above its melting point.
• In solution: Usually dissolved in water.

Substances like glucose or ethanol, which dissolve in water but do not form ions, are non-electrolytes.

Electrolytes vs. Non-Electrolytes

When a substance is placed in water, three main behaviors are possible with regard to conduction:

1. Non-electrolytes
• Dissolve as neutral particles (molecules).
• No or negligible ions are present.
• Solution does not conduct electricity measurably.
• Example: Solutions of sugar (glucose), urea, many organic compounds without ionizable groups.
2. Strong electrolytes
• Dissolve and form ions almost completely.
• High concentration of ions → good conduction.
• Examples:
• Most soluble salts: NaCl, KNO$_3$
• Strong acids: HCl, HNO$_3$, HClO$_4$
• Strong bases: NaOH, KOH.
3. Weak electrolytes
• Dissociate into ions only partly.
• A significant fraction remains as neutral molecules.
• Moderate to low conduction.
• Examples:
• Weak acids: CH$_3$COOH (acetic acid)
• Weak bases: NH$_3$ (ammonia in water).

The difference between strong and weak electrolytes does not refer to “strong” or “weak” current, but to the degree of ion formation in solution.

Dissociation and Solvation

In aqueous solutions, electrolytic conduction is made possible by two key processes:

1. Dissociation (for ionic or ionizable substances)
• The substance separates into ions.
• Example: Sodium chloride in water
  $$
  \text{NaCl(s)} \rightarrow \text{Na}^+(aq) + \text{Cl}^-(aq)
  $$
2. Solvation (hydration in water)
• Water molecules surround the ions, stabilizing them in solution.
• Ions are written with $(aq)$ to indicate they are hydrated:
  $$
  \text{Na}^+ + \text{H}_2\text{O} \;\to\; \text{Na}^+(aq)
  $$

For molecular substances that act as acids or bases, ion formation can happen by reaction with water. Example: Ionization of acetic acid:
$$
\text{CH}_3\text{COOH} + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}_3\text{O}^+
$$

The more ions are produced, the better the solution conducts electricity.

Conductivity of Electrolyte Solutions

Conductivity describes how well a solution conducts electric current. It depends on:

1. Concentration of ions
• More ions → more charge carriers → higher conductivity.
2. Charge of the ions
• Ions with higher charge (e.g. Mg$^{2+}$, Al$^{3+}$) can carry more charge per particle.
3. Mobility of ions
• Smaller, less strongly solvated ions usually move more easily.
• Viscosity of the solvent and temperature also influence mobility.

Specific conductivity (conductance per length and area)

The specific conductivity (or simply conductivity) $\kappa$ of a solution is defined so that:

• High $\kappa$ → solution conducts well.
• Low $\kappa$ → solution conducts poorly.

You do not need the detailed formula here; conceptually:

• $\kappa$ increases with the number and mobility of ions present.

Molar conductivity (conceptual)

For comparing different electrolytes, it is often helpful to consider the conductivity per mole of dissolved substance (molar conductivity). Conceptually:

• If two solutions contain the same concentration of electrolyte:
• The one with stronger dissociation (strong electrolyte) typically has higher molar conductivity.
• Weak electrolytes have lower molar conductivity at the same nominal concentration because fewer ions are present.

Details of molar conductivity and its concentration dependence are treated in more advanced treatments of electrochemistry.

Factors Affecting Electrolytic Conduction

1. Nature of the electrolyte

• Strong electrolytes: High conductivity at a given concentration.
• Weak electrolytes: Lower conductivity because only a fraction of the molecules are ionized.

At the same concentration, 0.1 mol·L$^{-1}$ HCl(aq) will conduct much better than 0.1 mol·L$^{-1}$ CH$_3$COOH(aq) because HCl is almost fully ionized, whereas acetic acid is not.

2. Concentration of the electrolyte

• For dilute to moderate concentrations:
• Increasing concentration usually increases conductivity, because there are more ions.
• At higher concentrations:
• Interactions between ions (ion pairing, crowding) can reduce mobility, so conductivity does not increase proportionally and can even rise more slowly.

This makes the relationship between concentration and conductivity non-linear, especially at higher concentrations.

3. Temperature

Increasing temperature typically:

• Increases ionic mobility (lower viscosity of water, faster motion).
• Slightly changes the dissociation equilibria.

Thus, most aqueous electrolyte solutions become better conductors as temperature increases.

4. Solvent properties

• Dielectric constant: Water has a high dielectric constant, which favors separation of ions (dissociation).
• Viscosity: More viscous solvents hinder ionic motion; conductivity is lower.

Therefore, the same electrolyte can have very different conductivities in different solvents.

Conduction in Molten Salts vs. Solutions

Electrolytes can conduct in two distinct liquid forms:

Molten salts (fused salts)

• Example: Molten NaCl above its melting point.
• Composition: Only ions (Na$^+$ and Cl$^-$), no solvent.
• Conduction: Ions move through the liquid lattice.
• Often used in high-temperature electrolysis (e.g. production of metals).

Aqueous solutions of electrolytes

• Example: NaCl(aq) at room temperature.
• Composition: Ions + solvent molecules (e.g. water).
• Conduction: Ions move through the liquid, strongly influenced by solvation and solvent properties.

The fundamental mechanism—movement of ions—is the same, but details such as mobility and required temperature differ.

Simple Experimental Observations

In a beginner laboratory, differences between electrolytes and non-electrolytes can be demonstrated with:

• A conductivity tester (e.g. a lamp or LED with electrodes dipped into the liquid).

Observations:

• Distilled water: Very poor conductor (lamp stays dark or very dim).
• Salt solution (NaCl(aq)): Lamp lights brightly (good conduction).
• Sugar solution (glucose): Lamp remains dark (no significant conduction).
• Dilute acetic acid solution: Lamp glows weakly (weak electrolyte).
• Molten NaCl (at high temperature, in suitable apparatus): Conduction occurs; electrolysis products can be observed at electrodes.

These experiments clearly show that:

• The presence of ions is crucial for conduction in liquids.
• Not all substances that dissolve in water produce ions.
• The kind and degree of ion formation determine how well a solution conducts.

Role of Electrolytes in Electrochemical Processes

In electrochemical cells and electrolytic setups:

• An electrolyte provides the ionic connection between electrodes.
• Conduction in the external circuit is electronic (through wires).
• Conduction inside the cell or solution is ionic (through the electrolyte).

Electrolytes thus:

• Complete the electrical circuit via ion movement.
• Enable charge balance when redox reactions transfer electrons at the electrodes.
• Influence cell performance through their conductivity and ionic composition.

Details of electrodes, electrode potentials, and full electrochemical cells are covered in other chapters of this section. Here, it is sufficient to see electrolytes as the ion-conducting medium that makes such processes possible.