Question

Why must \(\mathrm{NaCl}\) be melted before it is electrolyzed to give \(\mathrm{Na}\) and \(\mathrm{Cl}_{2}\) ? Write the anode, cathode, and overall cell reactions for the electrolysis of molten \(\mathrm{NaCl}\).

Short answer

Melting \(\mathrm{NaCl}\) is necessary because the ions must be free to move to conduct electricity. Anode reaction: \(2\mathrm{Cl}^- \rightarrow \mathrm{Cl}_{2} + 2e^-\), Cathode reaction: \(\mathrm{Na}^+ + e^- \rightarrow \mathrm{Na}\), Overall cell reaction: \(\mathrm{Na}^+ + \mathrm{Cl}^- \rightarrow \mathrm{Na} + \mathrm{Cl}_{2}\).

Step-by-step solution

Understanding the Necessity of Melting NaCl

Solid \(\mathrm{NaCl}\) is an ionic crystal where \(\mathrm{Na}^+\) and \(\mathrm{Cl}^-\) ions are locked in a rigid lattice structure. To be electrolyzed, these ions need to be free to move so that they can conduct electricity and travel to their respective electrodes. Melting \(\mathrm{NaCl}\) provides the necessary mobility for \(\mathrm{Na}^+\) and \(\mathrm{Cl}^-\) by disrupting the crystal lattice, thus allowing the electric current to pass through and chemical reactions to take place at the electrodes.

Identifying the Electrodes

In the electrolysis of molten \(\mathrm{NaCl}\), the \(\mathrm{Na}^+\) ions move towards the cathode, where they gain electrons (reduction), and \(\mathrm{Cl}^-\) ions move towards the anode, where they lose electrons (oxidation). The cathode is the site of reduction and the anode is the site of oxidation.

Writing the Cathode Reaction

At the cathode, the \(\mathrm{Na}^+\) ions are reduced to \(\mathrm{Na}\) metal. The half-reaction is: \[\mathrm{Na}^+ + e^- \rightarrow \mathrm{Na}\]

Writing the Anode Reaction

At the anode, \(\mathrm{Cl}^-\) ions are oxidized to \(\mathrm{Cl}_{2}\) gas. The half-reaction can be written as follows: \[2\mathrm{Cl}^- \rightarrow \mathrm{Cl}_{2} + 2e^-\] Note that two \(\mathrm{Cl}^-\) ions are needed for each molecule of \(\mathrm{Cl}_{2}\) produced.

Deriving the Overall Cell Reaction

By combining the half-reactions at the cathode and anode, we can write the overall cell reaction for the electrolysis of molten \(\mathrm{NaCl}\). The overall reaction is obtained by summing the two half-reactions, making sure to balance the number of electrons. The balanced overall cell reaction is: \[\mathrm{Na}^+ + \mathrm{Cl}^- \rightarrow \mathrm{Na} + \mathrm{Cl}_{2}\]

Key concepts

Electrochemical Cells

When we talk about electrolysis, we're delving into the workings of an electrochemical cell. An electrochemical cell is essentially a device capable of either generating electrical energy from chemical reactions or utilizing electrical energy to cause chemical changes. The latter is the case with the electrolysis of molten sodium chloride (NaCl).

During this process, the cell converts electrical energy, provided by an external power source, into chemical energy by inducing a chemical reaction that wouldn't occur spontaneously. This is an example of a non-spontaneous reaction harnessed to perform useful tasks, such as refining metals or producing chemicals like chlorine gas. The critical requirement for such a process is a conductive medium where ions are free to move, hence the melting of NaCl.

Ionic Lattice Structure

If we zoom in on a grain of table salt, we'll find an ionic lattice structure, a neat and tidy arrangement of alternating positively charged sodium (Na⁺) and negatively charged chloride (Cl⁻) ions. This arrangement is stable and strong, given that opposite charges attract, effectively locking the ions in place.

However, this stability is a double-edged sword; while it gives the salt crystal its rigidity, it also prevents the ions from moving freely. Thus, to dissolve this crystal fortress and free up the ions, NaCl must be heated to high temperatures where it melts. Once molten, the ions are liberated and can move - a mandatory state of affairs for electrolysis to transpire.

Reduction and Oxidation Reactions

At the heart of electrochemical cells are two pivotal reactions: reduction and oxidation, commonly referred to as redox reactions. In the simplest terms, reduction means gaining electrons, while oxidation involves losing them. Every electron given up by one species must be accepted by another, maintaining a balance.

For the electrolysis of molten NaCl, the sodium ions (Na⁺) undergo reduction, each gaining an electron to become neutral sodium metal (Na). Simultaneously, chloride ions (Cl⁻) are oxidized, losing electrons to form chlorine gas (Cl₂). These half-reactions occur at different electrodes, which in tandem drive the entire process.

Half-Reactions and Overall Cell Reaction

The beauty of electrochemistry lies in its neat bookkeeping. Each electron's journey is accounted for in half-reactions that detail the specific changes at the cathode and anode. For instance, the reduction of Na⁺ to Na at the cathode involves a single electron per sodium ion. Conversely, the oxidation of Cl⁻ into Cl₂ gas at the anode sees two chloride ions lose a total of two electrons.

When we tally up these half-reactions, we arrive at the overall cell reaction, which holds no excesses - the electrons lost equal the electrons gained. This meticulous accounting reflects the conservation of charge, an inviolable law in chemistry.

Conductivity in Molten Salts

Electrical conductivity in ionic compounds like NaCl is fascinating. As a solid, NaCl's ionic lattice structure is too rigid to allow free movement of ions, making it a poor conductor. But upon melting, these ions dissociate and are free to roam. Suddenly, molten NaCl is transformed into an electrical conductor, a dynamic playground for free ions.

This ability to conduct electricity in the molten state is the linchpin in the electrolysis process, allowing for the transport of Na⁺ to the cathode and Cl⁻ to the anode. As such, the essential pre-electrolysis melting of salt is all about creating an environment in which ions are no longer spectators, but active participants in the chemical ballet of electrolysis.