1.4 Electrochemical Systems and Components

Electrochemical cells are the backbone of many modern technologies. They convert chemical energy into electrical energy through redox reactions. Understanding their components and how they work is key to grasping the principles of electrochemistry.

These cells consist of electrodes, electrolytes, and a salt bridge. Electrons flow from the anode to the cathode, creating an electric current. The cell's voltage depends on the electrode materials and can be calculated using standard electrode potentials.

Electrochemical Cell Components

Structure and function of electrodes

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• Electrodes are conductive solid surfaces where redox reactions occur
  • Typically made of metals (copper, zinc) or carbon-based materials (graphite)
  • Provide a surface for electron transfer between the electrode and the electrolyte
• Two types of electrodes in an electrochemical cell
  • Anode undergoes oxidation, releasing electrons into the external circuit
  • Cathode undergoes reduction, accepting electrons from the external circuit
• Electrodes are connected to an external circuit (wires) to facilitate electron flow
• Electrode surface area affects the rate of the electrochemical reaction
  • Larger surface area allows for more reaction sites and faster reaction rates (porous electrodes)

Role of electrolytes in charge transfer

• Electrolytes are solutions that contain ions, allowing for electrical conductivity
  • Commonly aqueous solutions of salts (NaCl), acids (HCl), or bases (NaOH)
• Ions in the electrolyte carry charge between the electrodes
  • Cations (positively charged ions) move towards the cathode (Na+)
  • Anions (negatively charged ions) move towards the anode (Cl-)
• Electrolytes maintain charge balance in the cell
  • As electrons flow through the external circuit, ions migrate to maintain neutrality
• Concentration of the electrolyte affects the conductivity and reaction rate
  • Higher concentration generally leads to higher conductivity and faster reaction rates

Anodes vs cathodes in electrochemistry

• Anode is the electrode where oxidation occurs
  • Oxidation is the loss of electrons
  • Anode is considered the negative electrode
  • Example: In a zinc-copper cell, zinc is oxidized at the anode ($Zn \rightarrow Zn^{2+} + 2e^-$)
• Cathode is the electrode where reduction occurs
  • Reduction is the gain of electrons
  • Cathode is considered the positive electrode
  • Example: In a zinc-copper cell, copper is reduced at the cathode ($Cu^{2+} + 2e^- \rightarrow Cu$)
• In an electrochemical cell, electrons flow from the anode to the cathode through the external circuit

Components of electrochemical cells

• A standard electrochemical cell consists of:
  1. Two half-cells, each containing an electrode immersed in an electrolyte solution
  2. A salt bridge connecting the two half-cells
  3. An external circuit connecting the electrodes
• Half-cells are compartments that separate the oxidation and reduction reactions
  • Each half-cell contains an electrode and an electrolyte specific to the reaction
• Salt bridge is a connection between the two half-cells
  • Typically a tube filled with an electrolyte solution (potassium chloride)
  • Allows for the flow of ions between the half-cells to maintain charge balance
  • Prevents the mixing of the electrolyte solutions in the half-cells
• External circuit is the connection between the electrodes
  • Allows for the flow of electrons from the anode to the cathode
  • Can include a load (light bulb, resistor) to utilize the electrical energy generated by the cell

Electrochemical Reactions

Oxidation and reduction processes

• Oxidation is the loss of electrons
  • Occurs at the anode
  • Example: $Zn \rightarrow Zn^{2+} + 2e^-$
• Reduction is the gain of electrons
  • Occurs at the cathode
  • Example: $Cu^{2+} + 2e^- \rightarrow Cu$
• Oxidation and reduction reactions always occur simultaneously in an electrochemical cell
  • The number of electrons lost in the oxidation reaction must equal the number of electrons gained in the reduction reaction
• Redox reactions can be represented by half-reactions
  • Oxidation half-reaction: shows the species being oxidized and the electrons released
  • Reduction half-reaction: shows the species being reduced and the electrons consumed
• Overall cell reaction is the sum of the oxidation and reduction half-reactions
  • Electrons cancel out as they are transferred from the oxidized species to the reduced species

Standard electrode potentials and cell voltage

• Standard electrode potential ($E^0$) is the measure of the tendency of a half-reaction to occur under standard conditions
  • Standard conditions: 25℃, 1 M concentration for solutes, 1 atm pressure for gases
  • Measured relative to the standard hydrogen electrode (SHE), which has an assigned potential of 0 V
• Each half-reaction has a specific standard electrode potential
  • More positive $E^0$ values indicate a greater tendency for reduction
  • More negative $E^0$ values indicate a greater tendency for oxidation
• Standard cell potential ($E^0_{cell}$) is the difference between the standard electrode potentials of the cathode and anode
  • $E^0_{cell} = E^0_{cathode} - E^0_{anode}$
  • Positive $E^0_{cell}$ indicates a spontaneous reaction under standard conditions
• Nernst equation relates the cell potential to the concentrations of the reactants and products
  • $E_{cell} = E^0_{cell} - \frac{RT}{nF} \ln Q$
  • $R$: gas constant, $T$: temperature (K), $n$: number of electrons transferred, $F$: Faraday's constant, $Q$: reaction quotient

Gibbs free energy and cell potential

• Gibbs free energy ($\Delta G$) is a measure of the spontaneity of a reaction
  • Negative $\Delta G$ indicates a spontaneous reaction
  • Positive $\Delta G$ indicates a non-spontaneous reaction
• Relationship between Gibbs free energy and cell potential:
  • $\Delta G = -nFE_{cell}$
  • $n$: number of electrons transferred, $F$: Faraday's constant, $E_{cell}$: cell potential
• Under standard conditions:
  • $\Delta G^0 = -nFE^0_{cell}$
• A positive cell potential ($E_{cell} > 0$) corresponds to a negative Gibbs free energy ($\Delta G < 0$), indicating a spontaneous reaction
• A negative cell potential ($E_{cell} < 0$) corresponds to a positive Gibbs free energy ($\Delta G > 0$), indicating a non-spontaneous reaction