Unit 3: Electrochemistry

Learning Outcomes for Unit 3: Electrochemistry

Upon completion of this unit, students should be able to:

• Describe the electrolysis process in detail.
• Define oxidation-reduction (redox) reactions.
• Explain the roles and characteristics of oxidants and reductants.
• Utilize standard reduction potentials to predict oxidation-reduction reactions within an electrolytic cell.
• Predict products at the anode and cathode for the electrolysis of molten salts.
• Predict products at the anode and cathode for the electrolysis of aqueous metal salts when using graphite or platinum electrodes.
• Write accurate half-reaction equations.
• Write comprehensive cell reaction equations.
• Predict outcomes when reactive metals are employed as electrodes in the electrolysis process.
• Discuss the principles and applications of the electroplating process.

Introduction to Redox and Electrolysis

• Oxidation-Reduction Reactions: These processes involve the gaining and losing of electrons. Since electric current is defined as the flow of electrons, chemical energy can be converted to electrical energy, and electrical energy can used to drive chemical reactions that do not proceed spontaneously.
• Electrolysis Definition: The process in which electrical energy is used to produce a chemical change. It forces a non-spontaneous chemical reaction to occur.
• Environment for Electrolysis: Electrolysis only takes place in environments containing free charged ions. These ions serve as carriers of electrical charge between the two electrodes in a liquid or solution.
• Current in Electrodes: Within the electrodes themselves, current moves via free (delocalized) electrons found in the structure of the conductors.
• Electrolytes: A solution or liquid that conducts electricity. The most suitable electrolytes are aqueous solutions of ionic compounds or molten salts of ionic compounds. Solid ionic compounds are non-conductive because they lack free-moving charged particles or electrons.
• The Electrolytic Cell: A typical cell consists of:     * Two electrodes (made of chemically inert materials like platinum/graphite or chemically active metals).     * A suitable electrolyte.     * An AC or DC power source.
• Electrode Polarity and Terminology:     * The electrodes are dipped into the electrolyte and connected to the power source.     * Current flows out of the negative terminal and enters the positive terminal.     * Cathode: The electrode connected to the negative terminal; it is negatively charged (excess electrons) and is a "cation lover," attracting positively charged ions. Reduction occurs here.     * Anode: The electrode connected to the positive terminal; it is positively charged (deficiency of electrons) and is an "anion lover," attracting negatively charged ions. Oxidation occurs here.
• Ionic Behavior at Electrodes:     * At the anode (positive): Negatively charged ions deposit electrons and are oxidized.     * At the cathode (negative): Positively charged ions pick up electrons and are reduced.

Oxidation - Reduction (Redox) Reaction Concepts

• Redox Definition: A reaction process where electrons are gained or lost by the chemical species (ions) involved.
• Oxidation: The process of losing electrons. The species that loses electrons is oxidized and is called a reducing agent or reductant. In oxidation, the oxidation state number increases. Example: $Cu^{+}(s) \rightarrow Cu^{2+}(aq) + e^{-}$.
• Reduction: The process of gaining electrons. The species that gains electrons is reduced and is called an oxidizing agent or oxidant. In reduction, the oxidation state number decreases. Example: $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$.
• Half-Reactions: To determine agents, it is useful to break the global equation into two half-reactions.     * Example: Reaction between Calcium and Fluorine ($Ca(s) + F_2(g) \rightarrow CaF_2(s)$).     * Oxidation Half: $Ca \rightarrow Ca^{2+} + 2e^{-}$ ($Ca$ is the reductant).     * Reduction Half: $F_2 + 2e^{-} \rightarrow 2F^{-}$ ($F_2$ is the oxidant).

Rules for Writing and Balancing Reaction Equations

• Electrons in Half-Reactions: For reduction (gaining), electrons are written on the left-hand side. For oxidation (losing), electrons are written on the right-hand side.
• Combining Half-Reactions: The number of electrons gained must equal the number of electrons lost. If they differ, multiply the equations by integers so the electrons cancel out. The final balanced equation must not contain electrons.     * Example 3 Balancing:     * $5\times(Fe \rightarrow Fe^{2+} + 2e)$     * $2\times(MnO_4^{-} + 8H^{+} + 5e \rightarrow Mn^{2+} + 4H_2O)$     * Cell Reaction: $5Fe + 2MnO_4^{-} + 16H^{+} \rightarrow 5Fe^{2+} + 2Mn^{2+} + 8H_2O$

Standard Reduction Potentials

• Function of Reduction Potentials: These values ($E^{\circ}$) are used to predict which cation will be reduced and which anion will be oxidized when multiple species are present.
• Oxidizing Strength: Increases as the reduction potential becomes more positive (e.g., $F_2$ at $+2.87V$ is a very strong oxidant).

Table 1: Selected Standard Reduction Potentials

• Predicting Products at the Cathode: The cation with the higher (more positive) reduction potential will be reduced. Cations like $Li^{+}$, $K^{+}$, $Ca^{2+}$, $Na^{+}$, $Mg^{2+}$, $Al^{3+}$, $Zn^{2+}$, etc., have lower reduction potentials than water/hydrogen ions, so $H_2$ gas is produced instead of the metal. Cations like $Cu^{2+}$ and $Ag^{+}$ have higher potentials than water and will be reduced to their metal forms.
• Predicting Products at the Anode: The anion with the higher oxidizing strength (more reactive anion) will lose electrons. While the rule for cations is based on higher reduction potential value, for anions, those that oxidize more easily (like $Cl^{-}$, $Br^{-}$, $I^{-}$) involve different relative ease of oxidation.     * Anion Competition: $Br^{-}$, $Cl^{-}$, and $I^{-}$ oxidize more easily than $OH^{-}$. However, in very dilute solutions, $OH^{-}$ is preferentially oxidized to $O_2$ gas.     * Inert Anions: $SO_4^{2-}$, $NO_3^{-}$, $ClO_4^{-}$, $CO_3^{2-}$, and $PO_4^{3-}$ are not easily oxidized; in their presence, $OH^{-}$ will oxidize to produce $O_2$.

Electrolysis of Molten Salts

• Case Study: Molten Sodium Chloride ($NaCl$):     * Cell consists of a container for molten salt and inert electrodes (graphite or metal).     * Transmission of charge occurs via $Na^{+}$ and $Cl^{-}$ ions.     * Cathode (Reduction): $Na^{+}(l) + e^{-} \rightarrow Na(s)$. The sodium particles float to the surface.     * Anode (Oxidation): $2Cl^{-}(l) \rightarrow Cl_2(g) + 2e^{-}$. Chlorine gas bubbles out.     * Overall Cell Reaction: $2Na^{+}(l) + 2Cl^{-}(l) \rightarrow 2Na(s) + Cl_2(g)$.

Electrolysis of Aqueous Solutions

• Role of Water: In aqueous solutions, water can be oxidized or reduced. Water self-ionizes into hydronium ($H_3O^{+}$) and hydroxide ($OH^{-}$) ions ($H_2O \rightleftharpoons H^{+} + OH^{-}$).
• Aqueous Sodium Chloride (Brine):     * Ions present: $Na^{+}$, $H^{+}$, $Cl^{-}$, $OH^{-}$.     * Cathode: $2H^{+}(aq) + 2e^{-} \rightarrow H_2(g)$. $H^{+}$ is reduced because $Na^{+}$ is more reactive and prefers to stay in solution.     * Anode: $2Cl^{-}(aq) \rightarrow Cl_2(g) + 2e^{-}$. $Cl^{-}$ is more electronegative and withdraws electrons. (Note: in dilute solutions, $O_2$ is produced instead).
• Aqueous Sodium Sulphate ($Na_2SO_4$):     * Anode: $4OH^{-}(aq) \rightarrow 2H_2O(l) + O_2(g) + 4e^{-}$ ($OH^{-}$ oxidizes easier than $SO_4^{2-}$).     * Cathode: $2H^{+}(aq) + 2e^{-} \rightarrow H_2(g)$.
• Aqueous Copper (II) Sulphate ($CuSO_4$) with Platinum Electrodes:     * Ions: $Cu^{2+}$, $SO_4^{2-}$, $H^{+}$, $OH^{-}$.     * Cathode: $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$ ($Cu$ is less reactive than $H$).     * Anode: $4OH^{-}(aq) \rightarrow 2H_2O(l) + O_2(g) + 4e^{-}$.

Electrolysis of Dilute Acids (Acidified Water)

• Apparatus: Often carried out in a Hofmann Voltmeter.
• Process: $H_2SO_4$ dissociates to provide $H^{+}$ and $SO_4^{2-}$.
• Cathode: $2H^{+}(aq) + 2e^{-} \rightarrow H_2(g)$.
• Anode: $4OH^{-}(aq) \rightarrow 2H_2O(l) + O_2(g) + 4e^{-}$.
• Overall Reaction: $4H^{+}(aq) + 4OH^{-}(aq) \rightarrow 2H_2(g) + 2H_2O(l) + O_2(g)$.

Industrial Applications

• Production of Sodium Hydroxide: Concentrated brine is electrolyzed. As $H_2$ and $Cl_2$ leave as gases, the remaining ions are $Na^{+}$ and $OH^{-}$, forming $NaOH$.     * Uses of $NaOH$: Soap, detergents, paper, rayon, and bauxite purification.     * Uses of $Cl_2$: Bleach and water disinfectants.     * Uses of $H_2$: Fuel for pumping brine and heating the solution.
• Purifying Impure Metals (Electrolytic Refining):     * Blistered Copper (98% pure): Used as the anode. Pure copper is used as the cathode. Electrolyte is copper sulphate solution.     * Anode Reaction: $Cu(impure) \rightarrow Cu^{2+}(aq) + 2e^{-}$. The anode disintegrates/thins. Impurities drop as "slime" (containing Ag, Au, As, Fe).     * Cathode Reaction: $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(pure)$. The cathode thickens with 99.99% pure copper.     * Aluminum: Pure aluminum is obtained similarly from bauxite.
• Electroplating: Coating a metal with another (e.g., tin on steel, zinc on iron to prevent rusting).     * Object to be Plated: Becomes the cathode.     * Metal to be Deposited: Becomes the anode.     * Electrolyte: Must be a solution of the metal to be deposited (e.g., tin chloride for tin plating).

Questions and Discussion (Student Learning Activities Feedback)

• Testing Gases:     * Oxygen: Test with a glowing splint; if the gas produced at the anode burns with a bright flame/reignites, it is oxygen.     * Hydrogen: Test with a burning splint; if it produces a "pop" sound at the cathode, it is hydrogen.
• Indicator Changes: In the electrolysis of $NaCl$, if universal indicator is added, it turns blue. This occurs because as $H^{+}$ and $Cl^{-}$ are removed, $Na^{+}$ and $OH^{-}$ ions remain, making the solution alkaline.
• Reasoning for Copper Purity: Copper for electrical wiring must be 99.99% pure to minimize electrical resistance, which would otherwise cause electrical energy loss as heat.
• Reasoning for Acidification: Electrolytes like copper sulphate are often acidified to provide more free charged ions, speeding up the transmission of electrical charge.