CHM 101_ELECTROCHEMISTRY.pdf

# Introduction to electrochemistry and electrochemical cells Electrochemistry explores the fundamental connection between electricity and chemistry, focusing on how chemical reactions can either generate or consume electrical energy [3](#page=3). ### 1.1 Defining electrochemistry Electrochemistry is the scientific discipline that investigates the relationship between electricity and chemical reactions. It also examines the capacity of materials to facilitate the passage or transfer of electrical current [3](#page=3). ### 1.2 Electrochemical cells An electrochemical cell is a system composed of electrodes immersed in an electrolyte, where a chemical reaction either utilizes or produces an electric current [4](#page=4). #### 1.2.1 Voltaic (galvanic) cells Voltaic cells, also known as galvanic cells, are electrochemical cells where spontaneous chemical reactions produce electricity and supply it to an external circuit (#page=3, page=4) [3](#page=3) [4](#page=4). > **Tip:** Think of voltaic cells as energy generators, converting chemical energy into electrical energy through natural chemical processes. #### 1.2.2 Electrolytic cells Electrolytic cells are electrochemical cells in which an external source of electrical energy drives nonspontaneous chemical reactions (#page=3, page=4). In these cells, electrical energy is consumed to force a chemical change that would not occur naturally [3](#page=3) [4](#page=4). > **Example:** The electrolysis of a potassium iodide (KI) solution, which produces iodine and hydrogen gas, is an example of a process occurring in an electrolytic cell. This reaction is driven by the input of electrical energy [3](#page=3). --- # Electrical conduction and electrodes This topic explains the fundamental mechanisms of electrical conduction in metals and electrolytes, and defines the roles of electrodes (anode and cathode) in electrochemical cells. ### 2.1 Electrical conduction mechanisms Electric current is fundamentally the transfer of charge. This charge transfer can occur through two primary mechanisms: metallic conduction and ionic (electrolytic) conduction [5](#page=5). #### 2.1.1 Metallic conduction Metallic conduction is characteristic of metals and involves the flow of electrons. In this process, the atoms of the metal remain largely in place, and there are no discernible changes to the metal itself during conduction. Metals conduct electricity due to the presence of free, mobile electrons [5](#page=5). #### 2.1.2 Ionic conduction Ionic conduction, also known as electrolytic conduction, occurs when electric current is carried by the movement of ions through a solution or a pure liquid. In an electric field, positively charged ions (cations) migrate towards the negative electrode, while negatively charged ions (anions) move towards the positive electrode. This type of conduction relies on the presence of free, mobile ions [5](#page=5) [6](#page=6). > **Tip:** Both metallic and ionic conduction are important in the operation of electrochemical cells [5](#page=5). ### 2.2 Electrodes in electrochemical cells Electrodes serve as the surfaces where the critical half-reactions of oxidation and reduction take place within an electrochemical cell. These electrodes may or may not directly participate in the chemical reactions; those that do not react are termed inert electrodes [6](#page=6). #### 2.2.1 Anode and cathode Regardless of whether the cell is electrolytic or voltaic, electrodes are classified based on the electrochemical processes occurring at their surfaces [6](#page=6). * **Cathode:** The cathode is defined as the electrode where reduction occurs, meaning a species gains electrons [6](#page=6). * **Anode:** The anode is defined as the electrode where oxidation occurs, meaning a species loses electrons [6](#page=6). > **Tip:** It is important to remember that the cathode is always the site of reduction and the anode is always the site of oxidation. The sign (positive or negative) of these electrodes can vary depending on the type of electrochemical cell (electrolytic or voltaic) [6](#page=6). #### 2.2.2 Ion migration at electrodes In ionic conduction within a solution, the direction of ion movement is dictated by their charge and the applied electric field, leading them to the respective electrodes. Positively charged ions migrate towards the negative electrode, and negatively charged ions migrate towards the positive electrode. This movement of ions constitutes the electrolytic conduction [6](#page=6). --- # Voltaic (galvanic) cells and notation Voltaic cells, also known as galvanic cells, harness spontaneous oxidation-reduction (redox) reactions to generate electrical energy by physically separating the oxidation and reduction half-reactions [7](#page=7). ### 3.1 Construction and operation of voltaic cells A voltaic cell is constructed from two half-cells, each representing a half-reaction. A simple half-cell consists of a metal strip immersed in a solution containing its corresponding metal ions. For example, a zinc electrode comprises a zinc metal strip in a zinc salt solution, and a copper electrode consists of a copper metal strip in a copper salt solution [7](#page=7). These two half-cells are connected externally to allow electron flow and internally to permit ion flow, thus completing the electrical circuit. In a voltaic cell where zinc and copper are involved, zinc atoms lose electrons more readily than copper atoms. These electrons travel from the zinc electrode through the external circuit to the copper electrode. At the copper electrode, these electrons react with copper ions, causing copper metal to deposit. The overall process results in the reaction of zinc metal with copper ions to form zinc ions and copper metal, generating an electric current in the external circuit [7](#page=7). #### 3.1.1 Half-reactions and electrodes The two fundamental processes occurring in a voltaic cell are the half-reactions: * **Oxidation half-reaction:** A species loses electrons. The electrode where oxidation occurs is called the **anode**. For the zinc-copper cell, this is [9](#page=9): $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^{-}$ [9](#page=9). * **Reduction half-reaction:** A species gains electrons. The electrode where reduction occurs is called the **cathode**. For the zinc-copper cell, this is [9](#page=9): $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$ [9](#page=9). The definitions of anode and cathode apply to all electrochemical cells, including electrolytic cells [10](#page=10). The sum of the two half-reactions constitutes the overall **cell reaction**, which is the net reaction occurring in the voltaic cell. For the zinc-copper cell, the cell reaction is [10](#page=10): $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$ [10](#page=10). #### 3.1.2 The role of the salt bridge A crucial component of a voltaic cell is the salt bridge, which is a tube containing an electrolyte gel that connects the two half-cells. The salt bridge is essential because it allows the flow of ions between the half-cells but prevents the direct mixing of the electrolyte solutions [8](#page=8). If the ion solutions were to mix directly, a direct reaction would occur without the generation of an electric current. This would lead to a drop in voltage and a rapid depletion of the cell's energy. The salt bridge maintains electrical neutrality in each half-cell by allowing counter-ions to migrate, thereby sustaining the flow of electrons through the external circuit [8](#page=8) [9](#page=9). #### 3.1.3 Polarity of electrodes In a voltaic cell, the anode is assigned a negative (-) sign because electrons flow away from it. Conversely, the cathode is assigned a positive (+) sign because electrons flow towards it. This contrasts with electrolytic cells where the anode is positive and the cathode is negative [10](#page=10). > **Tip:** Remember that in voltaic cells, the anode is negative and the cathode is positive, driven by the spontaneous flow of electrons from oxidation to reduction. ### 3.2 Notation for voltaic cells A standardized notation system is used to represent voltaic cells and their components concisely. This notation helps in quickly understanding the cell's setup and reactions without needing a detailed diagram [11](#page=11). #### 3.2.1 Basic cell notation The general format for voltaic cell notation is: `Anode half-cell || Cathode half-cell` [11](#page=11). * The **anode** (oxidation half-cell) is always written on the **left** [11](#page=11). * The **cathode** (reduction half-cell) is always written on the **right** [11](#page=11). * A double vertical bar `||` represents the **salt bridge**, signifying the ionic connection between the two half-cells [11](#page=11). * A single vertical bar `|` represents a **phase boundary** between an electrode and its electrolyte solution [11](#page=11). For the zinc-copper voltaic cell with the cell reaction $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$ the notation is [11](#page=11): $Zn(s) | Zn^{2+}(aq) || Cu^{2+}(aq) | Cu(s)$ [11](#page=11). Here, $Zn(s) | Zn^{2+}(aq)$ represents the anode half-cell, and $Cu^{2+}(aq) | Cu(s)$ represents the cathode half-cell [11](#page=11). #### 3.2.2 Notation for electrodes involving gases When a half-reaction involves a gas, an inert material such as platinum (Pt) is used as an electrode. This platinum serves as the terminal and provides a surface for the half-reaction to occur [12](#page=12). For instance, the hydrogen electrode in an acidic solution, where hydrogen gas bubbles over a platinum plate, can be represented as a cathode: $H^{+}(aq) | H_{2}(g) | Pt$ [12](#page=12). The corresponding cathode half-reaction is: $2H^{+}(aq) + 2e^{-} \leftrightarrow H_{2}(g)$ [12](#page=12). To represent the hydrogen electrode as an anode, the notation is simply reversed: $Pt | H_{2}(g) | H^{+}(aq)$ [12](#page=12). #### 3.2.3 Notation for complex half-cells * A **comma (,)** is used to separate different ions or species present in the **same solution** within a half-cell [13](#page=13). Here are examples of notation for various electrodes: * For the reduction of chlorine gas to chloride ions: $Cl_{2}(g) | Cl^{-}(aq) | Pt$ Corresponding cathode reaction: $Cl_{2}(g) + 2e^{-} \leftrightarrow 2Cl^{-}(aq)$ [13](#page=13). * For the reduction of iron(III) ions to iron(II) ions: $Fe^{3+}(aq), Fe^{2+}(aq) | Pt$ Corresponding cathode reaction: $Fe^{3+}(aq) + e^{-} \leftrightarrow Fe^{2+}(aq)$ [13](#page=13). * For the reduction of cadmium ions to cadmium metal: $Cd^{2+}(aq) | Cd(s)$ Corresponding cathode reaction: $Cd^{2+}(aq) + 2e^{-} \leftrightarrow Cd(s)$ [13](#page=13). #### 3.2.4 Including concentrations and pressures It is often necessary to specify the concentrations of solutions or ions, and the pressures of gases involved in the half-reactions. These values are enclosed in parentheses next to the relevant species in the cell notation [13](#page=13). For example, a cell with a zinc electrode in a 1.0 M zinc ion solution and a hydrogen electrode with hydrogen gas at 1.0 atm pressure in a 1.0 M hydrogen ion solution would be written as: $Zn(s) | Zn^{2+}(1.0 \ M) || H^{+}(1.0 \ M) | H_{2}(1.0 \ atm) | Pt$ [13](#page=13). > **Example:** To represent a voltaic cell where silver metal oxidizes to silver ions and copper(II) ions reduce to copper metal, with concentrations specified, the notation would be: > $Ag(s) | Ag^{+}(0.1 \ M) || Cu^{2+}(0.5 \ M) | Cu(s)$ [13](#page=13). #### 3.2.5 Practice Problems Review Problems 1-4 on page 14 offer practical application of writing cell notation and cell reactions. These problems test the understanding of how to translate electrode reactions into cell notation and vice-versa, and how to determine the overall cell reaction from the given notation. For instance, problem 1 requires writing cell notation from electrode reactions while problem 2 asks for the cell reaction given the notation. Problems 3 and 4 also involve deriving cell reactions from specific cell notations [14](#page=14). --- # Cell potential, electrical work, and standard potentials This section explores the concepts of cell potential, its relationship to electrical work, and the definition and application of standard potentials in electrochemical cells [15](#page=15). ### 4.1 Cell potential and electrical work The maximum potential difference between the electrodes of a voltaic cell is defined as the cell potential, also known as the electromotive force (emf), denoted as $E_{\text{cell}}$. This potential difference represents the electrical pressure between two points and can be measured using an electronic digital voltmeter [15](#page=15). The electrical work done in moving a charge through a conductor is given by the product of the charge and the potential difference [15](#page=15). $$ \text{Electrical work} = \text{charge} \times \text{potential difference} $$ The Faraday constant, $F$, quantifies the magnitude of charge on one mole of electrons, specifically $9.6485 \times 10^4$ Coulombs per mole of electrons ($96,485$ C/mol e⁻). A faraday is a unit of charge equivalent to $9.6485 \times 10^4$ C [15](#page=15). The electrical work ($w$) can be expressed as: $$ w = -F \times \text{potential difference} $$ For molar amounts of reactants as written in a cell equation, the maximum electrical work ($w_{\text{max}}$) that a voltaic cell can perform is: $$ w_{\text{max}} = -nFE_{\text{cell}} $$ Here, $n$ represents the number of moles of electrons transferred, and $F$ is the Faraday constant [15](#page=15). > **Example:** > A voltaic cell with the reaction $Hg_2^{2+}(aq) + H_2(g) \leftrightarrow 2Hg(l) + 2H^+(aq)$ has a cell potential of $0.650$ V. To calculate the maximum electrical work when $0.500$ g of $H_2$ is consumed, we first identify the half-reactions: $Hg_2^{2+}(aq) + 2e^- \leftrightarrow 2Hg(l)$ and $H_2(g) \leftrightarrow 2H^+(aq) + 2e^-$. The number of moles of electrons transferred ($n$) is $2$. The maximum work for the reaction as written is $w_{\text{max}} = -2 \text{ mol e}^- \times 96,485 \text{ C/mol e}^- \times 0.650 \text{ V} = -1.25 \times 10^5$ V.C or $-1.25 \times 10^5$ J. Since $1$ mol of $H_2$ is $2.02$ g, for $0.500$ g of $H_2$, the maximum work is $w_{\text{max}} = \frac{0.500 \text{ g} \times -1.25 \times 10^5 \text{ J}}{2.02 \text{ g}} = -3.09 \times 10^4$ J [16](#page=16). > **Practice Question:** > What is the maximum electrical work that can be obtained from $6.54$ g of zinc metal that reacts in a Daniell cell, whose cell potential is $1.10$ V? The overall cell reaction is $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$ [17](#page=17). ### 4.2 Standard cell potentials and standard electrode potentials The contributions of the anode (oxidation) and cathode (reduction) to the overall cell potential are known as oxidation and reduction potentials, respectively. The cell potential can be expressed as the sum of these potentials [18](#page=18): $$ E_{\text{cell}} = \text{oxidation potential} + \text{reduction potential} $$ A reduction potential measures the tendency of an oxidized species to gain electrons in a reduction half-reaction. The oxidation potential for a half-reaction is equivalent to the negative of the reduction potential for its reverse half-reaction [18](#page=18). $$ \text{Oxidation potential for a half-reaction} = -\text{reduction potential for the reverse half-reaction} $$ By convention, reduction potentials are tabulated and referred to as electrode potentials, symbolized by $E$. Consequently, the cell potential can be calculated as [18](#page=18): $$ E_{\text{cell}} = E_{\text{cathode}} + (-E_{\text{anode}}) $$ or more commonly: $$ E_{\text{cell}} = E_{\text{cathode}} - E_{\text{anode}} $$ where both $E_{\text{cathode}}$ and $E_{\text{anode}}$ are expressed as reduction potentials [18](#page=18). The **standard cell potential** ($E^0_{\text{cell}}$) is the emf of a voltaic cell operating under standard-state conditions. These conditions include solute concentrations of $1$ M, gas pressures of $1$ atm, and a specified temperature, typically $25^\circ C$ [18](#page=18). The **standard electrode potential** ($E^0$) is the electrode potential measured under these same standard-state conditions (1 M solute concentrations, 1 atm gas pressures, and usually $25^\circ C$) [18](#page=18). #### 4.2.1 Calculating cell potentials using standard potentials To determine the standard cell potential ($E^0_{\text{cell}}$) from half-reactions and their standard electrode potentials, one half-reaction must be reversed to represent oxidation. The half-reaction with the more negative standard electrode potential typically corresponds to the species acting as the stronger reducing agent and is thus reversed to occur at the anode. Reversing the half-reaction also reverses the sign of its electrode potential, converting it to an oxidation potential. The half-reactions are then balanced for electron transfer (multiplied by appropriate factors, which do not alter the electrode potentials) and added to yield the overall cell reaction. The standard cell potential is the sum of the standard oxidation potential and the standard reduction potential [20](#page=20). Alternatively, and more directly, the standard cell potential can be calculated using the standard reduction potentials of the cathode and anode: $$ E^0_{\text{cell}} = E^0_{\text{cathode}} - E^0_{\text{anode}} $$ > **Example:** > Consider a voltaic cell with the following reduction half-reactions and standard electrode potentials: $Cd^{2+}(aq) + 2e^- \rightarrow Cd(s); E^0_{Cd} = -0.40$ V and $Ag^+(aq) + e^- \rightarrow Ag(s); E^0_{Ag} = 0.80$ V. To form a spontaneous cell, the half-reaction with the more negative electrode potential ($Cd^{2+}/Cd$) must be reversed to act as the anode. Thus, the oxidation half-reaction is $Cd(s) \rightarrow Cd^{2+}(aq) + 2e^-$, with an oxidation potential of $-E^0_{Cd} = -(-0.40 \text{ V}) = 0.40$ V. The reduction half-reaction remains $Ag^+(aq) + e^- \rightarrow Ag(s)$ with $E^0_{Ag} = 0.80$ V. After balancing electrons and adding the half-reactions, the standard cell potential is $E^0_{\text{cell}} = 0.40 \text{ V} + 0.80 \text{ V} = 1.20$ V. Using the formula $E^0_{\text{cell}} = E^0_{\text{cathode}} - E^0_{\text{anode}}$, we get $E^0_{\text{cell}} = E^0_{Ag} - E^0_{Cd} = 0.80 \text{ V} - (-0.40 \text{ V}) = 1.20$ V. The cell diagram is $Cd(s) | Cd^{2+}(aq) || Ag^+(aq) | Ag(s)$ [20](#page=20) [21](#page=21). > **Example 2:** > For the cell reaction $Zn(s) | Zn^{2+}(aq) || H^+(aq) | H_2(g) | Pt$, the half-reactions and potentials are $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-; -E^0_{Zn} = 0.760$ V and $2H^+(aq) + 2e^- \rightarrow H_2(g); E^0_{H_2} = 0.00$ V. The cell potential is calculated as $E_{\text{cell}} = E^0_{H_2} - E^0_{Zn} = 0.00 \text{ V} - (-0.760 \text{ V}) = 0.760$ V [21](#page=21). > **Practice Question:** > Calculate the standard cell potential of the following voltaic cell at $25^\circ C$ using standard electrode potentials: $Al(s) | Al^{3+}(aq) || Fe^{2+}(aq)| Fe(s)$. What is the cell reaction [22](#page=22)? --- # Electrolysis and its applications This section details the principles and applications of electrolysis, including electrolytic cells, factors affecting ion discharge, molten and aqueous electrolysis, stoichiometric calculations based on Faraday's laws, and methods for corrosion protection. ### 5.1 Electrolytic cells An electrolytic cell is an electrochemical cell where an external electric current drives a non-spontaneous chemical reaction. The process occurring within this cell is called electrolysis, which involves the decomposition of an ionic compound when an electric current is passed through its aqueous solution or molten form via electrodes. This process effectively converts electrical energy into chemical energy [27](#page=27). #### 5.1.1 Components of an electrolytic cell * **Electrodes:** These are conductors that carry current into and out of the cell, typically made of metals like copper (Cu) or zinc (Zn), or graphite [27](#page=27). * **Anode:** This is the positively charged electrode where electrons leave the electrolyte, thus attracting negative ions (anions). Oxidation (loss of electrons) occurs at the anode [27](#page=27). * **Cathode:** This is the negatively charged electrode where electrons enter the electrolyte, attracting positive ions (cations). Reduction (gain of electrons) occurs at the cathode [27](#page=27). * **Electrolyte:** This is the compound that dissociates into ions when dissolved in water or in a molten state. Examples include mineral acids, alkalis, and salts [28](#page=28). * **Power Source:** A direct current (DC) source, such as a battery, is required to drive the electrolysis [28](#page=28). * **Anions:** These are negatively charged ions that migrate towards the anode during electrolysis [28](#page=28). * **Cations:** These are positively charged ions that migrate towards the cathode during electrolysis [28](#page=28). #### 5.1.2 The electrolytic process Electrolysis involves three key steps: 1. **Ionization of electrolyte and water:** The electrolyte and water dissociate to form ions, which then move randomly within the electrolyte. For example, copper(II) sulfate dissociates into $Cu^{2+}$ and $SO_4^{2-}$ ions, while water dissociates into $H^{+}$ and $OH^{-}$ ions [29](#page=29). $$CuSO_4 \rightarrow Cu^{2+} + SO_4^{2-}$$ $$H_2O \rightleftharpoons H^{+} + OH^{-}$$ 2. **Migration of ions to the electrodes:** Once the electrical circuit is completed, cations move towards the cathode, and anions move towards the anode [29](#page=29). 3. **Discharge of ions at the electrodes:** This step involves the gain of electrons (reduction) at the cathode and the loss of electrons (oxidation) at the anode, resulting in redox reactions [29](#page=29). ### 5.2 Molten electrolysis In molten electrolysis, ionic compounds are decomposed in their molten state. For example, molten copper(II) chloride ($CuCl_2$) dissociates into $Cu^{2+}$ and $Cl^{-}$ ions [30](#page=30). At the cathode: $$Cu^{2+} + 2e^{-} \rightarrow Cu(s)$$ At the anode: $$2Cl^{-} \rightarrow Cl_2(g) + 2e^{-}$$ Overall reaction: $$Cu^{2+} + 2Cl^{-} \rightarrow Cu(s) + Cl_2(g)$$ Similarly, molten sodium chloride (NaCl) can be electrolyzed: $$Na^{+} + e^{-} \rightarrow Na(s)$$ $$Cl^{-} \rightarrow \frac{1}{2}Cl_2(g) + e^{-}$$ Overall reaction: $$Na^{+} + Cl^{-} \rightarrow Na(s) + \frac{1}{2}Cl_2(g)$$ The electrolysis of molten NaCl is commercially significant for producing sodium metal in a Downs cell. This cell is designed to keep the reactive products separate. Adding calcium chloride to sodium chloride lowers the melting point from 801°C to approximately 580°C, reducing energy costs [30](#page=30). ### 5.3 Aqueous electrolysis When an aqueous solution of an ionic compound is electrolyzed, water can participate in the reactions at either or both electrodes. Water can be reduced to hydrogen gas or oxidized to oxygen gas [31](#page=31). Reduction half-reaction of water: $$2H_2O(l) + 2e^{-} \rightarrow H_2(g) + 2OH^{-}(aq)$$ Oxidation half-reaction of water: $$2H_2O(l) \rightarrow O_2(g) + 4H^{+}(aq) + 4e^{-}$$ For an aqueous solution of copper(II) sulfate ($CuSO_4$), both $Cu^{2+}$, $SO_4^{2-}$, and ions from water ($H^{+}$, $OH^{-}$) are present, and preferential discharge of ions occurs [31](#page=31). #### 5.3.1 Factors affecting preferential discharge of ions The preferential discharge of ions at the electrodes is influenced by several factors: * **Position in the activity series:** Ions lower in the electrochemical activity series are preferentially discharged over those above them in dilute solutions [32](#page=32). * **Cations:** The order of discharge (from easiest to hardest reduction) is approximately: $K^{+} > Na^{+} > Ca^{2+} > Mg^{2+} > Al^{3+} > Zn^{2+} > Fe^{2+} > Pb^{2+} > H^{+} > Cu^{2+} > Hg^{2+} > Ag^{+} > Au^{+} > Pt^{+}$ [32](#page=32). * **Anions:** The order of discharge (from easiest to hardest oxidation) is approximately: $SO_4^{2-} > NO_3^{-} > Cl^{-} > Br^{-} > OH^{-}$ [32](#page=32). * **Example:** If both $Cu^{2+}$ and $H^{+}$ ions are present at the cathode, $Cu^{2+}$ will be preferentially discharged in dilute solutions because it is lower in the activity series. The cathodic reaction would be $Cu^{2+} + 2e^{-} \rightarrow Cu(s)$ [32](#page=32). * **Concentration of electrolyte:** A high concentration of an ion can sometimes reverse the order of discharge, particularly if the ions are close in the activity series. For instance, in a concentrated NaCl or HCl solution, $Cl^{-}$ is discharged preferentially over $OH^{-}$ due to its much higher concentration. However, in dilute solutions, $OH^{-}$ might be discharged [33](#page=33). * **Nature of the electrode:** The material of the electrodes can influence which ion is discharged, as reactions tend to occur in a way that requires less energy. For example, using a mercury (Hg) electrode for the electrolysis of NaCl can lead to the formation of sodium amalgam ($NaHg$) instead of hydrogen gas, as it requires less energy than the discharge of $H^{+}$ [33](#page=33). #### 5.3.2 Examples of electrolysis of common electrolytes * Electrolysis of dilute $H_2SO_4$ (using inert electrodes) produces $H_2$ gas at the cathode and $O_2$ gas at the anode, similar to the electrolysis of water [34](#page=34). * Electrolysis of $CuSO_4(aq)$ using platinum electrodes results in the deposition of copper at the cathode and oxygen evolution at the anode (if the concentration of $Cu^{2+}$ is higher than that of $H^{+}$). If copper electrodes are used, the anode dissolves, replenishing $Cu^{2+}$ ions in the solution [34](#page=34). * Electrolysis of concentrated NaCl solution with inert electrodes produces hydrogen gas at the cathode and chlorine gas at the anode [34](#page=34). * Electrolysis of aqueous KI with inert electrodes produces hydrogen gas at the cathode and iodine at the anode [34](#page=34). * Electrolysis of aqueous $Na_2SO_4$ with platinum electrodes produces hydrogen gas at the cathode and oxygen gas at the anode [34](#page=34). ### 5.4 Applications of electrolysis Electrolysis has numerous industrial and practical applications: * **Purification of metals:** Impure metals like copper, silver, and gold can be purified by making them the anode in an electrolytic cell, leading to a purer metal deposited at the cathode [35](#page=35). * **Extraction of reactive metals:** Electrolysis is used to extract highly reactive metals such as sodium, potassium, calcium, magnesium, and aluminum from their ores [35](#page=35). * **Industrial preparation of gases:** Essential gases like hydrogen ($H_2$), oxygen ($O_2$), and chlorine ($Cl_2$) are produced industrially through electrolysis [35](#page=35). * **Industrial manufacture of compounds:** Important compounds such as sodium hydroxide (NaOH) are manufactured using electrolysis [35](#page=35). * **Electroplating:** This process involves coating one metal with a thin layer of another metal, primarily to protect against corrosion (e.g., plating iron with a less reactive metal) and to enhance aesthetic appeal (e.g., on cars, cutlery). The object to be plated is made the cathode, and the coating material is typically made the anode [35](#page=35). ### 5.5 Stoichiometry of electrolysis The quantitative aspects of electrolysis are governed by Faraday's Laws. #### 5.5.1 Faraday's laws of electrolysis * **First Law:** The mass of a substance liberated or deposited at an electrode during electrolysis is directly proportional to the quantity of electric charge passed through the electrolyte [36](#page=36). Mathematically: $m \propto Q$, or $m = ZQ$, where $m$ is the mass, $Q$ is the quantity of charge, and $Z$ is the electrochemical equivalent [36](#page=36). The quantity of electric charge ($Q$) can be calculated as: $Q = \text{electric current (A)} \times \text{time lapse (s)}$ [36](#page=36). One Faraday ($1F$) is equivalent to the charge of one mole of electrons and is equal to 96,485 Coulombs (C). The coulomb (C) is the SI unit of electric charge, equivalent to an ampere-second (A.s) [36](#page=36). * **Second Law:** For a given quantity of electric charge, the mass of any metal deposited is proportional to its equivalent mass (atomic mass divided by the charge on the metal ion) [37](#page=37). Alternatively, when the same quantity of electricity is passed through different electrolytes connected in series, the relative number of moles of elements deposited or liberated is inversely proportional to the charges on their respective ions [37](#page=37). Mathematically: $n \propto \frac{1}{e^{-}}$, where $n$ is the number of moles of the element deposited and $e^{-}$ is the charge on the ion of the element discharged [37](#page=37). For a general cathode reaction $M^{x+} + Xe^{-} \rightarrow M(s)$, if $1F$ of electricity is passed, it deposits $\frac{1}{x}$ moles of $M$. Therefore, passing $1F$ of electricity through solutions containing $Na^{+}$, $Cu^{2+}$, and $Al^{3+}$ would deposit 1 mole of Na, $\frac{1}{2}$ mole of Cu, and $\frac{1}{3}$ mole of Al, respectively [37](#page=37). #### 5.5.2 Examples of stoichiometric calculations > **Example 1:** Calculate the mass of copper metal deposited during the passage of 2.50 A of current through a solution of $CuSO_4$ for 50 minutes. (Atomic mass of Cu = 63.5) > **Calculation:** > Time in seconds, $t = 50 \text{ mins} \times 60 \text{ s/min} = 3000 \text{ s}$. > Charge, $Q = I \times t = 2.50 \text{ A} \times 3000 \text{ s} = 7500 \text{ C}$. > The reaction at the cathode is $Cu^{2+} + 2e^{-} \rightarrow Cu(s)$. This means 2 moles of electrons (2F) are required to deposit 1 mole of Cu (63.5 g). > Mass of Cu deposited = $\frac{\text{Charge}}{\text{Faraday's constant}} \times \frac{\text{Molar mass of Cu}}{\text{Number of electrons involved}}$ > Mass of Cu = $\frac{7500 \text{ C}}{96485 \text{ C/F}} \times \frac{63.5 \text{ g/mol}}{2 \text{ mol e}^{-}/\text{mol Cu}}$ $\approx 2.47 \text{ g}$ [38](#page=38). > **Example 2:** What volume of $O_2$ gas at STP is produced by the oxidation of water in the electrolysis of $CuSO_4$, 50 minutes after the passage of 2.50 A of current? > **Calculation:** > From Example 1, $Q = 7500 \text{ C}$. > The oxidation half-reaction for water is $2H_2O(l) \rightarrow O_2(g) + 4H^{+}(aq) + 4e^{-}$. This means 4 moles of electrons (4F) produce 1 mole of $O_2$. > Moles of $O_2$ produced = $\frac{\text{Charge}}{\text{Faraday's constant}} \times \frac{1 \text{ mol } O_2}{4 \text{ mol e}^{-}}$ > Moles of $O_2$ = $\frac{7500 \text{ C}}{96485 \text{ C/F}} \times \frac{1}{4} \approx 0.0194 \text{ mol}$. > Volume of $O_2$ at STP ($22.4 \text{ dm}^3/\text{mol}$) = $0.0194 \text{ mol} \times 22.4 \text{ dm}^3/\text{mol} \approx 0.435 \text{ dm}^3$ or $435 \text{ cm}^3$ [38](#page=38). ### 5.6 Corrosion Corrosion is a redox process where metals are oxidized by oxygen in the presence of moisture. This process occurs more readily in metals that are more active, especially at points of strain or imperfections [40](#page=40). #### 5.6.1 Mechanism of iron corrosion (rusting) Rusting is a common example of metal corrosion, typically occurring with iron and steel [40](#page=40). 1. **Anodic region:** At points of strain, iron acts as an anode and is oxidized to iron(II) ions ($Fe^{2+}$), forming pits [40](#page=40). $$Fe(s) \rightarrow Fe^{2+}(aq) + 2e^{-} \quad \text{(Oxidation, anode)}$$ 2. **Cathodic region:** Electrons released flow through the metal to areas exposed to oxygen. These areas act as cathodes where oxygen is reduced to hydroxide ions ($OH^{-}$) [40](#page=40). $$O_2(g) + 2H_2O(l) + 4e^{-} \rightarrow 4OH^{-}(aq) \quad \text{(Reduction, cathode)}$$ 3. **Formation of iron(II) hydroxide:** The $Fe^{2+}$ ions migrate and combine with $OH^{-}$ ions to form iron(II) hydroxide ($Fe(OH)_2$) [41](#page=41). 4. **Further oxidation:** Iron(II) ions are further oxidized by oxygen to iron(III) ions, which then form hydrated iron(III) oxides and hydroxides, commonly known as rust [41](#page=41). $$2Fe(s) + \frac{3}{2}O_2(g) + xH_2O(l) \rightarrow Fe_2O_3 \cdot xH_2O(s) \quad \text{(Rust)}$$ At the anode, pitting appears, while rust forms at the cathodic region [42](#page=42). #### 5.6.2 Corrosion protection Several methods can be employed to protect metals from corrosion: 1. **Plating:** Coating the metal with a thin layer of a less easily oxidized metal [42](#page=42). 2. **Galvanizing:** Coating steel with zinc, which is a more active metal and preferentially corrodes, protecting the steel [42](#page=42). 3. **Protective coatings:** Applying barriers like paint to isolate the metal from the environment [42](#page=42). 4. **Formation of protective film:** Allowing a natural oxide layer (e.g., on aluminum) to form and act as a protective barrier [42](#page=42). 5. **Sacrificial anode:** Connecting the metal to be protected directly to a more active metal (sacrificial anode) that will be preferentially oxidized [42](#page=42). --- ## Common mistakes to avoid - Review all topics thoroughly before exams - Pay attention to formulas and key definitions - Practice with examples provided in each section - Don't memorize without understanding the underlying concepts

| Term | Definition | |------|------------| | Electrochemistry | The branch of chemistry concerned with the relationship between electrical energy and chemical transformations. | | Electrolysis | A process where electrical energy is used to drive a nonspontaneous chemical reaction, typically involving the decomposition of a substance. | | Electrolytic cell | An electrochemical cell in which electrical energy from an external source causes nonspontaneous chemical reactions to occur. | | Voltaic cell (Galvanic cell) | An electrochemical cell in which spontaneous chemical reactions produce electricity and supply it to an external circuit. | | Electrode | A conductor on which oxidation or reduction half-reactions occur; it may or may not participate in the reaction. | | Cathode | The electrode at which reduction occurs, meaning species gain electrons. In voltaic cells, it is typically the positive electrode, and in electrolytic cells, it is typically the negative electrode. | | Anode | The electrode at which oxidation occurs, meaning species lose electrons. In voltaic cells, it is typically the negative electrode, and in electrolytic cells, it is typically the positive electrode. | | Metallic conduction | The conduction of electric current through metals via the flow of free electrons without movement of atoms. | | Ionic (electrolytic) conduction | The conduction of electric current by the motion of ions through a solution or a pure liquid. | | Salt bridge | A tube containing an electrolyte in a gel that connects the two half-cells of a voltaic cell, allowing ion flow while preventing the mixing of solutions. | | Cell potential (Ecell) | The maximum potential difference between the electrodes of a voltaic cell, also known as electromotive force (emf). | | Faraday constant (F) | The magnitude of electric charge on one mole of electrons, approximately $9.6485 \times 10^4$ C/mol e-. | | Standard electrode potential (E0) | The electrode potential of an electrode measured under standard-state conditions (1 M concentration for solutes, 1 atm pressure for gases, and a specified temperature, usually 25°C). | | Standard cell potential (E0cell) | The emf of a voltaic cell operating under standard-state conditions. | | Oxidation potential | The potential associated with an oxidation half-reaction. | | Reduction potential | The potential associated with a reduction half-reaction. | | Faraday's Laws of Electrolysis | Laws that describe the quantitative relationships between the amount of electricity passed through an electrolyte and the mass of substance liberated or deposited at the electrodes. | | Electrochemical equivalent (Z) | The mass of a substance liberated or discharged at an electrode by the passage of one coulomb of electric charge. | | Corrosion | A redox process where metals are oxidized by oxygen in the presence of moisture, leading to degradation of the metal. |