Chapter 3 Electrochemistry

1. What is electrochemistry? Discuss the types of electrochemical cells.

Answer: Electrochemistry is the branch of chemistry that deals with the study of the relationship between electrical energy and chemical changes. It involves reactions in which either electricity is produced from chemical reactions (in galvanic cells) or chemical reactions occur as a result of the passage of an electric current (in electrolytic cells).

Types of Electrochemical Cells:

1. Galvanic Cells (Voltaic Cells): These cells convert chemical energy into electrical energy. A common example is the Daniell cell, where zinc undergoes oxidation and copper undergoes reduction. In a galvanic cell, spontaneous chemical reactions produce electricity. The two half-reactions occur in separate compartments connected by a salt bridge to maintain charge balance.

   • Example: In the Daniell cell, the anode is zinc (Zn) and the cathode is copper (Cu).
   • Reaction: $$\text{Zn (s)} \rightarrow \text{Zn}^{2+} (aq) + 2e^-$$ $$\text{Cu}^{2+} (aq) + 2e^- \rightarrow \text{Cu (s)}$$

2. Electrolytic Cells: In electrolytic cells, electrical energy is used to drive non-spontaneous chemical reactions. The most common application is the electrolysis of water, where electric current is used to break water molecules into hydrogen and oxygen gas. Electrolytic cells consist of two electrodes placed in an electrolyte solution.

   • Example: Electrolysis of water produces hydrogen and oxygen gases.

The major difference between galvanic and electrolytic cells is that in galvanic cells, the cell potential is positive, while in electrolytic cells, the cell potential is negative and requires an external power source.

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2. Derive the Nernst equation and explain its significance.

Answer: The Nernst equation provides the relationship between the cell potential and the concentrations of the reactants and products in an electrochemical reaction. It allows us to calculate the potential of a half-cell or a full electrochemical cell under non-standard conditions.

Derivation of the Nernst Equation:

The Nernst equation can be derived from the Gibbs free energy equation. For a reaction:

$$\text{Red} + \text{Ox} \rightarrow \text{Red} + \text{Ox}$$

The change in Gibbs free energy $\Delta G$ is related to the electrical work $W$ done in the electrochemical reaction:

$$\Delta G = -nFE$$

where:

• $n$ is the number of moles of electrons involved,
• $F$ is Faraday's constant,
• $E$ is the cell potential.

For a reaction at non-standard conditions, the change in Gibbs free energy is also related to the concentration of the reactants and products:

$$\Delta G = \Delta G^\circ + RT \ln Q$$

where:

• $\Delta G^\circ$ is the standard Gibbs free energy,
• $R$ is the universal gas constant,
• $T$ is the temperature in Kelvin,
• $Q$ is the reaction quotient.

Equating these two expressions for $\Delta G$ :

$$-nFE = \Delta G^\circ + RT \ln Q$$

At equilibrium, $\Delta G^\circ = -nFE^\circ$ , so the equation becomes:

$$-nFE = -nFE^\circ + RT \ln Q$$

Dividing by $-nF$ :

$$E = E^\circ - \frac{RT}{nF} \ln Q$$

This is the Nernst Equation:

$$E = E^\circ - \frac{0.0591}{n} \log Q \quad \text{(at 25°C)}$$

Significance of the Nernst Equation:

• It allows the calculation of the cell potential under non-standard conditions.
• It shows how the potential of an electrochemical reaction is affected by the concentrations of the reactants and products.
• It helps predict the direction of a reaction and calculate equilibrium concentrations.

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3. How does the concentration of reactants and products affect the electrode potential?

Answer: The concentration of reactants and products significantly influences the electrode potential. According to the Nernst equation, the electrode potential varies with the concentrations of the ions involved in the half-reaction.

1. Effect of Reactant Concentration:

   • If the concentration of the reactant is increased, the electrode potential becomes more positive (favorable for reduction).
   • If the concentration of the reactant decreases, the electrode potential becomes more negative.

2. Effect of Product Concentration:

   • If the concentration of the product is increased, the reaction shifts towards the reactant side, leading to a decrease in electrode potential (more favorable for oxidation).
   • If the concentration of the product decreases, the electrode potential increases (more favorable for reduction).

For example, in a concentration cell, when two solutions with different concentrations of ions are connected, the side with a higher ion concentration will have a higher potential, and the current will flow from the higher concentration to the lower concentration.

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4. What are primary and secondary batteries? Discuss their characteristics with examples.

Answer: Primary Batteries: Primary batteries are non-rechargeable batteries. Once their chemical reactions are exhausted, they cannot be reversed, and the battery cannot be reused. These are typically used in devices with low energy consumption, such as remote controls or flashlights.

• Example: The Leclanché cell and Zinc-carbon battery are common examples of primary batteries.
• Characteristics:
  • Non-rechargeable.
  • Shorter lifespan compared to secondary batteries.
  • Cost-effective for low-energy applications.
  • Commonly used for household items.

Secondary Batteries: Secondary batteries are rechargeable batteries. They can be discharged and then recharged multiple times, making them more cost-effective in the long term.

• Example: The Lead-acid battery and Lithium-ion battery are examples of secondary batteries.
• Characteristics:
  • Rechargeable.
  • Longer lifespan than primary batteries.
  • Higher initial cost but more cost-effective over time.
  • Used in applications like automobiles, mobile phones, and laptops.

The main difference between primary and secondary batteries lies in their ability to be recharged and reused.

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5. What is the significance of the conductivity in electrolytic solutions? Derive the relation between the molar conductivity and conductivity.

Answer: Conductivity is a measure of the ability of a solution to conduct electricity. It depends on the concentration of ions in the solution and their mobility. In an electrolytic solution, conductivity increases with the increase in ion concentration.

Molar Conductivity ($\Lambda_m$ ) is defined as the conductivity of a solution divided by its molar concentration:

$$\Lambda_m = \frac{\kappa}{c}$$

where:

• $\kappa$ is the conductivity of the solution (S/m),
• $c$ is the molar concentration of the electrolyte (mol/L).

Relation between Conductivity and Molar Conductivity: As the concentration of an electrolyte decreases, the molar conductivity increases because the ions have more freedom to move and contribute to the current. The molar conductivity at infinite dilution ($\Lambda_m^\circ$ ) is a constant and is used to predict the behavior of electrolytes in dilute solutions.

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6. What is the function of the salt bridge in a galvanic cell?

Answer: The salt bridge is a crucial component in a galvanic cell that maintains electrical neutrality by allowing the flow of ions between the two half-cells. Without it, the build-up of charge in the cells would prevent the redox reactions from occurring. The salt bridge is typically made of a gel containing a salt like KCl or NaNO₃, which dissociates into ions.

The salt bridge allows the anions to move toward the anode and the cations toward the cathode, thereby balancing the charge in each half-cell. This helps maintain a continuous flow of current through the external circuit.

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7. Explain the concept of standard electrode potential and how it is measured.

Answer: The standard electrode potential (E°) is the potential of a half-cell relative to the standard hydrogen electrode (SHE), which is defined as 0 volts. It measures the tendency of a substance to gain or lose electrons when it is in contact with its ions in a solution.

Standard Conditions for measuring electrode potential:

• The solution concentration is 1 mol/L.
• The temperature is 298 K (25°C).
• The pressure is 1 atm.

The standard electrode potential can be measured by connecting the half-cell to the SHE through a voltmeter. The higher the standard electrode potential, the greater the tendency of the species to gain electrons and undergo reduction.

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8. Explain the working of a fuel cell with an example.

Answer: A fuel cell is an electrochemical cell that converts the chemical energy of a fuel (such as hydrogen or methanol) directly into electrical energy through a redox reaction. It works by combining the fuel with oxygen, producing electricity, water, and heat.

Example: Hydrogen Fuel Cell In a hydrogen fuel cell:

• Anode: Hydrogen gas ($H_2$ ) is oxidized to produce electrons and protons. $$H_2 \rightarrow 2H^+ + 2e^-$$
• Cathode: Oxygen gas ($O_2$ ) is reduced by accepting the electrons and combining with protons to form water. $$O_2 + 4H^+ + 4e^- \rightarrow 2H_2O$$
• The electrons flow through an external circuit from the anode to the cathode, generating electricity.

Fuel cells are efficient and environmentally friendly, emitting only water as a by-product.

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9. Discuss the relation between the cell potential and the spontaneity of a reaction.

Answer: The spontaneity of a reaction can be determined by the cell potential of an electrochemical cell. If the cell potential ($E$ ) is positive, the reaction is spontaneous; if $E$ is negative, the reaction is non-spontaneous.

The relationship is given by the Gibbs free energy equation:

$$\Delta G = -nFE$$

• If $E > 0$ , then $\Delta G < 0$ , meaning the reaction is spontaneous.
• If $E < 0$ , then $\Delta G > 0$ , meaning the reaction is non-spontaneous.

Thus, the cell potential provides direct insight into whether a given electrochemical reaction will occur spontaneously.

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10. What are the different types of conductors used in electrochemical cells? Discuss their significance.

Answer: In electrochemical cells, conductors are used to facilitate the flow of electrons. The types of conductors include:

1. Electrodes: These are the conductors that participate in the electrochemical reactions at the anode and cathode.

   • Anode: The electrode where oxidation occurs (loss of electrons).
   • Cathode: The electrode where reduction occurs (gain of electrons).

2. Salt Bridge: Contains ionic conductors (like KCl) that allow ions to flow between the two half-cells, completing the circuit without mixing the solutions.

The efficiency of electrochemical reactions relies on the conductivity of the conductors used, which ensures that electrons and ions move freely, enabling the redox reactions to proceed effectively.

11. Explain the concept of electrode potential. How is it related to the tendency of a species to gain electrons?

Answer: Electrode potential is the potential difference between an electrode and its surrounding solution, which is a measure of the tendency of a chemical species to either gain or lose electrons when it is in contact with its ions in a solution. Electrode potential is usually measured in reference to the Standard Hydrogen Electrode (SHE), which is assigned a potential of 0 volts.

• A positive electrode potential indicates that the species is more likely to gain electrons (reduction), and thus, the substance acts as a good oxidizing agent.
• A negative electrode potential suggests that the species has a lower tendency to gain electrons and is more likely to lose electrons (oxidation), acting as a good reducing agent.

Electrode potential reflects the reduction potential of a species. The more positive the electrode potential, the greater its ability to gain electrons and undergo reduction. Conversely, a more negative electrode potential indicates a greater tendency for the substance to lose electrons and be oxidized.

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12. Discuss the concept of the electrochemical series and its application in predicting redox reactions.

Answer: The electrochemical series is a list of elements or ions arranged according to their standard electrode potentials (E°). These potentials are measured under standard conditions: 298 K, 1 atm pressure, and 1 mol/L concentration of ions. The electrochemical series provides crucial information about the relative tendencies of different substances to gain or lose electrons.

• Application in Redox Reactions: The electrochemical series can predict the direction of redox reactions. The substance with a more positive electrode potential will act as an oxidizing agent, and the substance with a more negative electrode potential will act as a reducing agent. The species at the top of the series will be reduced (gain electrons), and the species at the bottom will be oxidized (lose electrons).

For example, if we have a reaction between zinc and copper ions:

• Zinc has a more negative potential than copper, so zinc will lose electrons (oxidize), and copper ions will gain electrons (reduce).
• The overall reaction is: $$\text{Zn (s)} + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu (s)}$$

This indicates that the reaction is spontaneous because the reduction potential for copper is more positive than that for zinc.

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13. Explain the process of electrolysis. What factors affect the efficiency of electrolysis?

Answer: Electrolysis is a process in which an electric current is used to drive a non-spontaneous chemical reaction. During electrolysis, the electrical energy is converted into chemical energy, resulting in the decomposition of ionic compounds into their constituent elements.

Example: Electrolysis of water leads to the production of hydrogen and oxygen gases:

$$\text{2H}_2\text{O} \rightarrow \text{2H}_2 (g) + \text{O}_2 (g)$$

Factors Affecting the Efficiency of Electrolysis:

1. Concentration of the Electrolyte: A higher concentration of ions increases the number of charge carriers, which enhances conductivity and improves efficiency.
2. Voltage Applied: The minimum voltage required for electrolysis to occur is known as the decomposition voltage. If the applied voltage is greater than the decomposition voltage, the reaction occurs more rapidly.
3. Nature of Electrodes: The type of material used for electrodes affects the ease with which they can accept or donate electrons. Inert electrodes like platinum are commonly used in electrolysis.
4. Temperature: Higher temperatures generally increase the rate of the reaction, but they may also increase the loss of energy due to heating.

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14. What is the difference between a galvanic cell and an electrolytic cell?

Answer: The main difference between galvanic cells and electrolytic cells lies in the nature of the reactions and the purpose of the cell.

1. Galvanic Cell (Voltaic Cell):

   • Spontaneous Reactions: Galvanic cells produce electricity from spontaneous chemical reactions. The chemical energy is converted into electrical energy.
   • Example: Daniell cell.
   • Direction of Current: The current flows from the anode (oxidation) to the cathode (reduction) externally.
   • Energy Source: It provides energy without the need for an external power source.

2. Electrolytic Cell:

   • Non-Spontaneous Reactions: Electrolytic cells require external electrical energy to drive non-spontaneous reactions.
   • Example: Electrolysis of water or sodium chloride.
   • Direction of Current: The current is applied externally to force electrons to flow from the cathode (reduction) to the anode (oxidation).
   • Energy Source: An external power source is necessary to initiate the process.

In summary, galvanic cells generate electricity from spontaneous reactions, while electrolytic cells consume electricity to drive non-spontaneous reactions.

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15. How is the efficiency of an electrochemical cell related to its cell potential?

Answer: The efficiency of an electrochemical cell is directly related to its cell potential. The higher the cell potential, the more electrical energy the cell can produce. The relationship between the cell potential and the Gibbs free energy is given by the equation:

$$\Delta G = -nFE$$

where:

• $\Delta G$ is the Gibbs free energy,

• $n$ is the number of moles of electrons,

• $F$ is Faraday’s constant,

• $E$ is the cell potential.

• A higher cell potential indicates a more spontaneous reaction, which results in a larger release of energy, making the cell more efficient in converting chemical energy into electrical energy.

• Conversely, a lower cell potential indicates less energy production, and the cell is less efficient.

The efficiency of an electrochemical cell is also influenced by factors like the concentration of ions, the nature of the electrodes, and temperature.

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16. What is the role of Faraday’s laws in electrolysis?

Answer: Faraday’s laws of electrolysis provide quantitative relationships that describe how much substance is deposited or dissolved at the electrodes during electrolysis.

1. First Law of Electrolysis: The mass of a substance deposited or dissolved at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

   $$m = \frac{M \cdot Q}{F \cdot z}$$

   where:

   • $m$ is the mass of the substance,
   • $M$ is the molar mass of the substance,
   • $Q$ is the charge passed (in coulombs),
   • $F$ is Faraday’s constant,
   • $z$ is the number of electrons involved in the reaction.

2. Second Law of Electrolysis: When the same amount of electricity is passed through different electrolytes, the amount of different substances deposited is directly proportional to their equivalent masses.

Significance: Faraday’s laws enable us to calculate the amount of a substance that will be deposited or dissolved during electrolysis, making them crucial for industrial applications like electroplating and electrorefining.

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17. What is the difference between oxidation and reduction in electrochemical reactions?

Answer: In electrochemical reactions, oxidation and reduction occur simultaneously and are opposite processes:

• Oxidation: This is the loss of electrons. In an electrochemical cell, oxidation occurs at the anode. For example, in the Daniell cell, zinc (Zn) is oxidized to form zinc ions (Zn²⁺) and releases electrons:

  $$\text{Zn} (s) \rightarrow \text{Zn}^{2+} (aq) + 2e^-$$

• Reduction: This is the gain of electrons. Reduction occurs at the cathode. In the Daniell cell, copper ions (Cu²⁺) gain electrons and are reduced to form copper metal:

  $$\text{Cu}^{2+} (aq) + 2e^- \rightarrow \text{Cu} (s)$$

Oxidation always involves the loss of electrons, and reduction always involves the gain of electrons. Both processes are essential for the flow of current in electrochemical cells.

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18. Discuss the concept of electrochemical equilibrium and the Nernst equation in this context.

Answer: Electrochemical equilibrium occurs when the forward and reverse reactions in an electrochemical cell reach the same rate, resulting in no net change in the concentrations of reactants and products. At this point, the cell potential becomes zero, and the cell is no longer capable of doing electrical work.

At equilibrium, the Nernst equation helps determine the equilibrium cell potential based on the concentrations of the reactants and products:

$$E = E^\circ - \frac{0.0591}{n} \log Q$$

where $Q$ is the reaction quotient. At equilibrium, the reaction quotient $Q$ becomes equal to the equilibrium constant $K$ , and the cell potential $E$ becomes zero:

$$E = 0 \quad \text{when} \quad Q = K$$

The Nernst equation allows the calculation of the cell potential at any concentration, not just standard conditions, and is essential in understanding how concentration affects the spontaneity and efficiency of electrochemical reactions.

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19. Explain the process of rusting of iron in terms of electrochemistry.

Answer: The rusting of iron is an electrochemical process that involves both oxidation and reduction reactions:

• Oxidation (at the anode): Iron (Fe) is oxidized in the presence of oxygen and water to form iron(II) ions ($\text{Fe}^{2+}$ ) and electrons: $$\text{Fe} (s) \rightarrow \text{Fe}^{2+} (aq) + 2e^-$$
• Reduction (at the cathode): The electrons released from the oxidation reaction flow to another part of the metal, where they reduce oxygen from the air in the presence of water to form hydroxide ions ($\text{OH}^-$ ): $$\text{O}_2 (g) + 2H_2O (l) + 4e^- \rightarrow 4\text{OH}^- (aq)$$
• The iron(II) ions ($\text{Fe}^{2+}$ ) then react with oxygen to form iron(III) oxide ($\text{Fe}_2\text{O}_3$ ), commonly known as rust.

The process of rusting can be mitigated by coating iron with paint, galvanizing it (coating with zinc), or using sacrificial anodes.

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20. What are concentration cells and how do they work?

Answer: A concentration cell is a type of electrochemical cell in which the electrodes are identical, but the electrolyte concentrations differ. The cell generates electrical energy due to the difference in the concentration of ions at the two electrodes.

Example: In a concentration cell involving copper:

• One half-cell contains a higher concentration of copper ions ($\text{Cu}^{2+}$ ) while the other contains a lower concentration.
• The copper ions in the more concentrated solution will be reduced at the cathode, while copper from the solid electrode in the less concentrated solution will be oxidized to increase the concentration of copper ions.

The cell potential for a concentration cell can be calculated using the Nernst equation:

$$E = \frac{0.0591}{n} \log \frac{[C_1]}{[C_2]}$$

where $C_1$ and $C_2$ are the concentrations of the electrolytes in the two half-cells.

21. Explain the relationship between cell potential and Gibbs free energy in electrochemical reactions.

Answer: The relationship between cell potential ($E$ ) and Gibbs free energy ($\Delta G$ ) is fundamental in understanding whether an electrochemical reaction is spontaneous or not. The equation that links them is:

$$\Delta G = -nFE$$

Where:

• $\Delta G$ is the Gibbs free energy (measured in joules),

• $n$ is the number of moles of electrons transferred in the reaction,

• $F$ is Faraday's constant (approximately 96485 C/mol),

• $E$ is the cell potential (measured in volts).

• If $\Delta G$ is negative (i.e., $\Delta G < 0$ ), the reaction is spontaneous, and the cell will produce electrical energy (positive cell potential, $E > 0$ ).

• If $\Delta G$ is positive (i.e., $\Delta G > 0$ ), the reaction is non-spontaneous, and the cell will not produce electrical energy (negative cell potential, $E < 0$ ).

This equation essentially tells us that a positive cell potential corresponds to a spontaneous reaction that can generate electrical work, while a negative cell potential indicates a non-spontaneous process.

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22. What is the significance of the Nernst equation in determining the potential of an electrochemical cell under non-standard conditions?

Answer: The Nernst equation is crucial in electrochemistry because it allows the calculation of the cell potential under any conditions, not just the standard conditions. The standard cell potential is measured when all species are in their standard states (1 M concentration, 1 atm pressure, 298 K temperature), but real-world conditions often differ. The Nernst equation accounts for these deviations and is given by:

$$E = E^\circ - \frac{0.0591}{n} \log Q$$

Where:

• $E^\circ$ is the standard electrode potential of the cell,
• $n$ is the number of electrons involved in the reaction,
• $Q$ is the reaction quotient, which reflects the ratio of concentrations of products to reactants.

The equation shows how the cell potential changes with concentration and temperature. It highlights that the cell potential decreases as the concentration of the products increases and the concentration of the reactants decreases, making the reaction less spontaneous. Therefore, the Nernst equation is invaluable in predicting the behavior of electrochemical cells under practical, non-ideal conditions.

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23. Discuss the applications of electrochemical cells in everyday life and industrial processes.

Answer: Electrochemical cells have numerous practical applications in daily life and industry, from powering devices to producing essential chemicals. Some of the key applications are:

1. Batteries: Electrochemical cells are the foundation of most types of batteries, such as:

   • Primary Batteries (non-rechargeable), like alkaline batteries, which power household devices.
   • Secondary Batteries (rechargeable), such as lithium-ion batteries used in mobile phones, laptops, and electric vehicles. These batteries use reversible reactions, allowing them to be recharged multiple times.

2. Fuel Cells: These are electrochemical cells that generate electricity by converting the chemical energy of a fuel (usually hydrogen) into electrical energy. Fuel cells are used in hydrogen-powered vehicles and for backup power systems, offering an alternative to traditional combustion engines.

3. Electroplating: Electrochemical cells are used in electroplating processes, where a thin layer of metal (such as gold, silver, or chromium) is deposited onto a surface to enhance its appearance or prevent corrosion. This process involves the reduction of metal cations from a solution onto the object being plated.

4. Electrorefining: In industries like copper refining, electrochemical cells are used to purify metals. In this process, the metal to be purified is used as the anode, and pure metal is deposited at the cathode, leaving impurities behind.

5. Electrolysis of Water: Electrochemical cells are used to split water into hydrogen and oxygen gases through electrolysis, providing a method for producing hydrogen fuel.

These applications showcase the importance of electrochemical principles in modern technology and industry.

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24. What is the function of salt bridges in galvanic cells?

Answer: A salt bridge in a galvanic cell is an essential component that maintains electrical neutrality by completing the circuit and allowing ion flow between the two half-cells. Its primary function is to:

1. Prevent the accumulation of charge: As oxidation occurs at the anode (releasing electrons) and reduction occurs at the cathode (accepting electrons), an imbalance of charge would develop in each half-cell. The salt bridge allows ions to flow and neutralize this imbalance.
2. Maintain electrical continuity: The salt bridge typically contains an inert electrolyte like potassium chloride (KCl) or potassium nitrate (KNO₃), which dissociates into ions that move between the two half-cells. These ions help balance the charges by migrating into the respective half-cell.
3. Allow ion movement: As the reaction progresses, positive ions (like $\text{Na}^+$ or $\text{K}^+$ ) move towards the cathode, while negative ions (like $\text{Cl}^-$ or $\text{NO}_3^-$ ) move towards the anode.

In summary, the salt bridge ensures the galvanic cell operates efficiently by maintaining charge balance and facilitating the flow of ions, which is essential for the cell to generate a potential difference and produce electrical energy.

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25. Calculate the cell potential for a Daniell cell under non-standard conditions, given the following information:

• Concentration of $\text{Cu}^{2+}$ is 0.1 M.
• Concentration of $\text{Zn}^{2+}$ is 0.01 M.
• Standard electrode potentials: $E^\circ_{\text{Cu}^{2+}/\text{Cu}} = +0.34 \, \text{V}$ , $E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76 \, \text{V}$ .

Answer: We can calculate the cell potential using the Nernst equation:

$$E = E^\circ - \frac{0.0591}{n} \log Q$$

For the Daniell cell, the reaction is:

$$\text{Zn} (s) + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu} (s)$$

The standard cell potential is:

$$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = 0.34 \, \text{V} - (-0.76 \, \text{V}) = 1.10 \, \text{V}$$

Now, calculate the reaction quotient $Q$ :

$$Q = \frac{[\text{Zn}^{2+}]}{[\text{Cu}^{2+}]} = \frac{0.01}{0.1} = 0.1$$

Substitute the values into the Nernst equation:

$$E = 1.10 - \frac{0.0591}{2} \log 0.1$$ $$E = 1.10 - \frac{0.0591}{2} (-1)$$ $$E = 1.10 + 0.02955 = 1.12955 \, \text{V}$$

So, the cell potential under the given non-standard conditions is approximately 1.13 V.