Exploring Electrochemical Cell Varieties - An In-Depth Study of Electrolytic and Galvanic cell

What is Electrochemical Cell

An electrochemical cell is a device that uses redox (reduction-oxidation) processes to turn chemical energy into electrical energy or vice versa. It is made up of two half-cells, which are compartments filled with electrodes and electrolytes. In each half-cell, one electrode oxidizes (loses electrons) while the other reduces (gains electrons).

The two half-cells are linked by a conductive channel known as the external circuit, which permits electrons to flow. An electric current is generated as a result of this electron transport. The voltage or potential difference between the two electrodes is measured in volts (V) and is what drives electron migration.

There are two main types of electrochemical cells: galvanic (voltaic) cells and electrolytic cells. In a galvanic cell, also commonly known as a battery, the redox reactions occur spontaneously. Chemical energy is converted into electrical energy, producing a current that can be used to power devices. Examples include alkaline batteries and lead-acid batteries used in cars.

On the other hand, an electrolytic cell uses an external power source to drive a non-spontaneous reaction. This external electrical energy input forces the redox reactions to occur in the desired direction. Electrolytic cells are employed in processes like electroplating, metal refining, and water splitting for hydrogen, electroplating, electrolysis of water, and electrolytic production of chemicals. They allow for the deposition of metals onto a surface, purification of metals, and can even facilitate the decomposition of water into hydrogen and oxygen gases.

Electrochemical cells are essential in a variety of applications. They find application in portable electronic gadgets, transportation, renewable energy systems, and industrial operations. Understanding the concepts and functioning of electrochemical cells is critical for progressing in domains like energy storage, electrochemistry, and materials research.

In conclusion, electrolytic cells employ external electrical energy to drive non-spontaneous reactions, whereas galvanic cells generate electrical energy through spontaneous reactions. Both forms of electrochemical cells are important in a variety of applications, ranging from industrial operations to portable power sources.

Advantages and disadvantages of Electrolytic and Galvanic cell

Advantages of Electrolytic Cells:

Wide range of applications: 

Electrolytic cells are used in a variety of industries for operations such as electroplating, electrolysis, and electrolytic refining. Electrochemical synthesis and electrochemical sensors are also applications.

 Accurate control:

 Electrolytic cells provide accurate control of material deposition or dissolution, making them perfect for sectors where precision is critical.

 Enormous quantities:

 Because electrolytic cells can create enormous amounts of the desired product, they are suitable for industrial-scale manufacturing.

 Purification: 

Electrolytic cells can be used to purify metals by selectively removing contaminants using electrolysis.

Disadvantages of Electrolytic Cells:

 High energy consumption:

High energy usage: 

The operation of electrolytic cells is powered by an external power source. Because of the significant energy consumption that results, they are less energy-efficient than other processes.

 Required Cost: 

Specialized tools and materials are needed to set up and maintain electrolytic cells, which may be expensive.

 Environmental issues:

 The electrolytic process occasionally produces waste byproducts or releases hazardous substances that must be handled carefully and disposed of.

 Limited uses: 

Electrolytic cells have a limited range of uses and may not be appropriate for all chemical reactions. They might not work well for making non-metallic elements or extremely reactive compounds, for instance.

Electrolytic Cell 

Advantages of Galvanic Cells:

 Portable and self-contained: 

Galvanic cells, commonly known as batteries, are portable and can provide a source of electricity 

without the need for an external power supply or continuous connection.

 Longer shelf life: 

Batteries have a relatively long shelf life as they can store energy for an extended period without significant discharge.

 Easy to use: 

Galvanic cells are ready to use and do not require any special installation or setup. They can be easily replaced when depleted.

 Versatile: 

Galvanic cells come in different sizes and configurations, making them suitable for a wide range of applications, from powering small electronic devices to large-scale applications like emergency power backup.

Disadvantages of Galvanic Cells:

 Limited lifespan: 

Batteries have a limited lifespan and will eventually become depleted and need to be replaced or recharged.

 Environmental impact: 

Improper disposal of batteries can lead to environmental pollution due to the presence of toxic metals like mercury and lead.

 Expenses: 

Batteries can be expensive, especially high-capacity ones or rechargeable batteries that require an additional charging system.

 Limited power output: 

Galvanic cells have limited power output compared to other power sources, making them unsuitable for high-power applications. Additionally, their voltage gradually decreases as they discharge, resulting in lower efficiency over time.

 Chemical reactions products: 

The chemical reactions within galvanic cells can generate heat or produce harmful gases, which can be hazardous if not handled properly.

Galvanic Cell 

Major Differences between Electrolytic cell and Galvanic cell

The following are the main variations:

 Energy Transformation:

 Galvanic cells convert chemical energy into electrical energy as a result of internal chemical processes, which enable the passage of electrons from the anode to the cathode. A spontaneous electric current results from this. However, in an electrolytic cell, a non-spontaneous chemical reaction is powered by electrical energy from an external power source. This indicates that the direction of electron flow has switched from cathode to anode.

 Purpose: 

Galvanic cells are typically used to produce and provide electricity, such in batteries. They are frequently employed in gadgets that need a portable and practical energy source. Conversely, electrolytic cells are employed in a variety of electrochemical procedures that call for particular chemical processes to take place. These procedures may involve chemical compound synthesis, electroplating, and electrolysis.

 Construction or Structure: 

Galvanic cells are made up of two distinct half-cells, each having a unique metal electrode submerged in an electrolyte solution. The half-cells are linked by an external circuit, which permits electron flow. The two electrodes of electrolytic cells, on the other hand, are normally submerged in an electrolyte solution. However, unlike galvanic cells, the two electrodes in an electrolytic cell are often made of the same material.

 Spontaneity of Reactions:

The chemical processes in a galvanic cell take place spontaneously or without the need for external energy. The reactions go in a direction that causes an energy release. The processes in an electrolytic cell are not spontaneous; instead, they need an external electrical energy source to propel them in the desired direction.

 Sign of the cell potential: 

The positive cell potential in a galvanic cell signifies that the reaction is spontaneous and is releasing energy. In contrast, a negative cell potential in an electrolytic cell denotes an energy-input-required, non-spontaneous response.

           Overall, the main difference between electrolytic cells and galvanic cells lies in their purpose, energy flow direction, and how the reactions are driven.

Redox Reactions - The Undeniable Force Behind Chemical Reactions

Introduction 

Oxidation and reduction defined based on the addition or removal of oxygen or hydrogen atoms in a chemical reaction. 

Let's explain each process separately:

Oxidation: 

1. Addition of Oxygen: 

Oxidation occurs when an atom or molecule gains oxygen atoms. For example:

   a) Methane (CH4) + Oxygen (O2) -> Carbon Dioxide (CO2) + Water (H2O)

      In this reaction, methane is oxidized because it gains oxygen atoms from O2, resulting in the formation of carbon dioxide.

   b) Sulfur (S) + Oxygen (O2) -> Sulfur Dioxide (SO2)

      Here, sulfur is oxidized as it bonds with oxygen from O2, forming sulfur dioxide.

2. Removal of Hydrogen: 

Oxidation also occurs when an atom or molecule loses hydrogen atoms. For example:

   a) Ethanol (C2H5OH) -> Acetaldehyde (CH3CHO) + Hydrogen (H2)

      In this reaction, ethanol is oxidized as it loses hydrogen atoms, resulting in the formation of acetaldehyde and hydrogen gas.

   b) Glucose (C6H12O6) -> Gluconic Acid (C6H12O7)

      In this case, glucose is oxidized by losing hydrogen atoms, resulting in the formation of gluconic acid.

Reduction:

1. Removal of Oxygen: 

Reduction occur when an atom or molecule loses oxygen atoms. For example:

   a) Copper Oxide (CuO) + Hydrogen (H2) -> Copper (Cu) + Water (H2O)

      Here, copper oxide is reduced as it loses oxygen atoms, leading to the formation of copper metal and water.

   b) Potassium Permanganate (KMnO4) + Sulfur Dioxide (SO2) + Water (H2O) -> Manganese Dioxide (MnO2) + Potassium Sulfate (K2SO4) + Sulfuric Acid (H2SO4)

      In this reaction, potassium permanganate is reduced as it loses oxygen atoms, resulting in the formation of manganese dioxide, potassium sulfate, and sulfuric acid.

2. Addition of Hydrogen: 

Reduction also occurs when an atom or molecule gains hydrogen atoms. For example:

   a) Nitrogen (N2) + Hydrogen (H2) -> Ammonia (NH3)

      In this reaction, nitrogen is reduced as it gains hydrogen atoms, forming ammonia.

   b) Silver Nitrate (AgNO3) + Copper (Cu) -> Silver (Ag) + Copper Nitrate (Cu(NO3)2) + Nitric Acid (HNO3)

      Here, silver nitrate is reduced as it acquires hydrogen atoms from copper, generating silver, copper nitrate, and nitric acid.

    These examples tell us how oxidation involves the addition of oxygen or the removal of hydrogen, while reduction involves the removal of oxygen or the addition of hydrogen in a chemical reaction.

Redox Reactions 

Oxidation and reduction are two important chemical processes that occur in redox reactions. These phrases refer to the exchange of electrons between organisms.

Oxidation: 

The loss of electrons or an increase in the oxidation state of an atom or molecule is referred to as oxidation. In other words, a rise in positive charge or a decrease in negative charge is involved. When a material oxidises, it is referred to as an oxidising agent because it aids in the oxidation of another substance. The oxidized material is frequently decreased during the process.

Reduction: 

Reduction, on the other hand, refers to the addition of electrons or a reduction in an atom's or molecule's oxidation status. It entails either a reduction in positive charge or an increase in negative charge. A reducing agent is a chemical that acquires electrons (undergoes reduction) and encourages the reduction of another molecule.

Oxidation and reduction usually occur concurrently and are interconnected in redox processes. The oxidized species loses electrons, which are acquired by the reduced species. The transfer of electrons causes a shift in the charge distribution among the atoms or molecules involved, resulting in chemical reactions. 

In the reaction of magnesium (Mg) with oxygen (O2) to generate magnesium oxide (MgO), for example:

2Mg + O2 -> 2MgO

Mg is oxidised as it loses two electrons, changing its oxidation state from 0 to +2. Oxygen, on the other hand, is reduced because it gains two electrons, changing its oxidation state from 0 to -2. The overall process comprises electron transfer from Mg to O2, which results in the production of MgO.

    The ideas of oxidation and reduction are fundamental in chemistry, with applications ranging from energy generation (as in batteries and fuel cells) to industrial processes and biological systems.

Examples of Redox Reactions:

1. Rusting of Iron: 

In the presence of water (H2O), iron (Fe) interacts with oxygen (O2) to generate iron(III) oxide (Fe2O3), also known as rust. In this process, iron oxidation and oxygen reduction occur simultaneously:   

   4Fe + 3O2 + 6H2O -> 2Fe2O3 . 6H2O

2. Combustion of Hydrocarbons: 

   

When hydrocarbons such as methane (CH4) or octane (C8H18) burn in the presence of oxygen, they burn, creating carbon dioxide (CO2) and water. The hydrocarbon is oxidised, and oxygen is reduced in this reaction:

   CH4 + 2O2 -> CO2 + 2H2O

3. Photosynthesis: 

 Photosynthesis is the process by which plants utilise sunshine to transform carbon dioxide and water into glucose (C6H12O6) and oxygen. This process includes carbon dioxide reduction and water oxidation:   

   6CO2 + 6H2O + sunlight -> C6H12O6 + 6O2

Difference between Oxidation and Reduction 

Advantages of Redox Reactions:

1. Energy Production: 

 Many redox processes are involved in energy production, such as fuel combustion in engines, electricity generation in batteries, and glucose oxidation during cellular respiration. These reactions enable humans to capture and use energy for a variety of reasons.

2. Metabolism and Respiration: 

Redox processes are critical in biological systems, especially metabolism and respiration. They are essential in the breakdown of dietary molecules, the extraction of energy, and the synthesis of biomolecules necessary for cell activity.   

3. Environmental Cleanup: 

Redox reactions are commonly employed in environmental cleaning operations including water and soil remediation. Certain toxins and pollutants, for example, can be transformed into less dangerous chemicals through reduction reactions by adding reducing agents.  

4. Industrial Processes: 

Redox reactions are often employed in a wide range of industrial activities. Reduction processes, for example, are used in the manufacturing of metals such as aluminum to remove these metals from their ores. Similarly, oxidation-reduction processes have a role in the synthesis of chemicals, polymers, and medicines.

Overall, redox reactions are important in many disciplines, including energy generation, biological activities, environmental remediation, and many industrial uses.   

Electrochemical series and standard Electrode potential

 Electrochemical series and standard electrode potential

Contents In this lesson are,

Introduction

Electrochemical series

An important concept of electrochemical series

Aspects of Electrochemical Series Applications

EMF calculation

Counting the degree of spontaneity in reactions

Gibbs Free Energy Calculation

Estimating a redox reaction's end result

Definition of Standard Electrode Potential

Standard Electrode Potential: Its Importance

Measurement of standard electrode potential

Uses of standard electrode potential

a) Redox Reactions' Spontaneity 

Introduction

Electrochemical series

In chemistry, the electrochemical series is also known as the active series. The periodic table's elements are organized in ascending order according to the electrode potential values they represent. The potential of different electrodes is observed using conventional hydrogen electrodes. Different ions are positioned in an electrochemical array according to their propensity for oxidation or reduction. Whether or not it is metallic. By carefully recording the voltage at the end of the standard hydrogen electrode and the half-cell attached to it, the value of the standard electrode potential is afterward determined.

In comparison to the Standard Hydrogen Electrode, the electrochemical series indicates how electropositive or electronegative the element/ion combination is. The name "half-cell" also applies to this combination. In the SHE, a metal that is more electropositive loses electrons more readily than hydrogen. A more electronegative material, however, has a greater ability to absorb electrons. Typically, an element that is more electronegative will absorb electrons from an element that is more electropositive. Thus, it can be claimed that the electrochemical series serves as a gauge of electronegative character.

Due to low reactivity, metals, like copper and gold, are referred to as "precious" metals and are used to manufacture coins and jewelry. A group of chemical elements grouped according to their standard electrode potentials is known as an electrochemical series. The potential of a cell with one electrode acting as the cathode and a standard hydrogen electrode (SHE) acting as the anode is known as electrode potential. Reduction always takes place at the cathode while oxidation always takes place at the anode.

An important concept of electrochemical series

By definition, hydrogen has an electrode potential of 0.00 (the Standard Hydrogen Potential, or SHE). In relation to it, all other potentials are defined.

High in the Electrochemical Series are the half-cells (element/ion pairs) having a very positive Electrode Potential. They are powerful oxidizing agents.

Reducing agents are the half-cells with negative electrode potential. The value is more negatively correlated with the decreasing power. Non-metals are electronegative, whereas metals are often electropositive.

The most reactive metals are those near the bottom. The non-metals at the top of the series, in contrast, are the most active. Reactivity is therefore lowest in the center. Metals towards the bottom of the series can reduce metals higher up.

Similar to metals, non-metals higher in the series have the ability to oxidize non-metals lower in the series.

Two half-cells are connected to each electrode in each electrochemical cell. One reaction involves oxidation and the other reduction in each half-cell. The oxidation potential and reduction potential are the respective potentials for each reaction.

The total of a cell's oxidative and reducing capacities is known as the cell EMF. It gauges how spontaneously the cell as a whole reacts. It serves as a gauge for how much work a cell can accomplish. By taking the half-cells' standard electrode potential values and adding them suitably, the electrochemical process aids in measuring the EMF cell.

Electrochemical Series

Aspects of Electrochemical Series Applications

a) EMF calculation

Two half-cells are connected to each electrode in each electrochemical cell. One reaction involves oxidation and the other reduction in each half-cell. The oxidation potential and reduction potential are the respective potentials for each reaction.

The total of a cell's oxidative and reducing capacities is known as the cell EMF. It determines how spontaneously the cell as a whole reacts. It serves as a measurement for how much work a cell can accomplish. By taking the half-cells' standard electrode potential values and adding them suitably, the electrochemical process aids in measuring the EMF cell.

E^ocell=E^ored– E^ooxi

where E^ored and E^oox represent the typical reduction potentials of the reducing and oxidizing half-cells, respectively.

b) Counting the degree of spontaneity in reactions

Reactive EMF cells are intimately correlated with the vitality or spontaneity of redox reactions:

The response is spontaneous if the cell EMF is positive; it is non-spontaneous if the cell EMF is negative. Therefore, by examining the reactants and products, we can determine whether a redox reaction can occur spontaneously. We formulate the equations for the half-reactions of reduction and oxidation. Then, adding in accordance with the electrochemical series, their standard electrode potentials. We can determine if a response is spontaneous based on the cellular EMF that results.

c) Gibbs Free Energy Calculation

Another indicator of a reaction's spontaneity is the Gibbs free energy (G⁰cell). The following is how it relates to the EMF unit (E unit).

G⁰cell =nFE⁰cell, where n is the number of involved electrons and F is the Faraday constant, which is equivalent to 96485 coulombs mol^-1.

Once more, based on the cellular EMF signal, we have:

• If the EMF source is positive, the reaction is spontaneous and the Gibbs free energy is positive; if the cell EMF is negative, the reaction is spontaneous and the Gibbs free energy is negative.

d) Estimating a redox reaction's end result

The ultimate product of the reaction can be calculated if using only the reactants, as shown below.

Using the electrochemical series, we put out the standard electrode potential values for each reactant. Then, we determine which has the greatest and least amount of potential for reduction. Once we know these numbers, we may make the following predictions about the outcome:

The cathode reduces the ion with the highest reduction potential, whereas the anode oxidizes the ion with the lowest reduction potential. The reaction's end result is provided to us by oxidized and reduced ions.

Standard Electrode Potential

A measurement of the potential for equilibrium is the standard electrode potential. The potential of the electrode is the difference in potential between the electrode and the electrolyte. The electrode potential is referred to as the standard electrode potential when unity represents the concentrations of all the species involved in a semi-cell.

Definition of Standard Electrode Potential

In an electrochemical cell, the standard electrode potential occurs at, for example, 298 K, 1 atm of pressure, and 1 M of concentration. The typical electrode potential of a cell is denoted by the symbol "E^ocell."

Standard Electrode Potential: Its Importance

Redox reactions, which are composed of two half-reactions, constitute the foundation of all electrochemical cells.

At the anode, there is an oxidation half-reaction that results in an electron loss.

At the cathode, a reduction event occurs that results in an electron gain. The anode to the cathode is where the electrons move as a result.

The difference in the individual potentials of each electrode causes an electric potential to develop between the anode and the cathode (which are dipped in their respective electrolytes).

With the aid of a voltmeter, the cell potential of an electrochemical cell can be determined. A half-individual cell's potential, however, cannot be precisely quantified on its own.

It's also critical to remember that this potential can alter in response to modifications in pressure, temperature, or concentration.

The requirement for standard electrode potential emerges in order to acquire the individual reduction potential of a half-cell.

With the use of a reference electrode known as the standard hydrogen electrode, it is measured (abbreviated to SHE). SHE has an electrode potential of 0 volts.

By connecting an electrode to the SHE and measuring the cell potential of the resulting galvanic cell, the standard electrode potential of the electrode can be determined.

An electrode's oxidation potential is the polar opposite of its reduction potential. As a result, an electrode's standard reduction potential can be used to define its standard electrode potential.

High standard reduction potentials are exhibited by good oxidizing agents, whereas low standard reduction potentials are exhibited by good reducing agents.

Ca2+ has a standard electrode potential of -2.87 V, while F2 has a standard electrode potential of +2.87 V. This suggests that Ca is a reducing agent while F2 is an excellent oxidizing agent.

Measurement of standard electrode potential

The Standard Hydrogen Electrode (SHE) is coupled to a metal (or non-metal) electrode that contains its ion. H2 and H+ ions make up SHE. Under normal circumstances, a certain value of voltage is seen across the electrodes depending on the type of metal and ions used. For the specific metal/ion pairing, this is known as the "standard electrode potential value."

Uses of standard electrode potential

a)Redox Reactions' Spontaneity

The Gibbs free energy, or "G^o," must be negative if a redox reaction occurs on its own. The following equation provides an explanation:

G^ocell = -nFE⁰cell

F is Faraday's constant, and n is the total number of moles of electrons created for every mole of product (approximately 96485 C.mol-1).

The following equation can be used to determine the E⁰cell:

E⁰cell = E⁰cathode – E⁰anode

As a result, the E⁰cell can be calculated by deducting the cathode's standard electrode potential from the anode's. Because both n and F have positive positive values and the G^o value must be negative, the E⁰cell must be positive for a redox reaction to be spontaneous.

This suggests that during an unplanned process,

Since E⁰cell > 0, it follows that E⁰cathode > E⁰anode.

Thus, the cathode and anode's typical electrode potentials can be used to estimate how spontaneously a cell response would occur. It should be noted that the cell's " G^o " in electrolytic cells is positive while the cell's "G^o" in galvanic cells is negative.

Electrolytic cell & Faraday’s law of Electrolysis

Electrolytic cell & Faraday’s  law of Electrolysis 

Electrolytic cell & Faraday’s  law of Electrolysis 

Contents

Electrolytic cell

Faradays law of electrolysis

Definition of Electrolysis

Faradays constant

Faradays First and second laws of electrolysis

Difference between Galvanic and Electrolytic cell

Electrolytic cell

By passing an electric current through the system, it is possible to create a cell that actually operates on a chemical process. They are known as electrolytic cells.

An electrolytic device that employs electrical energy to induce a non-spontaneous redox reaction is known as an electrolytic cell. Certain chemicals can be electrolyzed using electrolytic cells, which are electrochemical cells. For instance, water can be electrolyzed to create gaseous oxygen and gaseous hydrogen with the use of an electrolytic cell. To do this, the non-spontaneous redox reaction's activation energy barrier is overcome by leveraging the flow of electrons (into the reaction region).

In that they both need a salt bridge, have a cathode and anode side, and have a steady flow of electrons from the anode to the cathode, so, electrolytic cells are extremely similar to voltaic (galvanic) cells. But the two cells also differ dramatically from one another.

The following are the three essential parts of electrolytic cells:

i.Cathode

ii.Anode

iii.Electrolyte

The cathode and anode exchange electrons through a medium that is provided by the electrolyte. In electrolytic cells, molten sodium chloride and water with dissolved ions are two common electrolytes. An electrolyte, commonly an ionic chemical that has been dissolved or fused, is in contact with two metallic or electronic conductors (electrodes) that are held apart from one another. The electrodes become positively and negatively charged when connected to a source of direct electric current, respectively.

In the electrolyte, negative ions migrate to the positive electrode (anode) and transfer one or more electrons to it, creating new ions or neutral particles. In the same way, positive ions migrate to the negative electrode (cathode) and combine with one or more electrons, losing some or all of their charge and creating new, lower-charged ions or neutral atoms or molecules.

The two procedures combine to produce a chemical reaction in which the negative ions' electrons are transferred to the positive ions . The electrolysis of sodium chloride (common salt), which results in the formation of sodium metal and chlorine gas, is one example; the energy needed to drive the reaction forward is provided by the electric current. The manufacture of caustic soda and electrodeposition for metal plating or refinement are two other frequent uses of electrolysis.

Examples include are Downs Cell and Nelson cell.

Using an electrolytic cell, as shown below, it is possible to electrolyze molten sodium chloride (NaCl).

 Electrolytic cell

Molten sodium chloride, which comprises dissociated Na+ cations and Cl- anions, is used to saturate inert electrodes. The cathode accumulates electrons and creates a negative charge when an electric current is introduced into the circuit. Now, the sodium cations are directed to the cathode, which is negatively charged. As a result, metallic sodium is created at the cathode. The chlorine atoms are brought to the positively charged cathode at the same time. As a result, chlorine gas (Cl2) is produced at the anode (with loss of 2 electrons, finishing the process). Below are the relevant chemical formulae and the general cell reaction.

Thus, metallic sodium and chlorine gas can be produced by electrolyzing molten sodium chloride in an electrolytic cell.

The main use of electrolytic cells is to create oxygen and hydrogen gas from water. The technique of creating a thin protective layer of one metal on the surface of another metal, known as electroplating, is another noteworthy use of electrolytic cells. They are also employed in the process of removing aluminium from bauxite. It should be mentioned that electrolytic cells are virtually usually used in the industrial manufacture of high-purity aluminium, high-purity copper, and high-purity zinc.

Faraday's law of electrolysis

In 1833, Michael Faraday found that the amount of product generated or absorbed at an electrode during electrolysis and the amount of electrical charge Q that moves through the cell are always related in a straightforward way. Law illustrates the quantitative link between the amount of electrical charge or electricity passed and the substance collected at electrodes.

The half-equation, as an illustration

Ag⁺+e^–→Ag

Above equation Informs us that 1 mol of e- must be provided from the cathode in order for 1 mol of Ag+ to deposit at cathode as 1 mol of Ag.

Electrolysis

A chemical change is induced by electrolysis, which involves passing an electric current through an electrolytic solution to stimulate the passage of ions. A liquid that conducts electricity is known as an electrolyte, or often a salt solution of metal. Electrolysis is the use of electric current to trigger a chemical process that is not naturally occurring.

Faraday Constant (F)

 We may multiply the charge per mole of electrons by the Avogadro constant to get the charge per electron since the negative charge on a single electron is known to be 1.6022 10-¹⁹ C. The Faraday Constant is this number, denoted by the letter F,

F = 1.6022 × 10^–19 C × 6.0221 × 10²³ mol^–1 = 9.649 × 10⁴ C mol^–1

Faraday’s First Law of Electrolysis

It is stated that “The mass of a substance deposited at any electrode is directly proportional to the amount of charge passed.” Mathematically it can be written as

m ∝ Q          (i)

Here:

 “m” is the mass of a substance (in grams) deposited or liberated at an electrode. “Q” is the amount of charge (measured in coulombs) or it is the electricity passed during electrolysis

By converting the sign of proportionality in equation (i) it becomes as follows

m=ZQ

 where Z is the constant of proportionality. Measured in g/c stands for grams per coulomb. Alternatively, it is known as the electrochemical equivalent. Z is the mass of an object deposited at electrodes during electrolysis while passing one coulomb of charge.

Faraday’s Second Law   

It states that “the mass of a substance deposited at any electrode on passing a certain amount of charge is directly proportional to its chemical equivalent weight.” Or “when the same quantity of electricity is passed through several electrolytes, the mass of the substances deposited are proportional to their respective chemical equivalent or equivalent weight”. Mathematically it can be represented as follows

w ∝ E

 w = mass of the substance

E = equivalent weight of the substance

Second law is also written as follows

 w1/w2=E1/E2

The equivalent weight or chemical equivalent of a substance is defined as ratio of its atomic weight and its valency.

Equivalent weight=Atomic weight/Valency

 

Difference between Galvanic and Electrolytic cell

Difference between Galvanic and Electrolytic cell

Electrochemical Cell and Its Types, Galvanic cell

 Electrochemical Cell and Its Types, Galvanic cell

Electrochemical Cell and Its Types, Galvanic cell

Here you will learn about,

Electrochemistry

Electrochemical cell and its types

Galvanic cell

Gibbs Free Energy Calculation using EMF

Equilibrium Constant Calculation Using EMF

Nernst Equation

Finding Concentration cell potential using Nernst Equation

Definition of Electrochemistry

The field of study known as "Electrochemistry" combines the study of ionic solutions with that of solid-state electrons. Any material that will be used in electrochemistry requires essential measurements, depending on the uses, to confirm its susceptibility, conductance, responsiveness, interaction, consistency, and lifespan in a given medium.

The study of the correlation between electrical energy and chemical changes is the focus of the branch of chemistry known as electrochemistry. Electrochemical reactions are those in which electric currents are either generated or input. These responses can be roughly divided into two categories:

Electrical energy produces chemical change i.e., the electrolysis phenomenon

Chemical energy to electrical energy conversion. i.e., the production of electricity using redox reactions that occur spontaneously.

An oxidation or reduction reaction at a polarized electrode surface is the subject of electrochemistry, which studies the movement of electrons in such reactions. At a particular potential, each analyte is oxidized or reduced, and the current measured is proportional to concentration. This method is an effective approach to bioanalysis.

Galvanic cell

Galvanic, also known as Voltaic, and electrolytic cells are the two varieties of electrochemical cells. While electrolytic cells utilize non-spontaneous reactions and therefore need an external electron source, such as a DC battery or an AC power source, galvanic cells get their energy from spontaneous redox reactions. Anode and cathode, which can be formed of the same metal or two distinct metals, as well as an electrolyte, in which the two electrodes are submerged, make up both galvanic and electrolytic cells.

DC electrical power is usually generated by galvanic cells. A straightforward galvanic cell would just have one electrolyte separated from it by a semi-permeable membrane, or a more complicated one would have two distinct half-cells joined by a salt bridge. In order to balance the developing charges at the electrodes, the salt bridge contains an inert electrolyte like potassium sulphate, whose ions will diffuse into the half-cells.

Galvanic cell Diagram

The anode is where oxidation happens, and the cathode is where reduction happens. The anode is the negative terminal for the galvanic cell because the anode's reaction serves as the source of electrons for the current.

Voltage is an intense attribute, meaning it is independent of the system's size and material content. Since galvanic cells contain a positive EMF, we want to rearrange the equation so that it will result in a positive value when the other EMF is added.

 Example of Galvanic cell,

Galvanic cell Example

The two EMF readings for the zinc half-reaction are +0.382 V and +1.221 V. We simply sum them all together to obtain a rough estimate of 1.5 V, which represents the EMF of an alkaline AA battery.

Gibbs Free Energy Calculation using EMF

Let's say someone asks us to express the energy in additional thermodynamic terms. Let's apply the following equation, where n represents the number of electrons exchanged, E represents the EMF in its standard condition, and F represents the Faraday constant, which is 96,485 C/mol.

  

Gibbs Free Energy Calculation using EMF

Instead of joules, Gibbs free energy is typically stated in kilojoules. We can determine from the sign which way the reaction must change to achieve equilibrium. Accordingly, a system operating under normal circumstances would have to move to the right, transforming some reactants into products before coming to equilibrium. The magnitude shows us how far away from equilibrium the standard state is.

Equilibrium Constant Calculation Using EMF

Assume that in order to determine how favorable this reaction is; it is necessary to determine the equilibrium constant K under standard conditions. The high K value suggests that the reaction will proceed fully to completion and is particularly beneficial to the products. For the batteries, the reaction will proceed until G^o =0, or equilibrium, has been reached.

The value of ΔG equals zero when the reactants and products of the electrochemical cell are in equilibrium. The reaction quotient and the equilibrium constant (Kc) are the same at this point. Because Δ G = -nFE, the equilibrium cell potential is also 0.

The following equation is generated by substituting the values of Q and E into the Nernst equation.

0 = E⁰_cell – (RT/nF) ln K_c

The equation is changed by converting the natural logarithm into base-10 logarithm and replacing T=298K (standard temperature). 

E⁰_cell = (0.0592V/n) log K_c

The following equation created by rearranging this one.

log K_c = (nE⁰_cell)/0.0592V

As a result, the equilibrium constant's link to the standard cell potential is found. The value of E⁰_cell will be greater than 0 when Kc is greater than 1 (you know the value of Kc is directly related to  E⁰ because value of Kc present in log) , indicating that the equilibrium supports the forward reaction. Similarly, E⁰_cell will have a negative value when Kc is less than 1, indicating that the opposite reaction will likely be preferred.

Nernst Equation

“Nernst equation is an equation relating the capacity of an atom/ion to take up one or more electrons (reduction potential) measured at any conditions to that measured at standard conditions (standard reduction potentials) of 298K and one molar or one atmospheric pressure.”

Walther Hermann Nernst, a German chemist, developed the equation. The cell potential of an electrochemical cell at any given temperature, pressure, and reactant concentration is frequently determined using the Nernst equation.

The standard cell potential, temperature, reaction quotient, and the cell potential of an electrochemical cell are all related by the Nernst equation. The Nernst equation can be used to calculate the cell potentials of electrochemical cells even in unusual circumstances.

Nernst Equation

E_cell = Cell Potential Of The Cell

F = Faraday Constant

E⁰ = Cell Potential Under Standard Conditions

Product / Reactant =Q = Reaction Quotient

R = Universal Gas Constant

T = Temperature

N = Number Of Electrons Transferred In The Redox Reaction

Finding Concentration cell potential using Nernst Equation

Consider a concentration cell, a particular type of galvanic cell that consists of two identical half-cells of the same material that differ only in concentration. The sodium ion, potassium ion, or Calcium ion pumps in our cell membranes, the ATP synthase employed in energy production, and the concentration gradients in our nerve cells are all examples of concentration cells.

In addition to the Henderson-Hasselbalch equation, the thermodynamics equation, is where the Nernst equation originates. When a concentration cell tries to reach equilibrium, a little voltage is generated. The Nernst Equation can be used to determine the potential created by a concentration cell and is as follows:

Concentration cells Nernst equation

The standard state EMF is 0 for any concentration cell because the two half-cells have identical half-reactions.