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Introduction to electrochemistry

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Intros
Lessons
  1. What is electrochemistry?
  2. The electrochemical cell
  3. Redox reactions
  4. Introducing half-equations
  5. Reducing and oxidising agents
  6. The activity series
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Examples
Lessons
  1. Identify whether reduction or oxidation is taking place in the half-equations.
    State whether the half-equations below show reduction or oxidation, giving a reason for your answer.
    1. Au  \, →  \, Au+ + e-
    2. Na  \, →  \, Na+ + e-
    3. Cl2 + 2e-  \, →  \, 2Cl-
  2. Write simple half-equations showing reduction and oxidation taking place.
    Write a balanced half-equation for the following chemical processes.
    1. Bromide ions being oxidised back to elemental bromine.
    2. Sodium metal being oxidised.
    3. Oxygen gas being reduced to oxide ions.
Topic Notes
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Introduction to Electrochemistry

Electrochemistry is a fascinating branch of chemistry that explores the relationship between electrical and chemical processes. At its core are redox reactions, where electrons are transferred between species, leading to changes in oxidation states. These reactions form the basis of electrochemical cells, devices that convert chemical energy into electrical energy or vice versa. Electrochemistry has numerous practical applications, from batteries and fuel cells to corrosion prevention and electroplating. The introduction video serves as a crucial starting point for understanding these fundamental concepts. It provides a visual representation of electron flow, explains the structure of electrochemical cells, and demonstrates how redox reactions drive the generation of electricity. By grasping these basic principles, students can better appreciate the role of electrochemistry in modern technology and its potential for solving energy-related challenges. This foundation is essential for delving deeper into more complex electrochemical phenomena and their wide-ranging applications in science and industry.

Fundamentals of Redox Reactions

Reduction and oxidation, commonly referred to as redox reactions, are fundamental concepts in chemistry that describe the transfer of electrons between chemical species. These processes are essential in understanding various chemical reactions, from simple rust formation to complex biological processes.

Reduction is the process by which a chemical species gains electrons. This can be remembered using the mnemonic "OIL RIG," where "RIG" stands for "Reduction Is Gain" of electrons. When a substance is reduced, it experiences a decrease in its oxidation state. Reduction can also be defined in terms of hydrogen and oxygen: it occurs when a substance gains hydrogen atoms or loses oxygen atoms.

Oxidation, on the other hand, is the process by which a chemical species loses electrons. Following the "OIL RIG" mnemonic, "OIL" represents "Oxidation Is Loss" of electrons. During oxidation, a substance's oxidation state increases. In terms of hydrogen and oxygen, oxidation happens when a substance loses hydrogen atoms or gains oxygen atoms.

The concept of electron transfer is crucial in understanding redox reactions. When one species is oxidized (loses electrons), another must be reduced (gains electrons) simultaneously. This interdependence is a key characteristic of redox reactions. For example, in the rusting of iron, iron atoms are oxidized as they lose electrons to oxygen molecules, which are subsequently reduced.

It's important to note that reduction and oxidation always occur together. One cannot happen without the other, which is why they are collectively referred to as redox reactions. This interdependence ensures that electrons are conserved throughout the reaction.

The "OIL RIG" mnemonic is an excellent tool for remembering the basics of redox reactions. It encapsulates the essence of these processes: Oxidation Is Loss, Reduction Is Gain. This simple phrase helps students and professionals alike to quickly recall the fundamental principles of electron transfer in chemical reactions.

In more complex systems, such as electrochemical cells, the separation of reduction and oxidation reactions becomes more apparent. The anode is where oxidation occurs, releasing electrons that travel through an external circuit to the cathode, where reduction takes place. This flow of electrons can be harnessed to produce electrical energy, as seen in batteries.

Understanding redox reactions is crucial in many fields, including biochemistry, where they play a vital role in cellular respiration and photosynthesis. In these processes, electrons are transferred between molecules in a series of redox reactions, ultimately driving the production or consumption of energy in living organisms.

To further illustrate the concepts of reduction and oxidation, consider the reaction between zinc and hydrochloric acid. Zinc atoms lose electrons and are oxidized, forming zinc ions. Simultaneously, hydrogen ions in the acid gain these electrons and are reduced to form hydrogen gas. This example demonstrates how the loss of electrons by one species (oxidation of zinc) is always accompanied by the gain of electrons by another (reduction of hydrogen ions).

In conclusion, reduction and oxidation are inseparable processes that involve the transfer of electrons between chemical species. The "OIL RIG" mnemonic serves as a helpful reminder of their fundamental nature. By understanding these concepts, one can better comprehend a wide range of chemical phenomena, from simple metal corrosion to complex biological processes that sustain life.

Electrochemical Cells and Voltaic Cells

Electrochemical cells, also known as voltaic cells or galvanic cells, are devices that convert chemical potential energy into electrical energy through redox reactions. These cells play a crucial role in various applications, from powering portable electronics to large-scale energy storage systems. Understanding their structure and function is essential for anyone interested in chemistry, physics, or electrical engineering.

The basic structure of an electrochemical cell consists of several key components:

  1. Anode: This is the negative electrode where oxidation occurs. The anode material loses electrons during the electrochemical reaction, becoming oxidized.
  2. Cathode: The positive electrode where reduction takes place. The cathode material gains electrons, becoming reduced in the process.
  3. Electrolyte: A solution containing ions that allows for the flow of electric current between the electrodes.
  4. Salt Bridge: A device that connects the two half-cells, allowing for ion flow to maintain electrical neutrality.

The function of an electrochemical cell relies on the principle of spontaneous electron transfer from a material with a lower reduction potential to one with a higher reduction potential. This transfer of electrons generates an electric current that can be harnessed for various purposes.

Here's how the process works in detail:

  1. At the anode, a spontaneous oxidation reaction occurs, releasing electrons into the external circuit.
  2. These electrons flow through the external circuit, powering any connected devices.
  3. The electrons then reach the cathode, where they participate in a reduction reaction.
  4. To maintain electrical neutrality, ions flow through the salt bridge between the two half-cells.

The salt bridge is a critical component of the electrochemical cell. It serves two primary functions:

  1. Completing the electrical circuit by allowing ion flow between the half-cells.
  2. Preventing the direct mixing of the electrolyte solutions, which would lead to unwanted reactions.

The conversion of chemical potential energy into electrical energy in a voltaic cell is governed by thermodynamics. The cell potential, measured in volts, is determined by the difference in reduction potentials between the cathode and anode reactions. This potential difference drives the flow of electrons through the external circuit, generating usable electrical energy.

Electrochemical cells find applications in various fields, including:

  • Batteries for portable electronics and electric vehicles
  • Fuel cells for clean energy production
  • Electroplating in industrial processes
  • Corrosion prevention in metallic structures

Understanding the principles of electrochemical cells is crucial for developing more efficient and sustainable energy storage solutions. As technology advances, researchers continue to explore new materials and designs to improve the performance and longevity of voltaic cells, addressing global energy challenges and promoting a greener future.

In conclusion, electrochemical cells represent a fascinating intersection of chemistry and physics, demonstrating how fundamental scientific principles can be harnessed for practical applications. By converting chemical potential energy into electrical energy, these devices have revolutionized our ability to store and utilize power, playing a vital role in modern technology and energy systems.

Half Equations and Electrode Processes

Half equations are a fundamental concept in electrochemistry, representing the separate processes of oxidation and reduction that occur during redox reactions. These equations provide a detailed view of electron transfer at the molecular level, making them essential for understanding electrode processes and overall redox reactions.

Oxidation half equations show the loss of electrons from a species, while reduction half equations demonstrate the gain of electrons. For example, the oxidation of zinc can be represented as: Zn Zn² + 2e. Here, zinc atoms lose two electrons to form zinc ions. Conversely, a reduction half equations might show copper ions gaining electrons: Cu² + 2e Cu.

Balancing half equations is crucial for accurately representing the chemical changes occurring at electrodes. The process involves ensuring that the number of atoms and charges on both sides of the equation are equal. For complex balancing half equations, especially those involving oxygen or hydrogen, additional steps may be necessary. In acidic solutions, H ions and HO molecules can be added to balance oxygen and hydrogen atoms. In basic solutions, OH ions and HO molecules serve this purpose.

Consider the oxidation of iron(II) to iron(III) in acidic solution: Fe² Fe³ + e. This equation is already balanced in terms of atoms and charge. However, a more complex example like the oxidation of sulfite ions to sulfate ions requires additional steps: SO² + HO SO² + 2H + 2e. Here, water and hydrogen ions are added to balance the oxygen atoms and overall charge.

The significance of half equations in understanding redox reactions cannot be overstated. They allow chemists to break down complex redox processes into their constituent parts, making it easier to analyze and predict reaction outcomes. By combining oxidation and reduction half equations, we can construct full redox equations, providing a complete picture of electron transfer between species.

Half equations are particularly useful in electrochemistry for describing processes occurring at individual electrodes in electrochemical cells. For instance, in a zinc-copper cell, the zinc electrode undergoes oxidation (Zn Zn² + 2e), while the copper electrode experiences reduction (Cu² + 2e Cu). Understanding these separate processes helps in calculating cell potentials and predicting the direction of spontaneous reactions.

Moreover, half equations play a crucial role in industrial applications, such as electrolysis and electroplating. In electrolysis, applied electrical energy drives non-spontaneous redox reactions, and balancing half equations help in predicting the products formed at each electrode. For electroplating, half equations describe the deposition of metal ions onto a surface, allowing for precise control of the plating process.

Balancing half equations also introduces students to important chemical principles like the conservation of mass and charge. It reinforces the idea that electrons are neither created nor destroyed in chemical reactions but are transferred between species. This concept is fundamental to understanding more advanced topics in chemistry and materials science.

In conclusion, half equations are powerful tools for representing and analyzing redox reactions. By breaking down complex processes into manageable parts, they provide insight into electron transfer, help in balancing equations, and form the basis for understanding a wide range of electrochemical phenomena. Mastering the concept of half equations and electrode processes is essential for students and professionals alike in fields ranging from basic chemistry to advanced materials science and engineering.

Oxidizing and Reducing Agents

Oxidizing and reducing agents play crucial roles in redox reactions, which are fundamental to many chemical processes. Understanding these agents is essential for comprehending electron transfer in chemical reactions. Oxidizing agents, also known as oxidants, are substances that accept electrons from other species, causing them to be oxidized. Conversely, reducing agents, or reductants, are substances that donate electrons to other species, causing them to be reduced.

Oxidizing agents are characterized by their ability to gain electrons, thereby reducing themselves in the process. Common examples of oxidizing agents include oxygen, hydrogen peroxide, and halogens like chlorine and bromine. These substances are typically found on the right side of the periodic table, particularly in groups 16 and 17. Oxygen, for instance, is a powerful oxidizing agent due to its high electronegativity and tendency to form stable compounds with many elements.

Reducing agents, on the other hand, are substances that readily lose electrons, oxidizing themselves in the process. Examples of common reducing agents include hydrogen, carbon monoxide, and metals like sodium and magnesium. These substances are often found on the left side of the periodic table, particularly in groups 1 and 2. Alkali metals, such as sodium and potassium, are strong reducing agents due to their low ionization energies and tendency to form positive ions.

The position of elements in the periodic table provides valuable insights into their potential as oxidizing or reducing agents. Generally, elements on the left side of the periodic table (metals) tend to be good reducing agents, while elements on the right side (non-metals) tend to be good oxidizing agents. This trend is related to the electronegativity of elements, which increases from left to right across the periodic table.

Identifying oxidizing and reducing agents in chemical reactions requires careful analysis of the changes in oxidation states of the elements involved. The species that undergoes an increase in oxidation state is being oxidized and acts as the reducing agent, while the species that undergoes a decrease in oxidation state is being reduced and acts as the oxidizing agent. For example, in the reaction between zinc and hydrochloric acid (Zn + 2HCl ZnCl + H), zinc is oxidized (acts as the reducing agent) while hydrogen ions are reduced (act as the oxidizing agent).

Understanding the roles of oxidizing and reducing agents is crucial in various fields, including chemistry, biology, and materials science. These agents are essential in processes such as metal extraction, battery technology, and biological electron transport chains. By recognizing the characteristics and behaviors of oxidizing and reducing agents, scientists and engineers can better control and utilize redox reactions in numerous applications, from industrial processes to environmental remediation.

The Activity Series and Standard Reduction Potentials

The activity series is a powerful tool in chemistry that allows us to predict the likelihood of oxidation for different metals. This series is closely related to periodic table trends and plays a crucial role in understanding chemical reactivity. The activity series ranks metals based on their reactivity, with the most reactive metals at the top and the least reactive at the bottom. This ranking provides valuable insights into how readily a metal will undergo oxidation, losing electrons in the process.

Understanding the activity series is essential for chemists and students alike, as it helps predict the outcomes of various chemical reactions involving metals. For instance, a metal higher in the series can replace a metal lower in the series in a compound. This knowledge is particularly useful in fields such as metallurgy, corrosion prevention, and battery technology.

The activity series is closely linked to periodic table trends. As we move down a group in the periodic table, the reactivity of metals generally increases. This is because atoms become larger, and their outer electrons are farther from the nucleus, making them easier to lose. Conversely, as we move across a period from left to right, the reactivity of metals typically decreases. This trend is due to the increasing nuclear charge, which holds electrons more tightly.

While the activity series provides a qualitative ranking of metal reactivity, the standard reduction potential table offers a more quantitative approach to understanding electrochemical reactions. The standard reduction potential, measured in volts, indicates the tendency of a chemical species to be reduced (gain electrons). This table is a cornerstone of electrochemistry and is essential for calculating the voltage of electrochemical cells and predicting the spontaneity of redox reactions.

The standard reduction potential table lists half-reactions and their corresponding standard reduction potentials under standard conditions (1 M concentration, 1 atm pressure, and 25°C). The more positive the reduction potential, the greater the tendency for the species to be reduced. Conversely, a more negative reduction potential indicates a greater tendency for the reverse reaction (oxidation) to occur.

The relationship between the activity series and the standard reduction potential table is inverse. Metals high in the activity series (more reactive) have more negative reduction potentials, while those lower in the series (less reactive) have more positive reduction potentials. This correlation allows chemists to use both tools interchangeably, depending on whether they need qualitative or quantitative information about metal reactivity.

In practical applications, the activity series and standard reduction potentials are invaluable for designing corrosion protection systems, developing new battery technologies, and optimizing industrial processes involving metal extraction and purification. For example, in galvanic corrosion prevention, a more reactive metal (higher in the activity series) is used as a sacrificial anode to protect a less reactive metal from oxidation.

As students progress in their chemistry education, mastering the concepts of the activity series and standard reduction potentials becomes increasingly important. These tools not only aid in predicting chemical reactions but also provide a deeper understanding of the fundamental principles governing electron transfer processes in chemistry. By connecting these concepts to periodic table trends, students can develop a more comprehensive view of chemical reactivity and its relationship to atomic structure.

Applications of Electrochemistry

Electrochemistry plays a crucial role in our daily lives and various industries, with its principles underpinning numerous practical applications. Understanding redox reactions is key to appreciating these applications, as they form the basis of electrochemical processes. One of the most common and essential applications of electrochemistry is in batteries. From the small button cells in watches to the large lithium-ion batteries powering electric vehicles, these devices convert chemical energy into electrical energy through controlled redox reactions. The widespread use of portable electronic devices and the growing electric vehicle market have made battery technology a critical area of research and development.

Fuel cells represent another significant application of electrochemistry, offering a clean and efficient method of energy conversion. Unlike batteries, fuel cells continuously convert chemical energy from fuel into electricity as long as fuel is supplied. They are increasingly being used in various sectors, including transportation and stationary power generation, due to their high efficiency and low environmental impact. Hydrogen fuel cells, in particular, are gaining attention as a potential solution for sustainable energy production.

Electroplating is a widely used industrial process that relies on electrochemical principles. This technique involves depositing a thin layer of metal onto a conductive surface using electricity. Electroplating is used for various purposes, including improving corrosion resistance, enhancing aesthetic appeal, and increasing durability. It finds applications in industries ranging from automotive manufacturing to jewelry production. The process allows for precise control over the thickness and composition of the deposited layer, making it invaluable in many manufacturing processes.

Corrosion prevention is another critical application of electrochemistry. Corrosion, an electrochemical process itself, causes significant economic losses and safety concerns across various industries. Understanding the principles of electrochemistry allows for the development of effective corrosion prevention strategies. These include cathodic protection, where a sacrificial metal is used to protect the main structure, and the use of corrosion inhibitors that interfere with the electrochemical reactions causing corrosion. In the oil and gas industry, for example, these techniques are crucial for protecting pipelines and offshore structures from the corrosive effects of seawater and other harsh environments.

The importance of understanding redox reactions in these applications cannot be overstated. Redox reactions form the core of electrochemical processes, determining the efficiency and effectiveness of batteries, fuel cells, and electroplating techniques. By manipulating these reactions, scientists and engineers can develop more efficient energy storage solutions, improve industrial processes, and create more durable materials. As we continue to face global challenges related to energy and sustainability, the practical applications of electrochemistry will undoubtedly play an increasingly important role in shaping our technological future.

Conclusion

In this article, we've explored the fundamental concepts of electrochemistry, delving into the world of redox reactions and their applications. The introduction video serves as a crucial starting point, providing a solid foundation for understanding these complex processes. We've covered key topics such as oxidation and reduction, electrochemical cells, and the importance of electron transfer in various chemical reactions. The significance of electrochemistry extends far beyond the laboratory, with applications in energy storage, corrosion prevention, and environmental science. We encourage readers to further explore this fascinating field, as it continues to play a vital role in technological advancements and scientific breakthroughs. Whether you're a student, researcher, or simply curious about the world around you, the principles of electrochemistry offer valuable insights into the nature of chemical reactions and their practical applications. Continue your journey by exploring additional resources and discovering how redox reactions fundamentals shape our modern world.

Introduction to Electrochemistry

What is electrochemistry? The electrochemical cell

Step 1: Understanding Redox and Electrochemistry

Electrochemistry is a branch of chemistry that deals with the relationship between electrical energy and chemical changes. This chapter begins by exploring redox reactions, which are reactions involving the transfer of electrons between two species. The term "redox" is derived from the two key processes involved: reduction (gain of electrons) and oxidation (loss of electrons). Understanding these fundamental concepts is crucial as they form the basis of electrochemical reactions.

Step 2: Key Definitions and Basics

In this introductory section, we will cover some essential definitions and basic concepts. An electrochemical cell is a device that converts chemical energy into electrical energy through redox reactions. The two main types of electrochemical cells are galvanic (or voltaic) cells and electrolytic cells. In a galvanic cell, spontaneous redox reactions generate electrical energy, whereas in an electrolytic cell, electrical energy is used to drive non-spontaneous chemical reactions.

Step 3: Components of an Electrochemical Cell

An electrochemical cell consists of two electrodes: the anode and the cathode. The anode is where oxidation occurs (loss of electrons), and the cathode is where reduction occurs (gain of electrons). These electrodes are typically made of metals and are immersed in an electrolyte solution that facilitates the flow of ions. The electrolyte is a substance that contains free ions and conducts electricity.

Step 4: Writing Half Equations

To understand the processes occurring at each electrode, we write half equations. A half equation shows either the oxidation or reduction process separately. For example, in a zinc-copper electrochemical cell, the half equation for the oxidation of zinc at the anode is Zn Zn² + 2e, and the half equation for the reduction of copper at the cathode is Cu² + 2e Cu. These half equations help in identifying the oxidizing and reducing agents in the reaction.

Step 5: Identifying Oxidizing and Reducing Agents

In a redox reaction, the substance that gains electrons is called the oxidizing agent, and the substance that loses electrons is called the reducing agent. For instance, in the zinc-copper cell, zinc acts as the reducing agent because it loses electrons, and copper ions act as the oxidizing agent because they gain electrons.

Step 6: The Activity Series

The activity series is a table that ranks metals based on their tendency to be oxidized. Metals higher in the series are more likely to lose electrons and be oxidized compared to those lower in the series. This concept is useful in predicting the outcomes of redox reactions and understanding the reactivity of different metals.

Step 7: Conductivity in Electrolyte Solutions

Electrolyte solutions are crucial for the functioning of electrochemical cells. Strong acids and bases dissociate completely in water, producing ions that can conduct electricity. For example, hydrochloric acid (HCl) dissociates into hydrogen ions (H) and chloride ions (Cl) in water. The presence of these free ions allows the solution to conduct electricity, enabling the flow of electrons in the electrochemical cell.

Step 8: The Electrochemical Cell in Action

In an electrochemical cell, chemical potential energy is converted into electrical energy. This process involves the movement of electrons from the anode to the cathode through an external circuit, creating an electric current. The potential difference between the two electrodes can be measured using a voltmeter. The salt bridge, typically filled with an electrolyte like potassium nitrate, maintains electrical neutrality by allowing the flow of ions between the two half-cells.

Step 9: Measuring Potential Difference

The potential difference, or voltage, between the anode and cathode is a measure of the cell's ability to produce electrical energy. This potential difference is due to the difference in chemical potential energy between the two electrodes. By connecting a voltmeter across the electrodes, we can measure this voltage and determine the efficiency of the electrochemical cell.

Step 10: Summary of Electrochemical Cells

In summary, an electrochemical cell is a device that converts chemical energy into electrical energy through redox reactions. The key components include the anode, cathode, electrolyte solution, and salt bridge. By understanding the principles of redox reactions, half equations, and the activity series, we can predict and explain the behavior of different electrochemical cells. This knowledge is fundamental to various applications, including batteries, electroplating, and corrosion prevention.

FAQs

  1. What is the difference between oxidation and reduction?

    Oxidation is the loss of electrons by a species, resulting in an increase in its oxidation state. Reduction is the gain of electrons by a species, resulting in a decrease in its oxidation state. These processes always occur together in redox reactions, with one species being oxidized while another is reduced.

  2. How does an electrochemical cell work?

    An electrochemical cell converts chemical energy into electrical energy through redox reactions. It consists of two half-cells connected by a salt bridge. At the anode, oxidation occurs, releasing electrons that flow through an external circuit to the cathode, where reduction takes place. The salt bridge maintains electrical neutrality by allowing ion flow between the half-cells.

  3. What is the significance of the activity series in chemistry?

    The activity series ranks metals based on their reactivity, with more reactive metals at the top. It helps predict the outcomes of reactions involving metals, such as displacement reactions. Metals higher in the series can replace those lower in compounds. This information is crucial in fields like metallurgy and corrosion prevention.

  4. How are standard reduction potentials used in electrochemistry?

    Standard reduction potentials provide a quantitative measure of a species' tendency to be reduced. They are used to calculate cell potentials, predict the spontaneity of redox reactions, and determine the direction of electron flow in electrochemical cells. The more positive the reduction potential, the greater the tendency for the species to be reduced.

  5. What are some practical applications of electrochemistry?

    Electrochemistry has numerous practical applications, including:

    • Batteries and fuel cells for energy storage and conversion
    • Electroplating for coating metals
    • Corrosion prevention in metal structures
    • Water treatment and purification
    • Sensors for chemical and biological detection
    These applications rely on the principles of redox reactions and electron transfer to achieve their specific functions.

Prerequisite Topics for Introduction to Electrochemistry

Before diving into the fascinating world of electrochemistry, it's crucial to have a solid foundation in certain fundamental chemistry concepts. Two key prerequisite topics that play a significant role in understanding electrochemistry are the history and development of the periodic table and balancing chemical equations.

The periodic table is a cornerstone of chemistry, and its understanding is vital for electrochemistry. The arrangement of elements in the periodic table provides crucial information about their electronic structure and chemical properties. In electrochemistry, we often deal with redox reactions involving various elements, and knowing their position in the periodic table helps predict their behavior in electrochemical processes. The periodic table trends such as electronegativity, atomic radius, and ionization energy are particularly relevant when studying electron transfer reactions in electrochemistry.

Equally important is the skill of balancing chemical equations. Electrochemistry involves numerous reactions, often complex ones, that need to be properly balanced. When dealing with electrochemical cells or electrolysis processes, you'll frequently encounter half-reactions that must be balanced correctly. The ability to balance equations ensures that you can accurately represent the transfer of electrons in redox reactions, which is at the heart of electrochemistry. Moreover, balancing chemical equations is essential for stoichiometric calculations in electrochemical systems.

Understanding these prerequisite topics provides a strong foundation for grasping more advanced electrochemical concepts. For instance, knowledge of the periodic table helps in predicting the standard reduction potentials of elements, while proficiency in balancing equations is crucial when working with the Nernst equation or calculating cell potentials. As you delve deeper into electrochemistry, you'll find that these fundamental skills are constantly applied and built upon.

By mastering the periodic table and equation balancing, you'll be better equipped to understand the principles of electrochemistry, such as electrode processes, electrochemical cells, and the applications of electrochemistry in various fields like battery technology, corrosion prevention, and electroplating. These prerequisite topics serve as the building blocks for more complex electrochemical theories and applications, making them indispensable for anyone looking to excel in this field.

In conclusion, a solid grasp of the periodic table and the ability to balance chemical equations are crucial stepping stones on your journey to mastering electrochemistry. These foundational concepts will continually support your learning and help you navigate the intricacies of electrochemical phenomena with greater ease and understanding.

In this lesson, we will learn:

  • The definition of electrochemical cell, reduction and oxidation.
  • How to write half equations showing oxidation and reduction.
  • How to identify oxidizing and reducing agents in reactions.
  • The activity series as a predictor of relative oxidative strength.

Notes:

  • Electrochemistry is similar to acid and base chemistry. Think about some of the key points we saw in in the Acid-Base theory and Solubility equilibria chapters:
    • Acids and bases are ionic substances that make ionic solutions when dissolved.
    • Acids and bases react with each other and with metals to produce salts, which are also ionic substances that make ionic solutions when dissolved.
    • Ionic solutions conduct electricity.
    • Reactions between some acids and bases or salts can release lot of energy.

      The reactions that release energy can be set up in chemical ‘circuits’ called electrochemical cells. An electrochemical cell is a system that converts chemical potential energy into electrical energy. The electrical energy (electricity!) is produced by the movement of electrons from one of the reactants to another. See example diagram below.

      Because electrons are negatively charged, reactions in electrochemical cells will always result in at least two of the atoms or ions in the reaction changing their individual charge. One is the atom losing the electrons and one is the atom gaining these electrons.

  • Any reaction where an atom or ion changes its individual charge is called a redox reaction. The word redox comes from reduction and oxidation, the two opposite effects happening in a redox reaction.
    • Reduction is when any atom or compound gains electrons, hydrogen, or loses oxygen.
    • Oxidation is when any atom or compound loses electrons, hydrogen, or gains oxygen.
    Reduction and oxidation are complementary – one cannot happen without the other. This means that if there is a chemical being reduced in a reaction, there will be a chemical getting oxidized (that’s why the overall reaction is called redox!).

    There is a trick to remember what oxidation and reduction do in chemical reactions – OIL-RIG:
    • Oxidation Is Loss (OIL) of electrons or hydrogen.
    • Reduction Is Gain (RIG) of electrons or hydrogen.

  • Any redox reaction can be split up to show reduction and oxidation separately in a half-equation. To do this, you need to use electrons to show the atom either gaining or losing electrons. We use the symbol e- to do this.
    For example, the two-electron oxidation of copper can be written in a half-equation:

    Cu  \, →  \, Cu2+ +  \, 2e-

    The reduction of chlorine can be written in a half-equation:

    Cl2 + 2e-  \, →  \, 2Cl-

    In half-equations, electrons in the reactants show reduction, and electrons in the products show oxidation.

    The e- symbol for electrons is ONLY used in half-equations; do not use it in full equations. In full equations showing the ‘real’ process, those electrons are contained in the other chemical that is being reacted!

  • Chemicals that reduce other chemicals are called reducing agents, while chemicals that oxidize other chemicals are called oxidizing agents.
    An agent or agency is someone/something that does things – a reducing agent does reduction to other chemicals, so the reducing agent itself is oxidized! Do not get them confused!
    We learned in lessons like Electron configuration 2 and Forming ions that some atoms tend to gain or lose electrons to obtain a full outer shell of electrons. If you know this already, then you already have a good idea of what elements are reducing and oxidizing agents!
    • Most reducing agents are metals because metals, as we know, normally lose electrons in a reaction, becoming positive ions while the other species in the reaction gains those electrons.
    • Most oxidizing agents are electronegative non-metals that, as we know, tend to gain electrons in reactions by ‘removing them’ from other elements. Fluorine is an excellent oxidizing agent, as is oxygen (that’s why it’s called oxidation!) and chlorine – these are the most electronegative elements in the table.

  • The activity series is a basic list that shows the relative tendencies of metal species to be oxidised – that is, to go from the elemental metal to metal ion plus electron(s). This has a correlation with electronegativity, as group 1 and 2 metals are amongst the highest in the series, like calcium, potassium and lithium. Near the bottom are the heavier transition metals like mercury, platinum and gold.
    The activity series is based on standard electrode potential (AKA standard reduction potential) which we will use in calculations in the later lesson Calculating cell potential (Voltaic cells).