Unlock the Power of Chemical Cells
Dive into the world of chemical cells and discover how they convert energy. Learn about electrodes, electrolytes, and their applications in modern technology. Master the principles of electrochemistry with our engaging lessons.

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Now Playing:Chemical cells – Example 0a
Intros
  1. What is an electrochemical cell?
  2. Electricity and chemistry
  3. Key features of a chemical cell.
Examples
  1. Apply the idea of cell potential to predict the flow of electrons in chemical cells.
    1. Two solutions are connected to form an electrochemical cell:
      • Ag metal in AgNO3 solution
      • Sn metal in SnCl2 solution

        Draw a chemical cell showing the cathode, anode and the direction of electron flow in this electrochemical cell. Which metal is being oxidised and which is being reduced?

      • Two solutions are connected to form an electrochemical cell:
        • Ni metal in NiCl2 solution
        • Zn metal in Zn(NO3)2 solution

          Draw a chemical cell showing the cathode, anode and the direction of electron flow in this electrochemical cell.

Introduction to electrochemistry
Notes

In this lesson, we will learn:

  • To recall the basic function of an electrochemical cell.
  • To explain the role of the key features in an electrochemical cell in redox processes.

Notes:

  • We learned in Introduction to electrochemistry that electrochemistry is about chemical potential energy (the energy ‘locked in’ the bonds of all chemical substances) being converted into electrical energy (energy from the movement of electrons).

  • Converting chemical potential to electrical energy requires electrons to move FROM one reactant TO the other reactant – which is where we get the processes of reduction and oxidation from, and the ‘redox’ phrase so common in electrochemistry. Redox processes that produce electricity occur in electrochemical cells, or chemical cells, which work just like electrical circuits (like in physics) but they also use ions in solution to carry charge, not just electrical wiring.
    The ‘battery’ of an electrochemical cell is the energy in the bonds of the reactants that are getting broken up!

  • The two most important features of a chemical cell are the two electrodes where redox happens:
    • The anode is the positively charged electrode where oxidation happens.
      • It attracts the negatively charged ions in the cell.
      • The negatively charged ions lose electrons (they are oxidized) at the anode, and the electrons flow through to the cathode.
    • The cathode is the negatively charged electrode where reduction happens.
      • It attracts the positively charged ions in the cell.
      • The positively charged ions gain electrons (they are reduced) at the anode by the electrons that have been flowing to the anode from the cathode.
    • Chemical cells work by electrons flowing from the anode to the cathode.
    • Electrode is just a general term for either cathode or anode, a conducting material where the redox processes happen. They must conduct electricity to work, so if the reactants are not in the solid state (see electrolysis in a later lesson) they would be made of a metal like platinum, or graphite.

  • Other features in a chemical cell, or redox cell, are:
    • The salt bridge, which is used to connect the two half-cells. It contains a solution of highly soluble spectator ions (e.g. KNO3 or NaNO3) to carry current without taking part in the reaction.
    • Wire will connect the two electrodes and allow electrons to flow from the anode to the cathode. This completes the ‘circuit’ (just like a circuit in physics!) for a one-way flow of electrons.
    • A voltmeter can be added to measure a flow of electrons.

  • The real working of a redox cell can be split into the two ‘half-reactions’, the two separate redox processes of reduction and oxidation.
    For example with Ni metal (in NiCl2 solution) reacting with Pb metal (in Pb(NO3)2 solution), the two half-reactions for the metal atoms are:
    • Ni2+ + 2e- \rightleftharpoons Ni
    • Pb2+ + 2e- \rightleftharpoons Pb
    Redox works by one substance being oxidized, losing an electron(s), and one substance being reduced, gaining an electron(s) – the electron(s) that was lost by what got oxidized! So which substance will do which process?
    Together, chemists have studied many redox experiments and with their combined results, a table of ‘standard reduction potentials’ (how easily can this substance be reduced?) was formed – all good chemistry textbooks will have this table in the back. For our example, Pb is more likely to be reduced than Ni,1 so lead will be reduced and nickel will be oxidized in this cell.
    • Because lead has a greater tendency to be reduced, when the half-cells are connected via wiring and a salt bridge, electrons will flow toward the lead electrode. This upsets the equilibrium between lead metal and Pb2+ ions:

      Pb2+ + 2e-\enspace \rightleftharpoons \enspace Pb

      The increase in electrons available will shift the equilibrium to the right to counter the imbalance, which produces lead metal in a reduction reaction. This makes the lead electrode the cathode of this cell.
    • Nickel has a smaller tendency to reduce than lead does, so it will do the oxidation when these two half-cells are connected up. When connected, electrons will flow away from the nickel electrode and toward the lead electrode through the wire. This disturbs the nickel equilibrium we wrote:

      Ni2+ + 2e-\enspace \rightleftharpoons \enspace Ni

      The decreasing amount of electrons will shift the equilibrium to the left to counter to produce more electrons. In doing this, it converts nickel metal into Ni2+ ions and e- in an oxidation reaction. This makes the nickel electrode the anode of this cell.


      1 Source for data: ATKINS, P. W., & DE PAULA, J. (2006). Atkins' Physical chemistry. Oxford, Oxford University Press.
Concept

Introduction to Chemical Cells

Welcome to our exploration of chemical cells, a fascinating topic in electrochemistry. Our introduction video provides an engaging overview of this subject, setting the stage for a deeper understanding. Chemical cells, also known as electrochemical cells, are devices that convert chemical energy into electrical energy or vice versa. This lesson will delve into the basic function and key features of these remarkable systems. We'll examine how different materials interact to produce electricity, the role of electrolytes, and the importance of electrodes. You'll discover how chemical cells power everything from your smartphone to electric vehicles. By understanding the principles behind chemical cells, you'll gain insight into a technology that's crucial to our modern world. Whether you're a student, hobbyist, or just curious about science, this introduction to chemical cells will provide a solid foundation for further learning.

FAQs
  1. What is a chemical cell or galvanic cell?

    A chemical cell, also known as a galvanic cell, is an electrochemical device that converts chemical energy into electrical energy. It consists of two half-cells, each containing an electrode and an electrolyte, connected by a salt bridge or porous barrier. The spontaneous redox reaction between the two half-cells generates an electric current.

  2. What are the two types of chemical cells?

    The two main types of chemical cells are:

    • Galvanic (or Voltaic) cells: These spontaneously produce electricity from chemical reactions.
    • Electrolytic cells: These use electrical energy to drive non-spontaneous chemical reactions.
  3. What is the difference between electrolysis and chemical cells?

    Electrolysis uses electrical energy to drive non-spontaneous chemical reactions, while chemical cells (specifically galvanic cells) spontaneously produce electrical energy from chemical reactions. Electrolysis is a process that occurs in electrolytic cells, whereas chemical cells typically refer to galvanic cells that generate electricity.

  4. What is the principle of a chemical cell?

    The principle of a chemical cell is based on the separation of oxidation and reduction half-reactions. This separation allows electrons to flow through an external circuit, generating an electric current. The cell utilizes the difference in reduction potentials between two half-cells to produce electrical energy from chemical reactions.

  5. How do standard reduction potentials help in predicting cell behavior?

    Standard reduction potentials help predict cell behavior by indicating the relative tendency of species to be reduced. In a chemical cell, the half-reaction with the more positive reduction potential will be the cathode (reduction), while the one with the more negative potential will be the anode (oxidation). This information allows us to determine the direction of electron flow and calculate the overall cell potential.

Prerequisites

When delving into the fascinating world of chemical cells, it's crucial to have a solid foundation in certain prerequisite topics. These fundamental concepts not only enhance your understanding but also provide the necessary context for grasping the intricacies of chemical cells. Two key areas that are particularly relevant are electric potential and electric potential energy and calculating cell potential in voltaic cells.

Understanding electric potential and electric potential energy is fundamental to comprehending how chemical cells function. This concept is closely related to chemical potential energy, which plays a vital role in the operation of chemical cells. In a chemical cell, the difference in electric potential between two electrodes drives the flow of electrons, generating an electric current. By grasping the principles of electric potential, you'll be better equipped to analyze the energy transformations occurring within chemical cells.

Moreover, the ability to calculate cell potential in voltaic cells is an essential skill when studying chemical cells. This knowledge allows you to predict the direction of electron flow and the spontaneity of redox reactions. Understanding how concentration affects cell potential is particularly important, as it directly influences the efficiency and performance of chemical cells. By mastering this prerequisite topic, you'll be able to analyze and optimize various types of chemical cells, including batteries and fuel cells.

The interplay between these prerequisite topics and chemical cells is profound. For instance, the concept of electric potential energy helps explain why certain chemical reactions occur spontaneously in cells, while others require an external energy input. Similarly, understanding cell potential calculations enables you to predict the voltage output of different cell configurations, which is crucial for designing and improving energy storage devices.

By thoroughly grasping these prerequisite topics, you'll develop a more intuitive understanding of chemical cells. You'll be able to explain phenomena such as why certain materials are chosen for electrodes, how concentration gradients affect cell performance, and why some chemical cells are more efficient than others. This knowledge is not only academically valuable but also has practical applications in fields like renewable energy, materials science, and electrochemistry.

In conclusion, investing time in mastering these prerequisite topics will significantly enhance your comprehension of chemical cells. The concepts of electric potential and electric potential energy, along with the ability to calculate and analyze cell potentials, form the bedrock upon which a deeper understanding of chemical cells is built. By establishing this strong foundation, you'll be well-prepared to explore the complexities of chemical cells and their wide-ranging applications in modern technology and scientific research.