Redox Reactions

Redox Reactions

Introduction

  • Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between two substances.
  • One substance undergoes reduction (gains electrons), while the other undergoes oxidation (loses electrons).

Recognising a Redox Reaction

  • Oxidation can be remembered as ‘OIL RIG’ - Oxidation Is Loss, Reduction Is Gain.
  • Oxidation states can be used to determine if a redox reaction has occurred; if an element’s oxidation state changes during a reaction, it is a redox reaction.
  • The sum of the oxidation states in a neutral compound is zero; in an ion, it equals the ion’s charge.

Balancing Redox Reactions

  • Half-reactions are often used to help balance redox equations; one half-reaction will correspond to the oxidation process and the other to the reduction process.
  • The same number of electrons should be gained in the reduction half-reaction as are lost in the oxidation half-reaction.
  • If in an acidic solution, H+ ions are added to balance any difference in hydrogen, then water (H2O) is added to balance oxygen. In a basic solution, OH- ions are used instead.

Electrochemical Cells

  • Electrochemical cells, also known as galvanic or voltaic cells, involve redox reactions. They generate an electrical current from a spontaneous chemical reaction.
  • The substance being oxidised (losing electrons) is the anode, and the substance being reduced (gaining electrons) is the cathode.
  • Positive ions move towards the cathode and negative ions towards the anode.
  • Electrons move through an external circuit from the anode (negative electrode) to the cathode (positive electrode).
  • The potential difference between the anode and cathode is known as the cell potential or electromotive force (emf).

Electrolysis

  • Electrolysis is the process of driving a redox reaction in the non-spontaneous direction by using an external electrical source (opposite of electrochemical cells).
  • During electrolysis, oxidation occurs at the positive electrode (anode), and reduction occurs at the negative electrode (cathode).

Applications

  • Redox reactions are crucial in numerous applications including batteries, fuel cells, metal refining, and the synthesis of many chemicals.
  • They are key processes in natural systems such as photosynthesis, respiration, and the cycling of elements within ecosystems.
  • For instance, in industrial production of halogens from sea water: chlorine (Cl2) is formed at the anode while hydrogen (H2) and hydroxide ions (OH-) are formed at the cathode during electrolysis.

To excel it is important to understand not only how to balance redox equations and how electrochemical cells work, but also to understand where these reactions occur in real-life contexts. Don’t focus solely on drilling practice problems, make sure to deepen your understanding of the principles underlying redox reactions.