Chemistry · MU Sofia

5.1 Chemical redox processes

Explains redox reactions as electron transfer, defining oxidation, reduction, reducer and oxidant, and how oxidation numbers track electron flow. Works through copper oxide decomposition as half-reactions and covers standard electrode potential with the standard hydrogen electrode. Also describes galvanic and electrolytic cells, electrolysis of molten and aqueous sodium chloride, and the cell potential equation linked to Gibbs free energy.

Inside: 10 sections · 7 labelled pictures · 14 flashcards · 6 min

Topic 16 of 42 in the MU Sofia syllabus for Chemistry (Dental Medicine, Year 1). The free AceMed note covers it in 10 sections, about 682 words (6 min of reading), with 14 flashcards.

In this note

  1. Introduction
  2. Electron transfer. Half-reactions.
  3. Electrochemical redox potential
  4. Standard electrode potential
  5. Chemical current sources.
  6. Electrolysis
  7. High Yield
  8. Cell Potential Equation
  9. Relation with Thermodynamics
  10. Link between Gibbs free energy and potential:

Introduction

Chemical redox processes

  • Redox stands for reduction-oxidation. These are chemical reactions where electrons are transferred between different atoms or ions.
  • Key Concepts:
  • Oxidation: The loss of electrons by an element or ion.
  • Reduction: The gain of electrons by an element or ion.
  • Simultaneous Occurrence: Both oxidation and reduction happen concurrently in redox processes.

Electron transfer. Half-reactions.

  • Electron Exchange: The Core of Redox
  • Redox reactions, short for reduction-oxidation, are all about electron exchange. Remember these two key terms:
  • Oxidation: When an atom or ion loses electrons.
  • Reduction: When an atom or ion gains electrons.
  • Who's Who in the Electron Transfer Drama?
  • Reducer: The electron donor (loses electrons) and gets oxidized.
  • Oxidant: The electron acceptor (gains electrons) and gets reduced.
  • Why Do Electrons Move?
  • Electronegativity plays a role. Atoms with higher electronegativity are better at attracting and holding electrons. So, in a reaction:
  • The more electronegative atom accepts electrons and reduces its oxidation state (becomes more negative).
  • The atom that loses electrons becomes a positive ion and increases its oxidation state (becomes more positive).
  • Keeping Track: Oxidation Numbers
  • Imagine a conditional charge on each atom if all bonds were ionic. This is called the oxidation number. It helps track electron flow:
  • Oxidizers: Reduce their oxidation state (become more negative).
  • Reducers: Increase their oxidation state (become more positive).
  • Example: Decomposing Copper Oxide (CuO)
  • The reaction: CuO→Cu + O₂
  • Can be broken down into half-reactions:
  • Oxidation (Cu loses electrons): Cu⁰ > Cu²⁺ + 2e⁻ (oxidation state increases)
  • Reduction (O gains electrons): ½ O₂ + 2e⁻ > O²⁻ (oxidation state decreases)
  • What's in a Half-Reaction?
  • A half-reaction focuses on either the oxidation or the reduction happening in the full reaction.
  • It shows the specific atom or ion involved in that specific act (oxidation or reduction).
  • It balances the charge of the species involved (atoms or ions) using electrons (e⁻).
  • Balancing the Act: Electron Equality
  • The key rule: the number of electrons donated (by reducers) must equal the number received (by oxidants). This helps balance redox equations.
  • Examples of Redox Reactions:
  • Rusting of iron
  • Combustion of wood
  • Photosynthesis

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