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ELECTROCHEMISTRY

ELECTROCHEMISTRY. Presentation by: P.K. CHOURASIA K.V MANDLA, Jabalpur Region. OBJECTIVES

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ELECTROCHEMISTRY

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  1. ELECTROCHEMISTRY Presentation by: P.K. CHOURASIA K.V MANDLA, Jabalpur Region

  2. OBJECTIVES • Electrochemistry is the study of production of electricity from the energy released during a spontaneous reaction and use of electrical energy to bring about non spontaneous chemical reaction. • How Commercially generate electricity. • It provides many ideas for protection of metal surface from corrosion.

  3. Redox Reactions • Redox reaction are those involving the oxidation and reduction of species. • OIL – Oxidation Is Loss of electrons. • RIG – Reduction Is Gain of electrons. • Oxidation and reductionmust occur together. • They cannot exist alone. Chapter 18

  4. Redox Reactions • Reduction Half-Reaction: • Cu2+(aq) + 2 e– Cu(s) • The Cu2+ gains two electrons to form copper. Chapter 18

  5. Redox Reactions • Oxidation Half-Reaction: Zn(s)  Zn2+(aq) + 2 e–. • The Zn loses two electrons to form Zn2+. Chapter 18

  6. Electrochemical Cells • Electrodes: are usually metal strips/wires connected by an electrically conducting wire. • Salt Bridge: is a U-shaped tube that contains a gel permeated with a solution of an inert electrolyte. • Anode: is the electrode where oxidation takes place. • Cathode: is the electrode where reduction takes place. Chapter 18

  7. Electrochemical Cells • Convention for expressing the cell: • Anode Half-Cell || Cathode Half-Cell • Electrode | Anode Soln || Cathode Soln | Electrode • Zn(s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu(s) • Electrons flow from anode to cathode. Anode is placed on left by convention. Chapter 18

  8. Electrochemical Cells Overall: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) Left Electrode Chapter 18 Right Electrode

  9. Electrochemical Cells • Electrode potential: The tendecy of an element when it is placed in contact with its ions, to become positively or negatively charged by losing and gaining electrons. • The standard potential of any galvanic cell is the subtraction of the standard half-cell potentials for the oxidation and reduction half-cells. E°cell = E°reduction - E°oxidation • Standard half-cell potentials are always quoted as a reduction process. The sign must be changed for the oxidation process. Chapter 18

  10. Electrochemical Cells • The standard half-cell potentials are determined from the differencebetween two electrodes. • The reference point is called the standard hydrogen electrode (S.H.E.) and consists of a platinum electrode in contact with H2 gas (1 atm) and aqueous H+ ions (1 M). • The standard hydrogen electrode is assigned an arbitrary value of exactly 0.00 V. Chapter 18

  11. Electrochemical Cells Chapter 18

  12. Electrochemical Cells Chapter 18

  13. Electrochemical Cells • When selecting two half-cell reactions the more negative value will form the oxidation half-cell. • Consider the reaction between zinc and silver: • Ag+(aq) + e– Ag(s) E° = 0.80 V • Zn2+(aq) + 2 e– Zn(s) E° = – 0.76 V • Therefore, zinc forms the oxidation half-cell: • Zn(s)  Zn2+(aq) + 2 e–E° = – (–0.76 V) Chapter 18

  14. Electrochemical Cells • Q1. How does salt bridge complete the circuit? • Q2. What is electrode potential? • Q3. In electrochemical cell, why is anode negative and cathode positive? • Q4. Which factors are required to form an efficient cell? Chapter 18

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