redox

• oxidation- gain of oxygen or loss of hydrogen
• reduction- loss of oxygen or gain of hydrogen
• If the oxidation number increases or gets more

  positive, the species has been oxidized

  If the oxidation number decreases (or gets more

  negative), the species has been reduced
• An oxidizing agent is the species that is reduced. (Also known as oxidisers)

  A reducing agent is the species that gets oxidised. (Also known as fuels)
• A single displacement reaction is one where

  a metal acts as a reducing agent for another

  metal ion.
• A titration is a process to determine the

  concentration of a solution.

  Just like in an acid-base titration, a standard solution

  is added drop by drop to a known volume of the

  analyte and then a change in colour indicates the

  equivalence point.
• Redox titration: redox reaction between oxidising agnet and reducing agent. Electrons are transferred from reducing agent to oxidising agent.
• Electrolytic cells use an external source of electrical

  energy to bring about a redox reaction that would

  otherwise be non-spontaneous
• Electrolysis is often used to break down compounds
• The reactant in the process of electrolysis is

  present in the electrolyte (solution that has ions)
• • The power source pushes electrons towards the

  negative electrode, where they enter the electrolyte, which is the cathode.

  • Electrons are released at the positive terminal, the anode, and returned to the source.
  • The current is passed through the electrolyte, not by electrons, but by the ions as they are mobile and migrate to the electrodes.
  • The chemical reaction occurring at each electrode remove the ions from the solution and so enable the process to continue.
• The ions in the electrolyte migrate to the electrodes by attraction of opposite charges. So positive ions

  (cations) are attracted to the negative electrode, the

  cathode, while negative ions (anions) are attracted to the positive electrode, the anode.
• Voltaic cell

  • anode- oxidation occurs here (negative)
  • cathode- reduction occurs here (positive)

  Electrolytic cell

  • anode- oxidation occurs here (positive)
  • cathode- reduction occurs here (negative)
• There are two types:

  1. Voltaic Cells, which use spontaneous redox reactions to generate electricity.
  2. Electrolytic Cells, which use electricity to force non-spontaneous redox reactions.
• Voltaic cells use the fact that electrons flow

  spontaneously from reducing agents (more easily

  oxidized) to oxidizing agents (more easily reduced).
• Always put the stronger reducing agent on the left!
• • Electrons flow from anode to cathode through the external circuit.
  • Anions migrate from cathode to anode through the salt bridge.
  • Cations migrate from anode to cathode through the salt bridge.
• Voltaic (Galvanic) cells:

  • Voltaic cells convert energy from spontaneous,

  exothermic chemical processes to electrical energy.

  • Oxidation occurs at the anode (negative electrode) and reduction occurs at the cathode (positive electrode) in a voltaic cell.

  Electrolytic Cells:

  • Electrolytic cells convert electrical energy to chemical energy, by bringing about non–spontaneous process.
  • Oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative electrode) in an electrolytic cell.
• • A voltaic cell generates a potential difference known as the electromotive force (EMF)
  • As electrons tend to flow from the half-cell with the more negative potential to the half-cell with the more positive potential, the potential generated is called the cell potential or electrode potential and is given the symbol E.
• • Platinum is used as the conducting metal in the

  electrode because it is fairly inert and will not ionize.

  • It also acts as a catalyst for the reaction of proton reduction
• spontaneity of a reaction

  ∆GΘ = –nFEΘ
• • J = C x V
  • The negative sign in the equation indicates that E and ∆G have opposite signs
  • When Eθcell is positive, ∆GΘ is negative, reaction is spontaneous
  • When Eθcell is negative, ∆GΘ is positive, reaction is non-spontaneous
  • When Eθcell is 0, ∆GΘ is 0, reaction is at equilbrium
• Because of the direct relationship between Eθcell and

  ∆GΘ, the more positive the value of Eθcell the more

  energetically favorable is the reaction.
• • Metals with low EΘ values (mostly negative) are therefore the strongest reducing agents
  • Likewise, a non-metal is able to oxidise the ions of another non-metal that has a lower EΘ value.
  • Non-metals with high EΘ values are therefore the strongest oxidising agents.
• Factors Affecting Amount of Product:

  • The amounts of products at the electrodes depend on the quantity of electric charge passed through the cell.
  • Quantity of charge is quantified as:

  =>Q = I t

  Q = charge; unit coulomb (C)

  I = current; measured in amperes (A)

  t = time; measured in seconds (s)

  The charge carried by one mole of electrons, known

  as Faraday (F) is 96,500C
• Formuals:

  
• Electroplating is the process of using electrolysis to

  deposit a layer of a metal on top of another metal or

  other conductive object.
• Summary of cells:

  
• The electrolysis of concentrated sodium chloride produces chlorine gas and hydrogen gas via the oxidation of chloride ions and the reduction of hydrogen atoms in the water. In dilute solutions of sodium chloride, some of the water is also oxidised to produce oxygen gas, and in very dilute solutions of sodium chloride, only oxygen gas will be produced.
• Electroplating involves the use of an electrolysis reaction, driving a non-spontaneous electrochemical reaction through the use of a power supply. The process allows for a metal cation to be transferred within a solution onto the surface of a cathode, where it undergoes oxidation to form a solid metal plating.
• • In dilute solutions of sodium chloride, the water will be oxidised leading to the production of oxygen gas at the anode.In concentrated solutions, the chloride ions will be preferentially oxidised leading to the formation of chlorine gas
• • Changes in the concentration of the salt solution have a large impact on which species is oxidised during electrolysis. In dilute solutions, water molecules are oxidised to produce oxygen gas. Whereas, in concentrated solutions of sodium chloride, chloride ions are oxidised to produce chlorine gas. So to summarise, if the concentrated solution contains a halide ion (Cl−, Br− or I−) then the halogen (Cl2, Br2 or I2) is given off at the anode. However, if the halide ions are present in a very low concentration, oxygen gas will be given off at the anode instead.
• • It is essential to use electrodes which are highly conductive, have high melting points, and yet are not preferentially oxidised or reduced over the species in solution.
• • If a reactive electrode like copper is used instead of an inert electrode then the solid copper electrode will be preferentially oxidised at the anode and the electrode will produce copper(II) ions in solution instead of generating oxygen gas.
  • This means that the electrolytes pH would remain constant, as would the concentration of copper(II) ions in solution. Copper would still be plated onto the anode and the copper cathode would be slowly eroded.
• • In order to quantify whether one element will more readily undergo reduction or oxidation compared to another element, we need a standard value around which to base our measurements. This comes in the form of the standard hydrogen electrode  (SHE).
  • Platinum is used as it is inert and will not preferentially undergo reduction over hydrogen gas.