half-reactions

Examples of half-reactions in the following topics:

• Balancing Redox Equations

  • This reaction is split into two half-reactions, one that involves oxidation and one that involves reduction.
  • To do this, multiply the oxidation half-reaction by 3 and the reduction half-reaction by 2, so that each half-reaction has 6e-.
  • In acidic media, H+ ions and water are added to half-reactions to balance the overall reaction.
  • We need to multiply the reduction half-reaction by 5 and the oxidation half-reaction by 2.
  • In basic media, OH− ions and water are added to half reactions to balance the overall reaction.

• Balancing Redox Equations

  • Every balanced redox reaction is composed of two half-reactions: the oxidation half-reaction, and the reduction half-reaction.
  • We can split this reaction into two half-reactions.
  • The oxidation half-reaction looks as follows:
  • First, we need to split this reaction into its two half-reactions.
  • For this reaction, we can multiply the first half-reaction by 3:

• Predicting Spontaneous Direction of a Redox Reaction

  • This means that Li would be written as the reduction half-reaction when compared to any other element in this table.
  • On the other hand, Fe would be written as the oxidation half-reaction when compared to any other element on this table.
  • The relative reactivities of different half-reactions can be compared to predict the direction of electron flow.
  • Half-reaction equations can be combined if one is reversed to oxidation in a manner that cancels out the electrons.
  • Predict the direction of electron flow in a redox reaction given the reduction potentials of the two half-reactions

• Half-Life

  • The half-life of a reaction is the amount of time it takes for the concentration of a reactant to decrease to one-half of its initial value.
  • To find the half-life of the reaction, we would simply plug 5.00 s-1 in for k:
  • Thus the half-life of a second-order reaction, unlike the half-life for a first-order reaction, does depend upon the initial concentration of A.
  • As initial concentration increases, the half-life for the reaction gets longer and longer.
  • The half-life of a reaction is the amount of time it takes for it to become half its quantity.

• Zero-Order Reactions

  • Unlike the other orders of reaction, a zero-order reaction has a rate that is independent of the concentration of the reactant(s).
  • Zero-order reactions are typically found when a material that is required for the reaction to proceed, such as a surface or a catalyst, is saturated by the reactants.
  • This is the integrated rate law for a zero-order reaction.
  • The half-life of a reaction describes the time needed for half of the reactant(s) to be depleted, which is the same as the half-life involved in nuclear decay, a first-order reaction.
  • For a zero-order reaction, the half-life is given by:

• Carbon-Carbon Bond Formation

  • The number of generally useful and well tested reactions for effecting carbon-carbon bond formation, ideally in a regio and stereospecific fashion, is relatively small, compared with reactions used to modify functional groups.
  • It should also be noted that more than half these reactions involve carbonyl reactants.
  • Moreover, in the half century since Woodward's reserpine synthesis was carried out, this "toolkit" has been expanded to include an assortment of new, tolerant and selective carbon-carbon bond forming reactions.
  • Many of the essential reactions used here were not known at the time of Corey's work.
  • The preparation of isogeranic acid (top equation) makes use of a transition metal coupling reaction.

• The Rate Law

  • The rate law for a chemical reaction relates the reaction rate with the concentrations or partial pressures of the reactants.
  • For the general reaction$aA + bB \rightarrow C$ with no intermediate steps in its reaction mechanism, meaning that it is an elementary reaction, the rate law is given by:
  • A smaller rate constant indicates a slower reaction, while a larger rate constant indicates a faster reaction.
  • What is the reaction order?
  • The reaction is first-order in hydrogen, one-half-order in bromine, and $\frac{3}{2}$-order overall.

• Electrochemical Cell Notation

  • A positive cell potential indicates that the reaction proceeds spontaneously in the direction in which the reaction is written.
  • The anode half-cell is described first; the cathode half-cell follows.
  • Within a given half-cell, the reactants are specified first and the products last.
  • A double vertical line ( || ) represents a salt bridge or porous membrane separating the individual half-cells.
  • A typical arrangement of half-cells linked to form a galvanic cell.

• Electrocyclic Reactions

  • An electrocyclic reaction is the concerted cyclization of a conjugated π-electron system by converting one π-bond to a ring forming σ-bond.
  • The reverse reaction may be called electrocyclic ring opening.
  • Once again, the number of curved arrows that describe the bond reorganization is half the total number of electrons involved in the process.
  • The sterospecificity of this reaction is demonstrated by closure of the isomeric trans,cis,cis-triene to trans-5,6-dimethyl-1,3-cyclohexadiene, as noted in the second example.
  • This mode of reaction is favored by relief of ring strain, and the reverse ring closure (light blue arrows) is not normally observed.

• Alkylidene Reactions

  • Many structural features of a metathesis catalyst may be changed and adjusted to suit the type of reaction desired.
  • A few of the resulting catalytic complexes are shown in the bottom half of the second illustration.
  • Diagrams 5, 6, 7, and 8 above show examples of some applications of metathesis reactions.
  • RCM reactions # 1 and 2 show two ring closures, and reaction # 3 demonstrates a double RCM in which an alkyne serves to transfer the metal carbene to a new location.
  • A carbene olefination may serve to terminate metathesis, as in reaction # 4, but is not catalytic.