Oxidation and Reduction Chapter 20

1
Oxidation and ReductionChapter 20
2
Types of Chemical Reactions

• Type I ions or molecules react with no apparent
  change in the electronic structure of the
  particles.
• Type II ions or atoms undergo changes of
  electronic structure. Electrons may be
  transferred from one particle to another. On the
  other hand, the sharing of the electrons may be
  somewhat changed.
• Type II reactions involving electron changes are
  called oxidation-reduction reactions.
• It is these "redox" reactions which we will now
  discuss.
• Before we indicate what oxidation-reduction
  reactions are, we will briefly indicate what they
  are not.

3
What Redox is NOT

• In the BaSO4 reaction in Table 26-1, the
  substances are all ionic.
• Since there is no change in the charge of these
  ions in the reaction, there are no electron
  changes.
• This reaction is not an oxidation-reduction
  reaction.
• The production of a (BaS04) is nearly always a
  result of a non-redox reaction.
• Most acid-base reactions are also the non-redox
  type.
• Since nearly every other kind of reaction is an
  oxidation-reduction reaction, redox reactions are
  important in the laboratory.
• They are also important in life processes and in
  industry.

4
Oxidation

• The term oxidation was first applied to the
  combining of oxygen with other elements.
• There were many known instances of this behavior
• Iron rusts
• Carbon burns
• In rusting, oxygen combines slowly with iron to
  form Fe2O3.
• In burning, oxygen unites rapidly with carbon to
  form CO2.
• Observation of these reactions gave rise to the
  terms "slow" and "rapid" oxidation.
• Chemists recognize, however, that other
  nonmetallic elements unite with substances in a
  manner similar to that of oxygen.
• Hydrogen, antimony, and sodium all burn in
  chlorine, and iron will burn in fluorine.
• Since these reactions were similar, chemists
  formed a more general definition of oxidation
• Electrons were removed from each free element by
  the reactants O2 or Cl2.
• Thus oxidation is defined as the process by which
  electrons are apparently removed from an atom or
  ion.

5
Reduction

• A reduction reaction was originally limited to
  the type of reaction in which ores were "reduced"
  from their oxides.
• Iron oxide was "reduced" to iron by carbon
  monoxide.
• Copper(II) oxide could be "reduced" to copper by
  hydrogen.
• In these reactions, oxygen is removed, and the
  free element is produced.
• The free element can be produced in other ways
• An iron nail dropped into a copper(II) sulfate
  solution causes a reaction which produces free
  copper.
• An electric current passing through molten sodium
  chloride produces free sodium.
• The similarity between oxidation and reduction
  reactions led chemists to formulate a more
  generalized definition of reduction.
• By definition, reduction is the process by which
  electrons are apparently added to atoms or ions.

6
OIL RIGthe Texas Definition

• Oxidation is Loss (of electrons), Reduction is
  Gain (of electrons)

7
Oxidizing and Reducing Agents

• In an oxidation-reduction reaction, electrons are
  transferred.
• All the electrons exchanged in an
  oxidation-reduction reaction must be accounted
  for.
• It seems reasonable, therefore, that both
  oxidation and reduction must occur at the same
  time in a reaction.
• Electrons are lost and gained at the same time
  and the number lost must equal the number
  gained.
• The substance in the reaction which gives up
  electrons is called the reducing agent. The
  reducing agent contains the atoms which are
  oxidized (the atoms which lose electrons).
• Zinc is a good example of a reducing agent. It is
  oxidized to the zinc ion, Zn2
• The substance in the reaction which gains
  electrons is called the oxidizing agent. It
  contains the atoms which are reduced (the atoms
  which gain electrons).
• Dichromate ion, Cr2072-, is a good example of an
  oxidizing agent. It is reduced to the chromium
  ion, Cr3
• If a substance gives up electrons readily, it is
  said to be a strong reducing agent. Its oxidized
  form, however, is normally a poor oxidizing
  agent.
• If a substance gains electrons readily, it is
  said to be a strong oxidizing agent. Its reduced
  form is a weak reducing agent.

8
Redox of nails and copper
9
Oxidation Numbers

• How is it possible to determine whether an
  oxidation-reduction reaction has taken place?
• We do so by determining whether any electron
  shifts have taken place during the reaction.
• To indicate electron changes, we look at the
  oxidation numbers of the atoms in the reaction.
• The oxidation number is the charge an atom
  appears to have when we assign a certain number
  of electrons to given atoms or ions.
• Any change of oxidation numbers in the course of
  a reaction indicates an oxidation-reduction
  reaction has taken place.
• Oxidation numbers are assigned according to the
  apparent charge of the element (aka, valence!)
• For example, suppose iron, as a reactant in a
  reaction, has an oxidation number of 2.
• If iron appears as a product with an oxidation
  number other than 2, say 3, or 0, then a redox
  reaction has taken place.

10
Determining Oxidation Numbers

• For all compounds, whether covalent, polar
  covalent, or ionic, we treat as ionic for
  counting electrons and for oxidation-reduction
  reactions.
• Rule 1 Sum of the oxidation numbers of all the
  atoms in the chemical species equals the charge
  on the species.Neutral compounds Sum of
  oxidation numbers 0Ionic species Sum of
  oxidation numbers charge of the ion
• Rule 2 In Binary Compounds, the more
  Electronegative (EN) element is assigned to have
  a negative oxidation number. (See EN trends.)
• Rule 3 Atoms may have only certain oxidation
  numbers. The range is
• Maximum oxidation number possible Group
  number.Minimum oxidation number possible
  (Group number - 8) (this number will be negative)

11
Determining Oxidation Numbers (Cont)

• Atoms which will have known oxidation numbers
  are
• Atoms as Elements Ex. H2, O2, P4, FeOxidation
  number 0
• Monoatomic IonsCations Ex. Na, Al3 (main
  group metals)Oxidation number Group
  NumberAnions Cl-, O2- Oxidation number Group
  Number - 8
• HydrogenCombined with Nonmetals Ex. NH3, H2O,
  HClOxidation number 1Combined with Metals
  Ex. NaH, CaH2 (hydrides)Oxidation number -1
• Oxygen (Unless O22-, peroxide)Oxidation number
  -2

12
CO (Sum will equal 0 since it is a neutral
molecule)

• O will have a -2 ox. number.
• 1 C 1 O 0(C?) (-2) 0C? 2
• Oxidation number of C in CO is 2Oxidation
  number of O in CO is -2 (known)
• Check ox. number to see if it falls within
  range2 is in between the maximum value of C,
  4, (Gr) and the minimum value of C, - 4, (Gr -
  8).So okay.

13
Cr2O72- (Sum of all oxidation numbers will equal
-2 since it is an ion.)

• 2 Cr 7 O -22(Cr?) 7(-2) -22(Cr?)
  (-14) -22(Cr?) 12Cr? 6
• Oxidation number of each Cr in Cr2O72- is
  6Oxidation number of each O in Cr2O72- is -2
  (known)
• Check ox. number to see if it falls within
  range6 is the maximum value that Cr can have
  (Gr). So okay.

14
CS2 (Sum will equal 0 since it is a neutral
molecule)

• C will have the positive oxidation number since
  it is less EN than SS will have a -2 charge
  since it is Gr 6, (6 - 8 -2)
• C 2 S 0(C?) 2 (-2) 0(C?) (-4) 0C?
  4
• Oxidation number of C in CS2 is 4Oxidation
  number of each S in CS2 is -2 (known)
• Check ox. number to see if it falls within
  range4 is the maximum value that C can have,
  (Gr). So okay.

15
HNO3(aq)  H3AsO3(aq)  ?  NO(g)  H3AsO4(aq)  H2O
(l)

• Step 1  Try to balance the atoms by inspection.
• The H and O atoms are difficult to balance in
  this equation. You might arrive at the correct
  balanced equation using a trial and error
  technique, but if you do not discover the correct
  coefficients fairly quickly, proceed to Step 2.
• Step 2  Is the reaction redox?
• The N atoms change from 5 to 2, so they are
  reduced. This information is enough to tell us
  that the reaction is redox. (The As atoms, which
  change from 3 to 5, are oxidized.)
• Step 3  Determine the net increase in oxidation
  number for the element that is oxidized and the
  net decrease in oxidation number for the element
  that is reduced.
• As  3 to 5     Net Change 2
• N  5 to 2      Net Change -3
• Step 4  Determine a ratio of oxidized to
  reduced atoms that would yield a net increase in
  oxidation number equal to the net decrease in
  oxidation number.
• As atoms would yield a net increase in oxidation
  number of 6. (Six electrons would be lost by
  three arsenic atoms.) 2 N atoms would yield a net
  decrease of -6. (Two nitrogen atoms would gain
  six electrons.) Thus the ratio of As atoms to N
  atoms is 32.
• Step 5 To get the ratio identified in Step 5,
  add coefficients to the formulas which contain
  the elements whose oxidation number is changing.
• 2HNO3(aq)  3H3AsO3(aq) ?  NO(g)  H3AsO4(aq) 
  H2O(l)
• Step 6  Balance the rest of the equation by
  inspection.
• 2HNO3(aq)  3H3AsO3(aq)   ?   2NO(g)  3H3AsO4(aq)
    H2O(l)

16
Cu(s)  HNO3(aq)  ?  Cu(NO3)2(aq)  NO(g)  H2O(l)

• The nitrogen atoms and the oxygen atoms are
  difficult to balance by inspection, so we will go
  to Step 3.
• The copper atoms and some of the nitrogen atoms
  change their oxidation numbers. These changes
  indicate that this reaction is a redox reaction.
  We next determine the changes in oxidation number
  for the atoms oxidized and reduced.
• Cu  0 to 2      Net Change 2
• Some N  5 to 2     Net Change -3
• We need three Cu atoms (net change of 6) for
  every 2 nitrogen atoms that change (net change of
  -6). Although the numbers for the ratio
  determined in Step 5 are usually put in front of
  reactant formulas, this equation is somewhat
  different.  Because some of the nitrogen atoms
  are changing and some are not, we need to be
  careful to put the 2 in front of a formula in
  which all of the nitrogen atoms are changing or
  have changed. We therefore place the 2 in front
  of the NO(g) on the product side. The 3 for the
  copper atoms can be placed in front of the Cu(s).
  
•      3Cu(s)  HNO3(aq)  ?  Cu(NO3)2(aq)  2NO(g) 
   H2O(l)
• We balance the rest of the atoms, being careful
  to keep the ratio of Cu to NO 32.
•     3Cu(s)  8HNO3(aq)  ?  3Cu(NO3)2(aq)  2NO(g) 
   4H2O(l)

17
Summary

• 1. An oxidation-reduction reaction involves an
  apparent transfer of electrons from one particle
  to another.
• 2. Oxidation is the process by which electrons
  are apparently removed from an atom or group of
  atoms.
• 3. Reduction is the process by which electrons
  are apparently added to atoms or groups of atoms.
  
• 3. Any substance in a reaction which loses
  electrons is a reducing agent.
• 4. Any substance in a reaction which gains
  electrons is an oxidizing agent.
• 5. If a substance gives up electrons readily, it
  is a strong reducing agent. Its oxidized form is
  usually a poor oxidizing agent.
• 6. If a substance acquires electrons readily, it
  is a strong oxidizing agent. Its reduced form is
  usually a poor reducing agent.
• 7. Oxidation number is the charge an atom appears
  to have when we assign a certain number of
  electrons to that atom.
• 8. Six rules for assigning oxidation numbers
• a. The oxidation number of any free element is
  O.b. The oxidation number of any single-atom ion
  is equal to the that ion.c. The oxidation number
  of hydrogen is usually 1.
• d. The oxidation number of oxygen in most
  compounds is 2-.
• e. The sum of the oxidation numbers of all the
  atoms in a particle equal the apparent charge of
  that particle.
• f. In compounds, elements of Group IA and Group
  IIA have an oxidation number numerically equal to
  their group in the periodic table.
• 9. In all chemical reactions, charge, number and
  kind of atoms, and number of electrons are
  conserved. Knowing these quantities, you can do a
  redox equation.
• 10. Redox reactions are more easily balanced by
  splitting the equation into half-reactions.