Dr Phil King F207

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Introduction to Inorganic Chemistry Module 06510
Dr Phil KingF207
Oxidation States
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Why Are Oxidation States Important?

• A large proportion of chemical reactions involve
  changes in oxidation states.
• Combustion
• Corrosion
• Photosynthesis
• Metabolism of Food
• Extraction of Metals from Ores

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The Nature of Oxidation An Historical Perspective
The Phlogiston Theory

• Pioneered by Becher (1635-1682) and Stahl
  (1660-1734).
• Every flammable substance contains a substance
  called phlogiston.
• During combustion, phlogiston is given off into
  the air.
• Different materials contain different amounts of
  phlogiston
• Charcoal has lots Rocks have none.
• Fires go out when the air has taken up all the
  phlogiston it can.
• Wood ash is lighter than wood so must have lost
  phlogiston.

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The Nature of Oxidation An Historical Perspective

• On burning in air, metals were considered to lose
  phlogiston to the air.

• The calx was considered to be the pure elemental
  substance (i.e., the zinc without the phlogiston).

• Loss of phlogiston explains why the calx (zinc
  oxide) appears to be lighter than the metal.

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The Nature of Oxidation An Historical Perspective

• Addition of phlogiston to the calx would be
  expected to regenerate the metal.

• Heating the calx (metal oxide) with pure
  phlogiston (charcoal) does indeed regenerate the
  metal.

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The Nature of Oxidation An Historical Perspective

• Priestley (1733-1804).

• Obtained a new air by exposing mercuric oxide
  to sunlight.

• The new air allowed candles to burn more
  brightly than normal air.
• He called it dephlogisticated air (i.e. air that
  lacks phlogiston).

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The Nature of Oxidation An Historical Perspective

• Lavoisier (1743-1794).

• Experimented on the combustion of phosphorus and
  of sulfur.
• Devised weighing scales that could determine
  weight to within 0.0005g.
• Proved that upon burning in air materials
  actually gained in weight.
• Reducing the burned materials with charcoal led
  to loss of weight.
• The weight of a sealed container containing an
  oxide and charcoal remained the same before and
  after heating (law of conservation of mass).

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The Nature of Oxidation An Historical Perspective

• The term oxidation was coined to describe the
  reaction of a substance with oxygen.
• The oxidation reaction below is used in fireworks
  to produce white sparks.

• The reverse reaction is called a reduction
  reaction.

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The Nature of Oxidation An Historical Perspective

• In the reaction below, magnesium atoms each lose
  two electrons to form Mg2 and the oxygen atoms
  each gain two electrons to form O2-.

• The same pattern of reactivity can be seen for
  the reaction of magnesium with chlorine.

• It seems sensible to regard both reactions as
  oxidation reactions even though only one of them
  involves the gain of oxygen atoms.
• Oxidation can be defined as the loss of electrons.

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Definitions of Oxidation and Reduction Reactions.

• In modern chemistry we us more general
  definitions of oxidation and reduction.
• Oxidation Increase in oxidation number.
• Reduction Decrease in oxidation number.

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Assigning Oxidation Numbers

• Oxidation numbers (states) are theoretical values
  used to simplify electron bookkeeping.
• They are extremely useful. Some examples of
  their usage are listed below
• Determining the shapes of molecules
• Balancing redox equations
• Predicting the reactivity of main group compounds
• Predicting the reactivity of transition metal
  complexes
• Determining the magnetic properties of molecules
  and complexes

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Assigning Oxidation Numbers

• Oxidation numbers can be assigned using a simple
  set of rules.

• Rule 1 The oxidation number of an element
  uncombined with other elements is 0.
• In F2 the oxidation number of fluorine is 0.
• In Na metal the oxidation number of sodium is 0.
• In O2 the oxidation number of oxygen is 0.

• Rule 2 The sum of the oxidation numbers of all
  atoms in a species is equal to the total charge
  of the species.
• For a neutral molecule such as CO2 the sum of the
  oxidation numbers of carbon and the two oxygen
  atoms equals 0.
• For a charged ion such as NH4 the sum of the
  oxidation numbers of nitrogen and the four
  hydrogen atoms equals 1.

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Assigning Oxidation Numbers

• Oxidation numbers can be assigned using a simple
  set of rules.

• Rule 3 The oxidation number of hydrogen is 1
  when in combination with non-metals and -1 when
  in combination with metals.
• In HCl the oxidation number of hydrogen is 1
• In LiH the oxidation number of hydrogen is -1.

• Rule 4 The oxidation numbers of elements in
  Groups 1 and 2 are 1 and 2, respectively.
• Li, Na, K, Rb, Cs and Fr (always)! have oxidation
  numbers of 1
• Be, Mg, Ca, Sr, Ba and Ra always have oxidation
  numbers of 2

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Assigning Oxidation Numbers

• Oxidation numbers can be assigned using a simple
  set of rules.

• Rule 5 The oxidation number of the halogens
  (Group 17) is -1, except when in combination with
  oxygen or another halogen higher in Group 17.
  Fluorine in a compound is always -1.
• In HCl the oxidation number of chlorine is -1.
• In IO3- the oxidation number of the iodine atom
  is 5.
• In ICl2- the oxidation number of each chlorine
  atom is -1 (as Cl occurs higher up Group 17 than
  I). The oxidation number of the iodine atom is,
  therefore, 1.
• Note that the sum of the oxidation numbers is
  equal to the overall charge on the ion 2(-1)
  (1) -1.

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Assigning Oxidation Numbers

• Oxidation numbers can be assigned using a simple
  set of rules.

• Rule 6 The oxidation number of oxygen is -2
  (unless in combination with fluorine).
• In H2O the oxidation number of the oxygen atom is
  -2.
• In IO3- the oxidation number of each oxygen atom
  is -2, therefore, the oxidation number of the
  iodine atom is 5.
• Note that the sum of the oxidation numbers of
  the oxygen and iodine atoms equals the total
  charge on the ion 3(-2) (5) -1.

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Assigning Oxidation Numbers

• What is the oxidation number of the underlined
  element in each of the following?

• C60
• H3PO4
• AlH3

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Redox Reactions

• Consider the oxidation reaction shown below.

• The magnesium is oxidised from Mg0 to Mg2 (it
  loses two-electrons).
• Each of the two chlorine atoms is reduced from
  Cl0 to Cl- (a total gain of two-electrons).
• A reaction in which one species is reduced and
  another oxidised is called a redox reaction.
• The total number of electrons being lost in a
  reaction must equal the total number of electrons
  being gained.

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Redox Reactions

• A redox reaction can be described in terms of
  individual half-reactions for the oxidation and
  reduction processes.

• Magnesium is being oxidised, therefore the
  oxidation half-reaction can be written as
• Mg Mg2 2e-

• Chlorine is being reduced, therefore the
  reduction half-reaction can be written as
• Cl2 2e- 2Cl-

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Balancing Redox Equations

• Redox equations can be balanced by splitting the
  equation into two half-equations and balancing
  the numbers of electrons in each.

Oxidation Half-Equation Fe2
Fe3 e- Reduction Half-Equation Sn4 2e-
Sn2

• To balance the numbers of electrons lost and
  gained we must multiply the oxidation
  half-equation by 2.
• 2Fe2 2Fe3 2e-

Adding the two half-equations together gives the
balanced redox equation.
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Balancing More Complicated Redox Equations

• The reduction half-equation is

• We have oxygen atoms on the left but none on the
  right. We balance the number of oxygen atoms by
  adding the appropriate number of water molecules
  to the side lacking oxygen atoms.

• To balance the hydrogen atoms in the water
  molecules on the right hand-side of the equation
  we add H ions to the left hand-side.

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Balancing More Complicated Redox Equations

• Our reduction half-equation is

• Our oxidation half-equation is

• We must multiply the oxidation process by 5 to
  balance the electrons

• Adding the half-equations together affords the
  balanced redox equation

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The booster rockets of the space shuttle use
solid fuels consisting of aluminium powder
(fuel), ammonium perchlorate (oxidizing agent and
fuel) and iron(III) oxide (catalyst).
What is the oxidation number of the underlined
element in each of the following compounds? NH4
NO ClO4- Al Al2O3 AlCl3 The reaction is a
redox reaction. How many electrons are being
lost in oxidation process(es) and gained in
reduction process(es)? Perchlorates (ClO4-) are
powerful oxidising agents. Why do you think this
is so?