R-AECHE-46

Oxidation is Loss, Reduction is Reduced OILRIG

We need to identify when a substance is oxidised or reduced, this is assisted via the use of oxidation numbers.

Electrical energy is produced from potential chemical energy.

There are many types of batteries,

• secondary
• primary
• fuel cells

Oxidisation Numbers

Shows number of electrons lost, gained, or shared unequally.

• Chlorine: Covalent bond, electrons shared equally, therefore oxidisation number is 0.
• $HCl_{(aq)}$: $H$: +, $Cl$: -
  • Shared unequally, thus $Cl$ is -1 (more electronegative), $H$ is +1.
• Simple Ion: Oxidisation number same as charge of ion!
  • Only mono-atomic ions?
  • Single atom
• $MgCl_{2}$: $Mg$: +2, $Cl$: -1
• $AlCl_{3}$: $Al$: +3, $Cl$: -1
• $Al_{(s)}$: 0, although electrons are delocalised, the metal is still neutral; electrons aren't transferred to another substance.
• $Br_{2}$: 0, Shared equally.
• $H_{2}O$: $H$: +1, $O$: -2, $H_{2}O$: 0
• $H_{2}O_{2}$: $H$: +1, $O$: -1
  • Drawing Lewis diagram, one sees that Oxygen bonds to 1 Hydrogen and 1 Oxygen. The Oxygen-Oxygen bond is equal, and thus there is only 1 electron being unequally shared.
  • This is the only case where oxygen has an oxidisation number of -1.

Rules of Oxidisation Number:

• IF you have uncombined element/element combined with itself, oxidation number/state is 0.
• IF you have simple ions, the oxidation number/station is charge of ion.
• For neutral compounds, the sum of the oxidation states is 0
• Elements with fixed oxidation states:
  • Hydrogen (except with metal hydrides): +1 (-1)
  • Oxygen (except in peroxides): -2 (-1)
  • Group 1 elements: +1
  • Group 2 elements: +2
  • Aluminium: +3
  • Halides: -1

Examples: $CuO$ - $Cu$: +2, $O$: -2 Note that this is Copper (II) Oxide.

• The (II) in this case is its oxidation number

Sodium Chlorate (I).

• $Na$: +1
• $ClO$: -1
  • Cl: +1
  • O: -2
• In this case, the (I) is for the Cl!!

Sodium Chlorate (?)

• ClO3: -1

• Cl: +5

• O:-2

  $\therefore$ The oxidation state for Cl is (V)

$SO^{-2}_{3}$: $S$: +4

Zinc + Copper (ii) Sulfate ->

• Molecular eq
  • $Zn_{(s)} + CuSO_{4(aq)} \to ZnSO_{4(aq)} + Cu_{(s)}$
• Ionic eq
  • $Zn_{(s)} + Cu^{2+}{(aq)} \to Zn^{2+}{(aq)} + Cu_{(s)}$
• What is oxidised? What is reduced?
  • Zinc is oxidised, Copper in $CuSO_4$ solution(or $Cu^{2+}$ ions) is reduced.

$Mg_{(s)} + 2H^{+}{(aq)} \to Mg^{2+}{{(aq)}} + H_{2(g)}$

Mg is the reducing agent, $H_{2}SO_4$ (NOT $H^{+}$ CAN WRITE $H^{+}$ IN $H_{2}SO_4$) solution is the oxidising agent.

• $Mg_{{(s)}} \to Mg_{(aq)}^{2+} + 2\overline{e}$
• $2H^{+}{(aq)} + 2\overline{e} \to H{2(aq)}$

$Cu^{2+}$ in $CuCl_{2}$ is the oxidising agent, $SO_{2}$ is the reducing agent

$H_2SO_4$ is oxidising agent, $HI$ is reducing agent.

c) $HI + H_2SO_4 \to H_2S + 4H_2O + 4I_2$

• $2I^{-} \to I_{2} + 2\overline{e}$
• $10H^{+} + SO_{4}^{2-} + 8\overline{e}\to H_{2}S +4 H_{2}O$

$MnO_{4}^{-} + 8H^{+} + 5\overline{e} \to Mn^{2+} + 4H_{2}O$ $5Fe^{2+} \to 5Fe^{3+} + 5\overline{e}$ $MnO_{4}^{-} + 8H^{+} + 5Fe^{2+} \to Mn^{2+} + 4H_{2}O + 5Fe^{3+}$

Galvanic Cells

• $Zn_{(s)} + CuSO_{4(aq)} \to ZnSO_{4(aq)} + Cu_{(s)}$
• Spontaneous reaction: happens at room temperature, on its own
• $Zn_{(s)} \to Zn^{2+}_{(aq)} + 2\overline{e}$
• $Cu^{+}{(aq)} + 2\overline{e}\to Cu{(s)}$

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