Chapter 19 oxidation reduction reactions
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Chapter 19 Oxidation - Reduction Reactions. 19.1 Oxidation and Reduction. Oxidation – Reduction . Most substances have the same oxidation number as their individual charge (the more electronegative element 1 st )

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Chapter 19 oxidation reduction reactions

Chapter 19Oxidation - Reduction Reactions

19.1 Oxidation and Reduction


Oxidation reduction
Oxidation – Reduction

  • Most substances have the same oxidation number as their individual charge

  • (the more electronegative element 1st)

  • All oxidation numbers in a compound must add up to equal the total charge on the compound

  • All single elements have a oxidation number of zero

  • All single ions have the same oxidation number as their charge

  • Oxygen has a charge of -2, except with F (+2), or peroxides (-1)

  • Hydrogen is usually +1, unless with a metal (-1)

  • Rules for assigning oxidation numbers


Oxidation
Oxidation

  • Reactions in which the atoms or ions of an element experience an increase in oxidation state

  • A species whose oxidation number increases is oxidized

Oxidation Half- Reaction

Fe → Fe+2 + 2e-


Reduction
Reduction

  • Reactions in which the oxidation state of an element decreases.

  • A species that undergoes a decrease in oxidation state is reduced.

Reduction Half- Reaction

Cu+2 + 2e- → Cu

Cu+2 + 2e- → Cu

+

Fe → Fe+2 + 2e-

Overall Rxn

Fe + Cu+2→ Fe+2 + Cu


Oxidation reduction1
Oxidation - Reduction

AgCl(aq) + Na(s) Ag(s) + NaCl(aq)

+1

-1

0

0

-1

+1

Charge increased from 0 to +1, so e- were lost

Charge reduced from +1 to 0, so e- were gained

The sodium was oxidized

The silver in silver chloride was reduced

Oxidation is loss, reduction is gain

OIL RIG


Leo the Lion goes Ger

Lose Electrons Oxidation

Gain Electrons Reduction

Electron Loss Means Oxidation


Oxidation reduction2
Oxidation - Reduction

H2O(l) H2(g) + O2(g)

+1

-2

0

0

Charge increased from -2 to 0, so e- were lost

Charge reduced from +1 to 0, so e- were gained

The hydrogen in water was reduced

The oxygen in water was oxidized


Example
Example

  • Household Bleach removes stains through a redox reaction:

    Stain molecules (s) + OCl-(aq) → colorless molecules (s) + Cl-(aq)

  • Determine the oxidation numbers of oxygen & chlorine in OCl- .

OCl-

-2 +1


Chapter 19 oxidation reduction reactions1

Chapter 19Oxidation - Reduction Reactions

19.2 Balancing Redox Equations


Example1
Example

  • Balance the following reaction:

    I- + MnO4- + H+ → MnO2 + I2 + H2O

2I- + MnO4- + 4H+ → MnO2 + I2 + 2H2O

But this reaction is balanced for mass not charge!

A half-reaction system has to be used to balance for charge.


Half reaction method
Half-Reaction Method

  • Write the formula equation then ionic equation

  • Assign oxidation numbers. Exclude anything with an ox. # of zero, or that doesn’t change ox. #

  • Write the ½ rxn for oxidation

  • Balance the atoms

  • Balance the charge (w/ electrons)

  • Write the ½ rxn for reduction

  • Balance the atoms

  • Balance the charge (w/ electrons)

  • Use coefficients to ensure the # of e- lost in ox. equals the # of e- gained in red.

  • Combine both ½ rxns and cancel (like Hess’s Law)

  • Combine ions to form initial compounds.


Example2
Example

Half-Reactions:

2MnO4- + 8H+ + 6e- → 2MnO2 + 4H2O

6I- → 3I2 + 6e-

Now try to balance the following reaction:

I- + MnO4- + H+ → MnO2 + I2 + H2O

Overall Balanced Equation:

6KI + 2KMnO4+ 8HCl → 2MnO2 + 3I2 + 4H2O + 8KCl


Chapter 19 oxidation reduction reactions2

Chapter 19Oxidation - Reduction Reactions

19.3 Oxidizing and Reducing Agents


Reducing agent
Reducing Agent

  • Substance that has the potential to cause another substance to be reduced.

  • They lose electrons; are oxidized.

Fe + Cu+2→ Fe+2 + Cu

Iron causes Copper to become reduced, so it is the Reducing Agent

Fe → Fe+2 + 2e-


Oxidizing agent
Oxidizing Agent

  • Substance that has the potential to cause another substance to be oxidized.

  • They gain electrons; are reduced.

Fe + Cu+2→ Fe+2 + Cu

Cu+2 + 2e- → Cu

Copper causes Iron to become oxidized, so it is the Oxidizing Agent


Disproportionation autooxidation
Disproportionation/Autooxidation

  • A process by which a substance acts as both an oxidizing and reducing agent


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