chapter 22 redox. the meaning of oxidation and reduction oxidation numbers balancing redox equations

Chapter 22 REDOX

Chapter 22 REDOX

The Meaning of Oxidation and Reduction Oxidation Numbers Balancing Redox Equations

Ch 22.1 The Meaning of Oxidation and Reduction

Oxygen in Redox Reactions Electron Transfer in Redox Reactions Corrosion

Oxygen in Redox Reactions

Oxidation – the combination of an element with oxygen to produce oxides

Oxygen in Redox Reactions

Burning

Oxygen in Redox Reactions

Bleaching

Oxygen in Redox Reactions

Rusting

Reduction

Reduction – the loss of oxygen from a compound

Redox Reactions

Reduction and Oxidation always occur together

2Fe2O3(s) + 3C(s) 4Fe(s) + 3CO2(g)

reduction oxidation

Electron Transfer in Redox Reactions

Oxidation – loss of electrons, gain oxygen Reduction – gain of electrons, loss of

oxygen “LEO the lion goes GER” LEO – Lose electrons oxidation GER – Gain electrons reduction

Electron Transfer in Redox Reactions

Oxidation: Mg Mg2+ + 2e-

Loss of electrons Reduction: S + 2e- S2-

Gain of electrons

Corrosion

Corrosion

2Fe(s) + O2(g) + 2H2O(l) 2Fe(OH)2(s)

4Fe(OH)2(s) + O2(g) + 2H2O(l) 4Fe(OH)3(s)

Corrosion of iron

Corrosion

Some metals completely corrode Iron

Some metals form a protective coating Aluminum

Some metals do not corrode at all Gold

Chapter 22.3 Balancing Redox Reactions

Identifying Redox Reactions Using Oxidation Number Changes Using Half Reactions

Identifying Redox Reactions

Two types of reactions: REDOX – electrons are transferred Everything else: single replacement, double

replacement, combustion, …. NO transfer of electrons

Identifying Redox Reactions

REDOX – the oxidation number of an element changes N2(g) + O2(g) 2NO(g)

Using Oxidation Number Changes

Fe2O3(s) + CO(g) Fe(s) + CO2(g)

Step 1 – Assign oxidation numbers to all atoms in the equation

Fe2O3(s) + CO(g) Fe(s) + CO2(g)

+3 -2 +2 –2 0 +4 -2

Using Oxidation Number Changes

Step 2 – Identify which atoms are oxidized and reduced

Fe2O3(s) + CO(g) Fe(s) + CO2(g)

+3 -2 +2 -2 0 +4 -2

Iron – reduced, Carbon - oxidized

Using Oxidation Number Changes

Step 3 – Use a bracket line to connect the atoms undergoing oxidation and one to connect the lines undergoing reduction

Fe2O3(s) + CO(g) Fe(s) + CO2(g)

-3

+2

Using Oxidation Number Changes

Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients

Fe2O3(s) + CO(g) Fe(s) + CO2(g)

3 x (+2) = 6

2 x (-3) = - 6

Using Oxidation Number Changes

Step 5 – Finally make sure the equation is balanced for both atoms and charge

Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

FRH – Flameless Ration Heater

Mg + H2O Mg(OH)2 + H2 + Heat

Problem: Mg forms a coating from corrosion – MgO – which is not water soluble, prevents the above reaction from happening

Solution: Add NaCl and Fe to the mix, breaks down the MgO and allows the reaction to happen

Chapter 23 Electrochemistry

Electrochemical Cells Half Cells and Cell Potentials Electrolytic Cells

Electrochemical Cells

The Nature of Electrochemical Cells Voltaic Cells Dry Cells Lead Storage Batteries Fuel Cells

The Nature of Electrochemical Cells

Zn(s) + Cu2+(aq) Zn2+

(aq) + Cu(s)

The Nature of Electrochemical Cells

The zinc bar becomes copper plated Zinc loses electrons and dissolves slowly Copper gains electrons and becomes a

solid Oxidation: Zn(s) Zn2+

(aq) + 2e-

Reduction: Cu2+(aq) + 2e- Cu(s)

The Nature of Electrochemical Cells

Reference Table J Look at any two metals, the metal that is

higher on the table is the one that is more readily oxidized

Electrochemical Cell

Any device that converts chemical energy into electrical energy or electrical energy into chemical energy

REDOX reactions must occur If an electrochemical cell is to be used for

electrical energy, the two half reactions must physically be separated

Voltaic Cell

Alessandro Volta (1745 – 1827) First electrochemical cell

Voltaic Cell

Convert chemical energy into electrical energy

Half Cell – part of a voltaic cell, consists of a metal rod in a solution of ions

Salt Bridge – Separates half cells, tube containing a strong electrolyte (can also use a porous plate)

Voltaic Cell

Anode – the anode where oxidation occurs

Cathode – the cathode where reduction occurs

Dry Cell

A voltaic cell in which the electrolyte is a paste

Alkaline Battery

Improved dry cell, the Zinc electrode doesn’t corrode as fast

Lead Storage Batteries

A group of cells connected together

Fuel Cell

Voltaic cell in which a fuel substance undergoes oxidation

Do not have to be recharged Oxidation: 2H2(g) + 4OH-

(aq) 4H2O(l) + 4e-

Reduction: O2(g) + 2H2O(l) + 4e- 4OH-(aq)

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Hydrogen Refueling Stations

The diagram shows a voltaic cell with copper and aluminum electrodes immediately after the external circuit is completed.

1 Balance the redox equation using the smallest whole-number coefficients. [1]

2 As this voltaic cell operates, the mass of the Al(s) electrode decreases. Explain, in terms of particles, why this decrease in mass occurs. [1]

3 Explain the function of the salt bridge. [1]

Answers

3 Cu2+ (aq) + 2 Al(s) 3 Cu(s) + 2 Al3+

(aq)

Aluminum particles are losing electrons and becoming aluminum ions that are entering the solution.

It allows migration of ions, maintains neutrality, prevents polarization

Electrolytic Cells

An electrochemical cell used to cause chemical change through the application of electrical energy (electrical energy is added)

Differences

Voltaic (Galvanic) Cells Electrolytic Cells

Flow of electrons is spontaneous

Flow of electrons is pushed by an outside power source

Anode negative

Cathode positive

Anode positive

Cathode negative

Similarities

Voltaic (Galvanic) Cells and Electrolytic Cells

Electrons flow from anode to cathode Reduction – cathode Oxidation – anode

Electroplating is an electrolytic process used to coat metal objects with a more expensive and less reactive metal. The diagram below shows an electroplating cell that includes a battery connected to a silver bar and a metal spoon. The bar and spoon are submerged in AgNO3(aq).

Explain why AgNO3 is a better choice than AgCl for use in this electrolytic process. [1]

Explain the purpose of the battery in this cell. [1]

Acceptable responses include, but are not limited to: Silver nitrate produces more ions than silver

chloride in water. AgNO3 readily dissolves in H2O; AgCl dissolves

only slightly in H2O. Acceptable responses include, but are not

limited to: The battery provides the electrical energy

necessary for the reaction to occur.

The apparatus shown in the diagram consists of two inert platinum electrodes immersed in water. A small amount of an electrolyte, H2SO4, must be added to the water for the reaction to take place. The electrodes are connected to a source that supplies electricity.

What type of electrochemical cell is shown? [1] What particles are provided by the electrolyte that

allow an electric current to flow? [1]

Electrolytic or electrolysis. Acceptable responses include, but

are not limited to: Ions, charged particles, H3O+, SO4

2–

Because tap water is slightly acidic, water pipes made of iron corrode over time, as shown by the balanced ionic equation below:

2Fe + 6H+ 2Fe3+ + 3H2

Explain, in terms of chemical reactivity, why copper pipes are less likely to corrode than iron pipes. [1]

Acceptable responses include, but are not limited to: Copper is less reactive than iron. Cu below H2 on Table J