OXIDATION-REDUCTION REACTIONSREDOX

Oxygen Reactions• Early chemists saw oxidation only as the combination of

an element with oxygen to produce an oxide. • burning of fuel is an oxidation reaction that uses oxygen. • e.g. when methane (CH4), a component of natural gas,

burns in air, it oxidizes and forms oxides of carbon and hydrogen.

• one oxide of carbon is carbon dioxide, CO2.

• Not all oxidation involves burning. • iron turns to rust, it oxidizes to compounds such as iron(III) oxide (Fe2O3).

Not just Burning…

• liquid household bleach contains sodium hypochlorite (NaClO)• releases oxygen• oxidizes stains to a colorless form.

• Powdered bleaches may contain:• calcium hypochlorite (Ca(ClO)2), • sodium perborate (NaBO3), or • sodium percarbonate (2Na2CO3 · 3H2O2).

Bleaching stains:

• Hydrogen peroxide (H2O2) also releases oxygen when it decomposes.

• It is both a bleach and a mild antiseptic that kills bacteria by oxidizing them.

• (insert dumb bar joke here)

Reduction• opposite of oxidation. • Originally, reduction meant the loss of oxygen from a

compound. • The reduction of iron ore to metallic iron involves the

removal of oxygen from iron(III) oxide. • The reduction is accomplished by heating the ore with

carbon, usually in the form of coke.(what they make at Tonawanda Coke).

• 2Fe2O3 + 3C 4Fe + 3CO2

What is Redox?• Oxidation and reduction always occur simultaneously.

• No oxidation occurs without reduction, and no reduction occurs without oxidation.

• The substance gaining oxygen is oxidized, while the substance losing oxygen is reduced.

• Reactions that involve these processes are therefore called oxidation-reduction reactions.

• are also known as redox reactions.

Electron Transfer• Modern concepts of oxidation extend to include many

reactions that do not involve oxygen. • Oxygen is the most electronegative element next to

fluorine• It’s an “electron stealer” • As a result, when oxygen bonds with an atom of a

different element (other than fluorine), electrons from that atom shift toward oxygen.

• Redox reactions are currently understood to involve any shift of electrons between reactants.

New Definitions (must know)• Oxidation - a process that involves complete or partial

loss of electrons OR a gain of oxygen; it results in an increase in the oxidation number of an atom

• Reduction - a process that involves a complete or partial gain of electrons OR the loss of oxygen; it results in a decrease in the oxidation number of an atom

Memory aid:

• “LEO says GER” •Lose Electrons Oxidation •Gain Electrons Reduction

More confusing vocab…• oxidizing agent - the substance in a redox reaction that

accepts electrons • in the reaction, the oxidizing agent gets reduced • Reduced = gains electrons

• reducing agent - the substance in a redox reaction that donates electrons • in the reaction, the reducing agent gets oxidized • Oxidized = loses electrons

Corrosion• the electrochemical reaction between a material, usually a

metal, and its environment that produces a deterioration of the material and its properties.

• Iron, a common metal often used in the form of the alloy steel, corrodes by being oxidized to ions of iron by oxygen.

• Oxygen, the oxidizing agent, is reduced to oxide ions (in compounds such as Fe2O3) or to hydroxide ions.

• The following equations describe the corrosion of iron to iron hydroxides in moist conditions.

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

• Corrosion occurs more rapidly in the presence of salts and acids. • Cars rust faster in where there is salt on

roads • These substances produce

electrically conducting solutions that make electron transfer easier. • The salt doesn’t rust that car, it just

makes it easier to rust faster.• The corrosion of some metals can

be a desirable feature.• Verdigris

Corrosion Resistance• Gold and platinum are “noble metals” because they are very

resistant to losing their electrons by corrosion. • Other metals lose electrons easily but are protected from

extensive corrosion by the oxide coating formed on their surface. • Aluminium oxidizes to form a coating of very tightly packed

aluminium oxide particles. • coating protects the aluminium object from further corrosion. • Iron forms a coating when it corrodes, but the coating of iron

oxide that forms is not tightly packed. • Water and air can penetrate the coating and attack the iron

metal below it. • The corrosion continues until the iron object becomes only a pile

of rust.

Controlling Corrosion

• corrosion of a steel support pillar of a bridge or the hull of an oil tanker is much more serious and costly.

• To prevent corrosion, the metal surface may be coated with oil, paint, plastic, or another metal.

• These coatings exclude air and water from the surface preventing corrosion.

• If the coating is scratched or worn away, the exposed metal will begin to corrode.

Another Control method• one metal is “sacrificed,” allowed to corrode, in order to save

another metal. • e.g., to protect an iron object, a piece of magnesium (or

another active metal ) may be placed in electrical contact with the iron.

• when oxygen and water attack the iron object, the iron atoms lose electrons as the iron begins to be oxidized.

• magnesium is a better reducing agent than iron (more easily oxidized - reference table J), the magnesium transfers electrons to the iron, preventing their oxidation to iron ions.

• So, the magnesium is “sacrificed” by oxidation and protects the iron in the process.

More example of sacrifice metals• zinc and magnesium blocks are sometimes attached to

piers and ship hulls to prevent corrosion damage in areas submerged in water.

• Underground pipelines and storage tanks may be connected to magnesium blocks for protection.

• It is easier and cheaper to replace a block of magnesium or zinc than to replace a bridge or a pipeline.

Assigning Oxidation Numbers• positive or negative number assigned to an atom to

indicate its degree of oxidation or reduction. • a bonded atom’s oxidation number is the charge that it

would have if the electrons in the bond were assigned to the atom of the more electronegative element.

Rules For Oxidation Numbers (p. 1)• Monatomic ions have the charge of the ion, Halogens are almost always -1, iron (III) is Fe+3

• Hydrogen is +1 except when bonded to a metal (like NaH), then it is -1

• Oxygen is -2 except in a peroxide (H2O2), then it is -1 and when bonded to fluorine, then it is postive.

• Crack out reference tables – Periodic Table

Rules For Oxidation Numbers (p. 2)• Elements not bonded to anything are ZERO

• Neutral compound must total zero when adding oxidation numbers

• Polyatomic ions, use the chart in the reference tables.

Almost Always…• Alkali Metals are +1• Alkaline Earths are +2• Halogens are -1• Oxygen is -2• IF there is a polyatomic ion, use the chart in the reference table and figure out what the other ion would be.

Identifying Reactions• All chemical reactions can be assigned to one of two classes.

• #1 Redox - electrons are transferred from one reacting species to another. • single-replacement reactions, • synthesis reactions • decomposition reactions• combustion reactions

• #2 Other reactions - no electron transfer occurs. • double-replacement reactions • acid-base reactions