Chapter 20: Oxidation -Reduction reactions

Section 20.1 - The meaning of oxidation and reduction

Why does this happen to cars? Why is it less of an issue in Arizona compared to other states?

Has the statue of liberty always been green?

Oxidation reactions● Oxidation reactions were originally

thought to be reactions that involved the combination of an element with oxygen to produce an oxide

● This happens in a number of everyday reactions

● Whenever something is burned● Whenever something corrodes (rusts)● Whenever you bleach a substance● When hydrogen peroxide decomposes

Reduction reactions● Reduction reactions are the opposite of oxidation● Originally, this was believed to signify simply the loss of oxygen from a

compound● That is a good rule of thumb, but is not always the case● A common example is the reduction of iron ore ● Oxygen is removed, iron ore and carbon dioxide are formed ● This occurs when iron ore and carbon are heated together

The relationship between oxidation and reduction● These two processes always happen together● Oxidation does NOT happen without reduction, and reduction does not

happen without oxidation● In the previous example, iron(III) oxide is reduced and the carbon is oxidized ● These complementary oxidation-reduction reactions are commonly known as

redox reactions

Does oxidation always require oxygen?● No! ● Although originally

scientists thought that oxygen was necessary, now we know that instead a transfer of electrons is key in this style of reaction.

● This is because oxygen is highly electronegative

● OILRIG mnemonic

What does this look like in terms of electron transfer?

● When a metal and a nonmetal are reacted together electrons are transferred from the metal to the nonmetal

● When you react magnesium and sulfur, two electrons are transferred from a magnesium atom to a sulfur atom

● The magnesium atoms are more stable by the loss of electrons, and the sulfur is more stable by the gain of electron

● Magnesium is oxidised and sulfur is reduced

Oxidising and reducing agentsA reducing agent is a substance that undergoes oxidation and loses electrons

Magnesium is the reducing agent in the prior reaction

An oxidising agent is a substance that undergoes reduction and gains electrons

Sulfur is the oxidising agent in the prior reaction

Write in your notes and fill in the blanks….In a chemical reaction, when a substance ________ electrons it undergoes _______ and is also called the __________ agent.

Practice identifying oxidizing and reducing agent Page 662

● You need to identify the reactants and the products● Identify the changes in charges of substances over the course of the reaction● Establish which substances have gained electrons, and which substances lost

electrons

In covalent compounds, it is not so easy…..● When a metal and a nonmetal react and form ions, it is easy to identify the

transfer of electrons● In covalent compounds electrons do not transfer completely● In covalent compounds electrons are shared● We can refer to partial loss and partial gains of electrons● In polar molecules, such as water, the electrons are pulled towards the

oxygen - why?● The hydrogen undergoes partial loss of electrons and is oxidized

Definitions of oxidation and reduction

Decrease in oxidation numberIncrease in oxidation number

Gain of hydrogen by a covalent compound

Loss of hydrogen by a covalent compound

Loss of oxygenGain of oxygen

Shift of electrons toward an atom in a covalent bond

Shift of electrons away from an atom in a covalent bond

Complete gain of electrons (ionic reactions)

Complete loss of electrons (ionic reactions)

ReductionOxidationProcesses Leading to Oxidation and Reduction

Why is it so important to know about oxidation and reduction reactions?

● Corrosion is an example of a redox reaction● Each year corrosion costs the american economy billions of dollars ● Iron is a common example

○ Atoms are oxidized by oxygen to form ions of iron, a process that speeds up in the presence of water

● In this process the oxygen is reduced to oxide ion● What happens to these ions? ● Process speeds up in the presence of salts

Iron can be oxidized to form iron hydroxide or Iron Oxide

The oxide ions can react with either Iron ions, (formed from oxidation) to form Fe2O3, or can react with hydroxide ions to form Iron Hydroxide

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

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

This process speeds up in the presence of salts and acids, as they speed up the conductive process and facilitate electron transfer

Do all metals corrode at the same rate? ● Some metals have an inherent resistance to

corrosion○ Example - gold and platinum, these are called

noble metals

● Their electrons are tightly held, meaning it is hard for them to be oxidized

● Some metals oxidize very quickly - and form an oxide coating

○ This protects the aluminium object from further corrosion

● Iron can also form a coating, but this is not tightly packed together, and the metal can still be damaged

Iron(III) oxide Aluminum oxide

Water Water

Oxygen Oxygen

Why are zinc blocks attached to the hulls of ships● The zinc is used a sacrificial

metal on the ships to save the iron on the ship

● When oxygen and water “attack” the iron, the iron atom loses electrons

● Zinc and magnesium are better reducing agents than iron, and immediately transfer electrons back to the iron atoms

● Can also be used for bridges, pipeline, storage tanks...

How else can corrosion be prevented?

● Coating the surface● This could be with oil, paint, plastic,

or another metal● Air and water are kept away from

the metal● If the coating is scratched, or

removed at all, the exposed metal will begin to corrode

Key Points● Oxidation and reduction are complementary reactions involving the loss/gain

of electrons ● An oxidizing agent gains electrons and is reduced● A reducing agent loses electrons and is oxidized ● Corrosion is a common consequence of redox reactions● Corrosion can be prevented through coating the surface, or the use of

sacrificial metals

Oxidation numbersSection 20.2

How do you express the degree of oxidation or reductionAn oxidation number is a positive or negative number assigned to an atom to indicate the degree of oxidation or reduction

A number of rules exist surrounding the determination of oxidation numbers

A bonded atoms 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

Oxidation numbers and ionic compounds In ionic compounds the oxidation numbers will equal their ionic charge

Remember the overall charge of an ionic compound is 0

Metals will always have a positive charge

Non metals will always have a negative charge

Polyatomic ions can have a positive or negative charge

Oxidation numbers and molecular compounds In molecular compounds, due to the covalent bond there is no ionic charges associated with atoms

But, oxygen is reduced when hydrogen and oxygen form water

Oxygen is a highly electronegative element, much more so than in the hydrogen

The two shared electrons are shifted towards the oxygen

In terms of oxidation number, assume that the electrons have been transferred

Hydrogen would be +1, Oxygen would be +2

Further rules for oxidation numbersMany elements can have several oxidation numbers

In particular, you can use neutral compounds and polyatomic ions

Some elements can also have multiple oxidation numbers, such as Cr

Chromium metal - oxidation number 0

Potassium dichromate = K2Cr2O7 - oxidation number +6

Chromium (III) potassium sulfate decahydrate - CrK(SO4)2.12H2O - oxidation np. 3

Summary of rules

What is the oxidation number of each kind of atom in the following ions and compounds?

a. SO2 c. Na2SO4

b. CO32– d. (NH4)2S

How does this all relate to chemical reactionsDuring redox chemical reactions, the oxidation number of elements may change

If copper reacts with silver nitrate, the oxidation number of silver decreases from +1 to 0

Copper ions are also oxidized

Oxygen and reduction can be defined in terms of changes in oxidation number

The relationship between oxidation number and gemstonesGemstones get their color from impurities

Eg. Iron, chromium, copper

The oxidation number of the metal can also change the color of the gemstone

Example - iron in beryl

Describing redox reactions

Section 20.3

Identifying redox reactionsReactions are either a REDOX reaction or not

All reactions fall into one of these two categories

REDOX: Single replacement, combination, decomposition, and combustion

NON redox - double replacement or acid - base reactiosn

Using oxidation numbers to identify redox reactions

In a lightning storm, oxygen and nitrogen molecules react to form nitrogen monoxide

N2(g) + O2(g) -> 2NO(g)

Is this a redox reaction?

Yes! The oxidation number of nitrogen increase to 2, the number of oxygen decreases to -2

Other indicators of redox reactions

Color changes often signify that a redox reaction is taking place

Balancing redox equationsRedox reactions can be complex, and often cannot be balanced by trial and error

The fact that the number of electrons gained in reduction is equal to the number of electrons lost in oxidation can be used to balance an equation

Two methods:

Oxidation number changes

Using half equations

Using oxidation number changesIn this method you compare the changes in oxidation number

1) Start with an unbalanced equation2) Assign oxidation numbers to all atoms in the equation3) Identify which atoms are being oxidized and which are being reduced4) Use bracketing lines to connect atoms that undergo oxidation and those that

undergo reduction5) Use coefficients to make the total increase in oxidation number equal the total

decrease in oxidation number6) Make sure the atom is balanced for both atoms and charge

Worked example - 1) Fe2O3, 2) sample problem 20.5

Half reactionsA half reaction is an equation that shows just the oxidation or just the reduction that takes place in a redox reaction

In this method you write and balance the oxidation half reaction, then you write and balance the reduction half reaction

In the example:

S(s) + HNO3 -> SO2(g) + NO(g) + H2O (l)

S(s) + -> SO2(g) Oxidation half reaction

NO3-(aq) -> NO(g) Reduction half reaction

Balancing half reactions1) S(s) + -> SO2(g) Oxidation half reaction2) NO3

-(aq) -> NO(g) Reduction half reaction

In equation 1 S is being oxidised as 4 electrons are being removed

In equation 2 N is being reduced - charge changes from +5 to +2

Atoms must be balanced in each half equation, and then the charges so that electrons gained equal electrons lost

When each half reaction is balanced you combine them to form a balanced chemical equation