Oxidation-reduction Reactions Princess...

8/2/2019 Oxidation-reduction Reactions Princess...

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The term oxidation was originally used todescribe reactions in which an element

combines with oxygen.


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Example: The reaction between magnesiummetal and oxygen to form magnesium oxideinvolves the oxidation of magnesium.

The term reduction comes from the Latin

stem meaning "to lead back." Anything thatthat leads back to magnesium metaltherefore involves reduction.


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The reaction between magnesium oxide andcarbon at 2000C to form magnesium metal andcarbon monoxide is an example of the reduction ofmagnesium oxide to magnesium metal.

After electrons were discovered, chemists becameconvinced that oxidation-reduction reactionsinvolved the transfer of electrons from one atom toanother. From this perspective, the reactionbetween magnesium and oxygen is written asfollows.


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2 Mg + O2 2 [Mg2+][O2-]

In the course of this reaction, each

magnesium atom loses two electrons toform an Mg2+ ion.

Mg Mg2+ + 2 e-

And, each O2 molecule gains four electrons

to form a pair of O2- ions.

O2 + 4 e- 2 O2-


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Because electrons are neither created nordestroyed in a chemical reaction, oxidationand reduction are linked. It is impossible to

have one without the other, as shown inthe figure below.


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Chemists eventually extended the idea ofoxidation and reduction to reactions thatdo not formally involve the transfer of

electrons.Consider the following reaction.

CO(g) + H2O(g) CO2(g) + H2(g)

As can be seen in the figure below, the

total number of electrons in the valenceshell of each atom remains constant in thisreaction.


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What changes in this reaction is the oxidationstate of these atoms. The oxidation state ofcarbon increases from +2 to +4, while theoxidation state of the hydrogen decreases from

+1 to 0.

Oxidation and reduction are therefore bestdefined as follows. Oxidation occurs when theoxidation number of an atom becomeslarger. Reduction occurs when the oxidationnumber of an atom becomes smaller.


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The terms ionic and covalent describethe extremes of a continuum of bonding.There is some covalent character in even

the most ionic compounds and vice versa.


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It is useful to think about the compounds of the maingroup metals as if they contained positive and negativeions. The chemistry of magnesium oxide, for example, iseasy to understand if we assume that MgO containsMg2+ and O2- ions. But no compounds are 100% ionic.There is experimental evidence, for example, that the

true charge on the magnesium and oxygen atoms in MgOis +1.5 and -1.5.

Oxidation states provide a compromise between apowerful model of oxidation-reduction reactions basedon the assumption that these compounds contain ions

and our knowledge that the true charge on the ions inthese compounds is not as large as this model predicts.By definition, the oxidation state of an atom is thecharge that atom would carry if the compound werepurely ionic.


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For the active metals in Groups IA and IIA, thedifference between the oxidation state of the metalatom and the charge on this atom is small enough to beignored. The main group metals in Groups IIIA and IVA,however, form compounds that have a significantamount of covalent character. It is misleading, for

example, to assume that aluminum bromide containsAl3+ and Br- ions. It actually exists as Al2Br6 molecules.

This problem becomes even more severe when we turnto the chemistry of the transition metals. MnO, forexample, is ionic enough to be considered a salt that

contains Mn2+ and O2- ions. Mn2O7, on the other hand, isa covalent compound that boils at room temperature. Itis therefore more useful to think about this compoundas if it contained manganese in a +7 oxidation state, notMn7+ ions.


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Let's consider the role that each elementplays in the reaction in which a particular

element gains or loses electrons..


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When magnesium reacts with oxygen, themagnesium atoms donate electrons toO2molecules and thereby reduce the

oxygen. Magnesium therefore acts asa reducing agent in this reaction.

2 Mg+ O2 2MgOreduc

ing

agent


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2 Mg

+ O22

MgOoxidi

zing

agen

t

The O2molecules, on the otherhand, gain electrons frommagnesium atoms and therebyoxidize the magnesium. Oxygen is

therefore anoxidizing agent.


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Oxidizing and reducing agents thereforecan be defined as follows. Oxidizingagents gain electrons. Reducingagents

lose electrons.

The table below identifies the reducingagent and the oxidizing agent for some of

the reactions discussed in this web page.One trend is immediately obvious: Themain group metals act as reducing agentsin all of their chemical reactions


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TYPICAL REACTIONS OF MAIN GROUP METALS

Reaction Reducing

AgentOxidizing

Agent2 Na + Cl2 2 NaCl Na Cl22 K + H2 2 KH K H24 Li + O2 2 Li2O Li O22 Na + O2 Na2O2 Na O22 Na + 2 H2O 2

Na+ + 2 OH- + H2 Na H2O2 K + 2 NH3 2 K

+ +

2 NH2- + H2 K NH3

2 Mg + O2 2 MgO Mg O23 Mg + N

2Mg

3N

2 Mg N

2

Ca + 2 H2O Ca

2+ +

2 OH- + H2 Ca H2O2 Al + 3 Br2 Al2Br6 Al Br2Mg + 2 H+ Mg2++

H2 Mg H+


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Metals act as reducing agents in theirchemical reactions. When copper is heatedover a flame, for example, the surfaceslowly turns black as the copper metalreduces oxygen in the atmosphere to formcopper(II) oxide.


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If we turn off the flame, and blow H2 gasover the hot metal surface, the black CuOthat formed on the surface of the metal isslowly converted back to copper metal. Inthe course of this reaction, CuO is reducedto copper metal. Thus, H2 is the reducingagent in this reaction, and CuO acts as anoxidizing agent.


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An important feature of oxidation-reductionreactions can be recognized by examiningwhat happens to the copper in this pair ofreactions. The first reaction converts coppermetal into CuO, thereby transforming a

reducing agent (Cu) into an oxidizing agent(CuO). The second reaction converts anoxidizing agent (CuO) into a reducing agent(Cu). Every reducing agent is therefore linked,or coupled, to a conjugate oxidizing agent, and

vice versa.

Every time a reducing agent loses electrons, itforms an oxidizing agent that could gainelectrons if the reaction were reversed.


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Conversely, every time an oxidizing agent gains

electrons, it forms a reducing agent that could loseelectrons if the reaction went in the oppositedirection.

The idea that oxidizing agents and reducing agents

are linked, or coupled, is why they arecalledconjugateoxidizing agents and reducingagents.Conjugatecomes from the Latin stemmeaning "to join together." It is therefore used todescribe things that are linked or coupled, such asoxidizing agents and reducing agents.


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The main group metals are all reducing agents.They tend to be "strong" reducing agents. Theactive metals in Group IA, for example, give up

electrons better than any other elements inthe periodic table.

The fact that an active metal such as sodium isa strong reducing agent should tell ussomething about the relative strength of the

Na+ ion as an oxidizing agent. If sodium metalis relatively good at giving up electrons,Na+ ions must be unusually bad at picking upelectrons. If Na is a strong reducing agent, theNa+ ion must be a weak oxidizing agent.


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Conversely, if O2 has such a high affinity forelectrons that it is unusually good at acceptingthem from other elements, it should be able to

hang onto these electrons once it picks themup. In other words, if O2 is a strong oxidizingagent, then the O2- ion must be a weakreducing agent.

In general, the relationship between conjugateoxidizing and reducing agents can be describedas follows. Every strong reducing agent (suchas Na) has a weak conjugate oxidizing agent(such as the Na+ ion). Every strong oxidizingagent (such as O2) has a weak conjugatereducing agent (such as the O2- ion).


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We can determine the relative strengths of apair of metals as reducing agents bydetermining whether a reaction occurs whenone of these metals is mixed with a salt of theother. Consider the relative strength of iron

and aluminum, for example. Nothing happenswhen we mix powdered aluminum metal withiron(III) oxide. If we place this mixture in acrucible, however, and get the reaction startedby applying a little heat, a vigorous reaction

takes place to give aluminum oxide and molteniron metal.

2 Al(s) + Fe2O3(s) Al2O3(s) + 2 Fe(l)


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By assigning oxidation numbers, we canpick out the oxidation and reduction halvesof the reaction.


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Aluminum is oxidized to Al2O3 in thisreaction, which means that Fe2O3 must bethe oxidizing agent. Conversely, Fe

2

O3

isreduced to iron metal, which means thataluminum must be the reducing agent.Because a reducing agent is alwaystransformed into its conjugate oxidizingagent in an oxidation-reduction reaction,the products of this reaction include a newoxidizing agent (Al2O3) and a new reducingagent (Fe).


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Since the reaction proceeds in this direction, itseems reasonable to assume that the starting

materials contain the stronger reducing agentand the stronger oxidizing agent

In other words, if aluminum reduces Fe2O3 toform Al2O3 and iron metal, aluminum must be astronger reducing agent than iron.


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We can conclude from the fact that aluminumcannot reduce sodium chloride to form sodiummetal that the starting materials in this reactionare the weaker oxidizing agent and the weaker

reducing agent.

We can test this hypothesis by asking: Whathappens when we try to run the reaction in theopposite direction? (Is sodium metal strong enough

to reduce a salt of aluminum to aluminum metal?)When this reaction is run, we find that sodiummetal can, in fact, reduce aluminum chloride toaluminum metal and sodium chloride when thereaction is run at temperatures hot enough to meltthe reactants.


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3 Na(l) + AlCl3(l) 3 NaCl(l) + Al(l)

If sodium is strong enough to reduce

Al3+ salts to aluminum metal and aluminumis strong enough to reduce Fe3+ salts toiron metal, the relative strengths of thesereducing agents can be summarized as

follows.

Na > Al > Fe


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Electrochemical reactions involve the transfer ofelectrons. Mass and charge are conserved when balancingthese reactions, but you need to know which atoms areoxidized and which atoms are reduced during thereaction. Oxidation numbers are used to keep track ofhow many electrons are lost or gained by each atom.

These oxidation numbers are assigned using the followingrules:

The convention is that the cation is written first in aformula, followed by the anion.For example, in NaH, theH is H-; in HCl, the H is H+.

The oxidation number of a free element is always 0.The

atoms in He and N2, for example, have oxidation numbersof 0.

The oxidation number of a monatomic ion equals thecharge of the ion.For example, the oxidation number ofNa+ is +1; the oxidation number of N3- is -3.


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Exceptions include OF2, since F is more electronegativethan O, and BaO2, due to the structure of the peroxide ion,which is [O-O]2-.

The oxidation number of a Group IA element in acompound is +1.

The oxidation number of a Group IIA element in acompound is +2.

The oxidation number of a Group VIIA element in acompound is -1, except when that element is combinedwith one having a higher electronegativity.The oxidationnumber of Cl is -1 in HCl, but the oxidation number of Cl is+1 in HOCl.

The sum of the oxidation numbers of all of the atoms ina neutral compound is 0.

The sum of the oxidation numbers in a polyatomic ion isequal to the charge of the ion.For example, the sum ofthe oxidation numbers for SO4

2- is -2.
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