Transcript Reaction types summary

Classification of reactions
-
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Synthesis
Decomposition
Neutralisation
Precipitation
Double decomposition
Reduction Oxidation
Synthesis
Synthesis is the making of a compound by the
combination of simpler substances.
The term is also generally used to describe the
making of a new substance.
Example:
Iron + sulphur  Iron sulphide
Fe
+ S
 FeS
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Decomposition
The breaking down of one substance into simpler
substances.
Example:
Calcium carbonate  calcium oxide + carbon dioxide
CaCO3

CaO
+
CO2
This is often carried out by heating.
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Neutralisation
The reaction between an acid and a base to produce
a neutral salt and water.
Sulphuric acid + sodium hydroxide  Sodium sulphate + water
H2SO4
+
2NaOH

Na2SO4 + 2H2O
Only the hydrogen ions from the acid combine with the
hydroxide ions from the base. All of the other ions start and
end the reaction in the same ionic form. They are called
spectator ions (they “watch” the other ions reacting)
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Precipitation
This is the reaction between two solutions resulting
in the formation of an insoluble substance, which
consequently appears as a solid within the liquid.
It is apparent that a solid has been formed as the
liquid becomes turbid (not wholly transparent)
Example:
Copper sulphate + sodium hydroxide  copper hydroxide + sodium sulphate
The copper hydroxide is insoluble and appears as a
precipitate.
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Double decomposition
This is really a precipitation and is given the name
double decomposition to indicate that the ions have
“switched round”. In the previous example the
hydroxide ions seem to go from the sodium to the
copper and the sulphate ions seem to go from the
copper to the sodium.
In reality ions in solution are not related to one
another and so no such movement takes place. The
collision of the copper and the hydroxide ions
within the solution forms the insoluble substance,
the remaining ions are simply spectators and named
as such. To collect the sodium sulphate it would be
necessary to remove the water
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Oxidation-Reduction: A Reaction
Oxidation: When a substances loses electrons.
Reduction: When a substance gains electrons.
Consider:
Ca(s) + 2H+(aq)  Ca2+(aq) + H2(g).
The neutral Ca(s) has lost two e- to 2 H+ to become Ca2+.
We say Ca has been oxidized to Ca2+
At the same time 2 electrons are gained by 2 H+ to form H2 .
We say H+ is reduced to H2 .
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Redox Reaction with Air
Consider the reaction of Ca with O2:
2Ca(s) + O2(g)  2CaO(s)
Ca is easily oxidized in air.
On the left there is shiny Ca metal.
On the right we see a white powder – Calcium oxide.
Again, Ca(s) loses electrons and is oxidized to Ca+2
And the neutral O2 has gained electrons from the Ca to
become O2- in CaO.
We say O2 has been reduced to O2-.
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Electron Transfer and Terminology
Lose electrons: Oxidation
Gain electrons: Reduction.
GER
Leo
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It Takes Two: Oxidation-Reduction
In all reduction-oxidation (redox) reactions, one species is
reduced at the same time as another is oxidized.
Oxidizing Agent:
The species which causes oxidation is called the oxidizing agent.
The substance which is oxidized loses electrons to the other.
The oxidizing agent is always reduced
Reducing Agent:
The species which causes reduction is called the reducing agent.
The substance which is reduces gains electrons from the other.
The Reducing agent is always oxidized
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Oxidation of Metals with Acids
It is common for metal to produce hydrogen gas when
they react with acids. For example, the reaction
between Mg and HCl:
Mg(s) + 2HCI(aq)  MgCl2(aq) + H2(g) .
In this reaction, Mg is oxidized and H in HCl is
reduced.
Note the change in oxidation state for these species:
Mg0  Mg+2 in MgCl2
&
H+ in HCl  H0 in H2
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Redox reaction with Acid
It is possible for metals to be oxidized with salt:
Fe(s) + Ni(NO3)2 (aq)  Fe(NO3)2 (aq) + Ni (s) .
Molecular Equation
The net ionic equation shows the redox chemistry well:
Fe(s) + Ni+2(aq)  Fe2+(aq) + Ni (s)
Net ionic Equation
In this reaction iron has been oxidized to Fe2+ while the Ni+2
has been reduced to Ni0.
What determines whether the reaction occurs ?
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Competition For e- Transfer
Consider: Na, Mg, Al,
Metallic character decreases left to right.
Metal tend to give up electrons.
Now consider the reaction:
Na + AlCl3  ??? (NaCl + Al)
To determine if the reaction occurs, the question is to determine which
metal has a greater affinity for electrons (or which is willing to lose e- ).
Na is more willing to lose e- than Al
Al is more willing to accept e- (less metallic)
Conclude: The reaction occurs.
3Na + AlCl3  3NaCl + Al
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Summary
Redox OxidationReduction-
Oxidation/Reduction reaction
Lose electron (LEO)
Gain electron (GER)
Activity Series- Table showing elements’ relative
ease of oxidation.
Active Metal Metal which prefers to lose e- and
therefore prefers the oxidized form.
Noble Metal Metal which do not lose e- and
therefore prefers the zero state.
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