Download Oxidation-Reduction (Redox) Reactions

Oxidation-Reduction (Redox) Reactions
Reactions in which electrons are transferred from one reactant
to another
oxidation – the loss of electrons
reduction – the gain of electrons
Oxidation of zinc:
“LEO the lion
says GER”
Zn( s ) → Zn 2+ (aq) + 2e −
solid zinc has lost 2 electrons to form zinc ions;
zinc is oxidized
Reduction of copper ions:
Cu 2+ (aq) + 2e − → Cu ( s )
copper ions have gained 2 electrons to form solid copper;
copper ions are reduced
Redox Reactions
What happens if solid zinc is placed in a solution containing
copper ions?
Redox Reactions
What happens if solid zinc is placed in a solution containing
copper ions?
In this case, one species was able to be oxidized while
the other was able to be reduced
It’s important to have both species present; otherwise
where would the electrons come from?
Let’s look at the equations again:
Zn( s ) → Zn 2+ (aq) + 2e −
Cu 2+ (aq) + 2e − → Cu ( s )
oxidizing agent – a species that accepts electrons
(Cu2+)
reducing agent – a species that donates electrons
(Zn)
Redox Reactions
Oxidation and reduction reactions can be written separately
(half-reactions), but they must always occur together
The sum of the two half-reactions gives the overall redox equation:
Zn( s ) → Zn 2+ (aq) + 2e −
Cu 2+ (aq) + 2e − → Cu ( s )
Zn( s ) + Cu 2+ (aq) + 2e − → Zn 2+ (aq) + Cu ( s ) + 2e −
Zn( s ) + Cu 2+ (aq) → Zn 2+ (aq ) + Cu ( s )
Redox Reactions
What if solid copper was placed in a solution containing zinc
ions? Would the reverse reaction occur?
No! No reaction would occur at all!
Cu ( s ) + Zn 2+ (aq) → no reaction
We can use an activity series (see
Table 4.5) to predict which redox
reactions will occur
Any metal in the activity series will
only be oxidized by metals appearing
below it
Zinc can be oxidized by copper, but
copper can NOT be oxidized by zinc
Redox Reactions
Which of the following combinations will result in a redox
reaction?
Ca ( s ) + Al 3+ (aq) → ?
Cr ( s ) + Mn 2+ (aq) → ?
Cu ( s ) + Fe 2+ (aq) → ?
Ba( s ) + Co 2+ (aq) → ?
Pb( s ) + Na + (aq) → ?
Balancing Redox Reactions
Let’s consider the following reaction:
Ba( s ) + Co 2+ (aq) → ?
Write each half-reaction and then add them together
to generate the net reaction:
2+
Ba( s ) → Ba (aq) + 2e
−
Co 2+ (aq) + 2e − → Co( s )
Ba( s ) + Co 2+ (aq) + 2e − → Ba 2+ (aq) + Co( s ) + 2e −
Ba( s ) + Co 2+ (aq) → Ba 2+ (aq) + Co( s )
Match these
terms:
oxidation
reduction
oxidizing
agent
reducing
agent
Balancing Redox Reactions
Let’s consider the following reaction:
Ca ( s ) + Al 3+ (aq) → ?
Write each half-reaction and then add them together
to generate the net reaction:
Ca ( s ) → Ca 2+ (aq) + 2e −
Al 3+ (aq ) + 3e − → Al ( s )
Ca ( s ) + Al 3+ (aq) + 3e − → Ca 2+ (aq ) + Al ( s ) + 2e −
Electrons must be balanced in redox reactions!
Balancing Redox Reactions
Let’s consider the following reaction:
Ca ( s ) + Al 3+ (aq) → ?
3
2
(Al
(Ca(s) → Ca
3+
2+
(aq) + 2e −
(aq) + 3e − → Al ( s )
)
)
Match these
terms:
oxidation
reduction
3Ca ( s ) + 2 Al 3+ (aq) + 6e − → 3Ca 2+ (aq) + 2 Al ( s ) + 6e −
3Ca ( s ) + 2 Al 3+ (aq) → 3Ca 2+ (aq) + 2 Al ( s )
oxidizing
agent
reducing
agent
Redox Reactions
Sometimes redox reactions involve combination of reactants:
Na ( s ) → Na + + e −
Cl2 ( g ) + 2e − → 2Cl −
Match these
terms:
Na ( s ) + Cl2 ( g ) + 2e − → Na + + e − + 2Cl −
oxidation
reduction
2
2
2
Electrons must be balanced in redox reactions!
Na ( s ) → Na + + e −
Cl2 ( g ) + 2e − → 2Cl −
2 Na ( s ) + Cl2 ( g ) + 2e − → 2 Na + + 2e − + 2Cl −
2 Na ( s ) + Cl2 ( g ) → 2 NaCl ( s )
oxidizing
agent
reducing
agent
Oxidation Numbers
A complete transfer of electrons never occurs for combination reactions
which form molecular compounds
H 2 ( g ) + F2 ( g ) → 2 HF ( g )
No ions are shown in this reaction.
How can we tell which element is oxidized and which is reduced?
Oxidation numbers provide a method of “bookkeeping” by determining
the charge an atom would have if electrons were transferred completely
The oxidation number of a free element (monatomic or diatomic) is 0
The oxidation number of fluorine is always -1 (except as a free element)
The oxidation number of hydrogen is always +1 (when combined with
nonmetals) or -1 (when combined with metals)
Oxidation Numbers
Assign oxidation numbers to all reactants and products:
The oxidation number of a free element (monatomic or diatomic) is 0
The oxidation number of fluorine is always -1 (except as a free element)
The oxidation number of hydrogen is always +1 (when combined with
nonmetals) or -1 (when combined with metals)
The oxidation number of a monatomic ion is the charge on the ion
H 2 ( g ) + F2 ( g ) → 2 HF ( g )
0
0
+1 -1
The oxidation number of hydrogen increases from 0 to +1 Æ oxidation
The oxidation number of fluorine decreases from 0 to -1 Æ reduction
Oxidation Numbers
More oxidation number rules:
The oxidation number of a Group 1A or 2A metal is +1 or +2 (except as
free elements)
The oxidation number of oxygen is almost always -2 (exceptions can
occur with compounds involving fluorine, hydrogen or Group 1A
or 2A metals)
NaH
SO2
+1 -1
? -2
? -4
Use circles to represent the oxidation
number of an atom
Use squares to represent the total
contribution from that type of atom
The sum of all oxidation numbers within a compound or
ion must be equal to the charge on the compound or ion
Oxidation Numbers
Assign oxidation numbers for the atoms without rules:
SO2
? -2
To make the overall compound neutral,
the oxidation number for sulfur (overall)
must be +4
+4 -4
SO2
+4 -2
+4 -4
There is only 1 sulfur atom, so it must be +4
Oxidation Numbers
Let’s try another example:
CO3−2
Each oxygen is -2
? -2
There are three oxygen atoms, so the
total contribution is -6
? -6
CO3−2
+4 -2
+4 -6
The ion has an overall charge of -2, so carbon
must be +4 (-6 + 4 = -2)
Oxidation Numbers
Let’s try another example:
N 2O5
Each oxygen is -2
? -2
There are five oxygen atoms, so the
total contribution is -10
? -10
N 2O5
? -2
The molecule has an overall charge of 0,
so nitrogen must contribute a total of +10
+10 -10
N 2O5
+5 -2
+10 -10
There are two nitrogen atoms, so each
must contribute +5
Oxidation Numbers
Assign the terms oxidation, reduction, oxidizing agent and reducing
agent to the following redox reaction:
Fe( s ) + PtCl2 (aq) → FeCl2 (aq) + Pt ( s )
0
+2 -1
+2 -1
0
0
+2 -2
+2 -2
0
The oxidation number of iron increases from 0 to +2 Æ oxidation
The oxidation number of platinum decreases from 0 to -2 Æ reduction
The oxidation number of chlorine remains constant Æ spectator
oxidizing agent – a species that accepts electrons
(Pt+2)
reducing agent – a species that donates electrons
(Fe)
Oxidation Numbers
Assign the terms oxidation, reduction, oxidizing agent and reducing
agent to the following redox reaction:
Cr ( s ) + AuCl3 (aq) → CrCl3 (aq) + Au ( s )