Transcript Half-Reactions

Assigning Oxidation Numbers
Rules & Half-Rxn
8 Rules for Oxidation Numbers
1. of a free, uncombined element = 0.
Na He
O2
N2
S8
Cl2
2. of a monatomic ion = charge on
ion.
+2
-1
+3
Ca
= +2. Cl = -1. Al
P
= +3.
Remember: Ions occur in ionic
compounds: CaCl2, Al(NO3)3, etc.
8 Rules for Oxidation Numbers
3.
Fluorine is always -1.
CF4
4. Hydrogen is nearly always +1,
except when it’s bonded to a
metal. Then it’s -1.
H2O, HNO3, H2SO4
LiH CaH2 NaH
8 Rules for Oxidation Numbers
5.
Oxygen is nearly always -2 except
when it’s…
-Bonded to fluorine, where O is +2 OF2
-In the peroxide ion, where O is -1.
eg. H2O2 and Na2O2
O22-
8 Rules for Oxidation Numbers
6. The sum of oxidation numbers in a
neutral compound is 0.
H2O CO2 NO SO3
7. The sum of oxidation numbers
in a polyatomic ion = charge of
the ion.
Sum in SO42- = -2. Sum in NO3- = -1.
8 Rules for Oxidation Numbers
8. In covalent compounds, the
oxidation number of the more
electronegative atom is the
negative charge it would have if
it was an ion.
*NH3: N = -3, H = +1.
SiCl4: Si = +4, Cl = -1.
+4
+3
+2
+1
0
-1
-2
-3
-4
2) And if
you’re lucky
you strike oil &
it shoots up
1) You dig down
with an oil rig
Assign Oxidation Nos
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KCl
K = +1, Cl = -1
CaBr2 Ca = +2, Br = -1
CO
C = +2, O = -2
C = +4, O = -2
CO2
Al(NO3)3
Al = +3, O = -2, N = +5
Na3PO4
Na = +1, O = -2, P = +5
H2S
H = +1, S = -2
NH4+1
N = -3, H = +1
SO3-2
S = +4, O = -2
Electrons are Negative!
• Why do we use the word “reduced” when
electrons are gained?
Look at how the oxidation number
changes.
For example, if Cl gains an electron it
becomes Cl-1. The oxidation number
decreased from 0 to -1. The oxidation
number was reduced.
Agents
• When an element is
OXIDIZED it caused
something else to be
reduced so it is called
the REDUCING AGENT.
• Likewise, when an
element is REDUCED it
caused something else
to be oxidized so it is
called the OXIDIZING
AGENT.
Non-metals
metals
Vocabulary Interlude
• Oxidizing Agent: Is itself reduced.
Accepts electrons from something else –
aids oxidation for another species.
• Reducing Agent: Is itself oxidized.
• Loses electrons to something else – aids
reduction for another species.
Reduction Half-Reactions
Electrons are gained so
• I2 + 2e-  2Ithey
are
like
a
reactant!
• O2 + 4e-  2O-2
• Half-reactions must demonstrate
conservation of mass & conservation of
charge.
• # of atoms of each element on LHS equals “
“ “
“ “
“
“ RHS.
• Total charge on LHS = Total charge on RHS
Oxidation Half-reactions
• K  K1+ + 1e• Fe2+  Fe3+ + 1e• Cu  Cu2+ + 2e-
Electrons are lost so
they appear on the
product side!
• Total Charge on LHS = Total Charge on RHS
• # atoms LHS = # atoms RHS