Transcript Redox Reactions - Solon City Schools

Chemical
Reactions
Non-Redox Reactions
(Double Replacement)
Redox Reactions
Precipitation
2 solutions  Solid ppt
Synthesis
A + B  AB
Neutralization
Acid + Base  salt + H2O
Decomposition
AB  A + B
Single Replacement
A + BC  B + AC
Combustion
CxHy + O2  CO2 + H2O
Redox Reactions
Redox Equations:
At the conclusion of our time together,
you should be able to:
1. Define redox
2. Figure out oxidation numbers for any
element
3. Show the change in oxidation numbers in a
reaction
What’s the Point ?
REDOX reactions are important in …
• Purifying metals
(e.g. Al, Na, Li)
C3H8O + CrO3 + H2SO4 
Cr2(SO4)3 + C3H6O + H2O
• Producing gases
(e.g. Cl2, O2, H2)
• Electroplating metals
• Electrical production (batteries, fuel cells)
• Protecting metals from corrosion
• Balancing complex chemical equations
• Sensors and machines (e.g. pH meter)
What is Redox? – Page 27
 Involves a transfer of electrons.
 Redox reactions have a change in charge.
 REDOX stands for REDuction/OXidation
 Oxidation refers to a loss of electrons
 Reduction refers to a gain of electrons
 As a mnemonic remember LEO says GER
• Lose Electrons = Oxidation
• Gain Electrons = Reduction
Determination of Oxidation and Reduction
If oxidation # increased; substance oxidized

(reducing agent)
If oxidation # decreased; substance reduced

(oxidizing agent)
Review of Oxidation Numbers – Page 28
1. Write down complete reaction.
2. Assign an oxidation number to each element.
 We will see that there is a simple way to keep track
of oxidation and reduction
Oxidation Rules
 Any element by itself is zero.
 Oxidation # for ionic compounds are the same as the




charge.
Group 1 metals = +1, Group 2 metals = +2, F = -1.
O is -2, H = +1, and Cl = -1 in compounds.
The sum of oxidation numbers in a compound =
zero, sum of oxidation numbers in an ion = the
charge.
An ion by itself has that oxidation number. Na+ = 1+
The Remainder of Redox
 Step 3: Determine which element is being reduced.
This element’s oxidation # will decrease. Connect
them and mark # of electrons gained.
 Step 4: Determine which element is being oxidized.
This element’s oxidation # will increase. Connect
them and mark # of electrons lost.
Example
 0
 2K
+
0
F2

+1 -1
2KF
+
0
O2
+4 -2

CO2 +
 -4 +1
 CH4
+1 -2
H2O
Examples 1&2 – Page 29
 Al + O2  Al2O3
 FeO  Fe + O2
Example 3
 Mg + PbCl4  Pb + MgCl2
Review of Oxidation Numbers
1. Any element, when not combined with atoms of a
different element, has an oxidation # of zero.
(O in O2 is zero, Na by itself is zero)
2. Any simple monatomic ion (one-atom ion) has an
oxidation number equal to its charge
(Na+ is +1, O2– is –2)
3. The sum of the oxidation numbers of all of the
atoms in a formula must equal the charge written
for the formula. (if the oxidation number of O is –2,
then in CO32– the oxidation number of C is +4)
Review of Oxidation Numbers
4. In compounds, the oxidation # of IA metals is +1,
IIA is +2, IIIA is +3, Zn & Cd is +2, Ag is +1.
5. In ionic compounds, the oxidation # of a
nonmetal or polyatomic ion is equal to the charge
of its associated ion.
(MgCl2, Mg is +2, therefore Cl is –1)
6. F is always –1, O is always –2
(unless combined with F),
H is usually +1, except when it is
bonded to metals in binary compounds.
(ex. NaH, H oxidation # is –1 or when it’s in
elemental form H2, oxidation # is 0).
Example 4
 Ba(NO3)2 + Na3PO4  Ba3(PO4)2 + NaNO3