oxidation numbers positive oxidation number negative oxidation number - loses partial or total...
TRANSCRIPT
Oxidation Numbers
Positive oxidation number
Negative oxidation number
- Loses partial or total control of electrons in a bond
- Gains partial or total control of electrons in a bond
Example
Mg in the +2 state (Mg+2)
Has lost partial or total control of 2 e-
- Used to tell how many electrons an atom has lost or gained in a chemical reaction
Redox - When one atom loses an electron another must receive it
Na Na+ + 1e-
Cl + 1e- Cl-
1) Oxidation - The lose of an electron by an atom
- Causes an increase in oxidation number
2) Reduction - The gain of an electron by an atom
- Causes a decrease in oxidation number
If Fe loses three electrons, we write this as
Fe Fe+3 +3e-
If the electrons appear on the products side, they are given off
oxidation
If Fe+3 gains three electrons, we write this as
Fe+3 +3e- Fe If the electrons appear on the reactants side, they are gained
reduction
Lose Electrons Oxidation
Gain Electrons ReductionLEO goes GER
L E O
G E R
Indicate if the atoms shown have lost or gained electrons
Then write down if the atom is oxidized or reduced
A. H H+ + e-
B. Cl2 + 2e- 2 Cl-
C. Fe+2 + 2e- Fe
D. Cl+5 + 6e- Cl-
E. S-2 S+4 + 6 e-
Lost or gained e- Oxidized or reduced
Lost
Gained
oxidized
reduced
Lost
Gained
Gained
oxidized
reduced
reduced
Oxidizing Agent - Causes another substance to be oxidized- Is reduced
Ex. Sn+4 + 2 e- Sn+2
Sn+4 is the oxidizing agent
Reducing Agent- Causes another substance to be reduced
- Is oxidized
Ex. Na Na+1 + 1 e-
Na is reducing agent
Finding Oxidation States
The most common oxidation numbers are in the upper right corner of your periodic table
Pb
+2+4
Oxidation states
Using your periodic table, list the oxidation states of the following
Ca
Na
O
N
Examples
+2
+1 -3, -2, -1 and a bunch more!
-2
Rules for assigning oxidation numbers
1. Uncombined elements have an oxidation number of zero
Cu0 Mg0 S0
2.) The oxidation state of an ion is the same as its charge
Oxidation state of zero means the atom is not losing or gaining any electrons
Cu+ Mg+2 S-2
+1 oxid. state +2 oxid. state -2 oxid. state
Cl20
3) Group 1 metals in a compound always have a +1 oxidation state
4) Group 2 metals in a compound always have a +2 oxidation state
MgCl2 LiBr Na2O+2 +1+1
5) Hydrogen always has a +1 oxidation state in a compound
Exception - When H is attached to a group 1 or 2 metal, it has a -1 oxidation state
HCl H2O+1 +1
NaH CaH2
-1-1
6) Oxygen in a compound always has a -2 oxid. state
Exception - In a peroxide, oxygen becomes -1
Na2O2 H2O2
-1 -1
H2O HNO3
-2 -2
7) Halogens are usually -1 in compounds
SO4-2
+6 -2
-8+6 = -2
The sum of the one sulfur and the 4 oxygens should be -2
8) The sum of all oxidation numbers in a polyatomic ion must equal the charge of the ion
Examples - Assign the oxidation numbers to all elements in the following ions
NO3- CrO4
-2+6
-6+5
+5 -2-2
+6 -8= -1 = -2
Assigning Oxidation Numbers
- All of the oxidation numbers of all atoms in a neutral compound must add up to zero
For example, in H2O,
H2O+1 -2
The two hydrogens add up to +2,
+2 + -2 = 0
Since water is a neutral compound, all oxidation states must add up to zero
Neutral compounds
Practice - Assign oxidation numbers to the following
MgCl2 LiOH Na2S H2SO4 NaNO3
+2 -1 +1 -2 +1 +1 -2 +1 +6 -2 +1 +5 -2
+2 -2 +1 +1-2 +2 -2 +2 -8+6 +1 +5 -6
0 0 0 0 0
Cu(NO3)2 NaH NO2
+2 +5 -2 +1 -1 +4 -2
-12+10+2 +1 -1 +4 -4
Recognizing Redox Equations
- Not all reactions are redox
- Double replacement reactions are NOT
- Single replacement reactions are
- If a reaction is redox then the oxidation numbers of some of the elements must be different on either side of the equationIs the reaction redox?
LiOH + HCl H2O + LiCl
2H2 + O2 2 H2O
Mg + CuSO4 MgSO4 + Cu
In the reaction Mg + Cl2 --> MgCl2
There are 2 steps occurring
Mg loses 2 electrons
Each chlorine gains an electron
MgCl
Cl
Electrons are moving. If e- could move through a wire, this would be an electric current
Electrochemistry
Half reaction- Either the reduction or the oxidation portion of a redox reaction
Al+3 + 3e- --> Al
Ca --> Ca+2 + 2e-
reduction
oxidation
The reduction ½ reaction shows an atom or ion gaining e- and the oxidation number decreasing
The oxidation ½ reaction shows an atom or ion losing e- and the oxidation number increasing
Steps in Writing Half Reactions
2 Li + CaBr2 --> 2 LiBr + Ca1. Assign oxidation numbers
+2 -1-1 +10 0
2. For each atom that changes its state, write down the starting and ending oxidation states
3. Add electrons to balance the equationCa+2 + 2e- Ca0 Li0 Li+ + e-
-Notice that each set of half reactions contains oxidationand reduction
-For each half reaction, both sides have the same charge
OxidationReduction
Ca+2 --> Ca Li --> Li+
Which reaction shows conservation of charge?
Fe(s) Fe 2+ (aq) + e- Fe + 2 e- Fe 2+
Fe + 2 e- Fe 3+
Fe(s) Fe 2+ (aq) + 2 e-
Examples
Write out the half reactions
A. KBr + Na K + NaBr
B. Sr + MgO Mg + SrO
C. NiCl + CuO NiO + CuCl
K+ +e- --> K
Na --> Na+ + e-
Sr --> Sr+2 + 2e-
Mg+2 + 2e- --> Mg
Ni+1 Ni+2 + 1 e-
Cu+2 + 1e- Cu+1
Mg + CuSO4 MgSO4 + Cu
Write the ½ reaction representing oxidation.
Given the reaction
Write the reduction ½ reaction for the following
2 Al + 3 Cu+2 2 Al 3+ + 3 Cu
Mg Mg+2 + 2 e-
Cu+2 + 2e- Cu
For the following
Zn + Cr 3+ Zn 2+ + Cr
a.) Write the ½ reaction for the reduction.
b.) Write the ½ reaction for the oxidation
c.) Which species loses electrons?
d.) What happens to the number of protons in a Zn atom when it changes to Zn 2+ as the redox reaction occurs.
If one of the diatomic elements move to the zero oxidation state, we must write the reaction with two atoms
ExampleWe cannot write
Cl-1 --> Cl + e-
We must write
2 Cl- --> Cl2 + 2e-
Other examples
Br2 + 2 e- --> 2 Br-
2 H+ + 2e- --> H2
B. Electrochemical CellsElectrochemical cell - Produces electricityelectricity from a divided redox reaction
- Electrons are moved from one atom to another through a wire
Voltaic Cell -(battery)- A spontaneous chemical reaction produces a flow of e-
Anode - Is the metal in a voltaic cell that is oxidized ( It will dissolve)
Cathode - The metal in a voltaic cell where reduction will occur
- Positive ions in solution will be reduced an collect on the cathode
External conductor (Wires) -permit flow of e-
Salt Bridge - U-tube containing electrolytic solution of + and - ions
- Allows migration of ions
- Keeps compartments neutral
These two jars are called ½ cells
How to identify the anode and cathode and the direction of e- flow
1. Find out which metal is oxidized / reduced
An Ox Anode is Oxidized
Table J: Best RA will lose e- so is the oxidizing ½ reaction
Red Cat Reduction happens at the Cathode
A reaction will be spontaneous if a pure metal is on the top, and the ion is below
Examples - Determine which atom/ion pairs will react spontaneously
A. Cu+2 and Ni
B. Mn+3 and Zn
C. K and Ni+2
D. Co and Al+3
YES
NO
YES
NO
As long as the solid metal is on top, the redox will be spontaneous.
2. Write ½ reactions
LEO - Loss of electrons is oxidation
GER - Gain of electrons is reduction
LEO
An Ox
GER
RED CAT
Mg0 Mg+2 + 2e- Ni+2 + 2e Ni0
3) Choose sign for anode / cathode
- ( e- flow from – to + )
- (-) lost from anode gained by cathode (+)
- e- flow from Mg to Ni
As the reaction continuesMg Mg+2 + 2e- Ni+2 + 2e- Ni
Anode Cathode
Will decrease in mass Will increase in mass
Problem - Solutions must always be electrically neutral
As soon as we add -, we must add + chargesZn
Zn+2 Cu+2
Cu
SO4-2 SO4
-2
Each solution has a total charge of zero (neutral)
If we want to add more Zn+2 to solution, we need to add some - ions
If we want to remove some Cu+2 from solution, we need to add some + charges
d. Salt Bridge Tube containing a salt solution
Na+ Cl-
Porous plugs allow ions to move out of the tube
As Zn+2 goes into solution, 2 Cl- come out of the salt bridge
As Cu+2 is removed from the solution, 2 Na+ come out of the salt bridge
Solutions are kept neutral
Zn
Zn+2
Cu+2
Cu
Cu+2
Zne-e-
2e-
2e- 2e- 2e- 2e-
2e-
SO4-2 SO4
-2
SO4-2
Na+Na+ Cl- Cl-
Zn+2
Cl-Cl-
Cl-Cl-
Cl-Cl-
Now the Zn+2 solution is neutral
Na+Na+
Na+Na+
Na+Na+
Now the Cu+2 solution is neutral
E. Electrolytic Cell
Electrolytic cell An electric current is used to force a chemical reaction to occur
Electrolysis A chemical reaction which uses electricity to break apart a compound – nonspontaneous
1. Description
Only one cell is needed
Cell contains a power source
Battery pulls electrons off of one electrode
And puts them on another electrode+e-
+-
e-
e-e-
e-
e--
The battery decides what is + and - electrodes
If we add melted NaCl (MOLTEN NaCl) to the cell
Na+ and Cl- will be free to move
Na+ move to the negative electrode and gains electrons
This is reduction The negative electrode is the cathode
Cl- move to the positive electrode and loses electrons
This is oxidation The positive electrode is the anode
+
+-
-Na+ Cl-
Na+ + e- --> Na 2 Cl- --> Cl2 + 2e-
reduction
CATHODE
oxidation
ANODE
For electrolytic cells, the charges are opposite from galvanic cells
The “red cat” and “an ox” statements still work
2. Electroplating- Electrolytic cell is used to produce pure metals (Na, Mg)
Electroplating Layering a metal onto a surface using an electrolytic cell
Cell looks the same (one cell with a power source
+
+-
-
Put the object to be plated on the – electrode ( cathode)Put the layering metal on the + electrode
AgThis is supposed to be a copper ring!
+
+-
-Ag
The battery will pull electrons off of the Ag, turning it into Ag+
Ag --> Ag+ + e-
The Ag+ goes into solution and the e- go through the wire
e-
e-
e-e-
e-
e- Ag+
The Ag+ now move over to gain its own e- on the ring
The Ag+ gets metallically bonded to the copper
Ag+ + e- --> Ag
and reduction occurs at the cathode, which is negative
Again, oxidation occurs at the anode, which is positive
Reduction Oxidation
In electroplating, the same atom is oxidized and then reduced