Transcript Chemistry -- Oxidation

Chemistry
Oxidation-Reduction
(Redox)
I. Introduction
Acids donate _____ and bases accept ____
H+
H+
proton(s)
proton(s)
In other reactions, substances donate or
accept electrons
The flow of electrons from one substance to
the next is how batteries “make” electricity
II. Oxidation Numbers
Assigning oxidation numbers to substances
enable chemists to keep track of the
“transfer” of electrons
Review:
Ionic compounds: give and take e’s
Covalent compounds: share e’s
In ionic compounds – e’s are really
transferred
In covalent compounds – e’s are shared
unevenly but we pretend that the most
electronegative atom gets all the electrons
Oxidation Number: The charge that an atom
in a molecule would develop if the most
electronegative atom in the molecule took
the shared electrons (if the e transfer was
complete)
Let’s calculate ox #’s for HCl…
Chlorine is more electronegative
We pretend it takes all the e’s
That leaves H’s ox # as +1
Cl ox # as -1
III. Rules for Assigning Oxidation
Numbers
1. The oxidation number of any pure element is
zero. Thus the oxidation number of H in H2 is
zero.
2. The oxidation number of a monatomic ion is
equal to its charge. Thus the oxidation number
of Cl in the Cl- ion is -1, that for Mg in the Mg 2+
ion is +2, and that for oxygen in O2- ion is -2.
3. The sum of the oxidation numbers in a
compound is zero if neutral, or equal to the
charge if an ion.
4. The oxidation number of alkali metals in
compounds is +1, and that of alkaline earths in
compounds is +2. The oxidation number of F is
-1 in all its compounds.
5. The oxidation number of H is +1 in most
compounds. Exceptions are H2 (where H = 0)
and the ionic hydrides, such as NaH (where H =
-1).
6. The oxidation number of oxygen (O) is -2 in
most compounds. Exceptions are O2 (where O
= 0) and peroxides, such as H2O2 or Na2O2,
where O = -1.
• For other elements, you can usually use If no
other rules apply, assume ON is the same as the
charge taken on in an ionic compound (“the
charge it would like to be)
Examples
1. What are the following ON’s?
A. Iron in Fe?
0
B. Iron in Fe3+?
+3
C. Br in MgBr2?
-1
D. Br in HBrO3
+5
2. What is the ON of each atom in the
following?
A. H2
0
B. Ca(OH)2
Ca: +2 O: -2 H: +1
C. F2CO
F: -1
C: +4 O: -2
D. NaSO4Na: +1 S: +6 O: -2
E. C2H4
C: -2 H: +1
F. NCl3
N: +3 Cl: -1
G. MgCoBr4
Mg: +2 Co: +2 Br: -1
OYO’s
16.1 What is the oxidation number of N in each of
the following substances?
A. N
B. N3C. NO316.2 What are the oxidation numbers of all atoms
in the following compounds?
A. LiNH2 B. N2H2 C. Ca(NO2)2
D. CO2
E. BF4F. PO43G. ClNO
H. S8
III. Oxidation and Reduction
Oxidation: The process by which an atom loses an electron
or electrons
Reduction: The process by which an atom gains an
electron or electrons
LEO goes GER
Lose Electrons–Oxidation Gain Electrons- Reduction
OIL RIG
Oxidation Is Loss (of electrons), Reduction Is Gain (of
electrons)
Examples
1. S atom goes from -2 to +6. Oxidized or
reduced? How many electrons?
2. C atom goes from -2 to -4. Oxidized or
reduced? How many electrons?
3. An atom goes from +5 to +3. Oxidized or
reduced? How many electrons?
4. An atom goes from -6 to -1. Oxidized or
reduced? How many electrons?
IV. Recognizing Redox Reactions
Mg(s) + 2HCl(aq)→ H2(g)
0
+1 -1
0
Mg is oxidized
H is reduced
+ MgCl2(aq)
+1 -1
NaOH(aq) + HCl(aq) → H2O(l) +NaCl(aq)
+1 -2 +1
+1 -1
+1 -2 +1 -1
NOT a redox reaction!
2Mg(s) + O2(g)→ 2MgO(s)
0
0
+2 -2
Mg is oxidized O is reduced
Examples – Redox, or not?
A. Ca(s) + Cl2(g) → CaCl2(s)
0
0
+2 -1
Ca oxidized Cl reduced
B. Ag+(aq) + OH-(aq) → AgOH(s)
+1
-2 +1
+1 -2 +1
not a redox
C. 3IF5(aq) + 2Fe(s) →2FeF3(aq) + 3IF3(aq)
+5 -1
0
+3 -1
+3 -1
Fe oxidized I reduced
Thionin – Two-Faced Solution
Thio+ + 2Fe2+ + 2H+ ↔ ThioH2+ + 2Fe3+
Purple
colorless
light activated
Methylene Blue
O2(g)
O2(dissolved) + methylene blue
colorless
→ O2(dissolved)
→ methylene blue
blue
Glucose + OH-
→ glucoside
Glucoside + methylene blue
→ methylene blue + OH-
blue
colorless
** When a redox reaction occurs, there will
always be atoms that are oxidized and
atoms that are reduced
The thing that is oxidized is the reducing
agent (provides electrons)
The thing that is reduced is the oxidizing
agent (takes electrons)
OYO’s
16.5 For each of the following, determine
whether or not it is a redox reaction. If it is
a redox, identify what was oxidized, what
was reduced, & the oxidizing and reducing
agents.
A. Cu2+(aq) + Zn(s) → Zn2+(aq) + Cu(s)
B. 2H2O → 2H2 + O2
C. NaCl + AgF → AgCl + NaF
D. Ca(s)+ 2HNO3(aq)→Ca(NO3)2(aq)+ H2(g)
IV. How Batteries Work
The fact that atoms can lose and gain electrons in
a reaction is the bases for how a battery works
Some reactions (like acid-base) involve the
rearrangement of electrons
Others (like formation, decomposition, combustion
and displacement reactions) involve the transfer
of electrons (redox)
Displacement reactions: an ion (or atom) in a
compound is replaced by an ion (or atom) of
another element
CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
Cu2+(aq) + Zn(s) → Zn2+(aq) + Cu(s)
+2
0
+2
0
How could you get the electrons that Zn
loses to Cu to create a current?
You’d have to connect them!
A simple battery
“salt bridge”
– Permits the migration of ions.These keep the
Zn2+ solution from becoming too positive (as
negative electrons leave), and the other solution
from becoming too negative (as negative
electrons arrive)
-- porous barrier: does not allow solutions to mix,
only permits ions to flow to both solutions
Anode and Cathode
Anode: where oxidation occurs
-- negative in galvanic cells
Cathode: where reduction occurs
-- positive in galvanic cells
As electrons travel through the wire,
electricity is created
(This schematic is reversed from the one
before.)
Examples
Draw a diagram of the Galvanic cell,
indicating the electron flow, labeling the
anode and cathode, and indicating the
positive and negative sides of the Galvanic
cell for the following reactions:
A. 2Ag+(aq) + Mg(s)→ 2Ag(s) + Mg2+(aq)
B. I2(aq) + Fe(s) → 2I-(aq) + Fe2+
OYO’s
Draw a diagram of the Galvanic cell,
indicating the electron flow, labeling the
anode and cathode, and indicating the
positive and negative sides of the Galvanic
cell for the following reactions:
16.6 2H+(aq) + Mn(s) → H2(g) + Mn2+(aq)
16.7 3Cl2(aq) + 2Al(s) → 2Al3+(aq) + 6Cl-(aq)
V. Real Batteries
A. The Dry Cell… (no aqueous substances)
Anode:
Zn(s)→ Zn2+(aq) + 2eCathode:
0
+2
2NH4+ (aq) + 2MnO2(s) → Mn2O3(s) + 2NH3(aq) + H2O(l)
+4
+3
B. The Mercury Battery
Zn(Hg) + 2OH-(aq)→ ZnO(s) + H2O(l) + 2e0
+2
Cathode: HgO(s) + H2O(l) + 2e- → Hg(l) + 2OH-(aq)
+2
0
Anode:
Why do batteries stop working?
One of the reactants is used up.
The salt bridge degrades.
Rechargeable batteries
They “charge” by reversing the reactions to
re-create the original substances
Eventually stop working because the salt
bridge degrades
B. The Lead Acid Battery
Anode: Pb(s) + SO42-(aq) → PbSO4(s) + 2e0
+2
Cathode:
PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- → PbSO4(s) + 2H2O(l)
+4
+2
Lead-acid batteries are charged by your
car’s alternator.
Stop working because some of the PbSO4(s)
falls to the bottom of the battery and is
removed from the reaction
VI. Corrosion
All redox reactions are not constructive!!
2Fe(s) + 2H2O(l) + O2(g) → 2Fe(OH)2(s)
0
0
+2 -2
(Fe(OH)2 is further oxidized to Fe2O3 (rust)
Corrosion costs US Govt $276 Billion/year
OYO
16.8 Consider the lead acid battery:
Pb(s) + PbO2(s) + 4H+(aq) + 2SO42-(aq) →
2PbSO4(s) + 2H2O(l)
What is being oxidized and what is being
reduced?