Transcript Redox Notes_2015

Unit 12: Redox and Electrochemistry
- RB Topic 11
*What kind of chemical reaction is responsible for
burning, rusting, bleaching clothes, powering
your cell phone as well as your body, and giving
fireflies their glow?
• Oxidation-Reduction Reactions, also called
REDOX
I. What are Redox Reactions?
• involve the transfer (gain/loss) of
ELECTRONS (e–) between different atoms
Na + Cl  Na
OXidation
+
Cl
REDuction
•losing electrons
•charge goes up
•gaining electrons
•charge is reduced
•Ex:
•Ex:
*CONSERVATION OF CHARGE*
oxidation and reduction happen
simultaneously (at the same time) – you can’t
have one without the other
 charges on both sides of arrow
MUST BE EQUAL
Remember:
LEO the lion says GER
• LOSING ELECTRONS is OXIDATION
• GAINING ELECTRONS is REDUCTION
II. Determining if a reaction is a redox reaction
If the oxidation number (charge) of an
atom CHANGES from reactant to product, a
redox reaction has occurred
*Note:
Double replacement reactions are
NOT redox reactions
Single replacement reactions are
ALWAYS redox reactions
Ex:
CaCl2 + Na2S  CaS + 2NaCl
Ca + MgS  Mg + CaS
A. Assigning Oxidation Numbers
(listed on the Periodic Table)
1. If a species is all by itself (not combined with
a different element), its oxidation number is
ZERO
Ex:
Mg
O2
S2
Cu
H2
2. In a polyatomic ion, the oxidation #’s of the
elements add up to equal the CHARGE on the ion
*first element is always +
*last element is always -
NH4
+
–
ClO
SO3
2–
SO4
2–
3. In a compound, the oxidation numbers of the
elements add up to ZERO (compounds are
neutral)
*first element is always +
*last element is always –
NaCl
NH4Cl
H2SO3
More Examples:
a.)
What is the oxidation number of chlorine in
iron(III) chloride?
b.)
What is the oxidation state of chromium in
potassium chromate?
• Skill 1
B. Checking for Changes in Oxidation Numbers
– if they change…it’s REDOX
1. start by assigning ox. #s
2 Na (s) + Cl2 (g) → 2NaCl (s)
•
•
•
•
The oxidation # on Na changes from 0 to +1
Na loses electrons
Na is oxidized
Na is the reducing agent
2 Na (s) + Cl2 (g) → 2NaCl (s)
•
•
•
•
The oxidation # on Cl changes from 0 to –1
Cl2 gains electrons
Cl2 is reduced
Cl2 is the oxidizing agent
2Mg (s)
+
O2 (g)
→
2MgO (s)
•
•
•
•
The oxidation # on Mg changes from 0 to +2
Mg loses electrons
Mg is oxidized
Mg is the reducing agent
•
•
•
•
The oxidation # on O changes from 0 to –2
O2 gains electrons
O2 is reduced
O2 is the oxidizing agent
• Skill 2
Ticket
Given the following:
4 Mn + 7 O2  2 Mn2O7
(a) Assign oxidation numbers to each symbol in the
reaction
(b) Write the formula of the substance that is oxidized
(c) Write the formula of the substance that is reduced
III. Writing Half-Reactions
1. Write oxidation #s
2. Write in electrons gained (reactant) or lost
(product) and include COEFFICIENTS
3. Make sure # of atoms and total charge STAYS
SAME (conservation of MASS and CHARGE)
2 Na + Cl2
Reduction half-reaction:
shows the electrons that are
gained and the element being
reduced
Ex: Mg2+
+
2e–
→
THE
→ 2 NaCl
Oxidation half-reaction:
shows the element being
oxidized and the electrons
that are lost
Ex: 2I–

I2 +
2e–
Sample Problems:
Write reduction and oxidation halfreactions for each of the following:
(a)
2Fe3+ + Ni  2Fe2+ + Ni2+
(b)
Br2 + Hg 
2+
Hg
+
–
2Br
(c)
Cu +
+
2Ag
 2Ag +
2+
Cu
(d)
4+
Sn
+ H2  Sn +
+
H
• Skill 3
In the presence of oxygen, iron forms iron (III)
oxide, more commonly known as rust, through
a reduction-oxidation reaction
4Fe + 3O2  2Fe2O3
What causes fruit to turn brown?
• Oxidation of compounds in fruit after
exposure to oxygen in the air causes a change
of color (browning) of fruit
Oxidation (losing electrons) is a process that can
be very damaging to living things.
• results in the creation of free radicals, which
are chemical species that have unpaired
electrons (they are unhappy – they don’t have
a completed valence shell)
• free radicals will rip electrons away from other
compounds in living things, such as proteins
or molecules of DNA, causing damage which
can lead to cancer
• Antioxidants get rid of free radicals
– Vitamins A, C, and E are antioxidants
IV. Electrochemical Cells
A. Use REDOX reactions to convert
chemical energy into electrical energy
OR convert electrical energy into
chemical energy
B. Two types of Electrochemical Cells
1. Voltaic /Galvanic (battery)
2. Electrolytic Cells
• convert chemical 
electrical energy
• Use a chemical reaction to
produce RELEASE electricity
(moving electrons)
• convert electricity 
chemical energy
• use electricity to make a
chemical reaction to occur)
1. Voltaic (Galvanic) Cells = BATTERIES
a. SPONTANEOUS redox reactions
– NO energy put in (energy is released)
***Use Table J to determine what is
oxidized and what is reduced
Higher on Table J is OXIDIZED (LEO)
Lower on Table J is REDUCED (GER)
b. Diagram of a voltaic cell using Zn and Zn(NO3)2
with Cu and Cu(NO3)2
*Parts to Know*
Electrodes: the places where oxidation
or reduction happens
o Anode: the site of oxidation (An Ox)
o Cathode: the site of reduction (Red Cat)
Wire: electrons flow through the wire
from the anode (–) to the cathode (+)
Salt bridge: allows IONS to flow
between the electrodes to keep the
charges balanced
The electrodes must be separated in order to
produce an electric current (flow of electrons).
The energy present in the flowing electrons
(ELECTRICITY) is captured and used to power
other processes.
c. Voltaic cell problems:
1. Look on Table J and find which element is
higher – this element is OXIDIZED
(On Table J – electrons flow
DOWNHILL, spontaneously)
2. Under each beaker write “oxidation” or
“reduction”
3. Label the oxidation electrode
“ANODE” (An Ox)
4. Label the reduction electrode
“CATHODE” (Red Cat)
5. Place a (–) charge on the anode
6. Place a (+) charge on the cathode
7. Draw in the direction of electron flow (from
ANODE (–) to CATHODE (+))
8. Write the half reactions under the correct
beakers
9. The salt bridge allows for ions to flow
between the electrodes
Positive (+) ions flow toward the CATHODE
(to balance the negative electrons)
Negative (–) ions flow toward the ANODE
(to replace the electrons that are leaving)
b. Diagram of a voltaic cell using Zn and Zn(NO3)2
with Cu and Cu(NO3)2
Label the parts of this voltaic cell
How many of you have had cavities filled?
…if you do, you have the potential to make a
tiny battery IN YOUR MOUTH!
• Cavities are filled with a mixture of metals
including zinc (Zn), tin (Sn), copper (Cu), and
silver (Ag)
• If you bite down on a piece of aluminum foil,
the saliva in your mouth, the aluminum foil,
and the filling make a little voltaic cell that
produces a tiny current that travels through
your tooth to the nerve below the filling…it’s a
bit UNPLEASANT!
• Electrochemical Cells Lab
• Lemon battery lab
2. Electrolytic Cells
a. NONSPONTANEOUS redox reactions –
need to put in energy
b. The species more likely to lose electrons
is forced to gain electrons
(On Table J – electrons are moving
UPHILL, nonspontaneously)
c. Electrons still flow from the anode to the
cathode, but now the signs are reversed
The anode is (+) and the cathode is (-)
(electrons are being forced to travel to the
negative electrode)
d. Used for
1. Electrolysis: using electricity to
break apart (lyse) compounds into their
elements
Ex: Obtaining active elements from compounds
(a) complete and balance following electrolysis reaction.
(b) assign oxidation numbers to each element
___NaCl (l)
____H2O (l)
2. Electrolytic cells can also be used for
electroplating – putting a metal coating on
something
Comparing & Contrasting Voltaic and Electrolytic Cells
Voltaic Cells (Batteries)
Electrolytic Cells
• SPONTANEOUS
redox rxn
• nonspontaneous
redox rxn
• energy is RELEASED!
(RELEASES electricity)
• energy is needed
(uses electricity)
Comparing & Contrasting Voltaic and Electrolytic Cells
Voltaic Cells (Batteries)
Electrolytic Cells)
• redox reactions
• anode is where oxidation happens
(An Ox)
• anode loses mass
• cathode is where reduction happens
(Red Cat)
• cathode gains mass
• electrons flow through wire
from anode to cathode
p. 186 #s 37-41, 51-52, 65-69, 71-74, 76-86, 88-91
37.
38.
39.
40.
41.
3
2
3
2
4
65.
66.
67.
68.
69.
1
4
2
4
3
71.
72.
73.
74.
1
1
3
4
76.
77.
78.
79.
80.
81.
2
3
2
2
3
2
82.
83.
84.
85.
86.
4
1
4
2
1
51. 2
52. 2
91. The anode loses mass, the cathode gains mass