chapter 5 volatic battery,chemistry third grade for chemistry department english section explain in details, slide share document

1. G12A
Chemistry:
Chapter 5:
Electrochemistry
Lesson 1 : Voltaic Cell

2. Chapter 5: Electrochemistry Lesson 1: Voltaic Cell

3. Starter Activity

4. Electrochemistry
• A camera trap captured this image of a mountain lion. Camera
traps are a noninvasive way to study animals by using a sensor
that triggers a camera’s shutter when the animal approaches.
Batteries power both the camera and the sensor.
• Chemical energy can be converted to electric energy and
electrical energy to chemical energy.

5. Redox in Electrochemistry
Electrochemistry is the study of the redox processes by which
chemical energy is converted to electrical energy and vice versa.
REDOX
Reduction: Gaining electron
Oxidation: losing electron
Recall: Because any loss of electrons by one substance must be accompanied by a gain in
electrons by something else, oxidation and reduction always occur together.

6. Redox in Electrochemistry

7. Electrochemical cells
Apparatus that uses a redox reaction to produce electric energy
Oxidation: Zn(s) → Zn2+(aq) + 2e Reduction: Cu2+(aq) + 2e- → Cu
Oxidation half-reaction: electron lost Reduction half-reaction: electrons gained
Bulb doesn’t light up
What is missing here?

8. Redox in Electrochemistry
• What do you think would happen if you separated
the oxidation half-reaction from the reduction
half-reaction? Can redox occur?
• 1) There is no way for zinc atoms to transfer
electrons to copper(II) ion.
• 2) A positive charge builds up in one solution and
a negative charge builds up in the other. The
buildup of positive zinc ions on the left prohibits
the oxidation of zinc atoms. On the other side, the
buildup of negative ions prohibits the reduction of
copper ions.

9. Redox in Electrochemistry

10. Salt-Bridge
A salt bridge is a pathway to allow the passage of ions from one side to
another, so that ions do not build up around the electrodes.
Oxidation: Zn(s) → Zn2+(aq) + 2e Reduction: Cu2+(aq) + 2e- → Cu

11. Salt-Bridge
• the purpose of a salt bridge is not to move
electrons from the electrolyte; rather it's to
maintain charge balance because the
electrons are moving from one-half cell to
the other. Salt bridge prevents the diffusion
or mechanical flow of solution from one-half
cell to another.
• It maintains electrical neutrality within the
internal circuit.
• If no salt bridge were present, the solution in
one half cell would accumulate negative
charge and the solution in the other half cell
would accumulate positive charge as the
reaction proceeded, quickly preventing
further reaction, and hence production of
electricity.

12. Electrochemical Cell

13. Starter: Voltaic Cells
• Main Idea: In voltaic cells, oxidation takes place at the anode, yielding electrons that
flow to the cathode, where reduction occurs.
can you clap with one
hand?
You cannot, you need two
Similarly, in Voltic cells
there are two half-cells,
and both are required to
produce energy.

14. Invention of voltaic cell
Did you know, the anatomy
of a frog has led to this
revolutionary invention!!

15. Voltaic Cells

16. Voltaic cell

17. Voltaic cells
Building the simplest
Voltaic Cell
Let’s try to build a voltaic
cell. We place a copper
plate and a zinc plate in
solution of sodium chloride

18. Chemistry of voltaic cells
Half-cells: The cell in which either oxidation or reduction takes place
Zn(s)/ Zn2+(aq)
Oxidation
Cu(s)/ Cu2+(aq)
Reduction
The electrode where
oxidation takes place
is called the anode.
(An-Ox)
The electrode where
reduction takes place
is called the cathode
(red-cat)
Anode
Cathode
Electron flow from anode to cathode and current flow from cathode to anode

19. Voltaic cells and energy
• The roller coaster at the top of the track has high potential
energy relative to track below because of the difference in
height. Similarly, an electrochemical cell has potential energy
to produce a current because there is a difference in the
ability of the electrodes to move electrons from the anode to
the cathode.

20. Calculating Electrochemical Cell Potentials
• The tendency of a substance to gain electrons is its
reduction potential
• The reduction potential of an electrode cannot be
determined directly, because the reduction half-reaction
must be coupled with an oxidation half-reaction.
• Chemists decided to measure the
reduction potential of all electrodes
against one electrode:
 the standard hydrogen electrode (SHE)
Electric charge can flow between two points only when a
difference in electrical potential energy exists between the two
points.

21. Calculating Electrochemical Cell Potentials
• The measure of the ability of a species to gain or lose electrons
Electrode Potential (E).
• The measure of the ability of a species to gain or lose electrons at its
standard state (1M concentration, 1atm pressure and and
temperature 25°C)is called Standard Electrode Potential (E°).

22. Standard hydrogen electrode [SHE]
SHE can act as an oxidation half-reaction or a reduction
half reaction, depending on the half-cell to which it is
connected
NOTE:
Electrode
potential
of SHE is
always
Zero.

23. Standard hydrogen electrode [SHE]
Salt bridge
1M acid solution ,H+(aq)
H2(g) bubbles
Platinum electrode
H2(g) (at 1 atm)
Oxidation:
H2(g) → 2H+(aq) + 2e
Reduction:
2H+(aq) + 2e- → H2(g)

24. Electromotive Force (EMF):

25. Electromotive Force (EMF):
• Electrons generated at the anode, the
site of oxidation, are thought to be
pushed or driven toward the cathode
by the electromotive force (EMF).
• This force is due to the difference in
electric potential energy between the
two electrodes and is referred to as the
cell potential.

26. What is a Volt?
• A Volt is a unit used to measure cell
potential.
• The electric potential difference of
a voltaic cell is an indication of
the energy that is available to
move electrons from the anode to
the cathode.

27. Voltaic cell
Anode
Zinc
electrode
Copper
electrode
AnOx
Anode Oxidation
Negative
charge (-)
Cathode
Positive
charge (+)
RedCat
Reduction at cathode
Summary- Concept map
standard potential of a voltaic cell is the difference between the standard reduction
potentials of the half-cell reactions
Oxidation
half-cell
Reduction
half-cell

29. Cu2+(aq) + 2e- → Cu
E0 (v)=+0.3419
2H+(aq) + 2e- → H2(g)
E0 (v)=+0.000V
Half-Cell Potentials

30. Half-Cell
Potentials

31. Determining Electrochemical Cell Potentials
E
0
cell
= Overall standard cell potential
E
0
reduction
= standard half − cell potential for reduction
E
0
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛
= standard half − cell potential for Oxidation

32. Activity- Label the parts

33. Activity- Label the parts

34. Cu2+(aq) + 2e- → Cu
E0 (v)=+0.3419
2H+(aq) + 2e- → H2(g)
E0 (v)=+0.000V
Half-Cell Potentials

35. Determining Electrochemical Cell Potentials
E
0
cell
= Overall standard cell potential
E
0
reduction
= standard half − cell potential for reduction
E
0
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛
= standard half − cell potential for Oxidation
Recall

36. Electrochemical Cell Potentials
Cu2+(aq) + 2e- → Cu(s) (reduction half-cell reaction)
H2(g) + Cu2+(aq) → 2H+(aq) + Cu(s)
H2(g) | H+(aq) || Cu2+(aq) | Cu(s) E
0
cell
= +0.3419 V
This reaction can be written in a form called Cell notation
H2(g) → 2H+(aq) + 2e- (oxidation half-cell reaction)
Reactant Product
Oxidation half-cell Reduction half-cell
Reactant Product

37. 2H+(aq)) + Zn (s) → Zn2+(aq) + H2(g)
Zn(s) | Zn2+(aq) || H2(g) | 2H+(aq) E
0
cell
= -0.762 V
This reaction can be written in a form called Cell notation
2H+(aq) + 2e- → H2(g)(reduction half-cell reaction)
Reactant Product
Oxidation half-cell Reduction half-cell
Reactant Product
Zn(s) → Zn2+(aq) + 2e- (oxidation half-cell reaction)
Electrochemical Cell Potentials

38. The Notation for Voltaic Cells

39. Determining Electrochemical Cell Potentials
E
0
cell
= E
0
reduction
− E
0
oxidation
E
0
cell
= +0.3419 − (-0.762V)
= +1.104 V
0.000 V
Standard
hydrogen
electrode
+0.342 V
Cu2+ | Cu
electrode
-0.762 V
Zn2+ | Zn
electrode
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
E
0
cell
= 1.04 V
E
0
cell
= E
0
Cu2+(aq) | Cu(s)
− E
0
Zn(s) | Zn2+(aq)
Note: Low reduction potential value
for oxidation and high reduction
potential value for Reduction

40. Cell potential in class Example-1
Analyze the Problem
Known
Standard reduction potentials for the half-cells
E 0
cell
= E 0
reduction
− E 0
oxidation
Unknown
overall cell reaction = ?
E
0
Cell = ?
cell notation = ?
The following reduction half-reactions represent the half-cells of a voltaic cell.
I2(s) + 2e– → 2I–(aq) and Fe2+(aq) + 2e– → Fe(s) Determine the overall cell reaction
and the standard cell potential. Describe the cell using cell notation
I2(s) + 2e–→2I–(aq) (reduction half-cell reaction)
Fe(s) +2e– → Fe2+(aq) (oxidation half-cell reaction)
E 0
I2
= +0.536 V
E 0
Fe
= -0.447 V
Solve the Problem
E 0
cell
= E 0
reduction
− E 0
oxidation E
0
cell
= +0.536 − (-0.447) = +0.983V

41. Cell potential in class Example-1
Rewrite the iron half-reaction in the correct direction.
I2(s) + 2e–→2I–(aq) (reduction half-cell reaction)
Fe(s) → Fe2+(aq)+2e– (oxidation half-cell reaction)
I2(s) + Fe(s) → 2I–(aq) + Fe2+(aq)
Fe | Fe2+ || I2 | I–
Cell notation
Evaluate the answer
E
0
cell
= +0.536 − (-0.447) = +0.983V E
0
cell
positive so the reaction is spontaneous
Solve the Problem
Salt bridge

42. Cell potential Practice problems

43. Cell potential Practice problems Answers
1. 2. 3.
4.

44. Using Standard Reduction Potentials
• Another important use of standard reduction potential is to
determine if a proposed reaction under standard conditions will
be spontaneous.
• Electrons in a voltaic cell always flow from the half-cell with the lower
standard reduction potential to the half-cell with the higher reduction
potential, giving a positive cell voltage.
• How can standard reduction potentials indicate spontaneity?

45. Using Standard Reduction Potentials
To predict whether any proposed redox reaction will occur spontaneously:
• Step 1: write the process in the form of half-reactions and look up the reduction
potential of each
• Step 2: Use the values to calculate the potential of a voltaic cell operating with
these two half-cell reactions.
• Step 3: Evaluate if the reaction spontaneous or nonspontaneous
•  if the calculated potential is positive, the reaction is spontaneous.
•  If the value is negative, the reaction is not spontaneous.

46. Using Standard Reduction Potentials
To predict whether any proposed redox reaction will occur spontaneously:
• Step 1: write the process in the form of half-reactions and look up the reduction
potential of each
• Step 2: Use the values to calculate the potential of a voltaic cell operating with
these two half-cell reactions.
• Step 3: Evaluate if the reaction spontaneous or nonspontaneous
•  if the calculated potential is positive, the reaction is spontaneous.
•  If the value is negative, the reaction is not spontaneous.
The reverse of a
nonspontaneous reaction will
occur because it
will have a positive cell voltage,
which means that the reverse
reaction is spontaneous.

47. Using Standard Reduction Potentials –
Practice problems

48. Using Standard Reduction Potentials –
Practice problems Answers
9.

49. Voltaic cell
Anode
Zinc
electrode
Copper
electrode
AnOx
Anode Oxidation
Negative
charge (-)
Cathode
Positive
charge (+)
RedCat
Reduction at cathode
Summary- Concept map
standard potential of a voltaic cell is the difference between the standard reduction
potentials of the half-cell reactions
Oxidation
half-cell
Reduction
half-cell

50. Section Summary- Concept map-2
H2(g) 1 atm, Pt
Negative
oxidation
Positive
reduction
Standard
potential is
measured by
connecting
with SHE
Half
Cell
Either
oxidation or
reduction
E0 (v)=+0.00V
Cathode
reduction
Standard
hydrogen
electrode
(SHE)
Anode-
Oxidation
STD
conditions
1 atm
298k
1mol/dm-3