Cells - PowerPoint PPT Presentation

PPT – Cells PowerPoint presentation | free to view - id: 11ffea-MTNlZ

The Adobe Flash plugin is needed to view this content

Get the plugin now

View by Category
Title:

Cells

Description:

Cell Stoichiometry ... Cell Stoichiometry (cont'd) The mathematical relationship ... Cell Stoichiometry. ne = (4.7 C / s) x (7650 s) 96500 C / mole = 0.37 mole ... – PowerPoint PPT presentation

Number of Views:139
Avg rating:3.0/5.0
Slides: 70
Category:
Tags:
Transcript and Presenter's Notes

Title: Cells

1
Cells
• Electrochemical Cell Contains electrodes
dipping into electrolyte solutions. The
half-reactions occur at two different electrodes
and a flow of electrons (current) is involved.
• Voltaic (galvanic ) cells a spontaneous redox
reaction occurs and causes a flow of electrons.
• Electrolytic cells an external source of
electric current is used to drive an otherwise
nonspontaneous redox reaction.

2
Voltaic (Galvanic) Cells
• Consider the spontaneous redox reaction that
occurs if a piece of copper is immersed in a
solution of silver nitrate (occurring in one
beaker)
• Cu(s) 2 Ag(aq) ? Cu2(aq) 2 Ag(s)
• No useful work is done in this process (e.g., no
usable electric current). Some heat will be
released to the surroundings.
• How can we harness this spontaneous reaction in
such a way as to yield useful work?
• Answer Separate the half-reactions in two
containers.

3
Figure 17.2Galvanic Cells
4
• Ex Suppose we have a piece of Zn metal dipping
into a 1 mol/L solution of ZnSO4 in the left-hand
compartment and a piece of Cu metal dipping into
a 1 mol/L solution of CuSO4 in the right-hand
compartment.
• If we connect the metal electrodes with a wire,
• The piece of Zn decreases in mass over time.
• The piece of Cu gains mass over time.
• The concentration of the Zn2(aq) in the
left-hand compartment increases over time.
• The concentration of Cu2(aq) in the right-hand
compartment decreases over time.

5
• What is the salt bridge?
• It contains a solution of an electrolyte such as
NaNO3 in a gel, such that the ions are mobile but
will not mix rapidly with the other solutions.
• If we place a voltmeter between the two
electrodes, we obtain a reading of 1.10 V.
• Now, lets analyze more closely what is happening
in the cell.

6
Figure 17.6 A Zn/Cu Galvanic Cell
7
• At the left-hand electrode (anode half-reaction)
• Zn(s) ? Zn2(aq) 2 e oxidation
• At the right-hand electrode (cathode
half-reaction)
• Cu2(aq) 2 e ? Cu(s) reduction
• Add them to get the overall reaction
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• But now the electrons lost by the Zn anode flow
through the external circuit to the Cu cathode
(an electric current which we can make use of)

8
• Note By convention the anode is placed on the
left and the cathode on the right when picturing
a galvanic cell.
• You should be able to describe the half-reactions
at each electrode and the direction of movement
of all species (including electrons and the ions
of the salt bridge).

9
Inert Electrodes
• Sometimes a half-reaction will not involve a
solid substance
• MnO4-(aq) 8H (aq) e ? Mn2 (aq) 4 H2O
(l)
• In such cases, an inert metal such as Pt is used
as an electrode to which the wire is attached. A
example is shown in the following figure.

10
Voltaic Cells with Inert Electrodes
• In this example the one half reactions has no
solid metal so carbon solid is used as the
electrode.
• Do the analysis of this cell.

C(s)
Zn(s)
Porous Cup
Electrolyte
MnO4(aq) H (aq)
Zn 2(aq)
11
Figure 17.7A Schematic of a Galvanic Cell
12
• Voltaic Cells have two electrodes
• Cathode electrode
• (GERC)
• Gain of
• Electrons is the process of
• Reduction and occurs at the
• Cathode.
• Anode electrode
• (LEOA)
• Loss of
• Electrons is the process of
• Oxidation and occurs at the
• Anode.

13
Cathode vs. Anode
• Cathode
• GERC
• Is a solid.
• Is the location of reduction.
• Electrons are on the reactant side.
• Electrons are consumed.
• Has a charge.
• Anode
• LEOA
• Is a solid
• Is the location of oxidation.
• Electrons are on the product side.
• Electrons are produced.
• Has a charge.

14
• Cell Potential (or EMF or voltage)
• Think of this as the push that sends electrons
from the anode to the cathode.
• We measure (with a voltmeter) the potential
difference between the anode and cathode.
• ?Ecell Ecathode Eanode
• The superscript refers to standard-state
conditions and reversible cell operation.

15
• ?Ecell must be positive in order for a
spontaneous reaction to occur. Thus, the cathode
must be at a higher potential than the anode
(think of electrons rolling downhill).
• Ex for the Zn-Cu cell previously described, ?E
1.10 V.

16
• The Standard Hydrogen Electrode (S.H.E.)
• 2 H(1 mol/L) 2e ? H2(1 atm) E 0.00
V
• is used as the reference for all half-reactions

17
Analysis of a Voltaic Cell
• Write the two half cells including the voltages
(Change the sign on the oxidation half-reaction
from the one listed in your data booklet).
• Balance electrons and write the net reaction with
the net voltage. (dont change the voltage)
• Identify the cathode. (GERC)
• Identify the anode. (LEOA)

18
Analysis of a Voltaic Cell (contd)
• Net voltage is the sum of the reduction potential
and oxidation potential.
• Identify the electrodes and electrolytes.
• Identify the directions of electron movement in
the wire.
• Identify the movement of the anions and cations
in the solution.

19
Voltaic Cell Cu(s) / Cu2(aq) // Zn 2(aq) /
Zn(s)
Electrode
Salt Bridge
Zn(s)
Cu(s)
Cu2(aq)
Zn 2(aq)
Electrolyte
20
A Zn/H Galvanic Cell
21
2 H(aq) 2e- H2(g)
Zn (s) Zn2(aq) 2e-
2 H(aq) Zn (s)
H2 Zn2(aq)
The voltaic cell on the previous slide is fully
described with the following notation
Zn(s) Zn(aq) H(aq), H2(g) Pt(s) )
22
Electrons flow from the anode to the cathode in a
voltaic cell. (They flow from the electrode at
which they are given off to the electrode at
which they are consumed.) Reading from left to
right, this line notation therefore corresponds
to the direction in which electrons flow.
23
Line Notation For Voltaic Cells
• Voltaic cells can be described by a line notation
based on the following conventions.
• Single vertical line indicates change in state or
phase.
• Within a half-cell, the reactants are listed
before the products.
• A double vertical line indicates a junction
between half-cells.
• The line notation for the anode (oxidation) is
written before the line notation for the cathode
(reduction).

24
• In a galvanic cell, the electrode with the higher
E value will be the cathode.
• Ex In a cell using the Ni2Ni (0.25 V) and
AgAg (0.80 V) half-reactions, Ag will be the
cathode
• and the spontaneous reaction will be
• Ni(s) 2 Ag(aq) ? Ni2(aq) 2 Ag(s)
• ?E 0.80 (0.25) 1.05 V
• Important Notice that the sign of the potential
of the anode is used as found in the table (not
reversed). This is because we subtracted the
anode potential from the cathode potential.

25
Voltaic Cells
• Primary Cells
• Produce energy until one component is used up,
• Examples are the zinc chloride cell and the 9
volt
• Secondary Cells
• Store energy and may be recharged
batteries

26
Dry Cells
• The ordinary zinc carbon cell
• Anode (oxidation)
• Zn (s) ? Zn 2 (aq) 2e
• Cathode (reduction)
• 2MnO2 (s) NH4 (aq) 2e ? Mn2O2 (s) 2NH3
(aq) H2O (l)

27
Dry Cells
• A new cell produces about 1.5V
• Once reaction reaches equilibrium its flat

28
Dry Cell
Metal Cap ()
Mixture of Carbon Manganese Dioxide
CathodeCarbon Rod
Ammonium Chloride Zinc Chloride Electrolyte
Anode Zinc Case ()
29
Alkaline Cells (eg. Duracell)
• The ordinary zinc carbon cell
• Anode (oxidation)
• Zn (s) ? Zn 2 (aq) 2e
• Immediately reacts with OH ions in the
electrolyte to form zinc hydroxide
• Zn (s) 2OH (aq) ? Zn(OH)2 (s) 2e

30
Alkaline Cells
• Cathode (reduction)
• 2MnO2 (s) H2O(l) 2e ? MnO2 (s) OH (aq)
H2O (l)
• Five times the life of the dry cell

31
Alkaline Cell
Metal Cap ()
Cathode outer steel case
Potassium Hydroxide Electrolyte
Powdered Zinc
AnodeSteel or Brass
Mixture of Carbon Manganese Dioxide
Metal Base ()
32
Button Cells
• Used in very small applications like watches,
cameras etc.
• Two main types
• Mercury zinc and silver zinc
• Anode (Oxidation)
• Zn (s) 2OH (aq) ? Zn(OH)2 (s) 2e

33
Button Cells
• Cathode Reduction ()
• depends on the type of battery
• HgO(s) H2O (l) 2e ? Hg (l) 2OH (aq)
• Ag2O(s)H2O (l) 2e ? 2Ag (s) 2OH (aq)
• Produce an almost constant 1.35V

34
Button Cell
Metal Cap ()
Zinc Powder
Cathode outer container of nickel or steel ()
Electrolyte
Mercury Oxide
35
Secondary Cells
• Fuel Cells

36
• Car Batteries
• Also called storage batteries or accumulators
• Each cell produces 2 volts so typical 12 volt car
battery contains 6 cells
• Both electrodes are lead plates separated by some
porous material like cardboard

37
• Positive electrode is coated with PbO2 Lead (IV)
Oxide
• The electrolyte is a solution of 4M sulfuric acid

38
• Anode Oxidation ()
• Pb(s) SO4 2- ? PbSO4 (s) 2e
• Cathode Reduction ()
• PbO2(s) SO4 2- 4H 2e ? PbSO4 (s) 2H2O
(l)
• Overall Reaction
• Pb(s) PbO2(s) 2H2SO4 ? 2PbSO4 (s) 2H2O (l)

39
• Electrodes are nickel and cadmium
• Electrolyte is potassium hydroxide
• Reactions involve the hydroxides of the two metals

40
• Anode (Oxidation)
• Cd (s) 2OH (aq) ? Cd(OH)2 (s) 2 e
• Cathode (Reduction)
• NiO-OH (s) H2O (l) e ? Ni(OH)2 (s) OH
(aq)
• Overall Reaction
• Cd (s) NiO-OH(s) H2O(l) ? Cd(OH)2 (s)
Ni(OH)2 (s)

41
Fuel Cells
• Limitation of dry cells looked at so far is that
they contain reactants in small amounts and when
they reach equilibrium.
• Primary Cells are then discarded, secondary cells
are then recharged
• A cell that can be continually fed reactants
would overcome this and allow for a continual
supply of electricity

42
Fuel Cells
• Fuel cells transform chemical energy directly
into electrical energy
• 60 efficiency
• Space Program uses hydrogen and oxygen with an
electrolyte of potassium hydroxide

43
Fuel Cells
• Anode (Oxidation)
• H2(g) 2OH (aq) ? 2H2O (l) 2e
• Cathode (Reduction)
• O2(g) 2H2O(l) 4e ? 4OH(aq)
• Overall Equation
• H2(g) O2(g) ? 2H2O (l)

44
Hydrogen Oxygen Fuel Cell

Electrolyte
Oxygen Gas Inlet
HydrogenGas Inlet
Porous Anode
Porous Cathode
Water outlet
45
Some basic generalizations about Voltaic Cells
• The net voltage is always positive.
• The cathode is the location of reduction. (GERC)
• The anode is the location of oxidation.(LEOA)
• Electrons always move from the anode to the
cathode.
• Cations move to the cathode, anions move to the
anode.

46
Electrolytic Cells
• In an electrolytic cell, a non-spontaneous redox
reaction is forced to occur by applying an
external voltage to the cell
• The potential for an electrolytic cell is
negative. This is the voltage that must be
supplied to cause the reaction to occur.

47
(No Transcript)
48
• This is a diagram of a Hoffman apparatus which is
used to separate water into hydrogen and oxygen
by applying an electrical current.

49
• Predict the products in the electrolysis of water
• What evidence would confirm your prediction?

50
Electrolysis of Water
• Prediction
• Cathode
• 2H2O 2e- H2 2OH- E-0.83V
• Anode
• 2H2O O2 4H 4e- E-1.23V
• Net
• 2H2O 2H2 O2 E -2.06V

51
Electron flow in the circuit is opposite to the
conventional current flow. The reaction at the
cathode (tube A) is the reduction of water2
H2O(aq) 2e- H2 (g) 2 OH-(aq)Oxidation
takes place at the anode (tube B). The oxidation
of water according to the reaction 2 H2O (l)
4 H(aq) O2( g) 4e-has a standard
electrode potential of -1.23V, so the overall
reaction is therefore2 H2O (l) 2 H2 (g) O2
(g)
52
• Predict the products for the electrolyis of brine
(aqueous sodium chloride)

53
Strange Salt
• In an aqueous solution of sodium chloride, water
would be predicted to be both the SOA and the SRA
so the predicted reaction in an electrolytic cell
is the same as that found for the electrolysis of
water
• 2 H20 (l) 2 H2 (g) O2 (g)

54
Strange Salt (contd)
• However, the aqueous chloride ion is oxidized
preferentially before the water, so the two half
reactions and net reaction are
• 2 Cl-(aq) Cl2 (g) 2e-
• 2 H2O(aq) 2e- H2 (g) 2 OH-(aq)
• 2 Cl-(aq) 2 H2O(aq) Cl2 (g) 2 OH-(aq)

55
Applications of Electrolysis
• Production of reactive metals from their molten
salts (eg. Al from Al2O3 Na from NaCl)
• Electroplating objects to enhance their value
and/or corrosion resistance
• Electrorefining of metals

56
Aluminum production
57
Electroplating
58
Electrorefining of Copper
59
Voltaic vs. Electrolytic Cells
• Spontaneous
• Converts chemical energy to electrical energy
• Oxidation occurs at the anode reduction occurs
at the cathode
• Non-spontaneous
• Converts electrical energy to chemical energy
• Oxidation occurs at the anode reduction occcurs
at the cathode

60
Voltaic vs. Electrolytic Cells
• Positive potential
• SOA and SRA react
• SOA is above SRA on table
• Electrons move from anode to cathode
• Two separated half cells are used
• Negative potential
• SOA reacts with the SRA
• Electrons forced to move from anode to cathode by
external power source
• Generally, only one container is used

61
Cell Stoichiometry
• The amount (moles) reduced or oxidized in a cell
(whether voltaic or electrolytic) is directly
proportional to the moles of electrons
transferred.

62
Cell Stoichiometry (contd)
• The mathematical relationship is

where I current (amperes
or Coulombs/second) t time (seconds) F
ne It
F
63
Cell Stoichiometry (contd)
• What is the mass of copper that can be plated out
in an electrolytic cell if 4.7 Amperes are
delivered for 2.1 hours?
• Known I 4.7 C/s
• t 2.1 h x 3600 s/h 7560 s
• F 96500 C/mole
• Unknown moles of electrons (which will lead to
moles of copper

64
Cell Stoichiometry
• ne (4.7 C / s) x (7650 s)
• 96500 C / mole
• 0.37 mole
• molCu 0.37 mole x 1 molCu x 63.55 g
• 2 mole 1 mol
• 12 g Cu

65
Corrosion and Electrochemistry
• Corrosion is simply the oxidation of structures
by environmental substances
• In a relatively neutral environment, the OA is
the H2O / O2 combination
• O2(g) H2O(l) 4e- 4 OH-(aq)

66
Corrosion and Electrochemistry
• In an acidic environment, the OA is the H /O2
combination
• O2(g) 4 H(aq) 4e- 2 H2O(l)
• Notice that this OA is higher on the redox table,
so it is a more powerful OA

67
Protection From Corrosion
• The cost of corrosion means that steps must be
taken to minimize or eliminate its effects
• One of the easiest ways is to coat the metal
object (ie. Painting or coating with oil)
• Another way is to galvanize the metal (this is a
coating of zinc metal on the object

68
Protection From Corrosion
• Cathodic protection (also called sacrificial
anode) is a means of protecting an expensive
metal object by sacrificing an inexpensive metal
• Consider ships (made of steel, which is mainly
iron), floating on salt water, a wonderful
electrolyte. Why dont they corrode and sink to
the bottom of the ocean?

69
Protection From Corrosion (contd)
• The reason they do not is because of cathodic
protection (sacrificial anodes)
• A plate of zinc or magnesium metal is attached on
the inside hull of the ship. When the OA (O2/H2O)
takes electrons from the iron atoms at the
surface of the ship, electrons are conducted from
the zinc to replace the stolen electrons. In this
way, the iron is preserved in its elemental state
and the zinc is sacrificed (notice on the redox
table that zinc is a stronger RA than iron so it