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
About This Presentation



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
Provided by: conrad9


Write a Comment
User Comments (0)
Transcript and Presenter's Notes

Title: 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.

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
  • 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

Figure 17.2Galvanic Cells
  • 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
  • If we connect the metal electrodes with a wire,
    the following observations are made
  • 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.

  • 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.

Figure 17.6 A Zn/Cu Galvanic Cell
  • At the left-hand electrode (anode half-reaction)
  • Zn(s) ? Zn2(aq) 2 e oxidation
  • At the right-hand electrode (cathode
  • 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)

  • 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).

Inert Electrodes
  • Sometimes a half-reaction will not involve a
    solid substance
  • MnO4-(aq) 8H (aq) e ? Mn2 (aq) 4 H2O
  • 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.

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

Porous Cup
MnO4(aq) H (aq)
Zn 2(aq)
Figure 17.7A Schematic of a Galvanic Cell
  • 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.

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.

  • 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.

  • ?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.

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

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)

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.

Voltaic Cell Cu(s) / Cu2(aq) // Zn 2(aq) /
Salt Bridge
Zn 2(aq)
A Zn/H Galvanic Cell
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) )
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.
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
  • 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

  • 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
  • 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.

Voltaic Cells
  • Primary Cells
  • Produce energy until one component is used up,
    then discarded
  • Examples are the zinc chloride cell and the 9
  • Secondary Cells
  • Store energy and may be recharged
  • Examples are Ni-Cad cells and lead acid car

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)

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

Dry Cell
Metal Cap ()
Mixture of Carbon Manganese Dioxide
CathodeCarbon Rod
Ammonium Chloride Zinc Chloride Electrolyte
Anode Zinc Case ()
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

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

Alkaline Cell
Metal Cap ()
Cathode outer steel case
Potassium Hydroxide Electrolyte
Powdered Zinc
AnodeSteel or Brass
Mixture of Carbon Manganese Dioxide
Metal Base ()
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

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

Button Cell
Metal Cap ()
Zinc Powder
Cathode outer container of nickel or steel ()
Mercury Oxide
Secondary Cells
  • Lead Acid (Car Battery)
  • Nickel cadmium Cells
  • Fuel Cells

Lead Acid Battery
  • 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

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

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

Nickel Cadmium Cells
  • Often called Nicads
  • Electrodes are nickel and cadmium
  • Electrolyte is potassium hydroxide
  • Reactions involve the hydroxides of the two metals

Nickel Cadmium Cells
  • 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
  • Overall Reaction
  • Cd (s) NiO-OH(s) H2O(l) ? Cd(OH)2 (s)
    Ni(OH)2 (s)

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

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

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)

Hydrogen Oxygen Fuel Cell

Oxygen Gas Inlet
HydrogenGas Inlet
Porous Anode
Porous Cathode
Water outlet
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
  • Cations move to the cathode, anions move to the

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.

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

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

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

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
  • Predict the products for the electrolyis of brine
    (aqueous sodium chloride)

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
  • 2 H20 (l) 2 H2 (g) O2 (g)

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)

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

Aluminum production
Electrorefining of Copper
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

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

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

Cell Stoichiometry (contd)
  • The mathematical relationship is

where I current (amperes
or Coulombs/second) t time (seconds) F
Faradays constant 96500 Coulombs/molee
ne It
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

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

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)

Corrosion and Electrochemistry
  • In an acidic environment, the OA is the H /O2
  • 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

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

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?

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
    will react instead)
About PowerShow.com