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REDOX EQUILIBRIA
Oxidation and Reduction
Electrochemical Cells and Fuel Cells
History of Battery Development:
Back in the late 1700’s, Luigi Galvani observed that the behavior of two dissimilar
metals zinc and copper, in contact with the electrolytes (ions in solution) of tissues.
These 2 metals produced an electric charge that caused a frog’s leg to contract.
Galvani coined the term animal electricity to describe whatever it was that
activated the muscles of his specimens.
But batteries were around long before this discovery of animal electricity. A prehistoric battery is shown below:
Ancient Bagdad Batteries
Clay jar with iron rod surrounded by copper cylinder. When
filled with vinegar + an electrolytic, solution produces
1.1 volts DC (circa. 250 BC to 640 AD).
It is believed that the Parthians who ruled Baghdad (circa 250 BC) used batteries to electroplate silver
using the “Baghdad battery”. It is also believed that the Egyptians electroplated antimony onto copper
over 4300 years ago. So what causes the frogs leg to jump? Indeed, why do you jump? What caused all
this were electrons moving from one place to another.
To understand batteries, we need to know what the chemicals are doing; and they are only just transferring electrons.
Electron transfer reactions are redox reactions. The transfer of electrons can be shown by means of half-equations.
Oxidation is the - loss of electrons. (L.E.O) In the Bagdad battery copper ions gained e-: Cu2+ + 2e- Cu(s)
Reduction is the - gain of electrons (G.E.R).
Zn(s) Zn(s) + 2eRedox equations can be created by combining two, equations (as above) and balancing the electrons.
Eg
Fe2+  Fe3+ + eoxidation loses electrons (LEO), called a ½ reaction
+
MnO4 + 8H + 5e  Mn2+ + 4H2O
reduction gains electrons (GER), called a ½ reaction
Multiplying all coefficients in the oxidation reaction by 5:
5Fe2+  5Fe3+ + 5e
(means that 5 electrons are gained and five are lost)
overall combined equation: MnO4- + 8H+ + 5Fe2+  Mn2+ + 4H2O + 5Fe3+
A species which can accept electrons from another species is an oxidizing agent. Eg MnO4- is the oxidizing agent above
A species which can donate electrons to another species is a reducing agent.
Eg Fe2+ in the above reaction is the RA.
Oxidation Numbers
As you can see electrons are the key to understanding redox equations, so we must make sure we can always follow or
find out where the electrons start and where they go. So we have invented a few rules to make sure we can keep track of
electrons, and they are called oxidation numbers.
Rules for Oxidation Numbers.
1) Elements in their standard states, the oxidation number of each atom is zero: This is obviously so because it
has not reacted yet! It has not been oxidized or reduced. Ex: Cl2, S, Na and O2
2) The oxidation number of an atom is the charge on the atom if the bonding were ionic. In simple ions, the
oxidation number of the atom equals the charge on the ion: Ex: Na+, K+, H+ all have an oxidation number of
+1 ; whereas O2-, S2- all have an oxidation number of -2.
3) In neutral compounds, the sum of the oxidation numbers on the atoms is zero:
Eg SO3; oxidation number of S = +6, oxidation number of O = -2.
+6 + 3(-2) = 0
In polyatomic ions (Cr2O72- ). The sum of the oxidation numbers on atoms is equals the charge on the ion.
Eg : Cr2O72-; oxidation number of Cr = +6, oxidation number of O = -2.
2(+6) + 7(-2) = -2
element
oxidation state
exceptions
Group 1 metals
always +1
Group 2 metals
always +2
Oxygen
usually -2
except in peroxides and F2O
Hydrogen
usually +1
except in metal hydrides where it is -1
Fluorine
always -1
Chlorine
usually -1
except in compounds with O or F
Ok, let’s use what we now know and see if we can follow the paths of electrons in an equation. Hopefully you’ll
remember that we first split an equation into 2 reactions (called half (½) reactions. So let us work with the following
chemicals.
Manganate(VII) ions, MnO4-, oxidize hydrogen peroxide, H2O2, to oxygen gas. The reaction is done with potassium
manganate(VII) solution and hydrogen peroxide solution acidified with dilute sulphuric acid.
As with all equations we must keep both reactants and products balanced for charge and for atoms. To balance the charge
we will use electrons. If we need to balance oxygen’s we will use water (H2O has oxygen, this may seem strange at first,
but just go along for a little while). To balance hydrogen’s we will use H+).
So all you are allowed to add are:
1)
water
2)
hydrogen ions
3)
electrons
During the reaction, the manganate(VII), purple coloured ions are reduced to green manganese(II) ions.
Let's start with the hydrogen peroxide half-equation. What we know is: H2O2  O2
The oxygen is already balanced. What about the hydrogen?
All you are allowed to add to this equation are water, hydrogen ions and electrons.
If you add water to supply the extra hydrogen atoms needed on the right-hand side, you will mess up the oxygen’s again that's obviously wrong!
Add two hydrogen ions to the right-hand side.
H2O2  O2 + 2 H+
Now all you need to do is balance the charges. You would have to add 2 electrons to the right-hand side to make the
overall charge on both sides zero.
H2O2  O2 + 2 H+ + 2e- (loses e- , so is a reduction reaction)
Now for the manganate(VII) half-equation:
You know (or are told) that the manganate(VII) ions turn into manganese(II) ions. Write that down.
MnO4-  Mn 2+
The manganese balances, but you need four oxygen’s on the right-hand side. These can only come from water - that's the
only oxygen-containing thing you are allowed to write into one of these equations in acid conditions.
MnO4-  Mn 2+ + 4 H2O
By doing this, we've introduced some hydrogens. To balance these, you will need 8 hydrogen ions on the left-hand side.
8 H+ + MnO4-  Mn 2+ + 4 H2O
Now that all the atoms are balanced, all you need to do is balance the charges. At the moment there are a net 7+ charges
on the left-hand side (1- and 8+), but only 2+ on the right. Add 5 electrons to the left-hand side to reduce the 7+ to 2+.
5e- + 8 H+ + MnO4-  Mn 2+ + 4 H2O
(gains e- , so is an oxidation reaction)
A species which can accept electrons from another species is an oxidizing agent. Eg MnO4- is the oxidizing agent above
A species which can donate electrons to another species is a reducing agent.
Eg H2O2 is the Reducing Agent.
This is the typical sort of half-equation which you will have to be able to work out. The sequence is usually:
Balance the atoms apart from oxygen and hydrogen.
Balance the oxygen’s by adding water molecules.
Balance the hydrogen’s by adding hydrogen ions.
Balance the charges by adding electrons.
Combining the half-reactions to make the ionic equation for the reaction
The two half-equations we've produced are:
So to balance electrons we find the least common denominator.
So now we simplify
ELECTROCHEMICAL CELLS
So let’s follow the electrons now in actual working situations, in batteries.
1.
Consider a zinc rod immersed in a solution containing Zn2+ ions (eg ZnSO4)
zinc electrode
Zn2+ ions in solution
The Zn atoms on the rod can lose two electrons and move into solution as Zn2+ ions: Zn(s)  Zn2+(aq) + 2e
This process would result in an accumulation of NEGATIVE electrons charge on the zinc rod.
These electrons leave this rod and go to then second beaker, there a copper rod is immersed in a solution of copper ions
(eg CuSO4), due to the following processes: Cu2+(aq) + 2e Cu(s) - Reduction (rod becomes positive)
Electrochemical cells are spontaneous reactions that generates a current. There are two half-cells connected by a
wire, so that electrons flow from one metal (electrode - anode) to another through metal (electrode-cathode)an
external circuit and the ions flow through an internal cell connection (or salt bridge). The half-cell in which a halfreaction occurs with a loss of electrons (oxidation) is the anode. The second half-cell in which a half-reaction
occurs with a gain of electrons (reduction), is the cathode.
Oxidation-reduction reactions occur when electrons are given up by the substance being oxidized (the reducing agent) and
simultaneously gained by the substance being reduced (the oxidizing agent). Consider the following redox reaction to help
explain the electron transfer process:
Cu2+ + Zn  Zn2+ + Cu
In the diagram below, each beaker represents one of the two half cells for the above reaction. But, because there is no way
for electrons to move from one beaker to the other, a redox reaction cannot yet occur.
If Zn and Cu electrodes were connected as in shown in the following diagram, an oxidation-reduction reaction would
occur since electrons could flow through the external wire.
Creating an Electrochemical/ Voltaic /or Galvanic Cell
If two different electrodes are connected, the potential difference (Volts) between the two electrodes will cause a current
to flow between them. Thus an electromotive force (emf) is established and can generate electrical energy.
The circuit must be completed by allowing ions to flow from one solution to the other. This is achieved by means of a salt
bridge (salts are any ionic compound, eg: KCl).
The combination of two electrodes in this way is known as an electrochemical cell, and can be used to generate
electricity. The two components which make up the cell are known as half-cells.
-ve
V
+ve
Zn
Cu
Cu2+
Zn2+
ANODE (-ve)
CATHODE (+VE)
Thus electrons flow from the zinc electrode to the copper electrode.
(GERC) Reduction takes place at the copper electrode: Cu2+(aq) + 2e  Cu(s) and Cations to Cathode
(LEOA) Oxidation thus takes place at the zinc electrode: Zn(s)  Zn2+(aq) + 2e and Anions to Anode
The overall cell reaction is as follows: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
The sulphate ions (anions) flow through the salt bridge from the Cu2+(aq) solution to the Zn2+(aq) solution (anode), to
complete the circuit. The cell reaction including spectator ions can thus be written as follows:
CuSO4(aq) + Zn(s)  Cu(s) + ZnSO4(aq).
The external connection must be made of a metallic wire in order to allow electrons to flow. The salt bridge must be made
of an aqueous electrolyte to allow ions to flow oppositely and internally. By allowing 2 chemical reagents to be connected
electrically, but not chemically; chemical energy is thus converted into electrical energy.
Designing electrochemical cells
Half-cells do not necessarily have to consist of a metal immersed in a solution of its own ions.. If the half-reaction does
not contain a metal, an inert electrode must be used. Platinum, or carbon, is used, as they are inert. If a gas is involved, it
must be bubbled through the solution in such a way that it is in contact with the electrode. Examples are shown below:
ION SOLUTION HALF CELLS
a) Cr2O72-(aq) + 14H+(aq) + 6e  2Cr3+(aq) + 7H2O(l) (Note the reactants have no solids for the condiut of electrons
A platinum electrode is used, immersed in a solution containing Cr2O72-, H+ and Cr3+ ions:
Pt
mixture of
Cr2O72-, H+
and Cr3+
ION SOLUTION HALF CELLS
b)
2H+(aq) + 2e  H2(g)
A platinum electrode is used again, immersed in a solution containing H+ ions. Hydrogen gas is bubbled through
the solution, in contact with the electrode:
H2
Pt
H+
Standard conditions
The electrode potential depends on the conditions used, including temperature, pressure and concentration of reactants. It
is therefore necessary to specify the conditions used when measuring electrode potentials. These conditions are normally
set at a temperature of 298 K, a pressure of 1 atm and with all species in solution having a concentration of 1.0 mol/L.
Electrode potentials measured under these conditions are known as standard electrode potentials -denoted by Eo.
Reference electrodes (THE SHE)
The emf of electrochemical cells is easy to measure, but the individual electrode potentials themselves cannot actually
be measured at all; it is only possible to measure the potential difference between two electrodes.
It is therefore only possible to assign a value to a half-cell if one half-cell is arbitrarily allocated a value and all
other electrodes are measured relative to it. An electrode used for this purpose is known as a reference electrode.
The electrode conventionally used for this purpose is the standard hydrogen electrode.
H2
Pt
H+
The gas pressure is fixed at 1 atm, the temperature is 25oC and the H+ ions have a concentration of 1.0 moldm-3.
This electrode is assigned a value of 0.00V. Using this electrode, it is possible to assign an electrode potential to all other
half-cells.
Eg if the standard Zn2+(aq) + 2eZn(s) electrode is connected to the standard hydrogen electrode and the standard
hydrogen electrode is placed on the left, the emf of the cell is -0.76V.
The Zn2+(aq) + 2eZn(s) half-cell thus has an electrode potential of -0.76V.
Eg if the Cu2+(aq) + 2e  Cu(s) electrode is connected to the standard hydrogen electrode and the standard hydrogen
electrode is placed on the left, the emf of the cell is +0.34V.
The Cu2+(aq) + 2e  Cu(s) half-cell thus has an electrode potential of +0.34V.
The standard electrode potential of a half-reaction can be defined as follows:
"The standard electrode potential of a half-reaction is the emf of a cell where the left-hand electrode is the
standard hydrogen electrode and the right-hand electrode is the standard electrode in question".
The equation emf = ERHS - ELHS can be applied to electrochemical cells in two ways:
Conventional Representation of Cells
As it is cumbersome and time-consuming to draw out every electrochemical cell in full, a system of notation is used
which describes the cell in full, but does not require it to be drawn.
Half-cells are written as follows:
- the electrode is placed on one side of a vertical line.
- the species in solution, whether solid, liquid, aqueous or gas, are placed together on the other side of the vertical line.
- if there is more than 1 species in solution, and it is on different sides of the ½ equation, they are separated by a comma
Eg Fe3+(aq) + e  Fe2+(aq)
Pt Fe2+, Fe3+
When two half-cells are connected to form a full electrochemical cell, the cell is written as follows:
- the positive electrode is always placed on the right
- the two half-cells are placed on either side of two vertical broken lines (which represent the salt bridge)
- the electrodes are placed on the far left and far right
- on the left (oxidation)(LEO) is written, on the right (reduction) GER
Pt Fe2+, Fe3+ Ag+ Ag
Eg
Cell reaction: Ag+(aq) + Fe2+(aq)  Ag(s) + Fe3+ (aq)
Rechargeable and non-rechargeable cells
Electrochemical cells are the basis for all batteries. Batteries contain two separate half-cells. The solutions are connected
by a salt bridge which allows ions to flow through without allowing mixing of the solutions.
Reactions taking place in the half-cells are irreversible the battery is non-rechargeable and called a PRIMARY CELL
If the reactions taking place in the half-cells are reversible, the battery is rechargeable,& called a SECONDARY CELL.
Fuel Cells
A fuel cell is a cell in which a chemical reaction between a fuel and oxygen is used to create a voltage. The fuel and
oxygen flow into the cell continuously and the products flow out of the cell. Therefore the cell does not need to be
recharged. The most widely used fuel cell is the hydrogen-oxygen fuel cell:
A fuel cell, like a regular electrochemical cell, consists of two half-cells connected by a semi-permeable membrane. An
aqueous solution of sodium hydroxide is used as the electrolyte.
Oxygen is pumped into one of the half-cells:
O2(g) + H2O(l) + 4e- 4OH-(aq)
E0 = +0.40 V
Hydrogen is pumped into the other half-cell:
H2O(l) + 2e- H2(g) + 2OH-(aq)
E0 = -0.83 V (flip this)
The oxygen half-cell is more positive and therefore undergoes reduction. The H2 half-cell is more negative and oxidizes
O2(g) + H2O(l) + 4e-  4OH-(aq)
reduction
H2(g) + 2OH-(aq)  H2O(l) + 2eoxidation
___________________________________________________________
O2(g) + 2H2(g)  2H2O(l)
overall cell reaction, emf = 1.23 V
There are a number of advantages of fuel cells as a way of producing energy:
The hydrogen-oxygen fuel cell produces water as the only product.
-
Fuel cells are more efficient than combustion engines. Typically fuel cells are approximately 50% efficient but
combustion engines are approximately 20% efficient.
However there are also a number of limitations of fuel cells as a way of producing energy:
Hydrogen is flammable. It is therefore both difficult and dangerous to store and transport.
Fuel cells use toxic chemicals in their manufacture and Fuel cells have a limited lifetime
Electrolysis
Electrolysis: the process in which electrical energy is used to bring about a non-spontaneous chemical change. In both
types of cells, the electrode at which the reduction occurs is the cathode and oxidation occurs is the anode.
In an electrolytic cell, the flow of electrons is being ‘pushed’ by an outside source such as a battery. The cathode is called
the negative electrode of the electrolytic cell. The anode in the electrolytic cell is called the positive electrode.). We
SWITCH the signs
The differences between galvanic and electrolytic cells can be summarized in a table.
Galvanic/Voltaic Cells
chemical energy
Electrolytic Cells
electrical energy
electrical energy
chemical energy
two half-cells with separate electrolytes and a salt bridge
electrodes in the same electrolyte
chemical reaction is spontaneous Eo total is positive
chemical reaction is not spontaneous Eo total is
negative
anode - negative terminal : cathode - positive terminal
anode - positive : Cathode - negative
oxidation always occurs at the anode
oxidation occurs at the anode
Electrolysis of Molten Salts
Electrolysis is a process where electrons are forced through a chemical cell, thus causing a chemical reaction.
Reduction always takes place at the cathode. In the electrolysis of molten salt, NaCl, the cathode and anode reactions are:
Anode (oxidation): 2 Cl-  Cl2 + 2 eAnode oxidation
2 Cl-  Cl2 + 2 e-
Cathode (reduction): Na+ + e-  Na
Cathode reduction
2 Na+ + 2 e- -> 2 Na
+
2 Cl + 2 Na  Cl2 + 2 Na
If one mole of electrons (96485 C or 1 Faraday) passes from the anode to the cathode, one mole of Na (23 g) will be
deposited, and half a mole of chlorine gas Cl2 (or one mole of Cl atoms) will be collected at the anode.
Electrolysis of Water
Pure water does not conduct electricity, because the numbers of H+ and OH- ions are small (10-7 mol/L each). In the
presence of an acid, water can be decomposed.
A potential of -2.06 V is the standard cell potential for,
Pt | H2O, [H+] = 1 M | O2 || H2O [OH-] = 1 M | H2 | Pt
And when a potential greater than 2.06 V is applied, the following reactions take place.
V
Anode oxidation
Cathode reduction
H2O 4 H+ + 4 e + O2
-2.06 V
4 H2O + 4 e  2 H2 + 4 OH2 H2O  2 H2 + O2
The Hall-Heroult process
Aluminum (Al) is the third the most abundant elements on Earth crust, in the form of bauxite Al2O3. Because Al it is very
reactive, this metal remained unknown to mankind until 1827, bound up inside bauxite. In 1886, two young men working
in two continents apart electrolyzed molten cryolite Na3AlF6 (melting point 1000° C); their discovery is now known as the
Hall-Heroult process, which is a commercial process.
Faraday’s Stoichiometry
Electrolysis CREATES chemical reactions. The following examples illustrate the stoichiometry of electrolysis. The
charge on a mole of electrons is called the Faraday. The best estimate of the value of a Faraday, according to the National
Institute of Standards and Technology (NIST), is 9.65x104 coulombs per mole of electrons. The charge on an electron
based is 1.60 x 10-19 coulombs per electron. If you divide the charge on a mole of electrons by the charge on a single
electron you obtain a value of Avogadros number of 6.022 x 1023 particles per mole.
Problem :
Molecular masses can be determined through electroplating. Determine the molecular mass and identity of a +2 metal, X,
that plates 46.3g of X in 6.75 hours at a current of 2 A.
Cathodic Protection is used to control rusting of a metal surface by making it the cathode .The anode will be
destroyed/sacrificed. Galvanizing is a form of this cathodic protection whereby Fe (the usual metal we protect) is coated
with Zn, the sacrificed anode (this is called galvanizing) . Of course you could just paint the metal too!
Industrial Applications
From Mineral to Metal: Metallurgy and Electrolytic Refining
Metallurgy, the science of extracting and refining metals from their ores.
Processing Ore: Pyrometallurgy vs. Hydrometallurgy
Most metals are NOT found as free elements, they are bound to oxygen or carbon or sulfur, like FeCO3 (siderite) and
Al2O3.2H2O(bauxite). How these three are accomplished depends upon whether pyrometallurgy or hydrometallurgy is
used. Pyrometallurgy (or "pyromet" for short) uses extreme heat. Hydrometallurgy (or "hydromet") uses chemicals to
dissolve the metal from its ore. Long Harbour is a hydromet facility.
Copper Pyromet Versus Hydromet Pyrometallurgy of Copper
PYROMET (Heat): Copper ore is ground into a fine powder which is then added to a mixture of water, oil and detergent.
The oil preferentially sticks to the mineral particles and carries the mineral to the top. The foam, containing the mineral
(chalcopyrite (CuFeS2), is skimmed off the surface, this process is called floatation.
The second step is to heat the mineral in a furnace the copper and iron separate
2 CuFeS2(s) + 4 O2(g) + HEAT  Cu2S(l) + 2 FeO(l) + 3 SO2(g)
Equation 1
Electrolytic purification of copper.
Because Zn and Fe (& other trace minerals) are impurities; we need to purify them from the Cu. The final step in
obtaining highly pure copper metal involves using an electrolytic cell to refine the copper.
The voltage applied to the electrodes is just high enough to oxidize the copper atoms at the anode to Cu+2 . The Cu+2 enter
the solution and go to the cathode to be reduced to pure copper metal.
Hydrometallurgy of Copper
Once flotation is completed, the major difference with hydromet processes is leaching instead of roasting. Leaching refers
to the dissolving of the desired metal in ACID (sulfuric acid)
Equation 2
2 CuFeS2(s) + H2SO4(aq) + 4 O2(g)  2 CuSO4(aq) + Fe2O3(s) + 3/8 S8(s) + H2O(l)
Next, the iron (III) oxide and sulfur are filtered off. The solution is placed in an electrolytic cell to get pure copper.
Equation 5
cathode: Cu2+ (aq) + 2 e-  Cu(s)
Summary: Pyrometallurgy -efficient, inexpensive, proven technology readily available
Hydrometallurgy - fewer impurities in solution, no SO2(g) produced thereby limiting acid rain.
Rusting
A piece of bare iron left outside where it is exposed to moisture will rust quickly, especially near salt.
Considering the sketch of a water droplet, the oxidizing iron supplies electrons at the edge of the droplet to reduce
oxygen from the air. The iron surface inside the droplet acts as the anode: Fe(s) -> + Fe2+(aq) + 2eThe electrons move through the iron to the outside of the droplet where: O2(g) + 2H2O(l) + 4e- -> 4OH-(aq)
The rusting of unprotected iron in the presence of air and water is then inevitable because it is driven by an
electrochemical process. However, other electrochemical processes can offer some protection against corrosion.
For magnesium rods can be used to protect underground steel pipes by a process called cathodic protection.
Types of Batteries
Batteries are a set of galvanic cells connected in series
- The negative electrode of one cell is connected to the positive electrode of the next cell
- The total voltage for a set of cells is the sum of the individual voltages.
There are two general types of batteries: primary batteries and secondary batteries.
A. Primary battery: a disposable battery. 3 types of primary batteries:
1. Dry Cell Batteries 1.55 V.
A galvanic cell with the electrolyte contained in a paste
The cheapest AAA, AA, C, and D-size batteries are dry cells.
Dry cells stop producing electricity when the reactants are used up.
Dry cells consist of a Zn(s) anode and an inert graphite cathode.
The electrolyte is a paste of MnO2, ZnCl2, NH4Cl and “carbon black”,
C(s) (soot)
The half-cell reactions are:
LEOA: Zn (s)  Zn2+(aq) + 2 eGERC:
2 MnO2 (s) + H2O(l) + 2 e-  Mn2O3(s) + 2 OH- (aq)
Overall cell reaction: 2 MnO2(s) + Zn(s) + H2O(l)  Zn2+(aq) + Mn2O3(s) + 2 OH- (aq)
2. Alkaline Cell Battery
These are slightly more expensive than dry cell batteries. They are also improved and longer lasting.
The NH4Cl and ZnCl2 of dry cells are replaced by KOH (hence the name alkaline!)
3. Button Cell Battery
These are much smaller than alkaline batteries and are
commonly used in watches and pacemakers.
Two common types of button cell batteries both have a
zinc container that acts as the anode and an inert
stainless steel cathode.
B. Secondary battery: a rechargeable battery
Two types of secondary batteries are lead storage and
Ni-Cd batteries.
1. Lead Storage Batteries (Lead-Acid Batteries)
These batteries are typically found in your car and consist of 6 cells with a total potential of ~ 12 V. When in use the cells
in the battery operate like a galvanic cell – they partially discharge and release electricity. This supplied electricity
(external voltage) forces the reverse, non-spontaneous, cell reaction. The cell is now recharged.
Lead storage batteries have anodes consisting of powdered lead packed into a grid and cathodes consisting of powdered
PbO2 packed into another grid. The electrolyte is a 4.5 mol/L solution of H2SO4 (aq).
2. Ni-Cd (Nicad) Battery
These usually contain three cells in series with a total potential of ~ 1.4 V. The anode is pressed cadmium, the cathode is
NiO(OH) and the electrolyte is KOH or NaOH.