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```OXIDATION NUMBERS
If a redox equation is very complicated, it is sometimes hard to work out what has
been reduced and what the reducing agent is. When this happens you can use a sort
of ‘electronic book-keeping’, called oxidation numbers, to make the task easier.
Oxidation numbers are assigned to atoms, in elements, compounds or ions,
according to a set of rules.
Once you have assigned them you can easily decide whether oxidation or reduction
has taken place:
When an element increases its oxidation number it has been oxidised.
When an element decreases its oxidation number it has been reduced.
If no species changes its oxidation number, the equation does not represent a
redox reaction.
Rules for assigning oxidation numbers.
1]
The oxidation number of an atom in a free, uncombined element is zero. E.g.
H2, O2, Mg, Na, S8
2]
The oxidation number of an atom in a simple monatomic ion is equal to the
charge on the ion. E.g. O2- O.N. of oxygen = -2, Mg2+ O.N. of magnesium =
+2.
3]
In a polyatomic ion the sum of the oxidation numbers of the atoms in that ion
equals the charge on the ion. E.g. SO42- the O.Ns of sulfur and oxygen must
4]
In compounds, the sum of the oxidation numbers of all atoms in the
5]
The oxidation number of oxygen in compounds is –2 except in peroxides,
where it is –1. (The only peroxide you are likely to meet is hydrogen peroxide
H2O2.)
6]
The oxidation number of hydrogen in compounds is +1 except in the case of
hydrides where it is –1. (A hydride is a compound containing a reactive metal
and hydrogen e.g. lithium hydride, LiH).
Example:
Determine the oxidation number of each atom in the equation:
2Mg + O2
2MgO
Both magnesium and oxygen on the left are elements, so Mg and O both have an
oxidation number of 0 (rule 1).
After reaction the product is ionic. Magnesium will have an oxidation
number of +2, since it is in the form Mg2+ (rule 2). The oxidation number of oxygen
will be –2 (rule 5).
0
0
2
2
2 Mg 2  O 2  2 Mg O
Test Yourself:
1]
Write oxidation numbers for every atom in the following:
a. Hg
b. HNO3
c. NH3
d. MnO2
3+
e. Cr
f. H2O2
g. NO2
h. SO422]
Which of the following are redox reactions? Identify the
oxidised and reduced species in each case.
a. CaO(s) + SiO2(s)
CaSiO3(s)
b. 2Rb(s) + Br2(l)
c. 2HI(aq) + H2O2(aq)
d. H2SO4(aq) + 2KOH(aq)
2RbBr(s)
I2(s) + 2H2O(l)
K2SO4(aq) + 2H2O(l)
BALANCING REDOX EQUATIONS
Write the formula for
the reactant and its
product
Identify a reactant and
its product
Are atoms other than
hydrogen & oxygen
balanced ?
No
Balance atoms other
than hydrogen and
oxygen
Yes
Are the oxygen
atoms balanced ?
No
Balance oxygen atoms
molecules
No
Balance the hydrogen
ions
Repeat for other
reactant & product
Yes
Are the hydronen
atoms balanced ?
Yes
Are the charges
balanced ?
Yes
Are the number of
electrons in the two half
equations equal ?
Add two half equations to get
net ionic equation.
Yes
Balance the charges by
side with the greatest
positive charge.
No
No
Balance electrons by
multiplying by
appropriate coefficients
When the reaction is carried out in alkaline conditions we have to put in an extra step
to the balancing of the equation.
Once the H has been balanced by adding H+, add an equal amount of OH- ions to
both sides of the equation, to convert all H+ to H2O molecules. (Then simplify by
cancelling out water molecules.)
Write a balanced equation for the following redox reactions:
--
1]
Cr(OH)4 + Na2O2
2]
MnO4 + HSO3
--
--
--
--
CrO4 + OH
2--
MnO2 + SO4
(Basic Conditions)
(Basic Conditions)
OXIDISING AGENTS
You must know the species involved in each of these reactions, their colours and be able to write balanced half
equations for each reaction.
SUBSTANCE
Hydrogen peroxide
HALF-EQUATION
H2O2 + 2H+ + 2e
2H2O
colourless
Halogen, X2 (Cl2, Br2 or I2)
colourless
2X-
X2 + 2e
Cl2 (pale green)
colourless
Br2 (deep red but orange-red in solution)
I2 (purplish black but brown in aqueous
solution)
Permanganate ion (acid solution)
Permanganate ion (neutral solution)
Permanganate ion (strongly alkaline solution)
Dichromate ion (acid solution
Nitric acid (concentrated)
Nitric acid (dilute)
MnO4 - + 8H+ + 5e
Mn2+ + 4H2O
purple
colourless
MnO4 - + 4H+ + 3e
MnO2 + 2H2O
purple
black
MnO - + e
MnO42-
purple
green
Cr2O72- + 14H+ + 6e
2Cr3+ + 7H2O
orange
green
NO3 - + 2H+ + e
NO2 + H2O
colourless
brown
NO3 - + 4H+ + 3e
colourless
NO + 2H2O
colourless *
(*Instantly reacts with oxygen in the air to form
brown nitrogen dioxide gas)
Fe3+ + e
Iron (III) salt
Copper (II) salt
Iodate (or bromate) ion
Fe2+
yellow
pale green
Cu2+ + 2e
Cu
blue
pink / brown
IO3 - + 6H+ + 6e
I - + 3H2O
colourless
colourless
REDUCING AGENTS
SUBSTANCE
Zinc
Zn
silvery
Magnesium
Mg
silvery
Iron
Hydrogen
Sulfur dioxide
Iron (II) salt
Sulfite ion
Iodide ion
colourless
Mg2+ + 2e
colourless
Fe
Fe2+ + 2e
grey
pale green
H2
2H+ + 2e
colourless
colourless
SO2 + 2H2O
SO42- + 4H+ + 2e
colourless
colourless
Fe2+
Fe3+ + e
pale green
yellow
SO32- H2O
SO42- + 2H+ + 2e
colourless
Thiosulfate ion
HALF-EQUATION
Zn2+ + 2e
2S2O32-
colourless
S4O62- + 2e
colourless
colourless
2I -
I2 + 2e
colourless
brown
ELECTROCHEMICAL CELLS
All chemical reactions involve some sort of energy change. Normally this is an
exchange of heat (thermal energy) between the system and its surroundings. If the
reaction is exothermic the system is a source of thermal energy. In redox reactions
(electron –transfer reactions) the energy transfer can either be thermal energy or
electrical energy.
e.g. the reaction between zinc metal and a copper salt. If zinc metal is added to a
solution of copper ions, copper metal and zinc ions are formed and the solution
becomes warm:
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
H is negative
Carried out in this way the reaction takes place at the surface of the zinc metal
where electrons are transferred directly from zinc atoms to copper ions.
If, however, the reactants are kept apart, the electrons can be led through a wire
instead of being transferred directly. The energy appears as electrical energy. A
device which does this is called an electrochemical cell.
An early electrochemical cell was the Daniell cell. It was based on the reaction
between zinc and copper ions and was once used to operate railway signals,
telegraph relays and doorbells.
In the inner compartment: Zn(s)
Zn2+(aq) + 2e- oxidation
In the outer compartment: Cu2+(aq) + 2eOverall reaction: Zn(s) + Cu2+(aq)
Cu(s) reduction
Cu(s) + Zn2+(aq) redox
[The zinc electrode is coated with mercury to stop it reacting with the sulfuric acid.]
The potential of the Daniell cell is 1.10 volts.
In principle any oxidation-reduction reaction can be used to operate an
electrochemical cell, but many have difficulties associated with them.
Cells and Half Cells



An electrochemical cell consists of two half-cells connected by a salt bridge.
A half-cell, is an electrode and the couple it is in contact with, and is where the
reaction takes place; oxidation in one, reduction in the other.
Each half-cell (or electrode) consists of:
(a) the oxidised and reduced forms of one of the reactants, one
or both of them in water solution. In identifying a half-cell
both forms, in the order oxidised form / reduced form,
must be stated. E.g. Zn2+/ Zn, Cu2+/ Cu, Fe3+/ Fe2+ etc
(b) A piece of metal through which the electrons pass to or from
the external circuit.
A voltage or potential difference is generated between the metal and the solution.
The electrode potential of a metal, measured in volts, and is caused by the charge
built up on the metal rod due to the dynamic equilibrium established between the
metal atoms and the metal ions in solution.
M+(aq) + e
 M(s)
The position of the equilibrium, the charge and size of the electrode potential,
depends on:
 The nature or reactivity of the metal – the electrode potential is more negative
for more reactive metals as they have a greater tendency to lose electrons to
form stable ions.

The ionic concentration [M+] – increasing the concentration of [M+] will alter
the equilibrium in favour of the formation of M(s) , so the electrode potential
becomes more positive.

The temperature – changes in temperature affect exothermic and
endothermic reactions differently.
It is impossible to measure the true potential difference of a metal in a solution of its
ions because there is no complete circuit, so electrodes or half-cells must be
connected to a second (reference) electrode to give a comparison of energy
difference (voltage).
One half-cell will tend to lose electrons and the other half-cell will tend to gain these
electrons: this allows the relative reduction potential of each half-cell to be
calculated.
THE STANDARD HYDROGEN ELECTRODE
The hydrogen electrode is used as a standard or reference electrode to measure the
electrode potential of any half-cell. It is given the reference potential of 0.00 V.
2H+(aq) + 2e - 
H2(g)
Eo = 0.00 V
Eo is measured under standard conditions:
 All solutions are 1 mol L-1
 all gas pressures are 1 atm or 100 kPa
 temperatures are 25oC (298 K)
An inert electrode is used to carry the electrons to and from the hydrogen half-cell,
as neither hydrogen gas nor ions can carry electrons. Platinum is the metal used
because
 it is inert
 it catalyses the reaction that produces hydrogen
+
-2H (aq) + 2e
H2(g)
Once the standard hydrogen electrode is set up, and because it has a reference
potential of 0.00 V , it can be used to determine the electrode potential of all other
half-cells.
Pure H2(g) at 298K and
1 atm
Platinum electrode
Acid solution containing
1.0 mol L-1 H+(aq), 298 K
Holes in side of glass casing to allow
Hydrogen gas to escape so pressure is
Maintained at 1 atm.
Measuring the Electrode Potential of a Half-Cell
The half-cell whose Eo value is to be determined is set up with the hydrogen half-cell
to make a complete cell. A salt bridge is used to connect the two half-cells and the
potential difference of the cell is measured, using a voltmeter. Since the hydrogen
half-cell is a reference half-cell with an electrode potential of zero, the reading on the
voltmeter is solely due to the other half-cell.
Half-Cell Conventions
All half-cells consist of a species in its oxidised and reduced states and an electrode,
which takes electric current to or away from the half-cell.
Half-cells are written with the oxidised form of the species first.
There are three types of half-cells:
 A metal rod and its aqueous solution
 A gas and its aqueous ions
 Two aqueous ions in solution together
A Metal and its Aqueous Ions
For metal/metal ion half-cells the metal itself acts as an electrode.
The half-cell for a metal M(s) and its aqueous ions M+(aq) is written as: M+/M
(oxidised form / reduced form). The slash (/) indicates phase boundary between
reactants.
For the zinc / zinc ions half-cell, written as Zn2+/Zn (oxidised form first):
 The half-reaction is Zn2+(aq) + 2e Zn(s)
Eo = -0.76 V
 Zn2+(aq), the oxidised form of zinc can act as an oxidant
 Zn(s), the reduced form of zinc, can react as a reductant
 The solid zinc metal Zn(s) acts as the electrode.
Eo values are given for the left to right direction in equations. Because E o is negative
in the Zn2+| Zn half-cell, the reaction which usually occurs is:
Zn(s)
Zn2+(aq) + 2e –
A Gas and its Aqueous Ions
Gas / aqueous ion half-cells require an inert electrode to carry the current in or out of
the half-cell, because gases cannot form conducting electrodes. Platinum Pt, or
carbon C (in the form of graphite), are often used as electrodes.
The half-cell for a gas X2 (g) and its aqueous ions X- (aq) is written as:
X2 (g) , X –(aq)/ inert electrode. (The / is used to separate the conductor from the rest
of the half cell.)
For the chlorine/chloride half-cell, written as,
Cl2 (g), Cl –(aq) / Pt (or C):
 The half-reaction is Cl2 (g) + 2e2Cl-(aq) Eo = +1.36 V
 Cl2(g) is the oxidant
 Cl-(aq) is the reductant
Because Eo is positive in the Cl2, Cl-(aq) half-cell, the reaction which usually occurs
is:
Cl2(g) + 2e2Cl-(aq)
Two Types of Aqueous Ions
Ions can be oxidised and reduced to form other ions, so aqueous ion / aqueous ion
half-cells exist. Ions, like gases cannot form conducting electrodes, so an inert
electrode is used in these half-cells.
The half-cell for the oxidised for of an aqueous ion (M2+(aq)) and the reduced form of
the aqueous ion (M+(aq)) is written as:
M2+(aq), M+(aq) / inert electrode. (The comma is used to separate the two reactants
in the same solution and the / indicates a phase separation between the solution and
the electrode.)
For the iron (III) | iron (II) half-cell, written as
Fe3+(aq), Fe2+(aq) / Pt (or C):
 The half-reaction is Fe3+(aq) + eFe2+(aq)
3+
 Fe (aq) is the oxidant
 Fe2+(aq) is the reductant
Eo = +0.77 V
Because Eo is positive in the Fe3+(aq),Fe2+(aq) half-cell, the reaction that usually
occurs is:
Fe3+(aq) + eFe2+(aq)
Electrode Potential Values
The values of electrode potential provide information on:
 The strength of metals as reducing agents (high negative E o values)
 The strength of non-metals as oxidising agents (high positive Eo values)
 The prediction of the spontaneity of redox reactions. If the overall value of E o for a
cell is positive, the reaction will occur spontaneously.
 Calculating the EMF or voltage of electrochemical (galvanic cells)
Reducing Agents and Eo values
Electrode potentials involving metals indicate:
 How readily hydrated metal ions in solution gain electrons (are reduced) to form
metal ions. The more positive the Eo value, the more likely it is that the metal ion
will form the metal.
 How readily metal atoms release electrons (are oxidised) to produce hydrated
ions in solution. The most reactive metals (those that are the best reductants), are
those that have the most negative electrode reduction potential values.
Arranging the metals in order of their electrode potential (Eo) values produces the
electrochemical series. Electrode potentials are good indicators of metal
reactivities. The two exceptions are lithium and calcium, which are chemically less
reactive than their Eo values would indicate.
Electrochemical
series
Lithium
Li
Potassium K
Calcium
Ca
Sodium
Na
Magnesium Mg
Aluminium Al
Zinc
Zn
Iron
Fe
Pb
Copper
Cu
Silver
Ag
Gold
Au
Metal ion/atom
half-equation
Li+ + e  Li
K+ + e  K
Ca2++ 2e  Ca
Na+ + e  Na
Mg2+ + 2e  Mg
Al3+ + 3e  Al
Zn2+ + 2e  Zn
Fe2+ + 2e  Fe
Pb2+ + 2e  Pb
Cu2+ 2e  Cu
Ag+ + e  Ag
Au+ + e  Au
Electrode potential
Eo / V
-3.02
-2.92
-2.87
-2.71
-2.30
-1.66
-0.76
-0.47
-0.13
+0.34
+0.81
+1.50
Metals with negative Eo values can reduce aqueous hydrogen ions to hydrogen gas.
If the electrode potential is positive, the metal atoms do not reduce hydrogen ions
and so they do not react with acids.
Oxidising Agents and Eo values
Electrode potentials involving non-metals indicate how readily ions or molecules
accept electrons (are reduced) to form other species. The most reactive non-metals,
(those that are the best oxidants), have the most positive electrode potentials values.
The Eo values for the common oxidising agents give a good indication of their
strengths as oxidants.
Electrochemical series
Fluorine
Chlorine
Oxygen
Bromine
Iodine
Nitrogen
Sulfur
F
Cl
O
Br
I
N
S
Non-metal reduction
half-equation
F2 + 2e  2FCl2 + 2e  2ClO2 + 4H+ + 4e  2H2O
Br2 + 2e  2BrI2 + 2e  2IN2 + 6H+ + 6e  2NH3
S + 2H+ + 2e  H2S
Electrode potential
Eo / V
+2.87
+1.36
+1.23
+1.07
+0.54
+0.27
+0.14
SUMMARY OF REDOX THEORY
 Redox reactions involve competition for electrons. When a species is reduced it
gains electrons; when a species is oxidised it loses electrons.
 When electricity flows through a molten salt, or through a solution of an
electrolyte, the substance is split up in a chemical process called electrolysis. At
the anode, anions (negatively charged ions) are oxidised. At the cathode,
cations (positively charged ions) are reduced.
 The chemical changes at each electrode are summarised using half-equations.
 The overall cell reaction can be found by adding together the two half-equations.
 Electricity may be produced whenever two metals are immersed in conducting
solutions, forming an electrochemical cell.
 Each half-cell reaction has its own tendency to attract electrons, which is
measured by the electrode potential E of the half-cell.
 The half-cell against which all other electrode potentials are measured is the
standard hydrogen half-cell or standard hydrogen electrode.
 The standard electrode potential Eo of a half-cell is the electrode potential of
that half-cell relative to the standard hydrogen electrode under standard
conditions. Eo is a measure of the oxidising or reducing power of the species in it.
 The standard conditions for electrochemical measurements are: all solutions have
a concentration of 1 mol L-1 and all measurements are made at a temperature of
25oC and 101.3 kPa pressure.
 Standard electrode potentials can be used to predict the likelihood of a reaction
proceeding spontaneously. If the Eo value for a cell is positive the reaction will be
spontaneous.
Cell voltage- the potential
of the right-hand electrode
with respect to the lefthand electrode
Half-cell notation
E.g. zinc-copper cell
Zn(s)/ Zn2+(aq)  Cu2+(aq) /Cu(s) Eo = +1.10V
Left-hand
electrode
Right-hand
electrode
Electrolyte in
contact with left –
hand electrode
Salt
bridge
Electrolyte in
contact with right –
hand electrode
In this reaction the cell voltage is positive i.e. the right-hand electrode is positive with
respect to the left-hand electrode. Electrons flow from the zinc to the copper. By
convention the cell notation always refers to the reaction taking place from left to
right.
Eocell = ERHE - ELHE
The salt bridge:
Without a salt bridge the half-cell containing the zinc would slowly become positively
charged as the electrons left it. The copper half-cell would slowly become negatively
charged. With the salt bridge present ions are able to move in and out of the
solutions keeping both half-cells electrically neutral. The salt used in the salt bridge
is chosen so that it does not react with the ions in either half-cell.
Oxidising and Reducing agents:
The standard electrode potential of a half-cell is a measure of the oxidising or
reducing power of the species in it – i.e. their ability to compete for electrons. In
general the stronger an oxidising agent, the more positive its electrode potential. A
strong reducing agent has a large negative electrode potential.
When looking at a table of standard electrode potentials for a number of half
reactions, the substances to the left of the double arrows are oxidising agents,
becoming reduced when they react in the direction left to right. The best oxidising
agent have a large positive Eo value and are on the left of the equations. Substances
to the right of the double arrows are reducing agents, becoming oxidised when they
react in the direction right to left. The best reducing agents are on the right of the
equations and have a large negative Eo value.
Ag+(aq) + e -  Ag(s)
E.g.
Oxidising agent
 Reduced during left
to right reaction
 Also known as
oxidant
 The more positive
the Eo value, the
more powerful the
oxidising agent
Reducing agent
 Oxidised during right to
left reaction
 Also known as
reductant
 The more negative the
Eo value, the more
powerful the reducing
agent
Half-Cell conventions:
In a cell diagram the following conventions are used:
 A single / represents a phase boundary between reactants e.g. between a
solid and a solution
 A double vertical line represents a boundary between two solutions (usually
a salt bridge)
 A comma is used to separate two reactants in the same solution, or a gas
from an aqueous ion.
 The conductors are written on the extreme left and right of the cell diagram.
When there is no metal in a redox couple, an electrical connection is made using an
inert electrode such as platinum or carbon. The half-cell is then written as: e.g.
Fe3+(aq),Fe2+(aq) / Pt(s). The least oxidised species present is written next to the
electrode.
The hydrogen electrode is written as: 2H+(aq) , H2(g) / Pt(s).
Uses of Electrode Potentials:
 Calculation of predicted cell emfs (voltages)
 Prediction of preferred direction of redox reactions
 Comparison of relative strengths of oxidising and reducing agents
ELECTROLYSIS
CHEMICAL REACTIONS DRIVEN BY ELECTRICITY
You have already seen that when a spontaneous redox reaction occurs, an electric
current is produced.
If we use electric current from a battery or power supply and pass it through a
conducting solution redox reactions will be produced. This process is called
electrolysis.
During electrolysis electrical energy is converted to chemical energy. The
reactions that occur in electrolytic cells are essentially the opposite to those occurring
in electrochemical cells.
Electrolytic Cell
Electrical
Energy
Chemical
Energy
Electrochemical
Cell
Reactions in electrolytic cells would not normally happen without the application of
electrical energy and are called non-spontaneous reactions. Chemicals formed by
electrolysis are often difficult to obtain by other means and many useful materials are
manufactured by this process. Important applications of electrolysis include:






plating a thin film of metal on surfaces of other metals to improve appearance
or prevent corrosion. (Electroplating)
extraction of reactive metals such as sodium, magnesium and aluminium from
their ores
industrial preparation of sodium hydroxide, chlorine and hydrogen.
recharging of car batteries and other rechargeable cells such as Ni-Cads.
refining copper metal
increasing the thickness of the surface oxide layer on aluminium metal to
improve its resistance to corrosion.
In an electrolytic cell the reactions that occur and redox reactions. Consider the
electrolysis of a solution of zinc iodide. If an electric current is passed through a
solution of zinc iodide, a grey deposit is formed on the cathode (negative electrode)
and a brown colour forms at the anode (positive electrode). It is easy to show that
the grey deposit is zinc and the brown substance is iodine. The reactions that have
occurred are:
At the cathode:
At the anode:
Zn
2+
(aq)
+ 2e
--
2I (aq)
--
Zn(s)
I2(aq) + 2e
--
The cathode reaction is reduction and the anode reaction is oxidation.
In electrolysis
 Oxidation occurs at the anode
 Reduction occurs at the cathode
Questions
1. What is the main difference between electrochemical and electrolytic cells?
2. What happens during electrolysis ?
3. Give at least three practical uses of electrolysis.
COMPARISON OF ELECTROLYTIC AND ELECTROCHEMICAL
CELLS
Electrolytic cells
A non-spontaneous reaction.
Reactions are forced by an applied
voltage which must be greater than that
which the cell can produce.
Electrons flow from the positive
electrode.
Converts electrical energy to a
chemical reaction.
Electrodes are usually immersed in a
common electrode.
Electrochemical cells
A spontaneous reaction.
Electrons flow from the negative
electrode.
Converts a chemical reaction to
electrical energy.
Two half-cells are often used with
separate electrodes.
Electrolytic cell:
External voltage source
pushes electrons through circuit.
Electrons
absorbed
from the
external
circuit
(cathode)
Migration of negative ions
Electrons
released
to external
circuit
(anode)
Migration of positive ions
Electrochemical cell:
Electrons flow through metallic
conductor of external circuit
Electrons
absorbed
Electrons
released
(anode)
Migration of negative ions
Migration of positive ions
(cathode)
Similarities:
 The electrolyte is the substance that conducts electricity within the cell.
Electric charge is carried by anions to the anode and by cations to the
cathode.
 The electrode where oxidation occurs is called the anode.
 The electrode where reduction occurs is called the cathode.
 In the external circuit, electrons travel through the wire from the anode to the
cathode.
Key ideas of electrolysis:
The products of electrolysis depend on:
 The reactivity of the ions present
 The type of electrode used
 The concentration of the reactants
 The amount of current used
 The temperature
e.g. molten NaCl produces Na and Cl2
dilute aqueous NaCl produces H2 and O2 (as for H2O)
For electrolytic cells:
 Oxidation takes place at the anode (sign +ve)
 Reduction takes place at the cathode (sign –ve)
 Electrons flow from the anode to the cathode
 Electric current is used to cause a chemical reaction
 The applied voltage must be greater then the voltage of the cell
 If more than one reaction is possible the following rules apply
Metal ions:
Whether a metal is produced depends on its position in the reactivity series relative
to the position of water. In a concentrated solution the metal will form if it is below
water in the activity series. If it is above water, hydrogen gas forms. If there are 2
possible products, the one lower on the activity series will form.
Non-Metal ions:
A simple ion will always form its element. If a polyatomic ion is present, oxygen will
form in preference.
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