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Redox reaction, Electrolysis and Electrochemistry
group 8
1. Rini Kurniasih K1310069
2. Susi Cahyanti K1310076
3. Wahyu Nugroho K1310082
Mathematics Education
Training Teacher and Education Faculty
SEBELAS MARET UNIVERSITY (UNS)
2011
2
CONTENTS
I. TITLE .............................................................. 3
II. OBJECTIVES .............................................................. 3
III. BASIC THEORY ............................................................... 3
III.1 REDOX REACTION ............................................................... 3
III.2 ELECTROLYSIS ............................................................... 5
III.3 ELECTROCHEMISTRY ............................................................... 5
IV. EQUIPMENTS AND MATERIALS ......................................................... 8
IV.1 REDOX REACTION ............................................................... 8
IV.2 ELECTROLYSIS ............................................................... 9
IV.3 ELECTROCHEMISTRY ............................................................ 10
V. PROCEDURE ............................................................. 12
V.1 REDOX REACTION ............................................................. 12
V.2 ELECTROLYSIS ............................................................. 12
V.3 ELECTROCHEMISTRY ............................................................. 12
VI. OBSERVATION DATA ............................................................. 13
VI.1 REDOX REACTION ............................................................. 13
VI.2 ELECTROLYSIS ............................................................. 13
VI.3 ELECTROCHEMISTRY ........................................................... 13
VII. DATA ANALYZE ............................................................. 14
VII.1 REDOX REACTION ............................................................. 14
VII.2 ELECTROLYSIS ............................................................. 16
VII.3 ELECTROCHEMISTRY ........................................................... 19
VIII. CONCLUSION ............................................................. 21
IX. BIBLIOGRAPHY ............................................................. 22
3
PAPER OF BASIC CHEMISTRY 2 EXPERIMENT
I. TITLE
Redox Reaction, Electrolysis and Electrochemistry
II. OBJECTIVES
1. Knowing the redox reaction
2. Studying about chemistry reaction by elecrtric current
3. Knowing the oxidation reaction and reduction reaction zinc
metal with cuprum in solution and measure cell potential
III. BASIC THEORY
III.1 REDOX REACTIO N
Redox (shorthand for REDuction-OXidation) reactions describe all
chemical reactions in which atoms have their oxidation number (oxidation
state) changed. This can be either a simple redox process, such as the
oxidation of carbon to yield carbon dioxide (CO2) or the reduction of
carbon by hydrogen to yield methane (CH4), or a complex process such as
the oxidation of sugar (C6H12O6) in the human body through a series of
complex electron transfer processes.
The term comes from the two concepts of reduction and oxidation.
It can be explained in simple terms:
Oxidation is the loss of electrons or an increase in oxidation state
by a molecule, atom, or ion.
Reduction is the gain of electrons or a decrease in oxidation state
by a molecule, atom, or ion.
Though sufficient for many purposes, these descriptions are not
precisely correct. Oxidation and reduction properly refer to a change in
oxidation number — the actual transfer of electrons may never occur.
4
Thus, oxidation is better defined as an increase in oxidation number, and
reduction as a decrease in oxidation number. In practice, the transfer of
electrons will always cause a change in oxidation number, but there are
many reactions that are classed as "redox" even though no electron transfer
occurs (such as those involving covalent bonds).
Non-redox reactions, which do not involve changes in formal
charge, are known as metathesis reactions.
Reduction potential is used to calculate the standard electrode potential
(Eocell).
This is the equation most commonly seen in textbooks:
where:
Eocell is the standard electrode potential (in volts).
Eored is standard reduction potential of the reducing agent.
Eooxi is negative of the standard reduction potential of the
oxidizing agent.
though the following equation is generally more useful as one is usually
only given reduction potentials, not oxidation potentials :
or equivalently:
where:
Eocell is the standard electrode potential (in volts).
Eocathode is standard reduction potential of the reducing agent.
Eoanode is the standard reduction potential of the oxidizing agent.
5
III.2 ELECTROLYSIS
Electrolysis is the passage of a direct electric current through an
ionic substance that is either molten or dissolved in a suitable solvent,
resulting in chemical reactions at the electrodes and separation of
materials.
The main components required to achieve electrolysis are :
An electrolyte : a substance containing free ions which are the
carriers of electric current in the electrolyte. If the ions are not
mobile, as in a solid salt then electrolysis cannot occur.
A direct current (DC) supply : provides the energy necessary to
create or discharge the ions in the electrolyte. Electric current is
carried by electrons in the external circuit.
Two electrodes : an electrical conductor which provides the
physical interface between the electrical circuit providing the
energy and the electrolyte
Electrodes of metal, graphite and semiconductor material are
widely used. Choice of suitable electrode depends on chemical reactivity
between the electrode and electrolyte and the cost of manufacture.
III.3 ELECTROCHEMISTRY
Electrochemistry is a branch of chemistry that studies chemical
reactions which take place in a solution at the interface of an electron
conductor (a metal or a semiconductor) and an ionic conductor (the
electrolyte), and which involve electron transfer between the electrode and
the electrolyte or species in solution.
6
If a chemical reaction is driven by an external applied voltage, as in
electrolysis, or if a voltage is created by a chemical reaction as in a battery,
it is an electrochemical reaction. In contrast, chemical reactions where
electrons are transferred between molecules are called oxidation/reduction
(redox) reactions. In general, electrochemistry deals with situations where
oxidation and reduction reactions are separated in space or time, connected
by an external electric circuit to understand each process.
An electrochemical cell is a device that produces an electric current
from energy released by a spontaneous redox reaction. Electrochemical
cells have two conductive electrodes (the anode and the cathode). The
anode is defined as the electrode where oxidation occurs and the cathode is
the electrode where the reduction takes place. Electrodes can be made
from any sufficiently conductive materials, such as metals,
semiconductors, graphite, and even conductive polymers. In between these
electrodes is the electrolyte, which contains ions that can freely move.
The Galvanic cell uses two different metal electrodes, each in an
electrolyte where the positively charged ions are the oxidized form of the
electrode metal. One electrode will undergo oxidation (the anode) and the
other will undergo reduction (the cathode). The metal of the anode will
oxidize, going from an oxidation state of 0 (in the solid form) to a positive
oxidation state and become an ion. At the cathode, the metal ion in
solution will accept one or more electrons from the cathode and the ion's
oxidation state is reduced to 0. This forms a solid metal that
electrodeposits on the cathode. The two electrodes must be electrically
connected to each other, allowing for a flow of electrons that leave the
metal of the anode and flow through this connection to the ions at the
surface of the cathode. This flow of electrons is an electrical current that
can be used to do work, such as turn a motor or power a light.
A Galvanic cell whose electrodes are zinc and copper submerged
in zinc sulfate and copper sulfate, respectively, is known as a Daniell cell.
Half reactions for a Daniell cell are these:
7
Zinc electrode (anode): Zn(s)→ Zn2+
(aq) + 2 e-
Copper electrode (cathode): Cu 2+
(aq)+ 2 e-→ Cu(s)
A modern cell stand for electrochemical research. The electrodes
attach to high-quality metallic wires, and the stand is attached to a
potentiostat/galvanostat (not pictured). A shot glass-shaped container is
aerated with a noble gas and sealed with the Teflon block.
In this example, the anode is zinc metal which oxidizes (loses
electrons) to form zinc ions in solution, and copper ions accept electrons
from the copper metal electrode and the ions deposit at the copper cathode
as an electrodeposit. This cell forms a simple battery as it will
spontaneously generate a flow of electrical current from the anode to the
cathode through the external connection. This reaction can be driven in
reverse by applying a voltage, resulting in the deposition of zinc metal at
the anode and formation of copper ions at the cathode.
To provide a complete electric circuit, there must also be an ionic
conduction path between the anode and cathode electrolytes in addition to
the electron conduction path. The simplest ionic conduction path is to
provide a liquid junction. To avoid mixing between the two electrolytes,
the liquid junction can be provided through a porous plug that allows ion
flow while reducing electrolyte mixing. To further minimize mixing of the
electrolytes, a salt bridge can be used which consists of an electrolyte
saturated gel in an inverted U-tube. As the negatively charged electrons
flow in one direction around this circuit, the positively charged metal ions
flow in the opposite direction in the electrolyte.
A voltmeter is capable of measuring the change of electrical
potential between the anode and the cathode. Electrochemical cell voltage
is also referred to as electromotive force or emf.
A cell diagram can be used to trace the path of the electrons in the
electrochemical cell. For example, here is a cell diagram of a Daniell cell:
Zn(s) | Zn2+
(1M) || Cu2+
(1M) | Cu(s)
8
First, the reduced form of the metal to be oxidized at the anode (Zn) is
written. This is separated from its oxidized form by a vertical line, which
represents the limit between the phases (oxidation changes). The double
vertical lines represent the saline bridge on the cell. Finally, the oxidized
form of the metal to be reduced at the cathode, is written, separated from
its reduced form by the vertical line. The electrolyte concentration is given
as it is an important variable in determining the cell potential.
IV. EQUIPMENTS AND MATERIALS
IV.1 Redox Reaction
No. Name Picture Quantity
1. Measurement glass
1
2. Pippete
2
3. Reaction tube rack
1
4. Reaction tube
6
9
No. Name Quantity
1. Solution CuSO4 0.5M 4 ml
2. Solution ZnSO4 0.5M 4 ml
3. Solution FeCl3 4 ml
4. Metal Cu 2 piece
5. Metal Zn 2 piece
6. Metal Fe 2 piece
IV.2 Electrolysis
No. Name Picture Quantity
1. Funnel
1
2. Power supply
1
3. Carbon electrode
2
10
4. Pippete
2
5. Measuring glass
1
6. U pipe
1
No. Name Quantity
1. Solution KI 0.5 M sufficient
2. Solution NaCl 0.5 M sufficient
3. Solution of PP 2 drop
4. Solution of amylum 2 drop
IV.3 Electrochemistry
No. Name Picture Quantity
11
1. Funnel
1
2. Beaker glass
2
3. Multimeter
1
4. Pippete
2
5. Salt bridge
1
No. Name Quantity
1. Solution ZnSO4 0.5 M sufficient
2. SolutionCuSO4 0.5 M sufficient
12
3. Salt solution sufficient
4. Metal Zn 1 piece
5. Metal Cu 1 piece
V. PROCEDURE
V.1 Redox Reaction
1. Enter 2 ml CuSO4 solution into the reaction tube add Zn metal.
Let it for several minutes and note what happen? By same way
done to Fe metal.
2. Enter 2 ml ZnSO4 solution into the reaction tube, then add Cu
metal. Let it for several minutes and note what happen by same
way, done to Fe metal.
3. Repeat by some way to FeCl3 solution with Cu and Zn metal.
V.2 Electrolysis KI and NaCl solution
1. Enter KI solution 0.5 M into U tube until glue form top and U
tube.
2. Connect electrodes with source of electric current during less
more 5 minutes.
3. Note the change at anode.
4. Note the change occur at catode with use pipette drop, add 2 drop
indicator PP solution at catode space, observe which occur.
5. By use pipette drop, add any drop amylum solution into anode
space, observe which occur.
6. Done step 1 to 5 with use NaCl 0.5 M solution.
V.3 Electrochemistry Cell
1. Making a half cell . Enter 100 ml ZnSO4 0.5 M
solution into chemistry glass 250 ml. Place a stick of zinc scuff
into the glass.
13
2. Making a half cell . Enter 100 ml CuSO4 0.5 M
solution into the chemistry glass 250 ml. Plece a stick of cuprum
scuff into the glass.
3. Connect zinc scuff with negative pole of multimeter and connect
cuprum scuff with positive pole of voltmeter.
4. Connect both cell with salt bridge. Read voltmeter then note the
result.
VI. OBSERVATION DATA
VI.1 Redox Reaction
1. CuSO4+Zn there are bubble
2. CuSO4+Fe
3. ZnSO4+Cu
4. ZnSO4+Fe
5. FeCl3+Zn there are bubble and become rusty
6. FeCl3+Cu
VI.2 Electrolysis
1. KI 0.5 M
a. Change on anode (+) : yellow-brown
b. After add amylum to anode :black
c. Change on cathode (-) : bubbles
d. After add PP to cathode : purple
2. NaCl 0.5 M
a. Change on anode (+) : light yellow
b. After add amylum to anode : light yellow
c. Change on cathode (-) : bubbles
d. After add PP to cathode : purple
VI.3 Electrochemistry
toward = 0.6 volt
14
VII. DATA ANALYZE
VII.1 Redox Reaction
1. Cu+ZnSO4 solution
cathode : Eored= -0.76 volt
anode : Eooxi= -0.34 volt
Eocell= -1.1 volt
Eocell is negative, therefore reaction can not happen and reactivity
Zn>Cu
2. Cu+FeCl3 solution
cathode : Eored= -0.072 volt
anode : Eooxi= -0.34 volt
Eocell= -0.412 volt
Eocell is negative, therefore reaction can not happen and reactivity
Fe>Cu
3. Fe+CuSO4 solution
cathode : Eored= 0.34 volt
anode : Eooxi= 0.44 volt
Eocell= 0.78 volt
Eocell is positive, where Cu is oxidator (doing reduction) and Fe is
redactor (doing oxidation) and the reactivity Fe>Cu so the reaction is
spontan.
4. Fe+ZnSO4 solution
cathode : Eored= -0.76 volt
anode : Eooxi= +0.44 volt
Eocell= -0.32 volt
15
Eocell is negative, therefore reaction can not happen and reactivity
Zn>Fe.
5. Zn+ CuSO4 solution
cathode : Eored= 0.34 volt
anode : Eooxi= 0.76 volt
Eocell= 1.1 volt
Eocell is positive, where Cu is oxidator (doing reduction) and Fe is
redactor (doing oxidation) and the reactivity Zn>Cu so the reaction is
spontan.
6. Zn+ FeCl3 solution
cathode : Eored= -0.072 volt
anode : Eooxi= 0.76 volt
Eocell= 0.6888 volt
Eocell is positive, where Cu is oxidator (doing reduction) and Fe is
redactor (doing oxidation) and the reactivity Zn>Fe so the reaction is
spontan.
Based on the experiment of redox reaction, the reaction can be occur only
CuSO4 solution add Zn metal and FeCl3 solution add Zn metal.
In the experiment we get the difference which is gotten by calculation.
From experiment Zn added in CuSO4 solution there is changes Zn become
brown and there are stratified or powder also there are bubbles. It is match
with theory hence occur reaction. While for Fe entered in CuSO4 there are
no changes. It is indicated that not occur reaction. Through the experiment,
metal Cu and Fe not react in the ZnSO4 caused not occurring changes. It is
same with metal Cu in the FeCl3 solution. In the ordered volta metals “ Li-
K-Ba-Ca-Na-Mg-Al-Zn-Fe-Ni-Sn-Pb-H-Cu-Hg-Ag-Pt-Au ” only can
reduct other metal which exist in the right side and can not reduct metal in
the left side. Metal which has law reduction potential will have reactivity.
According to volta cell set Zn-Fe-Cu.
The differences from calculating or theory with the experiment caused by:
16
1. Practicant do not have much time in observing reaction, hence reaction
which must occur can not be observed. Practicant can not see bubbles
or color changes.
2. Inaccurate when metal Fe, Cu, Zn started to react.
3. The instrument not clean.
4. Metal Fe, Zn, and Cu firstly can not be sandpaper it influence the
reaction.
Reaction happen spontaneously if potential cell (Eocell) which resulted
from the reaction between metal with a solution have positive value.
Potential cell is the difference between cathode potential with anode
potential. At redox reaction in anode happen oxidation reaction and at
cathode happen reduction reaction. Reactivity metals based on calculation
in CuSO4 solution Zn>Fe, in ZnSO4 solution Fe>Cu and in FeCl3 solution
Zn>Cu.
VI.2 Electrolysis
Reaction :
cathode :
anode :
17
In KI solution we get ion K+ and I
- and water molecule, possibility
reaction which occur in the left side is reduction K+ ion or reduction of
water molecule.
Electrolyte solution can conduct electric current. Conduction
electric through solution accompanied a reaction which is called
electrolysis. Electrolysis is expansion process of electrolyte in the
solution form or break by electric current direct. These reaction can take
place because the influence of electric current. So, at electrolysis occur
changes from electric energy become chemistry energy. If the electric
flow through ion compound melt, so the ion compound will be
expansion cation reduction in cathode white anion oxidation in anode.
If which flowed by electric is electrolyte solution, so reaction which
occur not cation and anion, may be water or the electrode.
Reaction at cathode:
Eo= -2.92 volt
E
o= -0.83 volt
Because the reduction potential of water is higher so water able to be
reducted hence in cathode form H2 gasses.
Reaction in anode oxidation:
Eo= -0.54 volt
E
o= -0.83 volt
Because potential I- higher than potential of water oxidation, so at
anode form iodine which in this experiment have yellow-brown, so
oxidation ion I- will be able to occur. So, at electrolysis KI solution
occur reaction with the result H2 gases, OH-, KI. Reaction
is oxidation reaction, we can conclude that in the right side
18
is anode because at anode occur oxidation reaction. While at reaction of
is reduction reaction and can be
conclude at this part look cathode.
Electrolysis KI reaction with C completely:
anode :
cathode :
When flowed by electric current KI solution occur at anode has orange
or yellow color rather brown which firstly has clear rather yellow.
Based on the experiment result which gotten at anode space which
added with amylum solution that occur changes there are change color
become brown. While at cathode space which added by PP solution the
color change from pure to purple and there exist bubbles. This show
that the solution has base properties. The changes at anode if mix with
iodine. At anode appear bubbles gass that is H2 gases.
Electrolysis NaCl 0.5 M solution
Reaction :
anode :
cathode :
Reaction which occur is below
E
o= -0.83 volt
19
Eo= -2.71 volt
At electrolysis of NaCl solution with electrode grafit appear much
bubble in the cathode (-), while in the anode (+) the solution change the
color. In the cathode there exist bubbles gases caused at this part occur
reduction H2O reaction. Like which occur at the experiment at anode
space which given amylum occur changes become much bubbles gases
of Cl2 and cathode space which added PP appear bubbles H2 gases just
little, not occur change color. From theory we know that in anode from
chlorine gases and cathode result hydrogen gases. Ion OH- which occur
react with ion Na+ hence result NaOH which can be crystaled because
the potential reduction of water is higher than Na. Therefore, water
more able to reduct hence in cathode form H2 gases. At this experiment
not show that the solution have base properties. The difference between
theory and experiment caused by less accurate doing experiment,
inaccurate when drop amylum ad PP, the instrument not clean and soon.
VI.3 Electrochemistry
If an electrolyte solution and metal which act as electrode
connected with salt bridge, so will have voltage. Voltage (different
putential) Eocell can be calculated by formula .
20
Based on the constanta, the magnitude of standart reduction
potential (Eocell)
Eo= +0.34 volt
Eo= -0.76 volt
The reaction
anode : Eooxi= 0.76 volt
cathode : Eored= 0.34 volt
Eocell= 1.1 volt
Eo Cu bigger than E
o Zn, hence Cu faced reduction and Zn
oxidation. At anode occur oxidation reaction that is and
cathode occur reduction reaction . Change number (+)
which formed in anode equivalent with negative ion which formed in
cathode. Negative ions flow from salt bridge to anode, because formed
positive ion. Because negative charge formed in cathode, so positive ion
flow from salt bridge toward electrode (cathode). Salt bridge function to
arrange the equilibrium of ions solution. Metal which faced reduction is
metal which faced oxidation is metal which has Eo cell smaller or which
more difficult to reduct.
In ordered volta properties more left, the Eo reduction value smaller
means more difficult for metal to reduct (more able to oxidated) and more
right the Eo value bigger means more able to reduct. From the
electrochemistry experiment making a half cell by
entering 100 ml ZnSO4 0.5 M into beaker glass also made a half cell
entered 100 ml CuSO4 0.5 M into beaker glass and
connected both beaker glass with salt bridge. Positive pole at voltmeter
dipped in CuSO4 solution and negative pole at ZnSO4 and Eo cell at the
experiment we obtain 0.6 volt. It is different with the theory, the caused by
inaccurate to read voltmeter, instrument not clean and soon.
21
VIII. CONCLUSION
1. Redox is the term used to label reactions in which the acceptance of an
electron (reduction) by a material is matched with the donation of an
electron (oxidation). Redox reaction happen because there is difference
of potential of each element can happen if Eocell positive (E
ocell>0).
2. Order of reactivity in the experiment Zn>Fe>Cu.
3. Electrochemistry is the study of solutions of electrolytes and of
phenomena occurring at electrodes immersed in these solutions.
4. Cell potential can be calculate with formula
Magnitude of cell tension from electrochemistry
Zn(s) | Zn2+
|| Cu2+
| Cu(s)
theory : 1.1 volt
experiment : 0.6 volt
5. Electrolysis is chemical reaction when lectricity is passed through an
liquid solution of an ion or an electrolyte.
Electrolysis KI
anode :
cathode :
Electrolysis NaCl
anode :
cathode :
22
IX. BIBLIOGRAPHY
Anshory, irfan.2000.Kimia.Jakarta:Erlangga
Brady,james.1994.Kimia Universitas.Jakarta:Erlangga
Cliff.D.A et al.Chemistry Astructural View Laboratory Manual Edisi3.
G.Wulfsberg.2000.Inorganic Chemistry.University Science
Book,Sausolito,CA.
Keenan,klemfelter/.1986.Kimia Universitas.Jakarta:Erlangga
Kuswati,tine maria.2005.Sains Kimia.Jakarta:Bumi Aksara
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2011)
Surakarta, May 17th
2011
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