analisis elektrokimia 02
DESCRIPTION
analisis elektrokimiaTRANSCRIPT
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Analisis
Elektrokimia
Oleh
Moh. Hayat
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Elektrokimia
Merupakan cabang Ilmu Kimia yang membahas mengenai sifat kelistrikan dan pengaruhnya terhadap zat-zat kimia.
Sebagian besar mempelajari perubahan kimiawi yang disebabkan oleh adanya arus listrik dan timbulnya kelistrikan karena adanya reaksi kimia.
Kajian elektrokimia meliputi banyak fenomena (mis: elektroporesis, korosi); peralatan (display elektrokromik, sensor elektroanalitik, baterai, fuel cell); dan teknologi (elektroplating, produksi aluminium)
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Elektroanalitik
Kumpulan metode analitik kuantitatif yang didasarkan pada sifat-sifat listrik larutan analit ketika ditempatkan dalam sel elektrokimia.
Teknik elektroanalitik mampu memberikan pengukuran dengan limit deteksi yang rendah dan mampu memberikan banyak informasi, stoikhiometri dan kecepatan transfer muatan, kecepatan transfer massa, kecepatan dan konstanta kesetimbangan reaksi kimia.
Pemanfaatan elektrokimia untuk analisis
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Pengantar Elektrokimia
Elektrokimia didasarkan pada reaksi redoks.
Sifat khas dari reaksi redoks adalah adanya transfer elektron dan dapat ditempatkan dalam suatu sel elektrokimia.
Dalam sel tersebut, oksidator dan reduktor ditempatkan dalam sel terpisah.
Kedua sel tersebut dihubungkan dengan jembatan garam yang mengisolasi pereaksi namun memungkinkan transfer muatan listrik.
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Sel Elektrokimia
Suatu sel elektrokimia terdiri dari dua konduktor yang disebut elektroda, yang masing-masing direndam dalam larutan elektrolit.
Dalam kebanyakan kasus, larutan yang digunakan untuk masing-masing elektroda merupakan elektrolit yang berbeda, dan harus dipisahkan untuk menghindari reaksi.
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Common Components
Electrodes: conduct electricity between cell and
surroundings
Electrolyte: mixture of ions involved in reaction or
carrying charge
Salt bridge: completes circuit (provides charge balance)
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H+ MnO4
- Fe+2
Connected this way the reaction starts
Stops immediately because charge builds up.
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H+ MnO4
- Fe+2
Salt Bridge allows current to flow
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H+ MnO4
- Fe+2 e-
Electricity travels in a complete circuit
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H+ MnO4
- Fe+2
Porous Disk
Instead of a salt bridge
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Electrodes
Anode:
Oxidation occurs at the anode
Cathode:
Reduction occurs at the cathode
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Reducing Agent
Oxidizing Agent
e-
e-
e- e-
e-
e-
Anode Cathode
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Types of cells
Voltaic (galvanic) cells:
Energy released from spontaneous redox reaction can be transformed into electrical energy.
Electrolytic cells:
Electrical energy is used to drive a nonspontaneous redox reaction source to drive a nonspontaneous reaction
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Fundamentals of Electrochemistry
1.) Electrical Measurements of Chemical Processes Redox Reaction involves transfer of electrons from one species to another.
- Chemicals are separated
Can monitor redox reaction when electrons flow through an electric current
- Electric current is proportional to rate of reaction - Cell voltage is proportional to free-energy change
Batteries produce a direct current by converting chemical energy to electrical energy. - Common applications run the gamut from cars to ipods to laptops
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Fundamentals of Electrochemistry
Basic Concepts
1.) A Redox titration is an analytical technique based on the transfer of electrons between analyte and titrant Reduction-oxidation reaction
A substance is reduced when it gains electrons from another substance
- gain of e- net decrease in charge of species - Oxidizing agent (oxidant)
A substance is oxidized when it loses electrons to another substance - loss of e- net increase in charge of species - Reducing agent (reductant)
(Reduction)
(Oxidation)
Oxidizing Agent
Reducing Agent
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Fundamentals of Electrochemistry
Basic Concepts
2.) The first two reactions are known as 1/2 cell reactions Include electrons in their equation
3.) The net reaction is known as the total cell reaction No free electrons in its equation
4.) In order for a redox reaction to occur, both reduction of one compound and oxidation of another must take place simultaneously Total number of electrons is constant
cell reactions:
Net Reaction:
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Fundamentals of Electrochemistry
Basic Concepts
5.) Electric Charge (q) Measured in coulombs (C) Charge of a single electron is 1.602x10-19C Faraday constant (F) 9.649x104C is the charge of a mole of
electrons
6.) Electric current Quantity of charge flowing each second through a circuit
- Ampere: unit of current (C/sec)
Fnq Relation between charge and moles:
Coulombs moles emol
Coulombs
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Fundamentals of Electrochemistry
Basic Concepts
7.) Electric Potential (E) Measured in volts (V) Work (energy) needed when moving an electric
charge from one point to another - Measure of force pushing on electrons
qEworkG Relation between free energy, work and voltage:
Joules Volts Coulombs
Higher potential difference
Higher potential difference requires more work to lift water (electrons) to higher trough
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Fundamentals of Electrochemistry
Basic Concepts
7.) Electric Potential (E) Combining definition of electrical charge and potential
qEworkG Fnq
nFEG Relation between free energy difference and electric potential difference:
Describes the voltage that can be generated by a chemical reaction
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Fundamentals of Electrochemistry
Basic Concepts
8.) Ohms Law Current (I) is directly proportional to the potential difference (voltage)
across a circuit and inversely proportional to the resistance (R) - Ohms (W) - units of resistance
9.) Power (P) Work done per unit time
- Units: joules per second J/sec or watts (W)
R
EI
IEt
qEP
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Galvanic Cells
1.) Galvanic or Voltaic cell Spontaneous chemical reaction to generate electricity
- One reagent oxidized the other reduced - two reagents cannot be in contact
Electrons flow from reducing agent to oxidizing agent - Flow through external circuit to go from one reagent to the other
Net Reaction:
Reduction:
Oxidation:
AgCl(s) is reduced to Ag(s) Ag deposited on electrode and Cl-
goes into solution
Electrons travel from Cd electrode to Ag electrode Cd(s) is oxidized to Cd2+
Cd2+ goes into solution
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Galvanic Cells
1.) Galvanic or Voltaic cell Example: Calculate the voltage for the following chemical reaction
G = -150kJ/mol of Cd
V.C
J.
mol
C.)mol(
JE
nF
GEnFEG
77707770
1064992
10150
4
3
Solution: n number of moles of electrons
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Galvanic Cells
2.) Cell Potentials vs. G Reaction is spontaneous if it does not require external energy
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Galvanic Cells
2.) Cell Potentials vs. G Reaction is spontaneous if it does not require external energy
Potential of overall cell = measure of the tendency of this reaction to proceed to equilibrium
Larger the potential, the further the reaction is from equilibrium and the greater the driving force that exists
Similar in concept to balls sitting at different heights along a hill
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Galvanic Cells
3.) Electrodes
Cathode: electrode where reduction takes place
Anode: electrode where oxidation takes place
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Galvanic Cells
4.) Salt Bridge Connects & separates two half-cell reactions Prevents charge build-up and allows counter-ion migration
Two half-cell reactions
Salt Bridge
Contains electrolytes not involved in redox reaction. K+ (and Cd2+) moves to cathode with e- through salt bridge (counter balances charge build-up NO3
- moves to anode (counter balances +charge build-up) Completes circuit
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Zn|ZnSO4(aZN2+ = 0.0100)||CuSO4(aCu2+ = 0.0100)|Cu anode
Phase boundary Electrode/solution interface
Solution in contact with anode & its concentration
Solution in contact with cathode & its concentration
2 liquid junctions due to salt bridge
cathode
Galvanic Cells
5.) Short-Hand Notation Representation of Cells: by convention start with anode on left
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Batteries are Galvanic Cells Car batteries are lead storage batteries.
Pb +PbO2 +H2SO4 PbSO4(s) +H2O
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Batteries are Galvanic Cells Dry Cell
Zn + NH4+ +MnO2
Zn+2 + NH3 + H2O + Mn2O3
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Batteries are Galvanic Cells Alkaline
Zn +MnO2 ZnO+ Mn2O3 (in base)
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Batteries are Galvanic Cells NiCad
NiO2 + Cd + 2H2O Cd(OH)2 +Ni(OH)2
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Corrosion
Rusting - spontaneous oxidation.
Most structural metals have reduction potentials that are less positive than O2 .
Fe Fe+2 +2e- E= 0.44 V
O2 + 2H2O + 4e- 4OH- E= 0.40 V
Fe+2 + O2 + H2O Fe2O3 + H+
Reactions happens in two places.
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Water
Rust
Iron Dissolves- Fe Fe+2
e-
Salt speeds up process by increasing conductivity
O2 + 2H2O +4e- 4OH-
Fe2+ + O2 + 2H2O Fe2O3 + 8 H+
Fe2+
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Preventing Corrosion
Coating to keep out air and water.
Galvanizing - Putting on a zinc coat
Has a lower reduction potential, so it is more easily oxidized.
Alloying with metals that form oxide coats.
Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.
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Electrolysis
Use Faradays Laws to evaluate the number of moles of a substance oxidised or reduced by passage of charge (current over a given period of time = I.t) through an electrode
Faraday: Q (charge) = nF N=number of moles of electrons
F=constant of 96500 Coulomb/mole
Example (try it): What current is needed to deposit 0.500g of chromium from a solution containing Cr3+ over a one hour period (MW for Cr=52)? (Ans=0.77A)
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Running a galvanic cell backwards.
Put a voltage bigger than the potential and reverse the direction of the redox reaction.
Used for electroplating.
Electrolysis
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1.0 M Zn+2
e- e-
Anode Cathode
1.10
Zn Cu 1.0 M Cu+2
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1.0 M Zn+2
e- e-
Anode Cathode
A battery >1.10V
Zn Cu 1.0 M Cu+2
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Calculating plating
Have to count charge.
Measure current I (in amperes)
1 amp = 1 coulomb of charge per second
q = I x t
q/nF = moles of metal
Mass of plated metal
How long must 5.00 amp current be applied to
produce 15.5 g of Ag from Ag+
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Calculating plating
1. Current x time = charge
2. Charge Faraday = mole of e-
3. Mol of e- to mole of element or compound
4. Mole to grams of compound
Or the reverse if you want time to plate
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Calculate the mass of copper which can be deposited by the passage of 12.0 A for 25.0 min through a solution of copper(II) sulfate.
How long would it take to plate 5.00 g Fe from an aqueous solution of Fe(NO3)3 at a current of 2.00 A?
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Other uses
Electrolysis of water.
Separating mixtures of ions.
More positive reduction potential means the reaction proceeds forward.
We want the reverse.
Most negative reduction potential is easiest to plate out of solution.