Analisis

Elektrokimia

Oleh

Moh. Hayat

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)

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

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.

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.

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)

H+ MnO4

- Fe+2

Connected this way the reaction starts

Stops immediately because charge builds up.

H+ MnO4

- Fe+2

Salt Bridge allows current to flow

H+ MnO4

- Fe+2 e-

Electricity travels in a complete circuit

H+ MnO4

- Fe+2

Porous Disk

Instead of a salt bridge

Electrodes

Anode:

Oxidation occurs at the anode

Cathode:

Reduction occurs at the cathode

Reducing Agent

Oxidizing Agent

e-

e-

e- e-

e-

e-

Anode Cathode

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

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

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

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:

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

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

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

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

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

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

Galvanic Cells

2.) Cell Potentials vs. G Reaction is spontaneous if it does not require external energy

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

Galvanic Cells

3.) Electrodes

Cathode: electrode where reduction takes place

Anode: electrode where oxidation takes place

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

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

Batteries are Galvanic Cells Car batteries are lead storage batteries.

Pb +PbO2 +H2SO4 PbSO4(s) +H2O

Batteries are Galvanic Cells Dry Cell

Zn + NH4+ +MnO2

Zn+2 + NH3 + H2O + Mn2O3

Batteries are Galvanic Cells Alkaline

Zn +MnO2 ZnO+ Mn2O3 (in base)

Batteries are Galvanic Cells NiCad

NiO2 + Cd + 2H2O Cd(OH)2 +Ni(OH)2

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.

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+

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.

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)

Running a galvanic cell backwards.

Put a voltage bigger than the potential and reverse the direction of the redox reaction.

Used for electroplating.

Electrolysis

1.0 M Zn+2

e- e-

Anode Cathode

1.10

Zn Cu 1.0 M Cu+2

1.0 M Zn+2

e- e-

Anode Cathode

A battery >1.10V

Zn Cu 1.0 M Cu+2

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+

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

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?

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.