31399_analisis Elektrokimia 02.pdf

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 masingmasing 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)

• 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

 Electricity

travels in a complete circuit

H+ MnO4-

e-

Fe+2

Instead

of a salt bridge

H+ MnO4-

Porous Disk

Fe+2

Electrodes • Anode: Oxidation occurs at the anode • Cathode: Reduction occurs at the cathode

e-

ee-

e-

Anode eReducing Agent

Cathode eOxidizing Agent

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. -



Can monitor redox reaction when electrons flow through an electric current -



Chemicals are separated

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” 

3.)

Include electrons in their equation

The net reaction is known as the total cell reaction 

No free electrons in its equation

½ cell reactions: Net Reaction: 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

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 Relation between charge and moles:

q  nF Coulombs

moles

Coulombs mol e 

6.) Electric current  Quantity of charge flowing each second through a circuit Ampere: unit of current (C/sec)

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 -

Relation between free energy, work and voltage:

Measure of force pushing on electrons

G  work  E  q Joules

Volts

Higher potential difference requires more work to lift water (electrons) to higher trough

Coulombs

Higher potential difference

Fundamentals of Electrochemistry Basic Concepts 7.) Electric Potential (E) 

Combining definition of electrical charge and potential

G  work  E  q Relation between free energy difference and electric potential difference:

q  nF

G  nFE

Describes the voltage that can be generated by a chemical reaction

Fundamentals of Electrochemistry Basic Concepts 8.) Ohm’s 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

E I R 9.) Power (P)  Work done per unit time -

Units: joules per second J/sec or watts (W)

q P  E  EI t

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

Reduction: Oxidation: Net Reaction:

AgCl(s) is reduced to Ag(s) 2+ Electrons Cd(s) is oxidized travel from to Cd Cd Ag deposited on electrode and Cl2+ Cdgoesgoes solution electrode tointo Ag electrode into solution

Galvanic Cells 1.) Galvanic or Voltaic cell 

Example: Calculate the voltage for the following chemical reaction

G = -150kJ/mol of Cd

n – number of moles of electrons

Solution:

G   nFE  E  

G nF

 150  10 3 J E  0.777 J  0.777 V C C   ( 2 mol ) 9.649  10 4  mol  

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

Anode: electrode where oxidation takes place

Cathode: electrode where reduction takes place

Galvanic Cells 4.) Salt Bridge  

Connects & separates two half-cell reactions Prevents charge build-up and allows counter-ion migration

Salt Bridge  Contains electrolytes not involved in redox reaction.

TwoCd half-cell reactions 2+) moves  K+ (and to cathode with e through salt bridge (counter balances –charge build-up  NO3- moves to anode (counter balances +charge build-up)  Completes circuit

Galvanic Cells 5.) Short-Hand Notation 

Representation of Cells: by convention start with anode on left

Phase boundary Electrode/solution interface anode

Zn|ZnSO4(aZN2+ = 0.0100)||CuSO4(aCu2+ = 0.0100)|Cu

Solution in contact with anode & its concentration

2 liquid junctions due to salt bridge

cathode

Solution in contact with cathode & its concentration

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 +2eEº= 0.44 V • O2 + 2H2O + 4e- 4OH- Eº= 0.40 V • Fe+2 + O2 + H2O Fe2O3 + H+ • Reactions happens in two places.

Salt speeds up process by increasing conductivity Water Fe2+

Iron DissolvesFe  Fe+2

Rust

eO2 + 2H2O +4e-  4OH-

Fe2+ + O2 + 2H2O  Fe2O3 + 8 H+

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 Faraday’s 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)

Electrolysis • Running a galvanic cell backwards. • Put a voltage bigger than the potential and reverse the direction of the redox reaction. • Used for electroplating.

1.10 e-

e-

Zn 1.0 M Zn+2 Anode

1.0 M Cu+2 Cathode

Cu

e-

e-

A battery >1.10V

Zn 1.0 M Zn+2 Cathode

1.0 M Cu+2 Anode

Cu

Calculating plating • • • • • • •

Have to count charge. Measure current I (in amperes) 1 amp = 1 coulomb of charge per second q=Ixt 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 e3. 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.