HCM INTERNATIONAL UNIVERSITY

SCHOOL OF BIOTECHNOLOGY

Course: ANALYTICAL CHEMISTRYLecturer: Dr. NGUYEN TUAN KHOI

MEMBERS OF GROUP

T ôn Th ị H ồng Th ảo

Nguy ễn Vi ệt Th ư

Nguy ễn Ng ọc Y ên Nhi

Đ oàn T ây Nguy ên

Nguy ễn Duy Trung

Nguy ễn Đ ức Thanh Long

V ũ Ng ọc C ư ơng

Nguyễn Vũ Nh ất Th ịnh

Ph ạm Nguyễn Huệ Nh ân

Thái Văn Ch í

Nguyễn Th ị Phương Thùy

Trần Đỗ Ngọc Oanh

V õ Hoàng Lâm

Nguyễn Th ị Thu Cúc

Trương Thị Ngọc Nhi

Lê Trần Khánh Trang

Đỗ Vân Khanh

Outline I. IntroductionII. Electrochemical cells

a. Galvanic cellsb. Electrolytic cells

III. Electrochemical cell applicationsa. Batteryb. Corrosionc. Electrolysis

IV. Electrochemical methods

a. Nernst equation

b. Potentiometry

c. Coulometry

d. Voltammetry

I. INTRODUCTION• Electrochemistry is the study of

reactions in which charged particles (ions or electrons) cross the interface between two phases of matter, typically a metallic phase (the electrode) and a conductive solution, or electrolyte. This reaction is simple oxidation-reduction process.

I. INTRODUCTION

Redox reaction(reduction-oxidation reactions)

• Are reactions that mention to the transfer of electrons between species.

• Describe all chemical reactions in which atoms have oxidation number change.Ox1 + red2 red1 +

ox2

II. Electrochemical cell

• Transform energy from chemical reaction to electrical energy or vice versa.

• An electrochemical cell includes: – Two electrodes: half redox reactions occur

oAnode: oxidation reaction occuroCathode: reduction reaction occur

– Electrolyte solution(s)

II. Electrochemical cell

• Conditions for generating electricity flow:

�̶ The electrodes must be externally connected by a metal wire to permit electron flow.

�̶ The electrolyte solutions are in contact to allow movement of ions.

II. Electrochemical cell

• There are two types of electrochemical cells:

– Galvanic cells (or Voltaic cells): spontaneous reactions occur.

– Electrolytic cells: nonspontaneous reactions occur (electrical energy supply).

Cu

1.0 M CuSO4

Zn

1.0 M ZnSO4

Cu plates out or deposits on electrode

Zn electrode erodes or dissolves

cathode half-cellCu+2 + 2e- Cu

anode half-cellZn Zn+2 + 2e-

-+What about half-cell reactions?

What about the sign of the electrodes?

What happened at each electrode?

Why?

a. GALVANIC CELL

a. Galvanic cells

Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)

anode cathode

HALF REACTION

Salt bridge

Anode (Ox) : Zn(s) = Zn2+ + 2e

Cathode (Red) : Cu2+ + 2e = Cu (s)

Net reaction : Zn (s) + Cu2+ = Zn2+ +Cu (s)

b. ELECTROLYTIC CELL

• cell potential: cell potential: Electrons flow from one electrode to the other in one direction, there is a potential difference between the electrodes..

• Cell potential is calculated in voltage (V) by the formula:

E cathode: reduction potential (V) E anode: oxidation potential (V)

CELL POTENTIAL

Cell potential (E cell) = E cathode – E anode

PRIMARY BATTERY

Primary battery has long been known as dry-cell. It cannot be recharged. It’s widely used to power flashlight and some other similar devices.

The first practical battery consisted of a stack of small electrical cell, each consisting of a silver plate and a zinc plate separated by a sheet of cardboard which had been soaked in salt water

A TYPICAL STRUCTURE OF A PRIMARY

BATTERY

The electrode reactions are Zn → Zn2+ + 2e–

2 MnO2 + 2H+ + 2e– → Mn2O3 + H2O

SECONDARY BATTERIES

A secondary or storage battery is capable of being recharged. Its electrode reactions can proceed in either direction.

During charging, electrical work is done on the cell to provide the free energy needed to force the reaction in the non-spontaneous direction.

Fuel cell

Conventional batteries supply electrical energy from the chemical reactants stored within them. When these reactants are consumed, the battery is "dead". An alternative approach would be to feed the reactants into the cell as they are required, so as to permit the cell to operate continuously. In this case the reactants can be thought of as "fuel" to drive the cell, hence the term fuel cell.

MODERN HYDROGEN-OXYGEN FUEL CELL

anode:  H2(g) → 2 H+ + 2e– E° = 0 v

cathode:   ½ O2 + 2 H+ + 2e– → H2O(l) E° = +1.23 v

net:   H2(g) + ½ O2(g) → H2O(l) E° = +1.23 v

ELECTROCHEMICAL CORROSION

Corrosion is the deterioration of materials by chemical processes. Of these, the most important by far is electrochemical corrosion of metals, in which the oxidation process M → M+ + e– is facilitated by the presence of a suitable electron acceptor

Sacrificial coating

One way of supplying this negative charge is to apply a coating of a more active metal

a very common way of protecting steel from corrosion is to coat it with a thin layer of zinc

Cathodic protection

A more sophisticated strategy is to maintain a continual negative electrical charge on a metal, so that its dissolution as positive ions is inhibited. The entire surface is forced into the cathodic condition.

ELECTROLYSIS

Electrolysis refers to the decomposition of a substance by an electric current

ELECTROLYSIS OF WATER

• Water is capable of undergoing both oxidation and reduction• Pure water is an insulator and cannot undergo signifigant

electrolysis without adding an electrolyte.• Electrolysis of a solution of sulfuric acid or of a salt such as

NaNO3 results in the decomposition of water at both electrodes:

• cathode:  H2O + 2 e– → H2(g) + 2 OH– E =+0.41 v ([OH–] = 10-7 M)

• anode:   2 H2O → O2(g) + 4 H+ + 2 e– E° = -0.82 v

• net:  2 H2O(l) → 2 H2(g) + O2(g) E = -1.23 v

ELECTROLYSIS

THE CHLORALKALI INDUSTRY

• The electrolysis of brine is carried out on a huge scale for the industrial production of chlorine and caustic soda (sodium hydroxide). Because the reduction potential of Na+ is much higher than that of water, the latter substance undergoes decomposition at the cathode, yielding hydrogen gas and OH–.

• 2 NaCl + 2 H2O → 2 NaOH + Cl2(g) + H2(g)

Schematic diagram of a cell for the production of chlorine

A modern industrial chloralkali plant

ELECTROLYTIC REFINING OF ALUMINUM

• The Hall-Hérault process takes advantage of the principle that the melting point of a substance is reduced by admixture with another substance with which it forms a homogeneous phase.

• The net reaction is2 Al2O3 + 3 C → 4 Al + 3

CO2

NERNST EQUATION

Nernst equation allows one unknown concentration to be determined from a measurement of the cell voltage.

aOx + ne- ↔ bRedE = E0 – (2.3026RT/nF)log

([Red]b/[Ox]a)

E: the reduction potential at the specific concentrationn: the number of electronsR: the gas constant (8.3143 V coul deg-1 mol -1)T: the absolute temperatureF: the Faraday constant (96487 colul eq-1)

At 25oC, the value of 2.3026RT/F is 0.05916

ELECTROCHEMICAL ANALYSIS

a. Potentiometry

b.Coulometry

c. Voltammetry

a. POTENTIOMETRY

• Potentiometry passively measures the potential of a solution

between two electrodes, affecting the solution very little in the

process. The potential is then related to the concentration of

one or more analytes.

• In potentiometry, there are no current, or only negligible

current flows, so the compound in the solution remain

unchanged. It is used for measure the cell potential and for

determine the analytical quantity of interest. Potentiometry is

a useful quantitative method.

Potentiometric titration

Mechanism

Reference electrode

Indicator electrode

Solution

Mobilities of ions

(Constant potential) (Change in potential)

Difference in potential

Ecell= Eind - Eref

Electrode

PotentiometryElectrode

Indicator Electrode

Reference Electrode

Metal

Membrane

Reference Reference electrodeselectrodesReference Reference electrodeselectrodes

Calomel Reference Electrodes

Calomel Reference Electrodes Silver/ Silver

Chloride Reference Electrodes

Silver/ Silver Chloride Reference

Electrodes

MembraneMembraneMetallicMetallic

Indicator electrodesIndicator electrodesIndicator electrodesIndicator electrodes

Application

Used in pH meter, by using glass electrode

In environment, used to analyse ion -CN-, F-, NH3, and NO3- in water and in wastewater.

• Coulometry: electrochemical method based on the quantitative oxidation or reduction of analyte

- Measure amount of analyte by measuring amount of current and time required to complete reaction

• charge = current (i) x time in coulombs

Q = ite

b. COULOMETRY

APPLICATION OF COULOMETRY

COULOMETER:

• A coulometer is a device used for measuring the quantity of electricity required to bring about a chemical change of the analyte.

• It is usual practice in coulometry to substitute the ammeter

c. Voltammetry

• Measures current as a function of applied potential under conditions that keep a working electrode polarized

• Include 3 electrodes1. Working electrode:

which the analyte is oxidizes or reduce

2. Counter electrode: which is often a coil of platinum wire or a pool of mercury.

3. Reference electrode: potential remains constant (Ag/AgCl electrode or calomel)

c. Voltammetry

Application of voltammetry in

Diabetes diagnostic

• Diabetes is a serious disease and is the fourth leading cause of death by disease in US. Its causes are unknown, and there is no cure.

Testing blood: A Crucial

Tool