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HCM INTERNATIONAL UNIVERSITY SCHOOL OF BIOTECHNOLOGY. APPLICATIONS OF ELECTROCHEMISTRY. Course: ANALYTICAL CHEMISTRY Lecturer: Dr. NGUYEN TUAN KHOI. MEMBERS OF GROUP. Ph ạm Nguyễn Huệ Nh ân Thái V ă n Ch í Nguyễn Th ị Ph ươ ng Thùy Trần Đỗ Ngọc Oanh V õ Hoàng Lâm

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HCM INTERNATIONAL UNIVERSITYSCHOOL OF BIOTECHNOLOGY

APPLICATIONS OF

ELECTROCHEMISTRY

Course: ANALYTICAL CHEMISTRY

Lecturer: Dr. NGUYEN TUAN KHOI


Members of group
MEMBERS OF GROUP

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

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


Outline
Outline

I. Introduction

II. Electrochemical cells

  • Galvanic cells

  • Electrolytic cells

    III. Electrochemical cell applications

  • Battery

  • Corrosion

  • Electrolysis

    IV. Electrochemical methods

  • Nernst equation

  • Potentiometry

  • Coulometry

  • Voltammetry


I introduction
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 introduction1
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 cells


Ii electrochemical cell
II. Electrochemical cell

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

  • An electrochemical cell includes:

    • Two electrodes: half redox reactions occur

      • Anode: oxidation reaction occur

      • Cathode: reduction reaction occur

    • Electrolyte solution(s)


Ii electrochemical cell1
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 cell2
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).


a. GALVANIC CELL

What about the sign of the electrodes?

What about half-cell reactions?

-

+

Why?

cathode half-cell

Cu+2 + 2e- Cu

anode half-cell

Zn  Zn+2 + 2e-

Cu

plates out or deposits on electrode

Zn electrode erodes

or dissolves

What happened at each electrode?

Cu

Zn

1.0 M CuSO4

1.0 M ZnSO4


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

a. Galvanic cells

HALF REACTION

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

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

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

Salt bridge

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

cathode

anode



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)


III.

Electrochemical cell

applications


Primary battery
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
A TYPICAL STRUCTURE OF A PRIMARY BATTERY

The electrode reactions are

Zn → Zn2+ + 2e–

2 MnO2 + 2H+ + 2e– → Mn2O3 + H2O


Secondary batteries
SECONDARY BATTERIES

A secondaryorstorage 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
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.



Electrochemical corrosion
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
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
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

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


Electrolysis of water
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



The chloralkali industry
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)



Electrolytic refining of aluminum
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 is

    2 Al2O3 + 3 C → 4 Al + 3 CO2


IV.

Electroanalytical

methods


Nernst equation
NERNST EQUATION

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

aOx + ne- ↔ bRed

E = E0 – (2.3026RT/nF)log ([Red]b/[Ox]a)

E: the reduction potential at the specific concentration

n: the number of electrons

R: the gas constant (8.3143 V coul deg-1 mol -1)

T: the absolute temperature

F: the Faraday constant (96487 colul eq-1)

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


Electrochemical analysis
ELECTROCHEMICAL ANALYSIS

  • Potentiometry

  • Coulometry

  • Voltammetry


A potentiometry
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.



Mechanism
Mechanism

Ecell= Eind - Eref

Difference in potential

Reference electrode

Indicator electrode

(Constant potential)

(Change in potential)

Mobilities of ions

Solution


Electrode
Electrode

Metal

Indicator Electrode

Membrane

Potentiometry

Electrode

Reference Electrode


Reference electrodes

Calomel Reference Electrodes

Silver/ Silver Chloride Reference Electrodes


Indicator electrodes

Metallic

Membrane


Application
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.


b. COULOMETRY

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


Application of coulometry
Application of coulometry


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

COULOMETER:


C voltammetry
c. Voltammetry electricity required to bring about a chemical change of the

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


C voltammetry1
c. Voltammetry electricity required to bring about a chemical change of the

  • Include 3 electrodes

    • Working electrode: which the analyte is oxidizes or reduce

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

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


Application of voltammetry in diabetes diagnostic
Application of voltammetry in electricity required to bring about a chemical change of the 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


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