Applied electrochemistry

Applied electrochemistry

Academic year 2020-2021

By Armanini Linda

Sommario

• Introduction ............................................................................................................................. 3
• Basic concepts ......................................................................................................................... 4
• The electrodes ........................................................................................................................ 5
• The electrolyte ....................................................................................................................... 7
• Overpotentials ..................................................................................................................... 12
• Foundamental laws ................................................................................................................ 12
• Electrodeposition, electrowinning and electrorefining ..................................................... 18
• Electrodeposition ............................................................................................................. 18
• Electrowinning ................................................................................................................. 23
• Electrorefining ................................................................................................................. 25
• Electrophoresis ....................................................................................................................... 25
• Gel electrophoresis .............................................................................................................. 26
• Isoelectric focusing ............................................................................................................. 27
• Electrophoretic deposition.................................................................................................. 27
• Electroless plating.................................................................................................................. 29
• Galvanic displacement (or deposition) ................................................................................ 29
• Autocatalytic deposition ..................................................................................................... 30
• Barrel plating....................................................................................................................... 31
• Ni, NiP and NiP-composite ...................................................................................................... 32
• Energy storage systems .......................................................................................................... 33
• Batteries ............................................................................................................................... 34
• Zinc carbon battery (non rechargeable) .................................................................................. 38
• Alkaline battery (non rechargeable) ........................................................................................ 38
• Zinc-based button battery (non rechargeable) ........................................................................ 38
• Li batteries (non rechargeable) ................................................................................................. 39
• Lead acid battery (rechargeable) ............................................................................................. 39
• NiCd battery........................................................................................................................... 40
• NiMH (nickel metal hydrides) battery ................................................................................... 40
• Lithium-ion battery............................................................................................................... 40
• State of art for Li-ion batteries ........................................................................................... 43
• Redox flow batteries (RFB).................................................................................................... 46
• Manufacturing..................................................................................................................... 51
• Electrochemical capacitors................................................................................................. 51
• Fuel cells .................................................................................................................................. 53
• Hydrogen production .......................................................................................................... 55
• Electrolytic process ............................................................................................................. 55
• Photoelectrochemical water splitting ............................................................................... 57
• Electrochemical sensors ........................................................................................................ 60
• Chlor-alkali plant ................................................................................................................ 61
• Electrocoagulation / electrooxidation.............................................................................. 63
• Synthesis of semiconductors.................................................................................................. 64

Introduction

The first evidence of electrochemistry can be dated back to 1791, with the famous experiment performed by Galvani on frogs: he observed that it was possible to have a flow of charges inside the body of frogs, inducing muscular contraction. Eight years later, in 1799, Volta built the first electrochemical device which was made of zinc and copper electrodes and a saline solution working as electrolyte (the Volta pile). The following year the first example of electroplating process was conducted by Luigi Prugnatelli, who plated gold onto silver exploiting Volta’s pile. The mathematical basis for some electrochemical phenomena was derived by Faraday in 1833, which can be considered the actual birth of electrochemistry.

Electrochemistry is that part of physical chemistry that deals with transformation of matter (chemistry) through the shift and exchange of electronic charges (electro). The foundamental difference from standard REDOX reactions is that this exchange of charges takes place at the interface between an ionic and an electronic conductor, as in the typical example of the reaction between a metal (electronic conductor) and a saline solution (ionic conductor). A big difference between these two types of conductors lays in their behaviour at high temperature: for electronic conductors the resistivity increases with the temperature, while for ionic conductors the mobility of the ions increases with the temperature, as well as the conductivity:

↓ Electronic conductors ↑ ↓ ↑ Ionic conductors

The ionic conductor is the so-called electrolyte, which can be both solid and liquid. The former case is less common because for solid materials very high temperatures are required to promote the flow of ions. It can also be both organic and inorganic.

Talking about the electrodes, they must be made of a conductive material, so they can be made of both conductive and semi-conductive material, and typically they are metallic. For sure, the circuit works only if it closed, so the configuration represented in the picture on the left is necessary: two interfaces, the electrodes, are placed in a conductive solution, the electrolyte, and they are externally connected by a wire closing the circuit. One electrode works as cathode and it is subjected to a reduction reaction, while the other one is the anode, which undergoes oxidation. Considering the following scheme, the forward reaction is oxidation, while the backward one is reduction: − ↔ + 3

The electrochemical device is able to work in two configurations: it can both adsorb and generate current. If current is forced to circulate, the device is called electrolytic cell and an external power supply is needed for the reaction to occur. On the other hand, a power generator is characterized by spontaneous reactions leading to the circulation of current in the external wire, which can be collected and used in other devices. From the thermodynamics point of view, all of this can be related to Gibbs free energy thanks to Faraday’s ∆ < 0 ∆ > 0 law: means spontaneous reaction, while means non-spontaneous reaction. The device is not at an equilibrium state, as current is forced to pass or is alternatively spontaneously generated, so also kinetics must be considered. The kinetic contributions are always against the system, so they hinder both the charging and discharging processes.

Oxidation and reduction reactions lead to a change in the valence of the metallic part: in the example provided, iron moves from a valence 2 to a valence 3 through oxidation, so it loses one electron. The opposite is valid for reduction reactions. + −2 + ↔ + 2 + 2

Another example frequent in corrosion is the oxidation of metallic iron to form rust (anode), while water and oxygen reduce giving OH: + −3 + 4 = + 8 + 8

The partial reactions are the following:

• Anode: 2+ − − − → + 2
• Cathode: + 2 + 4 → 4

Other processes that can be described with electrochemical reactions are lead corrosion in lead acid batteries (oxidation) and copper electrodeposition (reduction):

• − + + −() () () + + = + 2 + 2
• 42+ 42− − 42−( ) () () + + 2 = + ()

In the second case, copper does not interact with the anode as it precipitates in the metallic form on the substrate, so other reactions take place at the anode.

Basic concepts

The typical set-up of a cell is constituted by two electrodes immersed in an electrolyte and externally linked by a wire closing the circuit in which current circulates. For sure, electrodes must be conductive, so conductors or semiconductors must be exploited. The system can be exposed to air or sealed, which is the case of organic electrolytes or electrodes that can undergo unwanted oxidation due to atmospheric oxygen or moisture. The external wire can be connected to a power supply if the cell is forced to work as an electrolytic cell, or to a load if it works as a Galvanic cell (spontaneous evolution). In any case, the external circuit should be optimized in order to minimize the electric resistances induced by the wire and by the joints, so energy loss. This external contribute can be neglected only at the laboratory scale setup, while it must be considered at the industrial level.

The electrodes

The processes taking place at the electrode, so oxidation or reduction reactions, can be schematized in three classes (they are all referred to the cathode):

• Electrons accumulate at the interface between the electrode and the electrolyte and they induce the reduction of the ions in solution. So, in this case the electrolyte contains ions both before and after the reaction.
• The ions in the electrolyte are reduced to a zero-valent state: metallic ions are deposited and the interface changes with time, as it is progressively covered by copper.
• All the electrons are used to reduce hydrogen which is adsorbed in the atomic form and then released in form of gaseous hydrogen. This phenomenon is called hydrogen evolution and it can be very dangerous because the adsorbed hydrogen can easily migrate in the metal causing hydrogen embrittlement. If this phenomenon becomes very important, the gaseous layer may occupy the entire interface and prevent the passage of current. Hydrogen evolution can be limited, but water-based electrolytes are typically exploited, so it is almost impossible to completely avoid it. This reaction is in general considered a side reaction, unless when the goal is to generate hydrogen. The same discussion can be done for the anode, but in this case oxygen evolution takes place due to oxidation of ⁻.

In general, at the electrode more than one reaction takes place at a time. A typical example is copper electrodeposition, for which hydrogen evolution occurs as a side reaction. The competition between the two reactions leads to a decrease in the cathodic efficiency as a part of the current is wasted for the side reaction. An efficiency above 60% is considered high, below 50% is low. For chromium plating, we have = 10%, which is very inefficient but there are no alternatives.

At the interface electrode/electrolyte, electrons are present due to the polarization of the electrode, and the ions in the electrolyte are electrostatically attracted. In a certain range of polarization, there may be no exchange of charges at the interface (non-reactive interface), so they accumulate and form an electrochemical double layer behaving as a sort of capacitor (electrolytic capacitor). The dielectric constant can be defined based on the environment generated at the interface and the possibility of storing energy generated at the interface (typical of supercapacitors) can also enhance the charge/discharge rate. The capacity can be improved by increasing the area of the interface.

It is possible to describe this behaviour through three models:

• Helmholtz model is the easiest one and it considers only the presence of the ionic double layer.
• Gouy-Chapman model introduced the concept of diffusion layer, that is a gradient in the concentration of ions in the electrolyte.
• Gouy-Chapman-Stern model is a combination of the previous ones, considering also the contribute of the solvent molecules.

So, a reactive interface acts as a sort of resistor, while double layers behave as capacitors: this means that it is possible to model the electrode-electrolyte interface as a theoretical electrical circuit, that is a sequence of capacitors and resistors. In this way, if the parallelism device/physical system is correct, the interface can be described with the classical laws used for electrical circuits and the experimental evidence (for example the electrochemical impedance analysis). This is the case of the electrochemical impedance spectroscopy.

For what concerns the anode, it is important to distinguish between inert and sacrificial electrodes. The first class of anodes does not participate in any electrochemical reaction, so it acts only as a carrier of charges and the anodic reaction may be oxygen evolution (for water-based electrolytes) or the oxidation of the organic molecules of the solvent (for non-water-based electrolytes). Sacrificial anodes, on the other hand, are used for the flow of electrons, but they also participate in the electrochemical reaction: they are dissolved, and the metallic ions generated are released in the electrolyte. Sacrificial anodes can be exploited for metal plating and they maintain a constant concentration of metallic ions in solution, which is an advantage because the cell is always working at the nominal condition. Inert anodes may be useful to avoid the contamination of the electrolyte with the metallic ions, to enhance oxygen evolution or to oxidise unwanted organic contaminants in solution.

Zinc is very used as sacrificial anode and for cathodic protection against corrosion, MMO (mixed metal oxide) is constituted by a core titanium net covered by mixed metal oxides able to favour oxygen evolution and it is an example of inert anode, as well as platinum and graphite. Examples of sacrificial anodes are on the other hand the lithium ones (in non-rechargeable batteries) or copper ones used in electroplating.

The electrolyte

The electrolyte is an ionic conductor which can be water-based or not. Some examples of this last category are molten salts (they are obtained at very high temperature to induce the melting of the salts, so no solvent is exploited; they are used in the production of aluminium and in the recovery of metals from end-of-life materials), polymer electrolytes (combination of ions coordinated with polymeric species, which is solid or gel-like and this can solve security issues) or solid oxides (like yttria stabilized zirconia). For solid electrolytes, the working mechanism is still ion migration, which is enhanced by high temperature, increasing diffusion at the solid state.

The main parameter that must be considered in choosing the right electrolyte is conductivity, which is higher for liquid electrolytes (higher mobility of the ions) than in solid ones (only high temperatures can induce an adequate mobility):

−1 = [ ] 7 −1− 10 Conductivity is much smaller than the one of electrodes, which can be made of metals (10³) or semiconductors (10⁴). Note that they do not have an Ohmic junction, for which there is a linear correlation between current and voltage, but a Schottky one, in which the characteristic curve is more complicated and losses may be important.

Conductivity depends on the following parameters:

• Carrier charges, so valence of the system. The higher the number of carriers, the higher the conductivity.
• Concentration of the ionic species.
• Mobility of the carriers. This is the reason why solid electrolytes have a lower conductivity. It can be increased with the temperature (it is the opposite for electronic conductors like metals, for which conductivity decreases with temperature). Mobility depends on the nature of the salt itself, so its size, molar weight, charge, and solvation shell.
• Viscosity is fundamental as it is related to the interactions between the ions.

The molar conductivity can be computed considering both the contributions of cations and anions:

= +

For sure an increase in concentration of the salt is followed by an increase in both the anionic and cationic contributes as they are balanced.