Electrolytic and Electrochemical Cells

The redox reaction in an electrolytic cell is nonspontaneous. Electrical energy is required to induce the electrolysis reaction. An example of an electrolytic cell is shown below, in which molten NaCl is electrolyzed to form liquid sodium and chlorine gas. The sodium ions migrate toward the cathode, where they are reduced to sodium metal. Similarly, chloride ions migrate to the anode and are oxided to form chlorine gas. This type of cell is used to produce sodium and chlorine. The chlorine gas can be collected surrounding the cell. The sodium metal is less dense than the molten salt and is removed as it floats to the top of the reaction container.

An electrolytic cell decomposes chemical compounds by means of electrical energy, in a process called electrolysis; the Greek word lysis means to break up. The result is that the chemical energy is increased. Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and bauxite into aluminium and other chemicals

An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. When driven by an external voltage applied to the electrodes, the electrolyte provides ions that flow to and from the electrodes, where charge-transferring, or faradaic, or redox, reactions can take place. Only for an external electrical potential (i.e. voltage) of the correct polarity and large enough magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided undoes the effect of spontaneous chemical reactions.

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3. Note that the site of oxidation is still the anode and the site of reduction is still the cathode, but the charge on these two electrodes are reversed. Anode is now + charged and the cathode has a – charged.

4. The conditions under which the electrolyte cell operates are very important. The substance that is the strongest reducing agent (the substance with the highest EHYPERLINK “http://www.saskschools.ca/curr_content/chem30/modules/module8/reduction.html”0HYPERLINK “http://www.saskschools.ca/curr_content/chem30/modules/module8/reduction.html” value in the table) will undergo oxidation. The substance that is the strongest oxidizing agent will be reduced. If a solution of sodium chloride (containing water) was used in the above system, hydrogen would undergo reduction instead of sodium, because it is a stronger reducing agent that sodium.

Electrochemical cell

An electrochemical cell is a device capable of either deriving electrical energy from chemical reactions, or facilitating chemical reactions through the introduction of electrical energy. A common example of an electrochemical cell is a standard 1.5-volt “battery”. (Actually a single “Galvanic cell”; a battery properly consists of multiple cells.

Half-cells

The Bunsen cell, invented by Robert Bunsen.

An electrochemical cell consists of two half-cells. Each half-cell consists of an electrode, and an electrolyte. The two half-cells may use the same electrolyte, or they may use different electrolytes. The chemical reactions in the cell may involve the electrolyte, the electrodes or an external substance (as in fuel cells which may use hydrogen gas as a reactant). In a full electrochemical cell,, species from one half-cell lose electrons (oxidation) to their electrode while species from the other half-cell gain electrons (reduction) from their electrode. A salt bridge (i.e. filter paper soaked in KNO3) is often employed to provide ionic contact between two half-cells with different electrolytes-to prevent the solutions from mixing and causing unwanted side reactions. As electrons flow from one half-cell to the other, a difference in charge is established. If no salt bridge were used, this charge difference would prevent further flow of electrons. A salt bridge allows the flow of ions to maintain a balance in charge between the oxidation and reduction vessels while keeping the contents of each separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the Bunsen cell (right).

Equilibrium reaction

Each half-cell has a characteristic voltage. Different choices of substances for each half-cell give different potential differences. Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions-when equilibrium is reached the cell cannot provide further voltage. In the half-cell which is undergoing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. Similarly, in the reduction reaction, the further the equilibrium lies to the ion/atom with the more negative oxidation state the higher the potential.

Electrode potential

The cell potential can be predicted through the use of electrode potentials (the voltages of each half-cell). The difference in voltage between electrode potentials gives a prediction for the potential measured.

Cell potentials have a possible range of about zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts, because the very powerful oxidizing and reducing agents which would be required to produce a higher cell potential tend to react with the water.

Electrochemical cell types

Main types

Cells are classified into two broad categories,

Primary cells irreversibly (within limits of practicality) transform chemical energy to electrical energy. When the initial supply of reactants is exhausted, energy cannot be readily restored to the electrochemical cell by electrical means.[1]

Secondary cells can be recharged; that is, they can have their chemical reactions reversed by supplying electrical energy to the cell, restoring their original composition.

Primary electrochemical cells

Primary electochemical cells can produce current immediately on assembly. Disposable cells are intended to be used once and discarded. Disposable primary cells cannot be reliably recharged, since the chemical reactions are not easily reversible and active materials may not return to their original forms.

Common types of disposable cells include zinc-carbon cells and alkaline cells. Generally, these have higher energy densities than rechargeable cells, but disposable cells do not fare well under high-drain applications with loads under 75 ohms (75 Ω).

Secondary electrochemical cells

Secondary electrochemical cells must be charged before use; they are usually assembled with active materials in the discharged state. Rechargeable electrochemical cells or secondary electrochemical cells can be recharged by applying electric current, which reverses the chemical reactions that occur during its use. Devices to supply the appropriate current are called chargers or rechargers.

The oldest form of rechargeable cell is the lead-acid cell.[5] This electrochemical cell is notable in that it contains a liquid in an unsealed container, requiring that the cell be kept upright and the area be well ventilated to ensure safe dispersal of the hydrogen gas produced by these cells during overcharging. The lead-acid cell is also very heavy for the amount of electrical energy it can supply. Despite this, its low manufacturing cost and its high surge current levels make its use common where a large capacity (over approximately 10Ah) is required or where the weight and ease of handling are not concerns.

An improved type of liquid electrolyte cell is the sealed valve regulated lead acid (VRLA) cell, popular in the automotive industry as a replacement for the lead-acid wet cell. The VRLA cell uses an immobilized sulphuric acid electrolyte, reducing the chance of leakage and extending shelf life.[6] VRLA cells have the electrolyte immobilized, usually by one of two means:

• Gel cells contain a semi-solid electrolyte to prevent spillage.
• Absorbed Glass Mat (AGM) cells absorb the electrolyte in a special fibreglass matting

Other portable rechargeable cells are (in order of increasing power density and cost): nickel-cadmium cells (NiCd), nickel metal hydride cells (NiMH), and lithium-ion cells(Li-ion).[7] By far, Li-ion has the highest share of the dry cell rechargeable market.[8] Meanwhile, NiMH has replaced NiCd in most applications due to its higher capacity, but NiCd remains in use in power tools, two-way radios, and medical equipment.[8]

Electrochemical Cells

Galvanic and Electrolytic Cells

Oxidation-reduction or redox reactions take place in electrochemical cells. There are two types of electrochemical cells. Spontaneous reactions occur in galvanic (voltaic) cells; nonspontaneous reactions occur in electrolytic cells. Both types of cells contain electrodes where the oxidation and reduction reactions occur. Oxidation occurs at the electrode termed the anode and reduction occurs at the electrode called the cathode.

Electrodes & Charge

The anode of an electrolytic cell is positive (cathode is negative), since the anode attracts anions from the solution. However, the anode of a galvanic cell is negatively charged, since the spontaneous oxidation at the anode is the source of the cell’s electrons or negative charge. The cathode of a galvanic cell is its positive terminal. In both galvanic and electrolytic cells, oxidation takes place at the anode and electrons flow from the anode to the cathode.

Galvanic or Voltaic Cells

The redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries. Galvanic cell reactions supply energy which is used to perform work. The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus that allows electrons to flow. A common galvanic cell is the Daniell cell, shown below.

An extremely important class of oxidation and reduction reactions are used to provide useful electrical energy in batteries. A simple electrochemical cell can be made from copper and zinc metals with solutions of their sulfates. In the process of the reaction, electrons can be transferred from the zinc to the copper through an electrically conducting path as a useful electric current.

An electrochemical cell can be created by placing metallic electrodes into an electrolyte where a chemical reaction either uses or generates an electric current. Electrochemical cells which generate an electric current are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells. In other electrochemical cells an externally supplied electric current is used to drive a chemical reaction which would not occur spontaneously. Such cells are called electrolytic cells.

Electrolytic Cells

The concept of reversing the direction of the spontaneous reaction in a galvanic cell through the input of electricity is at the heart of the idea of electrolysis. See for a comparison of galvanic and electrolytic cells. If you would like to review your knowledge of galvanic cells (which I strongly suggest) before learning about electrolytic cells,

Electrolytic cells, like galvanic cells, are composed of two half-cells–one is a reduction half-cell, the other is an oxidation half-cell. Though the direction of electron flow in electrolytic cells may be reversed from the direction of spontaneous electron flow in galvanic cells, the definition of both cathode and anode remain the same–reduction takes place at the cathode and oxidation occurs at the anode. When comparing a galvanic cell to its electrolytic counterpart, as is done in , occurs on the right-hand half-cell. Because the directions of both half-reactions have been reversed, the sign, but not the magnitude, of the cell potential has been reversed. Note that copper is spontaneously plated onto the copper cathode in the galvanic cell whereas it requires a voltage greater than 0.78 V from the battery to plate iron on its cathode in the electrolytic cell.

You should be asking yourself at this point how it is possible to make a non-spontaneous reaction proceed. The answer is that the electrolytic cell reaction is not the only one occurring in the system-the battery is a spontaneous redox reaction. By Hess’s Law, we can sum the ΔG of the battery and the electrolytic cell to arrive at the ΔG for the overall process. As long as that ΔG for the overall reaction is negative, the system of the battery and the electrolytic cell will continue to function. The condition for ΔG being negative for the system (you should prove this for yourself) is that Ebattery is greater than – Ecell.

An electrolytic cell is an electrochemical cell in which the energy from an applied voltage is used to drive an otherwise nonspontaneous reaction. Such a cell could be produced by applying a reverse voltage to a voltaic cell like the Daniell cell.

If a voltage greater than 1.10 volts is applied as illustrated to a cell under standard conditions, then the reaction

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

will be driven by removing Cu from the copper electrode and plating zinc on the zinc electrode.

Electrolytic processes are very important for the preparation of pure substances like aluminum and chlorine.

Electrochemical vs Electrolytic cells

Spontaneous vs non spontaneous

Electrochemical cells are powered by redox reactions which are spontaneous. These spontaneous redox reactions produce electrical energy that is harnessed in a battery. The reverse reaction in each case is non spontaneous and requires electrical energy to occur. The general form of the reaction can be written as:

Spontaneous ———->

Reactants

Products

+

Electrical Energy

<----------- Non spontaneous

When chemists reverse the electrochemical cell they need to provide electrical energy in order for the redox reaction to work in reverse. Cells created in this way are termed electrolytic cells. These cells are widely used to produce certain metals like sodium and aluminum from their oxides or ores, and also to electroplate gold and silver onto rings and other jewellery. How exactly does and electrolytic cell work?

Electrolysis Demonstration

Electrolysis is the process in which electrical current is used to cause a redox reaction to occur. Click on the batteries to perform a simple lab on electrolysis.

Comparison of Electrolytic and Electrochemical cells

Electrolysis of Water

During the early history of the earth, hydrogen and oxygen gasses spontaneously reacted to form the water in the oceans, lakes, and rivers we have today. That spontaneous direction of reaction can be used to create water and electricity in a galvanic cell (as it does on the space shuttle). However, by using an electrolytic cell composed of water, two electrodes and an external source emf one can reverse the direction of the process and create hydrogen and oxygen from water and electricity. shows a setup for the electrolysis of water.

The reaction at the anode is the oxidation of water to O2 and acid while the cathode reduces water into H2 and hydroxide ion. That reaction has a potential of -2.06 V at standard conditions. However, this process is usually performed with [H+] = 10-7 M and [OH-] = 10- 7 M, the concentrations of hydronium and hydroxide in pure water. Applying the Nernst Equation to calculate the potentials of each half-reaction, we find that the potential for the electrolysis of pure water is -1.23 V. To make the electrolysis of water occur, one must apply an external potential (usually from a battery of some sort) of greater than or equal to 1.23 V. In practice, however, it is necessary to use a slightly larger voltage to get the electrolysis to occur on a reasonable time scale.

Pure water is impractical to use in this process because it is an electrical insulator. That problem is circumvented by the addition of a minor amount of soluble salts that turn the water into a good conductor (as noted in ). Such salts have subtle effects on the electrolytic potential of water due to their ability to change the pH of water. Such effects from the salts are generally so small that they are usually ignored.

Electroplating

Electroplating allows the production of metal coatings of such desirable commodities as silver and gold. People make fortunes gold or silver plating junk metal (usually aluminum) because they can sell gold plated necklaces for a comparable price to the real thing (or even pass them off as being solid gold). That’s how electrochemistry can be used to rip you off! In our discussion of electroplating, we will discuss how you can set up a cell for electroplating, how you can calculate the amount of precious material consumed, and various other calculations you can perform with electroplating. In terms of the variety of electrochemistry problems possible to ask, this section, perhaps rivaled by Thermodynamics, is the richest.

The setup for electroplating is quite simple and the entire cell is usually conducted in a single solution as shown in .

The gold from the anode is oxidized and dissolves in solution as Au3+. The electrons arriving at the aluminum glasses frame cathode reduce the Au3+ in solution to Au (s) on the surface of the frame cathode. We can calculate how long we should have our glasses frame in solution if we want a certain amount of gold to be plated.

Let’s assume it takes 1.0 g of gold to provide an adequate coating for our glasses and also assume that we are using an emf sufficient to produce 10 amperes (A) of current (1 A = 1 coulomb per second). how long it will take to plate that 1.0 g of gold.

As you can see from the , such a problem only involves the use of unit cancellation. To calculate the time needed to deposit a certain amount of material, you need to start with the amount, converted to moles. Then, multiply by the number of electrons consumed in the reduction (in this case 3). Using the definition of a faraday, 96500 C per mole of electrons, you can convert between moles and charge. Finally, by using the definition of an ampere, 1 C per second, you can convert the amount of charge required to deposit the material into a time in seconds. There are various ways of phrasing this same problem such as “how much gold is deposited in 146 seconds at 10 A” or “what current is required to deposit 1.0 g of gold in 146 seconds.” Don’t be fooled by those permutations of the same problem, they all boil down to simple unit cancellation which you have been doing since you learned how to do stoichiometry. Also note that in these problems, you do not need to know the cell potential. Students often try, incorrectly, to use the cell potential somewhere in that calculation. Furthermore, you need only know the number of electrons transferred–you could solve the same problem without even knowing what material was being plated (as long as you know its molar mass).

Galvanic cells compared to electrolytic cells

In contrast, a battery or Galvanic cell, converts chemical energy into electrical energy, by using spontaneous chemical reactions that take place at the electrodes. Each galvanic cell has its own characteristic voltage (defined as the energy release per electron transfer from one electrode to the other). A simple galvanic cell will consist only of an electrolyte and two different electrodes. (Galvanic cells can also be made by connecting two half-cells, each with its own electrode and electrolyte, by an ion-transporting “bridge,” usually a salt bridge; these cells are more complex.) The electrodes typically are two metals, which naturally have different reaction potentials relative to the electrolyte. This causes electrons of one of the electrodes to preferentially enter the solution at one electrode, and other electrons to leave the solution at the other electrode. This generates an electric current across the electrolyte, which will drive electric current through a wire that makes an exterior connection to each of the electrodes. A galvanic cell uses electrodes of different metals, whereas an electrolytic cell may use the same metal for cathode and anode.

A rechargeable battery, such as a AA NiMH cell or a single cell of a lead-acid battery, acts as a galvanic cell when discharging (converting chemical energy to electrical energy), and an electrolytic cell when being charged (converting electrical energy to chemical energy).

Anode and cathode definitions depend on charge and discharge

Michael Faraday defined the cathode as the electrode to which cations flow (positively charged ions, like silver ions Ag+), to be reduced by reacting with (negatively charged) electrons on the cathode. Likewise he defined the anode as the electrode to which anions flow (negatively charged ions, like chloride ions Cl−), to be oxidized by depositing electrons on the anode. Thus positive electric current flows from the cathode to the anode. To an external wire connected to the electrodes of a battery, thus forming an electric circuit, the cathode is positive and the anode is negative.

Consider two voltaic cells, A and B, with the voltage of A greater than the voltage of B. Mark the positive and negative electrodes as cathode and anode, respectively. Place them in a circuit with anode near anode and cathode near cathode, so the cells will tend to drive current in opposite directions. The cell with the larger voltage discharges, making it a voltaic cell. Likewise the cell with the smaller voltage charges, making it an electrolytic cell. For the electrolytic cell, the external markings of anode and cathode are opposite the chemical definition. That is, the electrode marked as anode for discharge acts as the cathode while charging and the electrode marked as cathode acts as the anode while charging

Uses of Electrolytic cells

Electrolytic cells like the one above are used to commercial important metals from their ores. Aluminum is produced in this manner. The cells can also be used to electroplate gold and silver onto base metal such as iron. In these cases the object to be electroplated is attached to the – terminal of the battery and behaves as a cathode in the reaction.

Research activity:

Aluminum and copper

Aluminum and copper are two metals of great importance to society.

Using the internet to research these two metals, create a learning display that discusses the means of production, refining, uses, history and the impact use of these metals have had on our environment.