Galvanic and Electrolysis Study Notes PDF

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UNIT 3 CHEMISTRYAOS 1: What are theoptions for energyproduction?BATTERIES, FUEL CELLS ANDELECTROLYSISKey Knowledge Galvanic cells as a source of energy  Galvanic cells as primary cells and as portable or fixed chemical energy storage devices that can produce electricity (details of specific cells n...

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UNIT 3 CHEMISTRY AOS 1: What are the options for energy production? BATTERIES, FUEL CELLS AND ELECTROLYSIS

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Key Knowledge Galvanic cells as a source of energy 

Galvanic cells as primary cells and as portable or fixed chemical energy storage devices that can produce electricity (details of specific cells not required) including common design features (anode, cathode, electrolytes, salt bridge and separation of half-cells) and chemical processes (electron and ion flows, halfequations and overall equations).



The comparison of the energy transformations occurring in spontaneous exothermic redox reactions involving direct contact between reactants (transformation of chemical energy to heat energy) compared with those occurring when the reactants are separated in galvanic cells (transformation of chemical energy to electrical energy).



The use of the electrochemical series in designing and constructing galvanic cells and as a tool for predicting the products of redox reactions, deducing overall equations from redox half-equations and determining maximum cell voltage under standard conditions.

Fuel cells as a source of energy 

The common design features of fuel cells including use of porous electrodes for gaseous reactants to increase cell efficiency (details of specific cells not required).



The comparison of the use of fuel cells and combustion of fuels to supply energy with reference to their energy efficiencies (qualitative), safety, fuel supply (including the storage of hydrogen), production of greenhouse gases and applications.



The comparison of fuel cells and galvanic cells with reference to their definitions, functions, design features, energy transformations, energy efficiencies (qualitative) and applications.

Rechargeable batteries 

The operation of rechargeable batteries (secondary cells) with reference to discharging as a galvanic cell and recharging as an electrolytic cell, including the redox principles (redox reactions and polarity of electrodes) and the factors affecting battery life with reference to components and temperature (no specific battery is required). 2

Production of chemicals by electrolysis 

Electrolysis of molten liquids and aqueous solutions using different electrodes.



The general operating principles of commercial electrolytic cells, including basic structural features and selection of suitable electrolyte (molten or aqueous) and electrode (inert or reactive) materials to obtain desired products (no specific cell is required).



The use of the electrochemical series to explain or predict the products of an electrolysis, including identification of species that are preferentially discharged, balanced half-equations, a balanced ionic equation for the overall cell reaction, and states.



The comparison of an electrolytic cell with a galvanic cell with reference to the energy transformations involved and basic structural features and processes.



The application of stoichiometry and Faraday’s Laws to determine amounts of product, current or time for a particular electrolytic process.

HOMEWORK DUE DATE

CHEC K (√)

Textbook Reading

Heinemann Chemistry 2 Chapter 5 pg. 126-146 Chapter 6 pg. 152-159

Textbook Questions

Chapter 9 pg. 240-256 5.1 (pg. 132) q. 1, 3, 4 5.2 (pg. 138) all questions 5.3 (pg. 141) all questions 5.4 (pg. 147) all questions 6.1 (pg. 160) all questions 9.1 (pg. 245) all questions 9.2 (pg. 251) all questi

Textbook Chapter Review Questions

9.3 (pg. 257) all questi Chapter 5 Review q. 1, 18, 19, 22-26 Chapter 6 Review - all q

Chapter 9 Review q. 1SPONTANEOUS REDOX REACTIONS

When a relatively strong oxidant and relatively strong reductant are present in the same beaker, there is a direct transfer of electrons between the two. The reaction is exothermic i.e. chemical energy is transformed into thermal (heat) energy. For example, when zinc metal is placed into a solution of copper (II) sulfate, the following spontaneous reaction takes place: Oxidation: Zn

(s)

 Zn2+ (aq) + 2e-

Reduction: Cu

2+

-

(aq)

+ 2e  Cu

Overall equation: Zn (aq)

+ Cu

(s)

+ Cu2+

The zinc metal is oxidized and a brown deposit of copper metal forms on the zinc.

(s) (aq)

 Zn2+

(s)

GALVANIC CELLS A galvanic cell harnesses the energy produced by spontaneous redox reactions by separating the two half reactions so that the transfer of electrons occurs via an external circuit (a wire), rather than directly between the reductant and the oxidant. In this way, chemical energy is transformed into electrical energy. Batteries are examples of galvanic cells (sometimes referred to as voltaic cells). They provide a convenient source of electricity to power electronic

An electrochemical cell is a general term to describe a

devices, such as mobile phones, computers,

cell in which chemical

hearing aids, calculators etc. They range in size,

energy is converted to

from small button cells used to power small

electrical energy or vice

devices, such as watches, to huge batteries used to

versa (i.e. it refers to both

power lighthouses. Batteries are a more expensive

galvanic and electrolytic cells

source of energy than fossil fuels, however, their

(to be covered later)).

convenience and portability offset their cost. A galvanic cell involves separating the oxidation and reduction half reactions into two half cells so that they take place in different locations. A half cell must contain an electrode (something that conducts electrons) and an electrolyte (a solution or molten substance that contains free moving ions). Here is a typical diagram of a galvanic cell set-up:

I metal, it is used as the electrode in the oxidation half cell. It must be placed in an electrolyte with which it won’t react d l ti

usually a metal ion and so it forms the 4 electrolyte solution. If its conjugate reductant is a metal, then it is d th

containing its conjugate oxidant (i.e. its metal ions) is used.

electrode to ensure no reaction takes place between the electrode and electrolyte.

In a galvanic cell: 

The electrode at which oxidation takes place is termed the anode and the electrode at which reduction takes place is termed the cathode. The following mnemonic can be used to remember this:

An OIL RIG Cat



The anode is the negative electrode and the cathode is the positive electrode.



In order for electrons to flow between the reductant and the oxidant, the two half cells must be connected by an external conducting wire. The electrodes and the wire that joins them comprises the external circuit of the galvanic cell.



The direction of electron flow is from the anode (where electrons are lost via oxidation) TO the cathode (where electrons are gained via reduction).



A galvanic cell will not produce electricity without a salt bridge. Without a salt bridge, the solution in the oxidation half cell would accumulate positive charge and the solution

The salt bridge is usually a piece of filter paper soaked in an unreactive electrolyte (an aqueous ionic compound)

in the reduction half cell would accumulate negative charge. Thus, the salt bridge

such as potassium nitrate

completes the internal circuit of the cell by

(KNO3). The electrolyte needs to be unreactive so that it

neutralizing any growing charge in the

does not react with anything

solutions. This is achieved by: -

in either half cell.

Anions (negative ions) from the salt bridge migrating to the half cell that contains the anode.

-

Cations (positive ions) from the salt bridge migrating to the half cell that contains the cathode.

Half Cells Involving Metal Reactants and Products 

If the metal anode in a galvanic cell acts as the reductant, then the electrode will eventually corrode (disintegrate) because metal atoms from the electrode lose electrons to become aqueous metal cations. If the metal anode corrodes to the point where it is no longer in contact with the solution in the half cell, then the galvanic cell will stop working.



If metal ions in the reduction half cell serve as the oxidant then the cathode will grow in size

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because the metal ions gain electrons to become metal atoms, which deposit on the surface of the cathode. Half Cells Involving Aqueous Reactants and Products Some half reactions involve only aqueous species. E.g. Fe2+(aq)  Fe3+(aq) + eIf the reductant is aqueous, then an inert (unreactive) anode must be used so that it does not participate in or interfere with the oxidation reaction taking place in the half cell. The electrode still needs to be able to transfer electrons, however, and so a conductive material, such as platinum or graphite (carbon) is commonly used. In these half cells, the reductant and its conjugate oxidant are both present in the electrolyte solution. Half Cells Involving Gaseous Reactants Some half reactions involve gases. E.g. H2

(g)

 2H+(aq) + 2e-

Such reactions must also be carried out on the surface of an electrochemically inert conductor such as platinum. A tube is placed over the electrode so that the gas can be bubbled into the solution. Its conjugate redox nonmetal ion (e.g. H+) is in the electrolyte solution. ELECTROCHEMICAL SERIES Just as the electrochemical series can be used to predict whether or not a spontaneous redox reaction will take place between two chemicals, it can be used to determine the anode and cathode reactions in a galvanic cell. We know that the chemicals listed on the left hand side are oxidants and those on the right hand side are reductants. We also know that oxidation takes place at the anode and reduction takes place at the cathode. Therefore, because the half equation that is higher in the electrochemical series contains the stronger oxidant, it will serve as the reduction half cell and contain the cathode. The half equation that is 6

lower in the electrochemical series contains the stronger reductant and so it will undergo oxidation at the anode.

Note that the electrochemical series only predicts whether a reaction is possible. It gives NO indication of the reaction rate. Therefore, whilst possible, a reaction may be too slow to actually occur. A catalyst may be needed to lower the activation energy or the conditions changed (i.e.

WORKED EXAMPLE If we had a galvanic cell consisting of a Cu/Cu2+ half cell and a Zn/Zn2+ half cell, which electrode would act as the anode and which would act as the cathode?

The Cu/ Cu2+ half cell is higher in the electrochemical series and so at the cathode we the reduction of Cu2+: Cu2+(aq) + 2e-  Cu

(s)

The Zn/Zn2+ half cell is lower in the electrochemical series and so zinc acts as the anode. Recall that the bottom equation must be written in reverse to represent oxidation taking place. So, at the anode we have: Zn

(s)

 Zn2+ (aq) + 2e-

Potential Difference If a chemical has a large reduction potential, it means that it readily gains electrons (i.e. undergoes reduction). Similarly, if a chemical has a large oxidation potential, it means that is readily loses electrons (i.e. undergoes oxidation). The electrode potentials (Eo values) listed on the electrochemical series give us an indication of the reduction and oxidation potentials of each half reaction. Those with a positive electrode potential have a greater reduction potential (i.e. are more likely to be reduced), whereas those with a negative

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electrode potential have a greater oxidation potential (i.e. are more likely to be oxidised). Thus, for a spontaneous redox reaction to occur, the Eo value of the oxidant must be higher than the Eo value of the reductant. When an oxidant and reductant are connected via an external circuit in a galvanic cell, the difference in electrode potentials between the two half cell’s is measured using a voltmeter. This is known as the cell’s potential difference (also known as electromotive force (emf) or voltage) and is a measure of the degree to which electrons are “pushed” through the circuit. The higher the potential difference, the stronger the push of the electrons through the circuit and the greater the energy of the electric current produced. It can be calculated using the formula below:

Cell potential difference = Eoreduction half cell (cathode)– Eooxidation half cell (anode) (higher half cell)

(lower half cell)

WORKED EXAMPLE What is the potential difference of a galvanic cell constructed from a Mg2+/Mg half cell and Fe3+/Fe2+ half cell? At the cathode:

Fe2+ (aq)  Fe3+ (aq) + e-

Eo = + 0.77 V

At the anode:

Mg2+ (aq) + 2e-  Mg

Eo = – 2.37 V

(s)

Cell potential difference = 0.77 – (-2.37) = 0.77 + 2.37

= 3.14 V

Overall equation: Mg2+ (aq) + 2Fe2+ (aq)  2Fe3+ (aq) + Mg

(s)

Eo = 3.14 V

How are Electrode Potentials Determined? Chemists determined the standard electrode potentials (Eo values) for each half equation in the electrochemical series by connecting each half cell to a reference standard electrode, known as the standard hydrogen electrode (SHE), which is based on the half cell reaction: 2H+ (aq) + 2e-  H2 (g)

Eo = 0.00 V

The reference electrode is known as the standard hydrogen electrode and the electrode potentials of each half cell are 8

known as standard potentials because they only apply at standard laboratory conditions (SLC): 

1 M concentration of solutions



pressures of 1 atm



a temperature of 25⁰C (298 K)

If the temperature, pressure or concentration deviate from these conditions, the electrode potentials will differ from those that are listed on the electrochemical series.

The standard hydrogen electrode is connected to the negative terminal of a voltmeter in each case and is given an arbitrary electrode potential of 0 volts. Therefore, when connected to another half cell, the voltage measured by the voltmeter is the electrode potential of that half cell. A negative voltage means that the half cell to which the SHE is connected is more readily oxidised (i.e. has a higher oxidation potential). The direction of electron flow is from the other half cell to the SHE (anode  cathode). A positive voltage means that the half cell to which the SHE is connected is more readily reduced (i.e. has a higher reduction potential). The direction of electron flow is from the SHE to the other half cell. e-

CATHODE

e-

e-

ANODE

ANODE

e-

CATHODE

Determining Relative Reducing and Oxidising Strength Experimentally The relative strengths of reductants and oxidants can be determined by constructing galvanic cells from different combinations of half cells and using a voltmeter to find the polarity of each electrode. 

The stronger reductant will serve as the negative electrode (anode), where oxidation takes place.



The stronger oxidant will undergo reduction at the positive electrode (cathode).

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For example, if we wanted to determine whether copper is a stronger reductant than silver, then we can set up a galvanic cell consisting of a Ag+/Ag half cell connected to a Cu2+/Cu half cell. There are two ways in which the galvanic cell could be connected to the voltmeter: either the Cu2+/Cu half cell is connected to the positive terminal and the Ag+/Ag is connected to the negative terminal or vice versa, as pictured below.

If the electrodes are connected correctly, then the voltmeter will give a positive reading. This reading reflects the potential difference of the cell. If the voltmeter gives a negative reading, it means that the electrodes must switch terminals. The terminals to which the electrodes are connected reflect the polarity of the electrodes.

In the above example, a positive voltage reading is obtained when the Cu2+/Cu half cell is connected to the negative terminal of the voltmeter. This means that copper serves as the negative electrode (anode) and is a stronger reductant than silver. The silver electrode is the positive electrode (cathode) in this galvanic cell because it is connected to the positive terminal of the voltmeter. This means that silver ions are a stronger oxidant than copper ions.

However, if a Cu2+/Cu half cell is connected to a Pb2+/Pb half cell, copper acts as the cathode this time and the lead acts as the anode, meaning lead is the stronger reductant. + 0.39 V

Pb(s) electrode

Cu(s) electrode

10

Pb2+(aq)

Cu2+(aq)

Pb (s) Pb2+ (aq) + 2e-

Cu2+ (aq) +2e- Cu (s)

In this way, we can order oxidants and reductants according to their oxidising and reducing strengths, respectively, as shown below. Ag+ (aq) + e- Ag (s) Increasing oxidising strength

Cu2+ (aq) + 2e- Cu (s)

Increasing reducing strength

Pb2+ (aq) + 2e- Pb (s) Note that the potential difference (voltage) of a galvanic cell measured experimentally is often less than the theoretical voltage calculated from the electrode potentials listed on the electrochemical series. For example, in the above galvanic cell, the voltage is shown to be 0.39 V. However, based on the electrode potentials on the electrochemical series, the potential difference of the cell is expected to be 0.47 V: Cu2+ (aq) + 2e-  Cu

(s)

Eo = +0.34 V

Pb2+ (aq) + 2e-  Pb

(s)

Eo = -0.13 V

Cell potential difference = 0.34 – (-0.13) = 0.47 V This is because the electrode potentials listed on the electrochemical series only apply at standard laboratory conditions (25oC, 1 atm and 1 M) and so if there are differences in these conditions, this will cause changes to the potential difference of the cell. Other factors that may affect the cell’s voltage include the purity of the electrodes. If the electrode is impure or has a surface coating, such as an oxide layer, then the voltage will be affected. The salt bridge can also offer some internal resistance, thus causing the voltage to drop.

BATTERIES AND FUELS CELLS Galvanic cells that are used commercially exist as three main types: primary (nonrechargeable) cells, secondary (rechargeable) cells and fuel cells. Primary and secondary cells are examples of batteries and are generally called dry cells, because the electrolye used takes the form of a paste, rather than an aqueous solution. The paste electrolyte has enough moisture to allow

A battery is a package of one or more galvanic cells that stores chemical energy for later conversion to electrical energy.

ions to migrate throughout the cell, but is solid enough that the battery will function without leaking. This means it can operate in any 11

orientation, making it ideal for use in portable devices...