Electrochemical Series 3b5k6o

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SUBMITTED TO-BY-KARANPREET KAUR

PRESENTED RAHUL RANJAN ROLL NO. – A-14 SECTION-

B6005 REGD.NO11010047

ACKNOWLEDGEMENT The below presented term paper was really an interesting and learning experience for me. But the completion of this project was quite a difficult task too. It came to be possible because of the help and of my friends, teachers and last but not the least seniors who played the major role in guiding me throughout my work. My fiends also gave a lot and helped me wherever I needed them. I would also like to thank my Chemistry Teacher (Nisha Mam) who showed belief on me and gave me the very interesting topic to work upon. I think I will be able to satisfy my mam with the work of mine.

- RAHUL RANJAN

INDEX: 1. INTRODUCTION 2. Electrode Potential 3. A note on hydrogen value 4. USES OF ELECTROCHEMICAL SERIES 5. Summarizing the electrochemical series 6. FIELDS USING ELECTROCHEMICAL SERIES 7. CATHODIC PROTECTION 8. CONCLUSION 9. REFERENCES

ABSTRACT: A series in which the metals and other elements are listed in the order of their chemical reactivity, the most active at the top and the less reactive or more “noble” metals at the bottom. In a broader sense such an activity series need not be limited to the metals but may be carried on through the electronegative (nonmetallic) elements as well. The electrochemical series as it applies to metals was first established by laboratory experiments in which the purpose was to determine which metals would displace others from solutions of their salts. By exhaustive experiments it becomes possible to draw up a complete list in the order of chemical activity, in which the metals at the top of the list are those which are found to give up their electrons most readily (that is, are the most electropositive elements). To obtain an accurate and reproducible activity series, it is best to use the electrode potential, or oxidation-reduction potential, which is defined as the voltage developed by a sample of pure metal immersed in a solution of one of its salts (at unit activity and at 25°C or 77°F) versus a hydrogen electrode immersed in hydrochloric or sulfuric acid of equivalent concentration

INTRODUCTION

Electrochemistry is a branch of chemistry that studies chemical reactions which take place in a solution at the interface of an electron conductor (a metal or a semiconductor) and an ionic conductor (the electrolyte), and which involve electron transfer between the electrode and the electrolyte or species in solution. In introductory chemistry, the reactivity series or electrochemical series is an empirical series of metals, in order of "reactivity" from highest to lowest. It is used to summarize information about the reactions of metals with acids and water, single displacement reactions and the extraction of metals from their ores. A serial arrangement of metallic elements or ions according to their electrode potentials determined under specified conditions; the order shows the tendency of one metal to reduce the ions of any other metal below it in the series. The electrochemical series is built up by arranging various redox equilibria in order of their standard electrode potentials (redox potentials). The most negative E° values are placed at the top of the electrochemical series, and the most positive at the bottom.

Electrode Potential In an electrochemical cell, an electric potential is created between two dissimilar metals. This potential is a measure of the energy per unit charge which is available from the oxidation/reduction reactions to drive the reaction. It is customary to visualize the cell reaction in of two half-reactions, an oxidation half-reaction and a reduction half-reaction. Reduced species -> oxidized species + ne-

Oxidation at anode

Oxidized species + ne- -> reduced species

Reduction at cathode

The cell potential (often called the electromotive force or emf) has a contribution from the anode which is a measure of its ability to lose electrons - it will be called its "oxidation potential". The cathode has a contribution based on its ability to gain electeons, its "reduction potential". The cell potential can then be written Ecell = oxidation potential + reduction potential If we could tabulate the oxidation and reduction potentials of all available electrodes, then we could predict the cell potentials of voltaic cells created from any pair of electrodes. Actually, tabulating one or the other is sufficient, since the oxidation potential of a half-reaction is the negative of the reduction potential for the reverse of that reaction. Two main hurdles must be overcome to establish such a tabulation 1. The electrode potential cannot be determined in isolation, but in a reaction with some other electrode. 2. The electrode potential depends upon the concentrations of the substances, the temperature, and the pressure in the case of a gas electrode.

In practice, the first of these hurdles is overcome by measuring the potentials with respect to a standard hydrogen electrode. It is the nature of electric potential that the zero of potential is arbitrary; it is the difference in potential which has practical consequence. Tabulating all

electrode potentials with respect to the same standard electrode provides a practical working framework for a wide range of calculations and predictions. The standard hydrogen electrode is assigned a potential of zero volts. The second hurdle is overcome by choosing standard thermodynamic conditions for the measurement of the potentials. The standard electrode potentials are customarily determined at solute concentrations of 1 Molar, gas pressures of 1 atmosphere, and a standard temperature which is usually 25°C. The standard cell potential is denoted by a degree sign as a superscript.

E°Ce ll

1. Measured against standard hydroden electrode. 2. Concentration 1 Molar 3. Pressure 1 atmosphere 4. Temperature 25°C

A note on hydrogen value: that each E° value shows whether the position of the equilibrium lies to the left or right of the hydrogen equilibrium.

That difference in the positions of equilibrium causes the number of electrons which build up on the metal electrode and the platinum of the hydrogen electrode to be different. That produces a potential difference which is measured as a voltage. Obviously if you connect one standard hydrogen electrode to another one, there will be no difference whatsoever between the positions of the two equilibria. The number of electrons built up on each electrode will be identical and so there will be a potential difference of zero between them.

According to standard oxidation potentials E° at 25°C (77°F). The electrochemical series as it applies to metals was first established by laboratory experiments in which the purpose was to determine which metals would displace others from solutions of their salts. By exhaustive experiments it becomes possible to draw up a complete list in the order of chemical activity, in which the metals at the top of the list are those which are found to give up their electrons most readily (that is, are the most electropositive elements). Such a list is shown in the table, where lithium exhibits the most reactivity as a metal. To obtain an accurate and reproducible activity series, it is best to use the electrode potential, or oxidation-reduction potential, which is defined as the voltage developed by a sample of pure metal immersed in a solution of one of its salts (at unit activity and at 25°C or 77°F) versus a hydrogen electrode immersed in hydrochloric or sulfuric acid of equivalent concentration.

USES OF ELECTROCHEMICAL SERIES: The electrochemical series tells us which metals are more active chemically than others and and which will replace others from a solution, such as is used in plating metals.  The biggest use is in creating batteries(Electrochemical Materials) which operate because of the differences in chemical activity between two metals. Zinc and Copper differ greatly in their chemical activity and when placed in a suitable environment (For example, placed in an acid solution(lemon) and connected with a wire or meter , will produce electricity which is a flow of electrons.

 The electrochemical series tells us which metal will most readily lose electrons to another substance, thereby producing the higher voltage or flow of electrons.

Summarizing the electrochemical series:Metals at the top of the series are good at giving away electrons. They are good reducing agents. The reducing ability of the metal increases as you go up the series.

Metal ions at the bottom of the series are good at picking up electrons. They are good oxidising agents. The oxidising ability of the metal ions increases as you go down the series.

Judging the oxidising or reducing ability from E° values:The more negative the E° value, the more the position of equilibrium lies to the left - the more readily the metal loses electrons. The more negative the value, the stronger reducing agent the metal is.

The more positive the E° value, the more the position of equilibrium lies to the right - the less readily the metal loses electrons, and the more readily its ions pick them up again. The more positive the value, the stronger oxidising agent the metal ion is.

FIELDS USING ELECTROCHEMICAL SERIES

Oxidizing and Reducing Strengths The electrochemical series helps to pick out substances that are good oxidizing agents and those which are good reducing agents. For example, a very high negative reduction potential of lithium electrode indicates that it is very difficult to reduce Li+ ions to Li atoms. Therefore, Li+ cannot accept electrons easily and so loses electrons to behave as a reducing agent. As the reduction potential increases (negative value decreases), the tendency of the electrode to behave as reducing agent decreases. Thus, all the substances appearing on the top of the series behave as good reducing agents. For example Li and K are good reducing agents while F- and Au are the poorest reducing agents. Similarly, substances at the bottom of the table have high reduction potential and they can be easily reduced. Therefore, they act as strong oxidizing agents. From the table we can conclude that H+ is a better oxidizing agent than Zn2+ while Cu2+ is a better oxidizing agent than H+; Fe is a better oxidizing agent than Cl2 and so on. All the substances appearing at the bottom of the table are good oxidizing agents.

Comparison of Reactivities of Metals

The relative ease with which the various species of metals and ions may be oxidized or reduced is indicated by the reduction potential values. The metals with lower reduction potential are not reduced easily but are easily oxidized to their ions losing electrons. These electrons would reduce the other metals having higher reduction potentials. In other words, a metal having smaller reduction potential can displace metals having larger reduction potentials from the solution of their salt. For example, copper lies above silver in the electrochemical series, therefore, if copper metal is added to AgNO3 solution, silver is displaced from the solution. In general a metal occupying higher position in the series can displace the metals lying below it from the solutions of their salts and so are more reactive in displacing the other metals. Thus, Li is the most electropositive element in solutions and fluorine is the most electronegative element.

Calculation of the EMF of the Cell The following steps determine the reduction potential of the cathode and anode:The two half-cell reactions are written in such a way that the reaction taking place at the left hand electrode is written as an oxidation reaction and that taking place at the right electrode is written as reduction reaction.The number of electrons in the two equations are made equal by multiplying one of the equations if necessary by a suitable number. However, electrode potential values (E°) are not multiplied.The electrode potentials of both the electrodes are taken to be reduction potentials and so the EMF of the cell is equal to the difference between the standard potential of the right hand side and the left hand side electrode. Ecell = ER - EL If the EMF of the cell is +ve, the reaction is feasible in the given direction and the cell is correctly represented, i.e., oxidation occurs at left electrode (anode) and reduction occurs at the right electrode (cathode). If it is -ve, the cell reaction is not feasible in the given direction and the cell is wrongly represented. Thus, to get positive value for the EMF the electrodes must be reversed.

Predicting the Liberation of Hydrogen Gas from Acids by Metals All metals having negative electrode potentials (negative E° values) show greater tendency of losing electrons as compared to hydrogen. So, when such a metal is placed in an acid solution, the metal gets oxidized, and H+ (hydrogen) ions get reduced to form hydrogen gas. Thus, the metals having negative E° values liberate hydrogen from acids.Metal having negative E° value.For example, metals such as Mg (E (Mg2+ Mg) = - 2.37 V),Zn (E (Zn2+ Zn) = - 0.76 V), Iron (E (Fe2+ Fe) = - 0.44 V) etc., can displace hydrogen from acids such as HCl and HSO4. But metals such as Copper, (E (Cu2+ Cu) = + 0.34V), silver (E (Ag+ Ag) = + 0.80V) and gold (E (Au3+ Au) = +1.42 V) cannot displace hydrogen from acids because of their positive reduction potential value.

Predicting Feasibility of a Redox Reaction

From the E° values of the two electrodes one can find out whether a given redox reaction is feasible or not. A redox reaction is feasible only if the species which has higher potential is reduced i.e., accepts the electrons and the species which has lower reduction potential is oxidized i.e., loses electrons. The electrochemical series gives the increasing order of electrode potentials (reduction) of different electrodes on moving down the table. This means that the species, which accept the electrons (reduced) must be lower in the electrochemical series as compared to the other which is to lose electrons. (oxidized). For example, from the electrochemical series E° value of Cu = +0.34 V and that of Ag = +0.80 V since the reduction potential of Ag is more than that of Cu, this means that silver has greater tendency to get reduced in comparison to copper. Thus, the reaction occurs more readily than the reaction .The reduction potential of copper is less than that of Ag, this means that copper will be oxidized or will go into solution as ions in comparison to Ag. Thus, the reaction, occurs more readily than Therefore, silver will be reduced and copper will be oxidized and the above reaction is not feasible. Rather the reverse reaction, can occur. Thus a metal will displace, any other metal, which occurs below it in the electrochemical series from its salt solution. When a metal having lower E° value is placed in a solution, containing ions of another metal having higher E° value, then the metal having lower E° value gets dissolved and the ions of the metal having higher E° value get precipitated.

CATHODIC PROTECTION Metal that has been extracted from its primary ore (metal oxides or other free radicals) has a natural tendency to revert to that state under the action of oxygen and water. This action is called corrosion and the most common example is the rusting of steel. Corrosion is an electro-chemical process that involves the age of electrical currents on a micro or macro scale. The change from the metallic to the combined form occurs by an“anodic” reaction: M → M+ + e(metal) (soluble salt) (electron) The anode and cathode in a corrosion process may be on two different metals connected together forming a bimetallic couple, or, as with rusting of steel, they may be close together on the same metal surface.

Regular maintenance showed that the Statue of Liberty suffered from galvanic corrosion

This corrosion process is initially caused by difference in natural potential in galvanic (bimetallic) couples ,Metallurgical variations in the state of the metal at different points on the surface, local differences in the environment, such as variations in the supply of oxygen at the surface (oxygen rich areas become the cathode and oxygen depleted areas become the anode).The principle of cathodic protection is in connecting an external anode to the metal to be protected and the ing of an electrical dc current so that all areas of the metal surface become cathodic and therefore do not corrode. In electro-chemical , the electrical potential between the metal and the electrolyte solution with which it is in is made more negative, by the supply of negative charged electrons, to a value at which the corroding (anodic) reactions are stifled and only cathodic reactions can take place.

Galvanic corrison in statue of liberty The most famous example of prevention of corrosion is the Ashoka Pillar at Sarnath. Built long long years ago, still it has not corroded just because of the coating of the metal with another metal. It shows that the process of coatng one metal with another more reactive metal to prevent corrosion had emerged centuries ago and is still in use because of its wide application.

CONCLUSION: The electrochemical series helps to pick out substances that are good oxidizing agents and those which are good reducing agents. As the reduction potential increases (negative value decreases), the tendency of the electrode to behave as reducing agent decreases. Thus, all the substances appearing on the top of the series behave as good reducing agents. The relative ease with which the various species of metals and ions may be oxidized or reduced is indicated by the reduction potential values. The metals with lower reduction potential are not reduced easily but are easily oxidized to their ions losing electrons. These electrons would reduce the other metals having higher reduction potentials. From the E° values of the two electrodes one can find out whether a given redox reaction is feasible or not. A redox reaction is feasible only if the species which has higher potential is reduced i.e., accepts the electrons and the species which has lower reduction potential is oxidized i.e., loses electrons.The electrochemical series gives the increasing order of electrode potentials (reduction) of different electrodes on moving down the table. This means that the species, which accept the electrons (reduced) must be lower in the electrochemical series as compared to the other which is to lose electrons. (oxidized).

Final conclusions are :1. The metal which has low standard reduction potential is more electropositive. 2. The metals which have low S.R.P. value are highly reactive. 3. The metal which has low S.R.P. value can replace other metal from its salt. 4. Metals which have low S.R.P. value are strong reducing agents. 5. Non-metals which have higher S.R.P. value and are strong oxidising agents. 6. Metals which have low S.R.P. value, their oxides are thermally stable. Metal which have high S.R.P. value, their oxides are thermally unstable. 7. Cation having higher S.R.P. value will discharge first at cathode and anion having low S.R.P. value discharge first at anode.

8. Metals which have low S.R.P. value have greater rate of corrosion.

REFERENCES

http://www.2classnotes.com/digital_notes.asp? p=Application_of_Electrochemical_Series 2.www.chemguide.co.uk/physical/redoxeqia/ec s.html 3.en.wikipedia.org/.../Standard_electrode_pot ential_(data_page) 4.en.wikipedia.org/wiki/Reactivity_series 5.www.tutorvista.com/.../electrochemicalmetals.php - United States 6.en.wikipedia.org/wiki/galvanic corrison 1.

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