HSC – Core Module 1: Production of Materials



HSC – Core Module 1: Production of Materials

4. Oxidation-reduction reactions are increasingly important as a source of energy

• Explain the displacement of metals from solution in terms of transfer of electrons

In many chemical reactions, metals donate electrons to positive ions (cations). Magnesium, for example, reacts with dilute acids such as hydrochloric acid and sulfuric acid. As the magnesium dissolves, it transfers its valence electrons to the hydrogen ions in the acid solution. The reaction of magnesium metal with hydrogen ions in the acid can be written symbolically as two separate half-equations. One half-equation shows the magnesium losing two electrons; the other shows the hydrogen ion gaining the electrons.

Reactive metals can also transfer their valence electrons to other cations. Thus a piece if zinc placed un copper sulfate solution quickly becomes coated in metallic copper. In this case the zinc transfers its valence electrons to the copper ions.

Reductants and oxidants

Reductant: An electron donor – also known as a ‘reducing agent’.

Reduction: The gain of electrons or decrease in oxidation state.

Oxidant: An electron acceptor; also called an oxidising agent.

Oxidation: the loss of electrons or increase in oxidation state.

Metals behave as reductants. When a metal reacts with other metal ions, we say they reduce the other metal ions by donating their electrons to these ions. The process of electron gain by the metal ions is called reduction.

Metal ions and hydrogen ions behave as oxidants. When metal ions react with other metals they remove the electrons from the metal. We say that the metal has been oxidised by the other metal ion. These types of reactions are commonly referred to as redox reactions.

Eg 1) The reaction between nickel metal and a solution of copper (II) ions is shown by the following net ionic equation. Discuss the process of oxidation and reduction with reference to this equation.

This equation shows that the solid nickel metal dissolves to produce nickel ions in solution. It also shows that the copper ions in solution are converted to solid copper metal. The nickel metal is the reductant that reduces the copper ions to copper metal. The copper ions are the oxidants that oxidise the nickel metal to nickel ions. This can be shown by the following half-equations.

Oxidation of Nickel:

Reduction of Copper (II) ions:

Eg 2) When a strip of magnesium metal is placed in a test tube containing dilute silver nitrate solution, the strip gradually becomes covered with a silvery-grey deposit of crystals. The crystals are separated and analysed and found to be silver metal. The balance net ionic equation is:

Write half equations for the oxidation of the reductant and the reduction of the oxidant

The magnesium is the reductant and is oxidised by silver ions to form magnesium ions. The silver ions are the oxidants and they are reduced by the magnesium to form silver metal. The net ionic equation is the sum of the two half equations written below.

Note: The reduction half equation is multiplied by 2 so that the electrons cancel out when the half-equations are added together.

• Account for changes in the oxidation state of species in terms of their loss or gain of electrons

The loss of gain of electrons is one way of describing redox reactions. Another way is to describe the reaction in terms of changes in oxidation states.

Oxidation state: A number given to an atom to indicate (theoretically) the number of electrons it has lost of gained (that is, its state of oxidation) also called the oxidation number

The oxidation state of an element is a measure of its degree of oxidation. Elements are assigned oxidation states according to the assumption that all bonds in the compounds are ionic. Although this is not true for many compounds, the concept is useful in explaining redox reactions.

Oxidation states of hydrogen and oxygen

Oxygen (in most oxide compounds) has an oxidation state of –II

Hydrogen (in most hydrogen compounds) has an oxidation state of +I

|Category |Oxidation state (OS) |Examples |

| | | |

|Elements (free) |0 | |

| | | |

|Simple Ions |Charge on the ion | |

| | | |

|Polyatomic Ions |Sum of the oxidation states of each element must | |

| |sum to the charge on the ion | |

| | | |

|Molecular Compounds |Sum of the oxidation states of each element must | |

| |sum to zero | |

The above table shows that all uncombined elements has zero oxidation. Their ions, however, have positive oxidation states or negative oxidation states depending on whether electrons have been lost or gained to form the ion. Thus copper ions have a higher oxidation stat ( ) than chloride ions ( ).

Oxidation states can be used to describe the oxidation and reduction processes in redox reactions

Generally:

Oxidation = Increase in oxidation state

Reduction = Decrease in oxidation state

Thus in a redox reaction, the oxidation states of the reductant increases and the oxidation state of the oxidant decreases.

Identify the relationship between displacement of the metal ions in solution by other metals to the relative activity of the metals

Redox reactions occurring between metal and metal ions are commonly called metal displacement reactions because one metal reacts and dissolves, while the other metal comes out of the solution as a solid deposit.

These displacement reactions can be used to create an activity series of metals related to their strength as reductants.

Consider the results of the following experiment in which four different metals (zinc, silver, magnesium and iron) are placed in test tubes containing copper (II) solution.

From the results obtained, it allowed us to establish an order of activity of these metals. Since the metals act as reductants, this order is also related to their strength as reductants. This order is from most active to least active:

Magnesium > Zinc > Iron > Silver

By comparing the reactions of a much wider variety of metal reductants with various metal ion oxidants, the relative ease of oxidation can be ranked for common metals.

Generally:

- Active metals are strong reductants

- Inactive metals are weak reductants

If we experimentally compare the strength of metal ions as oxidants, then we obtain a result that is in the reverse order. In this case the cations of inactive metals are much stronger oxidants than the cations of active metals.

Generally:

- Cations of inactive metals are strong oxidants

- Cations of active metals are weak oxidants.

A displacement reaction will only occur if the metal and the metal ion can successfully donate and accept electrons.

A spontaneous redox reaction will occur if the reductant is higher in the reduction half equation than the oxidant.

• Describe and explain galvanic in terms of oxidation/reduction reactions

In 1791, Luigi Galvani investigated the connected between metals and electricity using frog muscles. He observed that if two different metals were inserted into frog leg muscle, an electric current was generated.

Alessandro Volta discovered that a frog’s leg was not needed to generate electricity. He found that a bowl of salt water could be substituted and an electric current generated. He also found that by using several bowls of salt water and alternating copper and zinc electrodes in series generated a high voltage. This was the first working battery.

Volta later refined this battery by using metallic zinc and copper coins arranged in a vertical pile and separated by cardboard soaked in salt water. These ‘voltaic piles’ could be made to generate considerable currents by increased the number of coins in the pile.

These early devices are examples of electrochemical cells.

Galvanic cells are readily constructed from combinations of metals (electrodes) and electrolytes.

The cell is constructed from two half cells. One half-cell is the oxidation half cell; the other is the reduction half-cell.

Galvanic cells consisted of two half cells-linked by a metallic conductor (externally) and a salt bridge (internally).

In a galvanic cell there are two conducting terminals called electrodes. One electrode is called the anode and the other the cathode.

When the cells is operating, electrons leave the negative anode of the oxidation half-cell and travel to the positive cathode of the reduction half-cell through a conducting wire in response to a potential difference between the half cells.

Electrons and charges will not flow unless a complete circuit is present.

Note: Voltmeters are not necessary – they are only connected to measure the potential difference in the circuit.

An electrolyte or electrolyte gel is required for ions to move through the internal circuit as electrons move through the external circuit. In the electrolyte, cations move towards the reduction half-cell and anions move towards the oxidation half-cell. Solutions of nitrate salts such as potassium nitrate are commonly used as a salt bride (or ion bridge) in simple galvanic cells.

• Define the terms anode, cathode, electrode and electrolyte to describe galvanic cells

Galvanic Cell: An arrangement of electrodes and electrolytes in which a redox reaction causes a flow of electricity.

Half Cell: Either the oxidation or reduction half of an electrochemical cell

Electrode: The metallic conducting plates of a galvanic cell

Anode: The electrode at which oxidation takes place. This electrode is negative in a galvanic cell

Cathode: The electrode at which reduction takes place. This electrode is positive in a galvanic cell.

Electrolyte: A substance that releases ions when in solution of when melted and that carries an electric current.

Salt bridge: An electrolyte or electrolyte gel that joins two half-cells in a galvanic cell and allows movement of ions to maintain a balance of charges.

• Solve problems and analyse information to calculate the potential requirement of names electrochemical processes using tables of standard potentials and half equations.

Galvanic Cell Notation

A simpler way of expressing this galvanic cell notation is to use a cell diagram. Consider a call in which one half-cell is composed of a zinc electrode in a zinc nitrate solution and the other is made from a copper electrode in a copper (II) nitrate solution. The cell notation for this galvanic cell is:

This is composed of redox couples (the oxidant-reductant pair in a half equation)

The salt bridge is denoted by the || double vertical lines.

Nitrate ions are spectator ions and are involved in maintaining charge balance in the electrolytes.

It is conventional to show the oxidation half-cell at the left side of the cell diagram. The negative terminal of the voltmeter is connected to the anode and the positive terminal to the cathode. A positive voltage will then be registered.

In this example, the zinc is the more active metal and is the site of oxidation. The zinc nitrate solution is called the anolyte. The less active copper electrode is the site of reduction. The copper (II) nitrate solution is called the catholyte.

Standard Electrode Potentials

To obtain reliable and reproducible results when constructing galvanic cells, it is important to specify a set of standard conditions. These are:

- Electrolyte concentration = 1.0 mol/L

- Standard temperature = 25 C

- Standard pressure = 100 kPa

In addition, a standard reference half-cell is used to compare and test other half cells. This standard is the standard hydrogen half-cell . This half cell consists of an inert platinum electrode covered in platinum black, which is a highly porous form of platinum metal powder. This electrode is placed in a 1.0 mol / L solution of hydrogen ions and pure hydrogen gas is bubbled over the surface of the electrode under conditions of standard temperature and pressure.

The reduction half equation for this half-cell is:

The standard hydrogen half-cell is assigned a half-cell potential of zero volts. Thus, when a test half-cell is connected to the standard hydrogen half cell in a galvanic cell, the measure voltage (or standard cell potential) is equal to the half cell potential of the test couple.

The standard cell potential ( ) is defines as the sum of the standard half cell reduction potential and the standard half cell, oxidation potential.

The value of the standard cell potential may be positive or negative, depending on whether the test cell acts as a site of oxidation or a site of reduction.

Rules for predicting a spontaneous redox reaction:

• gather and present information on the structure and chemistry of a dry cell or lead-acid cell and evaluate it in comparison to one of the following:

▪ button cell

▪ fuel cell

▪ vanadium redox cell

▪ lithium cell

▪ liquid junction photovoltaic device (eg the Gratzel cell)

← in terms of:

▪ chemistry

▪ cost and practicality

▪ impact on society

▪ environmental impact

Batteries

Dry cells of Alkaline cells consist of only one galvanic cell. Technically, they are cells rather than batteries.

The Dry Cell

The dry cell (or the Leclanche cell) was first developed in 1866 by Georges Leclanche and has remained one of the most common and reliable sources of portable electric power.

Part of their appeal relies on their relative cheapness.

The central positive cathode consists of an inert graphite rod surrounded by graphite and manganese dioxide powder. The negative anode consists of the zinc casing of the cell. Between the two electrodes is an aqueous electrolyte paste containing ammonium chloride together with more powdered graphite and manganese dioxide.

An improved version of the normal dry cell is the more expensive alkaline cell. This 1.5V cell is used in higher drain appliances. The zinc anode is powdered so that it can deliver a higher, faster current for a longer time. Instead of the acidic paste, the electrolyte consists of a 7-molar potassium hydroxide solution. The cell has a life that is 5 times longer than the dry cell due to the larger amount of MnO2 present in the cell.

The Chemistry of a Dry Cell

|Dry Cell feature |Comment |

|Voltage |1.5 V |

|Anode (-) |Zn Casing |

| | |

|Anode Half Equation | |

|Cathode (+) |MnO2 , C |

| | |

|Cathode Half Equation | |

|Electrolyte |Aqueous paste of ammonium Chloride (26% w/w) |

|Other information |The MnO2 ensure that hydrogen atoms released from the ammonium ions during reduction are converted to water. As the |

| |cell operates, ZnCl2 forms in the anode; this contributes to the conductivity of the paste. The ammonia that forms at |

| |the cathode complexes with the zinc ions to form a stable complex ion |

|Cost / Practicality |The materials of the cell are inexpensive and cheap to replace even though the cell is non-rechargeable. These cells |

| |have low energy density. The voltage falls during use, due to a drop in electrolyte concentration around the cathode |

| |and the time required for the Mn2O3 to diffuse away from the cathode. |

| |These cells have a short shelf life as the zinc is attached by the ammonium chloride. As well, the zinc casing oxidizes|

| |during cell discharge and the acidic paste can leak out through cracks and corrode other components. |

|Effect on society |They are used in low-drain appliances such as torches, remote controls, LCD calculators and battery powered toys that |

| |do not require high currents. |

|Effect on Environment |Battery components are weakly acidic and non-toxic. There are few environmental consequences on disposal. |

The Lead-Acid Storage Battery

Unlike the dry cell, lead-acid car storage batteries can be recharged by the application of an external current. The battery us used during start-up to provide the energy for the car’s starter motor. The battery is gradually recharged during driving using electrical energy from the car’s alternator.

The diagram below shows the structure of a lead-acid storage battery. Six cells, each supplying about 2V are arranged in series to produce a battery with a voltage of about 12V. In each cell, the anode is composed of a porous lead sheet, and the cathode is made of a lead sheet coated in compressed lead (IV) oxide. Each electrode is in the form of a grid to increase the surface area. The electrolyte in each cell is sulfuric acid. In the battery, thin, perforated fiberglass sheets separate each electrode.

|Battery Features |Comments |

|Voltage |About 12 Volts |

|Anode (-) |Pb Sheet |

|Anode half-equation | |

|Cathode (+) |PbO2 powder on lead sheet |

|Cathode half-equation | |

|Electrolyte |Sulfuric Acid (35% w/w) |

|Other information about the battery |The life of the battery is limited by various factors including: |

| |a) the lead sulfate disintegrating from the electrode surrounding surface affects the ability of the battery to be recharged |

| |as it ages |

| |b) slow corrosion of the lead anode |

| |c) internal short circuiting. |

| |During discharge, the density of the electrolyte drops from 1.26 g/cm3 to 1.1g/cm3 as hydrogen ions are consumed. A battery |

| |hydrometer can be used to measure the density of the electrolyte and this is used as a measure of the state of charge of the |

| |battery. |

|Cost and practicality |The batteries are expensive due to lead content. The battery lasts many years and can recharged over and over again. |

| |Batteries can also be recharged externally to the car using a suitable transformer. Care must be taken to avoid too-rapid |

| |recharging as explosive hydrogen gas can form. These batteries are heavy and this limits their portability. They have the |

| |lowest energy density of most commonly used rechargeable batteries. |

|Effect on society |This is an important battery for car start up motors as it provides high currents over short periods. Also, it is useful |

| |storage battery to remote locations. It can be recharged by connecting to solar panels or electric generators. Useful for |

| |emergency lighting. |

|Effect on the environment |Lead acid batteries are recycled to retrieve the lead. Lead is toxic to organisms in the environment. It causes anaemia in |

| |humans. The electrolyte is highly acidic and can cause severe damage if spilled. Sealed lead-acid cells prevent acid fumes |

| |from causing corrosion |

Lithium Ion batteries

Lithium is an active metal, so batteries that use lithium must contain no water, but use non-aqueous solvents instead. Batteries that utilize lithium ions rather than lithium metal are much safer, and it is this type of battery technology that powers mobile phones and laptops. This type of battery has been used in the probes sent to Mars by NASA.

|Battery Feature |Comment |

|Voltage |3.6 Volts |

|Anode (-) |C (graphite) containing Li ions bound the crystal layers |

|Anode Half Equation | |

|Cathode (+) |Lithiated metal oxides eg ( ); contains Li ions bound to the lattice |

|Cathode Half Equation | |

|Electrolyte |LiPF6 or LiBF6 in an organic solvent |

|Other information about the battery |Microporous polyethylene membranes separate ions and electrodes. During discharge the lithium ions travel from the anode to the |

| |cathode. During recharge the lithium ions travel from the cathode back to the anode. |

|Cost and practicality |Expensive. Operates over a wide temperature range. Rechargeable. High energy density. Low maintenance. No memory effect. Fragile |

| |and so needs a protective circuit to limit peak voltages. Low self-discharge. Small and Light. Long recharging times |

|Effect on society |Major uses include laptop and mobile phone batteries as well as satellite batteries. This has had a significant impact on |

| |communications technology. |

|Effect on the environment. |The components are environmentally safe as there is no free lithium metal. |

The Vanadium Redox Battery

The vanadium is a redox flow battery system invented at UNSW in 1985.

The vanadium redox battery uses two tanks to store vanadium electrolytes which are pumped through a battery stack where the chemical energy is stored in the solutions is converted to electrical energy. The anolyte contains vanadium ions in the +II and +III oxidation states. The catholyte contains vanadium ions in the +IV and +V oxidation states. The potential difference between these two half-cells leads to electron flow in the external circuit.

Diagram on Back

|Battery |Comment |

|Voltage |1.26 Volts |

|Anode (-) |Bipolar electrodes composed of graphite impregnated polypropylene or graphite felt |

|Anode | |

|half-equation | |

|Cathode (+) |Bipolar electrodes composed of graphite impregnated polypropylene or graphite felt |

|Cathode half-equation | |

|Electrolytes | |

| | |

|Other information about |The technical benefits of this battery include: |

|the battery |Separate storage tanks can be used to store the chemical energy. Those tanks are separated from the battery. |

| |The system is safe because there is less of risk of electrolytes suddenly mixing and releasing energy |

| |The solutions act as coolants as they are pumped through the stack of cells inside the battery. |

| |Vanadium solutions are present in both compartments, so contamination is less of a problem |

| |There is a very high-energy efficiency as the electrochemical reactions are readily available. |

|Cost and practicality |High cost until fully commercialized. Vanadium is a plentiful metal and this will help to reduce the costs. The battery can be recharged|

| |at a low voltage (eg 2V) but can deliver a high voltage on discharge. |

| |Very little explosive hydrogen is generating in recharging, so natural ventilation is sufficient. |

| |Because there are no changes to the electrodes, the cells can also be recharged by replacing the spent solutions with fresh ones. (old |

| |ones are regenerated at a more convenient place) |

|Effect on Society |Some important applications of this technology include: |

| |Electrochemical storage of solar energy collected by photovoltaic cells |

| |Replacement of lead-acid storage batteries used to power up diesel engines, especially in remote areas; lead-acid storage batteries are |

| |durable but present environmental problems in disposal as lead is a toxic metal |

| |Electrochemical storage of wind energy collected by wind generators operating in low wind areas; this is made possible because the |

| |vanadium battery can be recharged at low voltages, acting as an emergency backup battery system. |

| |Power source for electric vehicles. |

|Effect on the |The solution can be indefinitely recycled and waste is minimized. |

|environment. | |

Extra Information for Conquering Chemistry.

The area of science that is concerned with the relations between chemistry and electricity is called electrochemistry.

Displacement reactions

A displacement reaction is a reaction in which a metal converts the ion of another metal to the neutral atom.

Electricity from Redox Reactions

Redox reactions involve the transfer of electrons from one reactant to another.

An electric current is a flow of electrons through a wire.

We can make redox reactions generate electricity by arranging for the oxidation and reduction of half reactions to occur at a different locations, and by providing a wire for the electrons to flow through. This occurs in all batteries we use.

Consider the following:

A copper strip of metal is suspended in a beaker of copper nitrate solution, and a spiral of silver wire in a beaker of silver nitrate solution. The two solutions are connected by a U-tube filled with a solution of potassium nitrate held in place by plugs of cotton wool. This U-tube which makes electrical contact between two solutions is called a salt-bridge; in order to make electrical contact it must contain some conducting substances such as a solution of potassium nitrate.

If a voltmeter is now connected across the pieces of metal, it is found that the silver wire is about 0.5 volt positive with respect to the copper.

The electron flow is from the copper through the external circuit to the silver wire.

If the reaction is left for a period of time the following changes will occur:

1) Metallic silver deposits on the silver wire – as can be seen by inspection.

2) Some of the copper strip dissolves, can be confirmed by weighing

3) The concentration of silver ions in the right beaker falls

4) The concentration of copper ions in the left beaker increases.

In other words electricity has been produced by a chemical reaction.

What is the purpose of the Salt Bridge?

If we remove the salt bridge, the voltage falls to zero and no current flows. Hence the salt bride is necessary to make this and other galvanic cells operate.

If the only changes were the ones already described, then the beaker containing copper nitrate would end up with an excess of positive copper ions, and the silver nitrate beaker would end up with an excess of negative nitrate ions.

It is impossible to have an imbalance of positive and negative ions in any solution, hence other changes must have occurred.

A closer inspection of the solutions shows that the nitrate concentration in the copper nitrate solution also increased, and the silver nitrate solution decreased in both silver and nitrate concentration.

This implies that there has been a migration of nitrate ions away from the solver nitrate solution through the salt bridge and into the copper nitrate solution.

The purpose of the salt bridge is to allow this migration to occur.

Both positive and negative ions migrate through the salt bridge. Not only do the nitrate ions migrate from right hand beaker to the left hand one, but also positive ions migrate from the right hand to the left. Copper ions move into the salt bride and push the potassium ions out into the silver nitrate solution. This migration preserves electrical neutrality in both solutions (half-cells).

The electrolyte used in the salt bridge must be one which does not react with any of the ions in the two solutions it is connecting.

When a galvanic cell produces electricity.

1) One electrode reactions liberates electrons which flow out of the metal of the electrode and into the external circuit

2) These electrons flow through the metallic conductor of the external circuit to other electrode

3) The reaction at the other electrode consumes these electrons

4) Ions migrate through the solutions and connecting salt bridge to maintain electrical neutrality.

An electric current through a metallic conductor is a flow of electrons. Through a conducting solution it is a migration of ions; negative ions move through the solution in the direction that completes the circuit for the electrons.

Positive ions move in the opposite direction. The ‘connection’ between the ions and the electrons is made by the electrode reactions which occur where the metallic electrode meets the solution.

The negative terminal of a galvanic cell is by definition the electrode from which electrons flow out into the external circuit.

Oxidation occurs at this negative electrode to provide the electrons for the external circuit. The positive terminal of the battery draws electrons back to the cell from the external circuit. Hence the reaction which occurs at the positive terminal to do this is reduction.

A galvanic cell is an ‘electron pump’; it pumps electrons out of the negative terminal into the external circuit and ‘sucks’ them back into the positive terminal. It can do this because a redox reaction is occurring in the cell.

Why are some cells non-rechargeable?

Some commercial galvanic cells can be used once only and then have to be thrown away. This is because if we attempt to recharge them, we do not simply reverse the reaction but instead bring about a different reaction. For example for a dry cell, the reaction needed is:

However when a current is passed through this cell the reaction that occurs is:

This is not only ineffective for recharging the cell, but is quite dangerous (because of the explosive nature of hydrogen).

Standard Electrode Potentials

The electromotive force or EMF of a galvanic cell is the potential difference (voltage) across the electrodes of the cell when a negligibly small current is being drawn. It is the maximum voltage that the cell can deliver.

Calculating EMF’s for Redox reactions and cell

Just as half reactions can be added to form complete redox reactions, so too standard electrode potentials can be combined to calculate EMF’s of complete reactions.

For example

First we break it down into its half equations:

Since the required reaction is made by the first equation,, it is made by adding the two halves. Thus it follows that:

Or in symbols:

Note: the half equation is double what is written on the standard potential table.

Doubling the half equation does not alter the

Since it known as the energy per electron. It does not matter how many electrons are in the reaction – the energy per electron is unaltered

In order to calculate the EMF of a galvanic cell:

1) Write the half reactions

2) Then follow the above procedure.

Also, when considering galvanic cells with electrodes A and B:

Where EMF is the voltage of electrode A relative to electrode B.

When comparing two reduction half equations, the one with the larger standard electrode potential has the greater tendency to occur. As it occurs, it drives the other half reaction in the reverse direction.

We can generalize from these examples:

The greater the standard electrode potential, the greater is the oxidizing strength of the oxidised form (left hand side) of the redox half equation.

Alternatively, the algebraically smaller the standard electrode potential, the greater is the reducing strength of the reduced form (right-hand side) of the redox half reaction.

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