OXIDATION AND REDUCTION - BISSFacultyPD

OXIDATION AND REDUCTION - BISSFacultyPD

TOPIC 9 REDOX PROCESSES 9.2 ELECTROCHEMICAL CELLS ESSENTIAL IDEA Voltaic cells convert chemical energy to electrical energy and electrolytic cells convert electrical energy to chemical energy. NATURE OF SCIENCE (4.5) Ethical implications of research the desire to

produce energy can be driven by social needs or profit. INTERNATIONALMINDEDNESS Research in space exploration often centers on energy factors. The basic hydrogen-oxygen fuel cell can be used as an energy source in spacecraft, such as those first engineered by NASA in the USA. The International Space Station is a good example of a multinational project involving the international scientific

community. THEORY OF KNOWLEDGE Is energy just an abstract concept used to justify why certain types of changes are always associated with each other? Are concepts such as energy real? UNDERSTANDING/KEY IDEA 9.2.A Voltaic cells convert energy from spontaneous, exothermic chemical

processes to electrical energy. APPLICATION/SKILLS Be able to construct and annotate a voltaic cell. APPLICATION/SKILLS Be able to distinguish between electron and ion flow in a voltaic cell.

The redox reaction involves a transfer of electrons from the reducing agent (species oxidized) to the oxidizing agent (species reduced). When redox reactions occur in the same container, the chemical energy is released as heat so no useful work is

obtained. Chemists wish to harness this chemical energy. Voltaic Cells The key is to separate the oxidizing

agent from the reducing agent, thus requiring the electron transfer to occur through a wire. The current from the electrons in the wire can then be harnessed to provide useful work. A voltaic (galvanic) cell uses a spontaneous redox reaction to produce a current that can be used to do work. APPLICATION/SKILLS Be able to explain how a redox

reaction is used to produce electricity in a voltaic cell. The reaction in a voltaic cell is always a redox reaction which can be broken into two half-reactions. Oxidation occurs at the anode. Reduction occurs at the cathode. (My way to remember this is that

reduction involves the gain of electrons and electrons were discovered by the cathode-ray experiments.) If we want to harness the energy, we need to separate the reactions into half-cells.

Then the electrons are only allowed to flow through an external circuit. This creates the voltaic cell. The easiest way to make a half cell is to put a strip of metal into a solution of its own ions.

If you put a strip of zinc into a solution of zinc sulfate, the zinc atoms form ions by releasing electrons which make the surface of the metal negatively charged with respect to the solution. An equilibrium is set up between the metal and its solution of ions. This charge separation is the electrode potential. The more reactive a metal, the more negative its electrode potential in its half-cell.

If you connect the two half cells by an external wire, electrons will flow spontaneously from the electrode of higher potential to the one with lower potential. The half cells connected in this way are called electrodes. The electrode where reduction occurs is

called the cathode. The electrode where oxidation occurs is called the anode. To make a voltaic cell An electrical wire must connect to the metal electrodes in each half-cell. A voltmeter can be attached to the external circuit to record the voltage generated.

Electrons always flow from anode to cathode through the wire. A salt bridge completes the circuit. Without a bridge, no voltage will be generated. The bridge is a glass tube or such that contains an aqueous solution of ions such as KNO 3. This allows ion flow in the opposite direction so that potential difference can be maintained. Cations flow towards the cathode.

Anions flow towards the anode. Any two metal half-cells can be connected to make a voltaic cell. The direction of electron flow and voltage

generated will be determined by the difference in reducing strength of the 2 metals. In most cases this can be judged by the relative position of the metals in the reactivity series. Electrons will flow from the higher element to the lower element in the series. The greater reducing agent is always the anode. UNDERSTANDING/KEY IDEA 9.2.B Oxidation occurs at the anode (negative electrode) and reduction

occurs at the cathode (positive electrode) in a voltaic cell. In a voltaic cell, oxidation always occurs at the anode. Reduction always occurs at the cathode. In the voltaic cell, the anode has a

negative charge. The cathode has a positive charge. GUIDANCE For voltaic cells, a cell diagram convention should be learned. Sometimes it is convenient to use a shorthand notation rather than drawing out the whole cell. 1. A single vertical line represents a phase boundary. 2. A double vertical line represents the salt bridge.

3. Anode is on the left, cathode right. 4. Spectator ions are omitted. 5. If a half cell contains 2 ions, they are separated by a comma because they are in the same phase. salt bridge H2(g) H+(aq) Cu2+(aq) Cu(s) anode cathode 6. Always include phases in parentheses.

UNDERSTANDING/KEY IDEA 9.2.C Electrolytic cells convert electrical energy to chemical energy, by bringing about non-spontaneous processes. APPLICATION/SKILLS Be able to construct and annotate an electrolytic cell. APPLICATION/SKILLS

Be able to distinguish between electron and ion flow in an electrolytic cell. To review, a voltaic cell takes the energy of a spontaneous redox reaction and harnesses it to produce an electric voltage.

An electrolytic cell uses an external source of voltage to bring about a redox reaction that would otherwise be nonspontaneous. The reactant in an electrolytic cell is the electrolyte. This is usually a liquid either a molten ionic compound or a solution of an ionic compound.

As the electric current passes through the electrolyte, redox reactions occur at the electrodes removing the charges on the ions and forming products that are neutral elements. The ions are said to be discharged during this process.

Reactive metals such as aluminum, lithium, magnesium, sodium and potassium are found naturally in compounds such as Al2O3 and NaCl where they exist as positive ions. Extraction of the metal element alone therefore involves reduction of the ions. But remember the E of these metals are so low or negative that there are no good reducing

agents available to do this. This means that electrolysis is usually the only means by which the metals can be extracted from their ores. The longer lines in the power source represent the positive terminal and the shorter lines represent the negative terminal. Electrolyte Anions are

oxidized here. Cations are reduced here. UNDERSTANDING/KEY IDEA 9.2.D Oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative

electrode) in an electrolytic cell. The positive ions are attracted to the negative electrode (cathode) while the negative ions are attracted to the positive electrode (anode). At the anode (positive electrode),

negative ions lose electrons. (oxidation) At the cathode (negative electrode), positive ions gain electrons. (reduction) In both types of cells, electrons always flow from anode to cathode. Oxidation always occurs at the anode. Reduction always occurs at the cathode.

The difference is in the charge of the electrodes. Voltaic cells anode is neg, cathode is pos Electrolytic cells anode is pos, cathode is neg APPLICATION/SKILLS Be able to explain how current is conducted in an electrolytic cell.

The source of electric power is a battery or DC power source. This power source is indicated in the diagram with two vertical lines, one being longer than the other. The longer line represents the positive terminal. The shorter line represents the negative terminal.

The electrodes are connected to the power supply but are never touching. Electrodes are made of a conducting substance such as metal or graphite. Electric wires connect the power supply to the electrodes.

The power source pushes the electrons towards the negative electrode where they enter the electrolyte. This is the cathode. Electrons are released at the positive terminal, the anode, and returned to the source. The current is not passed through the electrolyte by electrons but by the ions that are mobile and migrate to the electrodes.

The chemical reactions occurring at each electrode remove the ions from the solution and enable the process to continue. APPLICATION/SKILLS Be able to deduce the products of the electrolysis of a molten salt.

When the electrolyte is a molten salt, the only ions present are from the salt. There is no solvent. It is easy to predict the products because the anion migrates to the anode and the cation migrates to the cathode. EXAMPLE

Describe the reactions that occur during the electrolysis of molten lead II bromide. Write an equation for the overall reaction and comment on any observable changes. First: Deduce the ions and the appropriate electrodes. PbBr2(l) Pb2+(l) + 2Br-(l) cathode

anode Deduce the half equations. Anode: Br- is oxidized to Br2 2Br- Br2 + 2eCathode: Pb2+ is reduced to Pb Pb2+ + 2e- Pb Overall reaction: Pb2+ + 2Br- Br2 + Pb Observations: A brown liquid Br2 with a strong smell at the anode and the appearance of a grey metal Pb at the cathode.

Most ionic compounds have very high melting points so it is costly to supply and maintain the energy for electrolysis. Sometimes another compound is added to lower the melting point to make it more economical.

You may have more than one ion attracted to the same electrode, thus E data is important. Citations International Baccalaureate Organization. Chemistry Guide, First assessment 2016. Updated 2015. Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed. N.p.: Pearson Baccalaureate, 2014. Print. Most of the information found in this power point comes directly from this textbook. The power point has been made to directly complement the

Higher Level Chemistry textbook by Catrin and Brown and is used for direct instructional purposes only.

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