What are the most common type of reactions?
Oxidation-Reduction Reactions (Redox) among the most common and important
What processes are redox reactions involved in?
rusting of iron, manufacture and action of bleaches, respiration of animals
Oxidation
loss of electrons (electrons are products)
Reduction
gain of electrons (electrons are reactants)
When do oxidation-reduction reactions occur?
when electrons are transferred from the atom that is oxidized to the atom that is reduced
Describe the thermodynamics of redox reactions
produce energy in the form of heat or electricity- thermodynamically "downhill" (spontaneous)
Are redox reactions spontaneous?
yes
What is electricity used for?
to make nonspontaneous processes occur
Electrochemistry
study of the relationships between electricity and chemical reactions
20.1 Oxidation-Reduction Reactions
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How is it determined whether a given chemical reaction is an oxidation-reduction reaction?
done by keeping track of oxidation numbers of all the elements involved in the reaction
Oxidation Number
the charge (in general)
What must occur in any redox reaction?
both oxidation AND reduction
Oxidizing Agent/Oxidant
substance that makes it possible for another substance to be oxidized - removes electrons from another substance by acquiring them itself (therefore is itself reduced)
Reducing Agent/Reductant
substance that gives up electrons causing another substance to be reduced (this substance is oxidized in the process)
20.2 Balancing Oxidation-Reduction Equations
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Important to note when balancing a chemical equation
law of conservation of mass (amount of each element must be the same on both sides of the equation)
What is special about balancing redox reactions?
the gains and losses of electrons must be balanced
Half-reactions
equations that show either oxidation or reduction alone - provides a general method for balancing redox reactions
How is an acidic solution balanced?
1. Divide the equation into two incomplete half-reactions
2. Balance each half-reaction
a. first balance the elements other than H and O
b. next, balance the O atoms by adding H20
c. then balance the H atoms by adding H+
d. finally balance the charge by a
20.3 Voltaic Cells
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Voltaic/Galvonic Cell
device in which the transfer of electrons takes place through an external pathway rather than directly between reactants (energy released from spontaneous redox reaction can then be used to perform electrical work)
Electrodes
two solid metals that are connected by the external circuit
Anode
electrode at which oxidation occurs (more +, lose e)
Cathode
electrode at which reduction occurs (more -, gain e)
Half-cell
name of each of the two compartments of the voltaic cell
Which way do electrons flow in a voltaic cell?
towards the cathode
What purpose does the salt bridge serve?
keeps both half-cells electrically neutral (contains electrolyte solution whose ions will not react with other ions in the cell)
Migration patterns
anions migrate to the anode, cations migrate to the cathode
20.4 Cell EMF
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Why do electrons transfer spontaneously in a voltaic cell?
electrons flow from the anode of a voltaic cell to the cathode because of a difference in potential energy (like a waterfall) - difference in potential energy per electrical charge is measured in volts
Electromotive Force (emf)/Cell Potential
potential difference between two electrodes of a voltaic cell - provides the driving force that pushed electrons through the external circuit
Cell Voltage
cell potential or emf because it is measured in volts
Cell potential will be positive for...
any reaction that proceeds spontaneously
What does the mf depend on?
specific reactions that occur at the cathode and anode, the concentrations of reactants and products, and the temperature
Standard Conditions
1 M concentration, 1 atm pressure, 25 degrees Celsius
Standard emf/cell potential
E^o cell - emf under standard conditions (superscript ^o indicated standard-state conditions)
Calculating Standard Reduction Reaction Cell Potential
Ered(cathode) - Ered(anode)
How is a half-cell potential written?
as a reduction - E^o red (intensive property and therefore changing coefficient makes no difference)
Standard Reduction Potentials Chart
higher means cathode and reduction and higher reducing agent
What information is presented with the strength of the driving force?
more positive Ered, greater driving force of reduction
Ered cathode vs anode
cathode is more positive
What information is presented with the sign of E?
more positive Ered = greater tendency for the reactant to be reduce or oxidize another species
Most common oxidizing agents
halogens, O2, and oxyanions (MnO4-...)
Most common reducing agents
H2 and active metals
20.5 Spontaneity of Redox Reactions
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Voltaic cells use redox reactions that proceed...
spontaneously
Using E to determine spontaneity
Positive E = spontaneous
Free energy charge
Delta G = -nFE where n is the number of electrons transferred, F is Faraday's constant (1 F = 96,500 C/mol) - note: both n and F are positive numbers
Spontaneity determined by delta G
negative delta G = spontaneous
20.6 Effect of Concentration on Cell EMF
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Nerst equation
calculate emf at nonstandard conditions - helps to understand why the emf of a voltaic cell drops as it discharges
Relationship between concentrations and emf
[reactants] up = emf up, [products] up = emf down
Concentration cell
contains the same substances on either side with different concentrations
20.7 Batteries
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Battery
portable, self-contained electrochemical power source that consists of one or more voltaic cells
Labels on battery parts
cathode labeled +, anode labeled -
Primary Batteries
cannot be recharged, one time use - must be discarded after emf drops to 0
Secondary Batteries/cell
can be recharged from an external power source after its emf has dropped
Lead-Acid Battery
6 voltaic cells each producing 2 V - cathode of lead dioxide (PbO2) and anode of Pb - both immersed in sulfuric acid
What happens when a battery is recharged?
direction of the reaction is reversed - done in a car by a generator
What is the most common primary battery?
alkaline battery - zinc anode with MnO2 cathode
20.8 Corrosion
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Corrosion Reaction
spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound
Galvanized iron
iron coated with a thin layer of zinc
Cathodic protection
protecting a metal from corrosion by making it the cathode in a electrochemical cell
Sacrificial anode
the metal that is oxidized while protecting the cathode
20.9 Electrolysis
x
What are electrolysis reactions?
process in which electricity is used to cause nonspontaneous redox reactions to occur
Where does electrolysis take place?
electrolytic cells - consists of two electrodes in a molten salt or solution
What does electrolysis of molten salts require?
high temperatures due to the high melting points of ionic substances
Inert electrodes
did not undergo reaction but merely served as the surface where oxidation and reduction occurred
Active electrodes
participate in the electrolysis process
Electroplating
uses electrolysis to deposit a thin layer of one metal on another in order to improve beauty or resistance
Units for charge passing through electrical circuit
coulombs - 1 mol of electrons has a charge of 96,500 C or 1 faraday (F) therefore 1 F = 96,500 C/mol e
Calculating coulombs
coulombs = (amperes)(seconds)