As you watch the following video, answer the questions on the sheet provided: Decoding the Past - The Real Dr. Frankenstein
Oxidation-reduction reaction (Redox) Involves a transfer of electrons from the reducing agent to the oxidizing agent Oxidation Loss of electrons LEO OIL Reduction Gain of electrons GER RIG Oxidation number Oxidizing agent Reducing agent Half-reactions Overall reaction is split into two half-reactions, one involving oxidation and one reduction Work on the sample Review Problem
Balance the following oxidation-reduction reaction IN BASIC SOLUTION using the half-reaction method. Be sure to identify the oxidizing agent and reducing agent. HXeO 4 - (s) XeO 6 4- + Xe (g)
Luigi Galvani Italian physician who observed a frog’s leg twitch when it was touched with two different metals In attempting to explain what happened, Galvani thought that the animal tissue in the frog’s leg was the source of electricity Alessandro Volta Italian physicist who disputed Galvani’s hypothesis Resulting controversy resulted in discovery that electric currents could be produced by chemical reactions Volta used this discovery to create the first chemical battery
In general, all chemistry is electrical in the sense that it involves the behavior of electrons and other charged particles The term electrochemistry is reserved specifically for the study of the interchange of chemical and electrical energy
All of the following involve the principles of electrochemistry: Remote controls for TVs, DVD players, CD players, stereos Itty bitty teeny tiny batteries Calculators Silverware Metal-plated jewelry
Defined as a device in which chemical energy is changed to electrical energy Examples – batteries and fuel cells Name comes from the work of Volta and Galvani Uses a spontaneous redox reaction to produce a current that can be used to do electrical work
Electrons flow from one terminal to the other when the terminals are connected by an external circuit Terminals are called electrodes Anode Oxidation occurs here Electrons leave the cell here Cathode Reduction occurs here Electrons are accepted by the species being reduced and enter the cell here Electrodes are submerged in an electrolyte A salt solution that contains ions Electrolyte may be involved in the reaction or the ions may be used to carry the charge Can contain a salt bridge or a porous-disk connection in order to neutralize charge buildups in electrode compartments Completes the circuit!
Anode Oxidation occurs here Cathode Reduction Occurs here
AN OX Oxidation occurs at the anode RED CAT Reduction occurs at the cathode FAT CAT The electrons in a galvanic (voltaic) cell always flow From the Anode To the CATode
Recall that a galvanic cell consists of an oxidizing agent in one compartment that “pulls” electrons through a wire from a reducing agent in the other compartment The “pull” or driving force on the electrons is called the CELL POTENTIAL or electromotive force (emf) If pull occurs spontaneously, cell is a good battery!
A VOLTMETER is used to measure cell potential The unit of electrical potential is the volt (V) Defined as 1 joule of work per coulomb of charge transferred
The cell potential (always positive for a galvanic cell) and the balanced cell reaction is written somewhere on the diagram The direction of electron flow is given Obtained by inspecting the half-reactions and using the direction that gives a positive E 0 cell The anode and cathode are designated The nature of each electrode and the ions present in each compartment are labeled A chemically inert conductor such as Pt is required if none of the substances participating in the half-reaction is a conducting solid Example – Fe 2+ and Fe 3+
Reaction in a galvanic cell is always an oxidation- reduction reaction that can be broken down into two half-reactions Each half-reaction has a cell potential We can obtain the overall cell potential by summing the half- cell potentials! A cell will always run spontaneously in the direction that produces a POSITIVE cell potential Each potential is measured against a standard called the STANDARD HYDROGEN ELECTRODE (SHE) SHE consists of a piece of inert Platinum that is bathed by hydrogen gas at 1 atm SHE is assigned a potential of ZERO volts
Half-reaction cell potentials are listed in a convenient table! The values in the table correspond to REDUCTION half- reactions with all solutions at 1 M, all gases at 1 atm, and 25°C (298K) for all Cu 2+ + 2 e - → Cu E 0 = -0.34 V versus SHE SO 4 2- + 4 H + + 2 e - → H 2 SO 3 + H 2 O E 0 = 0.20 V versus SHE Symbol for standard conditions!
Elements that have the MOST POSITIVE reduction potentials are easily REDUCED In general, non-metals Elements that have the LEAST POSITIVE reduction potentials are easily OXIDIZED In general, metals Table can also be used to tell the strength of various oxidizing and reducing agents Another form of the activity series Metals having LESS POSITIVE reduction potentials are MORE active and will replace metals with more positive potentials
Calculate the cell voltage for the following reaction. Draw a diagram of the galvanic cell for the reaction and label completely See previous slide for requirements of a complete diagram Fe 3+ (aq) + Cu (s) → Cu 2+ (aq) + Fe 2+ (aq)
Combining thermodynamics, electrochemistry, and not to mention a bit of physics!
In chemistry, we often refer to processes as either spontaneous or nonspontaneous A spontaneous process is said to occur if it occurs without outside intervention Has nothing to do with the speed of the reaction To explore the idea of spontaneity, consider the following physical and chemical processes: A ball rolls down a hill but never spontaneously rolls back up the hill If exposed to air and moisture, steel rusts spontaneously. However, the iron oxide in rust does not spontaneously change back to iron metal and oxygen gas Heat flow always occurs from a hot object to a cooler one. The reverse process never occurs spontaneously
∆G for a spontaneous process represents the energy that is free to do useful work ∆G for a nonspontaneous process represents the minimum amount of work that must be expended to make the process occur
So far, we have described galvanic cells under standard conditions All solutions at 1 M, all gases at 1 atm, and 25°C (298K) for all What would happen to the cell potential if the solutions were not at 1M? Can be answered qualitatively in terms of Le Ch ȃ telier’s Principle
Virtually everything you encounter, including your own bodily processes and senses, is the result of one or more equilibrium reactions Chemical equilibrium can be defined as the condition where reactant and product concentrations remain constant Occurs when a forward reaction and its reverse reaction proceed at the same rate While concentrations do not change, products and reactants continue to interconvert at equal rates
If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance
An increase in reactant concentration will favor the forward reaction and thus, increase the driving force on the electrons E will increase An increase in product concentration will oppose the forward reaction and thus, decrease the driving force on the electrons E will decrease
Because cell potentials depend on concentration, we can construct galvanic cells where both compartments contain the same components but at different concentrations An increase in concentration of reactant will increase cell potential Increase in product concentration will decrease cell potential
Potential calculated from the Nernst equation is the maximum potential before any current flow has occurred As the cell discharges and current flows from anode to cathode, the concentration will change As reactants are being converted to products, E cell will decrease Eventually, the cell potential reaches zero Zero potential means reaction is at equilibrium (a dead battery) Also, Q = K (equilibrium constant) And ∆G = 0 as well