CHM 51Chapter 18-Electrochemistry
18.1 – Pulling the Plug on the Power Grid
Many chemical reactions involve the transfer of electrons between atoms or ions.
Electron transfer reactions
-All single displacement and combustion reactions
-Some synthesis and decomposition reactions
-The flow of electrons is associated with electricity.
Basic research into the nature of this relationship may, in the near future, lead to cheaper, more efficient ways of generating electricity.
-Fuel cell
18.2 – Balancing Oxidation-Reduction Equations
Oxidation is the process that occurs when
-the oxidation number of an element increases,
-an element loses electrons,
-a compound adds oxygen,
-a compound loses hydrogen, or
-a half-reaction has electrons as products.
Reduction is the process that occurs when
-the oxidation number of an element decreases,
-an element gains electrons,
-a compound loses oxygen,
-a compound gains hydrogen, or
-a half-reaction has electrons as reactants.
Redox reactions are reactions in which one species is reduced and another is oxidized. Therefore the oxidation state of the species involved must change
Oxidation Number (State): A value which indicates whether an atom is neutral, electron-rich, or electron-poor.
Rules for Assigning Oxidation Numbers
- An atom in its elemental state has an oxidation number of 0.
- A monatomic ion has an oxidation number identical to its charge
- An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion
- Hydrogen can be either +1 or –
- Oxygen usually has an oxidation number of -2
- Halogens usually have an oxidation number of -1
- The sum of the oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion
Writing half reactions
•*We generally split the redox reaction into two separate half-reactions—a reaction just involving oxidation or reduction.
–The oxidation half-reaction has electrons as products.
–The reduction half-reaction has electrons as reactants.
Example: Identify which reactant is oxidized and reduced. Then write the half oxidation and reduction reactions for
2Cs(s) + F2(g) 2CsF(s)
Balancing Redox Reaction using half equation in an Acidic Solution
Assign oxidation numbers to each atom from the given unbalanced equation
2.Split the equation into half-reaction
3.Complete and balance each half reaction
◦Balance all atoms except O and H
◦Balance oxygen atoms by adding H2O to one side of the equation
◦Balance hydrogen atoms by adding H+ ions to one side of the equation
◦Balance the number of electrons being transferred
4.Combine the half-reaction to obtained the final balanced equation
◦The electrons must be cancelled
◦Simplify the equation by reducing coefficients and canceling repeated species
◦After you’re done, double check your balanced equation
Example: Balance the following net ionic equation in acidic solution:
I1-(aq) + Cr2O72-(aq) Cr3+(aq) + IO31-(aq)
Additional steps for balancing equations in basic solution
Follow all steps from the acidic solution ( 1 4)
Add the desired number of –OH react with H+ ions in the reaction; add to both side of the reaction. This is a neutralization step
◦Simplify the equation by noting that H+ combines with –OH to give H2O
◦Cancel any repeating H2O and –OH ions and reduce reaction to the lowest coefficients
Example: Balance the following net ionic equation in basic solution:
MnO41-(aq) + Br1-(aq) MnO2(s) + BrO31-(aq)
Examples:
Balance the following net-ionic equation by the half-reaction method.
◦Cu(s) + NO3-(aq) Cu2+(aq) + NO2(g) Acidic condition
◦I-(aq) + MnO4-(aq) I2(aq) + MnO2(g) Basic Condition
18.3 Voltaic (or Galvanic) Cells: Generating Electricity form Spontaneous Chemical reaction
Galvanic Cells - Electrochemical Cell Components
•Electrons flow through a conductor in response to an electrical potential difference similar to water flowing downhill in response to a difference in gravitational potential energy.
•Electric current – the amount of electric charge that passes a point in a given period of time
–Whether as electrons flowing through a wire, or ions flowing through a solution
Redox Reactions and Current
•Redox reactions involve the transfer of electrons from one substance to another.
•Therefore, redox reactions have the potential to generate an electric current.
•To use that current, we need to separate the place where oxidation is occurring from the place where reduction is occurring.
Electrochemistry is the study of redox reactions that produce or require an electric current.
•The conversion between chemical energy and electrical energy is carried out in an electrochemical cell.
•Spontaneous redox reactions take place in a voltaic cell.
–Also know as galvanic cells
•Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy.
•Oxidation and reduction half-reactions are kept separate in half-cells.
•Electron flow through a wire along with ion flow through a solution constitutes an electric circuit.
•It requires a conductive solid electrode to allow the transfer of electrons.
Through external circuit
Metal or graphite
•Requires ion exchange between the two half-cells of the system.
Consider the reaction Zn(s) + Cu2+(aq) Zn2+(aq) + Cu (s)
Two conductors (anode and cathode)
Electrolytes solution: solution that each electrode is immersed in it
External circuit: provide a pathway for electron to move from one electrode to another
Salt Bridge: provide neutrality
Shorthand Notation for Galvanic Cells or Voltaic Cell
•Shorthand description of a voltaic cell
•Electrode | electrolyte || electrolyte | electrode
•Oxidation half-cell on the left; reduction half-cell on the right
•Single | = phase barrier
–If multiple electrolytes in same phase, a comma is used rather than |
–Often use an inert electrode
•Double line || = salt bridge
Example: Give short hand notation for a galvanic cell that employs the overall reaction
Pb2+(aq) + Ni(s) Pb(s) + Ni2+(aq)
Cell involving gas
◦Additional vertical line due to presence of addition phase
◦List the gas immediately adjacent to the appropriate electrode
◦Detailed notation includes ion concentrations and gas pressure
E.gCu(s) + Cl2(g) Cu2+(aq) + 2 Cl-(aq)
Cu(s)|Cu2+(aq)||Cl2(g)|Cl-(aq)|C(s)
Example: Given the following shorthand notation, sketch out the galvanic cell
Pt(s)|Sn2+,Sn4+(aq)||Ag+(aq)|Ag(s)
18.4 Standard Electrode Potential
•The difference in potential energy between the anode and the cathode in a voltaic cell is called the cell potential.
•The cell potential depends on the relative ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode.
Electromotive Force (emf): The force or electrical potential that pushes the negatively charged electrons away from the anode ( electrode) and pulls them toward the cathode (+ electrode).
It is also called the cell potential (E) or the cell voltage.
Eocellis the standard cell potential when both products and reactants are at their standard states:
◦Solutes at 1.0 MGases at 1.0 atm
oSolids and liquids in pure formTemp = 25.0oC
Standard Reduction Potentials
•A half-reaction with a strong tendency to occur has a large positive half-cell potential.
•When two half-cells are connected, the electrons will flow so that the half-reaction with the stronger tendency will occur.
Spotaniety of the reaction can be determined by the positive Eocellvalue
E°cell = E°ox + E°red
Note: Eocell is an intensive property; the value is independent of how much substance is used in the reaction
Ag+(aq) + e- Ag(s)Eored = 0.80 V
2 Ag+(aq) + 2e- 2 Ag(s) Eored = 0.80V
Standard Reduction Potentials
Examples
Of the two standard reduction half reactions below, write the net equation and determine which would be the anode and which would be the cathode of a galvanic cell. Calculate Eocell
Fe2+(aq) + 2e- Fe(s) Eored = -0.44 V
Al3+(aq) + 3e- Al(s) Eored = -1.66 V
Use tabulated standard electrode potentials to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25 °C. (The equation is balanced.)
Without calculating Ecell, predict whether each of the following redox reactions is spontaneous. If the reaction is spontaneous as written, make a sketch of the electrochemical cell in which the reaction could occur. If the reaction is not spontaneous as written, write an equation for the spontaneous direction in which the reaction would occur and sketch the electrochemical cell in which the spontaneous reaction would occur. In your sketches, make sure to label the anode (which should be drawn on the left), the cathode, and the direction of electron flow.
- Fe(s) + Mg2+(aq) Fe2+(aq) + Mg(s)
- Fe(s) + Pb2+(aq) Fe2+(aq) + Pb(s)
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