Chapter 17 Electrochemistry Flashcards
a redox reaction
One that entails changes in oxidation number (or oxidation state) for one or more of the elements involved.
the oxidation number of an neutral atom in an element is
+ex
Zero
ex: Cl2
Oxidation number
*ex with +/- on
The total number of electrons that an atom either gains or loses in order to form a chemical bond with another atom. The iron ion Fe3+, for example, has an oxidation number of +3 because it can acquire three electrons to form a chemical bond, while the oxygen ion O2− has an oxidation number of −2 because it can donate two electrons.
In an electronically neutral substance, the sum of the oxidation numbers is
0
oxygen ON is
-2 but -1 in peroxide (ANY PROXIDE)
Consequential to these rules, the sum of oxidation numbers for all atoms in a molecule is equal to
the charge on the molecule
What are the oxidation numbers?
the oxidation number for nitrogen is +5 and that for oxygen is −2, summing to equal the 1− charge on the molecule:What a
Oxidation
Lose electron
Reduction
Gain in electron
balancing equations for aqueous redox reactions steps:
- Write skeletal equations for the oxidation and reduction half-reactions.
- Balance each half-reaction for all elements except H and O
- Balance each half-reaction for O by adding H2O.
- Balance each half-reaction for H by adding H+.
- Balance each half-reaction for charge by adding electrons.
- If necessary, multiply one or both half-reactions so that the number of electrons consumed in one is equal to the number produced in the other.
- Add the two half-reactions and simplify.
- If the reaction takes place in a basic medium, add OH− ions the equation obtained in step 7 to neutralize the H+ ions (add in equal numbers to both sides of the equation) and simplify
galvanic cells (or voltaic cells) (5)
+ what is included in the setup of the cell
- those in which a spontaneous redox reaction takes place
- comprised of two half-cells, each containing the redox conjugate pair (one solid one aq ex: solid cu/cu2+aq and solid Ag/Ag2+)
- An external circuit is connected to each half-cell at its solid foil, meaning the Cu and Ag foil each function as an electrode.
- An external circuit is connected to each half-cell at its solid foil, meaning the Cu and Ag foil each function as an electrode.
- To keep the reactants separate while maintaining charge-balance, the two half-cell solutions are connected by a tube filled with inert electrolyte solution called a salt bridge
anode
What occurs here and which process?
the electrode at which oxidation occurs
positive end, where electrons leave
cathode
The electrode where reduction occurs
negative terminal or electrode through which electrons enter
The redox reactions in a galvanic cell occur only at
the interface between each half-cell’s reaction mixture and its electrode.
In the galvanic cell, positive charge build up at
the annode
In the galvanic cell, negative charge build up at
the cathode
Why a salt bridge is needed
Without the salt bridge, the solution in the anode compartment would become positively charged and the solution in the cathode compartment would become negatively charged,because of the charge imbalance. The attractive and repulsive forces will prohibit the flow of electrons within the cell.
Difference between Galvanic and electrolytic cells
Galvanic cells derives its energy from spontaneous redox reactions, while electrolytic cells involve non-spontaneous reactions and thus require an external electron source like a DC battery or an AC power source.
cell notations (3)
- The relevant components of each half-cell are represented by their chemical formulas or element symbols
- All interfaces between component phases are represented by vertical parallel lines; if two or more components are present in the same phase, their formulas are separated by commas
- By convention, the schematic begins with the anode and proceeds left-to-right identifying phases and
interfaces encountered within the cell, ending with the cathode
inert electrode (2)
Why we use this+how does it work
- This is required when neither member of the half-cell’s redox couple can reasonably function as an electrode, which must be electrically conductive and in a phase separate from the half-cell solution.
- simply provide or accept electrons to redox species in solution
active electrodes
Electrodes constructed from a member of the redox couple
When measured for purposes of electrochemistry, a potential reflects the
driving force for a specific type of charge transfer process, namely, the transfer of electrons between redox reactants.
cell potentials, Ecell (2)
What it is+formula
the difference in potential between two half-cells
standard cell potential, E°cell
a cell potential measured when both half-cells are under standard-state conditions (1 M concentrations, 1 bar pressures, 298 K)
standard hydrogen electrode (SHE) (3)
What it its signifigance/what is its conditions and compoents/formula
One particular half-cell that serve as a universal reference for cell potential measurements, assigning it a potential of exactly 0 V.
- A typical SHE contains an inert platinum electrode immersed in precisely 1 M aqueous H+ and a stream of bubbling H2 gas at 1 bar pressure, all maintained at a temperature of 298 K
standard electrode potential, E°X (EX) for a half-cell X is defined as
the potential measured for a cell comprised of X acting as cathode and the SHE acting as anode all under standard conditions
Good oxidizing agents have —- standard reduction potentials
high
good reducing agents have — standard reduction potentials.
low
In a reduction standard potential table, the highest values (most negative) represent the—–. This means that the more negative values are more likely to be the —-
Standard reduction potential measures the tendency for a given chemical species to be reduced.
stronger reducing agents.
anode
an increased value of E° corresponds to an increased driving force behind
the reduction of the species (hence increased effectiveness of its action as an oxidizing agent on some other species)
Negative values for electrode potentials are simply a consequence of assigning a value of 0 V to the SHE, indicating the reactant of the half-reaction is
a weaker oxidant than aqueous hydrogen ions.
E°cell is positive when E°cathode — E°anode, and so any redox reaction in which the oxidant’s entry is above the reductant’s entry is predicted to be —–.
> spontaneous
The spontaneous oxidation of copper by lead(II) ions is not observed,and so the reverse reaction, the oxidation of lead by copper(II) ions, is predicted to occur —
spontaneously
reversing the direction of a redox reaction effectively interchanges the — and the cell potential is now calculated…..
identities of the cathode and anode half-reactions, and so the cell potential is calculated from electrode potentials in the reverse subtraction order than that for the forward reaction
In the case of a redox reaction taking place within a galvanic cell under standard state conditions, essentially all the work is associated
transferring the electrons from reductant-to-oxidan
The work associated with transferring electrons is determined by
where n is the number of moles of electrons transferred, F is Faraday’s constant, and E°cell is the standard cell potential.
E° and K relating formula (2)
+what does this equation tell us?
This equation indicates redox reactions with large (positive) standard cell potentials will proceed far towards completion, reaching equilibrium when the majority of reactant has been converted to product
Nernst equation
What is it and when do we use it?
This equation describes how the potential of a redox system (such as a galvanic cell) varies from its standard state value
A convenient form of the Nernst equation for most work is one in which values for the fundamental constants (R and F) and standard temperature (298) K), along with a factor converting from natural to base-10 logarithms:
