CHEM10006 Study Guide - Final Guide: Standard Hydrogen Electrode, Nernst Equation, Standard Cell
12 Jun 2018
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Department
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Oxidation: addition of oxygen, loss of hydrogen, oxidation state
increases
Reduction: loss of oxygen, addition of hydrogen, oxidation state
decreases
REDCAT: Reduction at cathode, oxidation at anode
Electrons flow to negative electrode to positive
Salt bridge: Allows ion movement between half cells by
preventing build up of charge in the half cells that would stop
the reaction
Oxidation numbers: Method of keeping track of electron
movement in redox reactions, and identifying which species
under oxidation and reduction
1. Element: 0
2. Monatomic atom: same as charge
3. Covalent compounds with nonmetals: H is +1
4. Oxygen: -2 except in peroxides
5. Binary compounds: element with greater
electronegativity assigned a negative oxidation state
equal to the charge in it’s ionic compounds e.g. PF
5
,
Fluorine is -1
6. Sum of oxidation states = 0 for a neutral molecule
and charge for an ionic species
Concentration cells: Contain same element in each half cell
but of different concentrations; non standard concentrations
will even un until equilibrium, electrons flow from cell of lower
concentration to higher
equation
Acidic conditions: Balance excluding H and O then add H
2
O &
H
+
to balance, multiply half equations to make electrons equal,
combine, simplify
Basic conditions: Balance as if H
+
ions were present, add OH
-
to convert them to H
2
O, balance, cancel etc.
Galvanic cell: Voltaic cell, electrochemical cell that derives
electrical energy from spontaneous redox reactions taking
place within the cell
Daniell cell: Zinc-copper voltaic cell, Zn dissolves and Cu
2+
decreases, current flows because the copper ion has a greater
tendency to be reduced than the zinc ion
Standard ½ cell reduction potentials: Standard half cell
combined with standard hydrogen electrode, predicts cell
potential by knowing half cell reduction potentials
Electrochemical series: Ranking of half cells in order of their
reduction potential; a higher member can oxidise the reduced
form of a lower member (e.g. fluorine wants to be reduced so
has a high standard voltage of 2.87V); top left reacts with
bottom right
Standard ½ cell reduction potentials are dependent on the pH:
Nernst equation used to calculate potential at intermediate
pH’s
Stability field of water: Range of values of potential and pH for
which water is thermodynamically stable towards both
oxidation & reduction, stability important in understanding
chemistry occurring in natural waters
●Species with a potential more positive than the top
half reactions
More negative potential will be oxidized
Method 2: Write the reduction & oxidation half equations and
add
When given the standard half cell reduction potential, if you
reverse it to represent the species being oxidized, reverse the
sign of the potential
Nernst equation: Used to calculate cell potential outside of
standard conditions, there is a simplified version for 298k and
a different version used in biological contexts
Remember when dividing the concentration values of Q or K to
raise them to the power of the coefficient in the balanced
equation
At equilibrium the cell potential is 0, set Nernst equation to 0
to solve for equilibrium constant
Transition metals: D block elements in groups 3-12 and
periods 4-6, important biological roles, variable oxidation
states, incompletely filled D orbitals (except for Zinc),
characteristic colours due to electrons jumping between
orbitals, absorbing light, have unpaired electrons responsible
for their magnetic properties
Document Summary
Oxidation: addition of oxygen, loss of hydrogen, oxidation state increases. Reduction: loss of oxygen, addition of hydrogen, oxidation state decreases. Salt bridge: allows ion movement between half cells by preventing build up of charge in the half cells that would stop the reaction. Oxidation numbers: method of keeping track of electron movement in redox reactions, and identifying which species under oxidation and reduction. 1: monatomic atom: same as charge, covalent compounds with nonmetals: h is +1. Binary compounds: element with greater electronegativity assigned a negative oxidation state equal to the charge in it"s ionic compounds e. g. pf , Sum of oxidation states = 0 for a neutral molecule and charge for an ionic species. Concentration cells: contain same element in each half cell but of di(cid:1559)erent concentrations; non standard concentrations will even un until equilibrium, electrons (cid:712)ow from cell of lower concentration to higher equation. Acidic conditions: balance excluding h and o then add h 2 o &
