redox and electrochemistry Flashcards
oxidation
lose e-
reduction
gain e-
redox reactions
e- are transfrered and one molecule gains e- while one loses e-
- oxidation state of compounds changes
oxidation state
measured using oxidation numbers
- keep track of how e- are shared in a molecule
- model of electron distribution where e- is given to most eneg atom
oxidation state rule for pure atoms
pure elements in natural form like diatomics have an oxidation state of 0
ex: o2
what is oxidation state of monoatomic ion
its charge
ex: mg is 2+
what is true about the sum of oxidation states among atoms in a molecule
they add up to the overall charge of the molecule
oxidation state of F
-1, oxidizing agent
oxidation state of most halonges
-1
what is oxidation state of cholrine in HClO4
+7
oxidation state of hydrogen
usualyl 1+ unless bonded to a more electropositive atom like NAH and LIAlH
oxidation state of oxygen
usually -2 unless peroxide, then its -1
oxidation state of alkali metals
1+
oxidation state of alkali earth metals
2+
nonredox reactions
overall, the distribution of e- does not change
acid base, subsitution reactions, precipitation reactions , double displacement reactions
nonredox reactions
overall, the distribution of e- does not change
acid base, subsitution reactions, precipitation reactions , double displacement reactions
nonredox reactions
overall, the distribution of e- does not change
acid base, subsitution reactions, precipitation reactions , double displacement reactions
nonredox reactions
overall, the distribution of e- does not change
acid base, subsitution reactions, precipitation reactions , double displacement reactions
classic redox
single displacement
is combustion a type of redox
yes
something + 02 –> co2 + h2o
are combination reactions redox
yes
free elements –> compound
how to balance redox reactions
- split into half reactions
- balence non O and non H atoms
- balence ox
- balence H
- balence e-
- multiply so both half rxn has same amount of e-
- add/ cancel like terms
how to balance redox reactions
- split into half reactions
- balence non O and non H atoms
- balence ox
- balence H
- balence e-
- multiply so both half rxn has same amount of e-
- add/ cancel like terms
how to balance redox reactions
- split into half reactions
- balence non O and non H atoms
- balence ox
- balence H
- balence e-
- multiply so both half rxn has same amount of e-
- add/ cancel like terms
oxidation agent
is reduced and promotes oxidation of other reagnet
typically contain oxygen or other electronegative atoms
reducing agent
is oxidized and promotes reduction of other reagent
tend to have hydrogen
PCC
common oxidizing agent (weak)
NaBH4
common reducing agent
LiAlH4
common reducing agent
reduction potential
how likely something is to be reduced, higher value = more likely to be reduced
oxidation potential
reduction potential = - oxidation potential
standard potential of a cell
measured as reduction potential
what do all electrochemical cells have
a cathode and an anode
anode
electrode where oxidation happens
-surplus of e- generated that travel to the cathode
cathode
Electrode where reduction happens
galvanic cell
spontaneous redox reaction that generates a positive V difference
cell potential eqn
E cell = Ecathode - E anode
which electrode possesses the reduction potential in a galvanic cell
cathode (spontaneous rxn )
cell potential eqn for a galvanic cell
E cell = E cathode - E anode
another way of establishing the cell potential formula
E cell = reduction potential of cathode + oxidation potential of anode
Nerst Eqn
how electrical potential of a cell is effected by temperature and concentration of reactants
E’cell = E cell - RT/zf * lnQ
E’ cell
actual electrical potential of cell under given conditions
z in nertz eqn
moles of e- transfered
F in nertz eqn
faraday constant
q
reaction quotient
[products]/[reactants]
physiology nertz eqn
E’cell = standard potential - .05916/ z(logq)
Daniell cell
typical set up of a galvanic cell
- half reactions are carried out in two separated half cells
- electrodes are connected by a wire and a salt bridge (does not interfere with redox rxns
-
purpose of salt bridge
to prevent strong charge gradient buildup and hindering progress of rxn
Daniell cell
typical set up of a galvanic cell
- half reactions are carried out in two separated half cells
- electrodes are connected by a wire and a salt bridge (does not interfere with redox rxns
-
what happens to the zinc electrode (anode)
it becomes smaller as oxidation continues and e- are being stripped
what happens to the sulfur electrode (cathode)
will grow because of increased e-
galvanic cell notation
anode is on right || cathode is on left
charge of galvanic anode
- because it is the source of electrons
charge of galvanic cathode
positive because it has greater potential to be reduced
concentration cells
galvanice cells where the two half reations occur in the same cell
1. electrodes must be made out of same material
2. must be a concentration difference between the two regions of the cell
biological example of concentration cells
cell membrane with concentration difference between inside and outside of cell
electrolytic cells
apply energy to system and produce reaction in nonspontaneous direction
break down into components
which electrode is location of reduction in electrolytic cells
cathode, marked with a negative charge
anode of electrolytic cell
positive and where oxidation takes palce
portion of battery that provides charge
galvanic
portion of battery that allows for recharge
electrolytic
benedicts reagent
reduce cu2+ to cu , dark blue color change
connection between standard gibbs free energy and electric potential
ΔG= -nFEcell
n = moles of electrons moved
f is faradays constant
how does standard potential relate to Keq
nFEcell = RTlnkeq
if the standard redox potential of a cell is positive what happens to gibbes free enrgy
negative so spontaneous
trade off in galvanic cell
using spontaneous chemical reaction to provide electrical energy
half reactions of galvanic cell
zinc and copper
Zu –> Z 2+ + 2 e-
Cu 2+ + 2e- –> cu
half cell
seperated containers with metal electrodes
either cathode or anode
electrode on left of galvanic cell
zn
electrode on right of galvanic cell
cu
what allows e- to flow from zinc to copper in galvanic cell
the current carrying wire
zn electrode in galvanic cell
anode
negative
oxidation
cu electrode in galvanic cell
cathode
positive (e- move toward positive)
reduction
galvanic electrodes mnuemonic
An Ox and Red Cat
- anode, oxidation
- cathode, reduciton
electrolysis
uses external energy to power nonspontaneous (electrolytic ) reaction
- often used to split molecules –> make pure metals
requires voltage source
half reactions of electrolytic cells
2 na+ + 2 e- –> 2 Na (s) red
2 cl- –> cl2 + 2 e- (g) ox
does electrolytic reaction happen in 2 containers (like galvanic) or one?
only one, and there is no aqueous solution. only the ions na + and Cl-
anode of electrolytic cell
on left
is positive
still oxidation
2 cl - –> cl2 + 2 e-
ELECTRONS FLOW FROM ANODE TO BATTERY
cathode of electrolytic cell
Na+ +2 e- –> na
is negative
reduction
ELECTRONS ARE FLOWING FROM BATTERY TO CATHODE
what side of the battery is the electrolytic cathode attatched to
negative
