Oxidation Numbers, ox #
are _____ or _____ numbers assigned to each _______ or
____ assuming that the _____________ are transferred
________________ from the ______ electronegative
element to the _______ electronegative element.
________________ now mimic ________ systems.
ox# are written _______ followed by _________ and are
assigned using the following:
Oxidation Number Rules
The ox# of an ______ in it’s pure form, (eg. Pb(s), O2(g) )
is ____.
The ox# of a ____________ _____ is equal to the _________ on that ______.
eg. Cl1-
The algebraic ____ of the oxidation numbers in a ________ polyatomic compound is ___.
eg. Mg3N2 Fe(NO3)3
The most common ox# for O is ___, in peroxides it is ___,
and in compounds with fluorine it is ___. The most common ox# of H is ___, in metal hydrides it is
___. The algebraic ______ of the oxidation numbers in a
polyatomic _____ is equal to the _________ on the ____.
eg. Cr2O
7
2-
Assign the ox# to the most _______________ element __ Determine the ______ charge for that element Considering the charge on the ion or neutral compound,
find the total charge for the other element, then its ox #. In compounds of _______________, the ox# of the
______ electronegative atom is ______ and the ox# of the
______ electronegative atom is _______.
eg. CO2 In a covalent bond, the more ________________ atom
takes _____ the ____________ in the bond.
Energy Transfer Theory
Chemical reactions _________ with the ______________
of __________.
Often, this energy is _____ noticeable as it doesn’t
involve a ___________ change or the production of
________ or _______.
This is true of ____________ reactions, where the
___________ switch between entities with no
___________ change in _________.
To examine the energy ____________ we _________ the
chemical reaction into ___ separate reactions called _____
____________.
For:
Net ionic equation
Half reactions
The energy _________ is the electron ______________ –
______________
Redox reactions require the ______________ reaction
_________________________ to occur simultaneously
with the _______________ reaction
_________________________
Redox – ____________________________
_________ goes __________
______ _______
Balancing REDOX Reactions 1. Assign ___________ ____________ (ox. #) to ______ element in the
rxn and __________ the species being ____________ and ____________.
2. Write separate ______ ____________ (1/2 rxn), for the
______________ (ox.) and _____________ (red.) processes. 3. Balance the ___________ that undergo _____________ or
______________. 4. Add the total _________ gained or lost in _______ ox. or red. rxn.
5. Multiply each _________ so that once combined, the total # of e-s ______ out.
6. Combine the _________. 7. Balance the net ___________ (due to ions) using _____ ions to one
side (for acidic sol’n) or with _______ ions (for basic sol’n). 8. Add _______ to balance ____ and ____. 9. Check that the final eq’n is ________________.
Predicting Redox Reactions
Redox Rxn
occur due to the ____________ of e- from ________
substance to _______________.
is a ________________ of e- between _____________.
if ______ substance ________ e- from another, a
_______________________ rxn _____________. eg. Cu + 2Ag+ ↔ Cu2+ + 2Ag
Reducing agents (RA)
a substance that __________ e- to another substance and
undergoes __________________. Oxidizing agents (OA)
a substance that __________ e- to another substance and
undergoes __________________. eg. Cu + 2Ag+ ↔ Cu2+ + 2Ag
rxns are listed as ____________, S = Strong, W = Weak.
SOA OA1 + ne- ↔ RA1 WRA
WOA OA2 + ne- ↔ RA2 SRA spontaneous rxns occur only if the ______ is
____________ and to the _________ of the ______. For spontaneous rxn:
eg. Will a mixture of aqueous chromium (II) nitrate and tin (II) nitrate react and if so what is the overall equation?
List all species ________ determine the ________, _____________ on the _______ and the ________, ________ on the _________ on the REDOX table
eg. Create ½ rxns from the following reaction data and place in order of strongest to weakest OA:
Ion/metal Zn Pb Ag I1-
Zn2+
Pb2+
Ag+
I2
eg. Create ½ rxns from the following reaction data and place in order of strongest to weakest OA:
Cu2+ + Pb Cu + Pb2+ Ag + Br2 AgBr
Ag+ + Cu Ag + Cu2+ Au + Br2 no rxn
eg. Create ½ rxns from the following reaction data and place in order of strongest to weakest OA:
W + M2 W2+ + M1- Y2+ + Z Y + Z2+
X + M2 no rxn W + Y2+ no rxn
X2+ + Z X + Z2+
Standard Reduction Potentials ( also p 805) E° (volts)
F2 + 2 e- 2 F-1 +2.87
S2O82- + 2 e- 2 SO4
2- +2.01
Co3+ + e-1 Co2+ +1.81
Pb4+ + 2 e-1 Co2+ +1.80
H2O2 + 2 H+ + 2 e- 2 H2O +1.77
Au+ + e-1 Au +1.69
PbO2 + SO42- + 4 H+ + 2 e- PbSO4 + 2 H2O +1.69
MnO41- + 8 H+ + 5 e- Mn2+ + 4 H2O +1.51
Au3+ + 3 e- Au +1.50
Ce4+ + e-1 Ce3+ +1.44
ClO41- + 8 H+ + 8 e- Cl1- + 4 H2O +1.39
Cl2 + 2 e- 2 Cl- +1.36
2 HNO2 + 4 H+ + 4 e- N2O + 3 H2O +1.30
Cr2O72- + 14 H+ + 6 e- 2 Cr3+ + 7 H2O +1.23
O2 + 4 H+ + 4 e- 2 H2O +1.23
MnO2 + 4 H+ + 2 e- Mn2+ + 2 H2O +1.22
2 IO31- + 12 H+ + 10 e- I2 + 6 H2O +1.20
Br2 + 2 e- 2 Br-1 +1.07
AuCl41- + 3 e- Au + 4 Cl- +1.00
Hg2+ + 2 e- Hg +0.85
ClO1- + H2O + 2 e- Cl1- + 2 OH-1 +0.84
Ag+ + e- Ag +0.80
NO31- + 2 H+ + e- NO2 + H2O +0.80
Hg22+ + 2 e- 2 Hg +0.79
Fe3+ + e- Fe2+ +0.77
O2 + 2 H+ + 2 e- H2O2 +0.70
MnO41- + 2 H2O + 3 e- MnO2 + 4 OH-1 +0.60
I2 + 2 e- 2 I- +0.54
Cu+ + e- Cu +0.52
O2 + 2 H2O + 4 e- 4 OH- +0.40
Cu2+ + 2 e- Cu +0.34
SO42- + 4 H+ + 2 e- SO2 + 2 H2O +0.18
SO42- + 4 H+ + 2 e- H2SO3 + H2O +0.17
Sn4+ + 2 e- Sn2+ +0.15
Cu2+ + e- Cu+ +0.15
S + 2 H+ + 2 e- H2S +0.14
AgBr + e- Ag + Br-1 +0.07
2 H+ + 2 e- H2 +0.00
Fe3+ + 3 e- Fe -0.04
Pb2+ + 2 e- Pb -0.13
Sn2+ + 2 e- Sn -0.14
AgI + e- Ag + I-1 -0.15
Ni2+ + 2 e- Ni -0.26
Co2+ + 2 e- Co -0.28
H3PO4 + 2 H+ + 2 e- H3PO3 + H2O -0.28
Tl+ + e- Tl -0.34
PbSO4 + 2 e- Pb + SO42- -0.36
Se + 2 H+ + 2 e- H2Se -0.40
Cd2+ + 2 e- Cd -0.40
Cr3+ + e- Cr2+ -0.41
Fe2+ + 2 e- Fe -0.45
S + 2 e- S2- -0.48
Ga3+ + 3 e- Ga -0.53
Ag2S + 2 e- 2 Ag + S2- -0.69
Cr3+ + 3 e- Cr -0.74
Zn2+ + 2 e- Zn -0.76
Te + 2 H+ + 2 e- H2Te -0.79
2 H2O + 2 e- 2 OH-1 + H2 -0.83
Cr2+ + 2 e- Cr -0.91
Se + 2 e- Se2- -0.92
SO42- + H2O + 2 e- SO3
2- + 2 OH-1 -0.93
Te + 2 e- Te2- -1.14
Mn2+ + 2 e- Mn -1.18
V2+ + 2 e- V -1.19
Al3+ + 3 e- Al -1.66
Ti2+ + 2 e- Ti -1.75
Mg2+ + 2 e- Mg -2.37
Ce3+ + 3 e- Ce -2.48
Na+ + e- Na -2.71
Ca2+ + 2 e- Ca -2.87
Ba2+ + 2 e- Ba -2.91
Cs2+ + 2 e- Cs -2.92
Ra2+ + 2 e- Ra -2.92
K+ + e- K -2.92
Li+ + e- Li -3.00
Electrochemical Cells
(Galvanic, Voltaic, Electric)
the ___________ reactions are occurring _____________
and are _________ by a _______
the e- _________ occurs through this ___________ circuit
the ______ difference ______________ is manifested as
__________ __________.
X+Y+
K+ NO3-
Salt Bridge
As the cell proceeds:
Cathode Anode
mass of electrode
Solution electrical charge
Ions from salt bridge
Cell Shorthand Notation eg. For the Zn/Cu cell: ____________________________ Conventions: the │ notation indicates a phase ______________ where
the _______________ and _________________ are in
physical contact. the ║ notation represents the ______ __________ or
_________ ____________ if additional _____________ are ____________ or
specific ___________ are ___________, they are written
with the _____________ separated by a ___________ or
a _______________. if the cell has no _______ for the electrode, __________
electrodes _________ or __________ are used:
___________________________________________ standard cells are __________ at _________ all ______ sets of information – REDOX equation, cell
diagram and the cell notation are related and if ___ is
provided, the other ____ should be able to be _________.
Calculating Cell Potential, E°cell for each _____ ___________ the ____________
_______________ is listed on the chart on p. 805 or the
reference sheet. the complete ______ is a combination of the
____________ and _____________ half reactions - the
_______ of the two is the _______ potential, _________ this is the usable energy given by:
______________________________________________ as all half reactions are listed as ____________ rxns, the
oxidation reaction ________ the sign, so:
______________________________________________
eg. Calculate the E° cell for the following unbalanced
equation:
Br2 + Cu 2 Br¯ + Cu+
eg. a) Determine the anode, cathode and calculate the standard cell potential produced by a galvanic cell consisting of a Sn
electrode in contact with a solution of Sn2+ ions and a Cr
electrode in contact with a solution of Cr3+ ions.
b) Write the shorthand cell notation.
eg. Use complete half-reactions and potentials to predict whether the following reaction is spontaneous or non-spontaneous in aqueous solutions. If the cell is spontaneous, write the cell shorthand notation.
O2 + 2 SO
2 + 4 OH1- 2 SO
4
2- + 2 H2O
YOU MUST READ 9.6 Corrosion p.710 – 713 before the next lesson
Reactions of Metals with Water The _______ is the _______ and will be ____________.
The ______ and ___________ in it will be the __________.
Need to consider:
i) __________ ________ - _____ of the 3
ii) ________________ _________ this requires a very long ____
_____, or an _______________, as O2 is not very water ________
iii) ________ ______________ - as with ______ ______
• Metals _______ the _____________ _______ _______ will be
________________
That’s why:
• Not ___ metals _______ with __________ ________
eg. ones that _____: ones that ______:
• Not ___ metals _______ with _____________ _______
eg. ones that _____: ones that ______:
• Not ___ metals _______ with _________ ________
eg. ones that _____: ones that ______:
Metals with Metals • Alloys are _______ to change a metal’s ____________. • As _____ _______ undergo ____________ – rust,
corrosion, patina – _______ metals are _______ to stop
the ____________. eg. ___ or ___ need protection. • The ____________ metal, called a ___________ ______,
is ______ on the table _________ to the _______ in
_________. It will __________ the ____________ and
______ the _______.
eg. Which metals will ________ Zn?
eg. Which metals will ____ ________ Zn?
eg. a) Pick a metal that will protect Zn and determine the
cell potential if the cathode is oxygenated water.
b) Do the same for the case where Zn is not protected.
Cells and Spontaneity
Reactants Products
ΔEcell Cell Type Spontaneity
Electrolytic cells, ΔE < 0
requires _____________ source of ______ to ________
the ____ __________ within the cell.
e- are __________ from the _________ and __________
to the ____________ by the ________________. rxn is then the _____________ of the ________________
rxn. used to __________________ and to produce _________
__________. Solution rxns so _____+ could be the ____________
and/or the ____________ reaction.
eg. KI solution and battery
Cathode (red): Erº (V)
Anode (ox): Erº (V)
e- from the __________ (-ve) ____________ of the
_______________ reduce the ________
e- from the _____ flow to the ___________ (+ve) __________ of the ____________ to complete the _____________.
forms ______ and _______ from an ____________ sol’n!
versus
eg. Given the following molten system, AgCl, predict the products at each electrode. Assume inert electrodes and sufficient voltage to cause a reaction to take place. Consider all possible rxns.
eg. Given a 1.00 M solution of ZnI2 at 25°C, predict the anode and
cathode half cell reactions. What is the minimum voltage required for each cell to operate?
Stoichiometry of Cell Reactions
Basic Concepts 1. q = It where:
q = _____________________
I = _____________________
t = _____________________
1C = _______________
2. Faraday’s Law: the ________ of a substance
__________ or ______________ is __________ related to the ____________ transferred
3. Faraday = ___________________________
F = ________________
4.
eg.1 Calculate the amount of charge which passes
through an electrolytic cell with a current of 1.80A for 5.0 min.
eg. 2 Calculate the mass of Zn deposited by a
current of 2.50 A operating over 30.0 min in an electrolytic cell containing Zn(NO3)2(aq).
Free Energy & Electrochemistry ΔGº = ___________ and ΔGº = __________ and… ΔGº = ______________
eg. Calculate ΔGº and Ke for the electrolytic Zn/Al
cell, Zn │ Al3+│Al.
RT
0cell
EnF Keqln