OXIDATION NUMBERS
Review
1
When an atom becomes involved in a bond (and thus forms a compound) the atom is
either ______________________________ or _____________________________.
2.
When an atom loses electrons, it has been ____________________________________;
when an atom gains electrons, it has been ____________________________________.
3.
Do you remember burning the magnesium ribbon in our last lab? Use redox to
analyze this reaction.
3.
“The only kind of reaction in which an atom
truly loses an electron, and another truly
gains one is one that results in an
____________________ bond.”
BUT: The reaction which destroyed the Hindenburg was also a redox reaction!
2H2 + O2 → 2H2O These are molecules, not ionic compounds.
How can they have oxidation numbers if they don’t have ions?
Reminder: The charges on the ions in an ionic compound are whole numbers (1+, 2+
3+, 1-, 2-, 3-…) because electrons are
________________________________________
in an ionic compound.
In a molecular compound, there are either partial
charges (δ+ δ-) on the atoms or no charges (0)
because the atoms ______________________
_____________________________the electrons.
Concepts of Oxidation numbers
4.
The oxidation numbers in molecular compounds are not actual charges (like they are in
ionic compounds); they are assigned values to keep track of the ___________________.
They’re BOOKKEEPING!!
The oxidation numbers in molecular compounds are WHOLE
numbers, even though the actual charges are partial or zero.
PARTIAL
In a molecular bond, the atom with the greater ______________________________ is
assigned the negative oxidation number, and the other atom, which has the lesser
______________________________ is assigned a positive charge.
5.
A pure element that is all by itself always has an oxidation number of _______________
6.
The following elements always have the same oxidation number:
________________________________ always have an oxidation number of +1.
__________________________________ always have an oxidation number of +2
The oxidation number of aluminum is always _________________.
Fluorine always has an oxidation number of ____________________.
7.
The following elements often have this oxidation number, with some exceptions:
Hydrogen is usually _________________; however, when it is combined with a
metal, it is ___________. Oxygen is __________________, except in peroxide it is
________________.
8.
In binary compounds, halogens are generally _____________.
9.
In a compound, the sum of the oxidation numbers is ________________.
10.
In determining oxidation numbers in binary compounds, start with the element that is
more electronegative. This will normally be the LAST/SECOND* element in the formula.
Common Exceptions: ________________________ and __________________________
* Screencast says first element; it’s wrong!
Practice in Determining Oxidation Numbers
11.
Determine the oxidation numbers for the elements in each of the substances below.
A)
Phosphorous
B)
PCl3
C)
PCl5
D)
CO
E)
CO2
F)
H2O
G)
NH3
H)
LiNO3
I)
KMnO4
J)
SO42-
K)
PO31-
L)
H2Cr2O7
M)
NH4ClO4
N)
HClO3
O)
Ca(ClO2)2
P)
Bleach – (NaOCl)
Q)
Cl2
R)
AsCl3
S)
acetylene
(C2H2)
T)
hydrogen peroxide
(H2O2)
U)
sodium hydride
(NaH)
BALANCING REDOX EQUATIONS
Oxidation numbers can be used to balance difficult redox equations. For example: the space
shuttle is lifted off the earth by this tough redox reaction:
NH4ClO4 + Al → Al2O3 + HCl + N2 + H2O
This would take forever to balance, based on the
method we have used to date (which is called
_______________________________________)
However, the Law of Conservation of Mass says
that “matter is neither created nor destroyed in a
chemical reaction – just rearranged.” Electrons
are matter, so – however many electrons one
substance loses is the same number that another
substance must gain. This fact can be used to
balance tough redox reactions. This method is
called the ___________________________________________________ method.
Example:
__HNO3 + __H2S → __S + __NO + __H2O
Practice
13.
__H2Cr2O7 + __HI + __HCl → __CrCl3 + __I2 + __H2O
14.
__KCl + __KMnO4 + __H2O → __Cl2 + __MnO2 + __KOH
______________________________________________________
15.
__NH3 + __NO2 → __N2 + __H2O
Electrochemistry
16.
Electrochemistry is the study of the relationship between _________________________
________________________________ and ____________________________________
Put more precisely: electrochemistry is the study of the redox processes by which
______________________________ energy is converted to _______________________
energy and vice versa.
“Electrochemical processes are useful in ______________________, and critical in
_________________________________ functioning.” An example of the importance of
electrochemistry in our bodies: Every cell has a voltage across its plasma membrane.
(Voltage is electric potential energy that arises from the separation of opposite
charges.) This is particularly important in nerve cells, because it is changes in this
potential that allows electrical signals to flow along our nerve cells.
17.
A battery is a type of electrochemical cell, also known as a
___________________________________ or ______________________________ cell.
This type of cell involves a spontaneous chemical reaction, in which chemical energy is
converted to electrical energy. “A spontaneous process is a chemical or physical change
that occurs with _________________________________________________________.”
Example Copper wire is put into a solution of silver nitrate.
18.
We can tell that copper metal + silver nitrate is a
spontaneous reaction, and that the opposite reaction
(silver metal + copper nitrate) is non-spontaneous
because we actually tried these reactions. How can
we tell if these reactions are spontaneous or
non-spontaneous without physically doing them?
There are 2 different methods we can use:
A) Non-quantitative method: the_________________________________(see appendix)
B) Quantitative method for determining if redox reactions are spontaneous or not, and
also for calculating the E0 for a reaction/electrochemical cell, (e.g., a battery)
This method uses Standard Reduction Potentials. A reduction potential is a quantitative
measurement for the tendency of a substance to be reduced – i.e., how much it “wants”
to ________________ electrons. Since reduction can only occur when there is also
oxidation, the reduction potential of a substance must be measured versus another
substance. The standard reference potential is versus ___________________________.
2H1+ + 2e1- → H20
E0 = 0.000 volts
Standard conditions ( 0 ) are 1 molar for ____________, and 1 atmosphere for ________
The sum of the potentials of the two half reactions equals the potential of the
electrochemical cell, E0cell .
E0oxidation + E0reduction = E0cell
If E0cell > 0, the reaction in the electrochemical cell is _____________________.
If E0cell < 0, the reaction in the electrochemical cell is _____________________.
If E0cell = 0, the reaction in the electrochemical cell is at equilibrium.
Using the table on the next page, calculate the E0cell and state if the reaction is
spontaneous or not. NOTE: For the half reaction that is an oxidation, the potential will
have the opposite sign from the reduction potential. Also, it doesn’t matter if a half
reaction is multiplied by a coefficient; don’t multiply E0ox or E0red by the coefficient.
Example 1
E0ox
Sn + Cu2+ → Sn2+ +
E0oxidation + E0reduction = E0cell
=
Example 2
2Fe + 3Ca(NO3)2 → 2Fe(NO3)3 + 3Ca
E0ox
E0oxidation + E0reduction = E0cell
E0red
=
Practice in calculating E0cell and determining if cell is spontaneous or not
A) Mg + Pb2+ → Pb + Mg2+
B) 2Al + 3CoCl2 → 3Co + 2AlCl3
E0red
The structure of an electrochemical cell
19.
When we dipped Cu into AgNO3, the reaction was spontaneous; however, it didn’t
produce electricity for us to use, like a battery does. “Electricity involves the flow of
_____________________. If we want to derive electrical energy from this reaction, we
must force the electrons that leave the ________________________ to flow through a
wire in order to reach the __________________________. In an electrochemical cell,
the half-reactions are performed in separate ________________________. An
______________________ ( a strip of metal that participates in the reaction) is placed in
each half reaction vessel. The electrode in the oxidation vessel is called the
__________________; the electrode in the reduction vessel is called the __________________
How can you remember this?
Below is a picture of an electrochemical cell. The
voltage the cell puts out is theoretically
_____________; however, because of the resistance
in the ____________ It is actually less. The function
of the salt bridge is to prevent ___________________
from building up in either vessel. Over time, the electrodes will change in weight.
The __________________ will get heavier, and the ____________________ will get lighter.
The voltage will decrease over time; when E0cell becomes zero, the redox reaction is at
_______________________, and the battery is dead.
Electrolysis
It is possible for a non-spontaneous reaction to occur; however you must __________________
to happen. The method used to force a non-spontaneous redox reaction to happen is called
_________________________. In this, you put ___________________________ energy into
the non-spontaneous reaction, which forces it to occur; thus, you are converting
_________________________ energy into
_______________________ energy. For
example, in industry, putting electricity through
molten NaCl produces sodium metal
and chlorine gas. _____________________________
is another example; here one metal is “plated” onto
another. Everyday examples of this are silverware that
is silver-plated and chrome-plated car bumpers.
APPENDIX
The activity series is a non-quantitative
method for determining if a redox
reaction is spontaneous or nonspontaneous. Whichever metal is higher
on the activity series “wants” to be
oxidized more than the one that is lower.
Whichever halogen is higher on the
activity series “wants” to be reduced more
than the one that is lower.
Fluorine
Chlorine
Bromine
Iodine
Examples:
Al + FeCl3 →
Al0 + Fe3+ → spontaneous
Reactivity
Series for
Halogens
Fe + AlCl3 →
Fe0 + Al3+ → non-spontaneous
Sn + HCl →
Sn0 + H1+ → spontaneous
Ag + HCl →
Ag0 + H1+ → non-spontaneous
F2 + LiCl →
F0 + Cl1- → spontaneous
Cl2 + LiF →
Cl0 + F1- → non- spontaneous
Practice using the Activity Series
Determine if the following reactions will happen spontaneously or not.
A) Cu + ZnCl2 →
____________________ B) Zn + CuCl2 → ___________________
C) Ca + Pb(NO3) → ____________________ D) Pb + Ca(NO3)2 → __________________
E) Br2 + NaCl →
____________________ F) Cl2 + NaBr → ___________________
G) LiBr + F2 →
____________________ H) Br2 + LiF →
___________________
Extra example of calculating E0 and spontaneity
2Mn2+ + 8H2O + 10Hg2+
E0ox
Mn2+ + 4 H2O
E0red Hg2+ → Hg22+
→
2MnO41- + 16H1+ + 5Hg22+
→
MnO41- + 8H1+-
E0 = (-1.507)
E0 = (+0.7973)
E0oxidation + E0reduction = E0cell = (-1.507) + (+0.7973) = (-0.587 V) non-spontaneous
Practice
2SO42-
+
Co2+
→
Co
+
S2O82--