Transcript Redox Review 9.1

Warm-up 10/28
1.
2.
Why are hybrid orbitals important in
explaining molecular geometry?
State the hybridization of the central
atom in the following:
a.
b.
NH2CH2O
Pg. 356 #11
 Open-note
“quiz”
 Turn in on separate piece of paper
 When
finished:
a) turn quiz in on front table
b)start working on topic 9.1 review ws-have
the front side completed by tomorrow
Redox
Review
Key Review Questions
 What
are “redox” reactions?
 How
do you determine which element
has been oxidized and which reduced in
a redox reaction?
 How
do you determine the oxidizing
agent? Reducing agent?
“Redox Reaction”
 Type
of chemical reaction in which electrons are
transferred from one substance to another.

Oxidation
 Loss
of one or more electrons from a substance
 Ex: Fe2+ (aq)  Fe3+ (aq) + e
Reduction
 Gain
of one or more electrons by a substance
 Ex: 2 H+ (aq) + 2 e-  H2 (g)
Redox Reactions
 Oxidation
and reduction go together-- Whenever a
substance loses electrons and another substance
gains electrons
 Oxidation
Numbers are a system that we can use to
keep track of electron transfers
Oxidation Numbers Rules
Oxidation numbers always refer to
single atoms
The oxidation number of an
uncombined element is always 0
O2, H2, Ne
Zn
The oxidation number of Hydrogen
is usually +1 Hydrides are an
exception They are -1
HCl, H2SO4
The oxidation number of Oxygen is
usually -2 Peroxides are an exception
They are –1 In OF2 oxygen is a +2
H2O, NO2, et
Oxidation numbers of monatomic
ions follow the charge of the ion
O2-, Zn2+
The sum of oxidation numbers is
zero for a neutral compound. It is
the charge on a polyatomic ion
LiMnO4
SO42-
Warm-up 11/2
 Nitric
acid reacts with silver in a redox
reaction.
__ Ag(s) + __ NO3–(aq) + ____ → __ Ag+(aq) + __ NO(g) + ____
 Using
oxidation numbers, deduce the
complete balanced equation for the
reaction showing all the reactants and
products.
 (Total 3 marks)
Answer
 change
in oxidation numbers: Ag from 0
to +1 and N from +5 to +2;
Do not penalize missing charges on
numbers.
 balanced equation: 3Ag + NO3– + 4H+ →
3Ag+ + NO + 2H2O
 Award [1] for correct reactants and
product;
Award [3] for correct balanced equation.
Ignore state symbols. 3
This Week
 Lab
 9.1

Due Wednesday
Quiz Thursday
Review your notes regarding activity series,
redox titrations and Winkler method
(pg.218-225 in book)
Nomenclature (Naming)
 Ionic


Compounds
Resulting formula unit must be neutral
Binary compounds
 Metal,

Ex: Calcium chloride
 Writing

nonmetal-ide
formulas:
Determine the oxidation number (charge) of the
ions and make sure there are enough of each to
make a neutral compound
Metals with multiple ions
 Transition
metals can form more than one
type of ion (i.e. lose different amounts of
electrons)

Cu1+, Cu 2+
 We
use roman numerals to indicate the
charge

Cu1+ = “Copper I” ; Cu 2+ = “Copper II”
 Exceptions:
Ag1+, Zn2+, Cd2+, Al3+, Sc3+
Naming metals with multiple
ions
•
•
We can determine their charge by the ionic
formula, and the charge on the nonmetal
Ex. Copper (II) oxide (“copper two oxide”)
–
CuO
–
•
oxygen has a -2 charge, so it would only take
one Cu2+ to bond with Oxygen.
Ex. Copper (I) oxide
–
Copper 1+ -we would need two of these to
react with Oxygen so the formula would be:
•
Cu2O
Examples
 Lead
(II) hydroxide  Formula?
 Pb(OH)2
 Cadmium
 Cd(NO3)2
nitrate
10/30
 Please
write any materials you still need
for your IA
Balancing Redox Reactions
 Many
Redox rxns are complex and difficult to
balance .
 A systematic
approach to balancing these
reaction is required.
Balancing Redox Equations
1.
2.
3.
4.
5.
Divide the equation into 2 half reactions—one for
oxidation, one for reduction.
Balance each half reaction
1. Balance elements other than H and O
2. Balance O by adding H2O as needed
3. Balance H by adding H+ as needed.(acidic solution)
4. Balance charge by adding e- as needed.
Multiply half reactions by integers so that the # of e- lost in
one reaction = # of e- gained in the other reaction.
Add the two half reactions. Simplify by canceling species
that appear on both sides of the arrow.
Check your work. Make sure that both the atoms and
charges balance
Balancing Redox Equations 1
MnO41- (aq) + C2O42- (aq)  Mn2+(aq) + CO2 (g)


MnO41- = oxidizing agent (it is doing the oxidizing to the other substance, but
itself is being reduced).
C2O42- = reducing agent (it is doing the reducing to the other substance, but
itself is being oxidized).
Balancing Redox Equations 2
Cr2O72- (aq) + Cl1- (aq)  Cr 3+ (aq) + Cl2 (g)


Cr2O72- = oxidizing agent (it is doing the oxidizing to the other
substance, but itself is being reduced).
Cl- = reducing agent (it is doing the reducing to the other
substance, but itself is being oxidized).
Balancing Redox Equations 3
Cu (s) + NO31- (aq)  Cu 2+ (aq) + NO2 (aq)
Balancing Redox Equations 4
Mn 2+ (aq) + NaBiO3 (s)  Bi 3+ (aq) + MnO4 1- (aq) + Na 1+ (aq)
Warm-up (Super Fast!)
 Determine
the oxidation number of Mn in
the following:

MnSO4