(p. 583)
So lets set the stage for all this …
So when elements join together to
form compounds electrons are
transferred. eThis transfer of electrons is what gives
elements their + or – charge.
Oxidation States
These + and – charges we have talked
about so far are known as
FORMAL CHARGES.
Oxidation States are similar to this with
assigning + and – charges, but that’s
about where the similarity ends.
The + and – charges we assign for oxidation
states ARE NOT the charge of the element
from the periodic table!
Rules For Assigning Oxidation
Numbers
Oxidation Numbers
*****Elements by themselves not combined
with anything have an oxidation number of
0!!!!
Periodic Table
Lets Try It
Doesn’t Apply
Doesn’t Apply
+1 +7 -2
HClO4
0
So HClO4 is neutral (0) overall, so we add up what we have
Exception!
Because
the
halogen
(Cl) iswill
thework
last to
one
left
and
Cl get the
whatever
oxidation
number
make
rule 5neutral
doesn’t
apply,
because
the rules are a hierarchy,
HClO4
like
its suppose
to be
meaning first come first serve, the first 4 rules cancel
So 1 + charge,
8 – charges
(-2)(4)…therefore Cl has to have a number
out number
5
+7 to make HClO4 neutral overall.
Practice
-4 +1
CH4
+7
-2
MnO4-
***** Watch out
for the overall
charge of -1 here…
Sodium Sulfate
+1
+6 -2
Na2SO4
Oxidation and Reduction
Reactions
Oxidation: When
the charge of an
element increases
during a
reaction….its
charge becomes
more positive (+).
Reduction: When
the charge of an
element decreases
during a
reaction…its charge
becomes more
negative (-)
+3-----+5
+4 ------ +2
-3 ------ -1
+3 ----- -1
-2 ------ +2
-2 ------ -5
(P. 559)
Oxidizing Agent: The element in a chemical reaction that
gets reduced, causes the other element to be oxidized.
Reducing Agent: The element in a chemical reaction that
gets oxidized, causes the other element to be reduced.
+1-2
0
-2
+1
H: +1 --- 0
Reduction
Oxidation
Reduction
Reduction
Oxidation
0
0
+2 -1
Yes
Yes
No
Yes
+1 -2
+3
+1
+5
+5
Balancing ½ Reactions In Acids or
Bases
We’ve been assigning oxidation
numbers and putting in e- to balance
them so far.
Remember, For any chemical reaction
you can do TWO ½ reactions….one
for the Oxidation (becomes more +)
and one for the Reduction (more -).
Lets Refresh Our Memories
What Ion is present in an acidic
solution? What is the acidic Ion
H+ or H3O+
What Ion is present in a basic
solution? What is the basic Ion
OH-
+6
-2
3+
Cr2O72-  Cr3+
What!? Where did these
H2O come from?
Acidic
Solution
So add 7 H2O to the
opposite side.
So add 14 H+ to this side.
7 Oxygens
7 x 2 = 14
WHAT!? Where did these H+ come from?
So the first FOUR steps to balance all ½ reactions is
the same 4 steps. SO, try it out, try using the rules
to balance these first four reactions.
There’s two of each type (acid and base). So if you
do the four steps for an all the solutions, the acid
ones are done, but there are 2 more steps to add on
to the base one, so just leave some room on these
to do 2 more steps later.
Now Balancing in Basic Solutions
So like I said the first 4 steps are the same for
acidic and basic solutions.
So in your notes there is an example of a just basic
solution balancing. (All 6 steps)
Step 5: Add OH- to both sides of the equation to
equal the H+ you added in step 3.
Step 6: H+ + OH- = H2O So you cancel out as
many as possible if H2O is on both sides you
cancel it out.
Lets Complete One We already
Started.
Step 1: Use Coefficients to balance everything but H and O.
2e-+ 2H++ Cu2O(aq) 2 Cu(s) +H2O
Step 2: Balance O by adding H2O to opposite side.
Step 3: Balance H by adding H+ to opposite side.
Step 4: Add e- to more positive side to even out the charge on both
sides.
Lets Complete One We already
Started.
/
2H2O
2OH-+2e-+2H++Cu2O(aq) 2 Cu(s)+H/
O
+2OH
2
Step 5: Add OH- to both sides…as many as there is H+
Step 6: Combine H+ and OH- to make waters, cancel out if on both sides.
Cu2O(aq)+2e-+H20 (l) 2Cu(s)+2OH-(aq)
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