Have you ever drank from an aluminum can? Ever used a
flashlight? Use your calculator on a test recently? Enjoy
exercising? Are you alive?
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What do all of these things have in common?
Slide 5
ENERGY!!! And all forms of energy harnessing require an
understanding of Redox Reactions.
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Redox Notes Part I : Define oxidation and reduction.
Slide 7
Oxidation-Reduction Reactions Aka Redox Rxns Chemical reactions
that transfer electrons from one chemical to another at the same
time OXIDATION the LOSS of electrons in a chemical rxn REDUCTION
the GAIN of electrons in a chemical rxn
Slide 8
Redox is a tandem process Oxidation cannot happen without
reduction & vice versa Thus, you will have two types of agents
(chemicals) OXIDIZING AGENT substance that oxidizes another. It
gets reduced. REDUCING AGENT - substance that reduces another. It
gets oxidized.
Slide 9
Why can oxidation not happen without reduction? One substance
that donates the electrons needs a place for the electrons to
travel to. Electrons dont just vanish; they attach to another atom.
Therefore, one substance donates and the other accepts the
electron.
Slide 10
LEO the Lion and his Little es Meet Leos family. He loves them
very much.
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LEO goes hunting and loses his little es.
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LEO goes GER Listen to him ROAR!!
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LEOs pride searches everywhere
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And finally finds them playing at the watering hole
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LEO gains his little es back again.
Slide 16
LEO the lion goes GER L oss of E lectrons = O xidation G ain of
E lectrons = R eduction
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Example 1: Identify the movement of the electron(s) and then
label the reaction as oxidation or reduction: Al Al +3 + 3 e -
Slide 18
Example 2: Identify the movement of the electron(s) (If + -
gained e -. If - +lost e - ) and then label the reaction as
oxidation or reduction: 2Br - + Cl 2 Br 2 + 2Cl -
Slide 19
Atoms of elements, ions or compounds gain or lose outermost
electrons during reactions that form a new set of elements, ions or
compounds. REDOX
Slide 20
SC.912.P.8.10: Describe oxidation- reduction reactions in
living and non-living systems.
Slide 21
A transfer of electrons occurs during all single replacement or
Combustion reactions Sometimes in Double Replacement and
Decomposition reactions REDOX Reactions Energy
Slide 22
LEO the lion says GER Loss of Electrons is Oxidation _____ Gain
of Electrons is Reduction OIL RIG Oxidation Is Loss of electrons
causing the oxidation number to pump up to a higher value.
Reduction Is Gain of electrons causing a reduction in the value of
the oxidation number. Review definition of oxidation and
reduction.
Slide 23
Photosynthesis and cellular respiration are biological examples
of Redox reactions. Write the chemical equation for these
reactions. Living Systems
Slide 24
Fires, rusting, and metals reacting in acid are also examples
of RedOx reactions Lets see why these reactions are classified as
REDOX. Non-Living Systems
Slide 25
When unbonded elements react to form compounds, one of the
elements gains electrons (would that be the metal or the nonmetal?)
while the other loses electrons (would that be the metal or the
nonmetal?). Think about the formation of NaCl. 1. Write your
thoughts in your notebook. 2. Share your ideas with your partner.
3. Share with the class. Think-Pair-Share
Slide 26
Metals undergo Oxidation: loss of e- Ex Na Na + +e - Reduction:
gain of e- Ex Cl 2 + 2e - 2Cl - The Oxidizing Agent-substance
reduced: Cl 2 The Reducing Agent-substance oxidized: Na REDOX
Reaction Example
Slide 27
1) __ Mg + __ N 2 __ Mg 3 N 2 __ 2) __ Fe + __ O 2 __ Fe 2 O 3
__ (rust) 3) __ Ca + __ O 2 __ CaO __ 4) __ H 2 + __ O 2 __ H 2 O
__ Identify the REDOX reactions.
Slide 28
5) __CH 4 __+__ O 2 __CO 2 __ + __H 2 O__ 6) __ Cl 2 + __ KI __
KCl __ + __ I 2 7) __ CaCO 3 __ + __ HCl __ __ CaCl 2 __ + __ H 2
CO 3 __ __ __ Identify the REDOX reactions.
Slide 29
8) __ CH 4 __ + __ O 2 __ CO 2 __ + __ H 2 O __ 9) __ AgNO 3 __
+ __ Cu __ Cu(NO 3 ) 2 __ + __ Ag __ __ 10) __ C 6 H 12 O 6 __ + __
O 2 __ CO 2 __ + __ H 2 O __ __ __ Identify the REDOX
reactions.
Slide 30
Single Replacement Reactions are always redox reactions!
Combustion reactions are always redox reactions! Any time an
oxidation number changes (which means electrons are gained or lost)
during the reaction, a redox reaction is occurring. END OF
CHEMISTRY 1 PPT
Slide 31
You must assign oxidation numbers to all elements in the
reaction. Identify which elements oxidation number changed from
reactants to products. Some electron transfers are not as easy to
predict as metal with nonmetal. To figure out what is oxidized and
what is reduced you can follow a plan. ( Honors )
Slide 32
Oxidation number = number of e- gained or lost by an atom when
it forms an ion K + Br - EX: 2K + Br 2 2KBr Potassium is oxidzed
from 0 to +1 Bromine is reduced from 0 to -1 Changes in oxidation
number
Slide 33
1. The oxidation # of an uncombined atom = 0. 2. The oxidation
number of a monatomic ion is equal to the charge on the ion. 3. The
oxidation number of the most electronegative element in a molecule
is equal to the charge it would have if it were a ion with noble
gas configuration. 4. F is always -1 5. O is always -2 (except in
peroxides and when attached to F) Here are the Rules
Slide 34
6. H is always +1 (except when attached to more electronegative
metals, Li, Na, Ca, and Al 7. Group 1A, 2A, and 3A always have an
oxidation number equal to the group number (equal to the charge it
would have if it were a ion with noble gas configuration.) 8. Sum
of all oxidation numbers in a neutral compound is 0 9. If not
neutral, sum of all oxidation numbers is equal to the overall
charge on ion Rules cont
Slide 35
Neutral elements that are not bonded to any other element have
oxidation number of zero. Examples: Na (s), Cl 2(g), Hg (l) all
have oxidation numbers of zero. Group 1 metals that are bonded to
other elements have an oxidation number of +1 (positive one).
Examples: Na in NaCl is +1, Li in LiOH is +1 Group 2 metals that
are bonded to other elements have an oxidation number of +2
(positive two). What would Mg in MgO be? Examples
Slide 36
Group 2 metals that are bonded to other elements have an
oxidation number of +2 (positive two). Examples: Ca in CaCl 2 is
+2, Ba in Ba(OH) 2 is +2, Therefore, Mg in MgO would be +2 Oxygen
that is bonded to other elements has an oxidation number of -2
(negative two) unless it is in a peroxide or bonded to F. Examples:
O in CO 2 is -2, O in LiOH is -2, O in Na 3 PO 4 is -2 BUT WAIT
THERE ARE EXCEPTIONS Examples
Slide 37
Oxygen that is bonded to other elements has an oxidation number
of -2 (negative two) unless it is in a peroxide or bonded to F.
Examples: O in CO 2 is -2 O in LiOH is -2 O in Na 3 PO 4 is -2 What
is O in O 2 ? What is O in H 2 O? Exception: O in HOOH, hydrogen
peroxide, is -1 Examples
Slide 38
Halogens that are bonded to other elements have an oxidation
number of -1 (negative one) unless they are bonded to a more
electronegative element such as a halogen closer to the top of the
periodic table. WAIT! What happens to oxygen when it is bonded to
fluorine? Thats right, oxygen must be positive in this super rare
case! Fluorine is always -1, Why? What is O in O 2 ? zero. What is
O in H 2 O? -2
Slide 39
Halogens that are bonded to other elements have an oxidation
number of -1 (negative one) unless they are bonded to a more
electronegative element such as a halogen closer to the top of the
periodic table. Examples: Cl in NaCl is -1 Cl in PCl 5 is -1 Cl in
CaCl 2 is -1 Exception: Cl in ClF 5 is +5, Cl in ClBr 6 is -6 while
Br in ClBr 6 is +1 OK Lets face it, the larger halogens are only
easily predictable when bonded with a metal, otherwise, much
thinking is required! Examples
Slide 40
A compound has an overall charge of zero, which means all the
negative charges have to equal the positive charges. Examples: When
calculating the oxidation number of N in NO 2, use the rules above
to help you. You see that oxygen normally has an oxidation number
of -2 and there are two oxygen atoms. 2(-2) = -4. The total number
of negative charges is 4 negatives. The only other atom that is
present is nitrogen. That means the nitrogen is responsible for all
for the positive charge. X + -4 = 0. X = +4. Therefore, the
oxidation number on N in NO 2 is +4. The oxidation number of C in
CO (carbon monoxide) is +2. The oxidation number of C in CO 2
(carbon dioxide) is +4. The oxidation number of P in PCl 3
(phosphorous trichloride) is +3. The oxidation number of P in P 2 O
5 (diphosphorous pentoxide) is +5. Mn in MnO 2 is +4. Compounds are
Neutral
Slide 41
An ion has an overall charge equal to the charge of the ion.
That means the positive charges will NOT equal the negative
charges, but instead, when you add all the charges together the sum
will be equal to the charge of the ion. Example: The Mn in
permanganate ion, MnO 4 -, is =7 (Here is how: X + 4( - 2)= - 1) X
+ - 8 = - 1 add 8 to both sides. X = 7 Polyatomic Ions
Slide 42
H in HCl is +1 H in BH 3 is -1 Hydrogen bonded to a metal is
assigned -1, and hydrogen bonded to nonmetal is +1.
Slide 43
a) HCl b) KNO 3 c) OH - d) Mg 3 N 2 e) I 2 Assign oxidation
numbers to each element in the element, compound or ion. DO
NOW!
Slide 44
f) ClO 3 - g) Al(NO 3 ) 3 h) S 8 i) H 2 O 2 j) PbO 2 Assign
oxidation numbers to each element DO NOW!
Slide 45
k) NaHSO 4 l) SO 3 2- m) O 2 n) KMnO 4 o) LiH Assign oxidation
numbers to each element. More Practice!
Slide 46
p) Fe 2 O 3 q) SO 3 r) NH 4 + s) H 2 SO 4 t) Na Assign
oxidation numbers to each element in the compounds listed below.
More Practice!
Slide 47
Have you MASTERED THIS SKILL? Try it in the context of a
chemical equation!! Assign oxidation numbers to each element in the
compounds listed below. Assess yourself!
Slide 48
1) __ Mg + __ N 2 __ Mg 3 N 2 __ 2) __ Fe + __ O 2 __ Fe 2 O 3
__ 3) __ Ca + __ O 2 __ CaO __ 4) __ H 2 + __ O 2 __ H 2 O __
Assign oxidation numbers to each element
Slide 49
5) __CH 4 __+__ O 2 __CO 2 __ + __H 2 O__ 6) __ Cl 2 + __ KI __
KCl __ + __ I 2 7) __ CaCO 3 __ + __ HCl __ __ CaCl 2 __ + __ H 2
CO 3 __ __ __ Assign oxidation numbers to each element
Slide 50
8) __ CH 4 __ + __ O 2 __ CO 2 __ + __ H 2 O __ 9) __ AgNO 3 __
+ __ Cu __ Cu(NO 3 ) 2 __ + __ Ag __ __ Assign oxidation numbers to
each element
Slide 51
Fe + CuSO 4 Cu + Fe 2 (SO 4 ) 3 -omit spectator ions-ions that
dont change their O# Example: SO 4 stays at -2 Fe + Cu +2 Cu + Fe
+3 Split into half reactions and balance electron transfer Fe 3e -1
+ Fe +3 Cu +2 +2e -1 Cu Example: 6 electrons is the least common
multiple. Balance using Half Reactions
Slide 52
Split into half reactions and balance electron transfer 2(Fe 3e
-1 + Fe +3 ) 3(Cu +2 +2e -1 Cu) The iron lost a total of six
electrons as it was oxidized 2Fe 6e -1 + 2Fe +3 3Cu +2 +6e -1 3Cu
The balanced reaction is 2Fe + 3Cu +2 2 Fe +3 + 3Cu Balance using
Half Reactions Fe + CuSO 4 Cu + Fe 2 (SO 4 ) 3
Slide 53
1) __ Mg + __ N 2 __ Mg 3 N 2 __ 2) __ Fe + __ O 2 __ Fe 2 O 3
__ 3) __ Ca + __ O 2 __ CaO __ 4) __ H 2 + __ O 2 __ H 2 O __
Balance each reaction
Slide 54
5) __CH 4 __+__ O 2 __CO 2 __ + __H 2 O__ 6) __ Cl 2 + __ KI __
KCl __ + __ I 2 7) __ CaCO 3 __ + __ HCl __ __ CaCl 2 __ + __ H 2
CO 3 __ __ __ Balance each reaction.
Slide 55
8) __ CH 4 __ + __ O 2 __ CO 2 __ + __ H 2 O __ 9) __ AgNO 3 __
+ __ Cu __ Cu(NO 3 ) 2 __ + __ Ag __ __ Balance each reaction
Slide 56
Self Assessment. What happens during oxidation and reduction?
What types of reactions are also redox reactions? (Honors
Extension) How can you balance redox equations?
