Unit 2: Electrochemistry
Content Outline: Redox Reactions (2.1)
Associated AP Learning Objectives: 1.6, 2.26, 3.2, 3.8
As you go through this year in AP Chemistry, you must begin building skill of “seeing” a possible definition in a term…not memorizing all terms as there are way to many for you to memorize. Break words apart and think about a possible definition using the parts.
Important Concepts from previous units:
1) Electricity is associated with electron flow and ions. Can you see “flow of electrons” in the term “electricity”?
2) Electron quantities can be expressed in moles too. Moles are for expressing quantities of matter remember.
I. Law of Conservation of Matter
A. Matter is neither created nor destroyed; just transferred and transformed. Electrons are flowing from one element to another element and thereby transforming the charges.
II. Reduction-Oxidation Reaction (Redox Reaction)
A. Reduction (RIG) Reduction is gaining (“tion” means “process of”)
1. The process of a molecule gaining a negatively charged electron.
a. The charge of the molecule is reduced (made more negative) by accepting the negative electron.
b. The electron came from a molecule called the reducing agent.
Think of an “agent” as a molecule looking to reduce another molecule.
B. Oxidation (OIL) Oxidation is losing
1. The process of a molecule losing a negatively charged electron to another molecule.
a. The charge of the molecule is increased (made more positive) by losing the negative electron.
b. The electron was accepted by the oxidizing agent.
Think of an “agent” as a molecule looking to oxidize another molecule.
C. The Law of Conservation of Matter makes these always linked together…hence a Redox reaction.
D. Disproportionation
1. This is a molecule that can either be oxidized “agent or reduction “agent”
For example: H2O2(l) à H2O(l) + O2 (g).
Oxygen gas (O2) is oxidized…loses electrons.
Oxygen in H20 is reduced… gains electrons.
III. Oxidation States
A. These are “charges” assigned to elements, ions, polyatomic ions, and molecules.
B. Oxidation Rules: Please look at a Periodic Table as you progress and think about the location of elements as it relates to the Noble gases. Electrons are being transferred so as to ultimately behave “like” a Noble gas…in the lowest possible energy state.
1. A pure element is assigned charge of 0. For example, O2 = 0 charge.
2. A monatomic ion is assigned a charge equal to the ion strength. For example, Cl- = -1
3. The more electronegative element in a binary (2) compound is assigned the same charge it would have as a free ion. For example, H2O… Oxygen is -2 and H is +1.
4. Fluorine is always assigned a charge of -1.
5. Oxygen is always assigned a charge of -2, EXCEPT when with Fluorine then it is +1 or +2, OR in a Peroxide then it is assigned -1. For example, OF… O is +1 to balance F -1. H2O2… each O is assigned a -1 because each H is assigned +1. They must balance to zero, as it is not an ion.
6. Hydrogen is always assigned a charge of +1, EXCEPT when with a metal, then it is assigned a charge of -1. For example, LiH
7. Group One metals are assigned a charge of +1, Group 2 metals are assigned a charge of+2, Aluminum metal is always +3.
8. The sum of the oxidation numbers of all atoms in a polyatomic ion must equal that ion’s charge.
For Example, H2PO4- Each H is +1… so that equals +2. Each O equals -2… so it -8 total. -8 + 2 = -6. So phosphate must be +5 so that we have one negative left…-6 +5 = -1.
9. The sum of the oxidation numbers of all atoms in a neutral compound is 0. See H2O2 above.
IV. Balancing Redox reactions using the Half-Reaction method
A. One balanced reaction must be written for each reaction (Oxidation and Reduction) displaying the electron transfer and oxidation numbers.
1. You MUST pay close attention to charge of each element, vocabulary and look for ions.
B. Step 1: Write the formula equation for the problem, if not provided.
For example, H2S + HNO3 à H2SO4 + NO + H2O
Step 2: Take the equation and write it in ionic form, breaking apart compounds into ions or polyatomic ions. Do you still remember your common polyatomic ions? Remember this is electrochemistry… so we are talking about IONS (Charged particles).
H2S + H+ + NO3- à 2H+ + SO42- + NO2 + H2O
Step 3: Assign all elements/ions/polyatomic ions their charges using your rules above.
H2S + H+ + NO3- à 2H+ + SO42- + NO2 + H2O
+1 -2 +1 +5 -2 +1 +6 -2 +4 -2 +1 -2
Step 4: Delete elements that do not change charge values.
H2S+ NO3- à SO42- + NO2
Step 5: Write the half reaction for oxidation (lose of electrons)
H2S- à SO42- Sulfur lost 8 electrons to go from -2 to +6 in charge.
Step 6: Balance Oxygen using water.
H2S- + 4H2Oà SO42- + 10H+
Step 7: Balance Charges for this half reaction by adding electrons.
H2S- + 4H2Oà SO42- + 10H+ + 8 electrons
Step 8: Write the half reaction for the reduction (gain of electrons)
NO3- à NO2
Step 9: Balance the Oxygens with water You may need to add H+ to the opposite side in some reactions.
NO3- + 2 H+ àNO2 + H2O
Step 10: Balance Charges for this half reaction by adding electrons.
NO3- + 2 H+ + 1 electron àNO2 + H2O
Step 11: Balance electron transfer by adjusting coefficients in the half reactions.
(NO3- + 2 H+ + 1 electron àNO2 + H2O) x 8
Step 12: Combine half reactions.
H2S + 16H+ + 8NO3- + 4H20 + 8 electronsà 10H+ + SO42- + 8NO2 + 8H2O + 8 electrons
Step 13: Cancel and reduce like terms and right the remaining formula.
H2S + 16H+ + 8NO3- + 4H20 + 8 electronsà 10H+ + SO42- + 8NO2 + 8H2O + 8 electrons
Becomes
H2S + 8HNO3 à 8NO2 + 4H2O + H2SO4
V. Reduction and Oxidation strength series tables
A. You need to have a series table to predict reaction possibility.
B. On the oxidation series (losing electrons), reactions will occur with oxidizing agent molecules under but not above.
C. The higher on the table, the higher the reducing (lower oxidizing) strength.