Premed · Premed · General Chemistry 2
Lecture 13: Electrochemistry: Redox Reactions
General Chemistry II
Learning Objectives
By the end of this lecture, students will be able to:
- Define oxidation, reduction, oxidizing agent, and reducing agent
- Assign oxidation states to atoms in molecules and ions
- Identify redox reactions and the species being oxidized and reduced
- Balance redox equations using the half-reaction method in acidic and basic solutions
- Relate electron transfer to the concept of electrochemical cells
Lecture Content
I. Review of Oxidation and Reduction
Oxidation is the loss of electrons, corresponding to an increase in oxidation state. The mnemonic "OIL" (Oxidation Is Loss) helps remember this. Reduction is the gain of electrons, corresponding to a decrease in oxidation state, remembered by "RIG" (Reduction Is Gain). These two processes are inseparable: you cannot have oxidation without reduction, and vice versa.
The oxidizing agent (oxidant) is the species that gets reduced; it causes the oxidation of another species by accepting electrons from it. The reducing agent (reductant) is the species that gets oxidized; it causes the reduction of another species by donating electrons to it. Another helpful mnemonic is "LEO says GER": Lose Electrons = Oxidation, Gain Electrons = Reduction.
II. Assigning Oxidation States
The oxidation state, also called the oxidation number, is a bookkeeping tool used to track the transfer of electrons in chemical reactions. A set of rules, applied in order of priority, governs the assignment. Free elements have an oxidation state of 0. Monatomic ions have an oxidation state equal to their charge. Fluorine is always -1 in compounds. Oxygen is usually -2, with exceptions in peroxides (-1) and in OF2 (+2). Hydrogen is usually +1 when bonded to nonmetals and -1 when bonded to metals (in hydrides). Group 1A metals are always +1 and Group 2A metals are always +2 in compounds. Finally, the sum of all oxidation states must equal zero for a neutral compound or equal the charge for a polyatomic ion.
Applying these rules: in H2SO4, hydrogen is +1, oxygen is -2, and sulfur must be +6. In MnO4-, oxygen is -2 and manganese must be +7 to give an overall charge of -1. In Cr2O7^2-, oxygen is -2 and each chromium is +6.
III. Identifying Redox Reactions
A redox reaction is identified by checking whether any atoms undergo a change in oxidation state between reactants and products. The species whose oxidation state increases has been oxidized, while the species whose oxidation state decreases has been reduced. Not all reactions are redox reactions: acid-base neutralizations, simple precipitation reactions, and certain decomposition reactions involve no electron transfer.
For example, in the reaction 2Fe(s) + 3Cl2(g) -> 2FeCl3(s), iron goes from 0 to +3 (it is oxidized and serves as the reducing agent), while chlorine goes from 0 to -1 (it is reduced and serves as the oxidizing agent).
<image>A diagram showing the redox reaction between zinc metal and copper(II) sulfate solution. Panel A: Before reaction -- a zinc strip is placed in a blue CuSO4 solution. Oxidation states are labeled: Zn = 0, Cu = +2. Panel B: After reaction -- the zinc strip is partially dissolved and coated with reddish copper metal. The solution has become lighter. Arrows show: Zn -> Zn^2+ + 2e- (oxidation, labeled "reducing agent") and Cu^2+ + 2e- -> Cu (reduction, labeled "oxidizing agent"). The net equation Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s) is written below.</image>
IV. Balancing Redox Equations: The Half-Reaction Method (Acidic Solution)
The half-reaction method provides a systematic procedure for balancing redox equations. First, identify what is oxidized and what is reduced, and separate the overall reaction into two half-reactions. Then balance each half-reaction in four sub-steps: balance all atoms except oxygen and hydrogen, balance oxygen by adding H2O, balance hydrogen by adding H+, and balance charge by adding electrons. Next, equalize the number of electrons in both half-reactions by multiplying each by the appropriate factor. Finally, add the half-reactions together and cancel any species that appear on both sides. Verify that both atoms and charges are balanced.
As an example, consider MnO4-(aq) + Fe^2+(aq) -> Mn^2+(aq) + Fe^3+(aq) in acidic solution. The reduction half-reaction is MnO4- + 8H+ + 5e- -> Mn^2+ + 4H2O. The oxidation half-reaction is Fe^2+ -> Fe^3+ + e-, which must be multiplied by 5. Adding gives the balanced equation: MnO4- + 8H+ + 5Fe^2+ -> Mn^2+ + 4H2O + 5Fe^3+.
V. Balancing Redox Equations in Basic Solution
To balance a redox equation in basic solution, first complete the entire procedure as if the solution were acidic. Then convert to basic solution by adding OH- to both sides for each H+ present, combining H+ and OH- to form H2O, and canceling any water molecules that appear on both sides.
VI. Recognizing Common Oxidizing and Reducing Agents
Certain reagents appear frequently as oxidizing or reducing agents. Strong oxidizing agents, which are easily reduced, include KMnO4 (permanganate, purple, reduced to colorless Mn^2+ or brown MnO2), K2Cr2O7 (dichromate, orange, reduced to green Cr^3+), concentrated HNO3, H2O2, and the halogens in order of decreasing strength (F2 > Cl2 > Br2 > I2).
Strong reducing agents, which are easily oxidized, include the alkali metals (Li, Na, K), alkaline earth metals (Mg, Ca), active metals (Zn, Fe, Al), hydrogen gas, and carbon and carbon monoxide in metallurgical applications.
<image>A summary table of the half-reaction balancing method in acidic solution, presented as a flowchart. Step 1: "Separate into oxidation and reduction half-reactions." Step 2a: "Balance atoms other than O and H." Step 2b: "Balance O with H2O." Step 2c: "Balance H with H+." Step 2d: "Balance charge with e-." Step 3: "Multiply half-reactions so electrons cancel." Step 4: "Add half-reactions and simplify." Step 5: "Verify atoms and charge balanced." An additional branch from Step 5 leads to "For basic solution: add OH- to neutralize H+, form H2O, cancel duplicates." Each step includes a small example using the Cr2O7^2-/Fe^2+ reaction.</image>
VII. Disproportionation Reactions
A disproportionation reaction is a special type of redox reaction in which the same element is simultaneously oxidized and reduced. This is possible only when the element starts in an intermediate oxidation state. In the decomposition of hydrogen peroxide, 2H2O2 -> 2H2O + O2, the oxygen in H2O2 is in the -1 oxidation state. In the products, oxygen in water is -2 (reduced) while oxygen in O2 is 0 (oxidized). Another example is the disproportionation of Cu+: 2Cu+(aq) -> Cu(s) + Cu^2+(aq), where Cu+1 is both reduced to Cu0 and oxidized to Cu+2.
