Redox Reactions: MDCAT Chemistry notes
Redox Reactions MDCAT notes: oxidation and reduction, oxidizing and reducing agents, oxidation numbers, displacement reactions and electrochemical cells.
Oxidation and reduction
| Oxidation | Reduction |
|---|---|
| Loss of electrons | Gain of electrons |
| Increase in oxidation number | Decrease in oxidation number |
| Addition of oxygen / removal of hydrogen | Removal of oxygen / addition of hydrogen |
Oxidation and reduction always occur together, so the whole reaction is called a redox reaction.
- Oxidizing agent: accepts electrons and is itself reduced.
- Reducing agent: donates electrons and is itself oxidized.
Example: $\mathrm{H_2 + Cl_2 \rightarrow 2HCl}$. Hydrogen goes from 0 to +1 (oxidized, reducing agent); chlorine from 0 to −1 (reduced, oxidizing agent). Physical changes such as boiling, evaporating or dissolving NaCl are not redox reactions.
Oxidation number rules
- Free element: 0 (e.g. Na, $\mathrm{O_2}$, $\mathrm{Cl_2}$).
- Monatomic ion: equal to its charge.
- Oxygen: −2 (−1 in peroxides). Hydrogen: +1 (−1 in metal hydrides).
- Group 1 metals +1; group 2 metals +2; fluorine always −1.
- Sum in a neutral compound is 0; in an ion, equal to the charge.
Worked example: in $\mathrm{K_2Cr_2O_7}$, $2(+1) + 2x + 7(-2) = 0$, so $x = +6$. In $\mathrm{MnO_4^-}$, $x + 4(-2) = -1$, so $x = +7$.
Displacement reactions and the electrochemical series
A more reactive metal (higher in the activity series, more negative electrode potential) displaces a less reactive metal from a solution of its salt.
Order of some metals: K > Na > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au.
- $\mathrm{Fe + CuSO_4 \rightarrow FeSO_4 + Cu}$: copper is deposited on the iron.
- $\mathrm{Mg + FeSO_4 \rightarrow MgSO_4 + Fe}$: iron is precipitated.
- $\mathrm{Zn + CuSO_4 \rightarrow ZnSO_4 + Cu}$ is feasible, but Cu cannot displace Zn, and Zn or Cd cannot displace Mg.
Strength of halogens and halide ions
Oxidizing power: $\mathrm{F_2 \gt Cl_2 \gt Br_2 \gt I_2}$. The halide ions show the reverse as reducing agents: $\mathrm{I^- \gt Br^- \gt Cl^- \gt F^-}$. A larger halide ion holds its extra electron less tightly and gives it up more easily. Neutral halogen molecules are oxidizing agents, not reducing agents.
Electrochemistry and cells
Electrochemistry is the branch of chemistry dealing with the conversion of chemical energy into electrical energy and vice versa. All electrochemical cells are based on redox reactions.
| Galvanic (voltaic) cell | Electrolytic cell |
|---|---|
| Spontaneous redox reaction produces electricity | Electricity drives a non-spontaneous reaction |
| Two metals of different reactivity joined through an electrolyte (and salt bridge) | Electrodes dipped in an electrolyte connected to a battery |
| Example: Daniell cell (Zn / Cu) | Example: electrolysis of fused NaCl |
In every cell, oxidation occurs at the anode and reduction at the cathode. In a galvanic cell the more reactive metal is the anode (negative), and electrons flow through the external wire to the cathode.
Common MDCAT traps
- The oxidizing agent is reduced; the reducing agent is oxidized.
- Among $\mathrm{Cl_2}$, $\mathrm{Cl^-}$, $\mathrm{Br_2}$, $\mathrm{Br^-}$ the strongest reducing agent is $\mathrm{Br^-}$.
- Only a more reactive metal displaces a less reactive one; check the series order.
- A galvanic cell needs two different metals with an electrolyte, not direct contact.
Quick revision
- Oxidation: loss of electrons, increase in oxidation number.
- Electrochemical cells run on redox reactions.
- Iron in $\mathrm{CuSO_4}$ deposits copper.
- Anode: oxidation; cathode: reduction.
- Reducing power of halides: $\mathrm{I^- \gt Br^- \gt Cl^- \gt F^-}$.