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A reducing agent is the element or a compound in a redox (reduction-oxidation) reaction (see electrochemistry) that reduces another species. In doing so, it becomes oxidized, and is therefore the electron donor in the redox. For example consider the following reaction:
- 2Mg(s) + O2 → 2Mg2+(s) + 2O2-
The reducing agent in this reaction is magnesium. Magnesium donated its two valence electrons, has become an ion, and allows itself as well as oxygen to become stable.
Reducing agents need to be protected from air because they react with oxygen.
What makes a strong reducing agent?
Strong reducing agents easily lose (or donate) electrons easily. The atomic nucleus attracts its orbiting electrons. For elements whose atoms have relatively large atomic radii, the distance from the nucleus to the electrons is greater and its attraction for them weaker; these elements tend to be strong reducing agents. Also, elements that have a low electronegativity, “the ability of an atom or molecule to attract bonding electrons”1, and relatively small ionization energies serve as good reducing agents too. "The measure of a material to oxidize or lose electrons is known as its oxidation potential"2. The table below shows a few reduction potentials that could easily be changed to oxidation potential by simply changing the sign. Reducing agents can be ranked by increasing strength by ranking their oxidation potentials. The reducing agent will be the strongest when it has a more positive oxidation potential and will be a weak reducing agent whenever it has a negative oxidation potential. The following table provides the reduction potentials of the indicated reducing agent at 25° C. Also remember the useful mnemonic device, OILRIG which means, Oxidation Is Loss (of electrons) and Reduction Is Gain (of electrons).
| Oxidizing Agent | Reducing Agent | Reduction Potential (v) |
|---|---|---|
| Li+ + e- = | Li | -3.04 |
| Na+ + e- = | Na | -2.71 |
| Mg+2 + 2e- = | Mg | -2.38 |
| Al+3 + 3e- = | Al | -1.66 |
| 2H2O(l) + 2e- = | H2(g) + 2OH - | -0.83 |
| Cr+3 + 3e- = | Cr | -0.74 |
| Fe+2 + 2e- = | Fe | -0.41 |
| 2H+ + e- = | H2 | 0.00 |
| Sn+4 + 2e- = | Sn+2 | +0.15 |
| Cu+2 + e- = | Cu+ | +0.16 |
| Ag+ + e- = | Ag | +0.80 |
| Br2 + 2e- = | 2Br- | +1.07 |
| Cl2 + 2e- = | 2Cl- | +1.36 |
| MnO4-2 + 8H+ + 5e- = | Mn+2 + 4H2O | +1.49 |
In order to tell which is the strongest reducing agent, change the sign of its respective reduction potential in order to make it oxidation potential. The bigger the number the stronger a reducing agent it is.
For example if you were to list Cu, Cl-, Na and Cr in order, you get their reduction potential, change the sign to make it oxidation potential and list them from greatest to least. You well get Na, Cr, Cu and Cl-; Na being the strongest reducing agent and Cl- being the weakest one.
A few good common reducing agents include active metals such as potassium, calcium, barium, sodium and magnesium and also, compounds that contain the H- ion, those being NaH, LiAlH4 and CaH2.
Also, some elements and compounds can be both reducing or oxidizing agents. Hydrogen gas is a reducing agent when it reacts with non-metals and an oxidizing agent when it reacts with metals.
2Li(s) + H2(g) -->2LiH(s) hydrogen acts as an oxidizing agent because it accepts an electron donation from lithium, which causes Li to be oxidized.
Half Reactions 2Li(s)0 -->2Li(s)+1 + 2e-::::: H20(g) + 2e- --> 2H-1(g)
H2(g) + F2(g) --> 2HF(g) hydrogen acts as a reducing agent because it donates its electrons to fluorine, which allows fluorine to be reduced.
Half Reactions H20(g) --> 2H+1(g) + 2e-::::: F20(g) + 2e- --> 2F-1(g)
The Importance of Reducing and Oxidizing Agents
Reducing agents and Oxidizing agents are the ones responsible for corrosion, which is the “degradation of metals as a result of electrochemical activity”3. Corrosion requires an anode and cathode to take place. The anode is an element that loses electrons (reducing agent), thus oxidation always occurs in the anode, and the cathode is an element that gains electrons (oxidizing agent), thus reduction always occurs in the cathode. Corrosion occurs whenever there’s a difference in oxidation potential. When this is present, the anode metal will begin deteriorating given that there is an electrical connection and the presence of an electrolyte.
Common reducing agents
- Ferrous ion
- Lithium aluminium hydride (LiAlH4)
- Nascent hydrogen
- Potassium ferricyanide (K3Fe(CN)6)
- Sodium amalgam
- Sodium borohydride (NaBH4)
- Stannous ion
- Sulfite compounds
- Hydrazine (Wolff-Kishner reduction)
- Zinc-mercury amalgam (Zn(Hg)) (Clemmensen reduction)
- Diisobutylaluminum hydride (DIBAH)
- Lindlar catalyst
- Oxalic acid (C2H2O4)
Common reducing agents and their products
| Agent | Product |
|---|---|
| H2 Hydrogen | H+, H2O |
| metals | metal ions |
| C | CO2 carbon dioxide |
| hydrocarbons | CO2 carbon dioxide, H2O |
See also
External links
Sources
- 2,3http://www.siliconfareast.com/ox_potential.htm
- http://www.chemed.chem.purdue.edu/genchem/topicreview/bp/ch19/oxred_3.html
- http://www.members.aol.com/logan20/agents.html
- "Chemical Principles: The Quest for Insight", Third Edition. Peter Atkins and Loretta Jones pg. F76
- 1http://en.wikipedia.org/wiki/Electronegativity
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"Reducing agent"