Redox - C & EE 255B Prof. M. K. Stenstrom Winter 2011...

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C & EE 255B Prof. M. K. Stenstrom Winter 2011 BALANCING REDOX EQUATIONS Balancing redox (oxidation-reduction) equations is a simple and very useful technique of performing balances from empirical equations describing microbial stoichiometry. Each basic equation (synthesis or growth, respiration, and decay) can all be balanced and added together to describe a process. Moreover, the equations can be balanced for each type of metabolism (aerobic or oxic, anoxic, and anaerobic). This handout describes the techniques for balancing the equations and then shows some common examples. Be careful when using empirical redox techniques; the reactions can be balanced, but other information must be used to determine if the reaction actually occurs. RULES Each redox equation contains two parts -- the oxidation and reduction parts. Each is balanced separately. 1. The first rule is to balance the major atoms with known end products. The end products of the redox equations must be stated or determined from other sources. The redox equations give you no information about the actual end products. Common end products for carbon are CO 2 or cells. Other end products can occur as well. Major atoms are defined as all atoms except oxygen and hydrogen. 2. The next step is to balance the oxygen atoms by adding water (H 2 O) molecules. 3. Next balance the hydrogen with hydrogen ions (H + ). 4. Finally balance the charge with electrons (e - ). After these four steps one obtains a balanced half-reaction -- either the oxidation reaction or the reduction reaction. Oxidation reactions will produce electrons (electrons appear on the right-hand side of the equation). Reduction reactions will consume electrons (electrons appear on the left-hand side of the equation). SOME EXAMPLES Consider the oxidation of glucose, C 6 H 12 O 6 : Step 1. Balance the major atoms. In this case we will use CO 2 as the end product for carbon. C 6 H 12 O 6 --------> 6 CO 2 Step 2. Balance the oxygen with water: C 6 H 12 O 6 + 6 H 2 O --------> 6 CO 2 Step 3. Balance the hydrogen: C 6 H 12 O 6 + 6 H 2 O --------> 6 CO 2 + 24 H + Step 4. Balance the charge with electrons:
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2 C 6 H 12 O 6 + 6 H 2 O --------> 6 CO 2 + 24 H + + 24 e - (1) Note that the charge and all atoms are balanced. This is the balanced half-reaction. It is an oxidation reaction. Note that electrons are produced or are removed from the substrate (glucose). The oxidation state of carbon is zero on the left hand side and + 2 on the right hand side of the equation. For this reaction to occur it must be balanced by a second half-reaction. It contains a reactant that is called the hydrogen acceptor or electron acceptor. Metabolisms are distinguished by which electron acceptors are used. Aerobic metabolism uses oxygen, anoxic metabolism uses nitrate or nitrite or sulfate (generally in that order of declining preference), and anaerobic metabolism uses CO 2 or an organic molecule. Environmental engineers need to differential between anoxic and anaerobic metabolism, since the
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This note was uploaded on 02/02/2012 for the course CEE 255B taught by Professor Michaelstenstrom during the Fall '11 term at UCLA.

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Redox - C & EE 255B Prof. M. K. Stenstrom Winter 2011...

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