2: Metals in Bioinorganic Redox

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The formal oxidation state is the hypothetical charge that an atom would have if

all bonds to atoms of different elements were 100% ionic.

You have already used this when you calculated charge (oxidation state) of the metal in coordination complexes. The charge assigned to the metal is the same thing as its “oxidation state”, written with Roman numerals.

• Calculate the oxidation state on these complexes.

Redox and Reduction Potentials

Redox (reduction-oxidation) reactions are concerned with the transfer of electrons between species.

In simple terms:
Oxidation is the loss of electrons.

Reduction is the gain of electrons.

Metals can do redox more readily than main group elements. The metal ion interconverts between the oxidation states. For example iron, interconverts between Fe²⁺ and Fe³⁺ states in electron-transfer processes.

• How many electrons would a Cu⁺ ion (take / give up) to become Cu²⁺?

A reduction potential is the amount of energy involved in a redox change. A positive potential means energy is released; a negative potential means energy is consumed.

Reduction Potentials

Cu²⁺ + e^- -> Cu⁺ E⁰= + 0.153V

Fe³⁺ + e^- -> Fe²⁺ E⁰= + 0.771V

Mn³⁺ + e^- -> Mn²⁺ E⁰= + 1.51V

Co³⁺ + e^- -> Co²⁺ E⁰= + 1.84V

• Using the table above, how much energy would a Cu²⁺ ion (use up / give off) to become Cu⁺?

To go backwards, the sign of the reduction potential would change.

• Using the table, what would the reduction potential be for the reaction of Cu⁺¹ ion converting to Cu⁺²?

• Draw the reaction progress diagram showing the relative energy of the reaction of Cu⁺ proceeding to Cu²⁺.

Reduction Potentials in Reactions

Reduction potentials are used to predict if a given metal ion is able to drive a transformation. The reactions are broken into two parts: one reaction showing the donation of electrons and the other reaction showing the component accepting electrons. These are called half-reactions.

For example, the conversion of atmospheric oxygen to superoxide is shown here (a common cellular process).

O₂ +e^- -> O^2- E⁰ = - 0.45V

In order to drive this reaction, we need a metal that is willing to donate electrons.

M -> M⁺ + e^-

• Draw the Cu(I) to Cu(II) half-reaction in the direction that would potentially drive the formation of superoxide.

• What is the reduction potential for this reaction? ___________

• Add the reduction potentials for the two half-reactions.

  E⁰_red of O2: ____________

  E⁰_ox of Cu(I): ____________

  Net E⁰_rxn: _____________

• Would formation of copper (II) ion be able to convert oxygen into superoxide? Why or why not?

• List two ions from the reduction table at the beginning of the workbook that would be able to convert oxygen to superoxide.

Reduction Potentials in Reactions with Multiple Electrons

Suppose atmospheric oxygen is going to be converted into hydrogen peroxide. The balanced equation for that reaction is

O₂ + 2H⁺ +2e^- -> H₂O₂ E⁰ = +0.70V

The reduction potential is positive. The reaction would release energy.

• Draw the half reaction to convert a hydrogen peroxide molecule to atmospheric oxygen in the presence of protons.

• What is the E⁰ for this reaction? _____________

• Draw the half reaction for Cu(II)/(I) reaction that would balance this reaction.

• What is the E⁰ for this reaction? _____________

• Would the copper (II) ion be able to convert a hydrogen peroxide molecule to atmospheric oxygen in the presence of protons? Why or why not?

• Draw the half reaction for Mn(III)/(II) reaction that would balance this reaction.

• What is the E⁰ for this reaction? _____________

• Would this manganese reaction be able to convert a hydrogen peroxide molecule to atmospheric oxygen in the presence of protons?

• To balance the number of electrons in the reaction, how many manganese(III) ions would be needed to convert a hydrogen peroxide molecule to atmospheric oxygen in the presence of protons?

Ligand Effects

However, metals are not found as aqueous complexes in biological systems; they usually have ligands. Ligands can affect the reduction potential.

• Complete the table below by rating the ligands on the complexes as hard or soft bases.

Electrode Potentials of Cobalt and Copper Species

Cu(II)-Cu(I) Reactions	E^0(V)	HSAB ligands
Cu^2+ + 2 CN^- + e^- -> [Cu(CN)2]^-	+1.103
Cu^+2 + I^- + e^--> CuI	+0.86
Cu^+22+ Cl^- + e^- -> CuCl	+0.538
Cu(aq)^+2 + e^- -> Cu(aq)^+	+0.153
[Cu(NH3)4]^+2 + e^- -> [Cu(NH3)3]^+ + 2 NH3	-0.01

Miessler and Tarr, In Inorganic Chemistry, 3rd Edition, (2004) Pearson Education, NJ.

• Compare the reduction potentials of Cu(II) complexed with a hard ligand vs. Cu(II) complexed with a soft ligand. Which is more easily reduced? Which ligand stabilizes the higher oxidation state?

• Metals with ________ (hard / soft ) ligands tend to have higher reduction potentials.

• Propose a reason for this trend.

Effects of Biological ligands on reduction potentials

Nature has chosen metalloproteins because their redox potentials can be tunable. The reduction potential can vary due to a variety of factors.

Plastocyanin is an important copper-containing protein involved in electron- transfer of Photosystem I of photosynthesis.

The copper binding site is described as a ‘distorted tetrahedron’. The ligands are two nitrogen atoms (N1 and N2) from separate histidines and a sulfur (S1) from a cysteine. Sulfur, S2, from an axial methionine forms the apex bond. The Cu-S1 contact is shorter (2.13 Å) than Cu-S2 (2.9 Å).

• Draw this Cu binding site.

Bond Distances (Å)

	Cu(II)	Cu(I), pH 7	Cu(I), pH 3.8
Cu-S(Cys84)	2.13	2.17	2.13
Cu-S(Met92)	2.90	2.87	2.51
Cu-N(His37)	2.04	2.13	2.12
Cu-N(His87)	2.10	2.39	>4

Lippard and Berg, In Principles of Bioinorganic Chemistry. (1994) University Science Books, CA.

The elongated Cu-S2 bonding destabilizes the Cu(II) form and makes the reduction potential more positive.

• Explain why the long Cu-S2 bond makes the reduction potential more positive.

In the reduced form of plastocyanin, Cu(I), the Cu-S2 bond becomes 2.5 Å and the His-87 dissociates from the metal to become protonated.

• What would the new copper site geometry be?

• Why would Cu(I) prefer to bond to Met than to His? Hint: Consider Hard/Soft Acid/Base considerations.

Effects of Solvent Accessibility on Reduction Potentials

An excess of charge can be tolerated better in a solvent with a high dielectric constant (dielectric is one way to measure the polarity of a solvent).

• Explain why this observation (above) is true.

An increase in hydrophilicity makes it easier to build up charge. dielectric constants below, fill in the following statement.

‘Solvent’	Dielectric constant
Inner parts of protein	2-8
Water	80

• In a protein, a metal center and its corresponding charge is more likely to be located by _________.

Electron Transfer Protein Example

Min, et.al., Protein Sci. 2001 March; 10(3): 613–621.

Rubredoxins are iron-containing electron transfer proteins found in bacteria. The rubredoxin active site contains an iron ion coordinated by the sulfurs of four conserved cysteine residues.

• If the Fe is +2, what is the overall charge on this active site? _________

• If the Fe is +3, what is the overall charge on this active site? _________

Rubredoxin has two ways to stabilize charge changes.

1. Contraction of the Fe-Ligand bonds to greater electrostatic stabilization of the extra negative charge.

   

   • How do the shortened bonds in structure b stabilize the active site?

2. The reduced form has increased amounts of hydration around the redox site. It has six water molecules within an 8 Å radius around the Fe center, whereas the oxidized form has only one water molecule within the same radius.
   • Explain how these changes stabilize structure a.

Summary of Reduction Potentials and Bioinorganic Effects

• Choose the best answer for the following concepts:
  • Oxidation is the ________ (loss/gain) of electrons.

  • Reduction is the ________ (loss/gain) of electrons.

  • A _____________ (positive/negative) net reduction potential means that the reaction is favored.

  • Metals with ________ (hard / soft) ligands tend to have higher reduction potentials.

  • A longer bond in complex will __________ (increase / decrease) the reduction potential.

  • In a protein, a metal center and its corresponding charge is more likely to be located by _________ (hydrophobic / hydrophilic ) group.

  • In a protein, a complex with a negative charge will have a __________ (higher / lower) reduction potential than a complex with no charge.

• Define a half reaction:

Redox Practice Problems:

1. Suppose atmospheric oxygen is going to be converted into water. The balanced equation for that reaction is

   O₂ +4H⁺ +4e^- = 2H₂O E⁰ = +1.229V

   • Would copper(II) ions be able to drive the conversion of two waters to an atmospheric oxygen in the presence of protons? If so, how many would be needed to balance the reaction?

   • Would manganese(III) ions be able to drive the conversion of two waters to an atmospheric oxygen in the presence of protons? If so, how many would be needed to balance the reaction?

2. Electron-transfer rates have been studied in ruthenium-modified myoglobin. Reduced myoglobin (Ru⁺²) is a high-spin five-coordinate complex, whereas the oxidized form (Ru⁺³) binds a water molecule to form a low spin six- coordinate complex.
   • Draw d orbital splitting diagrams for the reduced form and the oxidized form. Explain why the oxidized form becomes low spin.

   • Why does the oxidized form bind an extra water molecule?

3. The following reactions as written are examples of oxidations/reductions (circle one).

   O₂ + 2H⁺ + 2e^- -> H₂O₂ E^o = + 0.281V

   Cu²⁺ + e^- -> Cu⁺ E^o = + 0.4 V (blue protein)

   Cu²⁺ + e^- -> Cu⁺ E^o = + 0.153V (water)

   • Provide an equation showing the formation of molecular oxygen from hydrogen peroxide, as well as its potential.

   • Suppose copper(II) were employed to assist in the formation of oxygen from hydrogen peroxide. The Cu(II) would act as a reducing/oxidizing agent (circle one).

   • The reaction of hydrogen peroxide to oxygen would be favored/disfavored in the presence of aqueous copper ion? Show why.

   • The reaction of hydrogen peroxide to oxygen would be favored/disfavored in the presence of blue copper protein? Show why.

   • In a blue copper protein, copper is bound by a sulfur from methionine, a sulfur from cysteine, and nitrogens from two histidines. Draw this center or ask for it.

   • Provide a reason for why its potential is so different from that of aqueous copper ion.

4. Prediction of reaction spontaneity.
   • Spontaneous processes have (positive/negative) ΔG’s.

   • Spontaneous processes have (positive/negative) E^o’s.

   This means that we can use the table to predict reaction spontaneity of redox reactions.

   • Will Cl₂ react spontaneously with Ni? Write a balanced equation. Support your conclusion with data from the table.

   • Will Sn²⁺ spontaneously react with Zn? Write a balanced equation. Support your conclusion with data from the table.

   • Will 1 M HCl dissolve (react with) Cu? Explain. Write a balanced equation.

   • Will Cr³⁺ react with Ag? Explain. Write a balanced equation. (Note that a balanced equation can be obtained by multiplying the half reactions by appropriate factors so that the number of electrons cancel when the equations are added. Multiplication does not affect the value of E^o.)

   • Assume that the Standard Reduction table is listed with the most positive half reaction at the top. A student notes that for any particular reduction half-reaction in the table, that reactant can oxidize any species (listed as a product) in a half reaction below it. Explain why this is true.