Question

Why acidic solution such as sulfuric acid is convinient to using for reduing agent

What is the coordination environment of V3+

Answer 1

Oxidizing and Reducing Agents

Common Oxidizing Agents and Reducing Agents

The Relative Strengths of Oxidizing and Reducing Agents

Common Oxidizing Agents and Reducing Agents

In looking at oxidation-reduction reactions, we can focus on the role played by a particular reactant in a chemical reaction. What is the role of the permanganate ion in the following reaction, for example?

2 MnO₄^-(aq) + 5 H₂C₂O₄(aq) + 6 H⁺(aq) 10 CO₂(g) + 2 Mn²⁺(aq) + 8 H₂O(l)

Oxalic acid is oxidized to carbon dioxide in this reaction and the permanganate ion is reduced to the Mn²⁺ ion.

Oxidation:			H2C2O4		CO2
			+3		+4
Reduction:			MnO4-		Mn2+
			+7		+2

The permanganate ion removes electrons from oxalic acid molecules and thereby oxidizes the oxalic acid. Thus, the MnO₄^- ion acts as an oxidizing agent in this reaction. Oxalic acid, on the other hand, is a reducing agent in this reaction. By giving up electrons, it reduces the MnO₄^- ion to Mn²⁺.

Atoms, ions, and molecules that have an unusually large affinity for electrons tend to be good oxidizing agents. Elemental fluorine, for example, is the strongest common oxidizing agent. F₂ is such a good oxidizing agent that metals, quartz, asbestos, and even water burst into flame in its presence. Other good oxidizing agents include O₂, O₃, and Cl₂, which are the elemental forms of the second and third most electronegative elements, respectively.

Another place to look for good oxidizing agents is among compounds with unusually large oxidation states, such as the permanganate (MnO₄^-), chromate (CrO₄^2-), and dichromate (Cr₂O₇^2-) ions, as well as nitric acid (HNO₃), perchloric acid (HClO₄), and sulfuric acid (H₂SO₄). These compounds are strong oxidizing agents because elements become more electronegative as the oxidation states of their atoms increase.

Good reducing agents include the active metals, such as sodium, magnesium, aluminum, and zinc, which have relatively small ionization energies and low electro-negativities. Metal hydrides, such as NaH, CaH₂, and LiAlH₄, which formally contain the H^- ion, are also good reducing agents.

Some compounds can act as either oxidizing agents or reducing agents. One example is hydrogen gas, which acts as an oxidizing agent when it combines with metals and as a reducing agent when it reacts with nonmetals.

2 Na(s) + H2(g) 2 NaH(s)

H2(g) + Cl2(g) 2 HCl(g)

Another example is hydrogen peroxide, in which the oxygen atom is in the -1 oxidation state. Because this oxidation state lies between the extremes of the more common 0 and -2 oxidation states of oxygen, H₂O₂ can act as either an oxidizing agent or a reducing agent.

The Relative Strengths of Oxidizing and Reducing Agents

Spontaneous oxidation-reduction reactions convert the stronger of a pair of oxidizing agents and the stronger of a pair of reducing agents into a weaker oxidizing agent and a weaker reducing agent. The fact that the following reaction occurs, for example, suggests that copper metal is a stronger reducing agent than silver metal and that the Ag⁺ ion is a stronger oxidizing agent than the Cu²⁺ ion.

Cu(s)	+	2 Ag+(aq)		Cu2+(aq)	+	2 Ag(s)
stronger reducing agent		stronger oxidizing agent		weaker oxidizing agent		weaker reducing agent

On the basis of many such experiments, the common oxidation-reduction half-reactions have been organized into a table in which the strongest reducing agents are at one end and the strongest oxidizing agents are at the other, as shown in the table below. By convention, all of the half-reactions are written in the direction of reduction. Furthermore, by convention, the strongest reducing agents are usually found at the top of the table.

The Relative Strengths of Common Oxidizing Agents and Reducing Agents

		K+ + e- K		Best
		Ba2+ + 2 e- Ba		reducing
		Ca2+ + 2 e- Ca		agents
		Na+ + e- Na
		Mg2+ + 2 e- Mg
		H2 + 2 e- 2 H-
		Al3+ + 3 e- Al
		Mn2+ + 2 e- Mn
		Zn2+ + 2 e- Zn
		Cr3+ + 3 e- Cr
		S + 2 e- S2-
		2 CO2 + 2 H+ + 2 e- H2C2O4
		Cr3+ + e- Cr2+
		Fe2+ + 2 e- Fe
		Co2+ + 2 e- Co
		Ni2+ + 2 e- Ni
		Sn2+ + 2 e- Sn
		Pb2+ + 2 e- Pb
		Fe3+ + 3 e- Fe
		2 H+ + 2 e- H2
		S4O62- + 2 e- 2 S2O32-
		Sn4+ + 2 e- Sn2+
		Cu2+ + e- Cu+
		O2 + 2 H2O + 4 e- 4 OH-
		Cu+ + e- Cu
		I2 + 2 e- 2 I-
oxidizing		MnO4- + 2 H2O + 3 e- MnO2 + 4 OH-
power		O2 + 2 H+ + 2 e- H2O2		Reducing
increases		Fe3+ + e- Fe2+		power
		Hg22+ + 2 e- 2 Hg		increases
		Ag+ + e- Ag
		Hg2+ + 2 e- Hg
		H2O2 + 2 e- 2 OH-
		HNO3 + 3 H+ + 3 e- NO + 2 H2O
		Br2(aq) + 2 e- 2 Br-
		2 IO3- + 12 H+ + 10 e- I2 + 6 H2O
		CrO42- + 8 H+ + 3 e- Cr3+ + 4 H2O
		Pt2+ + 2 e- Pt
		MnO2 + 4 H+ + 2 e- Mn2+ + 2 H2O
		O2 + 4 H+ + 4 e- 2 H2O
		Cr2O72- + 14 H+ + 6 e- 2 Cr3+ + 7 H2O
		Cl2(g) + 2 e- 2 Cl-
		PbO2 + 4 H+ + 2 e- Pb2+ + 2 H2O
		MnO4- + 8 H+ + 5 e- Mn2+ + 4 H2O
		Au+ + e- Au
		H2O2 + 2 H+ + 2 e- 2 H2O
		Co3+ + e- Co2+
Best		S2O82- + 2 e- 2 SO42-
oxidizing		O3(g) + 2 H+ + 2 e- O2(g) + H2O
agents		F2(g) + 2 H+ + 2 e- 2 HF(aq)

Fortunately, you don't have to memorize these conventions. All you have to do is remember that the active metals, such as sodium and potassium, are excellent reducing agents and look for these entries in the table. The strongest reducing agents will be found at the corner of the table where sodium and potassium metal are listed.

Predict whether the following oxidation-reduction reactions should occur as written:

(a) 2 Ag(s) + S(s) Ag₂S(s)

(b) 2 Ag(s) + Cu²⁺(aq) 2 Ag⁺(aq) + Cu(s)

(c) MnO₄^-(aq) + 3 Fe²⁺(aq) + 2 H₂O(l) MnO₂(s) + 3 Fe³⁺(aq) + 4 OH^-(aq)

(d) MnO₄^-(aq) + 5 Fe²⁺(aq)+ 8 H⁺(aq) Mn²⁺(aq) + 5 Fe³⁺(aq) + 4 H₂O(l)

Answer

(a) No. The S^2- ion is a better reducing agent than Ag and the Ag⁺ ion is a better oxidizing agent than S. Silver doesn't tarnish because it reduces sulfur; it tarnishes because it reacts with sulfur compounds in the presence of oxygen and the oxygen is reduced.

(b) No. Cu is a better reducing agent than Ag and the Ag⁺ ion is a better oxidizing agent than Cu²⁺ ions.

(c) No. MnO₄^- in base is not a strong enough oxidizing agent to oxidize Fe²⁺ to Fe³⁺.

(d) Yes. MnO₄^- in acid is a strong enough oxidizing agent to oxidize Fe²⁺ to Fe³⁺.

he total number of points of attachment to the central element is termed the coordination number and this can vary from 2 to as many as 16, but is usually 6. In simple terms, the coordination number of a complex is influenced by the relative sizes of the metal ion and the ligands and by electronic factors, such as charge which is dependent on the electronic configuration of the metal ion. These competing effects are described by the term ionic potential which is defined as the charge to radius ratio (q/r).

Based on this, it can be seen that the bigger the charge on the central ion, the more attraction there will be for negatively charged ligands, however at the same time, the bigger the charge the smaller the ion becomes which then limits the number of groups able to coordinate.

Coordination Number 2

This arrangement is not very common for first row transition metal ion complexes and some of the best known examples are for Silver(I). In this case we have a low charge and an ion at the right hand side of the d-block indicating smaller size

A method often employed for the detection of chloride ions involves the formation of the linear diamminesilver(I) complex.
The first step is:

       Ag+    +    Cl-      ?   AgCl (white ppt)

and to ensure that the precipitate is really the chloride salt, two further tests must be done:

      AgCl     +   2 NH3     ?   [Ag(NH3)2]+

and

     [Ag(NH3)2]+   +  HNO3   ?    AgCl (re-ppts)

The reaction of a bidentate ligand such as 1,2-diaminoethane with Ag(I) does not lead to chelated ring systems, but instead to linear two coordinate complexes. One reason for this is that bidentate ligands can NOT exist in trans arrangements, that is they can NOT span 180 degrees.

Coordination Number 3

Once again, this is not very common for first row transition metal ions. Examples with three different geometries have been identified:

Trigonal planar

Well known for main group species like CO₃^2- etc., this geometry has the four atoms in a plane with the bond angles between the ligands at 120 degrees.

Trigonal pyramid

More common with main group ions.

T-shaped

The first example of a rare T-shaped molecule was found in 1977 however since then several further examples have been reported.

Coordination Number 4

Two different geometries are possible. The tetrahedron is the more common while the square planar is found almost exclusively with metal ions having a d⁸ electronic configuration.

Tetrahedral, (T_d)

The chemistry of molecules centred around a tetrahedral C atom is covered in organic courses. To be politically correct, please change all occurrences of C to Co. There are large numbers of tetrahedral Cobalt(II) complexes known.

Square Planar, (D_4h)

These are much less common than tetrahedral and are included only because some extremely important molecules exist with this shape.

Coordination Number 5

Square pyramid, (C_4v)

Trigonal Bipyramid, (D_3h)

The structure of [Cr(en)₃][Ni(CN)₅] 1.5 H₂O was reported in 1968 to be a remarkable example of a complex exhibiting both types of geometry in the same crystal.
The reaction of cyanide ion with Ni²⁺ proceeds via several steps:

        Ni2+         +   2 CN-     ?   Ni(CN)2

        Ni(CN)2      +   2 CN-     ?   [Ni(CN)4]2-    
                                         orange-red

                                      log(?4) = 30.1

        [Ni(CN)4]2-  +     CN-     ?   [Ni(CN)5]3-
                                          deep red

Oxovanadium salts (Vanadyl, VO²⁺) often show square pyramidal geometry, for example, VO(acac)₂. Note that the Vanadium(IV) can be considered coordinatively unsaturated and addition of pyridine leads to the formation of an octahedral complex.

Coordination Number 6

Hexagonal planar

Unknown for first row transition metal ions, although the arrangement of six groups in a plane is found in some higher coordination number geometries.

Trigonal prism

Most trigonal prismatic compounds have three bidentate ligands such as dithiolates or oxalates and few are known for first row transition metal ions.

Octahedral, (O_h)

The most common geometry found for first row transition metal ions, including all aqua ions.
In some cases distortions are observed and these can sometimes be explained in terms of the Jahn-Teller Theorem.

Coordination Number 7

Three geometries are possible:
Not very common for 1st row complexes and the energy difference between the structures seems small and distortions occur so that prediction of the closest "idealised" shape is generally difficult.