Transition metals

"The transition elements are those elements having a partially filled d or f subshell in any common oxidation state."

The term "transition elements" most commonly refers to the d-block transition elements. The 2B elements zinc, cadmium and mercury do not strictly meet the defining properties, but are usually included with the transition elements because of their similar properties. The f-block transition elements are sometimes known as "inner transition elements". The first row of them is called the lanthanides or rare earths. The second row consists of the actinides. All of the actinides are radioactive and those above Z=92 are manmade in nuclear reactors or accelerators.

The general properties of the transition elements are

1. They are usually high melting point metals.
2. They have several oxidation states.
3. They usually form colored compounds.
4. They are often paramagnetic.

The transition elements include the important metals iron, copper and silver. Iron and titanium are the most abundant transition elements. Many catalysts for industrial reactions involve transition elements

(c) doc b 2(c) doc b 3(c) doc b 4(c) doc b 5(c) doc b 6(c) doc b 7(c) doc b 8(c) doc b 9(c) doc b 10(c) doc b

                                  2 to 9 are typical transition metal ion/compound coloured solutions

There are four transition series:

=>The first transition series: Scandium (Sc) through Copper (Cu):

=>The second transition series: Yttrium (Y) through Silver (Ag):

=>The third transition series: Lanthanum (La) to Hafnium (Hf) through Gold (Au))

=>The forth transition series which is incomplete: Actinium (Ac) to element 104 through element 109: 6d      

     subshell is filling, If elements 110 and 111 are found then this will complete this series.

Electronic Configuration of first transition series:

Group (column) number	3	4	5	6	7	8	9	10	11	12
atomic numberelement symbol	21Sc	22 Ti	23 V	24Cr	25Mn	26Fe	27Co	28Ni	29Cu	30Zn
electron configuration	3d14s2	3d24s2	3d34s2	3d54s1	3d54s2	3d64s2	3d74s2	3d84s2	3d104s1	3d104s2

The valence configuration for first series transition metals (Groups 3 - 12) is usually 3dⁿ 4s².

Exceptions:  The electron configurations for chromium (3d⁵ 4s¹) and copper (3d¹⁰ 4s¹) are exceptions.
This is because the 3d and 4s orbitals are very close in energy, and the energy of the 3d orbital drops going across the row.
For both chromium and copper the configuration which has more electrons in in the 3d subshell is of lower energy.  

At chromium  the difference in 3d and 4s orbital energies is of the order of the pairing energy (remember electron pairing requires energy).
The  3d⁵ 4s¹ configuration is of lower energy because this configuration has the maximum number of unpaired electrons for a d-subshell.

At copper (near the end of the transition series)  the 3d orbital energy has dropped so that 3d orbitals are actually lower in energy than 4s orbitals
This means the 3d¹⁰ 4s¹ configuration is of lower energy because it has more in 3d orbitals.

For the transition metal atoms, the total number of valence electrons equals the number of the column (group) in the periodic table (counting from the left).

For transition metal ions having charge ≥ +2, the number of d electrons equals the total number of valence electrons minus the charge on the ion.
This is because

Orbitals in the 3d and 4s subshells are of similar energy.
In transition metal atoms  the 4s subshell is of lower energy than the 3d subshell.
In transition metal ions of charge ≥ +2, 3d is of lower energy than 4s.
In transition metal ions of charge ≥ +2, all valence electrons in the d-subshell.
Therefore
Ni (Group 10) has 10 valence electrons and Ni²⁺ is d⁸
Fe (Group 8) has 8 valence electrons and Fe³⁺ is d⁵
Ti (Group 4) has 4 valence electrons and Ti³⁺ is d¹

*d and f – blocks elements show exceptional properties.

*d – block; the covalent radii of the elements decrease from left to half midway (to Cr) and the radii of the elements, from Cr to Cu are very close to one another until near the end where the size increase slightly, across a period; (Why?) because on passing from left to right, extra protons are placed in the nucleus and extra orbital electrons are added. The orbital electrons shield the nuclear charge incompletely (because shielding effect order s > p > d > f). Because of this poor shielding effect by d - electrons, the nuclear charge attracts all of the electrons more strongly: hence a contraction in size occurs.¹

            The radii of the elements, from Cr to Cu are very close to one another; this is due to the fact that the successive addition of d – electrons screens the outer electron (4s) from the inward pull of the nucleus. As a result of this, the size of the atom does not alter much in moving from Cr to Cu.²

            Near the end of the period, there is a slight increase in the atomic radii; this is due to the fact that near the end of the series, the electron - electron repulsions between added electrons in the same orbitals are greater than the attractive forces due to the increased nuclear charge. This result in the expansion of the electron cloud and thus the atomic radius increases.

Sc 1.44

Ti 1.32

V 1.22

Cr 1.17

Mn 1.17

Fe 1.17

Co 1.16

Ni 1.15

Cu 1.16

Zn 1.25

*The elements in the 3^rd group of the d-block show the expected increase in size Sc ----Y-----La.

*The elements in the 4^th to 12^th group of d – block the atomic radii of the first and second transition elements show expected increases but  radii of the second  and third transition series are almost same;  (why?) this is due to the fact in the atoms of the second transition series, the number of shells increases, their atomic radii are larger than those of the elements of the first transition series but in second and third transition metals between  lanthanum and hafnium there are 14 lanthanide elements are present, in which the antepenultimate 4f shell of electrons is filled. There is a gradual decrease in size of the 14- lanthanide elements from Ce to Lu. This is called lanthanide contraction. Thus the lanthanide contraction cancels almost exactly the normal size increases on descending a group of transition elements.

Ionic Radii:-

“The effective distance from the center of the nucleus of the ion upto which it exerts its influence on its electronic cloud is called ionic radii.”  Cation always has smaller radius and an anion is always larger than it parent atom.

*In case of isoelectronic ions, i.e., ions having same number of electrons but different nuclear charge, ionic radius decreases with increase in nuclear charge.

*Ionic radii follow the same trend as the atomic radii.

Oxidation states

The transition elements exhibit a large number of oxidation states. With the exception of few elements, most of these show variable oxidation states. These different oxidation states are related to the electronic configuration of their atoms. For example, the oxidation states exhibited by the transition elements of  the first series are listed in TABLE.

Different oxidation states of First transition series

Element		Outer electronic   configuration	Oxidation states
Sc	3d14s2		+2,+3
Ti	3d24s2		+2,+3,+4
V	3d34s2		+2,+3,+4,+5
Cr	3d54s1		(+1),+2,+3,(+4),(+5),+6
Mn	3d54s2		+2,+3,+4,(+5),+6,+7
Fe	3d64s2		+2,+3,(+4),(+5),(+6)
Co	3d74s2		+2,+3,(+4)
Ni	3d84s2		+2,+3,+4
Cu	3d104s1		+1,+2
Zn	3d104s2		+2

*Oxidation states with in the brackets are unstable.

Explanation

The existence of the transition elements in different oxidation states means that their atoms can lose different number of electrons. This is due to the participation of inner (n-1) d-electrons in addition to outer ns-electron because, the energies of the ns and (n-1) d-subshells are almost equal. For example, scandium has the electronic configuration of 3d¹4s². It exhibits an oxidation state of +2 when it uses both of its two 4s-electrons for bonding. It can also show oxidation state of +3 when it uses its two s-electrons and one d-electron. 

Important conclusion regarding oxidation states of transition elements

The examination of common oxidation states shown by different transition metals reveals the following facts :

(i)  The variable oxidation states of transition metals are due to participation of inner (n -1) d and outer  n s-electrons. The lowest oxidation state corresponds to the number of ns-electrons. For example, in the first transition series , the lowest oxidation states of Cr (3d⁵4s¹) and Cu(3d¹⁰4s¹) are +1 while for others , it is +2 (3d¹⁰4s²) 

(ii) Except scandium, the most common oxidation state of the first row  transition elements is +2 which arises due to loss of two 4s-electrons. This means that after scandium 3d-orbitals become more stable and therefore , are lower in energy than the 4s-orbitals. As a result, electrons first removed from 4s-orbitals.

(iii) For the first five elements, the minimum oxidation state is equal to the number of electrons in the s-orbitals and the other oxidation states are given by the sum of outer  s - and some or all d-electrons. The highest oxidation state is equal to the sum of the outer s (ns ) and (n -1) d-electrons. For the remaining five elements , the minimum oxidation state is given by the electrons in s-orbital while the maximum oxidation state is not related to their electronic configurations. The highest oxidation state shown by any transition metal is +8.

(iv) In the +2 and +3 oxidation states, the bonds formed are mostly ionic. In the compounds of higher oxidation states (generally formed wth oxygen and fluorine), the bonds are essentially covalent. Thus the bonds in +2 and +3 oxidation states are generally formed by the loss of two or three electrons respectively, while the bonds in higher oxidation states are formed by sharing of d-electrons. For example MnO₄^- (Mn in +7 ) state all the bonds are covalent.

(v)  Within a group, the maximum oxidation state increases with atomic number. For example, iron (Group 8) shows common oxidation states of +2 and +3 but ruthenium and osmium in the same group form compounds in the +4 , +6 and +8 oxidation states.

(vi) Transition metals also form compounds in low oxidation states such as +1 and 0 or negative. The common examples are [Ni(CO)₄], [Fe(CO)₅] in which nickel and iron are in zero oxidation state.

Complex formation tendency of transition metals:

Transition metal ions have very large vacant d- orbitals in their valence shells. As a result of this, lone pairs of electrons from ligands easily fill up the valence shells. When this happens, coordination compounds are formed. Note that this also takes place because d-orbitals have definite shapes and size as opposed to f-orbitals. That's why transition metals form coordinationcompounds while inner transition elements don't.

Magnetic properties of transition metals: 

The magnetic properties of a compound is a measure of the number of unpaired electrons in it. There are two main types of substances :

i) Paramagnetic substances : The substances which are attracted by magnetic field are called paramagnetic substances and this character arises due to the presence of unpaired electrons in the atomic orbitals. 

ii) Diamagnetic substances :  The substances which are repelled by magnetic field are called diamagnetic substances and this character arises due to the presence of paired electrons in the atomic orbitals.

            Most of the compounds of transition elements are paramagnetic in nature and are attracted by the magnetic field.

            The transition elements involve the partial filling of d-subshells. Most of the transition metal ions or their compounds have unpaired electrons in d-subshell (configuration from d¹ to d⁹) and therefore, they give rise to paramagnetic character. The magnetic character is expressed in Bohr magnetonsabbreviated as B.M. The magnetic moments of some ions of the first transition series are given in the following TABLE.

Magnetic moments of ions of first transition series.

Ion	Outer configuration	Number of unpaired electrons	Magnetic moment(mB)
Sc3+	3d0	0	0
Ti 3+	3d1	1	1.75
V3+	3d2	2	2.76
Cr3+	3d3	3	3.86
Cr2+	3d4	4	4.8
Mn2+	3d5	5	5.96
Fe2+	3d6	4	5.10
Co2+	3d7	3	4.4-5.2
Ni2+	3d8	2	2.9 - 3.4
Cu2+	3d9	1	1.8 - 2.2
Zn2+	3d10	0	0

The magnetic moments arise only from the spin of electrons. This can be calculated from the relation :

μ  = √n(n+2 B.M

Where n is the number of unpaired electrons and B.M represents Bohr magneton. It is clear from the Table that as the number of unpaired electron increases from 1 to 5, the magnetic moment and hence the paramagnetic character also increases. After d⁵ configuration, there is decrease in magnetic moment due to decrease in number of unpaired electrons. For example, d⁶ configuration has 4 unpaired electrons, d⁷ configuration has 3 unpaired electrons and so on. 

In addition to paramagnetic and diamagnetic substances, there are a few substances such as iron metal, iron oxide which are highly magnetic (about 1000 times more than ordinary metals). These are called Ferromagnetic substances.

Electronic Configuration of second transition series:

Second (4d) Transition Series (Y-Cd)

At. No.	39	40	41	42	43	44	45	46	47	48
Element	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd
Config.	4d15s2	4d25s2	4d45s1	4d55s1	4d55s2	4d75s1	4d85s1	4d105s0	4d105s1	4d105s2

I think this blog link contains all information related to your syllabus :) Click here =>

References:

1-http://hsc.csu.edu.au/chemistry/options/art/2769/ch984.htm#b1

2-http://www.ucc.ie/academic/chem/dolchem/html/comp/transmet.html

3-http://www.docbrown.info/page07/transition1.htm

4-http://www.newagepublishers.com/samplechapter/001210.pdf

5-http://www.bestchoice.net.nz/html/sa1/main/s148/p3960.htm

Walsh's Diagram
References:
1- http://mcgrady.chem.ox.ac.uk/lectures/bonding_in_molecule/lecture_3.pdf
2-http://www.pci.tu-bs.de/aggericke/PC4e/Kap_V/Walsh.htm
3-http://www.colby.edu/chemistry/PChem/notes/Walsh.pdf
4-http://www.chembio.uoguelph.ca/educmat/chm364_preuss/5_4%20Walsh%20Diagrams.pdf
5-http://www.dsf.unica.it/~sandro/Materiale_Corso/Molec/Demtroeder_polyatomic.pdf
6-http://users.physik.fu-berlin.de/~pascual/Vorlesung/SS06/Slides/AMOLI%20-%20L8a.pdf
7-

Determination of binary formation constants by: 

See Next => Determination of binary formation constants by pH-metry

Determination of binary formation constants by pH-metry (Irving and Rossotti method):http://notesgoogleblog.blogspot.in/2014/11/blog-post.html

See next => http://notesgoogleblog.blogspot.in/2014/11/blog-post.html

Click for other methods 
http://web.mst.edu/~tbone/Subjects/TBone/Datasheet.pdf

http://pac.iupac.org/publications/pac/pdf/1982/pdf/5412x2675.pdf

Chelate effect 

The chelate effect can be seen by comparing the reaction of a chelating ligand and a metal ion with the corresponding reaction involving comparable monodentate ligands. For example, comparison of the binding of 2,2'-bipyridine with pyridine or 1,2-diaminoethane (ethylenediamine=en) with ammonia.

Consider the following equilibria: 

For first reaction log K₁ = 10.6 ΔH°=-54 kJ mol^-1 ΔS°=+23 J K ^-1 mol^-1 while for second reaction log β₂ = 7.7 ΔH°=-46 kJ mol^-1 ΔS°=-8.4 J K ^-1 mol^-1

Although essentially the same copper-nitrogen bonds are formed in the two complexes, the formation of the en complex is distinctly favoured. Note in particular that the first reaction involves an increase in entropy. This is because two water molecules are released into solution when one en coordinates. The preference of chelate coordination is called the chelate effect and, as derived, is an entropic effect.

Thermodynamic origin of Chelate Effect:

It has been known for many years that a comparison of this type always shows that the complex resulting from coordination with the chelating ligand is much more thermodynamically stable. This can be seen by looking at the values for adding two monodentates compared with adding one bidentate, or adding four monodentates compared to two bidentates, or adding six monodentates compared to three bidentates.

A number of points should be highlighted from the formation constants in Table E4. In the first table, it can be seen that the ΔH° values for the formation steps are almost identical, that is, heat is evolved to about the same extent whether forming a complex involving monodentate ligands or bidentate ligands. What is seen to vary significantly is the ΔS° term which changes from negative (unfavorable) to positive (favorable). Note as well that there is a dramatic increase in the size of the ΔS° term for adding two compared to adding four monodentate ligands. (-5 to -35 JK^-1mol^-1). What does this imply, if we consider ΔS° to give a measure of disorder?

In the case of complex formation of Ni²⁺ with ammonia or 1,2-diaminoethane, by rewriting the equilibria, the following equations are produced.

Using the equilibrium constant for the reaction (3 above) where the three bidentate ligands replace the six monodentateligands, we find that at a temperature of 25° C: 

  ΔG∘=−2.303RTlog10K

=−2.303RT(18.28−8.61)

=−54 kJ mol−1

Based on measurements made over a range of temperatures, it is possible to break down the ΔG∘ term into the enthalpy and entropy components.

ΔG∘=ΔH∘−TΔS∘

The result is that: ΔH∘=−29kJmol−1

- TΔS° = -25 kJ mol^-1
     and at 25C (298K)
ΔS° = +88 J K^-1 mol^-1

Note that for many years, these numbers have been incorrectly recorded in textbooks. For example, the third edition of "Basic Inorganic Chemistry" by F.A. Cotton, G. Wilkinson and P.L. Gaus, John Wiley & Sons, Inc, 1995, on page 186 gives the values as:

ΔG° = -67 kJ mol-1
ΔH° = -12 kJ mol-1
-TΔS° = -55 kJ mol-1

The conclusion they drew from these incorrect numbers was that the chelate effect was essentially an entropy effect, since the TΔS° contribution was nearly 5 times bigger than ΔH°.

In fact, the breakdown of the ΔG° into ΔH° and TΔS° shows that the two terms are nearly equal (-29 cf. -25 kJ mol^-1) with the ΔH° term a little bigger! The entropy term found is still much larger than for reactions involving a non-chelating ligand substitution at a metal ion. How can we explain this enhanced contribution from entropy? One explanation is to count the number of species on the left and right hand side of the equation above.

It will be seen that on the left-hand-side there are 4 species, whereas on the right-hand-side there are 7 species that is a net gain of 3 species occurs as the reaction proceeds. This can account for the increase in entropy since it represents an increase in the disorder of the system.

An alternative view comes from trying to understand how the reactions might proceed. To form a complex with 6 monodentates requires 6 separate favorable collisions between the metal ion and the ligand molecules. To form the tris-bidentate metal complex requires an initial collision for the first ligand to attach by one arm but remember that the other arm is always going to be nearby and only requires a rotation of the other end to enable the ligand to form the chelate ring.

If you consider dissociation steps, then when a monodentate group is displaced, it is lost into the bulk of the solution. On the other hand, if one end of a bidentate group is displaced the other arm is still attached and it is only a matter of the arm rotating around and it can be reattached again.

Both sets of conditions favor the formation of the complex with bidentate groups rather than monodentate groups.  =>See Previous Post Also 

Stability of metal complexes with reference to thenature of metal ion :

The Irving-William series:

This series is obtained from a plot of long Kf against Z for the first-row transition elements. For any one set of ligands, i.e., no change in ligand type, the series shows a regular increase in stability constants from Sc to Cu, and a decrease to zinc. 

Ba²⁺ < Sr²⁺ < Ca²⁺< Mg²⁺ < Mn²⁺ < Fe²⁺ < Co²⁺ < Ni²⁺ < Cu²⁺ > Zn²⁺ 

This series can be understood as a consequence of electrostatic effects, of the first part of this series. Ligand-field interactions explain the increase in the value of K_f for Fe (II), d⁶; Co (II), d⁷; Ni (II), d⁸; and Cu (II), d⁹ – Cu (II) obtains increased stability from Jahn-Teller distortion when compared to Ni (II) complexes.
Stability of metal complexes with reference to the nature of  ligand:

 =>See Previous Post Also 
For detail study Click on these links: 
1- http://web.uvic.ca/~djberg/Chem324/Chem324-14.pdf
2 -http://bit.ly/10OzNcK (Important Link)

Trends in successive formation constants:

Normally, we will see for a series of stepwise constants the numerical order: 

K₁ > K₂> K₃ > …..> K_n

As a reaction proceeds fewer ligands are available in solution to coordinate, so this order is further enhanced. However, sometimes one of the higher-order constants is found to be significantly greater than an earlier one. This usually means that some significant structural change has occurred, such as the transition from high-spin to low-spin configurations. This occurs, for example in the step: 

The constant for this step is much greater than for the first two ligand-attachment reactions,accompanying the change from high spin d⁶ to low-spin d⁶. Another example is the conversion of Cd²⁺ to [CdBr₄]^2-, involving the change from octahedral to tetrahedral geometry. =>See Previous Post Also 

SEMESTER I

CHE5101 Inorganic Chemistry

Metal-LigandEquilibria in Solution 

Stability of co-ordination compounds 

The stability of compounds means in a most general sense, the compounds exist under suitable conditions may be stored for a long period of time. However when the formation of complexes in solution is studied, two types of stabilities, thermodynamic stability and kinetic stability are considered. In the language of thermodynamics, the equilibrium constants of a reaction are the measure of the heat released in the reaction and entropy change during reaction. The greater amount of heat evolved in the reaction, the most stable are the reaction products. Secondly, greater the increase in entropy during the reaction, greater is the stability of products. The kinetic stability of complexes refers to the sped with which transformation leading to the attainment of equilibrium will occur. Here we are mainly concerned with the thermodynamic stability of the complex compound. 

Thermodynamic stability vs. kinetic stability and their correlations: 

The most stable states of a kinetic reaction are those of the reactants, in which an input of energy is required to move the reaction from a state of stability, to that of reacting and converting itself to products. Kinetics is related to reactivity. In contrast, the most stable state of a thermodynamically favorable reaction is the products, because the reaction occurs spontaneously, without the need for energy to be added. Thermodynamics is related to stability.

Therefore, something that is unreactive will desire to stay in the form of reactants, which will require an input of energy to cause the reaction to go forward, converting reactants into products. This is illustrated in example #3 below. A reactive species does not require an input of energy to be converted from reactants to products, because its most stable and preferred state is that of the products. Instead, a thermodynamically favorable reaction requires energy to be converted from products back to reactants.An energy source moves the reaction forward (kinetics corresponds to movement). The same is for thermodynamically favorable reactions, except that the reaction must be stimulated backward from products to reactants. 

Cumulative or overall and stepwise constants: 

A cumulative or overall constant, given the symbol β, is the equilibrium constant for the formation of a complex from reagents. For example, the cumulative constant for the formation of AH₂ is given by

The stepwise constant, K, for the formation of the same complex from A^2- and H⁺ in two steps is given by:

For 1st equation:.........................

For 2nd equation.........................

Therefore β₁₂ = K₁K₂ and log β₁₂ = log K₁ + log K₂. For the dibasic acid log β₂ = pK₁ + pK₂. A cumulative constant or overall constant can always be expressed as the product of step-wise constants. Each stability constant should be defined by reference to an equilibrium expression, as in the examples above.

Further reading:

http://shodhganga.inflibnet.ac.in/bitstream/10603/3975/11/11_section%20c.pdf
http://www.hyperquad.co.uk/documents/HQfaq.pdf
http://chemwiki.ucdavis.edu/Physical_Chemistry/Equilibria/Chemical_Equilibria/Principles_of_Chemical_Equilibria/Kinetically_vs_Thermodynamically_Stable
http://classes.uleth.ca/200403/chem3820a/3820%20lecture%20chapter_9_part1_2004.pdf

Structure and Hybridization:

How we can understand this? “Easy” by solving question based on it.

SP

Example 1^st: BeF_2 (sp - Hybridization₎

Write valence shell electronic configuration.

Be: 1s², 2s²

1 - Write orbital configuration of valence shell only.

2 - Excite electron of 2s orbital into vacant p - orbital.

3 - 2s + 2p, having equal energy, now they are Hybrid orbital :) 

4 - Hybrid orbital forms covalent bond with surrounded atoms such as F. 

Thus, sp - Hybridization in BeF₂

*What happen with 1s² orbital electrons?

It remains as non-bonded electrons i.e. lone pair. Thus in this example one lone pair is present in 1s orbital.

*What causes excitation of electron?

Presence of two F – atom causes this excitation.

*How F – atoms combine with this sp – Hybrid orbital?

C₂H_{2  (}sp - Hybridization₎

Reference: http://www.mhhe.com/physsci/chemistry/carey5e/Ch02/ch2-3-3.html

C₂H₄ (sp² Hybridization)

Reference: http://www.mhhe.com/physsci/chemistry/carey5e/Ch02/ch2-3-2.html