Melting and Boiling Points
The transition metals have very high melting and boiling points.
The melting points of these metals rise to a maximum value and then decrease with increase in atomic number. However, manganese and technetium metals have abnormally low melting points.
Explanation: The high melting and boiling points of these metals are due to strong metallic bonds between the atoms of these elements. These metals have high enthalpies of atomization. The metallic bond is formed due to the interaction of electrons in the outermost orbitals The strength of bonding is roughly related to the number of unpaired electrons. In general, greater the number of valence electrons, stronger is the metallic bonding and consequently melting points are high.
Therefore, as we move along a particular series, the metallic strength increases up to the middle with increasing availability of unpaired electrons up to d5 configuration (eg., Sc has 1, Ti has 2, V has 3, Cr has 5 unpaired electrons) and then decreases with decreasing availability of unpaired d-electrons ( e.g.: Fe has 4, Co has 3 unpaired electrons and so on). Therefore, the melting points decrease after the middle because of increase of pairing of electrons.
Since the alkali and alkaline earth metals have only one or two outer electrons available, their melting points are relatively low in comparison to transition metals. Similarly, the elements of group 12 (zinc, cadmium and mercury) are quite soft with low melting points. Mercury is a liquid at room temperature and melts at – 38°C. These three elements behave typically because there are no unpaired electrons available for metallic bonding and, therefore, their melting points are low.
a) The first ionisation enthalpies of d-block elements are higher than those of s-block elements but are lesser than those of p-block elements. The ionisation enthalpies increases as we move across each series, though not quite regularly.
Explanation : The increase in ionisation enthalpy along a given transition series is due to the effect of increasing nuclear charge which would tend to attract the outer electron cloud with greater force. Ionisation enthalpy is expected to increase. However in case of transition elements , the addition of electron takes place to the last but one i.e. (n-1) d-subshell and this also increases the screening effect. With the increase in the electrons in (n-1) d-subshell , the outer electrons in ns-subshell are shielded more and more. Thus the effect of increasing nuclear charge is opposed by the additional screening effect of the nucleus and consequently , ionisation enthalpy increases from left to right but quite slowly among d-block elements.
The irregular trend in the first ionisation enthalpy of the first transition series elements is due to the fact that the removal of one electron alters the relative energies of 4s and 3d orbitals.There is a reorganisation of energy accompanying ionisation with some gains in exchange energy as the number of electrons increases in the dn configuration and from the transference of s-electrons into d-orbitals.
The value of chromium is lower because loss of on electron gives stable configuration (3d5).Zinc has high ionisation enthalpy because electron has to be removed from 4s orbital of stable (3d10 4s2) configuration.
The second and third ionisation enthalpies also increase along a period. However, the magnitude of increase in the second and third ionisation enthalpies for the successive elements, is much higher.
(a) Chromium and copper have exceptionally high ionisation enthalpy values than those of their neighbours. These exceptions are attributed to the extra stability of half filled and completely filled set of d-orbitals in chromium (3d5) and copper (3d10) respectively. After the loss of first electron, Cr and Cu acquire a stable configuration (3d5 and 3d10) and the removal of second electron disrupts the stability with considerable loss of exchange energy.
(b) The value of second ionisation enthalpy for zinc is correspondingly low because the ionisation involves the removal of an electron resulting stable 3d10 configuration.
(c) The third ionisation enthalpies is not complicated by the 4s orbital factor and shows high values for Mn2+ and Zn2+ because of stable( 3d5 and 3d10 ) electronic configurations. Similarly, the third ionisation enthalpy of Fe is very small because loss of third electron results in stable 3d5 configuration (IE3 for Fe < IE3 for Mn).
The third ionization enthalpies are quite high. There is break between the values for Mn (II) and Fe (II). Moreover, the high values for copper, nickel and zinc indicate that it is difficult to obtain oxidation state more than two for these elements.
The first ionisation enthalpies of third transition series are higher than those of first and second transition series.
Explanation : In the atoms of third transition series, there are filled 4f-orbitals. The 4f-orbitals have very poor shielding effect. As a result, the outer electrons have greater effective nuclear charge acting on the outer valence electrons. Therefore, their ionisation energies are higher.
The transition metals exhibit a large number of oxidation states in their compounds. 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 8.8 ahead.
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-electrons because, the energies of the ns and (n — 1) d-subshells are almost equal. For example, scandium has the electronic configuration 3d1 4s2 . It exhibits an oxidation state of +2 when it uses both of its 4s-electrons for bonding. It can also show oxidation state of +3 when it uses its two s-electrons and one d-electron. Similarly, the other atoms can show oxidation states equal to ns and (n — 1) d-electrons.
|Ti||+2 , +3 , +4|
|V||+2 , +3 , +4 , +5|
|Cr||+2 , +3, +4 , +5 , +6|
|Mn||+2 , +3, +4 , +5 , +6 , +7|
|Fe||+2 , +3, +4 , +5 , +6|
|Co||+2 , +3 , +4|
|Ni||+2 , +3 , +4|
|Cu||+1 , + 2|
Variable oxidation states of second and third transition series
The elements of second and third transition series also exhibit variable oxidation state.
Oxidation state of Second transition series
|Zr||+3 , +4|
|Nb||+2 , +3 , +4 , +5|
|Mo||+2 , +3 , +4 , +5 , +6|
|Tc||+ 2 , +4 ,+5|
|Ru||+2 , +3 , +4 , +5 , +6 , +7 , +8|
|Pd||+ 2 , +3 , +4|
|Ag||+1 , +2 + 3|
Oxidation State of third transition series
|Hf||+3 , +4|
|Ta||+2 , +3 , +4 , +5|
|W||+2 , +3 , +4 , +5 , +6|
|Re||-1 , +1 , +2 , +4 , +5 , +6 , +7|
|Os||+ 2 , +3 , + 4 , + 6 , +8|
|Ir||+2 , +3 , +4 , +6|
|Pt||+ 2, + 3 , +4, +5 , +6|
|Au||+ 1 , +3|
|Hg||+1 , +2|
The stability of a given oxidation state depends upon the nature of the elements with which the metal is combined. The highest oxidation states are found in compounds of fluorides and oxides because fluorine and oxygen are most electronegative elements.
a) The variable oxidation states of transition metals are due to participation of inner (n-1) d and outer ns 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 (3d1-104s2).
b) Except scandium, the most common oxidation state of the first row transition elements is +2 which arises due to loss of two 4s-electrons..After scandium 3d-orbitals become more stable and, therefore are lower in energy than the 4s-orbitals. As a result, electrons are first removed from4s-orbitals.
c) The elements which show the greatest number of oxidation states occur in or near the middle of the series.
For example: In the first transition series Manganese exhibits all the oxidation states from +2 to +7. The small number oxidation states at the extreme left hand side end (Sc, Ti) is due to lesser number of electrons to lose or share. At the extreme right hand side end (Cu, Zn), it is due to large number of d-electrons so that only a fewer orbitals are available in which the electron can share with others for higher valence.
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.
d) Scandium (II) is virtually unknown and titanium (IV) is more stable than Ti(III) and Ti(I). The only oxidation state of zinc is +2 in which no d-electrons are involved.
e) In the +2 and +3 oxidation states, the bonds formed are mostly ionic.In the compounds of higher oxidation states (generally formed with 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: In MnO–4 (Mn in +7) state all the bonds are covalent.
f) 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.
g) Transition metals also form compounds in low oxidation states such as +l and 0 or negative. The common examples are [Ni(CO)4], [Fe(CO)5] in which nickel and iron are in zero oxidation state.
h) The variability of oxidation states in transition elements arises because of incomplete filling of d-orbitals in such a way that their oxidation states differ by unity such as VII, VIII, VIV and VV. The variability of oxidation states of non-transition elements (p-block elements), where oxidation states normally differ by a unit of two such as Sn2+, Sn4+, In+ , In3+, etc.
i) Unlike p-block elements where the lower oxidation states are favoured by heavier members (due to inert pair effect), the higher oxidation states are more stable in heavier transition elements. For example: In group 6, Mo (VI) and W (VI) are found to be more stable than W (VI). Therefore, Cr (VI) in the form of dichromate in acidic medium is a strong oxidising agent whereas MoO3, and WO3, are not.
j) The transition elements show low oxidation states in some compounds or complexes having ligands such as CO, which not only form sigma bonds with the metal atoms but also have t-acceptor character .