Transition Metals

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  • Created on: 25-04-14 00:00
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Transition Metals and the d-Block Elements ­ Answers
1. In an atom of a d-block element, electrons go into the 4s subshell before the 3d subshell. When a d-block atom
forms an ion electrons are lost from the 4s subshell first before losing from the 3d subshell. The general
configuration for a d-block element is [Ar] 3dx 4s2, where x is the number of the column along the d-block.
2. In the fourth period, there are two exceptions to the 4s first rule. There is a slight gain in stability in having a full
or half-full set of d-orbitals. Thus chromium has the electron configuration [Ar] 3d5 4s1; copper has the
configuration [Ar] 3d10 4s1.
3. The d-block elements are those between the s-block and the p-block. All d-block elements have an outer
electron configuration of ndx, where x is any number from 1 to 10.
4. A transition metal has one or more unpaired d-electrons in one of its ions. E.g., titanium forms a Ti3+ ion, which has
one unpaired d-electron, and Ti2+, which has two unpaired electrons. Therefore, titanium is a transition metal.
5. Across the third period, the first ionisation energies generally
increase across, due to the increase in nuclear charge, without
an increase in the number of shielding electrons. However, in the
d-block, the outer electron that is removed in the formation of
an ion is a 4s-electron, and it is shielded by the inner
3d-electrons. Although the nuclear charge increases across the
block from scandium to zinc, the number of inner shielding
3d-electrons increases as well. This means that the first
ionisation energies of the d-block elements in period 4 are fairly
6. There is normally a big jump between successive ionisation
energies as a new quantum shell loses an electron. This electron is subjected to a much stronger pull from the
nucleus as it is much less shield. However, the energy of the 3d- and 4s- electrons are very similar in d-block
elements and so the big jump comes after all the 4s- and 3d- electrons have been removed.
7. Metallic bonding can be described as the attraction between the delocalised electrons and the positive ions
formed when the metal atoms lose their valence electrons into the cloud of delocalised electrons. The strength
of the bond determines the melting temperature of the metal. In the d-block metals, both the d-electrons and
s-electrons are used in bonding. Therefore, the melting temperatures of d-block metals are significantly higher
than those of s-block metals. The metals are also much harder and stronger.
8. The metals in the s-block and the non-transition d-block metals exist in only one oxidation state in their
compounds. However, the transition metals exist in different oxidation states, e.g. manganese has stable
compounds such as MnSO4 (+2), MnO2 (+4), K2MnO4 (+6) and KMnO4 (+7).
9. In most compounds in the +2 and +3 oxidation states, the transition metals are ionically bonded, but when the
transition metal is in an oxidation state of +4 or higher, it is covalently bonded, often in an anion. E.g. the
manganese atom in MnO4- ion is covalently bonded to the four oxygen atoms by one single and three double
10. The extra energy required to remove a third electron from an Fe2+ ion to form Fe3+ is small enough that it can be
recovered either from lattice energy or from hydration energy of the cation. However, this is not the case with
calcium, where removal of a third electron would have to be from an inner shell, requiring energy that cannot be
11. When d-block cations are dissolved in water, they become hydrated. The oxygen atom in a water molecule has a
lone pair of electrons that forms a dative covalent bond with an empty 3d- or 4p- orbital in the metal ion. E.g. the
hydrated chromium (III) ion, which has the formula [Cr(H2O)6]3+. Six dative covalent bonds form, so the hydrated
ion has the coordination number 6. This ion is known as a complex ion.
12. The coordination number is the number of near neighbouring atoms that are bonded to the central ion. E.g. In
[Cr(H2O)6]3+, the Cr3+ ion has six oxygen atoms as near neighbours.
13. The shape of a complex ion can be predicted using valence-shell electron-pair repulsion theory. E.g. in
[Cr(H2O)6]3+there are six dative bonds, each containing a pair of electrons. These six pairs of bonding electrons
repel each other to the position of minimum repulsion, which is also the position of maximum separation. The
shape therefore is octahedral.
14. Monodentate ligands use one lone pair of electrons to form a dative covalent bond with the d-block ion:

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Ammonia, NH3, and organic amines such as ethylamine form complexes. E.g. the complex of Cu2+ and
ammonia, which has the formula [Cu(NH3)4(H2O)2]2+.
Anions such as Cl- and CN- also form complexes, e.g.
[Fe(CN)6]4- is a complex between an Fe2+ ion and six CN- ions.
[CrCl4]- is a complex between a Cr3+ ion and four Cl- ions it is energetically unfavourable to fit six large
ligands around a small cation.…read more

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The energy difference between these two sets of d-orbitals in a typical complex ion is equal to the energy
of a photon in the visible region of the spectrum.
When white light shines through a solution of a complex ion of a transition metal, photons of a particular
frequency are absorbed and their energy promotes an electron from the lower energy level to the upper
energy level. This is called a d-d transition.…read more

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Chromium is unstable in the +2 state compared with the +3 state. The +2 state can be stabilised by formation of a
complex with ethanoate ions, CH3COO-. The hydrated chromium (II) ions undergo ligand exchange and a
precipitate of red chromium (II) ethanoate complex is formed:
2[Cr(H2O)6]2+ (aq) + 4CH3COO- (aq) [Cr2(CH3COO)4(H2O)2] (s) + 10H2O (l)
36.…read more

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E.g. Cu2+ + e- Cu+ E = +0.15 V
45. The feasibility of a redox reaction is indicated by the value of the standard potential of the cell, E cell. Feasibility
is worked out by:
1) Reverse the half equation that has one of the reactants on the RHS and changing the sign of its E value.
2) This half-equation is then added to the half-equation that has the other reactant on the LHS.…read more


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