Download Table showing examples of Complex ions with their bond

1
Metals and Non-Metals
Definitions and Periodic trends
Electronegativity is the power of a chemically bonded atom to attract electrons to itself.
Electronegativity decreases down the group but increase across a period due increased distance between the valence electron
and the nucleus i.e., greater atomic radius. Example of an electropositive (low electronegativity) element is Caesium (0.86)
and e.g. of high electronegative element is Fluorine (4.10), values in bracket are based on Aured-Rochow scale of
electronegativities.
Ionisation Potential (Energy) (I) is energy required to remove an electron from the system (atom, molecule, or ion) and
transport it to ∞ (infinity- because there has to be no attraction between electron and ion). The more closer the electron (i.e.
tightly bound and difficult to remove) to nucleus the higher its ionisation energy. The value of Ionisation energy (I) depends
on the orbital containing the electron and therefore on the quantum number of the shell (n). Greater the value of n, lower the
value of I; i.e. K (most difficult to remove) < L< M < N etc in a given shell or in terms of orbitals; i.e. s (most difficult to
remove) < p < d < f (most easily to remove) etc.
Alkali metals (group IA) have single electron in the outer orbital with higher value of n and therefore low I.
Noble gases with stable octet of electrons have high value of I. Group IIIA elements i.e. B, Al, have lower 1 st Ionisation energy
because electron removed is in p orbital (l= 1) not an s orbital (l=0), where l is the quantum number of an orbital (this is
discussed in chapter 3). Group VIA (6A) elements have rather lower I value.
Outer electron configuration of ns2 np3 with half filled p orbitals, i.e. ns2 np3 configuration has a special stability.
Ionisation potentials (energies) decrease down the group and increase across the period.
Ionisation Energy: Is the energy required to remove an electron in its ground state from an atom to infinity. Quantitatively
measured in kJmol—1
M(g) = M+ (g) + e— H = kJmol— ( - Change in, H – energy).
The first ionisation energy is the energy required to remove one electron from the parent atom. The second ionization energy
is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on.
Second ionisation energy > first ionisation energy
Moving from left to right across a period (decreasing atomic radius) - Ionisation energies increase
Moving down a group (increasing atomic radius) - Ionisation energy decreases.
Group IA elements have low ionisation energies because the loss of an electron forms a stable octet.
Oxidation: is the complete or partial removal of electrons from an atom, molecule or ion; i.e. oxidation is loss of electrons.
Oxidation is also addition of oxygen to an element or compound, or removal of hydrogen from the compound.
Oxidation is a reaction in which the oxidation number of an element is increased.
Reduction: is the complete or partial addition of electrons to an atom, molecule or ion; i.e. reduction is gain of electrons.
Reduction is also addition of hydrogen to an element or compound, or removal of oxygen from the compound.
Reduction is a reaction in which the oxidation number of an element is decreased.
Oxidising and Reducing agents.
An oxidising agent (or oxidant) is an atom or molecule or ion, which readily captures electrons.
A reducing agent (or reductant) is an atom or molecule or ion, which readily releases electrons.
Chemical reaction taking place in galvanic cells or during electrolysis are simple electron- transfer reactions, the function of
the metallic conductor (‘external circuit’) being merely to link the two electrodes.
Continued
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Transition Elements and structure
It is the availability of the d electrons, as well as the s electrons, which gives the transition elements their peculiar
characteristics;
i)
ii)
iii)
iv)
v)
vi)
vii)
Similarity of physical properties
Wide variation of oxidation state
Formation of coloured ions
Paramagnetism of ions
Catalytic activity
Formation of Interstitial compounds
Ready formation of Stable complexes
The d-block elements generally exhibit the characteristic properties of metals because the outer electrons (d) are loosely
bound and readily enter into metallic bond formation.
The metallic radius decreased in passing from Sc to Ni. Addition of electrons might be expected to result in an increase in
radius, but the electrons are being added to an inner orbital and it is the increased in nuclear charge in passing from Sc to Ni
that causes the decrease in radius. The decrease in radius coupled with as increasing relative atomic mass causes an increase
in density and a decrease in atomic volume in passing from Sc to Ni. There is increase in electronegativity in passing from Sc
to Ni.
Oxidation State (O.S.) With the exception of the end members of the series, the elements exhibit 2 or more oxidation states
due to the availability of unpaired ‘d’ electrons.
e.g. the Fe atom, has configuration; 1s 2 2s2 2p6 3s2 3p6 3d6 4s2 and may lose two 4s electrons to form the Fe2+ ion (O.S. of
+2). Under certain circumstances a further electron may be lost form the 3d sub-level, to give the Fe3+ ion (O.S. of +3). Also
‘d’ electrons may be used for covalent bonding as in the ferrates (e.g. K 2+ FeO42— ), where oxidation state is +6.
Mn, with the highest number of unpaired electrons shows wide range oxidation states and it is the only metal to exhibit an
oxidation state of +7.
Continued
3
Coloured ions All except Sc group and Zn group, d-block elements have coloured ions.
This is because each ion absorbs light of certain wavelengths only in the visible past of the spectrum, thus changing incident
white light into light whose hue () is composed of the colours complementary those which have absorbed. The d-level is split
into two when a complex ion is formed and difference in energy between these two levels is such that visible light can excite
an electron from the lower to the higher level; for instance Cu(H2O)42+and Fe(CN)63— absorb visible light. This is in contrast to
the electronic energy levels of ions of other elements, in which the energy levels are separated by distances such that
promotion of electrons from one level to another is only achieved by absorption of energy from the ultra-violet part of the
spectrum.
Colour of hydrated ions:
Complex ions Cations of the d-block elements are small and highly charged and therefore readily attract polar molecules or
ions to form complexes,
E.g. Hexacyanoferrrate (II) ion, Fe(CN)6 4— and [Fe(CN)6 3—]
Tetraammine Copper(II)ion, Cu(NH3)42+
Tetrahydroxozincate, Zn(OH)42—
Hexachlorostannate (IV), Sn Cl62—
Tetraaquocopper(II), Cu(H2O)42—
Diamminesilver(I), Ag(NH3)2+
HexaammineNickel (II), Ni(NH3)62+
Examples of Ligands
Fluoro F—
Chloro Cl—
Cyano CN—
Hydroxa OH—
Aquo H2O
Ammine NH3.
Carbonyl CO
Complex Anion. E.g. Si F62— Hexa fluoro Silicate(IV)
Complex Cation. E.g. Co(NH3)63+ - Hexa ammine Cobalt(III)
Neutral Complexes: If the charge on the central ion is equal and opposite to the charge on the ligand, a neutral complexes
results.
E.g. Cr (NH3)33+ Cl3— , Co(N3H)22+ Cl2—
Example of Soluble complex
Substance insoluble in H2O and soluble in NH3 solution is AgCl [Silver Chloride] and the complex formed is Ag(NH3)2+
[Diamminesilver(I)]. Complex is used in Photography.
Hybridisation in Complexes
Bond hybridisation
Geometrical Shapes
Example
d2sp3
Octahedral
Fe(CN)6 3—
sp3d2
Octahedral
Fe(F)6 3—
sp3
Tetrahedral
Zn(NH3)42+
TetraammineZinc (II) ion
dsp2
Square
Ni(CN)42—
Nickel tetracyanate (II) ion
Continued
Table showing examples of Complex ions with their bond hybridisation and geometrical shapes.
See Metals and non-metals reading lists for more details of structures