THE YARDS KNEW.

TRANSITION METALS

INTRODUCTION

A transition element (or d-block element) is one which forms some compounds in which there is an incomplete sub-shell of d electrons.

i.e. a transition element forms at least one ion with a partially-filled d-sub-shell.
The electronic configurations of these ten elements (the main transition elements) are basically similar.
i.e. [Ar] 3d^x 4s^y, where [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶
Electrons occupy the 4s orbitals first before filling the 3d orbitals. The 4s electrons are also removed first in the formation of positive ions.

For example, a typical transition element of the d block is iron, with the electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² or [Ar] 3d⁶ 4s², where 3d⁶ is an incomplete sub-shell of shell 3. Thus, the 4s orbital is filled with electrons before the 3d orbitals as it is at lower energy than the 3d orbitals. Once the electrons are actually in their orbitals, the energy order changes and the 4s orbital then behaves as the outermost, highest energy orbital. So, the 4s electrons are lost first when iron forms cations. Fe usually forms 2 ions, Fe²⁺ whose electronic structure is [Ar] 3d⁶ and Fe³⁺ whose electronic configuration is [Ar] 3d⁵ .

For Cr, the configuration [Ar] 3d⁵ 4s¹ is more stable than [Ar] 3d⁴ 4s².
For Cu, the configuration [Ar] 3d¹⁰ 4s¹ is more stable than [Ar] 3d⁹ 4s².

Both scandium (Sc) and zinc (Zn) are not typical transition metals because
- They have only one known oxidation state: Sc³⁺ and Zn²⁺ only.
- Their ions do not have partially filled d-sub-shell:
Sc³⁺ [Ar] 3d⁰
Zn²⁺ [Ar] 3d¹⁰
- The density of Sc (3 g cm^-3) is much lower than the others (mostly 7-8 g cm^-3); the melting point of Zn (420^oC) is much lower than the others (>1500^oC)

They are very similar in their physical properties due to the relatively small difference in effective nuclear charge.
This is because each additional electron enters the penultimate 3d shell where it provides a more effective shield between the nucleus and the outer 4s shell of electrons. Hence, although each successive nucleus has one more proton, this extra positive charge is partly shielded from the outer 4s electrons by the extra electron in an underlying 3d shell.
For iron, the electronic configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

electronic configuration: Iron

For Cr, the configuration [Ar] 3d⁵ 4s¹ is more stable than [Ar] 3d⁴ 4s².

Electronic configuarion: Chromium

For Cu, the configuration [Ar] 3d¹⁰ 4s¹ is more stable than [Ar] 3d⁹ 4s².

electronic configuration: Copper

d-BLOCK ELEMENTS AS METALS
The d-block elements can be said to be typical metals, being good conductors of heat and electricity, hard, strong, malleable, ductile, lustrous, and silver-white in colour, and generally they have much higher melting point and boiling point than the main group elements. However, there are some notable exceptions: Mercury has such a low melting point that it is a liquid at room temperature; copper is red-brown in colour, and gold is yellow.
These general physical properties of d-block metal, together with their fairly low chemical reactivity, make transition metals extremely useful as structural metals.

For example, iron is the most important structural metal ever used. Its great advantage of being much cheaper to produce outweighs its great disadvantage of suffering form worse corrosion than other more expensive transition metals. Another advantage is that it can be converted to steel, which is harder and has superior resistance to corrosion.

Titanium has a larger atomic radius than iron, and is therefore less dense. It does not corrode as iron does. Its combination of mechanical strength (similar as steel) and low density makes it attractive for use in aircraft components. The high cost of titanium has however limited its use to applications where no expense is spared. It is used in the construction of space shuttles, being batter able than steel to withstand the high temperature that are experienced when a space shuttle re-enters the earth’s atmosphere.

The similar atomic radii of the transition metals make it possible for atoms of one element to replace those of another element in the formation of alloys. The solid structure is slightly deformed but this is sufficient to modify the physical properties of the structural material, e.g. manganese is well-known for conferring hardness and wear resistance to its alloys, and chromium is responsible for conferring passivity on stainless steels. It is possible to make alloys containing transition elements in a wide range of composition, as a result of their similar atomic radii.

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COMPARISON OF SOME PHYSICAL AND CHEMICAL PROPERTIES OF d-BLOCK AND s-BLOCK METALS

Comparison of Melting Point and Hardness between d-Block and s-Block Metals

Melting point:

Transition elements have very high melting and boiling point due to the strong metallic bonding (since electrons from both 3d and 4s orbitals can contribute to the delocalized ‘sea of mobile electrons’ for metallic bonding). s-block elements (alkaline metals and alkaline earth metals) only have valence electrons from s orbitals for metallic bonding.

Explanation: The high melting points and the hardness of transition metals are caused by their strong metallic bonds. Metallic bonding is caused by the delocalization of valence electrons, leaving positive metal ions surrounded by a sea of mobile electrons. The more delocalized valence electrons and the smaller the radius of the ions, the stronger the metallic bond. Thus, across the period, the metallic bonds of transition metals are stronger than that of s-block elements due to the smaller atomic radii and the larger number of valence electrons.
More specifically, their strong metallic bonds result from their ability to release electrons from both outer and inner shells for bonding. For example, metals in the first series use both 3d and 4s electrons for bonding. This explains why they have higher melting points compared with s block elements, such as calcium (with lesser delocalized electrons, calcium has a weaker metallic bond and thus lower melting point).

Exception: Manganese has relatively low melting point due to stability of d⁵ arrangement of ½ filled shells and to the lower availability of such electrons for delocalization.
e.g. iron (d-block element) melts at 1538^oC and boils at 2860^oC, while calcium (s-block element) melts at 839^oC and boils at 1484^oC.

Comparison of Atomic Radii and Densities between d-Block and s-Block Metals

Density

Transition metals are hard and dense (except Ti (4.5 g cm^-3), Sc (3 g cm^-3) and Y) due to relatively small metallic radii (hence, small size) due to the increase in electronegativity and closed-packed structures in the metallic lattices.
Explanation: The almost constant metallic radius coupled with increased relative atomic mass also could account for this. The d-block elements are generally denser than the s-block elements. From Sc to Cu, there is a gradual increase in density.
In fact, all transition metals have a density greater than 5 g cm^-3. The densities of transition elements in period 5 are higher that in period 4, but lower than that in period 6. The transition element with highest density is iridium 22.61g cm^-3.
e.g. density of iron = 7.9 g cm^-3, density of calcium = 1.55 g cm^-3.

Atomic radius; ionic radius:

There is a little variation in atomic radii due to little variation in effective nuclear charge across series.

Explanation: Each electron enters the 3d shell where it provides more effective shielding between nucleus and the 4s valence electrons. This nullifies the effect of the additional proton in the nucleus.
The ionic radius of M²⁺ and M³⁺ show a decrease across the series. Because the 4s electrons are removed and the 3d sub-shell is now the outermost (the number of shells decreases) and the increase in effective nuclear charge is significant across the series of transition ions, it results in a decrease in ionic radius.
The s-block elements, such as calcium, have one or two valence electrons in orbitals outside filled shells of electrons. These filled shells are reasonably effective at shielding the valence electrons from the attraction of the nucleus. As a result, the atomic radii of the s-block metals are generally about 50% greater than those of the transition metals.
The atomic radius of the s-block metals increases from lithium to caesium (Group I) and from beryllium to barium (Group II). Each successive element in a group has an extra filled shell of electrons.

Comparison of Ionization Enthalpies between d-Block and s-Block metals
First ionization energy: (IE)

The first ionization energy is the energy required to remove 1 mol of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions.
M (g) -> M⁺ (g) + e^-
In moving across the series of metals from scandium to zinc, small increases in the values of the first ionization energies are observed. This is due to the build-up of electrons in the immediately underlying d-sub-shells that efficiently shields the 4s electrons from the nucleus and minimizing the increase in effective nuclear charge from element to element.

From Ti to Cu, the 4s electrons are removed before the 3d electrons. 1^st and 2^nd IE are almost constant.
Explanation: 4s electrons are removed. Building up of electrons in the immediately underlying d sub-shells efficiently shields the 4s electrons from increasingly positive nucleus. Although the effective nuclear charge is expected to increase across the series, the repulsive effect of the inner 3d electrons also increases due to the continuous addition of 3d electrons which shows an increase in shielding effect. Hence the increase in effective nuclear charge from element to element is insignificant.

The increase in 3^rd IE from Sc to Zn and 4^th IE from Ti to Zn are more rapid.
Explanation: The d electrons are now being removed. The remaining electrons provide relatively poor shielding effect so that there is significant increase in effective nuclear charge from element to element when considering the 3^rd and 4^th IE.
The s-block elements have relatively large atomic radii and therefore have low values of ionization energy. In each group, the atomic radius increases with increasing atomic number, and there is a corresponding decrease in the value of the ionization energy, as the electron being ionized is further from the nucleus.
Within a given period, the Group II ion M²⁺ is smaller than the corresponding Group I ion M⁺. This effect is due to the fact that the same number of electrons is under the control of more protons.

Some slight irregularities in the values of ionization enthalpies in the series are:
- 1^st IE of Zn is exceptionally greater than that of Cu. The atomic structure of Zn is probably more stable in having full 3d and 4s sub-shells.
- 2^nd IE of Cr is slightly greater than that of its preceding and succeeding neighbours, i.e. V and Mn. 2^nd IE involves the removal of an electron from a half-filled 3d sub-shell, which has extra stability. The case is similar for Cu (which possesses a full 3d sub-shell) and its neighbour Ni and Zn.
- 3^rd IE of Fe is smaller than its preceding and succeeding neighbours, i.e. Mn and Co. For Fe, removal of the third electron leaves a half-filled 3d sub-shell which has extra stability.

Comparison of Electronegativities and Reaction with Water between d-Block and s-Block Metals
Electronegativity:

Relatively, all the transition elements are more electronegative than s-block elements since there is an increase in effective nuclear charge and no changes in the number of shells. It leads to the increase in the electrostatic forces between nucleus and valence electrons. Therefore, generally, they show a low reactivity because the outer electrons are tightly held and so, are not easily given away.
e.g. The rusting of iron is a slow process requiring both water and air. Other elements in the First d-series often have an inherent oxide layer on its surface (e.g. Ti and Cr) that resists corrosion by water and air. Zinc and iron react with steam at elevated temperatures to yield hydrogen.
3Fe (s) + 4H₂O (l) -> Fe₃O₄ (s) + 4H₂ (g)

Scandium, the first element in the d-series, shows no typical d-block or transition metal properties. The metal is similar to calcium in its reactivity with water.
Sc (s) + 3H₂O -> Sc(OH)₃ (aq) + 3/2H₂ (g)

Calcium reacts with cold water vigorously.
Ca (s) + 2H₂O (l) -> Ca(OH)₂ (aq) + 2H₂ (g)

Comparison of Conductivity between d-Block elements and s-Block elements
Conductivity of the transition elements:
Transition metals are good conductors of heat and electricity compared to s-block elements because the conductivity of metals result from delocalization of the outer electrons in the metallic bond, which makes the electrons free to move under the influence of a potential difference.
Transition metal atoms also have the 3d electrons as part of their metallic bond, which makes them better conductors than calcium, which can only use its 4s electrons. In fact, copper and silver are the best metallic conductors at room temperature.

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VARIABLE OXIATION STATES OF TRANSITION METAL COMPOUNDS
Table 1: Common Oxidation States in the Chlorides and Oxides of the elements Scandium to Copper

Interpretation in terms of Electronic Structures and Successive Ionization Enthalpies

Because the 3d and 4s electrons are of similar energy, both levels of electrons are available for bonding. There is only a gradual increase in successive ionization enthalpies for the d electrons as the overall charge of the metal ion increases, and no serious increase until an inert gas configuration is broken into.
Thus, when elements of the first d-series react to form compounds, they can form ions of roughly the same stability by losing different number of electrons.

Starting from the element Sc, the number of possible of oxidation states increases roughly with atomic number. The maximum number of oxidation states occurs at Mn. Thereafter, the number of oxidation states decreases, reaching a minimum at Zn.

The stability of Mn(II) and Fe(III) ([Ar] 3d⁵) can be explained by the fact that this half filled subshell in these ions has a particular stability. Removal of an additional electron disturbs this stable configuration so that relatively higher energy is required.
The bold numbers represent the common oxidation states of elements. Therefore, they also show the common ions according to the variable oxidation states of each element (by removing each electron from 4s-sub-shell to 3d-sub-shell)
e.g: some common ions of Mn: Mn²⁺, Mn³⁺, MnO₄^-, MnO₄^2-.

*: Exception
It explains why Sc and Zn are not transition metals.

Predict from a given electronic configuration, the likely variable oxidation states of a transition metal:
In general, the more stable the oxidation states, the more commonly they appear in the compounds of the elements. For the first d-series, the common oxidation states for each element include +2 and +3 or both. At the beginning of the series, the +3 state is more stable and hence more common. Towards the end of the series, the +2 state appears more frequently.

Explanation:This is because across the period, there is an increase in the atomic number of each element. In this case, since the cations are formed by losing electrons of 4s- and 3d-sub-shells, we consider the shielding effect of [Ar] core which is constant (correctly, both argon core and 3d-sub-shell cause shielding effect). Across the period, the increase in the atomic number causes the increase in the electrostatic forces of attraction between the nucleus and the 3d- ad 4s-sub-shell electrons. Therefore, at the beginning of the series, the oxidation state of +3 is more common and towards the end of the series, +2 appears more frequently (starting with Mn).

For this d-block series, the highest oxidation state rises to +7 at manganese. This corresponds to the removal of all outer electrons outside the stable argon core of inner electrons in these elements. If a further electron is to be removed, it would require much more energy. Therefore, one of the most important oxidation states of these elements is that involving loss all the 3d and 4s electrons, i.e., +4 for Ti, +5 for V, +6 for Cr and +7 for Mn.
In order to predict the likely variable oxidation states of a transition metal, we have to consider the stability of the formed ions.

For example, Fe: [Ar] 3d⁶ 4s²
Fe²⁺: [Ar] 3d⁶
Fe³⁺: [Ar] 3d⁵
The most common oxidation states of Fe are +2 and +3. We can guess so because studying the electronic configuration of Fe, we realize that the two electrons at the 4s-sub-shell can be removed easily due to the highest energy level of the outermost shell Fe2+. The energy level of 3d electrons can be comparable to that of 4s electrons but the oxidation state of +3 is still more common compared to that of +4, +5, +6 and so on. This is because the stability of half filled sub-shell of 3d⁵.

Mn: [Ar] 3d⁵ 4s²
Mn²⁺: [Ar] 3d⁵
Mn³⁺: [Ar] 3d⁴
MnO₄^-
MnO₄^2-
The most common oxidation states of Mn is +2, +4, +6 and +7. We can guess so because studying the electronic configuration of Mn, we realize that the two electrons at the 4s-sub-shell can be removed easily due to the highest energy level of the outermost shell Mn²⁺. The ion of Mn²⁺ is really stable because of the stability of half filled sub-shell of 3d⁵. The appearance of the oxidation states of +2, +3, +4, +6 and +7 are just practically due to the different conditions (temperature, pressure) when the redox reactions occur.

Exceptions: The electronic configuration of Cu: [Ar] 3d¹⁰ 4s¹
Cu⁺ (aq) -> Cu²⁺ (aq) + Cu (s)
We recognize that theoretically, Cu⁺ is supposed to be more stable than Cu²⁺ due to the stability of the filled sub-shell 3¹⁰ but actually, it is opposite. Practically, chemists use the value of electrode potential to prove the stability of Cu²⁺. But, there is an acceptable hypothesis that explains the stability of Cu²⁺. This hypothesis states that the original electronic configuration of Cu is [Ar] 3d⁹ 4s². There is one electron moving from the high energy level of 4s-sub-shell to 3d-sub-shell to fully fill up the 3d-sub-shell and make it more stable. People realize that that moving electron should have high energy level (excited state) so that it can move from a sub-shell to another. Therefore, it can also be easily removed from the atom to form ion.
Similarly, Cr does not form the ion of Cr⁺. Instead, Cr²⁺ is more favourable.

In the higher oxidation states, the elements exist as covalently bonded oxo-compounds rather than as atomic ions. This is because the more the electrons are lost, the stronger the electrostatic forces between the positively charged nucleus and the negatively shared pair of electrons for example. (Covalent character in ionic bonding)
e.g. Cr⁶⁺ (aq) does not exist but CrO₄^2- (aq) ions are formed instead.
This is because the formation of Cr⁶⁺ (aq) ion requires a very high energy and the ion has such a high charge density that it would polarize adjacent molecules (e.g. water, oxygen) to give the chromate (VI) ion, CrO₄^2- (aq) instead. This explanation is also appropriate to the impossible existence of Mn⁷⁺.

CATALYSTS
In homogeneous catalyst, the catalyst is in the same physical state as the reactants.
- The catalytic activity of transition metals in homogeneous catalysis depends on their ability to exist in variable oxidation states.
e.g. Fe³⁺ (aq) is a homogeneous catalyst in the oxidation of I^- (aq) by S₂O₈^2- (aq).
S₂O₈^2- (aq) + 2I^- (aq) -> 2SO₄^2- (aq) + I₂ (aq)
- The uncatalysed reaction is slow due to high activation energy (since two negative ions, S₂O₈^2- and I^-, are involved and they repel each other).
- The catalysed pathway involves the following steps with low Ea:
Step 1: 2Fe³⁺ + 2I^- -> 2Fe²⁺ + I₂
Step 2: 2Fe²⁺ + S₂O₈^2- -> 2Fe³⁺ + 2SO₄^2-

Overall: S₂O₈^2- + 2I^- -> 2SO₄^2- + I₂
Both steps involve the interaction of oppositely charged ions (Fe³⁺ and I^- in step 1; Fe²⁺ and S₂O₈^2- in step 2), which attract one another strongly and hence, enhances the rate of reactions.

COMMON OXIDATION STATES AND THEIR INTERCONVERSIONS
Some redox reactions of vanadium ions with Zn in concentrated HCl

VO₂⁺ (aq) -> VO²⁺ (aq) -> V³⁺ (aq) -> V²⁺ (aq) (Zn in concentrated HCl)
Some redox reactions of ions of manganese in different oxidation states
Manganese(IV)
MnO₂ (s) + 4HCl (conc.) -> MnCl₂ (aq) + Cl₂ (g) + 2H₂O (l)
2MnO₂ (s) + 4OH^- (l) + O₂ (g) (from KNO₃ or KClO₃) -> 2MnO₄^2- (l) + 2H₂O (g)

Manganese(VI)
3MnO₄^2- (aq) + 4H⁺ (aq) -> MnO₂ (s) + 2MnO₄^- (aq) + 2H₂O (l)

Manganese(VII)
3MnO₄^2- (aq) + H⁺ (aq) (from CO₂ (aq)) -> MnO₂ (s) + 2MnO₄^- (aq) + H₂O (l)
2KMnO₄ (s) -> K₂MnO₄ (s) + MnO₂ (s) + O₂ (g)
4MnO₄^- (aq) +4H⁺ (aq) -> 4MnO₂ (s) + 2H₂O (l) +3O₂ (g)
5Fe²⁺ (aq) + MnO₄^- (aq) + 8H⁺ (aq) -> 5Fe³⁺(aq) + Mn²⁺ (aq) + 4H₂O (l)

48MnO₄^- + C₁₂H₂₂O₁₁ + 48OH^- -> 48MnO₄^2- + 12CO₂ + 35H₂O
3MnO₄^2- + 4H⁺ (from CO2 (aq))-> MnO₂ + 2MnO₄^- + H₂O
Interconvertion of oxidation states of chromium ions
2CrO₄^2- + H⁺ <-> Cr₂O₇^2- + OH^-
3H₂O₂ + Cr₂O₇^2- + 8H⁺ -> 2Cr³⁺ + 3O₂ + 7H₂O

2CrO₄^2- + H⁺ <-> Cr₂O₇^2- + OH^-
3H₂O₂ + Cr₂O₇^2- + 8H⁺ -> 2Cr³⁺ + 3O₂ + 7H₂O

SO YOU STILL WANT MORE,

COLOUR OF TRANSITION METAL IONS AND COMPLEX FORMATION BY THE d-BLOCK ELEMENTS

First of all, we have to take note that transition metal ions do not exist alone in a solution. Instead, they are hydrated ions which can be expressed as complexes of [M(H2O)_m]ⁿ⁺.
Transition metal complexes are often coloured due to the presence of incompletely filled d-orbitals in the metal ion.

In the presence of ligands, especially water in this case, the 3d orbitals pf the transition ,eta; become non-degenerate and split into two groups with slightly different energy (d-d* splitting). Since the d_z² and d_x²-y² orbitals have lobes pointing directly at the ligand point charges (where ligands usually get to form dative bonds), they experience greater repulsion. They are thus destabilised and are higher in energy compared to that of the d_xy, d_xz and d_xz orbitals which experience less repulsion. They are thus less destabilised and are lower in energy
When a electron from lower energy group is promoted to the higher energy group (d-d* transition), radiation in the visible region of the electromagnetic spectrum is absorbed. The light energy nor absorbed will be seen as the colour of the ions.

Compounds of non-transition elements are usually colourless because they have no d electrons and so, d-d* transition is not possible. (* : excited state of an electron)
Examples are the alkali and alkaline earth ions, the halides ions, Zn²⁺, Al³⁺, Bi³⁺ and Sc³⁺.

COMPLEX FORMATION BY THE d-BLOCK ELEMENTS

Complex formation is one of the most important properties of d-block metals. A complex is formed when a d-block metal is surrounded by other molecules or ions which form dative covalent bonds with the d-block metal. The molecules or ions which form the dative bonds are called ligands.

For example, the Cr³⁺ (aq) ion is often referred to simple as aqueous Cr³⁺ ions, implying a picture of simple Cr³⁺ ions surrounded by more or less mobile water molecules [Cr(H₂O)₆]³⁺ (aq) ([Cr(H₂O)₆]Cl₃). This is an example of a complex ion; the six molecules of water are ligands, each utilizing the lone pair of electrons from oxygen to form a dative covalent bond with the central Cr³⁺ ion.

SUMMARY QUESTIONS
1) What is meant by “Transition Metals”? Give an example.
2) Does it mean that all the elements in d-block are transition metals? Give an example to show your answer.
3) Explain why there is an exception in writing the electronic configuration of chromium and copper.
4) Why do almost the transition metals have the similar physical properties?
5) Explain the trend in melting point of the main transition metals across the period.
6) Explain the trend in density of the main transition metals (period 4) and compare to the densities of non-transition metals. And then, explain why Sc has low density.
7) Describe and explain the little variation in atomic and ionic radius of transition metals.
8) Describe the tendency of transition elements to have variable oxidation states. Give an example to prove it (redox reaction).
9) Explain why the transition metals generally show a low reactivity compared to that of s-block elements (Calcium vigorously reacts with cold water while iron slowly reacts with hot water).
10) Why do the transition ions exist as covalently bonded oxo-compounds rather than as atomic ions in the high oxidation states? (For example, Cr⁶⁺ does not exist but CrO₄^2- are formed instead)
11) What is the difference between homogeneous and heterogenous catalysts?
12) * Describe the ability of transition metals in forming complex.

Oh dear, now I'm stuffed.

GROUP VII

All the halogens have diatomic molecules (due to the high stability of diatomic molecules and the tendency to reach the stable noble gas configuration), which in the solid state are arranged in a simple molecular structure. Halogen molecules are held in place by weak intermolecular forces known as induced dipole-induced dipole attraction. These forces result from an asymmetric distribution of electrons within each halogen molecule. This produces an instantaneous dipole, which induces dipoles in neighboring molecules. Fluorine exists as a gas because of the small size of the molecule and the ‘tight-bound’ electron that lead to the low tendency to be polarized. Therefore, the intermolecular forces between molecules are weak. Such a weak attractive force between molecules is easily overcome, so the element has relatively low melting points and boiling points. In other words, the state of a halogen can be indicated.

STRUCTURAL CHARACTERISTICS

- The first ionization energies are high but decrease down the group. The decrease from fluorine to chlorine is much greater than the decrease between any other two consecutive elements.

- The electronegativities are relatively high but decrease down the group because the outer electrons become progressively better shielded from the nucleus as the atomic size increases. Thus, electrons in a covalent bond are attracted less and less to the halogen as its atomic number increases.

- The covalent radii increase down the group so that the value for iodine is nearly twice that for fluorine. This is because as the group is descended, an extra shell of electrons is added to the atom of each successive element.

PHYSICAL PROPERTIES

Appearance: All the halogens are coloured because of the absorption of incident radiation in the visible region of the spectrum. The absorbed radiation excites electrons to higher energy levels. In the small fluorine atom, the electrons are ‘tightly bound’, so only high frequency radiation is absorbed, i.e. light from the blue end of the spectrum; hence fluorine is yellow. As the size of the halogen atom increases, radiation of lower frequency is absorbed so that the colours change as indicated until, with iodine, only the extreme violet is not absorbed.

Non-metals: All are typical non-metals and are bad conductors of heat and electricity.

Electron affinity:
A (g) + e^- -> A^- (g)
The electron affinities indicate the ease of ion formation. Electron affinity decreases numerically from chlorine to iodine with increasing atomic number. This is because the outer electrons become more shielded from the nucleus as the atomic size increases, so the tendency to attract another electron decreases as the group is descended.

Oxidizing agents: All the halogens behave as oxidizing agents and in a reaction they are reduced. This is due to their high strong tendency to gain electron compared to that of the rest within the period (the number of atomic number increases which results in the increase in the effective nuclear charge while the shielding effect does not change and therefore, the attraction between the nucleus and the valence electrons is more significant).
Of the halogens, fluorine is the most powerful oxidizing agent. This can be explained by the relative size of their atoms and their ability to capture an electron. The fluorine atom is the smallest, with fewer inner-shield shielding electrons, so its nucleus has a greater attraction or ability to attract an electron.

The oxidizing power of halogens can be shown by the redox reaction of halogens with thiosulphate ion S₂O₃^2-. F₂, Cl₂ and Br₂ are able to oxidize S₂O₃^2- to sulphate ion SO₄^2-.
4Br₂ (aq) + S₂O₃^2- (aq) + 5H₂O (l) -> 8Br^- (aq) + 2SO₄^2- (aq) + 10H⁺ (aq)
The oxidizing power of I₂ is less than the others. Therefore, it just has sufficient to oxidize S₂O₃^2- into S₄O₆^2- with the presence of catalyst of Fe³⁺(aq)
I₂ (aq) + 2S₂O₃^2- (aq) -> 2I^- (aq) + S₄O₆^2- (aq)

DISPLACEMENT REACTIONS (the reactions between halogens and other halide ions)
Since fluorine is the most reactive halogen, in theory it can react with the halide ion of any of the other halogens. Fluorine becomes the fluoride ion and the free halogen (chlorine, bromine or iodine) is form from the halide ion. This is called displacement reaction.

Reactions of halogens with halide ions follow the order of relative oxidizing power:
F₂ > Cl₂ > Br₂ > I₂
F₂ (g) + 2X^- (aq) -> 2F^- (aq) + X₂ (g) (X = Cl, Br or I)

Similarly, chlorine will only displace Br^- and I^- from their solutions whereas bromine can displace I^- only:
Cl₂ (aq) + 2Br^- (aq) -> Br₂ (aq) + 2Cl^- (aq)
During this reaction, the colour of the aqueous becomes orange, which indicates that the element bromine has been produced. Chlorine also displace iodide ions to form iodine, and again an orange-brown solution is formed.
Cl₂ (aq) + 2I^- (aq) -> 2Cl^- (aq) + I₂ (aq)

In the same way, aqueous bromine displaces aqueous iodide ion to form iodine. (The presence of iodine in a solution can be confirmed by the addition of starch solution. A dark blue coloration is formed.
Br₂ (aq) + 2I^- (aq) -> 2Br^- (aq) + I₂ (aq)
The displacement reaction is an example of a redox reaction. This is because the halogen that reacts is reduced because it gains electrons to form halide ions, and halide ions are oxidized by the loss of electrons to form the halogen.

REACTIONS OF HALIDE IONS WITH SILVER IONS
Silver halide, AgX, is precipitated when halide ions react with aqueous silver nitrate.
Ag⁺ (aq) + X^- (aq) -> AgX (s)

The silver halides formed can be differentiated form one another by
- The colour of the ppt formed, and
- The reaction of the precipitate with aqueous ammonia.

SUMMARY QUESTIONS
2) Describe the halogens as a collection of diatomic non-metals showing a trend in radius, colour and electron affinity.
3) Why is there an irregular change in the bond energy of X-X (X = F, Cl, Br, I)?
4) Explain why halogens can be known as strong oxidizing agents?
5)* Explain the difference in the stability of silver halides in the reaction with ammonia solution.