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General Trend in Properties

Metallic Behaviour:

All the transition elements are metals. Similar to all metals the transition metals are good conductors of heat and electricity. Unlike the metals of Group-1 and group-2, all the transition metals except group 11 elements are hard. Of all the known elements, silver has the highest electrical conductivity at room temperature. Most of the transition elements are hexagonal close packed, cubic close packed or body centrered cubic which are the characteristics of true metals.

Figure 4.2 Lattice Structures of 3d, 4d, and 5d transistion metals

As we move from left to right along the transition metal series, melting point first increases as the number of unpaired d electrons available for metallic bonding increases, reach a maximum value and then decreases, as the d electrons pair up and become less available for bonding.

For example, in the first series the melting point increases from Scandium (m.pt 1814K) to a maximum of 2183 K for vanadium, which is close to 2180K for chromium. However, manganese in 3d series and Tc in 4d series have low melting point. The maximum melting point at about the middle of transition metal series indicates that d⁵ configuration is favourable for strong interatomic attraction. The following figure shows the trends in melting points of transition elements.

Variation of Atomic and Ionic Size:

It is generally expected a steady decrease in atomic radius along a period as the nuclear charge increases and the extra electrons are added to the same sub shell. But for the 3d transition elements, the expected decrease in atomic radius is observed from Sc to V, thereafter up to Cu the atomic radius nearly remains the same.

As we move from Sc to Zn in 3d series the extra electrons are added to the 3d orbitals, the added 3d electrons only partially shield the increased nuclear charge and hence the effective nuclear charge increases slightly. However, the extra electrons added to the 3d sub shell strongly repel the 4s electrons and these two forces are operated in opposite direction and as they tend to balance each other, it leads to constancy in atomic radii.

At the end of the series, d – orbitals of Zinc contain 10 electrons in which the repulsive interaction between the electrons is more than the effective nuclear charge and hence, the orbitals slightly expand and atomic radius slightly increases. Generally as we move down a group atomic radius increases, the same trend is expected in d block elements also. As the electrons are added to the 4d sub shell, the atomic radii of the 4d elements are higher than the corresponding elements of the 3d series.

However there is an unexpected observation in the atomic radius of 5d elements which have nearly same atomic radius as that of corresponding 4d elements. T is is due to lanthanoide contraction which is to be discussed later in this unit under inner transition elements.

Ionization Enthalpy:

Ionization energy of transition element is intermediate between those of s and p block elements. As we move from left to right in a transition metal series, the ionization enthalpy increases as expected. This is due to increase in nuclear charge corresponding to the filling of d electrons. The following figure show the trends in ionisation enthalpy of transition elements.

The increase in first ionisation enthalpy with increase in atomic number along a particular series is not regular. The added electron enters (n-1)d orbital and the inner electrons act as a shield and decrease the effect of nuclear charge on valence ns electrons. Therefore, it leads to variation in the ionization energy values.

The ionisation enthalpy values can be used to predict the thermodynamic stability of their compounds. Let us compare the ionisation energy required to form Ni²⁺ and Pt²⁺ ions.

For Nickel, IE₁ + IE₂
= (735 + 1753)
= 2490 kJmol^-1

For Platinum, IE₁ + IE₂
= (864 + 1791)
= 2655 kJmol^-1

Since, the energy required to form Ni²⁺ is less than that of Pt²⁺, Ni(II) compounds are thermodynamically more stable than Pt(II) compounds.

Oxidation State:

The first transition metal Scandium exhibits only +3 oxidation state, but all other transition elements exhibit variable oxidation states by loosing electrons from (n-1)d orbital and ns orbital as the energy difference between them is very small. Let us consider the 3d series; the following table summarizes the oxidation states of the 3d series elements.

At the beginning of the series, +3 oxidation state is stable but towards the end +2 oxidation state becomes stable. The number of oxidation states increases with the number of electrons available, and it decreases as the number of paired electrons increases.

Hence, the first and last elements show less number of oxidation states and the middle elements with more number of oxidation states. For example, the first element Sc has only one oxidation state +3; the middle element Mn has six different oxidation states from +2 to +7. The last element Cu shows +1 and +2 oxidation states only.

The relative stability of different oxidation states of 3d metals is correlated with the extra stability of half filled and fully filled electronic confiurations. Example: Mn²⁺(3d⁵) is more stable than Mn⁴⁺(3d³)

The oxidation states of 4d and 5d metals vary from +3 for Y and La to +8 for Ru and Os. The highest oxidation state of 4d and 5d elements are found in their compounds with the higher electronegative elements like O, F and Cl. for example: RuO₄, OsO₄ and WCl₆. Generally in going down a group, a stability of the higher oxidation state increases while that of lower oxidation state decreases.

It is evident from the Frost diagram (ΔG° vs oxidation number) as shown below. For titanium, vanadium and chromium, the most thermodynamically stable oxidation state is +3. For iron, the stabilities of +3 and +2 oxidation states are similar. Copper is unique in 3d series having a stable +1 oxidation state. It is prone to disproportionate to the +2 and 0 oxidation states.

Standard Electrode Potentials of Transition Metals

Redox reactions involve transfer of electrons from one reactant to another. Such reactions are always coupled, which means that when one substance is oxidised, another must be reduced. The substance which is oxidised is a reducing agent and the one which is reduced is an oxidizing agent. The oxidizing and reducing power of an element is measured in terms of the standard electrode potentials.

Standard electrode potential is the value of the standard emf of a cell in which molecular hydrogen under standard pressure (1 atm) and temperature (273 K) is oxidised to solvated protons at the electrode. If the standard electrode potential (E°), of a metal is large and negative, the metal is a powerful reducing agent, because it loses electrons easily. Standard electrode potentials (reduction potential) of few first transition metals are given in the following table.

In 3d series as we move from Ti to Zn, the standard reduction potential (\(E^{0}{ }_{M^{2+}} /_{M}\)) value is approaching towards less negative value and copper has a positive reduction potential. i.e., elemental copper is more stable than Cu²⁺.

There are two deviations., In the general trend, Fig shows that (\(E^{0}{ }_{M^{2+}} /_{M}\)) value for manganese and zinc are more negative than the regular trend. It is due to extra stability which arises due to the half filled d⁵ configuration in Mn²⁺ and completely filled d¹⁰ configuration in Zn²⁺.

Transition metals in their high oxidation states tend to be oxidizing. For example, Fe³⁺ is moderately a strong oxidant, and it oxidises copper to Cu²⁺ ions. The feasibility of the reaction is predicted from the following standard electrode potential values.

Fe³⁺(aq) + e^– ⇄ Fe²⁺ E° = 0.77V
Cu²⁺(aq) + 2e^– ⇄ Cu(s) E° = +0.34V

The standard electrode potential for the M³⁺/M²⁺ half-cell gives the relative stability between M³⁺ and M²⁺. The reduction potential values are tabulated as below.

The negative values for titanium, vanadium and chromium indicate that the higher oxidation state is preferred. If we want to reduce such a stable Cr³⁺ ion, strong reducing agent which has high negative value for reduction potential like metallic zinc (E° = – 0.76 V) is required.

The high reduction potential of Mn³⁺/Mn²⁺ indicates Mn²⁺ is more stable than Mn³⁺. For Fe³⁺/Fe²⁺ the reduction potential is 0.77V, and this low value indicates that both Fe³⁺ and Fe²⁺ can exist under normal conditions. The drop from Mn to Fe is due to the electronic structure of the ions concerned.

Mn³⁺ has a 3d⁴ configuration while that of Mn²⁺ is 3d⁵. The extra stability associated with a half filled d sub shell makes the reduction of Mn³⁺ very feasible (E° = +1.51V).

Magnetic Properties

Most of the compounds of transition elements are paramagnetic. Magnetic properties are related to the electronic configuration of atoms. We have already learnt in XI STD that the electron is spinning around its own axis, in addition to its orbital motion around the nucleus. Due to these motions, a tiny magnetic field is generated and it is measured in terms of magnetic moment. On the basis of magnetic properties, materials can be broadly classified as

• Paramagnetic Materials
• Diamagnetic materials, besides these there are ferromagnetic and antiferromagnetic materials.

Materials with no elementary magnetic dipoles are diamagnetic, in other words a species with all paired electrons exhibits diamagnetism. This kind of materials are repelled by the magnetic field because the presence of external magnetic field, a magnetic induction is introduced to the material which generates weak magnetic field that oppose the applied field.

Paramagnetic solids having unpaired electrons possess magnetic dipoles which are isolated from one another. In the absence of external magnetic field, the dipoles are arranged at random and hence the solid shows no net magnetism. But in the presence of magnetic field, the dipoles are aligned parallel to the direction of the applied field and therefore, they are attracted by an external magnetic field.

Ferromagnetic materials have domain structure and in each domain the magnetic dipoles are arranged. But the spin dipoles of the adjacent domains are randomly oriented. Some transition elements or ions with unpaired d electrons show ferromagnetism.

3d transition metal ions in paramagnetic solids often have a magnetic dipole moments corresponding to the electron spin contribution only. The orbital moment L is said to be quenched. So the magnetic moment of the ion is given by µ = g\(\sqrt{S(S+1)}\) µ_B

Where S is the total spin quantum number of the unpaired electrons and is µ_B Bohr Magneton. For an ion with n unpaired electrons S = \(\frac{n}{2}\) and for an electron g = 2.

Therefore the spin only magnetic moment is given by

The magnetic moment calculated using the above equation is compared with the experimental values in the following table. In most of the cases, the agreement is good.

Catalytic Properties

The chemical industries manufacture a number of products such as polymers, flavours, drugs etc., Most of the manufacturing processes have adverse effect on the environment so there is an interest for eco friendly alternatives. In this context, catalyst based manufacturing processes are advantageous, as they require low energy, minimize waste production and enhance the conversion of reactants to products.

Many industrial processes use transition metals or their compounds as catalysts. Transition metal has energetically available d orbitals that can accept electrons from reactant molecule or metal can form bond with reactant molecule using its d electrons. For example, in the catalytic hydrogenation of an alkene, the alkene bonds to an active site by using its π electrons with an empty d orbital of the catalyst.

The σ bond in the hydrogen molecule breaks, and each hydrogen atom forms a bond with a d electron on an atom in the catalyst. The two hydrogen atoms then bond with the partially broken π – bond in the alkene to form an alkane.

In certain catalytic processes the variable oxidation states of transition metals find applications. For example, in the manufacture of sulphuric acid from SO₃, vanadium pentoxide (V₂O₅) is used as a catalyst to oxidise SO₂. In this reaction V₂O₅ is reduced to vanadium (IV) Oxide (VO₂).

Some more examples are discussed below,

(i) Hydroformylation of Olefins

(ii) Preparation acetic acid from acetaldehyde.

(iii) Zeigler – Natta Catalyst

A mixture of TiCl₄ and trialkyl aluminium is used for polymerization.

Alloy Formation

An alloy is formed by blending a metal with one or more other elements. The elements may be metals or non metals or both. The bulk metal is named as solvent, and the other elements in smaller portions are called solute. According to Hume-Rothery rule to form a substitute alloy the difference between the atomic radii of solvent and solute is less than 15%.

Both the solvent and solute must have the same crystal structure and valence and their electro negativity difference must be close to zero. Transition metals satisfying these mentioned conditions form a number of alloys among themselves, since their atomic sizes are similar and one metal atom can be easily replaced by another metal atom from its crystal lattice to form an alloy. The alloys so formed are hard and often have high melting points. Examples: Ferrous alloys, gold – copper alloy, chrome alloys etc.

Formation of Interstitial Compounds

An interstitial compound or alloy is a compound that is formed when small atoms like hydrogen, boron, carbon or nitrogen are trapped in the interstitial holes in a metal lattice. They are usually non-stoichiometric compounds. Transition metals form a number of interstitial compounds such as TiC, ZrH_1.92, Mn₄N etc. The elements that occupy the metal lattice provide them new properties.

• They are hard and show electrical and thermal conductivity
• They have high melting points higher than those of pure metals
• Transition metal hydrides are used as powerful reducing agents
• Metallic carbides are chemically inert

Formation of Complexes

Transition elements have a tendency to form coordination compounds with a species that has an ability to donate an electron pair to form a coordinate covalent bond. Transition metal ions are small and highly charged and they have vacant low energy orbitals to accept an electron pair donated by other groups. Due to these properties, transition metals form large number of complexes. Examples: [Fe(CN)₆]^4-, [Co(NH₃)₆]³⁺, etc. The chemistry of coordination compound is discussed in unit 5.

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Coordination Compounds and Double Salts

When two or more stable compounds in solution are mixed together and allowed to evaporate, in certain cases there is a possibility for the formation of double salts or coordination compounds. For example when an equimolar solution of ferrous sulphate and ammonium sulphate are mixed and allowed to crystallise, a double salt namely Mohr’s salt (Ferrous ammonium sulphate, FeSO₄.(NH₄)₂SO₄.6H₂O) is formed.

Let us recall the blood red colour formation in the inorganic qualitative analysis of ferric ion, the reaction between ferric chloride and potassium thiocyanate solution gives a blood red coloured coordination compound, potassium ferrithiocyanate K₃[Fe(SCN)₆].

If we perform a qualitative analysis to identify the constituent ions present in both the compounds, Mohr’s salt answers the presence of Fe²⁺, NH₄⁺ and SO₄^2- ions, whereas the potassium ferrithiocyanate will not answer Fe³⁺ and SCN^–ions. From this we can infer that the double salts lose their identity and dissociates into their constituent simple ions in solutions, whereas the complex ion in coordination compound, does not loose its identity and never dissociate to give simple ions.

Double salts and coordination compounds are complex compounds. The difference between double salt and coordination compound is that a double salt contains two salts with different crystal structures whereas a coordination compound contains a central metal ion surrounded by molecules or ions known as ligands.

A salt is essentially composed of an anion and a cation. But the main difference between a double salt and a complex salt is that a double salt is a combination of two salt compounds whereas a complex salt is a molecular structure that is composed of one or more complex ions.

A complex salt is a compound composed of a central metal atom having coordination bonds with ligands around it. Do not completely dissociate into its ions in water. It cannot be easily analyzed by determining the ions in the aqueous solution.

Both double salt as well sas complexes are formed by the combination of two or more stabel compounds in stoichiometic reatio. However they differ in the fact that double salt disssociate ito simple ions cmpletely with dissolved in water. However complex ions do not simple ions completely.

A double salt is a salt that contains more than one cation or more than one anion. Other examples include potassium sodium tartrate, ammonium iron (II) sulfate (Mohr’s salt), and bromlite. The fluorocarbonates contain fluoride and carbonate anions.

Double salts and coordination compounds are complex compounds. The difference between double salt and coordination compound is that a double salt contains two salts with different crystal structures whereas a coordination compound contains a central metal ion surrounded by molecules or ions known as ligands.

A double salt is formed from a three-component system, comprising two separate salts and water, and at a given temperature this may be represented by a triangular diagram. Phase diagram of a three-component system.

Double salts are addition compounds which lose their identity in aqueous solution whereas complexes which are also addition compounds do not lose their identity in aqueous solution.

In chemistry, a double bond is a covalent bond between two atoms involving four bonding electrons as opposed to two in a single bond. Double bonds occur most commonly between two carbon atoms, for example in alkenes. Other common double bonds are found in azo compounds (N=N), imines (C=N), and sulfoxides (S=O).

A complex salt is a salt that contains one or more complex ions – ions with metal centers and different molecules attached. Complex salts include potassium ferricyanide (used to create dyes and in blueprint paper) and potassium argentocyanide (used in silver plating).

The molecularity of the chemical reaction is equal to the sum of the stochiometric coefficients of the reactants in the chemical equation of the reaction. It is also defined as the number of reactant molecules taking part in a single step of the reaction.

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Electronic Configuration – Detailed Explanation with Examples

We have already learnt in XI STD to write the electronic configuration of the elements using Aufbu principle, Hund’s rule etc. According to Aufbau principle, the electron first fills the 4s orbital before 3d orbital. Therefore filling of 3d orbital starts from Sc, its electronic configuration is [Ar]3d¹4s² and the electrons of successive elements are progressively filled in 3d orbital and the filling of 3d orbital is complete in Zinc, whose electronic configuration is [Ar]3d¹⁰4s².

However, there are two exceptions in the above mentioned progressive filling of 3d orbitals; if there is a chance of acquiring half filled or fully filled 3d sub shell, it is given priority as they are the stable configuration, for example Cr and Cu. The electronic configurations of Cr and Cu are [Ar] 3d⁵4s¹ respectively.

The extra stability of half filled and fully filled d orbitals, as already explained in XI STD, is due to symmetrical distribution of electrons and exchange energy. Note: The extra stability due to symmetrical distribution can also be visualized as follows. When the d orbitals are considered together, they will constitute a sphere.

So the half filled and fully filled configuration leads to complete symmetrical distribution of electron density. On the other hand, an unsymmetrical distribution of electron density as in the case of partially filled configuration will result in building up of a potential difference.

To decrease this and to achieve a tension free state with lower energy, a symmetrical distribution is preferred. With these two exceptions and minor variation in certain individual cases, the general electronic configuration of d – block elements can be written as [Noble gas] (n – 1)d^1-10 ns^1-2

To calculate an electron configuration, divide the periodic table into sections to represent the atomic orbitals, the regions where electrons are contained. Groups one and two are the s-block, three through 12 represent the d-block, 13 to 18 are the p-block and the two rows at the bottom are the f-block.

Electronic configuration, also called electronic structure, the arrangement of electrons in energy levels around an atomic nucleus. In terms of a more refined, quantum-mechanical model, the K-Q shells are subdivided into a set of orbitals (see orbital), each of which can be occupied by no more than a pair of electrons.

There are different orbital shapes (s, p, d, f) Each orbital can only hold 2 electrons max. There is a hierarchy, i.e. s orbitals will be filled before p orbitals which will be filled before d orbitals and so on. (s<p<d<f) (note, this is a general rule but there are exceptions).

The electron configuration is the standard notation used to describe the electronic structure of an atom. When assigning electrons to orbitals, we must follow a set of three rules: the Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule.

If you are given with the atomic number of an element you can find it’s period number and group number. The period number is related to the number of electron occupied shells in the element and the period number is linked to its valence electrons.

There are two main exceptions to electron configuration: chromium and copper. In these cases, a completely full or half full d sub-level is more stable than a partially filled d sub-level, so an electron from the 4s orbital is excited and rises to a 3d orbital.

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Position of D – Block Elements in the Periodic Table

We have already learnt the periodic classification of elements in XI std. the transition metals occupy from group – 3 to group – 12 of the modern periodic table.

Figure 4.1 – Position of d – block elements in the periodic table

D – Block elements composed of 3d series (4th period) Scandium to Zinc (10 elements), 4d series (5^th period) Yttrium to Cadmium (10 elements) and 5d series (6^th period) Lanthanum, Haffinium to mercury. As we know that the group-12 elements Zinc, Cadmium and Mercury do not have partially filled d-orbital either in their elemental state or in their normal oxidation states.

However they are treated as transition elements, because their properties are an extension of the properties of the respective transition elements. As per the IUPAC definition, the seventh period elements, starting from Ac, Rf to Cn also belong to transition metals. All of them are radioactive. Except Actinium; all the remaining elements are synthetically prepared and have very low half life periods.

The d-block elements are found in the middle of the period table. The d-block elements are called transition metals and have valence electrons in d orbital’s. The f-block elements,found in the two rows at the bottom of the periodic table, are called inner transition metals and have valence electrons in the f-orbital’s.

Transition elements are the elements that are found in Groups 3-12 (old groups IIA-IIB) on the periodic table (salmon-colored block in the middle of the table).

The Periodic Table provides a section for each of these groups of orbitals. The 10 electrons of the five d orbitals are filled by the elements found in the dropped central section of the table. This section is referred to as the d block elements, or the transition metals.

The d-block of the periodic table contains the elements of the groups 3-12 in which the d orbitals are progressively filled in each of the four long periods. The f-block consists of elements in which 4 f and 5 f orbitals are progressively filled. They are placed in a separate panel at the bottom of the periodic table.

The d-block elements are called transition elements because they exhibit transitional behaviour between s block and p-block elements. Their properties are transitional between highly reactive metallic elements of s block which are ionic in nature and the elements of p-block which are covalent in nature.

According to Aufbau principle , electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled. Therefore, 3d orbital is higher in energy than 4s. And hence electrons fill up in 4s before filling up in 3d.

The p sublevel has 3 orbitals, so can contain 6 electrons max. The d sublevel has 5 orbitals, so can contain 10 electrons max. And the 4 sublevel has 7 orbitals, so can contain 14 electrons max. In the picture below, the orbitals are represented by the boxes.

D-Block Elements:

Elements in which the last electron enters any one of the five d-orbitals of their respective penultimate shells are called d-block elements. The Importance of d-block Transition Metals. Transition metals, for the most part, are good conductors. They are also malleable, ductile, lustrous, and sliver-white in color. The d-block metals, and some of it’s key alloys, shaped the Bronze Age, Iron Age, and most importantly the steel age.

The p-block elements are found on the right side of the periodic table. They include the boron, carbon, nitrogen, oxygen and flourine families in addition to the noble gases. The noble gases have full p-orbital’s and are nonreactive.

In chemistry and atomic physics, the main group is the group of elements (sometimes called the representative elements) whose lightest members are represented by helium, lithium, beryllium, boron, carbon, nitrogen, oxygen, and fluorine as arranged in the periodic table of the elements.

Chlorine is in group 17 of periodic table, also called the halogens, and is not found as the element in nature – only as a compound. The most common of these being salt, or sodium chloride, and the potassium compounds sylvite (or potassium chloride) and carnallite (potassium magnesium chloride hexahydrate).

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Group 18 (Inert gases) Elements

Occurrence:

All the noble gases occur in the atmosphere.

Physical Properties:

As we move along the noble gas elements, their atomic radius and boiling point increases from helium to radon. The first ionization energy decreases from helium to radon. Noble gases have the largest ionisation energy compared to any other elements in a given row as they have completely filled orbital in their outer most shell. They are extremely stable and have a small tendency to gain or lose electrons. The common physical properties of the group 18 elements are listed in the Table.

Physical Properties of group 18 Elements

Properties of Inert Gases:

Physical Properties:

Noble gases are monoatomic, odourless, colourless, tasteless, and non-inflammable. They are highly unreactive. They are non-metallic in nature.

Chemical Properties:

Only the xenon and krypton show some chemical reactivity. Xenon fluorides are prepared by direct reaction of xenon and fluorine under different conditions as shown below.

When XeF₆ is heated at 50°C in a sealed quartz vessel it forms XeOF₄.

When the reaction is continued the following reaction takes place.

2XeOF₄ + SiO₂ → 2XeO₂F₂ + SiF₄
2XeO₂F₂ + SiO₂ → 2XeO₃ + SiF₄

On hydrolysis with water vapour XeF₆ gives XeO₃

XeF₆ + 3H₂O → XeO₃ + 6HF

When XeF₆ reacts with 2.5 M NaOH, sodium per xenate is obtained.

2XeF₆ + 16NaOH → Na₄XeO₆ + Xe + O₂ + 12NaF + 8H₂O

Sodium per xenate is very much known for its strong oxidizing property. For example, it oxidises manganese (II) ion into permanganate ion even in the absence of the catalyst.

5XeO₆^4- + 2Mn²⁺ + 14H⁺ → 2MnO₄^– + 5XeO₃ + 7H₂O

Xenon reacts with PtF₆ and gave an orange yellow solid [XePtF₆] and this is insoluble in CCl₄.

Xenon difluoride forms addition compounds XeF₂.2SbF₅ and XeF₂.2TaF₅. Xenon
hexa fluorides forms compound with boron and alkali metals. Eg: XeF₆.BF₃, XeF₆MF, M-alkali metals.

There is some evidence for existence of xenon dichloride XeCl₂.

Krypton form krypton difluoride when an electric discharge is passed through Kr and flourine at 183°C or when gases are irradiated with SbF₅ it forms KrF₂.2SbF₃.

Structures of Compounds of Xenon:

Uses of Noble Gases:

The inertness of noble gases is an important feature of their practical uses.

Helium:

1. Helium and oxygen mixture is used by divers in place of air oxygen mixture. This prevents the painful dangerous condition called bends.
2. Helium is used to provide inert atmosphere in electric arc welding of metals.
3. Helium has lowest boiling point hence used in cryogenics (low temperature science)
4. It is much less denser than air and hence used for filing air balloons.

Neon:

Neon is used in advertisement as neon sign and the brilliant red glow is caused by passing electric current through neon gas under low pressure.

Argon:

Argon prevents the oxidation of hot filament and prolongs the life in filament bulbs

Krypton:

Krypton is used in florescent bulbs, flash bulbs etc. Lamps filed with krypton are used in airports as approaching lights as they can penetrate through dense fog.

Xenon:

Xenon is used in florescent bulbs, flash bulbs and lasers. Xenon emits an intense light in discharge tubes instantly. Due to this it is used in high speed electronic flash bulbs used by photographers.

Radon:

Radon is radioactive and used as a source of gamma rays. Radon gas is sealed as small capsules and implanted in the body to destroy malignant i.e. cancer growth.