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scc-education,classification-of-elements,scc,sharma sir,CHEMICAL BONDING AND MOLECULAR STRUCTURE
Periodic table is a remarkable demonstration of the fact that the chemical elements are not a random cluster of entities but instead display trends and lie together in families.
At present 114 elements are known.. With such a large number of elements it is very difficult to study individually the chemistry of all these elements and their innumerable compounds individually. To ease out this problem, scientists searched for a systematic way to organise their knowledge by classifying the elements.
1. Dobereiner’s Law of Triads
The German chemist, Johann Dobereiner in early 1800’s was the first to consider the idea of trends among properties of elements According to this law there is a similarity among the physical and chemical properties of several groups of three elements (Triads).. In each case, the middle element of each of the Triads had an atomic weight about half way between the atomic weights of the other two (Table). Also the properties of the middle
element were in between those of the other two members.
Table showing examples of Dobereiner’s Triads
Limitations of Dobereiner’sLaw of Triads
It works only for a few elements, it was dismissed as coincidence.
The Newland’s Law of Octaves

The English chemist, John Alexander New-lands in 1865 gave this law.--
He arranged the elements in increasing order of their atomic weights and noted that every eighth element had properties similar to the first element (Table). The relationship was just like every eighth note that resembles the first in octaves of music.
Limitations of Newlands’s Law
This law is true only for elements up to calcium. Although his idea was not widely accepted at that time, he, for his work, was later awarded Davy Medal in 1887 by the Royal Society, London, because he was the first person who gave the idea of periodicity that is used in modern periodic table also.
Russian chemist, Dmitri Mendeleev (1834-1907) and the German chemist, Lothar Meyer (1830-1895). Working independently, both the chemists in 1869 proposed that on arranging elements in the increasing order of their atomic weights, similarities appear in physical and chemical properties at regular intervals. Lothar Meyer plotted the physical properties such as atomic volume, melting point and boiling point against atomic weight .
Mendeleev’s periodic law--
The properties of the elements are a periodic function of their atomic weights
( Repetition of properties after a regular time interval is known as periodicity.)
Structure of the Mendeleev’s periodic table—
1. He arranged elements in horizontal rows and vertical columns of a table in order of their increasing atomic weights in such a way that the elements with similar properties occupied the same vertical column or group.
2. Horizontal rows are known as periods, and they further divided in to odd and even series.
3. Groups were devided in to sub groups
Applications of Mendeleev’s periodic table.—
1.Study of elements become very easy.
2. Some of the elements did not fit in the scheme of classification if the order of atomic weight was strictly followed. He ignored the order and also correct the atomic weights. e.g, iodine with lower atomic weight than that of tellurium (Group VI) was placed in Group VII along with fluorine, chlorine, bromine because of similarities in properties
3. He forecasted about some undiscovered elements. On the basis of properties of elements and their gradual difference he proposed that some of the elements were still undiscovered and, he left several gaps in the table, e.g. both gallium and germanium were unknown at the time Mendeleev published his Periodic Table. He left the gap under aluminium and a gap under silicon, and called these elements Eka- Aluminium and Eka-Silicon. Mendeleev
predicted not only the existence of gallium and germanium, but also described some of their general physical properties. These elements were discovered later.
Mendeleev’s Predictions for the Elements Eka-aluminium (Gallium) and Eka-silicon (Germanium)
Drawbacks of Mendeleev's periodic table—
  1. He could not justify the position of isotopes. i.e. different atomic positions but they have same position in periodic table, 1H1, 1H2 and 1H3
  2. He could not justify the position of isobars, i.e. same atomic masses but elements having different positions in the periodic table, e.g. 20Ca40 and 19K39.
  3. There is no position for f-block elements in the main body of periodic table.
Position of hydrogen is controversial.
Note--We must bear in mind that when Mendeleev developed his Periodic Table, chemists knew nothing about the internal structure of atom.
In 1913, the English physicist, Henry Moseley modified. Mendeleev’s Periodic Law on the basis of his certain This is known as the Modern Periodic Law and can be stated as :
The physical and chemical properties of the elements are periodic functions of their atomic numbers.
Note--We know atomic number is equal to the nuclear charge (i.e., number of protons) or the number of electrons in a neutral atom. It shows the significance of quantum numbers and electronic configurations in periodicity of elements..
THE PRESENT FORM OF THE PERIODIC TABLE or long form” of the Periodic Table--
  1. It has 7 horizontal rows are called periods, hence there are altogether seven periods.
  2. 18 vertical columns known as groups. Elements having similar outer electronic configurations in their
atoms are arranged in vertical columns, referred to as groups or families. According to International Union of Pure and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18 replacing the older notation of groups IA … VIIA, VIII, IB … VIIB and 0.
3.The period number corresponds to the highest principal quantum number (n) of the elements in the period. The first period contains 2 elements. The subsequent periods consists of 8, 8, 18, 18 and 32 elements, respectively. The
seventh period is incomplete and like the sixth period would have a theoretical maximum (on the basis of quantum numbers) of 32 elements.
4. In this form of the Periodic Table, 14 elements of both sixth and seventh periods (lanthanoids and actinoids, respectively) are placed in separate panels at the bottom*.
Advantage of modern periodic law over Mendeleev’s periodic law—
  1. Position of isotopes and isobars became clear.
  2. Position of lanthanides and actinides are justified.
But position of hydrogen is still controversial.
The naming of the new elements had been traditionally the privilege of the discoverer (or discoverers) and the suggested name was ratified by the IUPAC. In recent years this has led to some controversy. For example, both American and Soviet scientists claimed credit for discovering element 104. The Americans named it Rutherfordium whereas Soviets named it Kurchatovium. To avoid such problems, the IUPAC has made recommendation that until a new element’s discovery is proved, and its name is officially recognized, a systematic nomenclature be derived directly from the atomic number of the element using the numerical roots for 0 and numbers 1-9. These are shown in Table. The roots are put together in order of digits which make up the atomic number and “ium” is added at the end. 
The distribution of electrons into orbitals of an atom is called its electronic configuration. An element’s location in the Periodic Table reflects the quantum numbers of the last orbital filled.
(a) Electronic Configurations in Periods
The period indicates the value of n for the outermost or valence shell. The number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled. The first period (n = 1) starts with the filling of the lowest level (1s) and therefore has two elements — hydrogen (ls1) and helium (ls2) when the first shell (K) is completed. The second period (n = 2) starts with lithium and the third electron enters the 2s orbital. The next element, beryllium has four electrons and has the electronic configuration 1s22s2. Starting from the next element boron, the 2p orbitals are filled with electrons when the L shell is completed at neon (2s22p6). Thus there are 8 elements in the second period. The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital. Successive filling of 3s and 3p orbitals gives rise to the third period of 8 elements from sodium to argon. The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital. Now you may note that before the 4p orbital is filled, filling up of 3d orbitals becomes energetically favourable and we come across the so called 3d transition series of elements. This starts from scandium (Z = 21) which has the electronic configuration 3d14s2. The 3d orbitals are filled at zinc (Z=30) with electronic configuration 3d104s2 . The fourth period ends at krypton with the filling up of the 4p orbitals. Altogether we have 18 elements in this fourth period. The fifth period (n = 5) beginning with rubidium is similar to the fourth period and contains the 4d transition series starting at yttrium (Z = 39). This period ends at xenon with the filling up of the 5p orbitals. The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, 5d and 6p orbitals, in the order — filling up of the 4f orbitals begins with cerium (Z = 58) and ends at lutetium (Z = 71) to give the 4f-inner transition series which is called the lanthanoid series. The seventh period (n = 7) is similar to the sixth period with the successive filling up of the 7s,
5f, 6d and 7p orbitals and includes most of the man-made radioactive elements. This period will end at the element with atomic number 118 which would belong to the noble gas family. Filling up of the 5f orbitals after actinium (Z = 89) gives the 5f-inner transition series known as the actinoid series. The 4fand 5f-inner transition series of
elements are placed separately in the Periodic Table to maintain its structure and to preserve the principle of classification by keeping elements with similar properties in a single column.
(b) Groupwise Electronic Configurations
Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties. For example, the Group 1 elements (alkali metals) all have ns1 valence shell electronic configuration as shown below. Thus it can be seen that the properties of an element have periodic dependence upon its atomic number and not on relative atomic mass.
. The elements in a group or family, exhibit similar chemical behaviour. This similarity arises because these elements have the same number and same distribution of electrons in their outermost orbitals.
We can classify the elements into four blocks viz., s-block, p-block, d-block and f-block
The s-Block Elements---Elements in which last electron enter in to the s-sub shell are known as s-block elements.
--The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1ns2 outermost electronic configuration belong to the s-Block Elements. Hence general electronic configuration is ns1to2.
Important properties of s-Block Elements are as ---1. They are all very reactive metals with low ionisation enthalpies.
2. They lose the outermost electron(s) readily to form 1+ ion (in the case of alkali metals) or 2+ ion (in the case of alkaline earth metals).
3. The metallic character and the reactivity increase as we go down the group.
4. Because of high reactivity they are never found pure in nature.
5. The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic.

Q. “He belongs to the s-block but its positioning in the p-block along with other group 18 elements.” Justify the position of He.
Ans. It is justified because it has a completely filled valence shell (1s2) and as a result, exhibits properties of other noble gases.
Q. What is the reason “position of hydrogen in periodic table is controversial”?
Ans. It has a lone s-electron and hence can be placed in group 1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 (halogen family) elements. Because it is a special case, we shall place hydrogen separately at the top of the Periodic Table
The p-Block Elements---- Elements in which last electron enter in to the p-sub shell are known as d-block elements.
The elements of Group 13 (aluminium family), Group 14 (Carbon family), Group 15 (Nitrogen family), Group 16 (Chalcogens), Group 17 (Halogens) and Group 18 (Inert Gases) which have ns2 np1, ns2 np2, ns2 np3 , ns2np4 , ns2 np5 , and ns2np6 outermost electronic configuration respectively.Hence general electronic configuration is ns2 np1 to 6
---s-Block Elements and p-block elements collectively are called the Representative 

Elements or Main Group
Elements. The outermost electronic
Important properties of p-Block Elements are as ---
1. This block has elements in all physical states and maximum elements are non-metals..
2.Noble gas (Gr. 18) element have completely filled valence shell with ns2np6 configuration so they show very low chemical reactivity..
3. Halogens (Group 17) andchalcogens (Group 16). are another important groups of non-metals. These two
groups of elements have high negative electron gain enthalpies and readily add one or two electrons respectively to attain the stable noble gas configuration.
4. The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group.
The d-Block Elements ----Elements in which last electron enter in to the d-sub shell are known as d-block elements. They also known as Transition Elements--- because these metals form a bridge between the chemically active metals of s-block elements and the less active elements of Groups 13 and 14 and they have last two orbits are incomplete.
----The elements of Group 3 to Group 12 in the centre of the Periodic Table belong to this group. Their general electronic configuration is ns1 or 2(n-1)d1 to 10 .

Important properties of d-Block Elements are as ---
1.They are all metals.
2.They mostly form coloured ions, exhibit variable valence (oxidation states), para-magnetism and often used as catalysts.
Q. What is the reason “ Zn, Cd and Hg do not show most of the properties of transition elements”.
Ans. Zn, Cd and Hg which have the electronic configuration, ns2 (n-1)d10 so due to complete d-orbitals they do not show most of the properties of transition elements.

The f-Block Elements---- 
Elements in which last electron enter in to the f-sub shell are known as f-block elements. And they also known as Inner-Transition Elements.
---The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids, Th(Z = 90) – Lr (Z = 103) are belong to this group. Their general electronic configuration is or outer electronic configuration is (n-2) f 1-14 (n-1) d 0–1 ns 2.

Important properties of f-Block Elements are as ---
1.They are all metals. Within each series, the properties of the elements are quite similar because they have same electronic configuration for last two orbits.
2. They show large number of oxidation states.
3. Actinoid elements are radioactive and not available in the nature. They are artificially prepared, hence their chemistry is not fully studied.
The elements after uranium are called Transuranium Elements.

Metals, Non-metals and Metalloids
Elements can be classify as on the basis of their properties.
Properties of Metals--. Metals comprise more than 78% of all known elements.
1.Position in the periodic table---They are on the left side of the Periodic Table.
2.Metals are usually solids at roomtemperature [mercury is liquid]
3. Metals usually have high melting and boiling points.; [gallium and caesium have very low melting points (303K and 302K, respectively)].
4.They are good conductors of heat and electricity.
5.They are malleable (can be flattened into thin sheets by hammering)
6.They are ductile (can be drawn intowires).
Properties of Non-Metals
1.They are located at the top right hand side of the Periodic Table.
2. Non-metals are usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions).
3.They are poor conductors of heat and electricity.
4. Most nonmetallic solids are brittle and are neither malleable nor ductile.
---- The elements become more metallic as we go down a group; the nonmetallic character increases as one goes from left to right across the Periodic Table.
The change from metallic to non-metallic character is not abrupt as shown by the thick zig-zag line in Fig. The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) bordering this line and running diagonally
across the Periodic Table show properties that are characteristic of both metals and nonmetals. These elements are called Semi-metals or Metalloids.
There are some patterns in the physical and chemical properties of elements as we descend in a group or move across a period in the Periodic Table, e.g.
1. Reactivity of the elements---within a period, chemical reactivity tends to be high in Group 1 metals, lower in elements towards the middle of the table, and increases to a maximum in the Group 17 non-metals.
Likewise within a group of representative metals (say alkali metals) reactivity increases on moving down the group, whereas within a group of non-metals (say halogens), reactivity decreases down the group.
Q. Why do the properties of elements follow some trends? And how can we explain periodicity?
Ans. Atomic structure and properties of the atom, i.e. in terms of number of electrons and energy levels of an atom
we can explain the periodic trends in certain physical and chemical properties of elements.
2.Trends in Physical Properties----Physical properties of elements such as melting and boiling points, heats of
fusion and vaporization, energy of atomization, atomic and ionic radii, ionization enthalpy, electron gain enthalpy
and electronegativity. etc. which show periodic variations.
(a) Atomic Radius
Q. It is very difficult to measure the size/radius of an atom, why?
Ans. 1. Because the size of an atom (~ 1.2 Å i.e., 1.2 × 10–10 m in radius) is very small.
2.The electron cloud surrounding the atom does not have a sharp boundary.
So the determination of the atomic size cannot be precise or there is no practical way by which the size of an individual atom can be measured.
(However, an estimate of the atomic size can be made)
Covalent Radius”---Half of the distance between the atoms (center to center) in the combined state of a non-metallic element when their atoms are bound together by a single bond in a covalent molecule e.g. the bond distance in the chlorine molecule (Cl2) is 198 pm and half this distance (99 pm), is taken as the atomic radius of chlorine.
Metallic Radius”--- Half of the inter nuclear distance separating the metal cores in the metallic crystal.
e.g. The distance between two adjacent copper atoms in solid copper is 256 pm; hence the metallic radius of copper is assigned a value of 128 pm.
Atomic Radius word refers to both covalent or metallic radius depending on whether the element is a non-metal or a metal. Atomic radii can be measured by X-ray or other spectroscopic methods.
Trend in periodic table--- It is in terms of nuclear charge and energy level.
The atomic size generally decreases across a period, because within the period the outer electrons are in the same valence shell and the effective nuclear charge increases as the atomic number increases resulting in the increased attraction of electrons to the nucleus.
Within a family or vertical column of the periodic table or in a group the atomic radius increases regularly with atomic number because the principal quantum number (n) increases and the valence electrons are farther from the nucleus.
Because the inner energy levels are filled with electrons, which serve to shield the outer electrons from the pull of the nucleus. Consequently the size of the atom increases.

Note--- The atomic radii of noble gases are not considered here. Being monoatomic, their van der Waals radii (non-bonded radii) values are very large.

Q. When we go left to right across the period atomic radii first decreases and then abruptly increases for inert gases. Explain well.
Ans. In case of other elements of period we consider metallic and covalent radii i.e. in bonded form but for inert gases we consider vander wall’s radii i.e. in non-bonded form.
(b) Ionic Radius
The removal of an electron from an atom results in the formation of a cation, whereas gain of an electron leads to an anion.
The ionic radii is the distances between cations and anions in ionic crystals.
In general, the ionic radii of elements exhibit the same trend as the atomic radii.
A cation is smaller than its parent atom because it has lesser electrons than its atom while its nuclear charge remains the same, e.g. the atomic radius of sodium is 186 pm compared to the ionic radius of 95 pm for Na+.
The size of an anion will be larger than that of the parent atom because it has more electrons than its atom while its nuclear charge remains the same. The addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge.
e.g.the ionic radius of fluoride ion (F ) is 136 pm whereas the atomic radius of fluorine is only 64 pm. On the other hand,
Isoelectronic species--Atoms and ions which contain the same number of electrons, we call them isoelectronic species..e.g. O2–, F, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges. The cation with the greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. Anion with the greater negative charge will have the larger radius, e.g. O2–> F> Na+> Mg2+
(c) Ionization Enthalpy [ΔH I.E.]
It is the quantitative measure of the tendency of an element to lose electron.
Ionization Enthalpy. is the energy required to remove an electron from an isolated gaseous atom (X) in its ground state.
First ionization enthalpy is the energy required to remove first electron from an isolated gaseous atom (X) in its ground state.
for an element X . the first enthalpy change (ΔiH1) can be depicted as--
Units of ionization enthalpy is kJ mol–1.
Second ionization enthalpy is the energy required to remove second electron from an monovalent cation (X+). in its ground state. the second enthalpy change (ΔiH2) can be depicted as--
Energy is always required to remove electrons from an atom and hence ionization enthalpies are always positive.
The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom. In the same way the third ionization enthalpy will be higher than the second and so on.
Ionization Enthalpy is maximum at the noble gases which have closed electron shells and very stable electron configurations. And minimam occur at the alkali metals and their low ionization enthalpies can be correlated with their high reactivity.
First ionization enthalpy generally increases as we go across a period and decreases as we descend in a group.
To understand these trends, we have to consider two factors :
  1. the attraction of electrons towards the nucleus, and
  2. the repulsion of electrons from each other.
The effective nuclear charge experienced by a valence electron in an atom will be less than the actual charge on the nucleus because of “shielding” or “screeningof the valence electron from the nucleus by the intervening
core electrons. For example, the 2s electron in lithium is shielded from the nucleus by the inner core of 1s electrons.
As a result, the valence electron experiences a net positive charge which is less than the actual charge of +3. In general, shielding is effective when the orbitals in the inner shells are completely filled. e.g.alkali metals which have a lone ns-outermost electron preceded by a noble gas electronic configuration.
Anomalous behavior of some elementsCase-IBoron (Z = 5) has slightly lesser ionization enthalpy than that of beryllium (Z = 4) even though the former has a greater nuclear charge,--
---When we consider the same principal quantum level, an s-electron is attracted to the nucleus more than a p-electron. In beryllium, the electron removed during the ionization is an s-electron whereas the electron removed during ionization of boron is a p-electron. The penetration of a 2s-electron to the nucleus is more than that of a 2p-electron; hence the 2p electron of boron is more shielded from the nucleus by the inner core of electrons than the 2s electrons of beryllium. Therefore, it is easierto remove the 2p-electron from boron compared to the removal of a 2s- electron from beryllium. Thus, boron has a smaller first ionization enthalpy than beryllium.
Case-II-- Oxygen (Z = 8) has slightly lesser ionization enthalpy than that of Nitrogen(Z = 7) even though the former has a greater nuclear charge--
---because in the nitrogen atom, three 2p-electrons reside in different atomic orbitals (Hund’s rule) whereas in the oxygen atom, two of the four 2p-electrons must occupy the same 2p-orbital resulting in an increased electron-electron repulsion. Consequently, it is easier to remove the fourth 2p-electron from oxygen than it is, to remove
one of the three 2p-electrons from nitrogen i.e. half filled orbitals are most stable (Hund’s rule)
(d) Electron Gain Enthalpy
When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change during this process is known as the Electron Gain Enthalpy egH). represented by equation as--
Electron gain enthalpy provides a measure of the ease with which an atom adds an electron to form anion Depending on the element, the process of adding an electron to the atom can be either endothermic or exothermic. For many elements energy is released when an electron is added to the atom and the electron gain enthalpy is -ve e.g. group 17 elements (the halogens) have very high negative electron gain enthalpies because they can attain stable noble gas electronic configurations by picking up an electron,
---but noble gases have large positive electron gain enthalpies because the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration.
(Note--The variation in electron gain enthalpies of elements is less systematic than for ionization enthalpies.) Electron gain enthalpy becomes more negative with increase in the atomic number across a period because the effective nuclear charge increases from left to right across a period.
Electron gain enthalpy to become less negative as we go down a group because the size of the atom increases and the added electron would be farther from the nucleus.
Electron gain enthalpy of O or F is less negative than that of the succeeding element S and Cl.
---This is because when an electron isadded to O or F, the added electron goes to the smaller n = 2 quantum level and suffers significant repulsion from the other electrons present in this level due to small size of n = 2 orbit than
n = 3 orbit.. For the n = 3 quantum level (S or Cl), the added electron occupies alarger region of space and the electron-electron repulsion is much less.
(e) Electronegativity
A qualitative measurement of the ability of an atom in a chemical compound to attract shared electron pair to itself is called electronegativity.
Unlike ionization enthalpy and electron gain enthalpy, it is not a measureable quantity.
However, a number of numerical scales of electronegativity of elements viz., Pauling scale, Mulliken-Jaffe scale, Allred-Rochow scale etc. widely used is the Pauling scale. Linus Pauling,
It is maximum for F > O > N.
The electronegativity of any given element is not constant; it varies depending on the element to which it is bound, electronegativity also increases in a period as the atomic radius.decreases and the nucleus charge increases and its values decrease with the increase in atomic radii down a group. The trend is similar to that of ionization enthalpy.
Fig: The periodic trends of elements in the periodic table
Hence Non-metallic elements have strong tendency to gain electrons. Therefore, electronegativity is directly related to that non-metallic properties of elements or electronegativity is inversely related to the metallic properties of elements. Thus, the increase in electronegativities across a period there is an increase in non-metallic properties (or decrease in metallic properties) of elements.
Periodic Trends in Chemical Properties
In this section we shall study the periodicity of the valence state shown by elements and the anomalous properties of the second period elements (from lithium to fluorine).
(a) Periodicity of Valence or Oxidation States
It can be understood in terms of their electronic configurations. The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals and / or equal to eight minus the
number of outermost electrons as shown below.
Nowadays the term oxidation state is frequently used for valence.
Let us consider the two oxygen containing compounds: OF2 and Na2O. The order of electronegativity of the three
elements involved in these compounds is F > O > Na. F in OF2 molecule. Being highest electronegative element, is
given oxidation state –1. and oxygen exhibits oxidation state +2. In Na2O, oxygen being more electronegative shows oxidation state –2. and sodium is given oxidation state +1. Thus, the oxidation state of an element in a particular compound can be defined as the charge acquired by its atom on the basis of electronegative consideration from other atoms in the molecule. There are many elements which exhibit variable valence. This is particularly characteristic of transition elements and actinoids.
(b) Anomalous Properties of Second Period Elements
i) Diagonal relationship
The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) i.e. elements of II period differs in many respects from the other members of their respective group and they show close similarity with diagonally present right hand side elements of III period This sort of similarity is commonly known as diagonal relationship in the periodic properties. e.g. lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals, form compounds with pronounced covalent character; the other members of these groups predominantly form ionic compounds. In fact the behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively.
Q.What are the reasons for the different chemical behaviour of the first member of a group of elements in the s- and p-blockscompared to that of the subsequent members in the same group?
Ans. The anomalous behaviour is due to (a) their small size (b) large charge/radius ratio and (c) high electro-
negativity of the (d) The first member of each group has only four valence orbitals (2s and 2p) available for bonding, whereas the second member of the groups have nine valence orbitals (3s, 3p, 3d).
(i) This is the reason the maximum covalency of the first member of each group is 4 (e.g., boron can only formF4-] whereas the other members of the groups can expand theirvalence shell to accommodate more than four pairs of electrons e.g., aluminium forms [AlF6-3]
(ii) The first member of p-block elements displays greater ability to form pπ – pπ multiple bonds to itself (e.g.C = C, C ≡ C, N = N, N ≡ N ) and to other second period elements (e.g., C = O, C = N, C ≡ N, N = O) compared to subsequent members of the same group.
Periodic Trends and ChemicalReactivity
We know the periodicity is related to electronic configuration and all chemical and physical properties are a dependent on the electronic configuration of elements. Now We try to explore relationships between these fundamental properties of elements with their chemical reactivity.
1. The atomic and ionic radii, as we know, generally decrease in a period from left to right. As a consequence, the ionization enthalpies generally increase and electron gain enthalpies become more negative across a period. In other words, the ionization enthalpy of the extreme left element in a period is the least and the electron gain enthalpy of the
element on the extreme right is the highest negative (note : noble gases having completely filled shells have rather positive electron gain enthalpy values). This results into high chemical reactivity at the two extremes and the lowest in the centre. Thus, the maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of a cation and at the extreme right (among halogens) shown by the gain of an electron forming an anion. This property can be related with the reducing and oxidizing behaviour of the elements. And it can be directly related to the metallic and non-metallic character of elements. Thus, the metallic character of an element, which is highest at the extremely left decreases and the non-metallic character increases while moving
from left to right across the period. The chemical reactivity of an element can be best shown by its reactions with oxygen and halogens.
Elements on two extremes of a period easily combine with oxygen to form oxides. The normal oxide formed by the element on extreme left is the most basic (e.g., Na2O), whereas that formed by the element on extreme right is the most acidic (e.g., Cl2O7). Oxides of elements in the centre are amphoteric (e.g., Al2O3, As2O3) or neutral (e.g., CO,
NO, N2O). Amphoteric oxides behave as acidic with bases and as basic with acids, whereas neutral oxides have no acidic or basic properties.
Among transition metals (3d series), the change in atomic radii is much smaller as compared to those of representative elements across the period. The change in atomic radii is still smaller among inner-transition metals
(4f series). The ionization enthalpies are intermediate between those of s- and p-blocks. As a consequence, they are less electropositive than group 1 and 2 metals.
In a group, the increase in atomic and ionic radii with increase in atomic number generally results in a gradual decrease in ionization enthalpies and a regular decrease in electron gain enthalpies in the case of main group elements. Thus, the metallic character increases down the group and non-metallic character decreases. This trend can be related with their reducing and oxidizing property


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