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Physical Chemistry

Language of Chemistry Stoichiometry Atomic Structure Oxidation and Reduction Periodic Table Acid and Base Electronic Theory of Valency Equivalent Weight and Atomic Weight

Physical Chemistry-II

Non-Metals

Inorganic Chemistry

Periodic Table

Introduction:

-The periodic table is an arrangement of all the elements in accordance with their increasing atomic number and recurring chemical properties.

-They are arranged in a tabular arrangement wherein a row is a period and a column is a group.

-Elements are arranged from left to right and top to bottom in the order of their increasing atomic numbers. Thus,

Elements in the same group will have the same valence electron configuration and hence, similar chemical properties.
Whereas, elements in the same period will have an increasing order of valence electrons. Therefore, as the energy level of the atom increases, the number of energy sub-levels per energy level increases.

-The first $94$ elements of the periodic table are naturally occurring, while the rest from $95 t o 118$ have only been synthesized in laboratories or nuclear reactors.

Mendeleev Periodic Table:

Dmitri Ivanovich Mendeleev $(1834 - 1907)$ in $1869$ published his periodic table.
He widely referred as the father of the periodic table put forth the first iteration of the periodic table similar to the one we use now.
Mendeleev stated, “The physical and chemical properties of the elements are a periodic function of their atomic weights.’’
His arrangement of the table was in the form of vertical columns and horizontal rows, which were named as groups and periods.
Table was divided into $7$ horizontal rows (periods) in which Period $1$ : has $2$ elements; Period $2 an d 3$ : has $8$ ; Period $4 an d 5$ : has $18$ elements; Period $6$ : has $32$ elements.
Table was divided into $8$ vertical columns (groups) and a zero group.
Each group $(I - V II)$ was further divided into two sub groups.
It is differ from modern periodic table in one aspect. i.e. Mendeleev modeled his periodic table on the basis of increasing atomic mass, whereas, the modern periodic law is based on the increasing order of atomic numbers.
Correction of faulty valency and thus atomic weight of some elements were done. e.g. Corrected: $(B e = 4.5 \times 3 = 13.5)$ Previously: $(B e = 4.5 \times 2)$
Even though Mendeleev’s periodic table was based on atomic weight, he was able to predict the discovery and properties of certain elements. Mendeleev also predicted elements like:

i) $e ka - S i l i co n \to G er mani u m$ $(G e)$ ii) $e ka - A l u mini u m \to G a ll i u m$ $(G a)$ iii) $e ka - B oro n \to S c an d i u m$ $(S c)$

During his time only around half of the elements known to us now were known, and most of the information known about the elements were inaccurate.

Demerits of Mendeleev Periodic Table

-Position of Isotopes, Isobars, Lanthanides and Actinides not fixed.

-Dissimilar elements placed in the same group i.e. alkali metal elements and coinage metals $(C u, A g, A u)$ .

-Anomalous Pairs $(a t o mi c wt . o f C o > a t o mi c wt . o f N i)$

-Position of Hydrogen.

Modern Periodic Table:

Moseley in 1913, founded that the atomic number is the fundamental property of an atom.
It states, "The physical and chemical properties of element are the periodic function of their atomic number".
In this table, chemical properties have been correlated with electronic configuration of elements.

Long Form Periodic Table:

Constructed by Bohr.
It is an arrangement of $144$ elements (which also includes theoretical elements which are predicted but yet to be discovered) into periods and groups.
The groups have elements with similar electronic configuration in their valence shell so they exhibit similar chemical properties.
There are a total of $18$ groups in the modern periodic table.
The elements are arranged in $7$ periods in the periodic table chart. The period number indicates the highest principal quantum number $(n)$ of the elements in that particular period. The first period has $2$ elements and the rest periods have $8, 8, 18, 18 an d 32$ elements, respectively. The $14$ elements of both sixth and seventh period are positioned in the two bottom rows in the periodic table. They are called Lanthanides and Actinides respectively.

Elemental Electronic Configuration-Periods:

-The consecutive periods in the Periodic Table are connected with the filling of the next higher principal energy level.

In the first period, there are two elements whose valence shell has the lowest principal energy level i.e. $n = 1$ . The elements in the first period are hydrogen $(1 s^{1})$ and helium $(1 s^{2})$ .

In the second period, $(n = 2)$ , $2 s$ and $2 p$ orbitals are filled. For example, the first element in the second period is Lithium $(Z = 3)$ and here the third electron enters the $2 s$ orbital. Second period has $8$ elements.
The third period $(n = 3)$ has $8$ elements and valence electrons are present in $3 s an d 3 p$ orbitals.
The fourth period $(n = 4)$ starts at Potassium, where the valence electron fills the $4 s$ orbital. Next $3 d$ orbitals are filled since they are energetically favourable than $4 p$ . Such elements are called $3 d$ transition elements. Then $4 p$ orbitals are filled and hence the fourth period has $18$ elements.
The fifth period $(n = 5), 4 d$ the transition series starts at yttrium $Z = 39$ .

.This period ends with the filling up of the $5 p$ orbitals. It also has $18$ elements.
The sixth period $(n = 16)$ has $32$ elements and here the electrons fill $6 s, 4 f, 5 d, an d 6 p$ orbitals. The elements with valence electrons filling the $4 f$ orbitals are called the lanthanide series.
The seventh period $(n = 7)$ element electrons fill in $7 s, 5 f, 6 d, an d 7 p$ orbitals and they contain the man-made radioactive elements. The filling up of the $5 f$ orbitals give rise to the $5 f -$ inner transition series known as the actinide series.

Elemental Electronic Configuration-Groups:

Elements in a group have similar valence shell electronic distribution and hence similar chemical properties.
The Group $1$ elements are also called alkali metals and have $n s^{1}$ valence shell electronic configuration. For example, Lithium $(Z = 1)$ has $1 s^{2} 2 s^{1}$ valence shell electronic configuration.
The elements are further classified into s-block, p-block, d-block, and f-block depending on the orbitals in which valence electrons are filled. The two exceptions are Hydrogen and Helium.

The s-Block Elements:

Group 1 (alkali metals) and Group 2 (alkaline earth metals) elements which have $n s^{1} an d n s^{2}$ valence shell electronic configurations are known as s-Block elements.
They lose the outermost electron to form $1 + i o n or 2 + i o n$ for alkali and alkaline metals, respectively.
They are thus reactive with low ionisation energy. As we go down the group, the reactivity and metallic character increases.

The p-Block Elements:

The p-Block Elements consist of elements of groups $13 t o 18$ .
The outer shell configuration varies from $n s^{2} n p^{1} t o n s^{2} n p^{6}$ in each period.
All the last period elements are noble elements and its outer orbitals are completely filled by electrons.
They have very low reactivity. The group $16$ (Chalcogens) and group $17$ (halogens) have high electron gain energies and can add one or two electrons to attain a stable outermost configuration.

The d-Block Elements (Transition Elements):

The d-block elements consist of Group $3 t o 12$ in the Periodic Table.
These elements are characterised by the filling of inner d orbitals by electrons.
The general outer electronic configuration of the d-block is $(n - 1) d^{(1 - 10)} n s^{(0 - 2)}$ .
They are all metals and form coloured ions.
They exhibit variable oxidation states and para-magnetism. However, $Z in c, C a d mi u m, an d M erc u ry$ have electronic configuration, $(n - 1) d^{10} 10 s^{2}$ and they do not behave like transition elements.

The f-Block Elements (Inner-Transition Elements):

The last two row elements down of the periodic table are called Lanthanides and Actinides
They have valence shell electronic configuration $(n - 2) f^{1} - 14 (n - 1) d^{(0 - 1)} s^{2}$ .
The last electron is filled in the f-orbital.
They are all metals and in each series the properties of the elements are similar.
The actinides can have a large number of oxidation states. Hence their chemistry is complicated.
These elements are radioactive.

Periodic Properties:

i) Atomic Radii:

It is the distance of nucleus to outermost shell.
Factors affecting atomic radii:

a) Effective nuclear charge is inversely proportional to atomic radii.(due to contraction)

b) Shielding effect is directly proportional to atomic radii(repulsion causes distance increase variation of atomic radii)

Variation of atomic radii:

i) Along group:

-Going from top to bottom, atomic number increases, size increases and hence atomic radii increases.

ii) Along period:

-Going from left to right , atomic radius decreases due to increase in effective nuclear charge.

iii) Decreases with number of bonds.

c) In isoelectric ions:

-Ions having same no. of electrons is called isoelectric ions.

-Atomic radius is inversely proportional to no. of protons.

-e.g. $O^{--} > P^{--}, N a^{+}, M g^{++}, S^{--}, C l^{-}, e t c$ .

-Radius of isoelectronic ions is determined by $z / e$ ratio, smaller the $z / e$ ratio, greater the size.

d) On atoms and ions:

-Decreases in no. of electrons increases effective nuclear attraction and hence decreases the size. So, $N a^{+} < N a, M g^{++} < M g$ and increase in number of electrons decreases effective nuclear attraction and hence size increases.

-E.g. $O^{--} > O^{-}; F^{-} > F, e t c$

ii) Ionization Potential:

-The amount of energy required to release most loosely bounded electron or outermost electron of an isolated atom and the amount of energy required to release the second electron from same isolated atom is called second ionization potential.

-Generally $I . P_{2} > I . P$ because energy required is higher.

Factors affecting I.P:

I.P is inversely proportional to size because less is the size more tightly electron is held and more energy is required.
Shielding effect is inversely proportional I.P. Greater the shielding effect (no. of $e^{-}$ in inner shell) more is the repulsion and hence less is I.P.
Half-filled and completely filled orbitals have more I.P as they are more stable than other.
More nuclear attraction(charge) causes decrease in size and hence causes increase in I.P.

Variation of I.P:

a) Along group:

-Decreases from top to bottom because of gradual increase in atomic size.

b) Along period:

-I.P goes on increasing along period due to decrease in size caused by shielding effect.

-Increases in Zero group.(Noble gas)

Note:

- $I . P . o f B e > I . P o f B$ (stable configuration)

- $I . P . o f N i t ro g e n > I . P o f O x y g e n$ (half-filled configuration)

- $C s$ has least I.P. among all metals.

iii) Electro-Negativity:

-Tendency of atom in a compound to attract a shared pair of electron toward itself.

Factors affecting electronegativity:

Atomic size is inversely proportional to electro-negativity. Greater the size less is the tendency to attract shared pair of electron toward itself.
Increase of effective charge increases the effective nuclear attraction and E.N. increases.
Greater is the $+ v e$ charge greater is the electro-negativity.i .e. $F e < F e^{++} < F e^{+++}$ and $X^{+} > X > X^{-}$ .

Variation of E.N:

a) Along group:

-Decreases from top to bottom as at. size increases from top to bottom.

b) Along period:

-Increases from left to right due to decrease in shielding effect.

iv) Electron affinity(E.A):

-Amount of energy releases when an electron is added to an isolated neutral atom.

-It is the capacity to form $- v e$ ion.

Factors affecting E.A:

Atomic size is inversely proportional to E.A.
Nuclear charge is directly proportional to E.A.
Half filled and full filled orbitals are very stable. As a result no tendency to add electron on it.

Variation of E.A.

a) Along group:

-Decreases from top to bottom as at. size increases from top to bottom.

b) Along period:

-Increases from left to right.

Diagonal Relationship:

Some elements resembles with properties of elements one step down in next group which is called diagonal relationship.

Elements of $2 n d$ period: $L i ↘ B e ↘ B ↘ C$

Elements of $3 r d$ period: $N a M g A l S i$