Classification of Elements and Periodicity in Properties
The Modern Periodic Table s a good tabular show of all compound elements. It is an important research tool in chemistry as well as other areas of technology. It could be used in several ways, despite the fact that generally speaking, we believe of the periodic graph as an symbol of chemistry. In this article there are illustrations of using the periodic schedule to help you learn and clear your doubts about the distinct elements. As soon as you know the various brands for every element, you can certainly make them much more common.

Genesis of Periodic Classification Dobereiner’s Triads
In 1829, Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and the third elements. A few triads proposed by him are listed.

Limitations of Dobereiner’s Triads
The triads given by Dobereiner were helpful in grouping some elements with similar characteristics together, but he could not arrange all the elements known at that time into triads.

• Newlands’ Law of Octaves
John Newlands proposed the law of octaves by stating that when elements are arranged in order of increasing atomic masses, every eighth element has properties similar to the first. Newlands called it law of octaves because similar relationship exists in the musical notes also.
This can be illustrated as:

Limitations of Newlands’ Law of Octaves

This classification was successful only up to the element calcium. After that, every eighth element did not possess the same properties as the element lying above it in the same group.

When noble gas elements were discovered at a later stage, their inclusion in the table disturbed the entire arrangement.

• Mendeleev’s Periodic Table

Mendeleev’s Periodic Law: The physical and chemical properties of the elements are a periodic function of their atomic masses.

Mendeleev arranged the elements known at that time in order of increasing atomic masses
and this arrangement was called periodic table.
Elements with similar characteristics were present in vertical rows called groups. The horizontal
rows were known as periods.

Description of Mendeleev’s Periodic Table

(i) In the periodic table, the elements are arranged in vertical rows called groups and horizontal rows known as periods.
(ii) There are nine groups indicated by Roman Numerals as I, II, III, IV, V, VI, VII, VIII and zero. Group VIII consists of nine elements which are arranged in three triads. The zero group contains elements belonging to inert gases or noble gases and elements present have zero valency.
(iii) There are seven periods (numbered from 1 to 7) or, horizontal rows in the Mendeleev’s periodic table.
Importance of Mendeleev’s Periodic Table

This made the study of the elements quite systematic in the sense that if the properties of one element in a particular group are known, those of others can be predicted.
This helped to a great extent in the discovery of these elements at a later stage.
Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and properties.

Defects in Mendeleev’s Periodic Table

Hydrogen has been placed in group IA along with alkali metals. But it also resembles Halogens of group VII A in many properties. Thus, its position is the Mendeleev’s periodic table is controversial.

Although the elements in the Mendeleev’s periodic table have been arranged in order of their atomic masses, but in some cases the element with higher atomic mass precedes the element with lower atomic mass.

The isotopes of an element have different atomic masses but same atomic number. As the Periodic table was framed on the basis of increasing atomic masses of the elements, different positions have been allotted to all the isotopes of a particular element.

According to Mendeleev, the elements placed in the same group must resemble in their properties. But there is no similarity among the elements in the two sub-groups of a particular group.

In some cases, elements with similar properties have been placed in different groups.

Lanthanoids and Actinoids were placed in two separate rows at the bottom of the periodic table without assigning a proper reason.

There is no proper explanation offered about why the elements placed in a group show resemblance in their properties.

• Modern Periodic Law

Physical and chemical properties of the elements are the periodic function of their atomic numbers.

• Present Form of the Periodic Table (Long form of Periodic Table)

The long form of periodic table, also called Modem Periodic Table, is based on Modern periodic law. In this table, the elements have been arranged in order of increasing atomic numbers.

• Structural Features of the Periodic Table

Blocks

The Periodic Table is divided into four zones called blocks :

s- Block, p- Block, d- Block and f- Block

Groups

The long form of periodic table also consists of the vertical columns called groups.
There are in all 18 groups in the periodic table.

Unlike Mendeleev periodic table, each group is an independent group.

Characteristics of groups:

All the elements present in a group have same general electronic configuration of the atoms.

The elements in a group are separated by definite gaps of atomic numbers (2, 8, 8,18, 18,32,32).

The atomic sizes of the elements in group increase down the group due to increase in the number of shells.

The Physical properties of the elements such as melting point, boiling point, density, solubility etc., follow a systematic pattern.

The elements in each group have similar chemical properties in general.

Periods

The Horizontal rows in a periodic table are known as periods.

There are in all seven periods in the long form of periodic table.

The first period is the shortest containing 2 Elements only.

The last two periods, 6th and 7th are the longest containing 32 elements each.

The first period resembles K shell with 2 elements.

The 2nd and 3rd periods resemble L shell with 8 elements each.

The 4th and 5th periods resemble M shell with 18 elements each.

The 6th and 7th periods resemble N shell with 32 elements each.

Characteristics of Periods:

In all the elements present in a period, the electrons are filled in the same valence shell.

The atomic sizes along a period decrease from left to right, in general.

s-Block Elements

General electronic configuration:

The general electronic configuration of elements in this block is

ns1-2

Characteristics of s-block elements:

All these elements produce basic salts.

Group 1 elements are called Alkali metals.

Group 2 elements are called Alkaline Earth Metals.

All the elements are soft metals.

They have low melting and boiling points.

They are highly reactive.

Most of them impart colours to the flame.

They generally form ionic compounds.

They are good conductors of heat and electricity.

p-Block Elements

General electronic configuration:

The general electronic configuration of elements in this block is

ns2np1-6

Characteristics of p-block elements:

The compounds of these elements are mostly covalent in nature.

They show variable oxidation states.

The non-metallic character of the elements increases along a period from left to right.

The reactivity of elements down a group decreases in general.

At the end of each period there is a noble gas element with a closed valence shell ns2 np6 configuration.

Metallic character increases down a group.

d-Block Elements

General electronic configuration:

The general electronic configuration of elements in this block is

(n -1) d1-10 ns0-2

The d-block elements are known as transition elements because they have incompletely filled d-orbitals in their ground state or in any of the oxidation states.

Characteristics of d-block elements:

They are all metals with high melting and boiling points.

The compounds of the elements are generally paramagnetic in nature.

They mostly form coloured ions, exhibit variable valence (oxidation states).

They are often used as catalysts.

f-Block Elements

General electronic configuration:

The general electronic configuration of elements in this block is

(n – 2) f1-14 (n -1) d0-1 ns2

They are known as inner transition elements because in the transition elements of d-block, the electrons are filled in (n – 1) d sub-shell while in the inner transition elements of f-block the filling of electrons takes place in (n – 2) f subshell, which happens to be one inner subshell.

Characteristics of f-Block elements:

The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids from the elements Ce (Z = 58) to Lu (Z = 71) and Actinoids from the elements Th (Z = 90) to Lr (Z = 103).

These two series of elements are called Inner Transition Elements (f-Block Elements).

They are all metals. Within each series, the properties of the elements are quite similar.

Most of the elements of the actinoid series are radio-active in nature.

• Metals

Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table.

Metals are solids at room temperature.

Metal usually have high melting and boiling points.

They are good conductors of heat and electricity.

They are malleable and ductile.

• Non-metals

Non-metals are located at the top right hand side of the Periodic Table.

Non-metals are usually solids or gases at low temperature with low melting and boiling points.

They are poor conductors of heat and electricity.

The non-metallic character increases along a period from left to right.

Most non-metallic solids are brittle and are neither malleable nor ductile.

• Metalloids

The elements such as Silicon, Germanium, Arsenic, Antimony and Tellurium show the characteristics of both metals and non-metals.

These elements are also called semimetals or metalloids.

• Noble Gases

These are the elements present in group 18.

Each period ends with noble gas element.

All the members are of gaseous nature.

Because of the presence of fully occupied orbitals, they have very little tendency to take part in chemical combination.

These are also called inert gases.

Xenon is the only exception to have formed three different fluorides.

• Representative Elements

The elements of group 1 or alkali metals , elements of group 2 or alkaline earth metals and elements of group 13 to 17 constitute the representative elements.

They are all elements of s-block and p-block.

• Transition Elements

The transition elements include, all the d-block elements.

They are present in the centre of the periodic table between s and p-block elements.

• Inner Transition Elements

Lanthanoids or Lanthanides, the fourteen elements after Lanthanum, and Actinoids or Actinides, the fourteen elements after Actinium are called inner transition elements.

They are also called f-block elements.

The elements after Uranium are also called transuranic elements.

Periodic Trends in Properties of Elements

Trends in Physical Properties

Atomic Radii:

It is defined as the distance from the centre of the nucleus to the outermost shell containing the electrons

Depending upon whether an element is a non-metal or a metal, three different types of atomic radii are used. These are:

Covalent radius
Ionic Radius
van der Waals radius
Metallic radius.

Covalent Radius:

It is equal to half of the distance between the centres of the nuclei of two atoms held together by a purely covalent single bond.

Ionic Radius:

It may be defined as the effective distance from the nucleus of an ion up to which it has an influence in the ionic bond.

van der Waals Radius:

Atoms of Noble gases are held together by weak van der Waals forces of attraction. The van der Waals radius is half of the distance between the centre of nuclei of atoms of noble gases.

Metallic Radius:

It is defined as half of the inter-nuclear distance between the two adjacent metal ions in the metallic lattice.

• Variation of Atomic Radius in the Periodic Table

Variation in a Period: Along a period, the atomic radii of the elements generally decreases from left to right.

Variation in a group: The atomic radii of the elements in every group of the periodic table increases as we move downwards.

• Ionic Radius

The ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals.

In general, the ionic radii of elements exhibit the same trend as the atomic radii.

Cation: The removal of an electron from an atom results in the formation of a cation. The radius of cation is always smaller than that of the atom.

Anion: Gain of an electron leads to an anion. The radius of the anion is always larger than that of the atom.

Isoelectronic Species: Some atoms and ions which contain the same number of electrons, we call them isoelectronic species.

For example, O2-, F–, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges.

Effective Nuclear Charge

The net nuclear charge on an electron of an atom is called the effective nuclear charge on that electron. The enthalpy of an atom depends on the effective nuclear charge on the incoming electron.
When the effective nuclear charge on the electron is high, it is more strongly attracted to the atom and the enthalpy will be high.

Oxidation Number

Oxidation number is the number of electrons an atom gains or loses in a chemical compound to form a chemical bond with another atom. It is the charge an atom is assumed to have when all other atoms are removed from it as ions.

The loss or gain of an electron is also known as the degree of oxidation.
The combination of the lose and gain of an electron is known as Redox reaction.
Oxidation number is also known as Oxidation state, which can be zero, positive or negative.

Examples: The oxidation number of atoms in O2, O3, P4, and S8 is zero.

Important Points for Determining Oxidation Number

All the elements in the elementary state or molecular states have oxidation number zero, e.g., He, Ne, Xe, Ar, H2 , N2, Cl2, S8, P4 etc.
The oxidation number of metal in amalgams is zero.
The oxidation number of metal in amalgams is zero.
The algebraic sum of the oxidation numbers of all the atoms in an uncharged (neutral) compound is zero.
If the coordinate bond is directed from more electronegative to the less electronegative atom then its contribution is zero for both the atoms.
In an ion, the algebraic sum is equal to the charge on the ion.
In all compounds. except for ionic metallic hydrides, the oxidation number of hydrogen is +1.
The oxidation number of alkali metals is +1 and for alkaline earth metals is + 2.
In case of a coordinate bond, it gives +2 value of oxidation number to the less electronegative atom.
Metal hydrides like NaH, MgH2, CaH2, LiH, etc have the oxidation number of hydrogen as -1.
As fluorine is the most electronegative element, it always has an oxidation number of -1 in all of its compounds.
In compounds containing oxygen, the oxidation number of oxygen is – 2 except in peroxides (-1) such as Na2O2, in OF2 and in O2 F2 (+2 and +1 respectively).
In case of a coordinate bond, it gives -2 values to the more electronegative atom when the coordinate bond is directed formless electronegative atom to more electronegative atom.
For p-block elements [Except F and O], the highest oxidation number is equal to their group number and lowest oxidation number is equal to the group number minus eight.
In transition elements, the lowest oxidation number is equal to the number of ns electrons and highest oxidation number is equal to the number of ‘ns’ and (n – l)d unpaired electrons.

Ionization Enthalpy

It is the energy required to remove an electron from an isolated gaseous atom in its ground state.

M (g) + I.E ⟶M+ (g) + e–

The unit of ionization enthalpy is kJ mol-1 and the unit of ionization potential is electron volt per atom.

Successive Ionization Enthalpies

If a gaseous atom is to lose more than one electron, they can be removed one after the other i.e., in succession and not simultaneously.

This is known as successive ionization enthalpy (or potential).

The example of successive Ionization Enthalpy for Nitrogen can be considered as follows:

• Variation of Ionization Enthalpies in the Periodic Table:

Variation of Ionization Enthalpy Along a Period

Ionization enthalpies are expected to increase along a period ,as the nuclear charge increases and the atomic size decreases from left to right.

Variation of Ionization Enthalpy in a Group

The ionization enthalpies of the elements decrease on moving from top to the bottom in any group.

The decrease in ionization enthalpies down any group is because of the following two factors.

There is an increase in the number of the main energy shells (n) on moving from one element to the other.

There is also an increase in the magnitude of the screening effect due to the gradual increase in the number of inner electrons.

• Electron Gain Enthalpy and Electron Affinity:

Electron gain enthalpy is also at times referred to as Electron affinity although in between them there is a very minute difference.

Electron gain enthalpy is defined as the amount of energy released when an electron is added to an isolated gaseous atom to form an anion. During the addition of an electron, energy can either be released or absorbed.

Electron affinity means a love for electron. It is the negative of the electron gain enthalpy.

Factors which affect electron affinity are atomic size and nuclear charge in general. As atomic radii increases, electron affinity increases. As nuclear charge increases, electron affinity increases. It decreases down a group and increases across a period.

Electron Gain Enthalpy is the energy released when an electron is added to an isolated gaseous atom so as to convert it into a negative ion. The process is represented as:

For majority of the elements the electron gain enthalpy is negative.

For example, the electron gain enthalpy for halogens is highly negative because they can acquire the nearest noble gas configuration by accepting an extra electron.

In contrast, noble gases have large positive electron gain enthalpies because the extra electron has to be placed in the next higher principal quantum energy level thereby producing highly unstable electronic configuration.

According to Thermodynamics, the relationship between electron gain enthalpy and electron affinity can be determined as follows.

Successive Electron Gain Enthalpies:

When an electron is added to a neutral isolated gaseous atom, the energy which is released is the first electron gain enthalpy of that element.

These are exothermic reactions in general.

X(g) +e ⁻ →X ⁻(g) + E

When an electron is added to an isolated gaseous anion with charge -1, then the released energy is called the second electron gain enthalpy of the element.

Electron gain enthalpy when done for another time, it is called successive electron gain enthalpy.

Successive electron gain enthalpy is positive for 2nd time because when electron is added for the 2nd time, it will experience repulsion from the electrons already present, so energy must be provided so that the atom can take up that electron.

Generally electrons from a gaseous atoms are lost in succession (i.e., one after another).

Similarly, these are also accepted one after the other, i.e., in succession.

After the addition of one electron, the atom becomes negatively charged and the second electron is to be added to a negatively charged ion.

But the addition of second electron is opposed by electrostatic repulsion and hence the energy has to be supplied for the addition of second electron.

Thus the second electron gain enthalpy of an element is positive.

For example, when an electron is added to oxygen atom to form O– ion, energy is released. But when another electron is added to 0¹⁻ ion to form O2- ion, energy is absorbed to overcome the strong electrostatic repulsion between the negatively charged 0– ion and the second electron being added.

Factors on which Electron Gain Enthalpy Depends

Atomic size: As the size of an atom increases, the distance between its nucleus and the incoming electron also increases and electron gain enthalpy becomes less negative.

Nuclear charge: With the increase in nuclear charge, force of attraction between the nucleus and the incoming electron increases and thus electron gain enthalpy becomes more negative.

Symmetry of the Electronic Configuration: The atoms with symmetrical configuration (having fully filled or half filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unstable.

In that case the energy will be needed and electron gain enthalpy (Δe𝔤H) will be positive. For example, noble gas elements have positive electron gain enthalpies.

Variation of Electron Gain Enthalpy Across a Period

Electron gain enthalpy becomes more negative with increase in the atomic number across a period.

Variation of Electron Gain Enthalpy in a Group

Electron gain enthalpy becomes less negative as we go down a group.

Non-metals usually have a more negative electron gain enthalpy than metals, as they are smaller in size than the metals in the same period.

The group of elements with the highest electron gain enthalpies are halogens, and chlorine has the highest enthalpies among halogens.

The above-mentioned orders are not always regular due to factors like stability of the atom, shielding effects, effective nuclear charge, and electronic repulsions.

Problems on Electron Gain Enthalpy (e.g.e) and Ionisation Enthalpy(I.E.)

Problem - 1:

The first ( $Δ H_{1}$ ) and second ( $Δ H_{2}$ ) ionisation enthalpies (in $k J m o l^{- 1}$ ) and the electron gain enthalpy ( $Δ_{e g} H$ ) (in $k J m o l^{- 1}$ ) of the elements I, II, III, IV and V are given below:

Element $Δ_{1} H_{1}$ $Δ_{1} H_{2}$ $Δ_{e g} H$
I 520 7300 -60
II 419 3051 -48
II 1681 3374 -328
IV 1008 1846 -295
V 2372 5251 +48
The most reactive metal and the least reactive non-metal of these are respectively.

I and V

V and II

II and V

IV and V

Solution:

Answer - C

I represents Li, II represents K, III represents Br, IV represents I, V represents He.

So, amongst these, II represents most reactive metal since Potassium is the most reactive metal of all and V represents least reactive non-metal, being a noble gas.
Also, II has the lowest value of first ionization enthalpy and V has positive electron gain enthalpy.

Problem - 2:

The first $\Delta _ i H_1$ and the second $\Delta _ i H_2$ ionization enthalpies (in $KJ mol ^{-1}$ )   $\Delta _{eg }H$ electron gain enthalpy (in $KJ mol ^{-1}$ ) of a few elements are given below:

Elements $\Delta H_1 \: \: \: \: \: \: \Delta H_2 \: \: \: \: \: \: \: \Delta _{eg }H$
I    520 7300 –60
II    419 3051 –48
III    1681 3374 –328
IV    1008 1846 –295
V 2372 5251 +48
VI 738 1451 –40
Why is  IV above the least reactive non-metal?

Solution:

The element IV has a high negative electron gain enthalpy $(\triangle_{eg}H)$ of -295
$KJ mol ^{-1}$ but its second ionization enthalpy( $\Delta _ i H_2$ ) is not so higher than the first ionization enthalpy $(\triangle_{i}H_{1})$ and hence is the least reactive non-metal.

Problem - 3:

$(I E)_{1}$ and $(I E)_{2}$ in kJ $m o l^{- 1}$ of K and Ca have been given as

$(I E)_{1}$
$(I E)_{2}$
K
419
3052
Ca
590
1145
The difference can be explained by:

stable inert gas configuration

hydrogen bonding

steric hindrance

van der wall effect

Solution:

Answer is A

$(I E)_{2}$ of K >>> $(I E)_{1}$ of K

since $K^{+}$ (after one electron is removed)has stable inert gas configuration as compared to corresponding values for Ca in which $C a$ ⁺ can lose second electron easily.

Problem - 4:

IE (in kJ $m o I^{- 1}$ ) of $L i$ and $B e$ are given below:

Z
$(I E)_{1}$
$(I E)_{2}$
$(I E)_{3}$
$L$
3
520.1
$X(1)$ 11810
$e$
4
899.3
$Y(2)$ 14810
From above table $X (1)$ is higher than $Y (2)$ . Justify.

Solution
The given statement is true.
X(1) is higher than Y(2).
The electronic configuration of Li is $1 s^{2} 2 s^{1}$ . It loses one electron to obtain stable electronic configuration of $1 s^{2}$ in which duplet is complete. It is electronic configuration of noble gas He. It is difficult to remove second electron from lithium as it will break the stable electronic configuration. Hence, X(1) has high value.

The electronic configuration of Be is $1 s^{2} 2 s^{2}$ . It loses one electron to obtain electronic configuration of $1 s^{2} 2 s^{1}$ . It is easy to remove second electron from beryllium as it will form the stable electronic configuration of noble gas He in which duplet is complete . Hence, Y(2) has low value.

Problem - 5:

Which of the following will have the most negative electron gain enthalpy and which the least negative?
P, S, Cl, F

Solution:

Halogens have the highest negative value of electron gain enthalpy. Here Cl and F both are halogens but electron gain enthalpy of Cl is more negative than F because of the very small size of F which causes inter-electronic repulsion among its electrons.
That is why Cl has the most negative value of electron gain enthalpy. P has the least negative one due to the larger size

Problem - 6:

Which of the following hast he most negative electron gain enthalpy?

$B r$

$S n$

$C a$

$M g$

Solution:

Generally, non-metals have more negative electron gain enthalpy than metals. This is because they gain electrons to complete their outermost orbital. Here, the rest elements are metals, Therefore, Br has the most negative electron gain enthalpy.

So, Correct option is A.

Problem - 7:

The electron gain enthalpy is least for?

$2 s^{2} 2 p^{5}$

$3 s^{2} 3 p^{5}$

$2 s^{2} 2 p^{4}$

$3 s^{2} 3 p^{4}$

Solution:

Electron gain enthalpy is defined as the amount of energy released when an electron is added to an isolated gaseous atom. During the addition of an electron, energy can either be released or absorbed .

So, as in option A, the 2p orbital needs only one electron to get isolated and stabilize .

Hence , option A is correct .

Problem - 8:

N have very high electron gain enthalpy.

True

False

Solution:

Correct option is B

N has low E.G.E. as its outermost shell is half-filled which makes it extra stable. Hence it has less tendency to accept an electron( e⁻).

Problem - 9:

Among halogens, the correct order of amount of energy released in electron gain (electron gain enthalpy) is:

$F > C l > B r > I$

$F < C l > B r > I$

$F > B r > C l > I$

$I > C l > B r > F$

Solution:

Correct option is B

Electron gain enthalpy becomes less negative on moving down a group due to increase in atomic size. However, the electron gain enthalpy of F is less negative than that of CI due to electron-electron repulsion in small-sized F atom.
Thus, the correct order is F<Cl>Br>I.
F, Cl, Br and I have ΔegH values(in kJmol⁻¹) -328, -349, -325 and -295 respectively.

Problem - 10:

Choose the correct order for the magnitudes of electron gain enthalpy.

$F > B r > O > S$

$F > S > B r > O$

$F > B r > S > O$

$S > F > B r > O$

Solution:

Correct option is C

E.A = F>Br>S>O
On moving from left to right along a period E.A increases due to more

Z_{e f f}

(

small size

),

but when small sized atom like Oxygen in its group faces more repulsion upon incoming electron. So, its

E . A

is less than

Sulphur .

Problem - 10:

Find the oxidation number of Mn is KMnO4 ?

Solution:

Let the oxidation number of Mn in KMnO4 be x.

We know that, the Oxidation number of K = +1 and of O = –2

And also,

(Oxidation number K) + (Oxidation number of Mn) + 4(Oxidation number of O) = 0

(+1) + (x) + 4(-2) = 0 or x = +7

Therefore, the oxidation number of Mn = +7

• Electronegativity

A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity.

Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity.

However, a number of numerical scales of electronegativity of elements viz, Pauling scale, Milliken- Jaffe scale, Allred Rochow scale have been developed.

The electronegativity of any given element is not constant; it varies depending on the element to which it is bound.

Along a Period
Electronegativity generally increases along a period, from left to right.

Down a Group
Electronegativity generally decreases down a group.

• Periodic Trends in Chemical Properties along a Period

Metallic character:

Decreases along a period, maximum on the extreme left (alkali metals).

Non-metallic character:

Increases along a period. (From left to right).

Basic nature of oxides:

Decreases along a period (From left to right).

Acidic nature of oxides:

Increases along a period (From left to right).

• Variation from Top to Bottom on Moving Down a Group

Metallic character: Generally increases because increase in atomic size and hence decrease in the ionization energy of the elements in a group from top to bottom.

Non-metallic character: Generally decreases down a group as electronegativity of elements decreases from top to bottom in a group.

Basic nature of oxides: Since metallic character or electro positivity of elements increases from top to bottom in a group, basic nature of oxides naturally increases.

Acidic character of oxides: Generally decreases as non-metallic character of elements decreases in going from top to bottom in a group.
Reactivity of metals. Generally increases down a group. Since tendency to lose electron increases.

Reactivity of non-metals: Generally decreases down the group, Higher the electro-negativity of non-metals, greater is their reactivity. Since electronegativity of non-metals in a group decreases from top to bottom, their reactivity also decreases.

Some Common Anomalies

1. Anomalous Properties of Second Period Elements

The first element of each of the group 1 (Lithium) and group 2 (Beryllium) and groups 13-17 (Boron to Fluorine) differs in many respect from the other members of their respective groups.

For example, lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals form compounds which have significant covalent character.

The other members of these groups, pre-dominantly form ionic compounds.

It has been observed that some elements of the second period show similarities with the elements of the third period placed diagonally to each other, though belonging to different groups.

This similarity in properties of elements placed diagonally to each other is called diagonal relationship.

Some Common Anomalies

2. Fluorine & Chlorine

The general expectation is that fluorine has more negative enthalpy than chlorine due to its small size and high electronegativity. But in reality it is the opposite, chlorine has more negative enthalpy than fluorine.

This is because fluorine has a small size, due to which its electron density is high, which leads to repulsion between electrons making it difficult for an electron to bond with the atom.

3. Carbon & Nitrogen

The electron gain enthalpy of carbon is more negative than that of nitrogen. This is because nitrogen has a half-filled p-subshell, which makes it more stable than carbon.

So, as nitrogen loses its stability on gaining an electron, it has a lower enthalpy than carbon, which becomes more stable on gaining an electron.

FAQs

Q-1. What is the basic difference in approach between Mendeleev's Periodic Law and the Modern Periodic Law?

Ans : The physical and chemical properties of elements are periodic functions of their atomic weights, according to Mendeleev's Periodic Law. The Modern Periodic Law, on the other hand, maintains that an element's physical and chemical properties are periodic functions of its atomic number.

Q-2.Why do elements in the same group have similar physical and chemical properties?

Ans :The quantity of valence electrons affects the physical and chemical properties of elements. The number of valence electrons in a group of elements is the same. As a result, physical and chemical properties of elements in the same group are comparable.

Q-3. Explain why cations are smaller and anions larger in radii than their parent atoms.

Ans : A cation possesses fewer electrons than its parent atom although has the same nuclear charge. That's why, a cation's nucleus attracts electrons more than its parent atom's nucleus. Hence, a cation is smaller than its parent atom in size. An anion, on the other hand, has one or more electrons than its parent atom, resulting in a higher electron repulsion and hence a reduction in effective nuclear charge. Hence, anions have a greater distance between their valence electrons and the nucleus than their parent atom. And this is why the radius of an anion is bigger than that of its parent atom.

Q-4. How does atomic radius vary in a period and in a group? How do you explain the variation?

Ans : Along a period, the atomic radius decreases from left to right. Because the effective nuclear charge increases as the atomic number increases from left to right along a period and the outer electrons present in the same valence shell are more attracted to the nucleus.
The atomic radius, on the other hand, tends to rise as you progress through the groups. This is due to increase of the distance between the nucleus and valence electrons as the primary quantum number (n) increases.

Q-5. What is the basic difference between the terms electron gain enthalpy and electronegativity?

Ans : The tendency of an isolated gaseous atom to accept an electron is measured by electron gain enthalpy, whereas the tendency of an atom in a chemical compound to attract a shared pair of electrons is measured by electronegativity.

Q-5. Would you expect the first ionization enthalpies for two isotopes of the same element to be the same or different? Justify your answer.

Ans : The number of electrons and protons (nuclear charge) in an atom determines its ionisation enthalpy. The protons and electrons in an element's isotopes are now the same. As a result, for two isotopes of the same element, the first ionisation enthalpy should be the same.

Q-6. What is electron gain enthalpy?

Ans: The energy released when an electron is added to an isolated gaseous atom is called electron gain enthalpy.

The energy released when the electron is added to a neutral isolated gaseous atom is first electron gain enthalpy, and when the electron is added to an isolated gaseous anion with -1 charge it is called second electron gain enthalpy and so on.

Q-7. Why is the second electron gain enthalpy of most atoms positive?

Ans: The second electron gain enthalpy of most of the atoms is positive due to electron-electron repulsions between already present electrons and the new one.

Q-8. Why are the electron gain enthalpies of all the noble gases positive?

Ans: Noble gases have completely filled shells and are the most stable elements, due to this if an electron is added to their atoms they lose their stability, as the stability is lost we need to provide energy in order to add an electron to the noble gases.

Q-9. What is the difference between electron gain enthalpy and electron affinity of an atom?

Ans: Electron gain enthalpy of an element is the energy released when an electron is added to an isolated gaseous atom, but electron affinity of an element is the energy released when an electron is added to a neutral isolated gaseous atom.

Q-10. Why the electron gain enthalpies of some of the elements of 2nd period i.e. O and F are less negative than the corresponding elements of the third period.

Ans: The elements of the second period have smallest atomic size among the elements in their respective group. As a result, there are considerable electron-electron repulsions within the atom itself and hence the additional electron is not accepted with the same ease as is the case with the remaining elements in the same group.

Q-11. What are the exceptions of Electron Affinity?

Ans: Exception in Electron Gain Enthalpy are as follows.

In the case of Chlorine and Fluorine, Chlorine has a higher negative electron gain enthalpy value.
In between Sulphur and Oxygen, Sulphur has a higher negative value than oxygen.

Q-12. What is Electronegativity?

Ans: Electronegativity of an element is the tendency of its atoms to attract the shared pair of electrons towards itself in a covalent bond. The electronegativity of any given element depends upon the following factors:

State of hybridization: An sp-hybridized carbon is more electronegative than sp2 hybridized carbon. This is, in turn, more electronegative than an sp3 hybridized carbon.
Oxidation state of the element: The electronegativity of an element increases with the oxidation state of the element.
Nature of the substituents attached to the atom: The carbon atom in CF3I acquires greater positive charge than in CH3I . The electronegativity values for the elements increases along a period from left to right and decreases down a group. As we move along the period from left to right, nuclear charge increases and atomic radius decreases. However, when we move down the group, atomic radius as well as cleaning effect increases.

Q-13. What are Polar and Non-polar Bonds?

Ans: A bond between two similar atoms is non-polar. This is because the shared pair of electrons is equally attracted by the two atoms as the electronegativity of the atoms is the same. For example: H2, Cl2, O2, N2

When the electronegativity of the two atoms forming a bond is different, the shared pair of electrons is attracted more towards the more electronegative element atom.
Higher the difference in electronegativity of the two binding atoms, the more is the dipole moment of the molecule. For example: HF, HCl, HBr, HI So, the more electronegative atom acquires a partial negative charge and the less electronegative atoms acquire a partial positive charge. As a result, two poles are developed and the molecule is said to be Polar.

Q-14. What are the major differences between metals and nonmetals?

Ans : The following table describes the major differences between metals and non-metals.

Metals Non-Metals
1 Metals have a greater tendency to losing electrons. Non-metals tend to show a lower tendency towards losing electrons.
2 Metals show a lower tendency towards gaining electrons. Non-metals show a greater tendency to gaining electrons.
3 Ionic compounds are formed by metals in general. Covalent compounds are formed by nonmetals in general.
4 Metal oxides are basic in nature. Nonmetallic oxides are acidic in nature.
5 Ionization enthalpies of metals are low. Ionization enthalpies of nonmetals are high.
6 Metals have a lower electronegative charge. They are a group of elements that are electropositive. Non-metals are electronegative in nature.

	Metals	Non-Metals
1	Metals have a greater tendency to losing electrons.	Non-metals tend to show a lower tendency towards losing electrons.
2	Metals show a lower tendency towards gaining electrons.	Non-metals show a greater tendency to gaining electrons.
3	Ionic compounds are formed by metals in general.	Covalent compounds are formed by nonmetals in general.
4	Metal oxides are basic in nature.	Nonmetallic oxides are acidic in nature.
5	Ionization enthalpies of metals are low.	Ionization enthalpies of nonmetals are high.
6	Metals have a lower electronegative charge. They are a group of elements that are electropositive.	Non-metals are electronegative in nature.

Q-15. The first ionization enthalpy (ΔiH₁) and the second (ΔiH2) ionization enthalpies and the (ΔegH) electron gain enthalpy in kJ mol- 1 of a few elements are given below.

Elements (ΔiH1) (ΔiH2) (ΔegH)
I 520 7300 −60
II 419 3051 −48
III 1681 3374 −328
IV 1008 1846 −295
V 2372 5251 48
VI 738 1451 −40

Elements	(ΔiH1)	(ΔiH2)	(ΔegH)
I	520	7300	−60
II	419	3051	−48
III	1681	3374	−328
IV	1008	1846	−295
V	2372	5251	48
VI	738	1451	−40

Which of the above elements is likely to be :

(a) The least reactive element ?

(b) The most reactive metal?

(c) The most reactive non-metal?

(d) The least reactive non-metal?
(e)The metal which can form a stable binary halide of the formula MX₂MX₂ (XX=halogen)?
(f)(f) The metal which can form a predominantly stable covalent halide of the formula MXMX (XX=halogen)?

Ans:

(a) The least reactive element ?
The least reactive element is likely to be element V. This is due to the fact that it has the highest initial ionization enthalpy (ΔiH₁)(ΔiH₁) and the highest positive electron gain enthalpy (ΔegH).(ΔegH).

(b) The most reactive metal.
Ans: Because it has the lowest first ionization enthalpy (ΔiH₁)(ΔiH₁) and a low negative electron gain enthalpy (ΔegH).(ΔegH). Element II is predicted to be the most reactive metal.

(c) The most reactive non-metal.

Ans: Element III, with a high first ionization enthalpy (ΔiH₁)(ΔiH₁) and the largest negative electron gain enthalpy (ΔegH).(ΔegH). is anticipated to be the most reactive nonmetal.

(d) The least reactive non-metal.
Ans: Because it has a very high first ionization enthalpy (ΔiH₂)(ΔiH₂) and a positive electron gain enthalpy (ΔegH).(ΔegH). element V is predicted to be the least reactive nonmetal.

(e) The metal which can form a stable binary halide of the formula MX₂MX₂ (XX=halogen)?
Ans: The negative electron gain enthalpy of element VI is low(ΔegH).(ΔegH). As a result, it is a metal. It also has the smallest second ionization enthalpy, (ΔiH₂).(ΔiH₂). As a result, a stable binary halide with the formula MX₂MX₂ (XX=halogen) can be formed.

(f) The metal which can form a predominantly stable covalent halide of the formula MXMX (XX=halogen)?
Ans :The first ionization energy of element I is low, whereas the second ionization energy is large. As a result, a primarily stable covalent halide with the formula MXMX (XX=halogen) can be formed.

Let's conclude here.

Wish you visit the next chapter on Chemical Bonding and Molecular Structure .

Bye.

@Subhas C Chakra

Element	$Δ_{1} H_{1}$	$Δ_{1} H_{2}$	$Δ_{e g} H$
I	520	7300	-60
II	419	3051	-48
II	1681	3374	-328
IV	1008	1846	-295
V	2372	5251	+48

	Z	$(I E)_{1}$	$(I E)_{2}$	$(I E)_{3}$
$L$	3	520.1	$X(1)$	11810
$e$	4	899.3	$Y(2)$	14810

Saturday, July 2, 2022

Class 11- Chemistry: Chapter 3: Classification of Elements and Periodicity in Properties

Classification of Elements and Periodicity in Properties

Periodic Trends in Properties of Elements

Trends in Physical Properties

Atomic Radii:

• Variation of Atomic Radius in the Periodic Table

• Ionic Radius

Effective Nuclear Charge

Oxidation Number

Important Points for Determining Oxidation Number

Ionization Enthalpy

Successive Ionization Enthalpies

• Variation of Ionization Enthalpies in the Periodic Table:

• Electron Gain Enthalpy and Electron Affinity:

Successive Electron Gain Enthalpies:

Problem - 1:

I and V

V and II

II and V

IV and V

Solution:

Answer - C

Problem - 2:

The first and the second ionization enthalpies (in ) electron gain enthalpy (in ) of a few elements are given below:Elements I 520 7300 –60II 419 3051 –48III 1681 3374 –328IV 1008 1846 –295V 2372 5251 +48VI 738 1451 –40Why is IV above the least reactive non-metal?

Solution:

The element IV has a high negative electron gain enthalpy of -295 but its second ionization enthalpy( ) is not so higher than the first ionization enthalpy and hence is the least reactive non-metal.

Problem - 3:

(IE)1​ and (IE)2​ in kJ mol−1 of K and Ca have been given as(IE)1​(IE)2​K4193052Ca5901145The difference can be explained by:

stable inert gas configuration

hydrogen bonding

steric hindrance

van der wall effect

Solution:

(IE)2​ of K >>> (IE)1​ of K

since K+ (after one electron is removed)has stable inert gas configuration as compared to corresponding values for Ca in which Ca⁺ can lose second electron easily.

Problem - 4:

IE (in kJ moI−1) of Li and Be are given below:Z(IE)1​(IE)2​(IE)3​L3520.1X(1)11810e4899.3Y(2)14810From above table X(1) is higher than Y(2) . Justify.

Problem - 5:

Which of the following will have the most negative electron gain enthalpy and which the least negative?P, S, Cl, F

Solution:

Problem - 6:

Which of the following hast he most negative electron gain enthalpy?

Br

Sn

Ca

Mg

Solution:

Problem - 7:

The electron gain enthalpy is least for?

2s22p5

3s23p5

2s22p4

3s23p4

Solution:

Problem - 8:

N have very high electron gain enthalpy.

True

False

Solution:

Problem - 9:

Among halogens, the correct order of amount of energy released in electron gain (electron gain enthalpy) is:

F>Cl>Br>I

F<Cl>Br>I

F>Br>Cl>I

I>Cl>Br>F

Solution:

Problem - 10:

Choose the correct order for the magnitudes of electron gain enthalpy.

F>Br>O>S

F>S>Br>O

F>Br>S>O

S>F>Br>O

Solution:

Problem - 10:

Solution:

• Electronegativity

• Periodic Trends in Chemical Properties along a Period

Some Common Anomalies

Some Common Anomalies

The element IV has a high negative electron gain enthalpy $(\triangle_{eg}H)$ of -295
$KJ mol ^{-1}$ but its second ionization enthalpy( $\Delta _ i H_2$ ) is not so higher than the first ionization enthalpy $(\triangle_{i}H_{1})$ and hence is the least reactive non-metal.

$(I E)_{1}$ and $(I E)_{2}$ in kJ $m o l^{- 1}$ of K and Ca have been given as

$(I E)_{1}$
$(I E)_{2}$
K
419
3052
Ca
590
1145
The difference can be explained by:

$(I E)_{2}$ of K >>> $(I E)_{1}$ of K

since $K^{+}$ (after one electron is removed)has stable inert gas configuration as compared to corresponding values for Ca in which $C a$ ⁺ can lose second electron easily.

IE (in kJ $m o I^{- 1}$ ) of $L i$ and $B e$ are given below:

Z
$(I E)_{1}$
$(I E)_{2}$
$(I E)_{3}$
$L$
3
520.1
$X(1)$ 11810
$e$
4
899.3
$Y(2)$ 14810
From above table $X (1)$ is higher than $Y (2)$ . Justify.

Which of the following will have the most negative electron gain enthalpy and which the least negative?
P, S, Cl, F

$B r$

$S n$

$C a$

$M g$

$2 s^{2} 2 p^{5}$

$3 s^{2} 3 p^{5}$

$2 s^{2} 2 p^{4}$

$3 s^{2} 3 p^{4}$

$F > C l > B r > I$

$F < C l > B r > I$

$F > B r > C l > I$

$I > C l > B r > F$

$F > B r > O > S$

$F > S > B r > O$

$F > B r > S > O$

$S > F > B r > O$

Elements (ΔiH1) (ΔiH2) (ΔegH)
I 520 7300 −60
II 419 3051 −48
III 1681 3374 −328
IV 1008 1846 −295
V 2372 5251 48
VI 738 1451 −40

(d) The least reactive non-metal?
(e)The metal which can form a stable binary halide of the formula MX₂MX₂ (XX=halogen)?
(f)(f) The metal which can form a predominantly stable covalent halide of the formula MXMX (XX=halogen)?

(a) The least reactive element ?
The least reactive element is likely to be element V. This is due to the fact that it has the highest initial ionization enthalpy (ΔiH₁)(ΔiH₁) and the highest positive electron gain enthalpy (ΔegH).(ΔegH).

(b) The most reactive metal.
Ans: Because it has the lowest first ionization enthalpy (ΔiH₁)(ΔiH₁) and a low negative electron gain enthalpy (ΔegH).(ΔegH). Element II is predicted to be the most reactive metal.

(d) The least reactive non-metal.
Ans: Because it has a very high first ionization enthalpy (ΔiH₂)(ΔiH₂) and a positive electron gain enthalpy (ΔegH).(ΔegH). element V is predicted to be the least reactive nonmetal.