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

 

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

1.  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.
2. The oxidation number of metal in amalgams is zero.
3. The oxidation number of metal in amalgams is zero.
4. The algebraic sum of the oxidation numbers of all the atoms in an uncharged (neutral) compound is zero. 
5. If the coordinate bond is directed from more electronegative to the less electronegative atom then its contribution is zero for both the atoms.
6. In an ion, the algebraic sum is equal to the charge on the ion.
7. In all compounds. except for ionic metallic hydrides, the oxidation number of hydrogen is +1. 
8. The oxidation number of alkali metals is +1 and for alkaline earth metals is + 2.
9. In case of a coordinate bond, it gives +2 value of oxidation number to the less electronegative atom.
10. Metal hydrides like NaH, MgH2, CaH2, LiH, etc have the oxidation number of hydrogen as -1.
11. As fluorine is the most electronegative element, it always has an oxidation number of  -1 in all of its compounds.
12.  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).
13.  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.
14. 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.
15. 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.