Periodic Properties

INTRODUCTION

In 1800, only 31 elements were known. By 1865, the number of identified elements had more than doubled to 63. 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 is also difficult to handle individually.

LAW OF TRIADS

Johann Döbereiner classified elements in group of three elements called triads.

In Döbereiner triad the atomic weight of the middle element is very close to the arithmetic mean of the other two elements

LAW OF OCTAVES

Since Döbereiner Law of triads worked only for few elements, it was dismissed.

Chancourtois arranged elements in order of increasing atomic weights and made a cylindrical table of elements.

John Newland arranged the elements in the increasing order of atomic weight and noted that the properties of the every eighth element are similar to the first one. This relationship is called as “Law of octaves”

MENDELEEV’S PERIODIC TABLE

According to Mendeleev’s periodic law the physical and chemical properties of elements are periodic functions of their atomic weights.

Merits of Mendeleev’s periodic table:

• Mendeleev’s periodic table was very helpful in remembering and studying the properties of large number of elements
• Mendeleev’s periodic table helped in correcting the atomic masses of some of the elements like gold, beryllium and platinum based on their positions in the periodic table
• Mendeleev could predict the properties of some undiscovered elements like scandium, gallium and germanium. By this intuition, he had left gaps for the undiscovered elements while arranging elements in his periodic table

 SUCCESS OF MENDELEEV’S WORK

Limitation of Mendeleev’s table

• Position of hydrogen is not correctly defined in periodic table. It is placed in group I though it resembles both group 1 and 17.
• In certain pairs of elements increasing order of atomic masses was not obeyed. For example argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1)
• Isotopes were not given separate places in the periodic table although Mendeleev's classification is based on the atomic masses.
• Some similar elements are separated and dissimilar elements are grouped together. For example copper and mercury resembled in their properties but had been placed in different groups. On the other hand lithium and copper were placed together although their properties are quite different.
• Mendeleev did not explain the cause of periodicity among the elements.

• Lanthanoids and actinoids were not given a separate position in the table.

Modern Periodic Table

Henry Moseley showed that the atomic number is a more fundamental property of an element than its atomic mass.

Mendeleev’s Periodic Law was, therefore, accordingly modified.

Modern Periodic Law:

The physical and chemical properties of the elements are periodic functions of their atomic numbers

 Nomenclature of elements of Atomic mass > 100

Electronic Configurations for Periods and groups

Electronic configuration of Zirconium (A. No. 40): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d²

Iodine I₅₃  = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁵

We can classify the elements into four blocks viz.,

s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons.

Hydrogen and Helium are exception to this rule

The s-Block Elements

The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns¹ and ns² outermost electronic configuration belong to the s-Block Elements.

The p-Block Elements

The outermost electronic configuration varies from ns²np¹ to ns²np⁶ in each period. The p-Block Elements comprise those belonging to Group 13 to 18 and these together with the s-Block Elements are called the Representative Elements or Main Group Elements.

The d-Block Elements (Transition Elements)

These elements have the general outer electronic configuration (n-1)d^1-10ns^0-2 .  These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. They are all metals.

The f-Block Elements (Inner-Transition Elements)

They have outer electronic configuration (n-2)f^1-14 (n-1)d^0-1ns²

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

METALS, NON METALS AND METALLOIDS

Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. Metals are usually solids at room temperature [mercury is an exception; gallium and caesium also have very low melting points (303K and 302K, respectively)]. Metals usually have high melting and boiling points. They are good conductors of heat and electricity. They are malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires). In contrast, non-metals are located at the top right hand side of the Periodic Table. In fact, in a horizontal row, the property of elements change from metallic on the left to non-metallic on the right. Non-metals are   usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions). They are poor conductors of heat and electricity.

Most non-metallic solids are brittle and are neither malleable nor ductile. The elements become more metallic as we go down a group; the non-metallic 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 zigzag line in periodic table.

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 non-metals. These elements are called Semi-metals or Metalloids

SHIELDING EFFECT

Shielding effect or screening effect: Due to the presence of electrons in the inner shells, the electron in the outer shell will not experience the full positive charge on the nucleus.

So due to the screening effect, the net positive charge experienced by the electron from the nucleus is lowered and is known as effective nuclear charge.

• Effective nuclear charge, Z_eff, experienced by an electron is less than actual nuclear charge , Z
• Electrons in the outermost shell are repelled (shielded) by electrons in the inner shells. This repulsion counteracts the attraction caused by the positive nuclear charge

Coulomb’s law:

 SHIELDING EFFECT

Effective nuclear charge, Z_eff, experienced by an electron is less than the actual nuclear charge, Z. „Electrons in the outermost shell are repelled (shielded) by electrons in the inner shells. This repulsion counteracts the attraction caused by the positive nuclear charge

Z_eff = Z – S (S = screening constant)

Shielding effect

1. Electrons in inner orbitals have greater shielding effect than electrons in same shell.
2. Shielding effect s > p > d > f

ATOMIC RADII

Periodicity

• As we move from left to right along the period, the effective nuclear charge "felt" by the outermost electron increases while the distance from the nucleus doesn't change that much (electrons are filling the same shell)
• Outermost electrons are attracted stronger by the nucleus, and the atomic radius decreases

IONIC RADII

When atom loses an electron, its radius always decreases

• Cations (positive ions) are always smaller than their respective neutral atoms.

When atom gains an electron, its radius always increases

• Anions (negative ions) are always larger than their  respective neutral atoms

 ISOELECTRONIC SPECIES

RADII OF ISOELECTRONIC IONS

IONIZATION ENERGY

• If sufficient energy is provided, the attraction between the outer electron and the nucleus can be overcome and the electron will be removed from the atom

FIRST IONIZATION ENERGY (IE₁)

• The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a 1⁺ ion

Na (g) + 496 KJ/ Mol → Na⁺ (g) + e^-

• Second ionization energy (IE₂)

The minimum amount of energy required to remove the 2nd electron from a gaseous 1+ ion

IE₁:    Ca(g) + 590 kJ/mol ® Ca⁺(g) + e^-

IE₂:  Ca⁺(g) + 1145 kJ/mol  ® Ca²⁺(g) + e^-

The 2nd electron "feels" higher nuclear charge (stronger attractive force) since the electron-electron repulsion has been decreased: For all atoms IE₂ > IE₁ 

Ionization Energy: Trends

• IE₁ increases from left to right along a period since the effective nuclear charge (Z_eff)  "felt" by the outermost electrons increases

• There are some exceptions to this general trend caused by additional stability of filled and half- filled subshells (orbitals with the same )
• IE₁ decreases as we go down a group since the outermost electrons are farther from the nucleus

IONIZATION ENERGY: PERIODICITY

IMPORTANT CONCLUSIONS

• Atoms of noble gases have completely filled outer shell, the smallest radii among the elements in the same period, and the highest ionization energies
• Atoms of metals, especially those to the left in the periodic chart, ionize easily forming cations and attaining the electron configuration of noble gases
• Atoms of non-metals, especially those to the right in the periodic chart, are very unlikely to loose electrons easily - their ionization energies are high

ELECTRON AFFINITY

• For most non-metals, it is much easier to achieve the stable electron configuration of a noble gas by gaining rather than losing electrons
• Therefore, non-metals tend to form anions
•  Electron affinity is a measure of an atom's ability to form negative ions
• The amount of energy released when an electron is added to an isolated gaseous atom to form an ion with a -1 charge.

                                                      Cl(g) + e^-   ® Cl^-(g) + 349 kJ/mol

Sign conventions for electron affinity

• If electron affinity < 0 (energy is released)

• If electron affinity > 0 (energy is absorbed)

Compare cation- and anion-forming processes:

ELECTRON AFFINITY: TRENDS

EA becomes more negative on going from left to right along a period

• There are some exceptions to this general trend caused by

additional stability of filled and half- filled subshells

(orbitals with the same l )

• EA becomes less negative as we go down a group because

the attraction of the outermost electrons to the nucleus weakens

ELECTRON AFFINITY: PERIODICITY

• Important conclusions

• Noble gases have completely filled outer shell and therefore zero electron affinity
• Non-metals, especially halogens, gain electrons easily forming anions and attaining the electron configuration of noble gases
• Metals are usually quite unlikely to gain electrons and form anions

Electron Gain Enthalpies (kJ mol^-1) of group17 elements

 

ELECTRONEGATIVITY

• Measures the tendency of an atom to attract electrons when chemically combined with another element

• If element "likes" electrons - high electronegativity (electronegative element)

• If element "dislikes" electrons - low electronegativity (electropositive element)

Sounds like the electron affinity but different

• Electron affinity measures the degree of attraction of an electron by a single atom forming an anion

• Electronegativity measures the attraction of electrons to the atom in chemical compounds

• The scale for electronegativity was suggested by Linus Pauling
• It is a semi-qualitative scale based on data collected from studying many compounds ​​​​​​​ 

ANOMALOUS PROPERTIES OF SECOND PERIOD ELEMENTS

The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) differs in many respects from the other members of their respective group. Behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties. What are the reasons for the different chemical behaviour of the first member of a group of elements in the s- and p-blocks compared to that of the subsequent members in the same group? The anomalous behaviour is attributed to their small size, large charge/ radius ratio, high Electronegativity and absence of d orbital