Chapter 3 : Classification of Elements and Periodicity in Properties
With time the number of identified elements increased.
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.
To ease out this problem, scientists searched for a systematic way to organise their knowledge by classifying the elements.
Classification of elements into groups and development of Periodic Law and Periodic Table.
Topics covered in this snack-sized chapter:
The elements are put in rows (called periods) by increasing atomic number.
Elements are put into columns (called groups) based on the way they react.
3.3 Modern Periodic Law and the present form of the Periodic Table arrow_upward
Mendeleev’s Periodic Table
Mendeleev arranged elements in horizontal rows and vertical columns of a table in order of their increasing atomic weights in such a way that the elements with similar properties occupied the same vertical column or group.
He proposed that some of the elements were still undiscovered and, therefore, left several gaps in the table.
Modern Periodic Table
Modern Periodic Law can be stated as follows:
“The physical and chemical properties of the elements are periodic functions of their atomic numbers.”
In modern periodic table:
- The elements are put in rows (periods) by increasing atomic number. The horizontal rows are called periods and are labeled from 1 to 7.
- Elements are put into columns (groups) based on the way they react. The vertical columns are called groups and are labeled from 1 to 18.
The naming of the new elements had been
traditionally the privilege of the discoverer (or discoverers) and the suggested name was ratified by the IUPAC.
3.4 Nomenclature of Elements with Atomic Numbers > 100 arrow_upward
Distribution of electrons into orbitals of an atom is called its electronic configuration.
An elementís location in the Periodic Table reflects the quantum numbers of the last orbital filled.
3.5 Electronic Configurations of Elements and the Periodic Table arrow_upward
(a) Electronic Configurations in Periods
The first period (n = 1) starts with the filling of the lowest level (1s) and therefore has two elements hydrogen (1s1
) and helium (1s2
) when the first shell (K) is completed.
The second period (n = 2) starts with lithium and the third electron enters the 2s orbital.
The L shell is completed at neon (2s2
). There are 8 elements in the second period.
The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital.
The third period consists of 8 elements from sodium to argon.
The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital.
It may be noted that 3d orbital is filled up first before 4p orbital.
3d transition series:
The fourth period ends at krypton with the filling up of the 4p orbitals.
- This starts from scandium (Z = 21) which has the electronic configuration 3d1
. The 3d orbitals are filled at zinc (Z=30) with electronic configuration 3d10
The fifth period (n = 5) beginning with rubidium is similar to the fourth period and contains the 4d transition series starting at yttrium (Z = 39). This period ends at xenon with the filling up of the 5p orbitals.
The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, 5d and 6p orbitals, in the order ó filling up of the 4f orbitals begins with cerium (Z = 58) and ends at lutetium (Z = 71) to give the 4f-inner transition series which is called the lanthanoid series.
The seventh period (n = 7) is similar to the sixth period with the successive filling up of the 7s, 5f, 6d and 7p orbitals and includes most of the man-made radioactive elements.
Filling up of the 5f
orbitals after actinium (Z = 89) gives the 5f-inner transition series known as the actinoid series.
(b) GroupWise Electronic Configurations
Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties.
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.
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.
3.6 Electronic Configurations and Types of Elements: s-, p-, d-, f-
The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1
outermost electronic configuration belong to the s-Block Elements.
They are all reactive metals with low ionization enthalpies. They lose the outermost electron(s) readily to form 1+ ion (in the case of alkali metals) or 2+ ion (in the case of alkaline earth metals).
3.6.1 The s-block Elements:
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.
At the end of each period is a noble gas element with a closed valence shell ns2
3.6.2 The p-Block Elements
These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterized by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements.
3.6.3 The d-Block Elements (Transition Elements)
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) are characterised by the outer electronic configuration (n-2)f1-14
Actinoid elements are radioactive. The elements after uranium are called Transuranium Elements.
3.6.4 The f-Block Elements (Inner-Transition Elements)
The elements can be divided into Metals and Non-Metals:
3.6.5 Metals, Non-metals and Metalloids
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 and ductile.
Non-metals are usually solids or gases at room temperature with low melting and boiling points.
They are poor conductors of heat and electricity. Most nonmetallic solids are brittle and are neither malleable nor ductile.
In a horizontal row, the property of elements change from metallic on the left to non-metallic on the right.
The elements become more metallic as we go down a group; the nonmetallic character increases as one goes from left to right across the Periodic Table.
The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) bordering this line and running diagonally across the Periodic Table show properties that are characteristic of both metals and nonmetals. These elements are called Semi-metals or Metalloids.
There are many observable patterns in the physical and chemical properties of elements as we descend in a group or move across a period in the Periodic Table.
Some physical properties of periodic trends are:
3.7.1 Trends in Physical Properties
(a) Atomic Radius:
The atomic size generally decreases across a period. It is because within the period the outer electrons are in the same valence shell and the effective nuclear charge increases as the atomic number increases resulting in the increased attraction of electrons to the nucleus.
(b) Ionic Radius:
The removal of an electron from an atom results in the formation of a cation, whereas gain of an electron leads to an anion.
The ionic radii 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.
A cation is smaller than its parent atom because it has fewer electrons while its nuclear charge remains the same. The size of an anion will be larger than that of the parent atom because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge.
(c) Ionization Enthalpy:
A quantitative measure of the tendency of an element to lose electron is given by its Ionization Enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state.
Across a period, increasing nuclear charge outweighs the shielding. Consequently, the outermost electrons are held more and more tightly and the ionization enthalpy increases across a period.
As we go down a group, the outermost electron being increasingly farther from the nucleus, there is an increased shielding of the nuclear charge by the electrons in the inner levels. In this case, increase in shielding outweighs the increasing nuclear charge and the removal of the outermost electron requires less energy down a group.
(d) Electron Gain Enthalpy:
Electron gain enthalpy provides a measure of the ease with which an atom adds an electron to form anion.
Electron gain enthalpy becomes more negative with increase in the atomic number across a period. The effective nuclear charge increases from left to right across a period and consequently it will be easier to add an electron to a smaller atom since the added electron on an average would be closer to the positively charged nucleus.
As we go down a group it become less negative because the size of the atom increases and the added electron would be farther from the nucleus.
A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity.
It increases across a period from left to right (say from lithium to fluorine) and decrease down a group (say from fluorine to astatine) in the periodic table.
Some chemical properties of periodic trends are:
3.7.2 Periodic Trends in Chemical Properties
(a) Periodicity of Valence or Oxidation States:
The valence of representative elements is usually (though not necessarily) equal to the number of electrons in the outermost orbitals.
Consider the two oxygen containing compounds: OF2
O. The order of electronegativity of the three elements involved in these compounds is
F > O > Na.
Each of the atoms of fluorine, with outer electronic configuration 2s2
, shares one electron with oxygen in the OF2
molecule. Being highest electronegative element, fluorine is given oxidation state -1. Since there are two fluorine atoms in this molecule, oxygen with outer electronic configuration 2s2
shares two electrons with fluorine atoms and thereby exhibits oxidation state +2.
(b) 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.
The anomalous behavior is attributed to their small size, large charge/ radius ratio and high electronegativity of the elements. In addition, the first member of group has only four valence orbitals (2s and 2p) available for bonding, whereas the second member of the groups have nine valence orbitals (3s, 3p, 3d).
As a consequence of this, the maximum covalency of the first member of each group is 4 (e.g., boron can only form [BF4
, whereas the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons e.g., aluminum forms [AlF6
The periodicity is related to electronic configuration, that is, all chemical and physical properties are a manifestation of the electronic configuration of elements.
The atomic and ionic radii generally decrease in a period from left to right, the ionization enthalpies generally increase and electron gain enthalpies become more negative across a period.
3.7.3 Periodic Trends and Chemical Reactivity