periodic classification

The Concept of Periodicity

Explain the concept of periodicity

Consider the electronic configuration of the first twenty elements of the periodic table shown in the table below.

Table 6.3: Electronic configurations of the first 20 elements

You will notice that elements in the same vertical columns (groups) have the same number of electrons in the outermost shells of their atoms. Because the outer electrons determine the chemical properties of an element, then the elements in each period tend to resemble each other closely in chemical behaviour. For instance, the noble gases, He, Ne and Ar show a chemical inertness which is characterised by the stable outer electron octet or duplet. Due to this reason, the compounds of the noble gases with other elements have not been found.

Attempts to classify elements by arranging them in order of increasing atomic weights shows that the properties of elements were periodic. This means elements with similar or comparable properties appear after a certain specific interval in a given arrangement. The occurrence of successive groups of elements showing strong chemical similarity in this way is called periodicity.

Therefore, periodicity is the repetition of similar chemical properties of elements after a certain specific interval in a given arrangement. The repetition in properties is due to repetition of similar electronic configuration of outermost shells of elements after certain intervals.

General Trends

This refers to change in some properties of elements across the periods and down the groups in the periodic table. These trends become more obvious if we leave aside the noble gases in Group 0. In this case, we shall concentrate our efforts on variations in the most important properties of the elements only. The following is a summary of the change in some properties of elements down the groups and across the periods.

The Change in Properties of Elements Across the Periods

Explain the change in properties of elements across the periods

Atomic and ionic size

The sizes of atoms and ions may be given in terms of atomic radius and ionic radius units respectively. The number of shells an atom or ion posses and the nuclear charge determines the size of an atom or ion. This is how the two properties vary along the period and down the group:

Atomic size

Along the period: Considering the normal elements only, the size of the atoms decrease from left to right across the period. This is because as atomic number increases across the period, the nuclear charge (due to increasing protons) increases and electrons in shells are pulled closer to the nucleus.

Ionic size

1. Positive ions (cations): Across the period; The ionic size does not change, i.e. remains the same, as you move across the period from either direction.
2. Negative ions (anions): A negative ion is larger compared to the corresponding neutral atom because on forming an ion, one or more electrons are added to the atom. The added electron(s) is/are repelled by the electron(s) already present in the outermost shell, hence leading to an increase in the size of an atom, even though no new shell is formed. Down the group and along the period: Ionic size increases down the group, and along the period, i.e. from left to right.

Atomic radii (singular: radius)

Along the period: In the period, atomic radii decrease from left to right with increase in the atomic number.

Electronegativity

Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a molecule. The electronegativity values of elements in group 0 (inert gases) is zero.

Along the period: Electronegativity increases while moving across the period from left to right in the periodic table.

Metallic character (or electropositivity)

Electropositivity is the tendency of an element to lose the valency electron(s) and donate the same to other elements (usually non-metallic elements). This process occurs during the formation of new substances e.g. molecules and compounds. Literally, such reactions occur between metals and non-metals whereby metals donate electrons and non-metals receive these electrons. So, metals are electropositive elements while non- metals are electronegative elements.

Along the period: Generally, metallic character decreases along the period from left to right.The gradation in metallic properties across the period is as follows: Metals → poor metals → metalloids → non-metals → noble gases

Chemical reactivity

Reactivity is the tendency of an element to lose or gain electrons in a chemical reaction.

Along the period: For metals, the reactivity decreases from left to right in a period while it increases for non-metals.

Ionization Energy or Ionization Potential (I.E or I.P)

This refers to the minimum amount of energy required to remove the most loosely bound electron from an isolated atom or ion in its gaseous state. The smaller the value of ionization energy, the easier it is to remove the electron from the atom.M(g) →M+(g) + e-

Along the period: It increases along the period from left to right with the increase in atomic number.

Electron affinity (Ea):

This is just opposite to I.E. It is defined as the amount of energy released when an extra electron is added to an isolated neutral atom in its gaseous state.

Along the period: The value increases along the period from left to right.

Density and melting point

The density of a substance is the ratio of its mass to its volume, while the melting point is the temperature at which a solid substance turns into liquid at standard atmospheric pressure.

1. Density-Across the period: Densities decrease across the period from left to right.
2. Meting point-Across the period: Melting points of elements decrease across the period from left to right.

The Change in Properties of Elements Down the Groups

Explain the change in properties of elements down the group

Atomic and ionic size

1. Atomic size-Down the group: Atomic size increases as you move down the group.
2. Ionic size- Positive ions (cations)-Down the group: On descending the group, the nuclear charge increases and the number of shells increase by one at each step so, the ionic size also increases. A positive ion is smaller than the corresponding neutral atom because on forming the ion, the metal atom loses both the valency electron(s) and the outermost shell. Valency electron(s) refer(s) to the electron(s) in the outer-most shell of an atom. Any further removal of electron(s) from the ion will decrease the ionic size further.Negative ions (anions)-Down the group and along the period: Ionic size increases down the group, and along the period, i.e. from left to right.

Atomic radii (singular: radius)

Atomic radius is the distance from the centre of the nucleus to the outermost shell (valency shell). Down the group: Atomic radii of elements increase down the group with increase in atomic size.

Electronegativity

Down the group: Electronegativity decreases while moving downwards in a group.

Metallic character (or electropositivity)

Down the group: Metallic character (electropositivity) increases down the group

Ionization Energy or Ionization Potential (I.E or I.P)

Down the group: It decreases gradually down the group.

Why is there a decrease in I.E as you go down the group? This is because electrons are held in their shells by their attraction to the positive nucleus, and as you go down the group, the size of the atom increases (increasing atomic radius). So, the outermost electron(s) of an atom gets further and further away from the attraction or pull of the positive nucleus, hence requiring little energy to remove from the atom.

Electron affinity (Ea)

Down the group: The value of electron affinity decreases down the group.

Density and melting point

1. Density-Down the group: Densities of elements increase down the group.
2. Meting point-Down the group: Melting points of elements decrease down the group as the elements become less metallic in nature.

Electronic Configuration to Locate the Positions of Elements in Periodic Table

Use electronic configuration to locate the positions of elements in periodic table

The modern periodic table is based on electronic configurations of the elements. Look at table 6.3 and study the electronic configurations of the first twenty elements and where they are placed in the periodic table.

Beryllium, magnesium and calcium have two electrons in the outer shell. These elements are in Group 2.

This pattern continues to Group 3, Group 4 and so on. The group number in the periodic table is the same as the number of electrons in the outermost shell. The halogens are the elements in Group 7. Bromine is one of the halogens. How many electrons does each bromine atom have in its outer shell?

As we move down each group, the number of shells increases by one at each step. Each atom of an element has one complete shell than the one above it.

As we move across each period, the outer shell is being filled by one electron at each step. Certain electronic configurations are found to be more stable than others are. The noble gases at the end of each period have full outer shells. They have stable duplet (2 electrons) or octet (8 electrons) in their outermost shells. This makes them more difficult to break up, and this fits well with the fact that they are so unreactive.

The outer electrons of an atom are mainly responsible for the chemical properties of an element. Therefore, elements in the same group will have similar chemical properties.