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PERIODIC LAW Periodic Table: This shows the arrangement or grouping of elements in order of increasing atomic number.

Periodic Law: Is the basic assumption behind the modern periodic table; it states that the properties of the elements are the periodic function of their atomic number.  


The electronic configuration  of an atom is the representation of the arrangement of the electrons distributed among the orbital shells and subshells. Commonly, the electronic configuration is used to describe the orbitals of an atom in its ground state


Orbital is the region of space around the nucleus where there is a high probability of finding electron. The four different types of orbitals are s, p, d, and f. These orbitals have different shapes and one orbital can hold a maximum of two electrons. The p-orbital has three degenerate orbitals, with a maximum of six electrons, d has five sub orbitals with a maximum of ten electrons and the f-orbital has seven sub-orbitals with maximum of fourteen electrons.

OrbitalsNumber of sub-orbitalsMax. no. of electronsShape
D510Double dumb-bell


1. Aufbau Principle: states that electrons are filled in their orbital in order of increasing energy level. The order is as follows:

1s < 2s < 2p < 3s < 3p < 4s < 3d <4p etc. 2. Pauli’s Exclusion Principle: states that two electrons in the same orbital of an atom cannot have same values for all quantum numbers. Usually the first electron in an orbital is assumed to enter with up-spin (clockwise ),↑ while the second electron enters the same orbital but with a down-spin (anti-clockwise ), ↓.

3. Hund’s Rule of Maximum Multiplicity: This rule states that electrons occupy each orbital singly first before pairing takes place in a degenerate orbital. Hence, no pairing in degenerate orbital until each orbital is singly occupied with parallel spin. EVALUATION:

  1. Draw a periodic table showing the 1st 20 elements
  2. Define periodic law
  3. Define periodic table.
  4. Write the electronic configuration of the following elements: Cu, Na, Al, and Cl

BLOCKS OF ELEMENTS From the position of the various elements on the periodic table and the electrons arrangements or configuration seem to stand out. Elements that have one and two electron(s) in their last s- orbital i.e. group 1 and 2 elements are called s-block elements. P-block elements have their last electron (s) in the p-orbital are called p-block elements.

Elements that have their last electrons in the d orbital are called d-block elements while f block elements have their last electrons in the f-orbital.

Moving across a particular period on the periodic table, two elements are present in the S block, six elements in the P block, ten elements in the d- block and fourteen elements in the f- block. This corresponds with the maximum number of electrons in the s, p, d, and f orbitals respectively. Also, S block elements are metals, P block elements are mostly non-metals d- block elements  are transition metals while f-block elements are lanthanides (rare earth metals) and actinides (heavy rare earth metals).

There is therefore a diagonal division of the elements into metals and non-metals as shown below. Metals are found on the left side of the thick boundary line and the non-metals on the right with metalloids occurring along the boundary line.  


  1. Paramagnetism: Transition elements exhibit paramagnetism because of the presence of unpair elecrons in the d-orbital. Paramagnetism is the ability of an element to align to the poles of a magnet.
  2. Variable oxidation states: Transition elements exhibit variable oxidation states because they can lose electrons from both 4s and 3d orbitals for bond formation e.g we have Fe2+ and Fe3+
  3. Complex ion formation: Transition elements form complex ions due to the presence of vacant or empty d- orbitals in their ions.
  4. Catalytic ability: The catalytic ability of transition metals is due to the fact that they exist in different oxidation states, hence they are used as catalyst. The partially filled d-orbital allows the exchange of electrons to and from molecules which enable them to act as catalyst.


  1. State the blocks of the following elements: H, F, Na, Ca, Fe and Mn
  2. State three properties of transition elements and explain any two of the properties.

FAMILIES OF ELEMENTS. Elements are arranged into groups or families and periods. Each group has been given a name to allow proper understanding during discussions on the periodic table. Names of the different group (1-8) on the periodic table are shown below.

Group (i) i.e. Li, Na, K etc. ——– alkaline metals

Group (ii) i.e. Be, Mg, Ca, etc. —— Alkaline earth metals

Group (iii) i.e. B, Al, Ga, etc. ——- Boron family

Group (iv) i.e. C, Si, Ge, etc. ——— Carbon family

Group (v) i.e. N, P, As, etc. ———- Nitrogen family

Group (vi) i.e. O, S, Se, etc. ——— Oxygen family or Chalcogens

Group (vii) i.e. Fl, Cl, Br, etc. —— Halogens

Group (viii) or group O i.e. He, Ne, Ar, etc. —— Rare, inert, noble gases.

Transition Elements: These are between groups 2 and 3 on periods 4 and 5.

Lanthanides (Rare-earth elements): these are fifteen elements La−Lu.

Actinides: These are found on the seventh period AC–Lr.

Artificial Elements: These are elements with atomic numbers 93 to 103. They are products formed during chemical reactions. They are unstable and disintegrate in a short time e.g. Plutonium (Pu) and curium (Cm)


1. Four elements P, Q, R and S, have 1, 2, 3 and 7 electrons in their outermost shell respectively. What is the element that is unlikely to be a metal?

2. The elements listed below belong to the same group in the periodic table; 9F, 17Cl, 35Br, 53I

(i) What is the electronic structure of the first-member?

(ii) What is the family name of the elements?

3. Which of the elements has the strongest oxidizing ability?


Some properties of the atom change along a group or across a period on the periodic table. Atomic radius which is measured of the size is one of such properties. The orbiting electrons in an atom are best represented by an electron cloud which has no distinct limit as the size of an action cannot be defined easily.

1. Atomic radius: This has been defined as the distance of closest approach to another identical atom in a given bonding situation. There are two types of atomic radii. Covalent radius and Van der Waals radius. Covalent radius is half the distance between two identical atoms which are not chemically bonded. For the two types of atomic radius two variations are noticeable:

(i) The atom radius increases down a group

(ii) The atomic radius decreases along a period.

This is because going down any group on the periodic table the number of valence electrons remains constant while the shells increase in size (radius) despite increase in nuclear charge. The atomic radius of potassium is greater than that of Sodium. The atomic radius of caesium is greater than that of rubidium.

Across a period, electrons are added to orbitals in the same shell, all the valence electrons are therefore at the same energy level. As atomic number increase the positive charge of the nucleus increases giving rise to greater attraction between the positive nucleus and negative electrons. This is turn result in contraction of the electrons cloud resulting in a smaller atom. Atomic radii therefore decrease across a given period on the periodic table.

2. Ionic Radius: Ions are formed by a loss or gain of electrons by an atom. A positive ion (cation) is smaller than the original metal atom because electrons are pulled in due to increase in effective nuclear charge. A negative ion (anion) is bigger than the corresponding non- metal atom because the effective nuclear charge is reduced. As we move across the second short period, the cationic radii decrease from sodium to aluminium while the anionic radii increase from phosphorous to chlorine.

3. Ionization energy: Ionization occurs when gaseous atom loses electrons from its outer most shell to become positively charged The energy required to do this is called ionization energy or ionization potential. First ionization energy of an element is the energy needed to remove one mole of electron(s) from one mole of atoms in the gaseous state. It is expressed in kilo-joules per mole of atoms ionized.

First ionization energy increase across the period with noble gases having the highest. As we go down he the group, the value of first ionization energy decreases.


First ionization energy KJMOL-1520500420400380


First ionization energy KJMOL-1496737577786101299912551521

Three factors that affect the ionization potential of an atom Ionization potential of an atom is affected by.

(i) Distance of the outer most electrons from the nucleus.

(ii) Size of the positive or effective nuclear change.

(iii) Screening effect of the inner electrons.

Moving from left to right across a period, there is a general rise in the first ionization energy. This is due to the fact that the nuclear charge is increasing across the period. This in turn causes a decrease in atomic radius that is a decrease in the distance of the outermost electrons from the nucleus. The screening effect is almost the same across the period. Down a group of the periodic table, ionization energy decreases because the nuclear charge on the outermost electron is reduced. The outermost electron are properly shielded from the effect of nuclear charge

4. Electron Affinity: is the energy released when an electron is added to gaseous atom in its lowest energy state. Its unit is kJmol-1 or electron volts (ev). Electron affinity increase across a period from left to right and decrease down the group on the periodic table.

Group 1 elements, alkali metals have the least tendency to add electrons to their neutral atoms.

Elements in groups VI and VII have greatest tendency to accept electron. Noble gases (group 8 or 0) have stable electronic configuration

5. Electronegativity: Electronegativity is the ability or power of that atom in a molecule to attract shared pair of electrons. It is more pronounced in heteronuclear molecules where two dissimilar atoms share one or more pairs of electrons.

Electronegativity increases across the period, i.e. going from left to right of the Periodic Table but decreases down the group i.e. going down the Periodic Table. The steady increase as one goes across the period is due to a steady increase in nuclear charge and decrease in atomic size. Consequently, the halogen atom, Fluorine, has the highest electronegativity in the period, due to the strong affinity for electrons. But down the group, the increase in atomic size due to screening effect of the inner shells of electrons decreases the nuclear attraction for shared electrons. The noble gases of group O are not assigned electronegativity values since they have completed shells of electrons.


  1. List three periodic properties of elements that generally increase the across the period of the Periodic Table.
  2. Explain the term electron affinity and discuss how it varies across the period and down the group of the Periodic Table.


Gradation in properties is not confined only to the elements, but it is also found in their compounds with increasing atomic number.

The extent of hydrolysis of the chlorides changes across the third period. Sodium chloride is not hydrolyzed at all in aqueous solution. The same applies to magnesium chloride although hydrated crystals undergo hydrolysis when heated given off HCl and leaving a basic salt. An aqueous solution of aluminium chloride shows appreciable hydrolysis and turns blue litmus red. The chlorides of silicon, phosphorus and Sulphur hydrolyze completely in water.

The general conclusion from the above is therefore as follows: From left hand side to right hand side across any period of representative elements, the metallic character, i.e. tendency to lose electron(s) decreases, and the non-metallic character, i.e. tendency to gain electron(s) increases. Also, as one goes across the period, ionic property decreases while covalent property increases.

DIAGONAL RELATIONSHIP Because metallic character increases down a group and decreases from left to right along a period, there exists a diagonal relationship between the chemical properties of the first member of a group and that of the second member of the next group as in the cases of lithium and magnesium on one hand, and beryllium and aluminium on the other (see the periodic table) EVALUATION         

  1. The atoms of four elements are represented as 20Q, 16R, 10S and 8 Which of the elements would be unreactive?
  2. Explain the meaning of the diagonal relationship the periods 2 and 3 of elements in the Periodic Table.
Lesson tags: Chemistry Lesson Notes, Chemistry Objective Questions, SS2 Chemistry, SS2 Chemistry Evaluation Questions, SS2 Chemistry Evaluation Questions First Term, SS2 Chemistry First Term, SS2 Chemistry Objective Questions, SS2 Chemistry Objective Questions First Term
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