L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 5:Electronic Configuration and the Periodic Table
5.1 Electronic Configurations of Atoms
The building up of electronic configuration of atoms are based on three principles, namely Aufbau principle, Hund’s rule and Pauli’s principle. These principles also govern the exact ground state electronic configuration of an element.
Note:
<1>Electrons occupying orbitals of the same energy level singly spin parallel:
Reason: to minimize electrical repulsion between like charges negative.
<2>Electrons occupying the same orbitals have opposite spins.
Reason: magnetic attraction which results from the opposite spins can counterbalance the electric repulsion between them.
5.2 Electronic Configurations in Relation to the Periodic Table
The electronic configuration of an atom can be expressed in two ways:
In ‘electron-in-box’ diagrams,
The building up of electronic configuration follow the three principles described previously. Also note that:
-Exactly half-filled subshells (e.g. np3 , nd5) attain extra stability because electrons are evenly distributed in different orbitals. This minimizes electronic repulsion.
-Full-filled subshells (e.g. np6, nd10 )attain extra stability because extra energy is needed to break the spin paired arrangement in a full filled subshell.
Name: ______Class No.:____Date: ______Marks:______
- Explain the following electronic configuration in the ground state of the atom:
a.Ca( Z = 20 ) 1s2 2s2 2p6 3s23p6 4s2 but not 1s2 2s2 2p6 3s23p6 3d2?
- Cr ( Z = 24 ) 1s2 2s2 2p6 3s23p63d5 4s1 but not 1s2 2s2 2p6 3s23p6 3d4 4s2 ?
c.Cu ( Z = 29)1s2 2s2 2p6 3s23p6 3d10 4s1 but not 1s2 2s2 2p6 3s23p6 3d9 4s2?
( 3 marks )
- Consider the following electronic configurations:
Fe2+ :1s2 2s2 2p6 3s23p6 3d6 , Fe3+ : 1s2 2s2 2p6 3s23p6 3d5
Mn2+ : 1s2 2s2 2p6 3s23p6 3d5 , Mn3+ : 1s2 2s2 2p6 3s23p6 3d4
Explain why
- Fe2+(aq) ion is readily oxidized to Fe3+(aq) ion?
- Mn3+(aq) ion is readily reduced to Mn2+(aq) ion?
( 4 marks )
5.3 The Periodic Table and The Atomic Properties of Elements
In the Periodic Table, the elements are arranged in such a way that
- element are in order of increasing atomic number , but not relative atomic mass.
- a new row (period) is stated when electrons start to enter a new principal energy level, and
- elements whose atoms have a similar outer electronic configuration are placed in vertical column. (group)
I.The structure of the Periodic Table
-Element in the Periodic Table can be sub-divided in several ways : group, period and block.
-Within any block, the final electron to be added to an atom enters a sub-level of the type shown by the block letter ( s, p, d or f). For example, Group I and II elements all have their outer electrons in the s-orbital and therefore placed in the s block.
The s block elements (Group I and II elements)
-Group I (Alkali Metals) and Group II (Alkaline Earth Metals) constitute the s-block.
-All the elements in this block are active metals.
-Group similarities and tends within the groups are generally clear.
The p block elements ( Group III to VIII elements)
-p block consists of the Halogens (Group VII), the Noble Gases (Group VIII) and the elements of group 3 to 6 (inclusive).
-Chemical behaviour in this block varies widely with the reactivity of the metals, metalloids and non-metals, and the comparatively lack of reactivity of noble gases.
Similarities within a group are shown by noble gases and halogens.
-Trends within a group are shown most dramatically by the group IV elements carbon to lead.
The d block elements (Transition Metals)
-The d-block lies between s and p blocks in the Periodic Table, d0block elements (or transition elements) are strictly defined as elements which have incomplete d electron-subshell when combined in compounds.
-Transition elements frequently have coloured compoundsand complex ions in various oxidation states.
-In general, transition metals have higher melting points, and are denser and harder than non-transition metals.
II.Ionization Enthalpies
Review: The first ionization enthalpy is
The magnitude of the ionization enthalpy of an element is determined by the attraction of the nucleus for the electron being removed. This attraction is dependent in turn upon:
The following graph illustrate the variation of the 1stI.E. of the first 19 elements in the Periodic Table:
The following trends can be observed:
- On passing from left to right across both periods, there is a general increase in the first ionization enthalpy.
- There is a drop in the first ionization enthalpy between Group II and Group III
(Be and B; Mg and Al)
- There is a drop in the first ionization enthalpy between Group V and VI (N and O; P and S)
- Across each period, ionization enthalpy reaches a maximum value at Group 0 (noble gases).
- There is a large drop in the first ionization enthalpy in moving from one period to another.
(He to Li; Ne to Na; Ar to K ) - The 1stI.E. decreases down a group.
III.Atomic Radii
The electron cloud of any atom has no definite limit. Therefore, the size of an atom cannot be defined simply and uniquely.
-In the case of atoms forming covalent molecules, covalent radii are used. The covalent radius of an atom is defined as half theinternuclear distance between two covalently bonded atoms in the molecule of the elements.
-In the case of metals, metallic radii are used. The metallic radius of an atom is defined as half the internuclear distance in a metallic crystal.
The following trends can be observed:
- On passing from left to right across a period, atomic radius of the elements decreases.
- The atomic radius increases down a group.
Name: ______Class No.:_____Date: ______Marks:______
Exercises
- The first and second I.E. of potassium and calcium are given below:
Ionization enthalpy / Potassium / Calcium
1stI.E. / kJ mol-1 / 419 / 590
2ndI.E. / kJ mol-1 / 3052 / 1145
Explain the difference in I.E. .
( 4 marks )
- Arrange fluorine, neon and sodium in order of increasing first ionization enthalpy. Explain the order you give.
( 4 marks )