Download chem 1411 – chapter 8

CHEM 1411 – CHAPTER – 7B
ATOMIC STRUCTURE & PERIODICITY
Classification of the Elements
In the modern Periodic table the classification of elements is based on Moseley’s Periodic law, which
states that the physical and chemical properties of elements are a periodic function of their atomic
numbers. Thus atomic number of elements is the basis of classification of elements.
In the periodic table, elements are mainly classified in to 4 groups.
1. Representative Elements (Main group Elements)
These are the elements of the groups IA through 7A. They have incompletely filled valence shells.
2. Noble gases
These are the elements of the 8A group. They have completely filled sub shells in the valence shell.
3. Transition Metals
They are the elements of the groups IB and 3B through 8B. All of them are metals with
incompletely filled (n-1) d sub shells. Elements of 2B have completely filled (n-1) d sub shells. But
are studied along with transition metals because their properties are similar to that of transition
metals.
4. Lanthanides and Actinides
Lanthanides constitute a group of elements with atomic numbers 58 – 71 and Actinides constitute
another group elements from atomic numbers 90 – 103. They have incompletely filled (n-2) f
orbital. All of them are metals and actinides are radioactive.
In the periodic table all the elements of a particular group have similar chemical properties because they
have same number of valence electrons.
Electron configuration of cations and anions (Refer Chapter 7)
Periodic variation in properties of elements
1. Effective nuclear Charge
The shielding effect of the inner electrons and increase in the size of the atom decreases the
electrostatic forces of attraction between the nucleus and the valence electrons. The shielding
effect of electrons decreases in the order s > p > d > f.This will reduce the effective nuclear charge.
Effective nuclear Charge of elements increases across a period and decreases down the group.
2. Atomic radius
It is half of the distance between the two nuclei of two adjacent atoms in a metal or a diatomic
molecule.
Atomic radius decreases across a period because of the increase in effective nuclear charge on the
valence electrons. But atomic radius increases down the group as the effective nuclear charge on
the valence electrons decreases due to the addition of more shells.
3. Ionic radius
Radius of a cation is always smaller than that of the corresponding atom because a cation is
formed by the loss of one or more electrons and this causes an increase in effective nuclear charge
on the valence electrons. On the other hand the radius of an anion is always larger than that of the
corresponding atom because an anion is formed by the gain of one or more electrons and this
causes a decrease in effective nuclear charge on the valence electrons.
4. Ionization energy (KJ / mol)
The minimum amount of energy required to remove an electron from the valence shell of a
neutral gaseous atom in the ground state is called the first ionization energy (I1)
The minimum amount of energy required to remove an electron from the valence shell of a mono
positive gaseous ion in the ground state is called the second ionization energy (I2)
I1 < I2 < I3 <….
In most cases the ionization energy of elements increases from left to right in the period because,
 Increase in effective nuclear charge
 Decrease in atomic radius
Both conditions require more energy to remove the electron from the valence shell
Ionization energy of elements decreases down the group because,
 Decrease in effective nuclear charge
 Increase in atomic radius
4. Electron Affinity
The energy released when an electron is added to the valence shell of a neutral gaseous atom in its
ground state is called the electron affinity.
 Electron affinity of elements increases from left to right in the period because of the
increase in effective nuclear charge on the valence electrons.
 Electron affinity of elements decreases down the group in the periodic table because of the
decrease in effective nuclear charge on the valence electrons.
Electron affinity of Fluorine is lesser than that of Chlorine because of its small size.
Diagonal Relationship
Some elements of the second and third periods which are diagonally positioned, show similarities
in their properties.
Ex. Li and Mg, Be and Al, B and Si
Diagonal relationship is due to the closeness of the charge densities (magnitude of charge / volume)
of their cations.
Variation in Chemical Properties of Representative Elements
Hydrogen (1s1)
Hydrogen is the lightest of all elements.
It forms H+ ion by losing electron. Hydrogen also form H- ion from ionic compounds like NaH and
CaH 2
It combines with Oxygen to form water.
Group IA , Alkali metals (ns1) (Oxidation state +1)
Alkali metals are highly reactive
They lose their shiny appearance when exposed to air as they combine with oxygen
They vigorously react with water to form the corresponding hydroxide and hydrogen
With O2, Li forms monoxide and other metals form monoxide and peroxide
Group 2A, Alkaline Earth Metals (ns2) (Oxidation state +2)
Reactivity of alkaline earth metals towards water and oxygen increases form Be to Ba
Be doesn’t react with water. Mg reacts slowly with steam. Other metals react with hot or boiling
water. The corresponding hydroxide and hydrogen are produced.
They react with oxygen to form oxides and with acids to produce salt and hydrogen
Group 3A elements (ns2 ,np1) (Oxidation state +3)
Boron is a metalloid and the rest are metals
B is not reactive towards O2 and H2O
Al has a protective coating of Al2O3 on its surface. Hence it is a self-protecting metal
Al reacts with O2 to form aluminum oxide and with acids to form salt and Hydrogen
Group 4A Elements (ns2,np2) ((Oxidation state +2 and +4)
Carbon is a non-metal, Si and Ge are metalloids and the rest are metals.
They do not react with water
Metallic elements of this group react with acids to form salts and hydrogen
+2 oxidation state becomes more stable down the group
Group 5A elements (ns2,np3) (Oxidation state –3, +3 and +5)
Nitrogen and Phosphorous are non-metals, As and Sb are metalloids and Bi is a metal
Nitrogen forms a series of oxides; Phosphorous forms trioxide and pent oxide
Group 6A elements (ns2, np4) (Oxidation state –2)
Oxygen, Sulfur and Selenium are non-metals. Te and Po are metalloids. Po is radioactive.
Oxygen is a diatomic gas (O2) whereas S and Se exist in the form of S8 and Se8 molecules.
O2 combines with a large number of elements to form oxides.
Sulfuric acid is formed when SO3 dissolves in water.
Group 7A elements ( ns2, np5) (Oxidation state -1)
They are non-metals and generally called halogens because they are salt producers.They are
highly reactive elements
React with H2 to form hydrogen halides. Aqueous solution of a hydrogen halide is called
hydrohalic acid.
Group 8A Elements (ns2,np6) (Noble gases)
They are mono atomic and the least reactive of all elements. The less reactivity is due to the
presence of fully filled valence sub-shells.
Oxides
The binary compounds of oxygen with other elements are called Oxides.
Types of oxides
1. Basic Oxides
Basic Oxides dissolve in water to form a base and react with acids to form salt and water.
Usually metals form basic oxides.
Ex. Na2O, CaO, MgO
2. Acidic Oxides
Acidic Oxides dissolve in water to form an acid and react with bases to form salt and water.
Nonmetals form basic oxides
Ex. CO2, SO2, P2O5
3. Amphoteric Oxides
Amphoteric Oxides show both acidic and basic properties. They react with both acids
and bases.
Ex. Al2O3, BeO, ZnO, Bi2O3
The basic nature of oxides decreases across a period and the acidic nature increases. Thus alkali
metal and alkaline earth metal oxides are strongly basic and the oxides of group 5A, 6A and 7A are
acidic.