MODERN PERIODIC LAW PERIODIC LAW The physical and chemical properties of elements are periodic function of their atomic numbers. The physical and chemical properties of elements are periodic function of their atomic weights.

 Periods: A horizontal row in the periodic table is called period. There are 7 periods and each period starts with different principal quantum number.  Groups: A vertical column in the periodic table is known as group. There are 18 groups in long form of periodic table.

 In multielectron atoms, the electrons in the outermost shell are screened from the nucleus by the inner electrons. As a result, the electron in the outermost shell does not experience the full charge of the nucleus.  The actual nuclear charge felt by electron is called the effective nuclear charge.  The effective nuclear charge can be determined by subtracting a screening constant for the inner electrons from the actual nuclear charge( atomic number). Z* =Z- S where Z* is effective nuclear charge Z is nuclear charge(atomic number) S is screening constant.

 Write the electronic configuration of the element in the following order and grouping: (1s), (2s,2p), (3s,3p), (3d), (4s,4p), (4d), (4f), (5s,5p), (5d), (5f),…………  Electrons in any group higher in the sequence than the electron under consideration contribute zero to screening constant.  Electrons in the same group contribute S= 0.35 each. ( if electron is present in 1s it will contribute S= 0.30).  For an electron in ns or np orbital, all electrons in (n-1) shell contribute S= 0.85 each and all electrons in (n-2) shell contribute S= 1.0 each.  For an electron in nd or nf orbital, all electrons in lower groups contribute S= 1.0 each.

 The term atomic radius means the distance from centre of the nucleus to the outermost shell of electrons.  An atom gets larger as the number of electronic shells increase; therefore the radius of atoms increases as you go down a certain group in the periodic table of elements.  In general, the size of an atom will decrease as you move from left to the right of a certain period.

 Ionization energy is the quantity of energy that an isolated, gaseous atom in the ground electronic state must absorb to discharge an electron, resulting in a cation. H(g) → H+(g)+e−  Ionization energies are dependent upon the atomic radius. Since going from right to left on the periodic table, the atomic radius increases, and the ionization energy increases from left to right in the periods and up the groups.

 Electron affinity is defined as the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion.  On moving down the group, electron affinity decreases.  On moving across a period, electron affinity increases.

 Electronegativity can be understood as a chemical property describing an atom's ability to attract and bind with electrons.  From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one.  From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.

 Ionic character: ionic character depends upon electronegativity difference between bonded atoms. The greater the difference in the electronegativities of the bonded atoms, the higher is the ionic character of the bond. As we move down the group the degree of ionic character in halides increases because of increase in electronegativity difference. HF > HCl > HBr > HI As a result of ionic and covalent character, their properties also change. For example, AlF3 is essentially ionic. AlCl3 has an intermediate character and AlBr3 and AlI3 exist as covalent dimers. Therefore, the melting point of these halides decrease from AlF3 AlI3. AlF3 and AlCl3are conductors in fused sate (ionic character) while AlBr3 and AlI3 are non-conductors (covalent character).

Oxidising and reducing character: Oxidation is a process in which electrons are lost and reduction is a process in which electrons are gained. Thus, an oxidising agent is a substance which gains electrons while reducing agent is a substance which loses electrons. Since the elements on the left of the periodic table have low ionization energy values and therefore, have good tendency to lose electrons. Consequently, they act as strong reducing agents. On the other hand the elements on the right hand side of the periodic table have great tendency to accept electrons and therefore they act as good oxidising agents. Now, since ionization energy values decrease on moving down a group therefore, reducing character of elements in a group increases down a group. Conversely, the oxidising character of elements decreases on moving down a group. As we move along a period from left to right, the ionization energy increases and consequently, reducing character decreases while oxidising character increases. LITHIUM IS THE STRONGEST REDUCING AGENT WHILE FLOURINE IS THE STRONGEST OXIDISING AGENT.

OXIDATION STATES: Like periodic properties, the oxidation states of elements are closely related to the number of electrons in the valence shell of their atoms. Oxidation states of different elements correspond to the number of electrons gained or lost by their atoms to acquire a complete shell of eight electrons, ns2 np6.  Variation of oxidation state in a period: as we move along a period, the positive oxidation state increases and negative oxidation state decreases. For example, in second period, the oxidation state increases from Li to C from +1 to +4 and then decreases from nitrogen to fluorine from -3 to -1.  Variation in a group : In a group, the maximum oxidation state shown by a p-block element is equal to the sum of its s and p electrons, which is the same as its group number.

acidic and basic character of hydroxides: The acidic or basic character of hydroxides and oxides depend upon the position of the element in the periodic table. E O------H There are two possibilities: 1.If the electronegativity of E is low, as in case of metals like Na, K, Mg and Ca etc. the electrons of E-O bonds are drawn more closely to the oxygen atom. This is because oxygen is highly electronegative element. As a result bond between E-O breaks giving OH- ions. Therefore such substance behaves as bases and readily react with acids like HCl to get neutralized. 2.If the electronegativity of E is high as in case of non-metals like F, Cl, N etc., the electrons of E-O bond are shared more equally between non-metal and the oxygen atom. As a result, the oxygen –hydrogen bond becomes weaker and cleaves forming a proton. Therefore, such substances behaves as acids and are neutralized by bases.

Reactivity of alkali and alkaline earth metals: As the value of ionization energy decreases down the group from Li to Cs, therefore, the reactivity of alkali metal increases from Li to Cs. All the elements are highly electropositive giving +1 ions. Because of the very high second ionisation energies of these elements, their oxidation state in compounds never exceeds +1. On the other hand, alkaline earth metals are in general less reactive than alkali metals. This is because of their relatively high ionization energies and high heat of atomization in comparison to alkali metals. The chemistry of this group is mainly dominated by +2 oxidation state.

Some of the general important chemical trends are discussed below: The anomalous properties of elements in first short period (from Li to F) are explained due to their peculiar atomic properties such as small size, low ionisation energy and high electronegativity values. Diagonal relationship between elements of different groups (such as Li – Mg, Be-Al, B-Si) are due to their similar atomic properties. Trends in bond type with the position of the element in the periodic table and with oxidation state for a given element. The variable oxidation states of transition elements. The stability of compounds and their trends. Trends in stability of coordination compounds and the electron donor power of various types of ligands.

 Electron affinity of F is less than that of Cl. Why?  What is shielding effect? Calculate the shielding constant and effective nuclear charge for 3d electrons of zinc(z=30).  Electronegativity of H and F are 2.1 and 4.0 respectively. What is the percentage of ionic character of H-F bond?  What is ionisation energy ? Write various factors on which it depends.  Explain why successive electron affinities have negative values.  Calculate effective nuclear charge for 3p electron in phosphorus atom ( z=15).  Define electronegativity and discuss paulings scale of electronegativity.  Define electron affinity. What are the factors that govern it?