Periodic table -II

Session : Periodic Table II Session

Session Objectives : Session Objectives Causes of periodicity Atomic size,ionic radii,trend in groups and periods Ionisation energy. Electron affinity Electronegativity Valency and its trend Anomalous behaviour of first element of group Diagonal relationship

Causes of periodicity : Causes of periodicity Repetition of similar valence shell configuration after regular interval.

Effective Nuclear Charge : Effective Nuclear Charge In a poly-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. The effective nuclear charge, Zeff, is found this way: Zeff = Z −  where Z is the atomic number and  is a screening constant, usually close to the number of inner electrons.

Sizes of Atoms : Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

Sizes of Atoms : Sizes of Atoms Bonding atomic radius tends to… decrease from left to right across a row due to increasing Zeff. increase from top to bottom of a column due to increasing value of n

Sizes of Ions : Sizes of Ions Ionic size depends upon: Nuclear charge. Number of electrons. Orbitals in which electrons reside.

Sizes of Ions : Sizes of Ions Cations are smaller than their parent atoms. The outermost electron is removed and repulsions are reduced.

Sizes of Ions : Sizes of Ions Anions are larger than their parent atoms. Electrons are added and repulsions are increased.

Sizes of Ions : Sizes of Ions Ions increase in size as you go down a column.Due to increasing value of n.

Sizes of Ions : Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

Illustrative example : Illustrative example Which one of the following is correct order of ionic size?(a) Ca2+ > K1+ > Cl- > S2- (b) S2- > Cl- > K+ > Ca2+ (c) Ca2+ > Cl- > K1+ > S2-(d) S2- > Ca2+ > Cl- > K+ Solution: Hence, answer is (b). Size of iso electronic species decreases with increasein nuclear charge, more interelectronic repulsion inS and Cl is the reason of their increased size.

Illustrative example : Rank each set of Ions in order of increasing size.a) K+, Rb+, Na+ b) Na+, O2-, F - c) Fe+2, Fe+3 Illustrative example Solution: a) since K+, Rb+, and Na+ are from the same group (1A), they increase in size down the group: Na+ < K+ < Rb+ b) the ions Na+, O2-, and F- are isoelectronic. O2- has lower Zeff than F-, so it is larger. Na+ is a cation, and has the highest Zeff, so it is smaller: Na+ < F- < O2- c) Fe+2 has a lower charge than Fe+3, so it is larger: Fe+3 < Fe+2

Illustrative example : Illustrative example Which has the smallest size?(a) Na+ (b) Mg2+(c) Al3+ (d) P5+ Solution: Size of isoelectronic species decreases with increase innuclear charge. Hence, answer is (d).

Ionisation energy : Ionisation energy Isolated gaseous atom Amount of energy required to remove an electron from the ground state of a isolated gaseous atom or ion.

Successive ionization energies : Successive ionization energies IE3 > IE2 > IE1 First ionization energy is that energy required to remove first electron. Second ionization energy is that energy required to remove second electron, etc. It requires more energy to remove each successive electron.

Successive ionization energies : Successive ionization energies When all valence electrons have been removed, the ionization energy takes a quantum leap.

Factors affecting values of ionisation energy : Factors affecting values of ionisation energy

Factors affecting values of ionisation energy : Factors affecting values of ionisation energy 3. Screening effect or shielding effect Combined effect of attractive and repulsive forces between electron and proton.

Factors affecting values of ionisation energy : 6. Stable Configuration Factors affecting values of ionisation energy

Trends in First Ionization Energies : Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron.For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

Trends in First Ionization Energies : Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron. As you go from left to right, Zeff increases. However, there are two apparent discontinuities in this trend.

Trends in First Ionization Energies : Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. Electron removed from p-orbital rather than s-orbital. Electron farther from nucleus. Small amount of repulsion by s electrons.

Trends in First Ionization Energies : Trends in First Ionization Energies The second occurs between Groups VA and VIA. Electron removed comes from doubly occupied orbital. Repulsion from other electron in orbital helps in its removal.

Illustrative example : Illustrative example In which of the following pairs thereis an exception in the periodic trendfor the ionization energy?(a) Fe – Ni (b) C – N(c) Be – B (d) O – F Solution: Hence, answer is (c). Since Be has stable configuration (2s2) as compared to B(2s2, 2p1).

Illustrative example : Illustrative example If the first ionization energy of helium is 2.37 kJ/mole, the first ionization energy of neon in kJ/mole is:(a) 0.11 (b) 2.37 (c) 2.68 (d) 2.08 Solution: Hence, answer is (d). Ionization energy decreases down the group.

Illustrative example : Illustrative example First ionisation energy of Be is more than Li but the second ionisation energy of Be is less than Li. Why? Solution:

Illustrative example : Using the Periodic table only, rank the following elements in each of the following sets in order of increasing IE. a) Ar, Ne, Rn b) At, Bi, Po c) Be, Na, Mg d) Cl, K, Ar a) Rn, Ar,Ne These elements are all noble gases and their IE decreases as you go down the group. b) Bi, Po, At These elements are all Period 6 elements and the IE increases from the left to the right. c) Na, Mg, Be These elements are close to each other, Be & Mg are in the same group, Be is higher than Mg & Na is next to Mg & lower in IE. d) K, Cl, Ar These elements bracket the noble gas Ar, and Cl would be lower than Ar and K would be lower yet! Solution Illustrative example