1 Periodic relationships among the elements

2 Periodic table - *The elements are arranged according to the no. of electrons. - *The horizontal raw is called (period) and the vertical column (group). - *It is divided into two blocks : A block : -It contains the representative elements or the main group elements. -The outer most electrons are in s or p orbitals. -It has 8 groups from I A to VIII A. -I A is called alkali metal, -II A is called alkaline earth, -VII A contains halogens and VIII A contain Noble gases. -The elements in the same group behave alike chemically because they have similar electron configuration.

3 B block : -It contains the transition and the inner transition elements whose outer most electrons are in d or f orbitals respectively - *The 1 st period contain H 1 and He 2. - *The 2 nd & 3 rd period contains 8 elements. - *The 4 th & 5 th period contains 18 elements. - *In the 6 th & 7 th period two f series elements branch from group IIIB which are the Lanthanide series or called the rare earth series (4f in the 6 th period) & the Actinide series (5f in the 7 th period). - Components - No. of period - 1s s 2p s 3p s 3d 4p s 4d 5p s 4f 5d 6p s 5f 6d 7p

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5 Valence electrons : - These are the electrons in the outer most principle quantum level of an atom. They are involved in bonding. Core electrons : - These are the inner electrons other than valence electrons. Eg. N 7 : 1s 2 2s 2 2p 3. Elements of the same group have the same valence electron.eg.Li 3, Na 11 Potassium (K 19 ) has the same chemistry of Li and Na because it occupies the same gp K : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 = [Ar] 4s 1. where Ar = Argon 18 Ca 20 = [Ar] 4s 2 Ti 22 = [Ar] 4s 2 3d 2 (transition element). 1s2s2p   Core Valence electrons electrons

6 Electronic configuration of ions : - In - ve ions (anions), we add electrons to the last orbital. - In + ve ions (cations), we remove electrons from the last orbital. eg: H 1 1s 1 + e H - 1s 2 F 9 1s 2 2s 2 2p 5 + e F - 1s 2 2s 2 2p 6 O 8 1s 2 2s 2 2p 4 + 2e O -2 1s 2 2s 2 2p 6 N 7 1s 2 2s 2 2p 3 + 3e N -3 1s 2 2s 2 2p 6 Na 11 1s 2 2s 2 2p 6 3s 1 – e Na + 1s 2 2s 2 2p 6 F -, O -2, N -3, Na + are called isoelectronic ions because they posses the same electron number. i.e same electron configuration as [Ne] Ca 20 [Ar] 4s 2 – 2e Ca +2 [Ar] Al 13 [Ne] 3s 2 3p 1 – 3e Al +3 [Ne] Mn 25 [Ar]4s 2 3d 5 – 2e Mn +2 [Ar]3d 5. N.B. In Mn 25 we remove es from outer shell 4s.

7 Ionization energy (IE): IE = +ve sign “It is the minimum energy required to remove an electron from an isolated atom in its ground state and in the gas state.” - The smaller the ionization energy, the more easily the electron can be removed. X (g) + E 1 X + (g) + e E 1 = 1 st ionization energy X + (g) + E 2 X +2 (g) + e E 2 = 2 nd ionization energy X +2 (g) + E 3 X +3 (g) + e E 3 = 3 rd ionization energy -After an electron is removed from neutral atom, the repulsion between the remaining electrons decreases as the nuclear charge remains constant. -Therefore more energy is needed to remove another electron from the + ve ion therefore 1 st IE < 2 nd IE < 3 rd IE.

8 Ionization energy in the periodic table : - It increases in the period from left to right, Gr Ans. : As the es are in the same energy level and the nuclear charge increase therefore the attraction to outer most shell es increase, so need high IE. i.e metals have low IE and non metals have high IE. - It decreases in the group on going from up to down, Gr. Ans.: As when we go down the group we add new shell so the outer most es become away from the nucleus so the attraction decrease. Therefore He has the highest 1 st IE among all elements. N.B. All properties behave like IE. increase in period from left to right and decrease in group from up to down eg. EN, EA, except radius. I.E decreases I.E increases

9 Electron affinity (E A ): “ It is the amount of energy released when an electron is added to an atom in its gaseous state.” X (g) + e X - (g) + Energy E A = - ve sign EA is always – ve i.e energy is released. The more - ve the value the greater the tendency for an atom to accept es. Electron affinity in the periodic table : - EA increase (become more – ve) in period from left to right. i.e metals have low E A and non metals have high E A. - Halogens have the most E A (most – ve) because accepting electron gives them the stable configuration of Noble gases. - E A decrease in group from up to down i.e Fluoride has the highest EA.

10 Electronegativity (EN):- Def.:- It is the ability of an atom to attract electrons towards it self in a molecule. - Electronegativity increases from left to right across each period.Gr Ans.: - Because in the same period the nuclear charge increase thus the nuclear attraction increase. - Electronegativity decrease in the group as we go from up to down. Gr Ans.: - Because atoms becomes larger and the inner es act as a shield so decrease the attraction between the + ve nucleus and the peripheral es. - Therefore the most electronegative elements are halogens (gp VII), oxygen, nitrogen and sulpher and the least electronegative elements are the alkali metals and alkaline earth metals (gp I & II).

11 Atomic radii : - In the period the radius decrease as we go from left to right. Gr. Ans.; As the nuclear charge increase thus the nuclear attraction increase and atomic size decrease. i.e radius of alkaline earth metal (IIA) is smaller than alkali metal (IA) eg. Li > Be - In the group the radius increase as we go from up to down. Gr Ans. : Because we add new shells. Eg Li < Na < K < Cs. Ionic radii : - Usually among isoelectronic ions anions are larger than cations. Eg. N -3 > O -2 > F - > Na + > Mg +2 > Al +3 All have the configuration of [Ne] : 1s 2 2s 2 2p 6

12 Bonds between atoms -Lewis symbols for elements :- Molecules are formed of atoms in which each atom tends to acquire a stable electron configuration of noble gases by forming bonds. - The number of valence electrons of any element is equal to its group number, which equal to the number of dots in the atom’s Lewis symbol. Group1A2A3A4A5A6A7A Symbol.X.X.X...X...X.. :X:. :X:... :X:...H.H.Be...B...C.. :N:. :O:. :F:.

13 I-Ionic bond :- - It is a type of bond which involve complete transfer of electrons from metals to non metals to form + ve and – ve ions respectively. - Ionic bond is formed between the most reactive metals eg. in gp I & II and the most reactive nonmetals eg. in gp VI & VII. -Oppositely charged ions are arranged in a symmetrical array which are held together by electrostatic attraction in all directions resulting in the formation of hard crystals of high melting point. 11 Na Na + + e 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 as [Ne] stable config. 17 Cl + e Cl - 1s 2 2s 2 2p 6 3s 2 3p 5 + e 1s 2 2s 2 2p 6 3s 2 3p 6 as [Ar] stable config.

14 II- Covalent bond :- -It is a type of bond in which the 2 es of the bond are shared between 2 atoms ( both atoms tends to gain es). It may be:- a- Single covalent bond :- If 2 es are shared eg. in Cl 2 molecule. (es revolve 50% around each atom) b- Double covalent bond :- If 4 es are shared eg. in O 2 molecule. c- Triple covalent bond :- If 6 es are shared eg. in N 2 molecule.

15 - Covalent bond may be between 2 similar atoms as above or 2 different atoms, specially different in electronegativity. It is called polar covalent bond. eg. in HCl Cl is more electronegative than H so it will attract es to itself forming electrical dipole H 2 O molecule :(it must be drawn bent) -To indicate the presence of an electrical dipole in the molecule, the notation δ +, δ - is used. This electric dipole which is responsible for the high boiling point of the substance more than other substance that has no electric dipole. - According to the electronegativity difference between the 2 bonded atoms the bond can be classified as : Ionic if the difference in EN > 1.6, covalent (EN = 0) or polar covalent (EN between 0.5 – 1.6)

16 III- Coordinate bond (dative bond):- - Only one atom gives or shares the 2 es of the bond (donor atom having free es pair) to the other atom (acceptor atom) eg. Ammonium ion NH 4 + Hydronium ion H 3 O +

17 Summary :- Coordinate bond Covalent bond Ionic bond Formed between es rich molecule and es deficient molecule eg. H 3 N: + H + Formed between 2 electronegative elements e.g. Cl 2 Formed between electropositive element (metal) and electronegative element (non metal) e.g. NaCl

18 Bonds between molecules I- Hydrogen bond :- -It is a weak bond between hydrogen and an electronegative element as (O, F and N). The electronegative atom attract strongly the electron pair forming the bond with the hydrogen atom eg. HF HF forms H-bond between 5 molecules only. N.B. 1-The melting and boiling point of the hydrides of gp IV elements increase regularly with increasing the molecular weight.i.e m.p of CH 4 < SiH 4 <GeH 4 < SnH 4 (Ge= germanium, Sn = stannus) 2- The melting and boiling point of the hydrides of N, O and F as in (NH 3, H 2 O, HF) are high in comparison with the hydrides of the corresponding elements in the same gp due to the formation of the hydrogen bond between hydrogen and these elements.

19 II- Molecular or Van der Waal’s bond :- -It is a weak force which arises due to the rotation of the es around the nucleus producing instantaneous dipole, not permanent dipole eg. in noble gases. -Liquification or solidification of noble gases can be explained by Van der Waal’s forces of attraction. -Van der Waal’s forces operate between individual molecules so called “ molecular crystals” E.g In iodine and naphthalene which are characterized by low melting points.