p block

If Co(+3) is bonded with O donor atom (ie. H2O & oxalato) then ligand acts as strong field ligand

If Cu(+2) is bonded with water then ligand acts as strong field ligand

If Ni(+4) is bonded with fluoride ion then ligand acts as strong field ligand

If Fe(+2) & Mn(+2) are bonded with NH3 then ligand acts as weak field ligand

3d series elements in +2 oxidation state form outer orbital complexes (weak field) upto NH3 in spectrochemical series

Atomic radius of Al > Ga due to poor screening this is called transition contraction. Radius of In ~ Tl

IE order B > Tl > Ga > Al > In . R : higher IE of Tl & Ga are due to poor shielding of f & d rescpectively . (W Shape graph)

EN : B (2.0) > Tl (1.8) > In(1.7) > Ga(1.6) > Al(1.5)

MP : B > Al > Tl > In > Ga ( B exceptionally high whereas Ga is exceptionally low)

Gallium with unusually Low MP(303K) could exist in liq. state in summer. It had High BP (2676K) makes it useful material for measring High Temperature.

BP : B > Al > Ga > In > Tl

Due to small size of B the sum of its first 3 IE are very high . This prevents it to form +3 ions and forces it to form only covalent compounds

IE : C > Si > Ge > Pb > Sn Reason : Small decrease in ∆f H from Si to Ge & slight increase in ,∆f H from Sn to Pb is the consequence of poor shielding effect & increase in size of atom.

EN : C (2.5) > Pb (1.9) > Si = Ge = Sn(1.8)

Density : Pb > Sn > Ge > C > Si ( Only C family has exception for density)

MP : C > Si > Ge > Pb > Sn. C has exceptionally high value.

There is consider increase in covalent radius from C to Si there after from Si to Pb a small increase is seen . This is due to presence of completely filled d & f orbitals in heavy members. Covalent radius : C < Si < Ge < Sn < Pb . But in this Si ~ Ge & Sn ~ Pb due to poor screening.

Catenation : C > Si > Ge ~ Sn. Pb has no catenation tendency .

Since sum of 1st 4 IE is very high , compunds in +4 are generally covalent.

The central metal atom in these halides undergoes sp3 & tetrahedral while Exceptions : SnF4 & PbF4 which are ionic in nature

PbI4 does not exist because Pb-I bond initially formed during the reaction does not release enough E to impair 6s2 e- & excite one of them to higher orbital to have 4 unpaired e- around Pb atom. (Pb+2 > Pb+4 stability)

As to Bi only a small increase in covalent radius due to completely filled d or f in heavy members.

IE is much greater than C family due to extra stable half filled electronic configuration.

EN :N(3.0) > P (2.1) > As (2.0) > Sb = Bi (1.9)

Ionic radius : As > P > N > Bi > Sb

MP : As > Sb > Bi > P > N . It 1st increase from N to As then decrease from Sb to Bi

BP in general increase from top to bottom in the group.

Except N all the elements show allotropy

In case of N , all oxidation states from +1 to +4 tend to disproportionate in acid solution.

In case of P nearly all intermediate oxidation states disproportionate into +5 & -3 both in alkali & acid. However +3 oxi state in case of As, Sb, Bi become increasingly stable wrt disproportionation.

N - N bond is weaker than single P - P bond due to high interelectronic repulsion of non - bonding e- , owing to small Bond length.

EA : S > Se > Te > Po > O due to compact natrure of O it has less EA than S.

MP : O < S < Se < Po < Te

BP : O < S < Se < Po < Te

MP & BP : H2O > H2Te > H2Se > H2S

Among hexahalides, hexaflourides are the only stable halides. SF6 is exceptionally stable for steric reason.

All elements except O form dichlorides and dibromides.

The dihalides are formed by sp3 & are tetrahedral. The well known monohalides are dimeric in nature. Ex : S2F2 , S2Cl2 ,S2Br2 , Se2Cl2 , Se2Br2 . These dimeric halides undergoe disproportination.

Basicity : NH3 > PH3 > AsH3 > SbH3 >_ BiH3

Due to High EN & small size of N, NH3 exhibits H-bonding in solid as well as liquid state. Due to this, it has hingher MP & BP than that of PH3

MP : NH3 > SbH3 >AsH3 > PH3 (1432)

BP : BiH3 > SbH3 > NH3 > AsH3 > PH3 (54132 order)

All the trihalides of these elements except those of N are stable, only NF3 is stable in case of N.

EA : Cl(-349) > F (-333) > Br (-325) > I (-296) R : EA of F < Cl due to small size of F atom , there are strong interelectronic repulsions in relatively small 2p orbitals of F . Thus incoming e- doesnt experience much attraction.

Bond dissociation enthalpy : Cl2 (242.6 Kj/mol) > Br2 ( 192.8) > F2 (158.8) > I2 (151.1) R : the relatively large e- - e- repulsion among the lp in F2 molecules as they are mich closer to each other than in Cl2. So enthalpy of F2 is smaller compared to Cl2.

The non metallic character, decreases down the group, I shows some distinctive metallic properties ex : it possesses metallic lustre & forms +ve ions like I+ , I+3 etc.

F oxidises H2O to O2 , Cl & Br react with H20 & form hydrohalic & hypohalous acids. The rxn of I with H20 is non - spontaneous. I- can be oxidised by O in acidic medium.

Most of rxn of F are exothermic ( due to small & strong bond formed by it with other elements). F forms only 1 oxoacid while other halogens form many.

X react with H to give HX but affinity of H decreases from F to I . They dissolve in H2O to form hydrohalic acids.

Acidic strength : HF < HCl < HBr < HI. The stability of these halides decrease down the grp due to decrease in bond dissociation enthalpy in the order HF > HCl > HBr > HI.

MP : HI > HF > HBr > HCl

BP : HF > HI > HBr > HCl

In this family higher oxides of halogen are more stable .( asked in mains & Adv)

Same oxidation state then I [ greater polarisability] > Cl[multiple bond formation] > Br [ lacks both ]

EA : positive electron gain enthalpy. Ne > Ar = Kr > Xe > Rn > He

Helium has lowest BP (4.2 K) of any known substance.

Helium has unusual property of diffusing through most commonly used laboratory materials such as rubber, glass or plastics.

Atomic radii : Li > Be > B > C > N > O > F < Ne ( for Ne due to vanderwaal )

Radius : H < He ; Li < Ne ; Na < Ar K > Kr ; Rb > Xe [ due to poor shielding Kr and Xe are smaller than 1A grp ]

IE : Li < Be > B < C < N > O < F < Ne

Ionisation energy ( for 3d series) Sc< Ti > V <Cr< Mn < Fe >Co> Ni <Cu < Zn

EA : Li > Be < B ; Na > Mg < Al Ne < N < Be < Li < C < O < F

D block elements have typical metallic properties except Zn , Cd , Hg , Mn

Transition metals are very hard except Zn, Cd ,Hg.

The atomic radii increases towards the end, this maybe attributed to e- e- repulsions. The pairing of e in d orbitals occur after d5. The repulsive interactions b/w paired e- in d orbitals become very dominant towards the end of period & cause expansion of e- cloud thus resulting in increased atomic size. Ionic radii also follow similar trends.

The radii of 3rd (5d) series are virtually the same as those of corresponding members of the 2nd series. This is due to lanthanoid contraction.

IE : Sc < Ti > V < Cr < Mn < Fe > Co > Ni < Cu < Zn

Mn has high metallic/ ionic radii than its adjacent elements Cr , Fe

Standard electrode potential : M2+/M : Only Cu has +ve ie. + 0.34V ; M3+/M2+ : Mn,Co, Fe has +ve value

Cu +1 compounds are unstable in aq.solutions & undergo disproportionation into +2 & 0

MnO4- : intense purple color due to charge transfer spectrum , diamagnetism along with T dependent paramagnetism, KMnO4 -> K2MnO4 + MnO2 + O2

Lutetium , Cerium , Gadolinium has 5d1 trick : lucegad ; 4f2 & 4f8 are absent and the nxt e- enters 5d1

Oxidation states : Noble gas configuration eg : Ce+4( f0)

A half filled f shell ex: Eu+2 & Tb+4 ( f7)

A completely filled f level eg : Yb+2 ( f14)