CHEMISTRY FORMULAS BOOKLET {GIST OF COLLEGE LEVEL CHEMISTRY}, Study notes of Chemistry

Download CHEMISTRY FORMULAS BOOKLET {GIST OF COLLEGE LEVEL CHEMISTRY} and more Study notes Chemistry in PDF only on Docsity! 1 & Short Formula (Chemistry) SHORT FORMULA CHEMISTRY eee lcs oo ATOMIC STRUCTURE onneeee taint ee 4KZe? Estimation of closest distance of approach (derivation) of a-particle : R= m.v2 1 Mo. The radius of a nucleus : R=R, (A)? cm he Planck's Quantum Theory : Energy of one photon=hv = d a 4 : Photoelectric Effect : hy = hy, + 3 ms Bohr’s Model for Hydrogen like atoms : h ae 1. mvr=n Or (Quantization of angular momentum) E. z? 27 - 2x” me* 2 Es-—t 2=2178% 108 2 watom = 1362 ev; Es n n n n 2 hn? 2 2 6 3. etx = = 9.529 x01 A 4, y= 2EZB = 2-18 x10" xz mis Zz 4n‘e*m z nh n De-Broglie wavelength : -tLh = mo p (for photon) Wavelength of emitted photon : 1 a1 1 ~ = y= R24) a n? ng No. of photons emitted by a sample of H atom: An(An+1) 2 2 { Short Formula (Chemistry) Heisenberg’s uncertainty principle : h h h Ax.Ap > in or m Ax.Av > an or AX.Av > tam Quantum Numbers : * Principal quantum number (n} = 1, 2, 3, 4.... too. . . _ nh Orbital angular momentum of electron in any orbit = On * Azimuthal quantum number (/) = 0. 1, ..... to (n-1). * Number of orbitals in a subshell = 2¢ + 1 Maximum number of electrons in particular subshell = 2 x (24 + 1) h h * Orbital angular momentum L = an E+) HAY S(E+1) |= | STOICHIOMETRY M f t if Li it a Relative atomic mass (R.A.M) = Mass oF one atom ofan eremen = Total Number of nucleons p* mass of one carbonatom = ¥-ma Number & Nolume at STP: > va Mole * | + mol. wt.) |x mol. wt. + At. wt. x At. wt. Density: | ; density of the substance Specific gravity = density of water at 4°C For gases: Molar mass of the gas PM Absolute density (mass/volume) = Molar volume ofthe gas RT eas PMgaset = Mgas M, Vapour density V.D.= dy, . PMi, at - Mu, 2 Myes = 2 V.D. § { Short Formula (Chemistry) Volume strength of H,0,: 20V HO, means one litre of this sample of HO, on decomposition gives 20 It. of O, gas at S.T.P. Valume,strength of H,O> Normality of HO, (N) = 56 Volume strength of HQ, Molarity of H,O,(M) = 1.2 Measurement of Hardness : mass of CaCO, Hardness in PPM = Tota! mass of water Calculation of available chlorine from a sample of bleaching powder : 3.55xxxV(mL) % of Cl, = Wwe where x = molarity of hypo solution and v = mL. of hypo solution used in titration. Temperature Scale : c-90 K-273 F-32 _ R-R(O) 700-0 373-273 212-32 — R(100) -R(O) where R = Temp. on unknown scale. Boyle's law and measurement of pressure : 1 At constant temperature, Va P PY, = PY, Charles law : Vi _ Ve At constant pressure, VaT = or Tr = Tr 1 2 Gay-lussac’s law: At constant volume, Pat Pi _ Po T, 7 — temp on absolute scale Ideal gas Equation : PV=nRT pv= {Rtore= 2 RTorPm=aRT m m Daltons law of partial pressure : p, = MRT p, = RT v v Total pressure = P,+ P,+P,+.... T and so on.  6 Short Formula (Chemistry) Partial pressure = mole fraction X Total pressure. Amagat’s law of partial volume : VEV, FV, FV, + Average molecular mass of gaseous mixture : Total mass of mixture nyM, +z Mz +ng My Total no. of moles in mixture ~ Ny +Nz +N Graham's Law : 1 Rate of diffusion rc< rT ; d=density of gas tr fd, ~My, YV.Dy Kinetic Theory of Gases : 1 — PV= 3 MN WU? Kinetic equation of gases a )_3 3 Average K.E. for one mole = N, 3m U? | = 2 KN,T = 3 RT w is } Root mean suqare speed 3RT Ule= Ta molar mass must be in kg/mole. Average speed U,=U,+U,+U, +... RT KT - Uo =yoy zy K is Boltzmman constant xo V aM mm Most probable speed _ f2RT _ [2kT Yes 1M Vm Vander wall’s equation : a 8a ap?’ 'c™ 27Rb 7 { Short Formula (Chemistry) Vander wall equation in virial form : Ze + St oF VB Freres Reduced Equation of state: (, 3 Poy Jov-=er, a0 THEREODY NAMIC Thermodynamic processes : 1. Isothermal process: T=constant dT=0 AT=0 2. Isochoric process : V = constant dV=0 AV=0 3. Isobaric process : P = constant dP=0 AP =0 4 Adiabatic process : q=0 or heat exchange with the surrounding = O(zero) IUPAC Sign convention about Heat and Work : Work done on the system = Positive Work done by the system = Negative 1* Law of Thermodynamics AU=WU,-U)=qtw Law of equipartion of energy : ft U 3 nRT {only for ideal gas) f AE = = mR (AT) 3 PRAT) where f = degrees of freedom for that gas. (Translational + Rotational) f=3 — formonoatomic =5 fordiatomic or linear polyatmic =6 fornon-linear polyatmic. Calculation of heat (q) : Total heat capacity : Aq dq = tte yp Or at ar Pe  1 Short Formula (Chemistry) Thermochemistry : Change in standard enthalpy AH? = He 2 - Ha = heat added at constant pressure. = C,AT. t Hroduets > Heactants => Reaction should be endothermic as we have to give extra heat to reactants to get these converted into products and if Freaucts < Hheadarts > Reaction will be exothermic as extra heat content of reactants will be released during the reaction. Enthalpy change of a reaction AH veacton = Ay causes ~ Hoeactanis AM? actions = HY roduess 7 HY pecans = positive - endothermic = negative - exothermic Temperature Dependence Of AH: (Kirchoff's equation) : For a constant volume reaction AH,° = AH,° + 4C, (T,-T,) where AC, = C, (products) — C, (reactants) For a constant volume reaction AES = AE? + facy aT Enthalpy of Reaction from Enthalpies of Formation : The enthalpy of reaction can be calculated by AH? = Iv, AH,° —Zv, AH). v, is the stoichiometric coefficient. > products >reactarts Estimation of Enthalpy of a reaction from bond Enthalpies : /Enthalpy required to \ ‘Enthalpy released to AH= break reactants into | form products from the | gasesous atoms (gasesous atoms ) Resonance Energy : AH? = AH? - AH rescnance ST +, experimesta callulated = AH® — AH? 6, calolLiatea ©, experimental CHEMICAL EQUILIBRIUM At equilibrium : (i) Rate of forward reaction = rate of backward reaction (i) Concentration (mole/litre) of reactant and product becomes constant. (ill) AG = 0.  Short Formula (Chemistry) WQ=k, Equilibrium constant (K) : rate constant of forward reaction K, = ‘ate constant of backward reaction Kp" Equilibrium constant in terrms of concentration (K.): Ke, fre Kp (aris)? Equilibrium constant in terms of partial pressure (K, ): _ Pel’ Pol" ° Pal’ Pol Equilibrium constant in terms of mole fraction (K,): XeXp K= + KAxB Relation between K, & K, : K, = KART)". Relation between K, & K, : K, = K,(P)" Ky AH 4 * log Kr = 5403R [i 7 1 | AH = Enthalpy of reaction Relation between equilibrium constant & standard free energy change : AG® = — 2.303 RT log K Reaction Quotient (Q): (cr Ipy* The values of expression Q = ° Ar Br Degree of Dissociation (a): a =no. of moles dissociated / initial no. of mofes taken = fraction of moles dissociated out of 1 mole. Note: % dissociation = @ x 100 Observed molecular weight and Observed Vapour Density of the mixture : molecular weight of equilibriummixture Observed molecular weight of A,(g) = totalno.of moles D-d _ M,-M, as = Mr=M, (n—1)xd (n= My  1 Short Formula (Chemistry) External factor affecting equilibrium : Le Chatelier's Principle: If asystem at equilibrium is subjected to a disturbance or stress that changes any of the factors that determine the state of equilibrium, the system will react in such a way as to minimize the effect of the disturbance. Effect of concentration : If the concentration of reactant is increased at equilibrium then reaction shift in the forward direction . If the concentration of product is increased then equilibrium shifts in the backward direction Effect of volume : If volume is increased pressure decreases hence reaction will shift in the direction in which pressure increases that is in the direction in which number of moles of gases increases and vice versa. If volume is increased then, for An > 0 reaction will shift in the forward direction An <0 reaction will shift in the backward direction An = 0 reaction will not shift. Effect of pressure: * If pressure is increased at equilibrium then reaction will try to decrease the pressure, hence it will shift in the direction in which less no. of moles of gases are formed. Effect of inert gas addition: (i) Constant pressure : If inert gas is added then to maintain the pressure constant, volume is increased. Hence equilibrium will shift in the direction in which larger no. of moles of gas is formed An > 0 reaction will shift in the forward direction An <0 reaction will shift in the backward direction An = 0 reaction will not shift. di) Constant volume : Inert gas addition has no effect at constant volume. Effect of Temperature : Equilibrium constant is only dependent upon the temperature. A AS' , and intercept = 1 AH If plot of énk vs T is plotted then it is a straight line with slope = — * For endothermic (AH > 0) reaction value of the equilibrium constant increases with the rise in temperature * For exothermic (AH < 0) reaction, value of the equilibrium constant decreases with increase in temperature * For AH > 0, reaction shiffts in the forward direction with increase in temperatutre * For AH < 0, reaction shifts in the backward direction with increases in temperature. * If the concentration of reactant is increased at equilibrium then reaction shift in the forward direction . * If the concentration of product is increased then equilibrium shifts in the backward direction Vapour Pressure of Liquid : Partial pressure of HO vapours Relative Humidity = ——*——* "2 Vapour pressure of H,Qat that temp. ‘"ts } 15 | Short Formula (Chemistry) RELATIVE STRENGTH OF TWO ACIDS : [H*] furnished by] acid cy {Ka [H*] furnished by Il acid = C202 ka oz (h) PH of a mixture of two weak acid(both monoprotic) solutions : Kar _ 4 (a, << 1) and (a, <<1) = Kap Op 4 = = ©KonCK [H*]= C0, + 0,0.,.= yCiKai + Cok a2 ** If water is again considered third weak acid in solution of two weak acid then TH] = JK a1Cy + KaoCo +Ky, Cy Kay, = 1074 = Ky, [H*] = yC,Kg, +02 Kay +1074 (i) PH of a mixture of weak acid(monoprotic) and a strong acid solution : if [SA] = C, and [WA] =C,, then [H*] from SA = C, and [H"] from WA = C, Let HA is a weak acid. ae Cy+ C7 +4Ky.Ce 2 ** If a strong acid of low conc is added in water then [H*] of solution can be calculated as C+ Ci +4K, (H1= 4 o SALT HYDROLYSIS : Salt of Type of hydrolysis k, h pH / ‘oni k k | | (a) weak acid & strong base anionic ea Hw 7+ = pk,+ = loge ka k,c 25° 2 (b) strong acid & weak base cationic Kw Kw 7- I k, - loge 9 ky kyo 2 PR 9 109 . kw kw 1 1 (c) weak acid & weak base both kak, Kaky 7+ 2 Pk, - 2 Pk, (d) Strong acid &strongbase = ~--=----1 do not hydrolysed------- pH=7 Hydrolysis of ployvalent anions or cations For [Na,PO,]=C. K,, * K,,= K, "te } 16 | Short Formula (Chemistry) i K,, * K,.= K, K., * K,, =K, Generally pH is calculated only using the first step Hydrolysis Kw Kas SIOH]= ch = JKyxe SIHI= JS 1 So pH = glPKw + PKas +logC] Hydrolysis of Amphiprotic Anion. (Cation is not Hydrolysed e.g. NaHCO,, NaHS, etc.) j pKa, —pK, pH (Hco3) = Maes me) &) Similarly for HPO, and HPO, amphiprotic anions. _{ PKay +PKay _[PKap | PKag PAu P03) -( 3 and PHupot) -| 2 ° H,PO, Ke, HPO,” —Ker_, HPO —Kss_, PO. ionisation. x K, re Kage Kye Kye 1 The pH of H,PO,= 5 (PK,, log) 1K, >> K,,2> Ky 1 PH of NaH,PO, = 3 (pK,, + pK.) 1 PH of Na, HPO, = 5 (PK, + PK,.) 1 PH of Na,PO, = 2 (pKw + pKa, +logC) .. Sec hydrolysis can neglect. BUFFER SOLUTION : (a) Acidic Buffer : e.g. CH, COOH and CH,COONa. (weak acid and salt of its conjugate base). [Salt] , PH= pK, + log [Acid] [Henderson's equation] (b) Basic Buffer: e.g. NH,OH + NH,Cl. (weak base and salt of its conjugate acid). [Salt] [Base] pOH = pK, + log Buffer capacity (index) : Total no. of molesof acid ‘alkali added perlitre Buffer capacity = Change in pH , dx (a+Ob-x) Buffer capacity = GApH = 2.303 asb “Tri 17 | Short Formula (Chemistry) INDICATOR : Hin = Ht + In” [Hin] in“ On] or [H']=K,, [in] [lonised form] PH= PK, +109 para] = PH= pk, + !09 TUnionised form] SIGNIFICANCE OF INDICATORS : © Extent of reaction of different bases with acid (HCI) using two indicators : Phenolphthalein Methyl Orange NaOH 100% reaction is indicated 100% reaction is indicated NaOH + HCI > NaCl + H,0 NaOH + HCl — NaCl + H,O Na,co,, 50% reaction upto NaHCO, 100% reaction is indicated stage is indicated Na,Co, + HCl > NaHCO, + NaCl Na,CO, + 2HCI > 2NaCl + HO + CO, NaHCco, No reaction is indicated 100% reaction is indicated NaHCO, +HCl > NaCl+H,O+ CO, o ISOELECTRIC POINT: IH'T= yKaiK,2 = PKar + PKao pH 2 SOLUBILITY PRODUCT : K,, = (XSF (ys =x ys CONDITION FOR PRECIPITATION : If ionic product K,p > Kgp precipitation occurs, ifK,, = Kg, saturated solution (precipitation just begins or is just prevented). cee bre ERECHRIST RSS ELECTRODE POTENTIAL For any electrode — oxidiation potential = — Reduction potential cel = R-P of cathode —R.P of anode ce! = RP. of cathode + O.P of anode E,,, is always a +ve quantity & Anade will be electrode of low R.P = SRP of cathode — SRP of anode. an Go Greater the SRP value greater will be oxidising power. “oot 20 ' Short Formula (Chemistry) FARADAY’S LAW OF ELECTROLYSIS : First Law : w= 7zq w=Zit 2Z= Electrochemical equivalent of substance Second Law: Wak W constant Wee o = = constan E, E> Ww _ ixtxcurrent efficiency factor E 96500 actual mass deposited/produced CURRENT EFFICIENCY = Theoritical mass deposited/produced Oo CONDITION FOR SIMULTANEOUS DEPOSITION OF Cu & Fe AT CATHODE 0.0591 4 0.0591 1 os = Pout teu S109 yar = Ere? ro 5 100 pga Condition for the simultaneous deposition of Cu & Fe on cathode. CONDUCTANCE : 1 a Conductance = =~ Resistance eo Specific conductance or conductivity : . Se 1 (Reciprocal of specific resistance) K= > K = specific conductance ca Equivalent conductance : , _ Kx1000 ; . Age Normality unit : -ohm~ cm? eq"* = Molar conductance : 3, — Kx1000 ; a . m = Molarity unit : -ohm~ cr’ mole: é Specific conductance = conductance x a KOHLRAUSCH’S LAW: Variation of 4,, /4,, of a solution with concentration : (i) Strong electrolyte dao = Ma — Be (ii) Weak electrolytes : ALS ae tae where 4 is the molar conductivity a 0 of cations obtained after dissociation per formula unit n_=No of anions obtained after dissociation per formula unit "O14 21 4 Short Formula (Chemistry) APPLICATION OF KOHLRAUSCH LAW: 1. Calculation of 2°, of weak electrolytes : 24 cHgcoonn . 2° cttgs00He) + etd et 2. To calculate degree of diossociation of a week electrolyte kK = oa? ea (1-o) 3. Solubility (S) of sparingly soluble salt & their K,, 1000 Dye = Ay = OK Solubility K,, = S% a IONIC MOBILITY : It is the distance travelled by the ion per second under the potential gradient of 1 volts per cm. It’s unitis cm?s ‘v * Absolute ionic mobility : AO oe We : doe Mh, oO ag = Fue : ag FF ™ a v —> speed (V/4)—potentialgradient He _ | He ‘. [ec]. v= (pc: Where t, = Transport Number of cation & t, = Transport Number of anion lonic Mobility w= Transport Number: SOLUTION & COLLIGATIVE PROPERTIES 1.. OSMOTIC PRESSURE : i) m=pgh Where, p = density of soln., h = equilibrium height di) Vont — Hoff Formula (For calculation of O.P.) n=CST n=CRT= 7 RT (ust like ideal gas equation) ., © = total conc. of all types of particles FOF Cyt Cg Fcc = (Mag =My + Vv Note: If V, mL of C, conc. + V, mL of C, conc. are mixed CV, +05Vy m=|"yaye Vi + Vp _ (Mi + 2V2 ne RT Type of solutions : (a) Isotonic solution — Two solutions having same O.P. 1, = K (at same temp.)  22 { Short Formula (Chemistry) (b) Hyper tonic— If x, > 7, — | solution is hypertonic solution w.r.t. 2"¢ solution. (c) Hypotonic — |I"* solution is hypotonic w.r.t. [9 solution Abnormal Colligative Properties : (In case of association or dissociation) VANT HOFF CORRECTION FACTOR ( ic exp/observed/actual/abnormal value of colligative property - Theoritical value of colligative property exp./abserved no. of particles / conc. observed molality Theoritical no. of particles ~ Theoritical molality theoretical molar mass (formula mass) exp erimental /observed molar mass (apparent molar mass) °o i->1 > dissociation. i<1 = association. exp Qo i= Theor 2m =iCRT RE (iC, + igCy + igCy.....) RT Relation between i & a (degree of dissociation) : 1=1+(n-1)a0 Where, n=x+y. Relation b/w degree of association B & i. i=1+ (\-")p 2. RELATIVE LOWERING OF VAPOUR PRESSURE (RLVP) : Vapour pressure : Poy <P Lowering in VP = P—P, = AP . . AP Relative lowering in vapour pressure RLVP = Pp Raoult's law : — (For non - volatile solutes) Experimentally relative lowering in V.P = mole fraction of the non volatile solute in solutions P-P. RIP = => sy P-Ps on Po ON M = (molality } x 7000 (M= molar mass of solvent) "ost 25 \ Short Formula (Chemistry) (b) Negative deviation @ P_ exp < x,p? + x,p%, . A A i) BBP AnB. strong force of altraction. ail) AH i= VE @™) AV ay = VE (1L+1L<2L) ® AS, = +ve wi) AG, =—ve eg. H,O + HCOOH H,0 + CH,COOH H,0+HNO, CH al CHCI,+CH,OCH, > C=O- -H-c~—Cl CHS ‘cl P°A > PIB XA=0 Immiscible Liquids : @ Poa =P, + Pe (i) P,=P,2X, =P,° — [Since, X,= 41] (ili) P,=P2X,=P,° — [Sinee, X, = 1]. . PR _ Ma | Pa _ WaMp (W) Prag = PAP WO) Be ) Bo Ma W _MaRT _ OgRT Pea AT: Pa Ta T, B.P. of solution is less than the individual B.P.'s of both the liquids. "oe t 26 \ Short Formula (Chemistry) Henry Law : This law deals with dissolution of gas in liquid i.e. mass of any gas dissolved in any solvent per unit volume is (l) proportional to pressure of gas in equilibrium with liquid. map m=kp weight of gas ™ —~ ‘Volume of liquid SOLID STATE Classification of Crystal into Seven System Crystal System Unit Cell Dimensions and Angles Cubic a=b=c;a=B=y= 90° Orthorhombic azbec;a=B=y= 90° Tetragonal a=bec:;a=B=7= 90° Monoclinic azbec;a=y=90° #8 Rhombohedral a=b=c;a=pPp=y7290° Triclinic azbecia#Ph#y7#90° Hexagonal a=bec;a= f= 90%y= 120° ANALYSIS OF CUBICAL SYSTEM Property sc Bcc a (i) atomic radius (r) = {38 2 4 di) No. of atoms per unit cell (Z) 1 2 i) ~~ C.No. 6 8 @w Packing efficiency 52% 68% wy) No. voids (a) octahedral (Z) (b) Tetrahderal (22) NEIGHBOUR HOOD OF APARTICLE : Simple Cubic (SC) Structure : Type of neighbour Distance nearest a (next)" ay2 (next)? ay3 Bravais Example Lattices SC, BCC, FCC Nacl SC, BCC, end centred & FCC S, sc, BCC Sn,Zno, SC, end centred s, sc Quartz sc H,BO, sc Graphite Fcc a 2/2 a= edge length 4 12 74% 4 8 no.of neighbours 6 (shared by 4 cubes) 12 (shared by 2 cubes) 8 (unshared) "ort 27 (ll) (uu) t \ Short Formula (Chemistry) Body Centered Cubic (BCC) Structure : Type of neighbour nearest (next) (next)? (next}* (next)* Distance 2r= an 2 =a = av2 = alt 2 ay3 Wi: Face Centered Cubic (FCC) Structure : Type of neighbour nearest (next)' next)’ (next)® (next)' Distance z{M DENSITY OF LATTICE MATTER (d} = yy, [3] no.of neighbours 12 24 no. of neighbours. _ (3x8 12= > (3x8> ol 4 \ 24 12 24 where N, = Avogadro’s No. M = atomic mass or molecular mass. IONIC CRYSTALS C.No. Oa ew e T Limiting radius ratio | | X 0.155 — 0.225 (Triangular) 0.225 -0.414 (Tetrahedral) 0.414 -0.732 (Octahedral) 0.732 — 0.999 (Cubic). EXAMPLES OF AIONIC CRYSTAL (a) Rock Salt (NaCl) Coordination number (6 : 6) Edge length of unit cell :- (c) Zinc Blende (ZnS) C.No. (4: 4) (b) CsCl C.No. (8 : 8) a ete) "30 4 30 | Short Formula (Chemistry) 2303 ba-x) tab) 9 ab—x) Now if ‘B’ is taken in large excess b > > a. = 2808, 8 =p (09 > k b' is very large can be taken as constant w= 22"3 tog? yn 2303 oa akb= lea >? = ¢ ax , k’ is psuedo first order rate constant METHODS TO DETERMINE ORDER OF A REACTION (a) Initial rate method : r= k [A}? [B]’ [C]* if [B] = constant [C] = constant then for two different initial concentrations of Awe have 7 ‘o,f TAgh Y fo = KIA 5 foo = KIASL" = (Rae (b} Using integrated rate law : It is method of trial and error. (c) Method of half lives : 1 forn® order reaction ty, % Ror (d) Ostwald Isolation Method : rate = k [Al* [B}? [C]° = k, [Al® METHODS TO MONITOR THE PROGRESS OF THE REACTION: fa Progress of gaseous reaction can be monitored by measuring total pressure at a fixed volume & temperature 2.303 or by measuring total volume of mixture under constant pressure and temperature. -. k = t log P,(n-1) TP) —P, {Formula is not applicable when n = 1, the value of n can be fractional also.} t (b) By titration method : _ 2.303 Mo 1. “aK, a-x«V, > k= t log vy, 2. Study of acid hydrolysis of an easter. _ 2.303 V..-Vo k= 1 log vV-\ © By measuring optical rotation produced by the reaction mixture : ew 22308 59 [ 902 = 19 o=  Short Formula (Chemistry) EFFECT OF TEMPERATURE ON RATE OF REACTION. K, +10 C.= K = 2 to 3 (for most of the reactions) t Arhenius theroy of reaction rate. I Enthalpy (H) Th hi Jd onthal SH, = Summation of enthalpies of reactants rosholdl enthal Sronety PY SH, = Summation of enthalpies of reactants DH = Enthalpy change during the reaction i Ea. = Energy of activation of the forward reaction = Ea Ea! Ea, = Energy of activation of the backward reaction SH. i Progress of reactor (or reaction ccovdinae) —>> E,>E, —> endothermic E,<E, —> exothermic AH = (E,—£,) = enthalpy change AH=E,-E,, E,, ‘eshold = E.. + E, = E, + E, Arhenius equation k= Ae ERT r=k [cone] dink _ Ey dT RT? Ea 1 log k= ("2303 R JT "199" If k, and k, be the rate constant of a reaction at two different temperature T, and T, respectively, then we have log K2 E, (=- OT,” 2803 R Ty Ina & Ink=InA- Ee slope =— Ee E,20 * Tow,.K oA Ink REVERSIBLE REACTIONS “ k=A eEa!RT =A gia RT k, =A, @ f kw Be AL) eu eine we Ky 7 AY “"30 4 32 \ Short Formula (Chemistry) __4AH {At In KF By +In Ay InK., Ink,. endothermic 4/T exothermic VT BI) . K, 5 - E,, ky, +E, ko [IC] ~ Kz ° ky +p (ii) REVERSIBLE 15’ ORDER REACATION ( both forward and backward } enh] it Ky (iii) SEQUENTIAL 15' ORDER REACTION IAl= [Ale x=a(t— e*") Ka 1 K, Kit ket = yt K, —Ky i ee (K,-K,) In Kp Dane CASE-I K,>>K, CASE IT: K, >> K,  35 { Short Formula (Chemistry) (n-1) d*? ns'* (d) f-Block elements General electronic configuration is (n - 2) f** (n— 1) d* ns*. All f-black elements belong to 3" group. Elements of f-blocks have been classified into two series. (1) 1* inner transition or 4 f-series, contains 14 elements ,.Ce to ,,Lu. (2). IInd inner transition or 5 f-series, contains 14 elements ,,Th to ,,,Lr. Prediction of period, group and block : oO Period of an element corresponds to the principal quantum number of the valence shell. oO The block of an element corresponds to the type of subshell which receives the last electron. o The group is predicted from the number of electrons in the valence shell or/and penultimate shell as. follows. (a) Fors-block elements ; Group no. = the no. of valence electrons (b) For p-block elements ; Group no. = 10 + no. of valence electrons (c) Ford-block elements ; Group no. = no. of electrons in (n — 1) d sub shell + no. of electrons in valence shell Metals and nonmetals : ¢ The metals are characterised by their nature of readily giving up the electron(s) and from shinning lustre. Metals comprises more than 78% of all known elements and appear on the left hand side of the periodic table. Metals are usually solids at room temperature (except mercury, gallium). They have high melting and boiling points and are good conductors of heat and electricity. Oxides of metals are generally basic in nature (some metals in their higher oxidation state form acid oxides e.g. CrO,). . Nonmetals do not lose electrons but take up electrons to form corresponding anions. Nonmetals are located at the top right hand side of the periodic table. Nonmetals are usually solids. liquids or gases at room temperature with low melting and boiling points. They are poor conductors of heat and electricity. Oxides of nonmetals are generally acidic in nature. Metalloids (Semi metals) : The metalloids comprise of the elements B, Si, Ge, As, Sb and Te. Diagonal relationship : 2" period Li Be B c 3” period Na My Al Si Diagonal relationship arises because of ; @ on descending a group, the atoms and ions increase in size. On moving from left to right in the periodic table, the size decreases. Thus on moving diagonally, the size remains nearly the same (Liz 1.23 A & Mg = 1.36 A; Lit = 0.76 A & Mg = 0.72 A) di) it is sometimes suggested that the diagonal relationship arises because of diagonal similarity in electronegativity values. (Li= 1.0 &Mg=1.2; Be=15&Al=1.5;B=2.0 & Si= 1.8) The periodicity of atomic properties : (i) Effective nuclear charge : The effective nuclear charge (Z.,) = Z—c, (where Z is the actual nuclear charge (atomic number of the element) and c is the shielding (screening) constant). The value of o i.e. shielding effect can be determined using the Slater's rules. (ii) Atomic radius : (A) Covalent radius : It is one-half of the distance between the centres of two nuclei (of like atoms) bonded by a single covalent bond. Covalent radius is generally used for nonmetals. "36 4 36 i Short Formula (Chemistry) (B) Vander Waal’s radius (Collision radius) : It is one-half of the internuclear distance between two adjacent atoms in two nearest neighbouring molecules of the substance in solid state. (Cy Metallic radius (Crystal radius) : Itis one-half of the distance between the nuclei of two adjacent metal atoms in the metallic crystal lattice. ¢ Thus, the covalent, vander Wall's and metallic radius magnitude wise follows the order, Foran ~ Feyysas * Maree way Variation in a Period Variation in a Group Ina period left to right In @ group top to bottom Nuclear charge (Z) increases by one unit Nuclear charge (Z) increases by more than one unit Effective nuclear charge (2.1) almost remains constant because of increased screening effect of inner shells electrons Effective nuclear charge (Zes) also increases But number of orbitals (n) remains constant But number of orbitals (n) increases. The effect of increased number of atomic shells overweighs the effect of increased nuclear charge. As a result of this the size of atom increases from top to 4 bottom in a given group z Hence atomic radii decrease with increase in atomic number in a period from left to right. As a result, the electrons are pulled closer to the nucleus by the increased Zer. Tne (iii) lonic radius : The effective distance from the centre of nucleus of the ion up to which it has an influence in the ionic bond is called ionic radius. Cation Anion It is formed by the lose of one or more electrons from| the valence shell of an atom of an element. Cations are smaller than the parent atams because, (i) the whole of the outer shell of electrons is usually] removed. Tt is formed by the gain of one or more electrons in thel valence shell of an atom of an element. Anions are larger than the parent atoms because (i) anion is formed by gain of one or more electrons in the} neutral atom and thus number of electrons increases but {ii} in a cation, the number of positive charges on the| nucleus is greater than number of orbital electrons} leading to incresed inward pull of remaining electrons} magnitude of nuclear charge remains the same. (ii) nuclear charge per electrons is thus reduced and the] electrons cloud is held less tightly by the nucleus leading to} the expansion of the outer shell. Thus size of anion is| increased causing contraction in size of the ion, lonisation Energy : lonisation energy (IE) is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a cation. Mig) —©0) , M@)te ; Mt (g) + IE, ——> M*(g) + &- M** (g) + IE, —> Mig) + e&- IE,, IE, &IE, are the Is, Ilr? & Ill ionization energies to remove electron from a neutral atom, monovalent and divalent cations respectively. In general, (IE), < (IE), <{IE),<..... (iv) * Factors Influencing lonisation energy (A) Size of the Atom : lonisation energy decreases with increase in atomic size. (B) Nuclear Charge : The ionisation energy increases with increase in the nuclear charge. (cy Shielding or screening effect : The larger the number of electrons in the inner shells, greater is the screening effect and smaller the force of attraction and thus ionization energy (IE) decreases. (D) Penetration effect of the electron : Penetration effect of the electrons follows the orders> p>d> f for, the same energy level. Higher the penetration of electron higher will be the ionisation energy.  Short Formula (Chemistry) () ) C©00000 0 Cc Electronic Configuration : If an ator has exactly half-filled or completely filled orbitals, then such an arrangement has extra stability. Electron Gain Enthalphy : (CHANGED TOPIC NAME) The electron gain enthalpy AH", is the change in standard molar enthalpy when a neutral gaseous atom gains an electron to form an anion. X(g) +e (Q)—>X @) The second electron gain enthalpy, the enthalpy change for the addition of a second electron to an initially neutral atom, invariably positive because the electron repulsion out weighs the nuclear attraction. Group 17 elements (halogens) have very high negative electron gain enthalpies (i.e. high electron affinity) because they can attain stable noble gas electronic configuration by picking up an electron. Across a period, with increase in atomic number, electron gain enthalpy becomes more negative As we move in a group from top to bottom, electron gain enthalpy becomes less negative Noble gases have large positive electron gain enthalpies Negative electron gain enthalpy of O or F is less than S or Cl. Electron gain enthalpies of alkaline earth metals are very less or positive Nitrogen has very low electron affinity (i) Electron affinity « (i) Electron affinity « Effective nuclear charge (,,) Atomic size (ii) Electron affinity « {iv) Stability of half filled and completely filled orbitals of a Screening effect ° subshell is comparatively more and the addition of an extra electron to such an system is difficult and hence the electron affinity value decreases. (VI) Electronegativity : (a) (b) (vil) Electronegativity is a measure of the tendency of an element to attract shared electrons towards itself in a covalently bonded molecules. Pauling’s scale : A=X,-X,= 0.208 JE - yEaaXEen E, , = Bond enthalpy/ Bond energy of A—B bond E,_, = Bond energy of A—A bond E,_, = Bond energy of B- B bond (All bond energies are in kcal / mol} A=X,-X,= 0.1017 Eng -yExa*Eos All bond energies are in kJ / mol Mulliken’s scale : IE+EA Say = Paulings's electronegativity x, isrelated to Mulliken’s electronegativity y,, as given below. %p= 1.35 (y)!2— 1.37 Mulliken’s values were about 2.8 times larger than the Pauling’s values Periodicity of Valence or Oxidation States : There are many elements which exhibit variable valence. This is particularly characteristic of transition elements and actinoids. “to 4 40 1 Short Formula (Chemistry) 1 Coordinate Bond (Dative Bond): The bond formed between two atom in which contribution of an electron pair is made by one of them while the sharing is done by both. + o>) ‘ (@ NH4 {ammonium ion) Se H a) -H Donor Acceptor os O (ii) ©, (@@zone) 66 or Jo ‘\ Other examples : H,SO,, HNO, , H,O* , N,O, [Cu(NH,),]** Formal Charge : Total numiver of vaierice Total number of non bonding election in the free atom |“) (ione pair) electrons fotal number of bending (shared) electrons structure -us 5 ov :Q Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species. Limitations of the Octet Rule : 1. The incomplete octet of the central atom LiCl, BeH, and BCL, AICI, and BF, 2. Odd-electron molecules nitric oxide, NO and nitrogen dioxide. NO, N=G O=N-O: 3. The expanded octet PF, H,SO. 10 slectrons around the P atom 12 electrons around the & atorii 12 alactrons around the S atom 4. Other drawbacks of the octet theory @ some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like XeF, , KrF, , XeOF, etc., (i) This theory does not account for the shape of molecules. (i) It does not explain the relative stability of the molecules being totally silent about the energy of a molecule “a1 41. 1 Short Formula (Chemistry) Valence bond theory (VBT) : H,(g) + 435.8 kJ mol- > Hig) + HQ) Distance of separation ~ Bond 435.8 Energy Bond Length 74nm Imemucies: cistanca-——> Orbital Overlap Concept according to orbital overlap concept, the formation of a covalent bond between two atoms results by pairing of electrons present, inthe valence shell having opposite spins Types of Overlapping and Nature of Covalent Bonds The covalent bond may be classified into two types depending upon the types of overlapping : (i) Sigma(c) bond, and (ii) pi (x) bond {i) Sigma (c) bond : This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. ®s-s overlapping ene 4 samt + ¢aam ———> teenage <ae as: aauar s—orbital s-crbital s-s overlapping @s-p overlapping: 8 ae eee Hl + Qe —— iid ate s-orbiiai p-ordiiai $-p orbital ®@ p-p overlapping : This type of overlap takes place between half filled p-orbitals of the two approaching atoms. p-orbital p-orbital p-p overlaping {ii) pi(z) bond : In the formation of x bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the intemuclear axis. oo 8 | p-orbital p-orbital p-e oveilaping Strength of Sigma and pi Bonds : In case of sigma bond, the overlapping of orbitals takes place to a larger extent. Hence, it is stronger as compared to the pi bond where the extent of overlapping occurs to a smaller extent.  Short Formula (Chemistry) Valence shell electron pair repulsion (VSEPR) theory : ® i) iid, W 0] “) PepS {iv) Determination of hybridisation of an atom in a molecule or ion: Steric number rule (given by Gillespie) : Steric No. of an atom = number of atom bonded with that atom + number of lone pair(s) left on that atom. The main postulates of VSEPR theory are as follows: The shape of a molecule depends upon the number of valence shell electron pairs [bonded or nonbonded) around the central atom. Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged. These pairs of electrons tend to occupy such positions in space that minimise repulsion and thus maximise distance between them. The valence shell is taken as a sphere with the electron pairs localising on the spherical surface at maximum distance from one another. A multiple bond is treated as if itis a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair. Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure. The repulsive interaction of electron pairs decreases in the order : lone pair (#'p) - lone pair (ép) > lone pair (ép) - bond pair (bp) > bond pair (bp) -bond pair (bp) Hybridisation : Salient features of hybridisation : The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised. The hybridised orbitals are always equivalent in energy and shape. The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement is obtained. Therefore, the type of hybridisation indicates the geometry of the molecules. Important conditions for hybridisation : The orbitals present in the valence shell of the atom are hybridised. The orbitals undergoing hybridisation should have almost equal energy. Promotion of electron is not essential condition prior to hybridisation. It is the orbital that undergo hybridization and not the electrons. Table-3 Steric number Types of Hybridisation Geometry 2 sp Linear 3 sp? Trigonal planar 4 sp? Tetrahedral 5 sp'd Trigonal bipyramidal 6 sp? d* Octahedral 7 sp® d® Pentagonal bipyramidal Hybridization Involving d-orbital : Type of ‘d’ orbital involved sped dz? sped? dx? - y? & dz? sp'd> dx? - y?, dz? & dxy dsp? dey "as 4 45 1 Short Formula (Chemistry) 1 {B) Intenmolecular H-Bonding : it is formed between two different molecules of the same or different compounds. (a) In water molecules Se Bt Be b+ SF Se Gt OF H-O..H-O..H-O..H-O Hae Hae Het He (by The hydrogen bonds in HF link the F atom of one molecule with the H-atom of another molecule, thus forming a zig-zag chain (HF), in both the solid and also in the liquid. Intermolecular forces (Vander Waal’s Forces) : Intermolecular attractions hold two or more molecules together. These are weakest chemical forces and can be of following types. (a) lon-dipole attraction : (b} Dipole-dipole attraction : (e) lon-induced dipole attraction : (d) Dipole-induced dipole attraction : (e} Instantaneous dipole- Instantaneous induced dipole attraction : (Dispersion force or London forces) So Strength of vander waal force -- molecular mass. oO van der Waal’s force « boiling point. Metallic bond : Two models are considered to explain metallic bonding: (A) Electron-sea model (B) Band model Some special bonding situations : (a) Electron deficient bonding: There are many compounds in which some electron deficient bonds are present apart from normal covalent bonds or coordinate bonds. These electron deficient bonds have less number of electrons than the expected such as three centre-two electron bonds (3c-2e) present in diborane B,H,, Al,(CH.),, BeH,(s) and bridging metal carbonyls. (b) Back Bonding : Back bonding generally takes place when out of two bonded atoms one of the atom has vacant orbitals (generally this atom is from second or third period) and the other bonded atom is having some non-bonded electron pair(generally this atom is from the second period). Back bonding increases the bond strength and decreases the bond length. For example, in BF, Vacant Filled 2p-orital 2p-orbital the extent of back bonding in boron trihalides. BF, > BCI, > BBr, "ae 4 46 1 Short Formula (Chemistry) COORDINATION COMPOUNDS Addition Compounds : They are formed by the combination of two or more stable compounds in stoichiometric ratio. These are (1) Double salts and (2) Coordination compounds Double salts : Those addition compounds which lose their identity in solutions eg. K,S0, , Al,(SO,)5 Coordination Compounds : Those addition compounds which retain their identity (i.e. doesn’t lose their identity) in solution are Fe(CN), + 4KCN > Fe(CN),. 4KCN or K, [Fe(CN)g] (aq.) ===> 4K* (aq.) + [Fe(CN),]* (aq.) Central Atom/lon : In a coordination entity-the atom/ion to which are bound a fixed number of ligands in a definite geometrical arrangement araund it. Ligands : The neutral molecules, anions or cations which are directly linked with central metal atom or ion in the coordination entity are called ligands. Chelate ligand : Chelate ligand is a di or polydentate ligand which uses its two or more donor atoms to bind a single metal ion producing a ring. Ambidentate Ligand : Ligands which can ligate through two different atoms present in it < a . MeN So nitrito-N : M<O—N=O nitrito-O M<SCN_thiocyanato orthiocyanato-S_ ; M+« NCS _ isothiocyanato orthiocyanato-N Coordination Number: The number of ligand donor atoms ta which the metal is directly attached. Oxidation number of Central Atom : The oxidation number of the central atom is defined as the charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom. [Fe(CN),]* is +3 and it is written as FeiIll). “a7 } 47 1 Short Formula (Chemistry) Denticity and Chelation : Table : 1 Common Monodentate Ligands Common Name JUPAC Name Formula methyl isocyanide methylisocyanide HNC triphenyl phosphine triphenyl phosphine/tripheny! phosphane | PPh: pyridine pyridine CcHeN (py) ammonia ammnine NH3 methyl amine methylamine MeNHy water aqua or aquo HO carbonyl carbonyl co thiocarbony! thiocarbonyl cs nitrosyl nitrosyl NO fluoro fluoro or fuorido* F chloro chloro or chloride* or bromo bromo or bromido* Br iodo iode or iodido* r cyano cyanide or cyanido-C* (C-bonded) ene isocyano isocyanido or cyanido-N* (N-bonded) NC- thiecyano thiocyanato-S(S-bonded) SCN" isothiocyano thiocyanato-N(N-bonded) NOS" cyanate (cyanate) eyanate-O (O-bonded) OCN isocyanate dsocyanate) eyanato-N (N-bonded) NCO" hydroxo, hydroxo or hydroxide” OH nitro nitrito-N (N-bonded} NOS nitrite: nitrita-O (O-bended) ONO™ nitrate nitrato: NOs amido amido NH imido imido NH? nitride nitrido| Ne azido azido Nae hydride hydride oxide oxido: peroxide peroxide superoxide ‘superoxido acetate acetato [sulphate sulphato thiosulphate thiosulphato sulphite sulphito hydrogen sulphite hydragensulphito sulphide sulphido or thio hydrogen sulphide hydrogensulphido or mercapto thionitrito thionitrito nitresylium nitrosylium or nitrosonium nitronium nitronium NO.* * The 2004 IUPAC draft recommends that anionic ligands will end with "so 4 50 1 Short Formula (Chemistry) 1 Werner's Theory : According to Werner most elements exhibit two types of valencies : (a) Primary valency and (b) Secondary valency. {a) Primary valency : This corresponds to oxidation state of the metal ion. This is also called principal, ionisable or ionic valency. It is satisfied by negative ions and its attachment with the central metal ion is shown by dotted lines. {b) Secondary or auxiliary valency : Itis also termed as coordination number (usually abbreviated as CN) of the central metal ion. It is non-ionic or non-ionisable (i.e. coordinate covalent bond type). In the modem terminology, such spatial arrangements are called coordination polyhedra and various possibilities are CN.=2 linear CO.N.=4 tetrahedral or square planar Effective Atomic Number Rule given by Sidgwick : Effective Atomic Number (EAN) = Atomic no. of central metal - Oxidation state of central metal + No. of electrons donated by ligands. Triangular octahedral. Valence bond theory : The model utilizes hybridisation of (n-1) d, ns, np or ns, np, nd orbitals of metal atom orion to yield a set of equivalent orbitals of definite geometry to account for the observed structures such as octahedral, square planar and tetrahedral, and magnetic properties of complexes. The number of unpaired electrons, measured by the magnetic moment of the compounds determines which d-orbitals are used. Table : 5 Coordiantion number of metal | Type of hybridisation | Shape of complex 4 sp’ Tetrahedral 4 dsp* Square planer 5 Trigonal bipyramidal 6 Octahedral 6 Octahedral Coordination Number Six : In the diamagnetic octahedral complex, [Co(NH,),]*, the cobalt ion is in +3 oxidation state and has the electronic configuration represented as shown below. [Co(NH,),F* tit d?sp* hybrid orbital The complex [FeF,]* is paramagnetic and uses out outer orbital or high spin or spin free complex. So, er orbital (4d) in hybridisation (sp°d?) ; it is thus called as [Fer I> WA a 4 spd? hybrid orbitals Coordination Number Four : Inthe paramagnetic and tetrahedral complex [NiCI,]*, the nickel is in +2 oxidation state and the ion has the electronic configuration 3d*. The hybridisation scheme is as shown in figure INCL Abb fiby 4 sp® hybrid orbitals Fees 6. Short Formula (Chemistry) Similarly complex [Ni(CO),] has tetrahedral geometry and is diamagnetic as it contains no unpaired electrons. The hybridisation scheme is as shown in figure. wieo — LAT ETAT sp* hybrid orbitals The hybridisation scheme for [Ni(CN),]* is as shown in figure. INKCN) > ALLARD dsp? hybrid orbitals It suffers from the following shortcomings : A number of assumptions are involved. There is no quantitative interpretation of magnetic data. It has nothing to say about the spectral (colour) properties of coordination compounds. It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds. It does not make exact predictions regarding the tetrahedral and square-planar structures of 4-coordinate complexes. It does not distinguish between strong and weak ligands. Magnetic Properties of Coordination Compounds : Magnetic Moment = Vl A(2+2) Bohr Magneton ; n = number of unpaired electrons For metal ions with upto three electrons in the d-orbitals like Ti**, (d') ; V°* (d?) ; Cre* (d°) ; two vacant d-orbitals are easily available for octahedral hybridisation. The magnetic behaviour of these free ions and their coordination entities is similar. When more than three 3d electrons are present, like in Cr? and Mn* (d*) ; Mn and Fe* (d*) ; Fe** and Co* (d*) : the required two vacant orbitals for hybridisation is not directly available (as a consequence of Hund’s rules). Thus, for d*, d> and d® cases, two vacant d-orbitals are only available for hybridisation as a result of pairing of 3d electrons which leaves two, one and zero unpaired electrons respectively. Crystal Field Theory : The crystal field theory (CFT) is an electrostatic model which considers the metal-ligand bond to be ionic arising purely from electrostatic interaction between the metal ion and the ligand (a) Crystal field splitting in octahedral coordination entities : Fnerqy + ¥ Lreigy coirespana 4 W Boyde ato absorntion. / “ average Qo +: onetyy (Ban cenier) , a ~N level Metal / Average enetdy We Om rye. fd + Shad, LO Pus Cus io. avoludon o1 enemy, Stee. fad spaing t l i i i ! x 1 dorbitats 7 of t erbiteir ip ephorivel | i i i | jor oclaheural complex i Fiée metal iain Figure showing crystal field splitting in octahedral complex. The crystal field splitting, Aj, depends upon the fields produced by the ligand and charge on the metal ion.  §2 { Short Formula (Chemistry) Ligands can be arranged in a series in the orders of increasing field strength as given below : F< Br < SCN’ < Cl < 8% < F-< OH-< C,0,2-< HO < NCS < edta*+ < NH, < en <NO,-< CN-<CO Such a series is termed as spectrochemical series. It is an experimentally determined series based on the absorption of light by complexes with different ligands. Calculation of Crystal Field stabilisation energy (CFSE) Formula: CFSE = [- 0.4 {n) tog + 0.6 (n’} 4] Ag+ *nP. where n & n’ are number of electron(s) in tog & e, orbitals respectively and A, crystal field splitting energy for octahedral complex. *n represents the number of extra electron pairs formed because of the ligands in comparison to normal degenerate configuration. ({b) Crystal field splitting in tetrahedral coordination entities : In tetrahedral coordination entity formation, the d orbital splitting is inverted and is smaller as compared to the octahedral field splitting. For the same metal, the same ligands and metal-ligand distances, it can be shown that A, = (4/9)Ap, Encigy Spleting of & ornita's Free metal ion. in, tetrahedires crystal flela. Figure showing crystal field splitting in tetrahedral complex. COLOUR IN COORDINATION COMPOUNDS : According to the crystal field theory the colour is due to the d-d transition of electron under the influence of ligands. We know that the colour of a substance is due to the absorption of light at a specific wavelength in the visible part of the electromagnetic spectrum (400 to 700 nm) and transmission or reflection of the rest of the wavelengths. Limitations of crystal field theory a) It considers only the metal ion d-orbitals and gives no consideration at all to other metal orbitals (such SS, P,, P, and p, orbitals) (2) It is unable to account satisfactorily for the relative strengths of ligands. For example it gives no explanation as to why H,O is a stronger ligand than OH in the spectrochemical series. Q) According to this theory, the bond between the metal and ligands are purely ionic. It gives no account on the partly covalent nature of the metal ligand bonds. (4) The CFT cannot account for the m-bonding in complexes. STABILITY OF COORDINATION COMPOUNDS : The stability of a coordination compound [ML,] is measured in terms of the stability constant (equilibrium constant) given by the expression, B, = [ML,VIM(H,), IL" for the overall reaction : M(H,O), + nL == ML, + nH, By convention, the water displaced is ignored, as its concentration remains essentially constant. The above overall reaction takes place in steps, with a stability (formation) constant, K,, K,, K,, ...... K, for each step as represented below :  55 { Short Formula (Chemistry) a a b A A b a } A a f M M \ M A b Al b Al a a b b ala aTb bTb Note: With [M(AA)b,]. only one form is possible. M(AA)abcd have six geometrical isomers. aii) M(AA),O, -Two geometrical isomers are possible. Ci cl i cl oN ao a 7 | > 7) NL cl XQ en Geometrical isomers (cis and trans) of [CoCI,{en).] Optical lsomerism : Acoordination compound which can rotate the plane of polarised light is said to be optically active. Octahedral complex : Optical isomerism is common in octahedral complexes involving didentate ligands. For example, [Co(en).,]** has d and ¢ forms as given below. oN tN z ! eh dextre mirror taeve dand ¢ of [Co(en),]* Square planar complex : Square planar complexes are rarely found to show the optical isomerism. The plane formed by the four ligating atoms and the metal ion is considered to be a mirror plane and thus prevents the possibility of chirality. Organometallic compounds Metal Carbonyls : Compounds of metals with CO as a ligand are called metal carbonyls. They are of two types. {a) Monomeric : Those metal carbonyls which contain only one metal atom per molecule are called monomeric carbonyls. For examples :[Ni(CO),] (sp, tetrahedral) ; [Fe(CO),] (dsp*, trigonal bipyramidal). {b) Polymeric : Those metal carbonyls which contain two or more than two metal atoms per molecule and they have metal-metal bonds are called polymeric carbonyl. For example : Mn, (CO),), Co.(CO),, etc. "get 56 i Short Formula (Chemistry) 1 The M— C x bond is formed by the donation of a pair of electrons from a filled d orbital of metal into the vacant antibonding n* orbital of carbon monoxide. Thus carbon monoxide acts as c donor (OC > M) and a n acceptor (OC <M). with the two interactions creating a synergic effect which strengthens the bond between CO and the metal as shown in figure. Synergic bonding Sigma (o) bonded organometallic compounds : (a) Grignard’s Reagent R—-—Mg-—X where Risa alkyl or aryl group and X is halogen. (b) (CH,), Sn, (C,H.), Pb, Al, (CH), Aly (CpH,), eto. Pie (z)-bonded organometallic compounds : These are the compounds of metal with alkenes, alkynes, benzene and other ring compounds. ELS AS Zeise’s salt: K [PtCl, (n?-C,H,)] / \ f A = iN Ferrocene and bis(benzene)chromium : Fe (n° -C,H,), and Cr(?-C5H,), Cr METALLURGY The compound of a metal found in nature is called a mineral. The minerals from which metal can be economically and conveniently extracted are called ores. An ore is usually contaminated with earthy or undesired materials known as gangue. {a) Native ores contain the metal in free state. Silver, gald, platinum etc, occur as native ores. {b) Oxidised ores consist of oxides or oxysalts (e.g. carbonates, phosphates, sulphates and silicates ) of metals. {c) Sulphurised ores consist of sulphides of metals like iron, lead, zinc, mercury etc. {d) Halide ores consist of halides of metals.  Short Formula (Chemistry) \Metak as # a2 eee a a Orestetusawasaeee aw Gomposition sess as eee ees Aluminium Bauxite |AIOy(OH)s3_2x [where 0 < X < 1] AO, Diaspore Al,O3.H20 Corundam ALLO; Kaolinite (a form of clay) [Als (OH)4 Siz05] Iron Haematite Fe,O3 Magnetite Fe,0, Siderite Feco, Iron pyrite Fes, Limonite Fe03.3H20 Copper Copper pyrite CuFeS, Copper glance Cus Cuprite Cu, Malachite CucO3.Cu(OH}2 Azurite 2CuCO3.Cu(OH)s Zinc Zinc blende or Sphalerite ZnS Calamine ZNO, Zincite ZnO. Lead Galena Pbs Anglesite PhSO, Cerrusite Pbco, Magnesium Carnallite KCI.MgCl, 6H, (K»MgCl, .6H,0) Magnesite MgCO,, Dolomite MgCO, CaCO, Epsomsalt (Epsomite) MgSO. 7H,0 Langbeinite KMg-(SOq)3 Tin Cassiterite (Tin stone) SnO; Silver Silver glance (Argentite) Ags Chlorargyrite (Horn silver) AgCl Metallurgy : (A) {B) The scientific and technological process used for the extraction/isolation of the metal from its ore is called as metallurgy. The isolation and extraction of metals from their ores involve the following major steps: Crushing and Grinding : The ore is first crushed by jaw crushers and ground to a powder. Concentration : The removal of unwanted useless impurities from the ore is called dressing, concentration or benefaction of ore “Sof 60 | Short Formula (Chemistry) (b) Electrolytic reduction (Hall-Heroult process) : 2Al,0,+ 3G ——> 4Al + 3C0, Cathode : AF (melt) + 3e- ——s Ald) Anode: C(s) + O? (melt) ——> CO(g) + 2e- C(s) + 20* (melt) —— CO, (g) + 4e- Metallurgy of some important metals Extraction of iron from ore haematite : Reactions involved : At 500 - 800 K (lower temperature range in the blast furnace) 3Fe,0,+CO ——> 2Fe,0,+ CO, Fe,O,+CO ——> 3Fe+4CO, Fe,O,+CO ——> 2Fe0+CO, At 900 - 1500 K (higher temperature range in the blast furnace): C+CO, ——> 2C0; FeO +CO ——>Fe+Co, Limestone is also decomposed tom CaO which removes silicate impurity of the ore as slag. The slag is in molten state and separates out from iron. CaCO, —+CaO+C0,; CaO+ SiO, —>CaSiO, 2. Extraction of copper : From copper glance / copper pyrite (self reduction) : 2CuFeS,+40,—+ Cu,S + 2FeO + 380, Cu,S + FeO + SiO, —> FeSiO, (fusible slag) + Cu,S (matte) 2FeS +30, ——> 2FeO +280, ; FeO+Si0, ——> Fesio, 20u,8 + 30, —— 2Cu,0+280,; 2Cu,O0 + Cu,S ——> 6Cu+ SO, (self reduction) 3. Extraction of lead : @ —— 2PbS(s) +30, (g) 2+ 2PbO (s) > 2P HY") + CO, @) A (i) —-3Pbsis) —T"_, Pps is) + 2PHO (s) —H*™ = apniy +80, @) air absence of air 4. Extraction of zinc from zinc blende : The ore is roasted in presence of excess of air at temperature 1200 K. 2ZnS + 30, —> 2Zn0 +280, The reduction of zinc oxide is done using coke. Zn0 + —SH18BK , Pn +00 5. Extraction of tin from cassiterite : The concentrated ore is subjected to the electromagnetic separation to remove magnetic impurity of Wolframite. SnO, is reduced to metal using carbon at 1200 — 1300°C in an electric furnace. The product often contains traces of Fe, which is removed by blowing air through the molten mixture to oxidise FeO which then floats to the surface. SnO, + 20 —> Sn + 200 2Fe + Oy > 2FeO 6. Extraction of Magnesium : From Sea water (Dow’s process) : Sea water contains 0.13% magnesium as chloride and sulphate. It involves following steps. @ Precipitation of magnesium as magnesium hydroxide by slaked lime. (b) Preparation of hexahydrated magnesium chloride. "31 61 {D) t \ Short Formula (Chemistry) The solution on concentration and crystallisation gives the crystals of MgCl,.6H,O. [o) Preparation of anhydrous magnesium chloride. (d) Electrolysis of fused anhydrous MgCl, in presence of NaCl. MgCl, == Mg* + 2CI- Atcathode: Mg +2e° ——> Mg(99% pure); Atanode : 2Cr — > Cl, + 2e Extraction of gold and silver (MacArthur-Forrest cyanide process) : (a) From native ores : Extraction of gold and silver involves leaching the metal with CN-. 4Au/Ag (8) + 8CN (aq) + 2H,O(aq) + O,(g) —— 4[Au/Ag (CN),|(aq) + 40H—{aq) 2[Au Ag (CN),/(aq) + Zn(s) ——> 2Au/Ag (s) + [Zn(CN),J* (aq) (b) From argentite ore : Ag,S (conc. ore) + 2NaCN “aR 2AgCN + NaS. 4Na,S + 50, + 2H,O ——> 2Na,SO, + 4NaQH + 2S AgCN + NaCN ——; NalAg(CN).] (soluble complex) 2Na[Ag(CN),] + Zn (dust) ——> 2Ag | + NaZn(CN),]. Purification or Refining of metals : Physical methods : These methods include the following processes : (I) Liquation process : This process is used for the purification of the metal, which itself is readily fusible, but the impurities present in it are not, used for the purification of Sn and Zn, and for removing Pb from Zn-Ag alloy. (Il) Fractional distillation process : This process is used to purify those metals which themselves are volatile and the impurities in them are nonvolatile and vice-versa. Zn, Cd and Hg are purified by this process. (II) Zone refining method (Fractional crystallisation method) : This process is used when metals are required in very high purity, for specific application. For example pure Si and Ge are used in semiconductors Chemical methods : These methods include the following methods () OXIDATIVE REFINING : This method is usually employed for refining metals like Pb, Ag, Cu, Fe, etc. In this method the molten impure metal is subjected to oxidation by various ways. i) POLING PROCESS : This process is used for the purification of copper and tin which contains the impurities of their own oxides. Green wood — Hydrocarbons > CH, 4CuO + CH, — 4Cu (pure metal) + CO, + 2H,O (Ill) ELECTROLYTIC REFINING : Some metals such as Cu, Ni, and Al are refined electrolytically. (IV) VAPOR PHASE REFINING : ti) Extraction of Nickel (Mond's process) : The sequence of reaction is H,Og) + > CO@) +H, Ni(s) +4 CO(s) _8°C_, [Ni(CO)1 @) INi (CO),}(g) 200°C _, Ni + 4CO@) (ii) Van Arkel-De Boer process : Impure Ti+ 21, 8-250, iy, 100" _, Ti +21, Tungsten fament  Short Formula (Chemistry) s-BLOCK ELEMENTS & THEIR COMPOUNDS Group 1 of the periodic table consists of the elements : lithium, sodium, potassium. rubidium, caesium and francium . The elements of Group 2 include beryllium, magnesium, calcium, strontium, barium and radium. Hydration Enthalpy : The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.Li- has maximum degree of hydration and for this reasons lithium salts are mostly hydrated e.g., LiCl .2H,O Physical properties : Allthe alkali metal are silvery white, soft and light metals. Because of the larger size, these element have low density. The melting and boiling point of the alkali metals are low indicating weak metallic bonding alkali metals and their salts impart characteristic colour to an oxidizing flame. Chemical Properties: The alkali metal are highly reactive due to their larger size and low ionization enthalpy. © Reactivity towards air : They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxide. © Reducing nature: The alkali metals, are strang reducing agents, lithium being the most and sodium the least powerful. © Solution in liquid ammonia: The alkali metals dissolve in liquid ammonia giving deep blue solution which are conducting in nature. M+ (x+y) NH, —— [M(NH, ), I + [e(NH,),F- The blue colour of the solution is due to the ammoniated electron and the solutions is paramagnetic. M-(am) + e+ NH, (-) —2°S284"9_ NH, (am) + 1/2 H,(g) In concentrated solution, the blue colour changes to bronze colour and becomes, diamagnetic. ANOMALOUS PROPERTIES OF LITHIUM (i) exceptionally small size of its atom and ion, and (ii) high polarising power (/.e., charge/ radius ratio }. The similarity between lithium and magnesium is particularly striking and arises because of their similar size: atomic radii, Li= 152 pm, Mg = 160 pm; ionic radii : Lit = 76 pm, Mg? = 72 pm. GROUP 2 ELEMENTS : ALKALINE EARTH METALS The first element beryllium differs from the rest of the member and shows diagonal relationship to aluminium. Hydration Enthalpies Hydration enthalpies of alkaline earth metal ions. Be? > Mg?* > Ca > Sr? > Ba**. The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals , e.g., MgCl, and CaCl, exist as MgCl, .6H,O and CaCl,. 6H,O while NaCl and KCI do not form such hydrates. Physical Properties The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. The melting and boiling point of these metals are higher due to smaller sizes. Because of the low ionisation enthalpies they are strongly electropositive in nature. The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence these elements do not impart any colour to the flame. Calcium, strontium and barium impart characteristic colour to the flame. "get 65 | Short Formula (Chemistry) 6. Potassium sulphate (K,SO,): 7. Quick Lime, Slaked Lime and Lime Water : Limestone 1000S, quick tine 22. siaked lire (Caco,) (Cad) [Ca(OH),] Suspension ~——— Suspended in water (Milk of lime) | Ca(OH), [ite Clear solution (Lime wateri Ca(OH), H,O_» C,H, Acetylene Heated in elecitic Lies GH, Acetylene + Coke Cat furnace at 2000°C oN H,O (CaO + 3C + Cad, + CO) Can, +C Ammonium . y. Nitroim compounds * NA¢Ammonia) Ferilizer) | Milk of line (white washing) NH, 40 [-—T> Lime water (for the detection of CO,) Ca —=> Ca(OH), __ CL Bleaching powd Quick ime Sluked [TTT TTT Bisaching powder Ime Na,Co, . Slaking with [> NaOH (Caustic-sada) NaOH solution sio WS Parle or 4paristwater Sodalime Slaked Silica {Amixtureof Ca(OH), and NaOH) lime cod absort numbar of cases good absorbent for numbar of gaset Motion {Building maisrai) 3Ca(OH), + 2Cl, ——> Ca(OCl),. Ca(OH),. CaCl,. 2H,O (bleaching powder}. "Sef 66 | Short Formula (Chemistry) p-BLOCK ELEMENTS & THEIR COMPOUNDS TRENDS IN PROPERTIES OF p-BLOCK ELEMENTS. Electronegativty, ionization enthalpy, oxidizing sower, sce Peso see dso tosh ss] ienigation euthalpy, oxidizing power. Cowalant radius, van der Waals’ radius, enthalpy of ato nization (pte group 14), rretallic charactor {A) GROUP 13 ELEMENTS : THE BORON FAMILY Oxidation state and trends in chemical reactivity : General Oxidation State = + 3. Reactivity towards acids and alkalies 2 Al(s) + 6 HCl(aq) ——-> 2 Al* (aq) + 6 Cr-(aq) + 3 H,(g) 2A\(s) + 2NaOH (aq) + 6H,O (1) ——> 2Na’ [AI(OH),} (aq) + 3H,(g) Sodium tetrahydroxoaluminate (III) Reactivity towards halogens 2E(s) + 3X, (g) > 2EX, (s) (X= F, Cl Br, I) BORON (B): Some Important Reactions of Boron and its compounds : Na BO. |  67 { Short Formula (Chemistry) ba $0,4H,0 aro. Les Race - NoDve | MED HO eo i Na3,0, + [EE OO Ne, NaBes 6,9, 4 Ke, 8 FR), called inca, cui50, 2 ks traohie ip iebeae} io /s aul Rau, cl BH.cl-———+ BH, gs ee - es FPBCL+HO! *° RCH,CH.OH PBR, we BH, ONHOMERES (4, BINH)) HBH H,BO, (BW). x 3 B,R,H, borazine) o Small amines such as NH,, CH,NH, and (CH.),NH give unsymmetrical cleavage of diborane H+ 2NH, —— [H,B (NH,).° + [BH oO Large amines such as (CH,),N and pyridine give symmetrical cleavage of diborane. 2(CH,),N + B,H, ——> 2H,B —— N(CH, ° BH, #2C0 —200°C: 20M _, 2BH..CO (borane carbonyl) {B) GROUP 14 ELEMENTS : THE CARBON FAMILY Carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb) are the members of group 14. Electronic Configuration = ns?np?. Oxidation states and trends in chemical reactivity Common oxidation states = +4 and +2. Carbon also exhibits negative oxidation states. In heavier members the tendency to show +2 oxidation state increases in the sequence Ge < Sn < Pb.  70 { Short Formula (Chemistry) ue When a compound like CH,SiCl, undergoes hydrolysis, a complex cross-linked polymer is obtained. er The hydrocarbon layer along the silicon-oxygen chain makes silicones water-repellent. COMPOUNDS OF LEAD: ye PROF}, PBTO, + GS ihusie fead carhonate} a | 2 PECL [ATO pcre, 4 (Yellow) ko Ki H.SIHC! Net 4 (Yellow) 9, (from airy Eo Or ° IPL(OHS a PEINO,), + PbO, PHS, + H,0 +0, Espoo eHore, | I T PhO + NO, +0. iPio( Ohi COMPOUNDS OF TIN: Soc, LHg,Chy em Hg (white ppt), (giey) Hol \ 140 SMOHICI (SCL [H.SNClJ — Sn{OH), {c) GROUP 15 ELEMENTS : THE NITROGEN FAMILY Electronic Configuration : ns* np* Atomic and lonic Radii : Covalent and ionic (in a particular state) radii increase in size down the group. Physical Properties: Allthe elements of this group are polyatomic. Metallic character increases down the group. The boiling points in general , increase from top to bottom in the group but the melting point increases upto arsenic and then decreases upto bismuth. Except nitrogen , all the elements show allotropy.  71 { Short Formula (Chemistry) Chemical Properties : Oxidation States and trends in a chemical reactivity : The common oxidation states of these elements are -3, +3 and +5. The stability of +5 oxidation state decreases. and that of +3 state increases (due to inert pair effect) down the group ; Bi* > Sb* > As*; Bi < Sb < As* Nitrogen exhibits +1, +2, +4 oxidation states also when it reacts with oxygen Anomalous properties of nitrogen : (i) The stability of hydrides decreases from NH, to BiH, which can be observed from their bond dissociation enthalpy. Consequently , the reducing character of the hydrides increases. Basicity also decreases in the order NH, > PH, > ASH, > SbH, 2 BiH, PROPERTIES OF HYDRIDES OF GROUP 15 ELEMENTS Property NHs PH; AsH; | SbH, BiH; Melting point / K 195.2 | 1395 | 1567 185 - Boiling point /K 2385 | 1855 | 2106 | 2546 290 (E-H)Distancefpm | 101.7 | 1419 | 1519 | 1707 - HEH angle (2) 107.8 93.6 91.8 91.3 - AH / kJ mor -46.1 13.4 66.4 145.1 278 AggeH(E—H)/ kd mor! | 389 322 297 255 - {ii) The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state. Their acidic character decreases down the group. The oxides of the type EO, of nitrogen and phosphorus are purely acidic, that of arsenic and antimony amphoteric and those of bismuth is predominantly basic. {iii) Nitrogen does not form pentahalide due to non — availability of the d-orbitals in its valence shell. Pentahalides are more covalent than trinalides. Halides are hydrolysed in water forming oxyacids or oxychlorides. PCI, + HO H,PO, +HCI; SbCl, + H,O — SbOCI ~ (orange) + 2HCI; BiCl, + HO—> BiOC! 1 (white) + 2HCI {iv} These elements react with metals to form their binary compounds exhibiting -3 oxidation state , such as, Ca,N, (calcium nitride) Ca_P, (calcium phosphide) and Na,As, (sodium arsenide). NITROGEN (N) AND ITS COMPOUNDS : NH, NHAG, (NFL).C1,0 pet coc ‘ ! N , CaCO, +NH <= CatN* ND MoM, ——o NILtMg(Cl 8, (ptrolien} "Fo 4 72 i Fe,0,xH,0 Short Formula (Chemistry) NHNO, Catl-8NH, NO (PDs MINGOH BW) sy - 8 “ or NOy “S HAAPO. CaO +H,0 — Ca(Oii), ;used far drying of NH, Oxides of Nitrogen Name Formula | Oxidation state Common methods of Physical appearance and| of nitrogen preparation chemical nature Heat Dinitragen oxide NO 41 NH.NO.——> N,O + 2 1.0 colourless gas , neutral Nitrogen{T) oxide } Nitrogen monoxide No 42 colourless gas , neutral {Nitragen<TT) oxide} (Nitric acid) Dinitrogen tioxir IN NO, +3 blue solid , acidic Nitrogen dioxide No. +4 brown gas, acidic [Nitrogen(T¥) oxide Oe 0 Dinitrogen tetracxide | N,O. +4 2NO, == NO, colourless solid / INitrogeniiV) oxide] Heat liquid , acidi«c NO, + P.0,——> NO, coger penioxide| NO, +5 [4 HNO,+P.0,, 4HPO,*2N,0, [Nitrogen (tv) oxide] NaOH | NK, ! Nas, ASE, @HSO, | Lil) distiiation, NH NaNO, + NHI NHANO, an dite HNO, je nto (disproportionate) dil HNO, NOC] ‘ [Fe(H,0).NOj" {brown ring) N,O+ NO, “ye } 75 t \ Short Formula (Chemistry) (D) GROUP 16 ELEMENTS : THE OXYGEN FAMILY {i) {ii) (iii) Electronic Configuratiot rns? npt. Atomic and lonic Radii : Due to increase in the number of shells , atomic and ionic radii increase from top to bottom in the group. The size of oxygen atoms is however, exceptionally small. Physical Properties : Oxygen and sulphur are non-metal, selenium and tellurium metalloids, whereas polonium is a metal. Polonium is radioactive and is short lived (Half-life 13.8 days). The melting and boiling points increase with an increase in atomic number down the group. Catenation : Tendency for catenation decreases down the group. This property is prominently displayed by sulphur (S,). The S—S bond is important in biological system and is found in some proteins and enzymes such as cysteine. Chemical Properties Oxidation states and trends in chemical reactivity : Elements of the group exhibit + 2, + 4, + 6 oxidation states but + 4 and + 6 are more common. Anomalous behaviour of oxygen : The anomalous behaviour of oxygen is due to its small size and high electronegativity. The absence of d orbitals in oxygen limits its covalency to four. Their acidic character increases from H,O to H,Te. The increase in acidic character can be understood in terms of decrease in bond (H-E) dissociation enthalpy down the group. Owing to the decrease in bond (H-E) dissociation enthalpy down the group , the thermal stability of hydrides also decreases from H, to H,Po. All the hydrides except water possess reducing property and this property increases from H_S to H,Te. PROPERTIES OF HYDRIDES OF GROUP 16 ELEMENTS [Propertys sae e eee baa oe ole ee MSs ae hee Sex a doo aybe: ae m.p./K 273 188 208 222 b.p/K 373 213 232 269 H-E distance/pm 96 134 146 169 HEH angle (°) 104 92 1 90 A.H/kJ mol * -286 -20 73 100 gigs (H-E)/kJ mort 463 347 276 238 Dissociation constant? 18x 10% 13x 107 1.3104 2.3x*10% Reducing property of dioxide decreases from SO, to TeO, ; SO, is reducing while TeO, is an oxidising agent. Oxides are generally acidic in nature. The stabilities of the halides decrease in the order F > Cl > Br>I. Sulphur hexafluoride SF, is exceptionally stable for steric reasons. The well known monohalides are dimeric in nature, Examples are S.F.,, S.Cl., S,Br,, Se,Cl, and Se_Br,. These dimeric halides undergo disproportionation as given below : 28e,Cl, —> SeCl,+ 38e.  76 { Short Formula (Chemistry) 1 OXYGEN (O,) AND ITS COMPOUNDS : 2- Ethyl anthrequinol sulphide ppt sulphide ppt {CuS, Cd8...As,$,} (208. NiS, Mn, CoS) ©, c0nG. H4$0, oF (NHLSO. Ry 8 & HNO, Fe.(S0,}, FeS + H,80, KS : NOS (ail) Ln cr 0 Z s Fecl, ONS, IFeecny x Ny. THO OHH.OL! FecyrS } nyeg mers (yallowforange coicuration) S+Wné Cr" +3 clo, 8 Men, $+0, 4.80, (Cu*HS0,) 2, ge oO keene) b YL £20," Cr" (green) H.S+0, un so MinQ, iH” Mine ~ woos m. 10% LO age 28 +O, 6 oN \ I, |_N&COUHO NaHSo. —NalO: N2,S0, ee HO ° ee, I Ais so, SO 4 Na,$.0, {hypo} HCI a as NY ac . R oO 80H (no SS bond) S0.{OH) Cl 4 (ohloresu!ahuric acid} letergent Oxo-acids of Sulphur 1. Suplhurous —_ acid series H (a)H,SO, sly) sulphurous acid S=O0 HO: 2. Sulphuric acid series f (a)H,SO, sw sulphuric acid HO—S—-OH CG 3. Peroxo acid series t i {ay H,SO, SMI) peroxomonosulphuric acid HO—s—O—OH o Caro, acid)  77 { Short Formula (Chemistry) ic-Chamber Brt80, $0, FeSO,4(NH,),80,+CO NaH + NaS + Na,S0,+ 1, | Na.$ + $9,+ Ne,CO, Na,SO, NaS, A A a e Coy 2g ay §0,+8 H e i, Ci Na SO, |———2-—+Na,8,0, + Nal NahiSO,+ Hole —— sa ee Je ano. Re 9 Na JAg(S,03} in excass of Ne,2,0,) (Fe(S,0,)] ags,o+ uct oF NaslAg(s:0.)) (Pink or violet) fave) (aluble} % S Al HO Q he Ag st Na.S.0. Fet+ 8.0, ihlack) Cu,8,0,d—— + Na, [Cu(8,0,).] {soluble complex) {E) GROUP 17 ELEMENTS : THE HALOGEN FAMILY Fluorine, chlorine, bromine, iodine and astatine are members of Group 17. Electronic Configuration : ns?np> Atomic and lonic Radii The halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge . Physical Properties Fluorine and chlorine are gases, bromine is a liquid whereas iodine is a solid. Their melting and boiling points steadily increase with atomic number. The X-X bond disassociation enthalpies from chlorine onwards show the expected trend : Cl-Cl > Br-Br>F-F>I1I-I.  80 { Short Formula (Chemistry) force them to react to form the compounds were unsuccessful for quite a few years. In March 1962, Neil Bartlett, then at the University of British Columbia, observed the reaction of a noble gas. First , he prepared a ted compound which is formulated as O,* PtF-. He , then realised that the first ionisation enthalpy of molecular oxygen (1175 kJ mol -') was almost identical with that xenon (1170 kJ mol-'). He made efforts to prepare same. type of compound with Xe* PtF,- by mixing Pt F, and Xenon. After this discovery, a number of xenon compounds mainly with most electronegative elements like fluorine and oxygen, have been synthesised 3 If Helium is compressed and liquified it forms He()) liquid at 4.2 K. This liquid is a normal liquid like any other liquid. But if it is further cooled then He(11) is obtained at 2.2 K, which is known as super fluid, because itis a liquid with properties of gases. It climbs through the walls of the container & comes out. It has very high thermal conductivity & very low viscosity. CLATHERATE COMPOUNDS : During the formation of ice Xe atoms will be trapped in the cavities (or cages) formed by the water molecules in the crystal structure of ice. Compounds thus obtained are called clatherate compounds. Clathrate provides a convenient means of storing radioactive isotopes of Kr and Xe produced in nuclear reactors. 1:5 Ket F2 a 2/3 Ba 2 [xor J [S8F] &| 2 Xe + HF Sy, 8 |B Ay Ts 0+ Xe + HE Ne HO KF + Xe +i XetF +0.+H.0 Ap Xe + XeO, +HF +6, 2 hy 0,6 + Xe hi Xe + HE er Xe + F [REI DXF Ta A, Sir exer « SOK ES MeO HF ng oti ee TT FSP XeO,F AS + Xe, eR DIRE aes é XeF, +0. F, icsy er j 4 XeF,+ Cs,.XeF] d-BLOCK ELEMENTS & THEIR COMPOUNDS The general electronic configuration of d-block elements is (n—1) d'- ns°? . where n is the outer most shell. General trends in the chemistry of transition elements. Metallic character : Nearly all the transition elements display typical metallic properties such as high tensile strength, ductility, malleability, high thermal and electrical conductivity and metallic lustre. With the exceptions of Zn,Cd, Hg and Mn, they have one or more typical metallic structures at normal temperatures. The transition elements (with the exception of Zn, Cd and Hg) are very much hard and have low volatility. Melting and boiling points : The melting and boiling points of the transition series elements are gernerally very high.  Short Formula (Chemistry) Density : The atomic volumes of the transition elements are low compared with the elements of group 1 and 2. Thisis because the increased nuclear charge is poorly screened the transition metals are high. Oxidation states : Most of transition elements show variable oxidation states. Participation of inner (n - 1) d-electrons in addition to outer ns-electrons because, the energies of the ns and (n — 1) d-subshells are nearly same. Different oxidation states of first transition series. Outer Element electronic Oxidation states configuration Se 3d'4s? +3 Ti 307457 +2, 43, +4 v 3d°4s? 42,43, 44,45 Cr 3d°4s' +2, +3, (+4), (+5), +6 Mn 3024s? 42,43, +4, (+5), +6, +7 Fe 30°45" +2, +3, (+4), (+5), (+8) Co 3d’4s* +2, +3, (+4) Ni 3a°4s? +2, 43, +4 Cu 3d"“4s' +1, 42 Zn 3d"“4s* +2 Characteristics of Oxides and Some lons of V and Cr OS. | Oxides Behaviour len Name cf lon Colour of fon Hydroxide +2 vO basic vanadi-m dl) violet (vanadous) a V0, basic vanadivm dll} green y “4 vo, amphoteric vor blue V0} brown 45 V0; amphoteric vor dioxovanadium (V]—_yellow vol orthovanadate | colourless +2 cra 1 basic cr chromium (II) light blue CHO), (ehromaus) +2 Cro 1 basic cr chromiur {Il} light blue CHOH). {chromous) +3 Cr,0, 7 amphoteric cr chromium (Ili) violet CrOH), [ crremic CrOH) chromite green a5 cro, acidie Cry chromate yellaw: CrO,{OH), H,Cr,0, Cr,0;, dichromate orange  Short Formula (Chemistry) Standard electrode potentials : The value of ionisation enthalpies gives information regarding the thermodynamic stability of the transition metal compounds in different oxidation states. Smaller the ionisation enthalpy of the metal, the stable is its compound. Electrode potentials : In addition to ionisation enthalpy, the other factors such as enthalpy of sublimation, hydration enthalpy, ionisation enthalpy etc. determine the stability of a particular oxidation state in solution The overall energy change is AH= A, HP +IE+ A. WH The smaller the values of total energy change for a particular oxidation state in aqueous solution, greater will be the stability of that oxidation state. The electrode potentials are a measure of total energy change. Qualitative, the stability of the transition metal ions in different oxidation states can be determined on the basis of electrode potential data. The lower the electrode potential ie., more negative the standard reduction potential of the electrode, the more stable is the oxidation state of the transition metal in the aqueous solution. 3 3 Fe Co Mm Gu 27 9} Catoulated values Thermochemical data (kJ mol") for the first row Transition Elements and the Standard Electrode potentials for the Reduction of M'toM Element (M)|® © Akg (M) bAgbhy: FAH: 1g 3H! (MY) de BRAY di 459 B61 A310 21885 21.63. v 19 648, 1370 1893 alle Sh 398 653. 1599 1929 20,99 Mo 279 718: 4510 1862 “1:18 Fe a6 FOZ 1560 A998 “o4e Co. 427 ter 1640 “2073 “0.26 Ni 431 36. 1750 2127 -0.25, Cu 339 J45, A960 “2127 034 Zn 139 908. 1730 22059 C76, Formation of Coloured lons : Most of the compounds of transition metals are coloured in the solid form or solution form. The colour of the compounds of transition metals may be attributed to the presence of incomplete (n— 1) d-subshell. Re Bana daa de a qd, Ste dap de d-d irarsition gesde cme: « ay we New dy de the The excess of other colours constituting white light are transmitted and the compound appears coloured. The observed colour of a substance is always complementary colour of the colour which is absorbed by the substance. Magnetic Properties : {i substances. Paramagnetic substances : The substances which are attracted by magnetic field are called paramagnetic "ssf 85 | Short Formula (Chemistry) Mn.©, {explosive all} KMin®, green} if KMnO, is in exce: if acid is in excess fy he, wo en Kin, + MnO, +0, Mn’ + Q, a oa MnQ, + ©, {if placed in sunl ght) Mn+00,<———-——-| ~S Fe” Ma’ + Fe" Mn+ SO? = Ma” Mi?'+ Mn, Min, 4 » HS Mr*+ § I in alkaline or neutral medium loy4vnO, — S+MnO, 80,7+ MnO, Mn, 5. Potassium dichromate (K,CrO.) [Fetr0. A |NA,CO, + air Na,Cr, is}+ F045) + CO,T +H,0 NH,Ci {NH,).Cr,0, 4 Nef + Cr,0,+ HO. {green} cPigrean) + I, Gr'igreen} +S Cr'(green) + Fe%—f2 > “6H,O + KHSO, + Crd, (bringh orange/red) H,80,+80.4H.0 CP '(qreen) + S07 <7 chrome alum Kel (s) conc. H,80. Nagi GH,GOOH C0, C1, BS» Na Cro, Se PcrO_f (yellow {deep red vapours} {CH,COO},Pa  Short Formula (Chemistry) QUALITATIVE ANALYSIS Charcoal Cavity Test: Observation Inference Incrustation or Residue Metallic bead Yellow when hot, white when cold None Zn** Brown when hot, yellow when cold Grey bead which Pb? marks the paper No characteristic residue Red beads or scales Cut White residue which glows on heating None Ba**,Ca?*, Mg** Black None Nothing definite-generally coloured salt Cobalt Nitrate Test : S.No. Metal Colour of the mass 1 Zinc Green 2. Aluminium Blue 3. Magnesium Pink 4. Tin Bluish - green Flame test : 2585 8 3 OSBdroPPanie == 9289 8h ee es ee OBR gee ee Ee Re HESRERPRENTRREER EERE R hee RRR Crimson Red / Carmine Red Lithium Golden yellow Sodium Violet/Lilac Potassium Brick red Calcium Crimson Strontium Apple Green/Yellowish Green Barium Green with a Blue centre/Greenish Blue Copper Borax Bead test : Colour in oxidising flame Colour in reducing flame Metal When Hot When Cold When Hot When Cold Copper Green Blue Colourless Brown red Iren Brown yellow Pale yellow/Yellow Bottle green Bottle green Chromium Yellow Green Green Green Cobalt Blue Blue Blue Blue Manganese ViolevAmethyst Red/Amethyst Grey/Colourless Grey/Colourless Nickel Violet Brown/Reddish brown Grey Grey "37 87 | Short Formula (Chemistry) Analysis of ANIONS (Acidic Radicals) : (a) DILUTE SULPHURIC ACID/DILUTE HYDROCHLORIC ACID GROUP : 1. CARBONATE ION (CO,7) : Dilute H,SO, test : A colourless odourless gas is evolved with brisk effervescence. CaCO, +H,SO, —> CaSO, +H,0+C0,T Lime water/Baryta water (Ba(OH).) test : CO, + Ca(OH), —> CaCO, . (milky) +H,0 CaCO, +CO,+H,O ——> Ca(HCO,), (soluble) A > CaCO, 1+H,0+CO, 2. SULPHITE ION (SO,”): 2° Dilute H,SOQ, test : CaSO, + H,SO, —— CaSO, +H,0 + SO, T ; SO, has suffocating odour of burning sulphur. Acidified potassium dichromate test : The filter paper dipped in acidified K,Cr,O, turns green. Cr,0,2- + 2H* +380, ——+ 2Cr* (green) +380, +H,0. Barium chloride/Strontium chloride solution : SO, + Ba?/Sr? ——+ BaSO./SrSO, | (white). White precipitate dissolves in dilute HCI. BaSO, | + 2H* ——> Ba?*+S0,T+H,0. 3. SULPHIDE ION (S?-} : ep Dilute H,S0, test : Pungent smelling gas like that of rotten egg is obtained. $2 + 2Ht —> H,ST Lead acetate test : (CH;COO),Pb + H,S > PbS | (black) + 2CH,COOH. Sodium nitroprusside test : Purple coloration is obtained. S$? +[Fe(CN), (NO) > [Fe(CN),NOS]* (violet). Cadmium carbonate suspension/ Cadmium acetate solution : Na, + CdCO, > CdS | (Yellow) + Na,CO, . NITRITE ION (NO, }: Dilute H,SO, test : NO," + Ht —> HNO, ; (2HNO, ——> H,O +N,0,); 3HNO, ——> HNO, + 2NO+H,O ;2NO +0, ——> 2NO, T Starch iodide test : 2NO,- + 3I-+4CH,COOH —> I, + 2NO 1+ 4CH,COO-+ 2H,0 Starch + I,- —— Blue (starch iodine adsorption complex) 5. ACETATE ION (CH,COO } Dilute H,SQ, test: (CH,COO),Ca+H,SO, —— 2CH,COOH (vinegar like smell) + CaSO, Neutral ferric chloride test : 6CH,COO- + 3Fee + 2H,O0 ——-+ [Fe,(OH).(CH,COO),]* (deep red/ blood red colouration) + 2H* [Fe(OH)(CH,COO},)* + 4H,0 —®2_, 3Fe(OH),CH,COO J (brownish red) + 3CH,COOH + Ht (b) CONC . H,SO, GROUP : 1. CHLORIDE ION (CI): e road Concentrated H,SO, test : Cl +H,SO, ——> HCl (colourless pungent smelling gas) + HSO,- NH,OH + HCl ——> NH,CIT (white fumes) + H,O Silver nitrate test : Cl +Agt —— AgCll (white) White precipitate is soluble in aqueous ammonia and precipitate reappears with HNO. AgCl + 2NH,OH ——> [Ag(NH,),ICI (Soluble) + 2H,0 ; [Ag(NH,)]Cl + 2H* ——> AgCl1+2NH,*. Chromyl chloride test : "so 4 90 | Short Formula (Chemistry) IIA Group (Hg”, Pb®, Bi®*, Cu, Cd) P group filtrate + dil, HCI + HS (g) Previpitate, Filtrats, meve for Il" group i | v Black precipitate (Hy8, PbS, CuS ar Bi,8,) Yellow precipitate (CdS), Dissolve in 69% HNO, and then add KON 1 igs re Ede HN! " Vv Dissolve precipitate in 50% HNO, ane filter Colours solution, [Ca(CNY P Pass H,Sig) \ ‘ Fil:rate Insoluble (Hg8} (Nitrates of Pb, Bi and Cu} Yellow precipitate (CdS). Dissolve in aquaregia + Sn, : Black orenipitate (Hg) Colourless solution of (Pb(NO,}, of BI{NO,),).Add dilate Blue colour solution {Cu(NO,),}, H.SO, and litle C.H.OH (it further decreases the Add ammonia solution solubility of PbSO.,) and filter | Intense blue colaur solution 1 [Cu(NH,),J(NO,),. Add KCN solution in excess, Fillrate (Bi(SO.\). White precipitete (Pb80,) ‘Add NH,QH in excess. Oe oa eons Colourless solution of [CufCN),]™ nel ¢ White precipitate of B(OH), Pass H.S (9) parts {Pb(CH,COO},) Dissolve in dilute HC! and divide itinte -wo parts, ‘ No black precipitate of Cus. Ipart + XI Ti parl + K,CrO, ¢ - - part + NaQH + Na.Sno, Il part + excess of water Yellow precipitate Yellow precipitate L (Pbr,) (PbCrO,). White turbidity (Bi0'Cr). Black precipitate {Bi}. IIB Group (As*, Sb*, Sn*, Sn*} Soluble in (NH,),S, as (NH,), AsS,, (NH,), SbS,, (NH,),Sn8,. + Add dilute HCI and filter. | | Insoluble (As,S, + some S). Soluble (SbCl, & SnCl,). Dissolve in conc. HNO, Divide into two parts. and divide into two parts. t r.OOO”™~CSS Add excess of water. | part + Ammonium molybdate IW part + NH,CL+ | red heal anal oa Fool Yell J NH.OH 7 Ns80. White precipitate . ellow precipitate {SbOCl). : ipils {{NH,},A80, .12 MeO, ) White precipitate J Black or grey precipiele (Ha) {(MgiNF,) AsO,) Add tartaric acid. 4 Precipitate is soluble. oh i Il? Group (AM, Cr°s, Fe**} I Group _ Boiloff Filtrate, > Short Formula (Chemistry) H,S + then add conc. HNO, (1-2) drops + NH,Cl (solid) + NH,OH —— Preripitate. Fi If gelatinous white prec pitats (ANOH),) { Dissoive in dilute HCI{AICI). | Add sodiar Fydroxice solution. | While precipitate dissolves in excess of precivitant (NaAlo.). Add NHGL and heat Gelatinous while of AKOH),. ite, move for (V" group. J { if reddish brown If green precipitate (-o(OH),), prosisitate (CHOH),) L | Dissoive in diiute Fuse the precipitate with HCI (FeCl) jusion mixture & divide the (KNO, + Na,CO,) solution into two parts. & extract with water (Na,CrQ,), Divide solution inte two part | P pari + CH,COGH + (CH,COO}.Pb I part + a” K,[Fe(CN}. KSCN. Prussian biue (aqyy Blood red (ag) precipitate of of Fe(SCN}, ‘e/FeiCN},),. part + WW" part +Bach. Yellow precipitate of PhCrO, Yellow orecipitate of Bara, cipitate  Short Formula (Chemistry) Iv!" GROUP (Zn*, Mn?*, Ni%*, Co?*} Ill Group filtrate + NH,OH (excess) & NH,CI, then pass HS Precipitate Filtrate, [ZnS {wnite), Mn (light pink), NiS (black) or CoS (black)] move for ¥" group. Acd dilute HCI and shake and filter. —— Black precipitate (CoS or NiS). (nde ac ) Dissolve in aquaregia and ¥ evacorate to dryness. Boil (to remove any H.S dissolved). Add BrJ/H,O + NaOH and filter. Blue residue Yellow residue Black precipitate Filtrate (CoCl,) (NCL) (Mno,}. (Na,Zno,). J J J ~ Turns pink in water. Tums green ir water. Dissolve in cone. Divide the solution into HNO, and then add two parts. Add diemethylglyoxime PbO, and heat. in ammonical solution. v L Violet-red Lpart + NH,OH + ipart + Red/Resy red precipitate (purple colour) CH,COOH + KNO,. NH,CNS(solid) 4(C,H.N,O,}, Ni} solution (HMnO,), L + amyl alcohol. Yellow precipitate 4 KK, [Co(No.,]}. Blue colour in Pass H.S through Uoart+ alcahol layer of Lpart. CH.COOH + (NH,), [Co (SCN), t Kre(CNj, White precipitate L (Zn8). White precipitate of variable composition. Add excess of K,'Fe(CN),] 4 Precipitate of composition 2n, ky, [Fe (GN), ¥" Group (Ba*, Sr*, Ca) : IV Group filtrate ——> Boil off H,S then add (NH,),CO, (aq), NH,OH & NH,CI(s) ! White precipitate Filtrate, (BaCo,, SrCO, or CaCO,). move for VI group. Dissolve in CH,COOH and divide into three parts and test in the sequence given below. part + Kem, I Part + (NH) SO, I part + nH 1),C,0, Yellow precipitate White precipitate White precipitate (BaCro, insoluble in CH,COOH). (Srso,). (Cac, O,). vith GROUP: MAGNESIUM ION (Mg2*) : Mg?* + NH, +HPO,? ——> Mg(NH,)PO, 4 (white) 5 Mg?" +6 CO,2 +'7 HO —> 2MgCO,, Mg(OH),. 5 H,0 | + 2HCO,- Titan Yelllow (a water soluble yellow dyestuff) : It is adsorbed by Mg(OH), producing a deep red colour or precipitate.  95 { Compendium (Chemistry) cl Cl ,OMgX e 7 H,0' Se=0+RMgX —3 SOC "5 Rc 0.Et Eto E10” “R 8. Alkyl Cyanides : RMgX + (CN), —> RCN + Mg(CN)x 9. Primary amines : H,0* RMgX + RONH, ——» RNHMgX + RH + Mg(OCH,)CI— > RNH, (40-90%) RMgX + CINH, ——> RNH, + Mgxcl teoseeaneeeossesnzenss AO GAIKYVIG cesses snenaunes sees ae Electrophilic addition reactions : Mechanism Step 1: Attack of the electrophile on m bond forms a carbocation. oN y Lor vi, lg 2 + -c- ™ [on the more substituted carbon Step 2: Attack by a nucleophile gives the product of addition. loge Lol -c-co” + Nu; —> -C-C- 1 i | E E Nu va: NA A” Lo e.g. (a) Addition of water Ze = q + HO H® “ore H OH (Markounikov orientation} (b) Addition of hydrogen halides (where HX = HCI, HBr, Hl) HX 7 = | R-c=c-R' #45 R-cH-cx-e —P4se-c-c-r (Markovnikov addition) 4 b ARR UR MARA a ge ROHR RNR RAN ceccnganecessanasnes HARMONIC COMPOUNAS «oo cc cascascesaunans Electrophilic aromatic substitution : {a) Bromination of Benzene : Bromination follows the general mechanism for electrophilic aromatic sub- Stitution. Bromine itself is not sufficiently electrophilic to react with benzene, but a strong Lewis acid such as. FeBr, catalyzes the reaction. Step 1 : Formation of a stronger electrophile. en r. & 3 Br-Br: + FeBr, ==» |:Br—Br Fees, Step 2 : Electrophilic attack and formation of the sigma complex. H. i H @ 9 4 a 5 H H H 8 o ‘Br Br FeBr, H Br ir 3 H H - oe - — H + FeBr, H H H H i yu H H H sigma complex "O58 4 96 1 Compendium (Chemistry) Step :2 Loss of a proton gives the products. H FeBr, 4 oe Br x _—_ + HBr + FeBr. H ory i bromobenzene Ts reacts / Ts oy ny ress, oN + \ ener) errepr, \, intemediate \ products [0.3 keatne! or Mee ennennnnnennncnnnnnnnnnn reaction cvardinate——-—-» (b) Nitration H6 - No, ees Hg NO, #82. H,0° + HS0°+°NO, _Electrophitic (hone? Os : > nitronium ian attack” ° NO, (x-complex) ont complex) I ws Coe > ho OU The electrophilic reagent, SO,, attacks the benzene ring to form the intermediate carbocation. 2H,S0, == SO, +H,0® + Hs0° @ ah ie H 4g SO.H > CK so® $0° Tanai? (d) Friedel Craft reaction : (c) Sulphonation : Alkylation mechanism oN Sb 3 0 R-Cl + AlCl, —> R TEI—AICl, —> R® + aici’ carbocation R @ © +s > (Ch a-complex  97 1 Compendium (Chemistry) R R (iii) Ch + AICl, —> Or HCI + AICI, Acylation mechanism Acylation of benzene may be brought about with acid chlorides or anhydrides in presence of Lewis acids. Step 1: Formation of an acylium ion. oO ho ll @ © @ . @ R-C-Cl + AIL —s R-C-C-Alc, == Big, + IR-C=Q ee R-C=0] v C a acylchloride complex acylium ion Step 2: electrophilic attack. 9 Il @ CN R H sigma complex Step 3 : Loss of a protan. Complexation of the product. o ll aeylbenzene sigma complex ° Il ea cate fIAICL ‘3. CH,CH, (2) HO acetyl chloride ethylbenzene CH.CH. protayl-Aestoprannae i70—S8D%) Note : Friedal - Crafts acylations are generally free from rearrangements and multiple substitution. They do not go on strongly deactivated rings. cOcH, eg. Oo) + cHcocl “45 or Hel Acetophenone Chemical Reactions of Benzene : Conc. HNO; H,80, + H,0 NO, Conc. H,S0, +803 SO, aid + HO