Periodic Trends in Properties of Elements

Chemistry CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES PERIODIC TRENDS IN PROPERTIES OF ELEMENTS ATOMIC PROPERTIES OR PERIODIC PROPERTIES OF ELEMENTS The physical properties of the elements can be divided into two categories : (i) Properties of individual atoms: Properties such as atomic and ionic radiate ionisation enthalpy, electron gain enthalpy, valency, screening effect, effective nuclear charge and electronegativity are the properties of individual atoms and directly dependent on their electronic configurations. (ii) Properties of group of atoms: Properties such as melting point, boiling point, heat of fusion, heat of vaporization, density, atomic volume, etc., are the properties of group of atoms together and are indirectly related to their electronic configurations. All these properties which are directly or indirectly dependent on the electronic configuration of the elements are called atomic properties. Since electronic configurations of the elements are periodic function of atomic numbers of the elements, therefore, these atomic properties are also a periodic function of atomic numbers of the elements. On account of this, atomic properties are also called periodic properties. Valency: It is defined as the combining capacity of the elements. The word valency is derived from an Italian word "Valentia" which means combining capacity. Old concept: Given by: Frankland Valency with respect to Hydrogen: Valency of H = 1 It is defined as the number of hydrogen atoms attached with a particular element. Valency IA IIA IIIA IVA VA VIA VIIA NaH MgH2 AlH3 SiH4 PH3 H2S H–Cl 3 2 1 2 3 4 1 Note: Valency w.r.t. H across the period increases up to 4 and then again decreases to 1. 1 Chemistry Valency with respect to oxygen: Valency of 'O' = 2 It is defined as twice the number of oxygen atoms attached with a particular atom. Valency IA IIA IIIA IVA VA VIA VIIA Na2O MgO Al2O3 SiO2 P2O5 SO3 Cl2O7 1 2 3 4 5 6 7 Note : Valency with respect to oxygen increases from 1 to 7 across the period. Valency w.r.t. 'O' is equal to the group number. New concept: This concept is based on the electronic configuration. According to this concept valency for IA to IVA group elements is equal to number of valence shell e– and from VA to zero group, it is – [8– (number of valence e–)]. Valency = No. of valence e– IA IIA IIIA ns1 ns2 Valency = (8– no. of valence e–) VA VIA VII 0 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2np6 Valence shell e–1 2 3 4 5 6 7 8 Valency 2 3 4 3 2 1 0 1 IVA (8 – 5) = 3 (8 – 8) = 0 Note : All the elements of a group have same valencies because they have same number of valence shell electrons. Screening effect () and effective nuclear charge (Zeff) (a) Valence shell e– suffer force of attraction due to nucleus and force of repulsion due to inner shell electrons. (b) The decrease in force of attraction on valence e– due to inner shell e– is called screening effect or shielding effect.(i.e. total repulsive force is called shielding effect.) 2 Chemistry (c) Due to screening effect. valence shell e– experiences less force of attraction exerted by nucleus.(i.e. total attraction force experienced by valence e– is called Zeff.) (d) There is a reduction in nuclear charge due to screening effect. Reduced nuclear charge is called effective nuclear charge. (e) If nuclear charge = Z, then effective nuclear charge = Z – (Where (Sigma)= Screening constant) So, Zeff = (Z – ) th Z.e (n-1)th shell (n-2) shell Force of attraction nth shell (valence shell) Ele c rep tronic ulsi on Test electron io n ic repuls Electron Slater's rule to know screening constant () (a) Screening effect (S.E.) of one e– of the 1s is 0.30. Ex. In He (1s2) Screening effect of one 1s e–. where = 0.30 Zeff = Z – = 2 – 0.30 = 1.7 (b) Screening effect of ns and np (Outermost orbit) electron is 0.35 (c) Screening effect of (n – 1) penultimate orbit s, p, d electrons is 0.85 (d) Screening effect of (n – 2) and below all the e– present in s, p, d, f is 1.0 3 Chemistry (Effective Nuclear charge of elements of second period) Element Electronic Z Configuration of ns & np (n–1) Total Effective electron orbital Screening nuclear Constant charge (a + b) Z*=Z –  3Li 1s2 2s1 3 – 0.85× 2=1.70 1.70 1.30 4Be 1s2, 2s2 4 1× 0.35=0.35 0.85× 2=1.70 2.05 1.95 5B 1s2,2s2,2p1 5 2× 0.35=0.70 0.85× 2=1.70 2.40 2.60 6C 1s2,2s2,2p1 6 3× 0.35=1.05 0.85× 2=1.70 2.75 3.25 7N 1s2,2s2,2p3 7 4× 0.35=1.40 0.85× 2=1.70 3.10 3.90 8O 1s2,2s2,2p4 8 5× 0.35=1.75 0.85× 2=1.70 3.45 4.55 9F 1s2,2s2,2p5 9 6× 0.35=2.10 0.85× 2=1.70 3.80 5.20 Periodic variation (a) From left to right in a period Zeff increases (I) That is why in a period Zeff increases by 0.65 and hence atomic size decreases considerably. (II) In transition series Z increase by + 1 but screening effect increases by 0.85 So Zeff is 0.15 (1– 0.85 = 0.15) [Because e– enters in (n – 1) orbit which has value of = 0.85] In transition series Zeff increases very less amount, by 0.15 from left to right and hence atomic size remains almost constant. Element Sc Zeff 3.00 3.15 (b) Ti V Cr Mn Fe Co Ni Cu Zn 3.30 3.45 3.60 3.75 3.90 4.05 3.70 4.35 From top to bottom in a group Zeff remain constant Element Li Na Zeff 1.30 2.20 K Rb Cs Fr 2.20 2.20 2.20 2.20 4 Chemistry ATOMIC RADIUS The average distance of valence shell e– from nucleus is called atomic radius. It is very difficult to measure the atomic radius because – (I) The isolation of single atom is very difficult. (II) There is no well-defined boundary for the atom. (The probability of finding the e– is 0 only at infinity). So, the more accurate definition of atomic radius is –  Half the inter-nuclear distance(d) between two atoms in a homoatomic molecule is known as atomic radius.  This inter-nuclear distance is also known as bond length. Inter-nuclear distance depends upon the type of bond by which two atoms combine. Based on the chemical bonds, atomic radius is divided into four categories – 1. Covalent radius 2. Ionic radius 3. Metallic radius 4. Vander waal radius 1. Covalent radius One half of the distance between the nuclei (internuclear distance) of two covalently bonded atoms in homoatomic molecule is called the covalent radius of that atom. The covalent bond must be single covalent bond. The covalent radius (rA) of atom A in a molecule A2 may be given as: rA = d AA 2 i.e. the distance between nuclei of two single covalently bonded atoms in a homoatomic molecule is equal to the sum of covalent radii of both the atoms dA–A = rA + rA In a heteroditopic molecule AB where the electronegativity of atoms A and B are different, the experimental values of internuclear distance d A-B is less than the theoretical values (rA + rB). According to Schomaker and Stevenson – DA-B = rA + rB – 0.09 x 5 Chemistry Where x is the difference of electronegativities of the atoms A and B. According to Pauling – If the electronegativities of the two atoms A and B are x A and xB respectively then DA-B = rA + rB – (C1xA – C2xB) C1 and C2 are the Stevenson's coefficients for atoms A and B respectively. 2. Metallic Radius Metal atoms are assumed to be closely packed spheres in the metallic crystal. These metal atom spheres are considered to touch one another in the crystal. One half of the internuclear distance between the two closest metal atoms in the metallic crystal is called metallic radius. Metallic > Covalent radius For example – Metallic radius and covalent radius of potassium are 2.3 Å and 2.03 Å respectively. 3. Van Der Wall's Radius or Collision radius The molecules of nonmetal atoms are generally gases. On cooling, the gaseous state changes to solid state. In the solid state, the nonmetallic elements usually exist as aggregations of molecules are held together by van der wall forces. One half of the distance between the nuclei of two adjacent atoms belonging to two neighboring molecules of a compound in the solid state is called van der walls radius. It may also be defined as half of the inter nuclear distance of two non-bonded neighboring atoms of two adjacent molecules. 2 × vander waal's radius d Van der Wall's radius = 1 × Internuclear distance between two successive nuclei of two 2 covalent molecules (d) 6 Chemistry Van der wall's radius > Metallic radius > Covalent radius The van der walls radius and covalent radius of chlorine atom are 1.80 Å and 0.99 Å respectively 4. Ionic Radius A neutral atom changes to a cation by the loss of one or more electrons and to an anion by the gain of one or more electrons. The magnitude of charge on cation and anion is equal to the number of electrons lost or gained respectively. The ionic radii of the ions present in an ionic crystal may be calculated from the inter molecular distance between the two ions. (a) Radius of Cation Radius of cation is smaller than that of corresponding atom. Reasons (I) During the formation of cation either one shell is removed or (II) After removing an electron effective nuclear charge increase. (b) Radius of an Anion Radius of an anion is invariably bigger than that of the corresponding atom. Reasons (I) The effective nuclear charge decrease in the formation of anion. Thus the electrostatic force of attraction between the nucleus and the outer electrons decreases as the size of the anion increases. (II) Interelectronic repulsion increases. Factors affecting atomic radius are (a) Atomic radius  1 Effective nuclear charge (Zeff) Li > Be > B > C > N > O > F (b)Atomic radius number of shells Li < Na < K < Rb < Cs (c) Atomic radius  Screening effect (d) Atomic size Magnitude of –ve charge < O– < O–2 (e) Atomic radius  Mn > Mn+2 > Mn+3 > Mn+4 7 Chemistry (f) Atomic radius  1 N — N > —N N— >N Bond order N Periodic variation of atomic radius (a) Across a period: It decreases from left to right in a period as nuclear charge increases Ex. (b) Li > Be > B > C > N > O > F < Ne In a group: It increases from top to bottom in a group as number of shell increases Ex. Li < Na < K < Rb < Cs Exceptions (a) Transition elements Sc Ti V Cr Mn Fe Co Ni Zeff  Screening effect Zeff > Screening effect (b) Lanthanide Contraction (I) Outermost electronic configuration of inner transition elements is (n –2) f1 – 14, (n–1)s2p6d0–1, ns2 (n = 6 or 7) (II) e– enters in (n – 2) f orbitals (III) Mutual screening effect of e– is very less, because of complicated structure of f-orbital (IV) Nuclear charge increases by one (+1) in lanthanides and actinides so atomic size of these elements slightly decreases. It is known as lanthanide contraction. Its effect is also observed in 5d transition series. Here Nuclear charge > Screening effect. (V) In Ist, 2nd and 3rd. transition series, Radii– 3d < 4d 5d (except IIIrd B) IIIB IVB size Sc Ti size increases increases Y Zr La Hf } Equal due to lanthanide contraction 8 Chemistry (c) Transition contraction IIIA B p > d > f (II) As subshell is more closer to nucleus so more energy will be required to remove e – in comparison p,d & f. to Ex. Be B 1s2, 2s2 Ionisation Potential Be 1s2, 2s2 2p1 > B After losing one e–, B attains electronic configuration of Be, so ionisation potential of B is more than Be. IIndonisation Potential of B > Be (e) Stability of half-filled and fully filled orbitals (I) Half-filled p3,d5, f7 or fully filled s2, p6, d10, f14 are more stable than others so it requires more energy. Ex. N O 1s2, 2s2 2p3 1s2, 2s2 2p4 Ionisation Potential order is O < N Because of half-filled p-orbitals in N, its ionisation energy (stability) is higher than O. ionisation potential order Na < Al < Mg (II) Because s-orbital in Mg is completely filled and its penetration power is also higher than p-orbital (Al). ionisation potential order Mg+ (2,8,1) < Al+ (2,8,2) < Na+ (2,8) Periodic variation of ionisation energy (a) Variation in period among the representative elements: Ionisation energies generally increases along the period because in moving left to right in a period the effective nuclear charge per outermost electron increases while the corresponding principal quantum number remain same. 12 Chemistry (b)Variation in a group among the representative elements: The ionisation energy generally decreases in moving from top to bottom because the size increases due to the increase of the principal quantum number. On the other hand, the effective nuclear charge Zeff for the outermost electron remains almost the same along the group. Application of ionisation potential (a) Metallic and non-metallic character Ionisation Potential Low (Na, K, Rb etc.) Metallic nonmetallic  Ionisation Potential High (F, Cl, Br etc.) Ionisation Potential  (b) Metallic property Reducing character Reducing character  (I) 1 1 Ionisation Potential IA group has minimum ionisation potential so they are strong reducing agents in gaseous state ( Li < Na < K < Rb < Cs) (II) IA group – In Aqueous state reducing character Li > K ~ Rb > Cs > Na As the degree of hydration is more in Li due to high charge density. (III) VIIA group has maximum ionisation potential so they are strong oxidising agents (F > Cl > Br > I) (c) Stability of oxidation states (I) If the difference between two successive ionisation potential > 16eV then lower oxidation state is stable. Ex. Na  Na+ Ist ionisation potential } 42.7 eV Na+ Na+2 IInd ionisation potential Difference of ionisation potential > 16 eV So Na+ is more stable. (II) If the difference between two successive ionisation potential < 11 then higher oxidation state is stable. Ex. Mg  Mg+ Mg+  Mg+2 Ist ionisation potential } 7.4 eV IInd ionisation potential 13 Chemistry Difference ofionisation potential < 11 eV So Mg+2 is more stable. Al  Al+ } 12.8 eV So Al+ is more stable Al+  Al+2 } 6.0 eV So Al+3 is more stable Al+2  Al+3 Al+ is stable only in gaseous state { Al+3 is stable in liquid and solid state. ELECTRON AFFINITY (a) The amount of energy released when an electron is added to the outermost shell of one mole of an isolated gaseous atom in its lower energy state. (b) The positive value of the electron affinity indicates that the process, i.e. X (g) + e– , is exothermic and the negative value indicates the process to be an endothermic one. Thus, the convention accepted in defining the electron affinity apparently contradicts the established convention in the thermodynamics. Thus, F(g) + e–  , H = – 328 kJ mol–1, and EA = 328 kJ mol–1 and N(g) + e–  , H = +31 kJ mol–1, and EA = –31 kJ mol–1 (c) Electron affinity just defined is actually first electron affinity since it corresponds to the addition of one electron only. In the process of adding further electron, the second electron will be added to gaseous anion against the electrostatic repulsion between the electron being added and the gaseous anion. Energy instead of being released is supplied for the addition of an electron to an anion. A(g) + e–  A(g)– + E1 ; A(g)– + e– + Energy supplied  A(g)2– 14 Chemistry Factors affecting electron affinity (A) Atomic size Electron Affinity  Electron Affinity  (B) Screening effect (C) Effective nuclear charge (Zeff) (d) 1 Atomic size 1 Screening effect Electron Affinity  Zeff Stability of completely filled or half-filled orbitals Electron affinity of filled or half filled orbital is very less or zero. Ex. (I) Elements (He, Ne, Ar , ....), Electron Affinity = Zero ( egH = + ve) (II) Elements (Be, Mg, Ca, ....), Electron Affinity = ~ Zero ( egH = + ve) (III) Elements (N, P, As, .... ), Electron Affinity = Very less Periodic variation of electron affinity (a) In a period: The effective nuclear charge increases and the size decreases with the increase of atomic number in a period. This is why, the electron affinity in general increases in a period. In fact, it reaches the climax for the Group VII A (i.e. halogens) elements. Electron affinities (kJ mol –1 ) IA IIA IIIA IVA VA VIA VIIA Noble gas Li Be B C N O F Ne 59.8 –60 27 122 –31 141 328 –90 Na Mg Al Si P S Cl Ar 53 –60 44 134 72 200 349 0 Ga Ge As Se Br 29 120 77 195 324 15 Chemistry (b) In a group: For the representative elements, in moving down in a group generally the electron affinity falls down with the increases of atomic number because the effective nuclear charge Zeff at periphery per electron remain almost same but the size gradually increases due to addition of new shell. ELECTRONEGATIVITY (EN) (a) The tendency of an atom to attract shared electrons towards itself is called electronegativity. (b) Electronegativity and Electron affinity both have tendency to attract electrons but electron affinity is for isolated atoms. Whereas electronegativity is for bonded atoms. (c) A polar covalent or ionic bond of A – B may be broken as (I) A – B  A– : + B+ (Electronegativity A > Electronegativity B) or (II) A – B  A+ + :B– (Electronegativity A < Electronegativity B) depending on their tendency to attract bonded electron. (d) There is no unit of electronegativity because it is a relative value (e) Electronegativity is property of a bonded atom not an isolated gaseous atom so no energy will be liberated or absorbed. Difference between electronegativity and Electron Affinity Electronegativity Electron Affinity Tendency of an atom in a molecule Energy released when an electron is added to to attract the bonded electrons neutral isolated gaseous atom Relative value of an atom Absolute value of an atom It regularly changed in a period or group It does not changes regularly It has no unit It is measured in eV/atom or KJ mol–1 or K Cal mole–1 Note: Small atoms are normally having more electronegativity than larger atoms. 16 Chemistry Pauling's Scale Pauling related the resonance energy (AB) of a molecule AB with the electronegativities of the atoms A and B. If xA and xB are the electronegativities of atoms A and B respectively then 0.208  AB = xA – xB if xA > xB or AB = 23.06 (xA – xB)2 AB = EAB(experimental) – EAB(theoretical) where EA–B is the energy of A–B bond. EAB (theoritical) = (EA  A  EB B ) In a purely covalent molecule, AB, the experimental and theoritical values of bond energy A–B are equal, So AB = 0 or 0 = 23.06 (xA – xB)2 or xA = xB In an ionic molecule AB, EAB(experimental) is more than EAB(theoritical) Pauling assumed the electronegativity value of fluorine 4 and calculated the electronegativity values of other elements from this value. Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7 Be 1.5 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Ra 0.9 B 2.0 Al 1.5 Ga 1.6 In 1.7 Tl 1.8 C 2.5 Si 1.8 Ge 1.8 Sn 1.8 Pb 1.8 N 3.0 P 2.1 As 2.0 Sb 1.9 Bi 1.9 O 3.5 S 2.5 Se 2.4 Te 2.1 Po 2.0 H 2.1 F 4.0 Cl 3.0 Br 2.8 I 2.5 In Pauling's scale, elements having almost same electronegativity areC = S = I = 2.5 N = Cl = 3.0 P = H = 2.1 Cs= Fr = 0.7 Be = Al = 1.5 17 Chemistry Mullikan’s Electronegativity According to Mullikan, the electronegativity of an element is the average value of its ionisation potential and electron affinity. or electro-negativity = Electron affinity + Ionisation potential 2 When both are expressed in electron volt Allred Rochon’s Electronegativity: The electronegativity of an element is the electrostatic force of attraction between the electron present on the circumference of outermost shell of this atoms and the atomic nucleus. If the distance between the circumference of outermost shell & the nucleus is r and the effective nuclear charge Zeff then Electron-negativity Zeff e 2 0.359 Zeff  r2 r2 = + 0.744  Zeff = Z -  Z = The actual number of charge present in the nucleus i.e. number of protons and  = Shielding constant. Factors Affecting electronegativity (A) Atomic size (B) Effective nuclear charge (Zeff) Electronegativity  (C) 1 Electronegativity  Zeff Atomic size Hybridisation state of an atom (D) sp > sp2 > sp3 s character 50% Electronegativity 3.25 Oxidation state Mn+2 < Mn+4 < Mn+7 33% 2.75 25% 2.5 O–2 < O–1 < O < O+1 < O+2 Fe < Fe+2 < Fe+3 (I) s-orbital is nearer to nucleus so by increasing s-character in hybridization state, Electronegativity also increases. (II) Electronegativity does not depend on filled or half-filled orbitals, because it is a tendency to attract bonded electron, not to gain or loss electron from outside. 18 Chemistry Periodic Variation (a) In a period: In moving from left to right in a period, the size gradually contracts and the effective nuclear charge increases. This is why, the electronegativity gradually increases with the increase of atomic number in a period up to the group VIIA elements. (b) In a group: In the representative elements, in moving down in a group, the size increase while Zeff per electron at the periphery remains almost constant. This is why, the electronegativity generally falls in a group with the increase of atomic number. Application of electronegativity (a) Metallic and non-metallic nature Low electronegativity  Metals High electronegativity  Non-Metals Metallic character increases down the group but decreases along a period. (b) Bond length ΔEN  1 Bond length Here EN = difference in electronegativities of bonded atoms Note: HF has minimum bond length as expected because of much difference in the electronegativities of H and F. (c) Acidic strength of hydrides NH3 < H2O < HF In a particular period acidic strength of hydride is depends on electronegativity of M in M – H bond. Higher the electronegativity of M greater the acidic strength of hydride. (d) Nature of bonds (I) According to Hanny & Smith formula Ionic % = 16 (XA – XB) + 3.5 (XA – XB)2 Here XA = Electronegativity of A XB = Electronegativity of B 19 Chemistry If XA – XB > 2.1 Ionic % > 50% i.e. Ionic bond If XA – XB < 2.1 Ionic % < 50% i.e. covalent bond (II) According to Gallis XA – XB > 1.7 Ionic XA – XB < 1.7 Covalent If XA = XB ; then A – B will be non-polar. Ex. H—H, F—F If XA > XB and difference of electronegativities is small then A —— B bond will be polar covalent (e) Nature of hydroxides (I) As per Gallis, In AOH if electronegativity of A is more than 1.7 (Non metal) then it is acidic in nature. (II) If electronegativity of 'A' is less than 1.7 (metal) then AOH will be basic in nature Ex. XA Nature NaOH ClOH 0.9 3.0 Basic (III) If XA – X0 > X0 – XH Acidic then AO bond will be more polar and will break up as A —— OH  A+ + OH– It shows basic nature (IV) If XA – XO < XO – XH A — O ——H  H+ + AO– Ex. It shows Acidic nature In NaOH XO – XNa (2.6) > XO – XH (1.4) So hydroxide is basic In ClOH XO – XCl (0.5) < XO – XH (1.4) So hydroxide is acidic 20 Chemistry (f) Nature of oxides : Consider an oxide AO If XA – XO > 2.3 Basic oxide If XA – XO = 2.3 Amphoteric oxide If XA – XO < 2.3 Acidic oxide (I) Along a period acidic nature of oxide increases. (II) Down the group basic nature of oxide increases Na2O MgO XA – XO > 2.3 Basic Al2O3 SiO2 XA – XO = 2.3 Amphoteric P4O10 SO3 Cl2O7 XA – XO < 2.3 Acidic ie. when in periodic table the distance between the element and oxygen increases, basic character increases. NO2 > ZnO > K2O  acidic character decreases (g) Hydrolysis products of AX Where A = Other element and X = Halogen (I) If electronegativity of X > Electronegativity of A then on hydrolysis product will be HX. Ex. (BCl3), Electronegativity of Cl > Electronegativity of B PCl3 + 3HOH 3HCl + H3PO3 PCl3. Cl2 + H2O  2HCl + POCl3 POCl3 + 3HOH  3HCl + H3PO4 21 Chemistry (II) If electronegativity of X < electronegativity of A then on hydrolysis product will be HOX (hypohalous acid)   O Ex. Cl  Cl Here electronegativity of O > Cl Cl2O So On hydrolysis – (h) Nomenclature of inorganic compounds Prefix —— less electronegative element Suffix —— More electronegative element DENSITY (a) In a group volume of an atom increases along with atomic weight but atomic weight increases more than atomic volume, So density increases in a group. (b) Factors affecting density are (I) Inner shell configuration (II) Atomic weight (III) Packing capacity (IV) Chemical bond Periodic Variation (a) In period density first increases till maximum and then decreases. (s-block to d-block increases, d-block to p-block decreases) (b) In group from top to bottom density increases regularly Ex. In VIIA group – F and Cl are gases (Low density) Br is liquid (density 3.19 g/cm–3) I is solid (density 4.94 g/cm–3) From s-block to p-block packing capacity and strength of bond increases. 22 Chemistry BOILING POINTS AND MELTING POINTS (a) It is a property of aggregate of atoms and not of a single atom. i.e. why it is a molecular property. (b) Melting point of a solid depend on – (I) Structure of solid (II) Chemical bond (III) Bond energy (c) Structure of solid made up – (I) atomic solid (II) molecular solid (III) metallic solid. Order of decreasing melting point is : Atomic solid > Metallic solid > Molecular solid (d) Chemical bond among particles in solid. Order of bond strength is – Covalent bond > Metallic bond > Vander waal bond (e) Bond energy – Covalent solids like diamond, SiO2 etc have only covalent bonds between atoms so their bond energy is higher than molecular or metallic solids. Periodic variation In period from left to right boiling point and melting point first increases then decreases. Alkali metals – Crystal structure BCC (low boiling point & melting point) Transition metals – ............"..... FCC (High boiling point & melting point)  Inert gases ............. Lowest boiling point & melting point (Vander waal force)  Transition elements .......... Highest melting point  Metals Non metals  W(Tungston) maximum melting point (3410°C)  Hg(Mercury) lowest melting point (  38°C)   Carbon (In the form of diamond) Highest melting point (3727°C)  Helium Lowest melting point ( - 270°C)  23 Chemistry In Group (a) In s-block elements boiling point & melting point decreases down the group. Li, Na (solid) Cs, Fr (liquid) It is due to weak cohesive energy of metallic bond in IA. (b) In d-block elements boiling point & melting point increases down the group (due to lanthanide contraction, Zeff increases and hence bond energy increases) (c) In p-block elements (I) From IIIA – IVA group boiling point & melting point decreases down the group and from VA to '0' group, boiling point & melting point increases down the group. (Atomic or molecular weight vander waal force) (II) boiling point and melting point of monoatomic molecules are lesser than diatomic molecules. '0' group < Halogens (III) Atomic solid non-metals like B, C and Si has higher boiling point and melting point due to strong covalent bond. (IV) Boiling point & melting point of molecular solids are less because of weaker vander waal force among molecules Ex. I2. Order of melting point Li > Na > K > Cs C > Si > Ge He < Ne < Ar < Rn Be > Ca > Sr > Ba > Mg C > B > Si Zn > Cd > Hg F2 < Cl2 < Br2 < I2 B > Al > Ga Cu > Au > Ag BASIC INORGANIC NOMENCLATURE Oxidation Number  It is an imaginary or apparent charge developed over atom of an element when it goes from its elemental free state to combined state in molecules.  It is calculated on basis of an arbitrary set of rules.  It is a relative charge in a particular bonded state.  In order to keep track of electron-shifts in chemical reactions involving formation of compounds, a more practical method of using oxidation number has been developed. 24 Chemistry  In this method, it is always assumed that there is a complete transfer of electron from a less electronegative atom to a more electronegative atom. Rules governing oxidation number The following rules are helpful in calculating oxidation number of the elements in their different compounds. It is to be remembered that the basis of these rule is the electronegativity of the element  Fluorine atom Fluorine is most electronegative atom (known). It always has oxidation number equal to –1 in all its compounds  Oxygen atom In general, and as well as in its oxides, oxygen atom has oxidation number equal to –2. In case of (I) peroxide (e.g. H2O2, , Na2O2 ) is –1, (II) super oxide (e.g. KO2) is –1/2 (III) ozonide (e.g. KO3) is –1/3 (IV) in OF2 is + 2 & in O2F2 is +1  Hydrogen atom In general, H atom has oxidation number equal to +1. But in metallic hydrides ( e.g. NaH, KH), it is –1.  Halogen atom In general, all halogen atoms (Cl, Br, I) have oxidation number equal to –1. But if halogen atom is attached with a more electronegative atom than halogen atom, then it will show positive oxidation numbers. 5 Ex. 5 K ClO3 HI O 3 , 7 , 5 HCIO 4 KBrO3 , 25 Chemistry  Metals (a) Alkali metal (Li , Na, K, Rb, .......) always have oxidation number +1 (b) Alkaline earth metal (Be , Mg , Ca .......) always have oxidation number +2. (c) Aluminium always has +3 oxidation number Note: Metal may have positive or zero oxidation number  Oxidation number of an element in free state or in allotropic forms is always zero 0 Ex. 0 0 0 O2 , S8 , P4 , O3  Sum of the oxidation numbers of atoms of all elements in a molecule is zero.  Sum of the oxidation numbers of atoms of all elements in an ion is equal to the charge on the ion  If the group number of an element in modern periodic table is n, then its oxidation number may vary from (n – 10) to (n – 18) (but it is mainly applicable for p-block elements) Ex. N- atom belongs to 15th group in the periodic table, therefore as per rule, its oxidation number may vary from –3 to +5 (N-3H3 , N-2O3 , N+4O2 , N-52 O5) -3 Or ( +2 N H 3 , NO +3 4 +5 N 2 O3 N O2 N 2 O5 , , , )  The maximum possible oxidation number of any element in a compound is never more than the number of electrons in valence shell. (But it is mainly applicable for p-block elements) 26 Chemistry List of common oxidation state of an element in Periodic Table 1 1 H +1 –1 18 2 He 3 Li +1 2 3 Be +2 13 5 B +3 –3 14 6 C +4 +2 –4 etc. 11 Na +1 12 Mg +2 13 Al +3 14 Si +4 –4 19 K +1 20 Ca +2 31 Ga +3 32 Ge +4 –4 33 As +5 +3 –3 34 Se +6 +4 –2 37 Rb +1 38 Sr +2 49 In +3 +1 50 Sn +4 +2 51 Sb +5 +3 –3 52 Te +6 +4 –2 55 Cs +1 56 Ba +2 81 Tl +3 +1 82 Pb +4 +2 83 Bi +5 +3 84 Po 3 21 Sc +2 +3 4 22 Ti +2 +3 +4 5 23 V +2 +3 +4 +5 6 24 Cr +2 +3 +4 +5 +6 7 25 Mn +2 +3 +4 +5 +6 +7 8 26 Fe +2 +3 +4 +5 +6 9 27 Co +2 +3 +4 +5 10 28 Ni +2 +3 +4 11 29 Cu +2 +3 12 30 Zn +2 15 16 7 8 N O +5 +2 +4 –1/2 +3 –1 +1 –2 –3 0 etc. 15 16 P S +5 +6 +3 +4 +1 +2 –3 –2 17 9 F –1 17 Cl +5 +7 +3 +1 –1 etc. 35 Br +5 +3 +1 –1 53 I +7 +5 +1 0 –1 85 At 10 Ne 18 Ar 0 36 Kr +4 +2 0 54 Xe +6 +4 +2 0 86 Rn 27 Chemistry Elements General Rule: The names of metals generally end in-ium or-um (examples are sodium, potassium, aluminum, and magnesium). The exceptions are metals that were used and named in ancient times, such as iron, copper, and gold. The names of nonmetals frequently end in-ine,-on, or -gen (such as iodine, argon, and oxygen.) Given the names of the constituent elements and common ions, most of the common inorganic compounds can be named using the rules presented below. Acids Acids are normally classified in two groups, hydracids and oxyacid’s Hydracids Hydracids are acids which contain hydrogen and a non-metal , but no oxygen. General Rule: The names of hydracids have the prefix hydro-(sometimes shortened to hydr-) and the suffix-ic attached to the stem based on the names of the constituent elements (other than hydrogen.) For example: HCl (made of hydrogen and chlorine) is hydrochloric acid; HBr (made of hydrogen and bromine) is hydrobromic acid; Hl (made of hydrogen and iodine) is hydroiodic acid; HCN (made of hydrogen, carbon, and nitrogen) is hydrocyanic acid; and H2S (made of hydrogen and sulfur) is hydro sulfuric acid. Oxoacids or Oxyacid’s The acids which contain hydrogen , oxygen and a metal or non-metal. In this case , more than one possibilities arise because of the presence of different number of oxygen atoms. An example of such an oxoacid series is as follows: HClO, HClO2, HClO3, HCIO4. All these contains same three elements but differ in the number of oxygen atoms present. General Rule - 1 If a class of acids contains only one member, its name is given the suffix-ic. For example, hydrogen , carbon and oxygen combine to form only one acid i.e. H 2CO3. It is called carbonic acid (carbonic acid.) 28 Chemistry General Rule - 2 If an acid series contains two acids, such as H 2SO4 and H2SO3, the acid containing more oxygen atoms is given the suffix -ic, while the acid with fewer oxygen atoms is given the suffix-ous. For example : H2SO4 is sulfuric acid , and H2SO3 is sulfurous acid Similarly, HNO3 is nitric acid and HNO2 is nitrous acid. In the case of an extensive acid series (such as HClO, HClO2, HClO3, HClO4), the acid with the one oxygen atoms lesser than -ous acid is given the prefix hypo- and the suffix -ous, and the acid with the one oxygen atom more than the -ic acid is given the prefix per and a suffix-ic. In the above example, HClO is hypochlorous acid HClO2 is chlorous acid, HClO3 is chloric acid, and HClO4 is perchloric acid. A very valid doubt may arise, “How to decide which one is -ic acid to begin with ? This is somewhat arbitrary and has no unique answer. It depends upon the name given to it by old chemists, when no systematic nomenclature existed. Now these names have becomes so popular that it is difficult to change them. Do not worry, as you go through these notes, you will get a sufficient idea regarding these names. Cations (Positive ions) Metal atoms with single positive charge Rule: Names of positive ions end in-ium if the ion has only one oxidation state (Only one level of net charge). For example, the positive ion of sodium is Na+ (sodium ion), and the positive ion of aluminum is Al3+ (aluminum ion). Metal atoms with more than one possible charge Rule: If the cation has variable valency (charge), charge is specified in roman numerals in round brackets immediately after the name of metal atom. For example, Sn2+ is written as tin (II) ion. Alternately, the less positive ion ends in -ous, and the more positive ion ends in-ic . For instance, the two positive ions of copper are Cu+ (cuprous) and Cu2+ (cupric). The oxidation state of a positive ion can also be designated by placing a Roman numeral after the name of the elements. These positive ions of copper can also be written as copper (I) and copper (II), respectively. 29 Chemistry Ions Name + Cu cuprous ion Sn Sn 2+ Stannous ion 4+ Stannic ion 3+ Fe Ferric ion 2+ Fe Ferrous ion General Rule-3 Suffix-nium is often used with cations containing non-metals. For example, the positive ion of ammonia is NH4+ (ammonium) and the positive ion of water (H2O) is H3O+ or H+ (hydronium). Remember these names ! + NO2 : nitronium NO+ : nitrosonium H3O+ : hydronium + From NH3 ammonia is derived NH4 : ammonium. Similarly. N2H4 : hydrazine  N2H5 : hydrazinium C6H5NH2 : aniline  C6H5NH3+ :anilinium C5H5N : pyridine C5H5NH+ : pyridinium Anions (Negative Ions) Anions can always be looked upon as ions derived from acids by removal of one or more protons. Accordingly, anions can be classified as follows. Anions derived from hydracids Rule : Names of negative ions from hydracids end in –ide. For example , Cl– (chloride) from HCl, and CN– (cyanide) from HCN. Following examples will give you a better insight in this nomenclature. It is also useful to remember them. 30 Chemistry Remember these names F– fluoride from HF Cl– chloride from HCl Br– bromide from HBr I– iodide from HI O2– oxide S2– sulphide N3– nitride P3– phosphide As3– arsenide Cu– cupride H– hydride Au– auride Anions derived from oxyacids Anion derived from an oxyacid by removal of one or more H + ions is termed as oxyanion. Rule: If the oxyacid is – ic acid, suffix - ate is used with oxy-anion. For example CO32– carbonate (from H2CO3) ZnO22– zincate SiO32– silicate Rule : If the oxyacid is - ous acid, suffix -ite is used with oxy-anion. For example NO2– (nitrite) is derived from HNO2 (nitrous acid), and SO32– (sulfite) is derived from H2SO3 (sulfurous acid) Rule : If the oxyacid has prefixes per-or hypo–, the oxyanion will have same prefixes. For example, ClO4– percholrate ion from HClO4, perchloric acid ClO– hypochlorite ion from HClO, hypochlorous acid Remember these names ! 31 Chemistry SO42– sulphate SO32– Sulphite NO3– nitrate, NO2– nitrite SnO32–stannate SnO22–stannite, PbO32–plumbate, PbO22–plumbite Anions containing replacable hydrogen ions Polyprotic acid. Any acid containing more than one replacable hydrogens is said to be a polyprotic acid. Replacable hydrogens. H atoms which can be lost as H+ in reactions with a base. H atoms connected to O atoms in oxyacid’s are all replacable. If all the replacable hydrogens are removed, we obtain the anions discussed in the sections above. However, in all the polyprotic acids it is always possible to remove less than the maximum number of replacable hydrogens . e.g. H3PO4 is triprotic. We can remove one, two or three H+ ions from it to generate H2PO4–, , HPO42– and PO43– . You are already familiar with phosphate ion, PO3–. The other two anions, H2PO4– and HPO42– still contain H atoms that are replacable. We consider their nomenclature in this section. Rule-1 : A prefix bi- (old notation) or hydrogen – (IUPAC notation) is attached to the name of anion. Rule -2 : For triprotic or higher acids, numerical prefixes (e.g. mono, bi, tri) are also used to indicate the number of replacable H atoms left in the sample). Ex. HCO3– is bicarbonate or hydrogen carbonate HSO3– bisulphite or hydrogen sulphite HS– bisulphide or hydrogen sulphide etc. when anion has – 3 charge , e.g. PO43– then following possibilities arise HPO42– monohydrogen phosphate, H2PO4– dihydrogen phosphate Miscellaneous Anions (To be committed to memory) NH2– amide, NH2– imide, N3– nitride, N3–azide, O22– peroxide, O2– superoxide, O3– ozonide CrO42– chromate, Cr2O72– dichromate, MnO42– manganate, MnO4– permanganate , CN– cyanide OCN– cyanate, SCN– thiocyanate, S2O32– thiosulphate, CH3COO– acetate, C2O42– oxalate [Fe (CN)6]3– ferricyanide ,[Fe(CN)6]4– ferrocyanide 32 Chemistry Difference between Atoms and ions We have been talking for quite sometime. The distinction between atoms and ions, though simple, must be very clear in your mind. The following are the points summarise this difference. Method of writing formula of an ionic compound In order to write the formula of an ionic compound which is made up of two ions (simple or polyatomic) having net charges x and y respectively, follow the following procedure. (I) Write the symbols of the ions side by side in such a way that positive ion is at the left and negative ion at the right as AB. (II) write their charges on the top of each symbol as AxBy. X (III) , i.e. formula AyBx. Now apply crisis- cross rule as A (IV) Y B Cancel out any common factor (or HCF). Ex. : 1. Calcium 2. Aluminium oxide 3. Potassium chloride = CaCl2 = Al2O3 phosphate = K3PO4 33 Chemistry 4. Magnesium nitride 5. Calcium oxide = Ca2O2 6. Ammonium sulphate = Mg3N2 = (NH4)2SO4 Cancelling the common factor, answer is CaO 1. Atoms Atoms are perfectly neutral 1. 2. In atoms, the number of portons is equal to the 2. number of electrons. Na (protons 11, electrons 11) Cl (protons 17, electrons 17) 3. Except noble gases, atoms have less than 8 3. electrons in the outermost orbit e.g. Na 2, 8, 1; Ca 2, 8,8, 2 Cl 2, 8, 7 ; S 2, 8, 6 Chemical activity is due to loss or gain or 4. sharing of electrons as to acquire noble gas configuration 4. Ions Ions are charged particles containing one or more atoms. In cations (positively charged ions), number of protons is more than the number of electrons. In anions (negatively charged ions) the no. of protons is less than the number of electrons. e..g. Na+ (protons 11, electrons 10). Cl– (protons 17, electrons 18) Ions have generally 8 electrons in the outermost orbit , i.e., ns3np6 configuration. Na+ , 2, 8; Cl– 2, 8, 8 ; Ca2+ 2, 8, 8 The chemical activity is due to the charge on the ion. Oppositely charged ions are held together by electrostatic forces. 34