1 Periodic Table • Lavoisier 1789 – 33(21) elements Traité Élémentaire de Chimie (1789) the 1st modern chemistry textbook • Dalton 1808 - 36 elements • Berzelius 1813-14 - 47 elements • Mendeleev 1869 - 63 elements • The last element discovered in Nature 1939 – 223Fr • Nuclear synthesis of new elements since 1940 • 2012 - 118 elements, named up to 112 2 Periodic Table 1829, Johann Wolfgang Döbereiner (1780-1849) Triades: Li, Na, K Ca, Sr, Ba S, Se, Te Cl, Br, I Jena, Germany 3 Periodic Table 1859, Jean-Baptiste Dumas (1800-1884) Quadruplets: F, Cl, Br, I; Mg, Ca, Sr, Ba 1863, Alexandre-Émile Béguyer de Chancourtois (1820-1886) Spiral 1864, William Odling (1829-1921) Groups of seven elements 1864, John Alexander Reina Newlands (1837-1898) Ordered elements according their atomic mass Law of octaves 4 5 Periodic Table 1869, 1871 Mendeleev – predicted properties of missing elements (Sc, Ga, Ge, Tc, Rh, Po, Hf), Rare gasses He, Ar Properties of elements are a periodic function of atomic mass (exceptions: Ar/K; Co/Ni; Te/I; Pa/Th) 1913 Moseley Corrected the periodic law : Properties of elements are a periodic function of atomic number 1870, Lothar Meyer (1830-1895) Periodicity of atomic volumes 6 Periodická tabulka prvků 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Periodic Table 2012 7 Periodic Properties of Atoms Atomic numbers - effective nuclear charges Oxidation numbers Atomic radii Ionistaion energies Electron affinity Electronegativity Polarizibility, polarization power Metallic – Semimetallic – Nonmetallic Properties 8 Groups and Periods Groups: repeating electron configuration governs similarity of chemical properties Periods: stepwise filling of electron shells and nuclear charge increase govern changes in properties 9 Rules for Filling Orbitals with Electrons The orbitals with lowest energy are filled first – Aufbau principle Only two electrons fit into one orbital with opposite spin – Pauli’s principle Maximum number of unpaired electrons in energetically degenerate atomic orbitals – Hund’s rule Occupation of orbitals by electrons can change order of their energies 10 11 Electronic Configurations of Main Group Elements Main Group Elements = s- and p-elements Fill s and p orbitals Oxidation number changes by 2 ns2 npx CO CO2 SO2 SO3 PCl3 PCl5 Alkali metals : ns1 Alkaline earths : ns2 Triels : ns2 np1 Tetrels : ns2 np2 Pniktogens : ns2 np3 Chalkogens : ns2 np4 Halogens : ns2 np5 Rare gasses : ns2 np6 very stable configuration 12 Main Group Compounds Oxidation number changes by 2 result of ns2 npx Diamagnetic = no unpaired electrons (except O2) Colorless 13 Electronic Configurations of Transition Elements Transition elements = d-Elements Fill (n−1)d and ns orbitals Oxidation number changes by 1 3d, 4d, 5d, 6d elements – 4th to 7th period (n−1)dx At least in one compound have incompletely filled d orbitals Zn (M2+ = d10) , Sc (M3+ = d10), recently Sc1+ Early transition elements – oxophilic, 3rd – 7th group, lack of d-electrons Late transition elements – chalkophilic, 8th – 12th group, excess of d- electrons 14 Transition Element Compounds Oxidation number changes by 1 result of (n−1)dx More oxidation states Paramagnetic Colored 15 Characteristic Oxidation Numbers of 3d-Elements Cu2+Cu+ Zn2+ Ni2+ Co3+Co2+ FeO4 2−Fe3+Fe2+ MnO4 −MnO4 2−MnO4 3−Mn4+Mn3+Mn2+ CrO4 2−Cr3+Cr2+ VO2 +VO2+V3+V2+ Ti4+Ti3+ Sc3+Sc+ 7654321 16 Change in Ordering of Energy Levels 4s/3d Ar [Ne] 3s2 3p6 (4s0) K [Ar] 4s1 (3d0 4p0) Ca [Ar] 4s2 (3d0 4p0) Sc [Ar] 3d1 4s2 (4p0) Ti [Ar] 3d2 4s2 (4p0) 17 Change in Ordering of Energy Levels 4s/3d Experimental results Increasing effective nuclear charge Screenig of electrons 18 High Stability of Half-Filled Orbitals Cu [Ar] 3d10 4s1 (4p0) Cr [Ar] 3d5 4s1 (4p0) 19 Electronic Configurations of Free and Bound Atoms Ni [Ar] 3d9 4s1 (4p0) free atom in vacuum Ni [Ar] 3d10 (4s0 4p0) in compounds, Ni(CO)4 20 Lanthanides and Actinides La = (Xe) 5d1 6s2 Lu = (Xe) 4f14 5d1 6s2 f-elements Ac = (Rn) 6d1 7s2 Lr (Rn) 5f14 6d1 7s2 21 Electronic Configurations of Lanthanides Xe [Kr] 4d10 5s2 5p6 E(4f) > E(6s) Cs [Xe] 6s1 4f0 5d0 Ba [Xe] 6s2 4f0 5d0 La [Xe] 4f0 5d1 6s2 transition metal Ce [Xe] 4f1 5d1 6s2 E(4f) < E(6s), E(5d) Pr [Xe] 4f3 6s2 Eu [Xe] 4f7 5s2 5p6 5d0 6s2 Gd [Xe] 4f8 5s2 5p6 5d0 6s2 Gd [Xe] 4f7 5s2 5p6 5d1 6s2 4f half filled Lu [Xe] 4f14 5d1 6s2 4f completely filled 22 23 Electronic Configurations of Actinides Rn [Xe] 4f14 5d10 6s2 6p6 E(5f) > E(7s) Fr [Rn] 7s1 Ra [Rn] 7s2 5f 0 6d0 Ac [Rn] 5f0 6d1 7s2 transition metal Th [Rn] 5f0 6d2 7s2 E(5f) < E(7s), E(6d) Pa [Rn] 5f2 6d1 7s2 U [Rn] 5f3 6d1 7s2 Np [Rn] 5f4 6d1 7s2 Pu [Rn] 5f6 6d0 7s2 Am [Rn] 5f7 6d0 7s2 Cm [Rn] 5f7 6d1 7s2 Bk [Rn] 5f8 6d1 7s2 Cf [Rn] 5f10 6d0 7s2 Es [Rn] 5f11 6d0 7s2 Fm [Rn] 5f12 6d0 7s2 Md [Rn] 5f13 6d0 7s2 No [Rn] 5f14 6d0 7s2 Lr [Rn] 5f14 6d1 7s2 24 Elektronová slupka Valenční sféra – atomové orbitaly, nejvzdálenější od jádra, zcela nebo zčásti zaplněné, které leží nad elektronovou konfigurací nejbližšího nižšího vzácného plynu Valenční sféra rozhoduje o fyzikálních a chemických vlastnostech Vnitřní elektrony – elektronové “jádro” – všechny nižší zcela zaplněné elektronové hladiny vzácných plynů, neúčastní se chemických reakcí 25 Octet Formation Ar [Ne] 3s2 3p6 Isoelectronic ions 26 Size of Atoms Atomic radii Approximation of atoms as inflxible koule, r = 10−10 m Covalent radius = half of interatomic distance between same atoms Diamond C-C distance = 1.54 Å Covalent radius = 0.77 Å 27 Size of Atoms Atomic radii increase down the group – filling of higher (n) orbitals by electrons, electrons are further from nucleus Filled d-orbitals: r(Al) > r(Ga) Al [Ne] 3s2 3p1 (3d0) Ga [Ar] 3d10 4s2 4p1 28 Radius of maximum electron density Atomic number, Z Filled d-orbitals: r(Al) > r(Ga) Al [Ne] 3s2 3p1 (3d0) Ga[Ar] 3d10 4s2 4p1 29 Size of Atoms Radius 30 Atomic Radii (pm) Radius increases 31 Covalent Radii, rcov (Å) r(2s) > r(2p) r(3s) ∼ r(3p) Atomic Number, Z 32 Size of Atoms Atomic radii decrease along the periods: electrons added to orbitals with the same n, increasing Z – positive nuclear charge contracts electron shell 33 Lanthanide Contraction Lanthanide Contraction: outer orbital is 6s, electrons filled to 4f, increasing Z, Radii decrease from La 169 pm to Lu 153 pm 34 Atomic Radii, pm 35 Atomic Radii, Å Atomic Radius, Z 36 Atomic Radii of Transition Elements, pm 37 Ionic Radii Ionic Radii Increase down the group Ionic Radii, Å 38 Ionic Radii Isoelectronic ions: N3− > O2− > F− > Na+ > Mg2+ > Al3+ With increasing Z and increasing positive charge radii decrease Cations smaller than neutral atoms Anions bigger than neutral atoms Fe2+ > Fe3+ Pb2+ > Pb4+ With increasing positive charge, radii decrease 39 Ionic and Atomic Radii, Å 40 Ionization Ionization = removal of an electron from an atom (or ion) Spent energy = always endothermic Electron further from nucleus is torn off most easily, wekly bound Removing second electron from a cation even more energetically demanding: Removing electron – decrease of e-e repulsion, radius decreases Cations smaller than neutral atoms Anions bigger than neutral atoms 41 Ionization Energy, IE IE = energy for removing the weakest bound electron from an atom in a gas phase (at 0 K) [kJ mol−1]. Electron bonding in a given orbital M(g) M(g) + + e− First IE IE1 > 0 [kJ mol-1] M(g) + M(g) 2+ + e− Second IE IE2 >> 0 [kJ mol-1] 1. IE < 2. IE < 3. IE < 4. IE < ....... Subsequent ionization is more demanding energetically : same Z, smaller number of e is held more tightly, charge separation is disadvantageous 42 Ionization Energy, IE [kJ mol−1] Uzavřené elektronové slupky 43 Cl [Ne] 3s2 3p5 44 Ionization Energy, IE (kJ mol−1) Atomic Number, Z High IE for rare gases Low IE for alkali metals 45 Ionization Energy, IE (kJ mol−1) Increases along a period Decreases down a group 46 First Ionization Energy, IE (kJ mol−1) Atomic Number, Z Mg-Al P-SN-O Be-B 47 Electron Affinity, EA EA = Electron affinity released (EA < 0) or absorbed (EA > 0) during attachment of an electron to an atom or an ion in gas phase (at 0 K). First EA generally < 0, except Be, N, ….. Why? Second EA always > 0, attachment of an e− to anion is unfavorable by energy, compensated by lattice energy release Oxides, O2− EA1(O) < 0 EA2(O) > 0 48 First Electron Affinity, EA (kJ mol−1) Energy released (EA < 0)- - - - - - - - Octet Alkali Metal Anions 49 First Electron Affinity, EA (kJ mol−1) Increases along a period Decreases down a group Energy released (EA < 0) 50 Pauling’s Electronegativity Ability of an atom to attract bonding electrons in covalent bond Dissociatin energy of a polar bond A−B is higher than an average of dissociation energies of nonpolar bonds A−A and B−B ED(AB) = {ED(AA) × ED(BB)}½ + Δ Δ = 96.48 (χA − χB)2 χF = 4.0 Pauling χF = 3.98 today’s value Linus Pauling (1901-1994) NP in Chemistry 1954, Peace 1963 51 Pauling’s Electronegativity Dissociation energy from experiments ED(F2) = 154.8 kJ mol−1 ED(Br2) = 192.5 kJ mol−1 ED(BrF) = 238.5 kJ mol−1 ED(BrF) = {ED(F2) × ED(Br2)}½ + Δ Δ = 96.48 (χA − χB)2 χF = 3.98 χBr = ? AB χχ − Δ = 48.96 Square root of energy?? 52 Pauling’s Electronegativity 70.4475218297HI 130.75396218367HBr 170.992122218432HCl 431.927077218565HF % ionicity χB - χA Δ½ ED(BB) kJ mol−1 ½ ED(AA) kJ mol−1 ED(A-B) kJ mol−1 A-B 53 Mulliken’s Electronegativity Orbital electronegativities – s, p, d, hybrid χM = 3.15 χP 2 EAIE M + =χ 54 Allred and Rochow’s Electronegativity Coulombic force – nucleus attracts bonding electrons 2 04 1 r eZ F eff πε = B r Z A eff AR += 2 χ 55 Electronegativity Increases along a period Decreases down a group 56 Mutual Polarization of Ions Mn+ Am- Polarization Polarizibility of an anion, an atom, or a molecule Polarization power of a cation 57 Polarizibility, α [m3] Deformation of electronic density in an atom or an ion by external electric field (other charged particles) Change of volume of electron cloud, α [m3] Value of α depends on strength of electron bonding by a nucleus, atomic size, number of electrons Soft atom (ion, molecule) = easy to deform Hard atom (ion, molecule) = resists deformation 58 Polarizibility of Atoms, 106 pm3 Cs Rb K Na Li Atom 52.9 43.7 41.6 24.4 24.0 α Xe Kr Ar Ne He Atom 1.322 1.419 1.329 1.027 α 3.99 2.46 1.62 0.39 0.20 α C(ar) C(2) C(3) C(4) Atom 5.530 3.465 2.317 0.321 0.408 α Br I Cl F H Atom 59 Polarization Power Increases with increasing charge Increases with decreasing radius q/r charge density Al3+ hard cation Cs+ soft cation 60 Metallic – Nonmetallic Properties 61 Metallic – Nonmetallic Properties 62 Metals Close-packed structures High coordination numbers (12) Large atoms Low ionization energies High polarisibility Omnidirectional metallic bonging 63 64 Nonmetals Covalent bonds strong and directional Good orbital overlap Small atoms High ionization energies, low polarizibility Weak vdW interactions 65 Metalloids - Semimetals Weaker covalent bonds vdW interactions Secondary bonds 66 Metalloids - Semimetals 0 0.5 1 1.5 2 2.5 3 12 13 14 15 16 17 18 B Al C Si Ge Sn Pb N P As Sb Bi O S Se Te Po F Cl Br I 0 0.5 1 1.5 2 2.5 3 12 13 14 15 16 17 18 B Al C Si Ge Sn Pb N P As Sb Bi O S Se Te Po F Cl Br I Nonmetals Semimetals Metals Group . IE r 67 16th Group O and S - nonmetals Se - nonmetallic and semimetallic modifictions (allotropes) Te - semimetallic Po - metallic with a rare structure 68 16th Group Po - metal 69 Te Te - semimetal 70 Se Gray selenium Red selenium Se8 nonmetal semimetal 71