History Of Periodic Table

History of Periodic Table 1869: Dmitri Mendeleev organized the periodic tablebased on atomic weights “Father of the Periodic Table” 1913: Henry Moseley rearranged the periodic tablebased on the positive charges in the nucleus Lead to the periodic law: the states that a periodicpattern appears in the physical and chemical propertiesof the element when they are arranged in order ofincreasing atomic number

Discovering the Periodic TableHLiAncient Times1894-1918Midd. PSClArCr Mn Fe Co Ni Cu Zn Ga Ge As Se BrKrNa MgKCa ScRb SrYCs Ba LaFrTiVHeZr Nb Mo Tc Ru Rh Pd Ag CdInSn Sb TeHfTlPb BiTaWRe OsIrPt Au HgIXePo At RnRa Ac Rf Db Sg Bh Hs MtCe Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb LuTh PaJournal of Chemical Education, Sept. 1989UNp Pu Am Cm Bk CfEs Fm Md No Lr

Periodic Trends A pattern where elemental characteristics changepredictably as you go across a period or down agroup You need to know the trends in: Metals v. Nonmetals Atomic Radius (Atomic Size) Ionization Energy Electronegativity Ion Formation

Groups of Elements1A1H123Be3472AAlkali earth metals6AOxygen groupTransition metals7AHalogens3ABoron group8ANoble gases4ACarbon group8AHe3A 4AB C5A 6A 7A 2N O F NeHydrogenInner transition metals5678910AlSiPSClAr8BK3B 4B 5B 6B 7B1B 2B 13 14 15 16 17Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br12202122Rb SrYZr Nb Mo Tc Ru Rh Pd Ag CdIn3940414249HfTaW727374376Nitrogen groupNa Mg1955A2ALi1141A Alkali metals38Cs Ba5556FrRa8788*W25432644Re Os757627282947304546IrPt Au HgTl77788179483180323334Sn Sb Te5051Pb Bi8283523536IXe5354Po At Rn848586105106107108109La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu57W24KrRf Db Sg Bh Hs Mt104*23185859Ac Th Pa89909160U92616263646566Np Pu Am Cm Bk Cf9394959697986768697071Es Fm Md No Lr99100101102103

Metals vs. Nonmetals Mainly divided into metals and nonmetals Metals: On the left-hand side (left of stair-step line) Non-metals: On the upper right-hand side Metalloids: On the stair-step line

Metals and Nonmetals123HHe12LiBeBC345Na Mg114K1957Ca ScOFNe678910AlSiPSClAr131415161718TiVCr Mn Fe Co Ni Cu Zn Ga Ge As Se BrKr23243536IXe5354202122Rb SrYZr Nb Mo Tc Ru Rh Pd Ag CdIn3940414249HfTaW72737437612N38Cs Ba5556FrRa8788*WNonmetals252627282930METALS4344Re Os7576474546IrPt Au HgTl77788179483180323334Sn Sb Te5051Pb Bi828352Po At Rn848586Rf Db Sg Bh Hs Mt104105106107108Metalloids109La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu575859Ac Th Pa89909160U92616263646566Np Pu Am Cm Bk Cf9394959697986768697071Es Fm Md No Lr99100101102103

Metals, Nonmetals, & Metalloids1Nonmetals2345Metals67MetalloidsZumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 349

Properties of Metals, Nonmetals,and MetalloidsMETALSmalleable, lustrous, ductile, good conductors of heatand electricityNONMETALSgases or brittle solids at room temperature, poorconductors of heat and electricity (insulators)METALLOIDS (Semi-metals)dull, brittle, semi-conductors (used in computer chips)

Orbitals Being 5p66sLa5d6p77sAc6dZumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 34567 1s4fLanthanide series5fActinide series

Electron Filling in Periodic Tablemetallic character increasesnonmetallic character increasesmetallic character increasesnonmetallic character increases

Melting PointsHe1HMg-259.2234LiBe180.51283650KCa ScRb Sr38.86 300098850770710BoC2000 - 3000oCAl660TiVCNOFY1500 1852 2487 2610 2127 2427 1966 1550920TaP1423 44.2420 29.78 960Zr Nb Mo Tc Ru Rh Pd Ag CdHfSiS119NeWRe OsIr961InAr-101 -189.6Kr817 217.4 -7.2 -157.2Sn Sb TeIXe321 156.2 231.9 630.5 450 113.6 -111.9Pt Au HgTlPb BiPo At Rn2222 2997 3380 3180 2727 2454 1769 1063 -38.9 303.6 327.4 271.3 254Ralph A. Burns, Fundamentals of Chemistry , 1999, page 1999ClCr Mn Fe Co Ni Cu Zn Ga Ge As Se Br1423 1677 1917 1900 1244 1539 1495 1455 1083Cs Ba La28.6-269.72027 4100 -210.1 -218.8 -219.6 -248.6Na Mg63.25650He0.126SymbolMelting point oC-71

Densities .811.14Na MgAlSiPS0.972.702.4 1.82w 2.07 1.557 1.402K0.865Ca ScTiV1.554.55.96Rb 74.76.48.410.28.67.37.3Cs Ba LaHfTaWPt Au HgTlPb Bi1.9013.116.619.38.0 – 11.9 g/cm3Mg1.74WAr3.1197.45.516.7Cl7.1Sn Sb Te3.51.11 1.204KrIn2.6YNeCr Mn Fe Co Ni Cu Zn Ga Ge As Se BrZr Nb Mo Tc Ru Rh Pd Ag Cd1.5361.74F11.512.5Re Os12.5Ir12.010.521.4 22.48 22.4 21.45 19.3 13.55 11.85 11.3412.0 – 17.9 g/cm36.79.86.1Po At Rn9.4 18.0 g/cm3SymbolDensity in g/cm3C, for gases, in g/L---4.4

Atomic Radius The distance from the nucleus to the outer egde of theelectron cloud Increases as you go down a group Example: Calcium atoms are larger than berylliumatoms Decreases as you go across a period Example: Fluorine atoms are smaller than Oxygen atoms Why? Down groups, you have more electron shells Across periods, you have the same number of electronshells, but more protons (positive). This pulls in theelectrons.

Atomic Radius vs. Atomic Number0.3CsRbatomic esLi0.15LaZnXeKr0.1ClF0.05HeH00102030atomic number405060

Atomic .171.711.751.46 1 Angstrom

Periodic Trends in Atomic RadiiLeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 175

Relative Size of AtomsZumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 350

Decreasing Atomic SizeAcross a Period As the attraction between the ( ) nucleus and the (–) valence electrons ,the atomic size . Greater coulombic attraction. From left to right, size decreases because there is an increase in nuclearcharge and Effective Nuclear Charge (# protons – # core electrons). Each valence electron is pulled by the full ENCLiBeB1s22s11s22s21s22s22p1(ENC 1)(ENC 2)(ENC 3)LiBeB

Ionic Radii The distance from the nucleus to the edge of theelectron cloud of an ion Increases as you go down a group Example: Calcium ions are larger than beryllium ions Decreases as you go across a period Example: Fluorine ions are smaller than Oxygen ions Why? Down groups, you have more electron shells Across periods, you have the same number of electronshells, but more protons (positive). This pulls in theelectrons.

Sizes of ions: electron repulsion Valence electrons repel each other. When an atom becomes aanion (adds an electron to itsvalence shell) the repulsionbetween valence electronsincreases without changing ENC Thus, F– is larger than F- 9 - Fluorine atomF21s 2s22p5 1e--9 9 --Fluorine ionFluorideF11s22s22p6

PbBi1.711.751.462.622.17Li1 Be2 0.60Na1 0.310.950.65K1 Cations: smallerthan parent atomsIIAMg2 Ca2 N31.71Al3 0.50Ga3 1.33Rb1 0.99Sr2 0.621.48Cs1 1.13Ba2 0.81Tl3 1.691.350.95In3 1-2.212.16 1 AngstromAnions: LARGERthan parent atoms

Electronegativity The ability of an atom to attract electrons towardsitself from a covalent chemical bond Fluorine is the most electronegative element Decreases as you go down a group Example: Chlorine has a lower electronegativity thanFluorine Increases as you go across a period Example: Fluorine has a higher electronegativity thanOxygen Why? Bigger atoms have a harder time pulling electrons in tothemselves

.53.0Na Mg1.23B4B5B6BKCa ScTiVCr Mn Fe Co Ni Cu Zn Ga Ge As Se Br0.81.01.31.51.61.61.71.61.8Rb SrYZr Nb Mo Tc Ru Rh Pd Ag CdInSn Sb Te0.81.21.41.61.81.92.22.22.21.71.71.8Cs BaLa*HfTaWRe OsIrPt Au HgTlPb BiPo At0.71.11.31.51.71.92.22.21.81.82.01.00.9yFrRa .01.91.92.42.1* Lanthanides: 1.1 - 1.3yActinides:1.3 - 1.5Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 373Below 1.02.0 - 2.41.0 - 1.42.5 - 2.91.5 - 1.93.0 - 4.02.8I2.52.2

Ionization Energy The amount of energy needed to remove an electronfrom an atom Decreases as you go down a group Example: Lithium has a higher ionization energy thansodium Increases as you go across a period Example: Chlorine has a higher ionization energy thanphosphorous Why? Electrons are easier to remove from large elements Electrons are harder to remove from atoms that almosthave their “happy” full shell of 8 electrons

HeNeArFirst Ionization energyKrHLiNaKRbAtomic number

Ionization Energies It takes more energy to remove the second electron from an atomthan the first, and so on. There are two reasons for this trend:1. The second electron is being removed from a positivelycharged species rather than a neutral one, so more energy isrequired.2. Removing the first electron reduces the repulsive forcesamong the remaining electrons, so the attraction of theremaining electrons to the nucleus is stronger. Energy required to remove electrons from a filled core is prohibitivelylarge and simply cannot be achieved in normal chemical reactions.Copyright 2007 Pearson Benjamin Cummings. All rights reserved.

Factors Affecting Ionization EnergyNuclear ChargeThe larger the nuclear charge, the greater the ionization energy.Shielding effectThe greater the shielding effect, the less the ionization energy.RadiusThe greater the distance between the nucleus and the outerelectrons of an atom, the less the ionization energy.SublevelAn electron from a full or half-full sublevel requires additionalenergy to be removed.Smoot, Price, Smith, Chemistry A Modern Course 1987, page 189

Nuclear charge increasesShielding increasesAtomic radius increasesIonic size increasesIonization energy decreasesElectronegativity decreasesSummary of Periodic TrendsShielding is constantAtomic radius decreasesIonization energy increasesElectronegativity increasesNuclear charge increases1A02AIonic size (cations)decreases3A 4A 5A 6A 7AIonic size (anions)decreases

History of Periodic Table 1869: Dmitri Mendeleev organized the periodic table based on atomic weights “Father of the Periodic Table” 1913: Henry Moseley rearranged the periodic table based on the positive charges in the nucleus Lead to the periodic law: the states that a periodic pattern appears in