Periodic Table The most useful tool in the Lab Early Organization • J.W. Dobereiner (1829) organized elements in triads ▫ Triad – three elements with similar properties (ex: Cl, Br, I) • J.R. Newlands (1864) organized elements in octaves ▫ Octave – repeating group of 8 elements Development of the PeriodiceTable • Dmitri Mendeleev taught chemistry in terms of properties. • Mid 1800’s - molar masses of elements were known. • Wrote down the elements in order of increasing mass. • Found a pattern of repeating properties. Mendeleev’s Table • Grouped elements in columns by similar properties in order of increasing atomic mass. • Found some inconsistencies - felt that the properties were more important than the mass, so switched order. • Also found gaps. • Must be undiscovered elements. • Predicted their properties before they were found. The Modern Periodic Table • Henry Moseley – British physicist • Arranged elements according to increasing atomic number • The arrangement today • Symbol, atomic number & mass The New Way • • • • • Elements are still grouped by properties. Similar properties are in the same column. Order is by increasing atomic number. Added a column of elements (noble gases) Weren’t found because they are unreactive. Organization • Horizontal rows = periods ▫ There are 7 periods ▫ Each period represents an energy level ▫ Every element in the same period has the same # of energy levels and the same core electron configuration Organization • Vertical column = group or family ▫ Similar physical & chemical prop. ▫ Same # of valence electrons ▫ Same common oxidation state ▫ Identified by number & letter • Horizontal rows are called periods • There are 7 periods • Group 1A are the alkali metals • Group 2A are the alkaline earth metals • Group 7A is called the Halogens • Group 8A are the noble gases The group B are called the transition elements These are called the inner transition elements, and they belong here • The elements in the A groups are 1A called the representative elements 8A 2A outer s or p filling 3A 4A 5A 6A 7A Lanthanides – the 4f orbital fills for these elements Actinide series – the 5f orbitals are being filled for these elements. Types of elements • Metals • Non-metals • Metalloids or semi-metals Metals • • • • • • • Good conductor of heat and electricity Malleable Ductile High tensile strength High luster Solid at room temperature React by losing electrons Nonmetals • Poor conductors of heat and electricity • React by gaining electrons • Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) Semi-Metals • Heavy, stair-step line • Metalloids border the line ▫ Properties intermediate between metals and nonmetals ▫ Learn the general behavior and trends of the elements, instead of memorizing each element property • B, Si, Ge, As, Sb, Te Families Group IA – alkali metals most reactive metals Silvery in appearance Soft Combine easily with non-metals Melting point is higher than the boiling point of water Have 1 valence electron Families • Group 2 – Alkaline Earth Metal Family ▫ Harder, stronger, denser, higher melting point, and less reactive than alkali ▫ Usually not found as free elements, but as compounds ▫ Have 2 valence electrons Families • Group 7 – Halogens ▫ Most reactive family ▫ Non-metals ▫ Have seven valence electrons • Group 8 – Noble Gas ▫ Inert, unreactive ▫ Have full set of valence electrons H 1 1s 1 Li 22s1 1s 3 Na 11 K 19 Rb 37 Cs 55 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 22s22p63s23p64s23d104p65s24d105p66s24f 1s Fr 87 145d106p67s1 S- block s1 s2 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • really should include He, but it fits better later. • He has the properties of the noble gases. 1s2 He 2 Ne 2 2 6 1s 2s 2p 10 1s22s22p63s23p6 Ar 18 1s22s22p63s23p64s23d104p6 Kr 36 1s22s22p63s23p64s23d104p65s24d105p6 Xe 54 1s22s22p63s23p64s23d104p65s24d10 Rn 6 2 14 10 6 5p 6s 4f 5d 6p 86 The P-block p1 p2 p3 p4 p5 p6 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals. Areas of the periodic table • Group A elements = s & p blocks • representative elements ▫ Wide range of phys & chem prop. Transition Metals -d block 1 d 2 d 3 d s1 5 d s1 5 6 7 8 10 10 d d d d d d • d orbitals fill up after previous energy level, so first d is 3d even though it’s in row 4. 1 2 3 4 5 6 7 3d F - block • inner transition elements f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 1 2 3 4 5 6 7 4f • f orbitals start filling at 4f 5f Atomic Size } Radius • First problem: Where do you start measuring from? • The electron cloud doesn’t have a definite edge. •Atomic Radius = half the distance between two nuclei of a diatomic molecule. Trends in Atomic Size Influenced by three factors: 1. Energy Level Higher energy level is further away. 2. Charge on nucleus More charge pulls electrons in closer. 3. Shielding effect (blocking effect) WHAT HAPPENS TO ATOMIC RADII? • Does a negative ion (anion) get larger or smaller? • Does a positive ion (cation) get larger or smaller? Trends in Ionic Size • Cations form by losing electrons. • Cations are smaller than the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration. Ionic size • Anions form by gaining electrons. • Anions are bigger than the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration. WHAT IS IONIZATION ENERGY? • The energy required to remove an electron • Which element has the highest ionization energy? Why? What determines Ionization Energy? • The greater the nuclear charge, the greater IE. • Greater distance from nucleus decreases IE • All the atoms in the same period have the same energy level. • But, increasing nuclear charge • So IE generally increases from left to right. Ionization Energy The energy required to remove the first electron is called the first ionization energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE. Driving Force • Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full orbitals. • Atoms behave in ways to achieve noble gas configuration. WHAT IS ELECTRONEGATIVITY? • The ability of an atom to pull off an electron. • Which element has the highest electronegativity? Why? Periodic Trend • Metals are at the left of the table. • They let their electrons go easily • Low electronegativity • At the right end are the nonmetals. • They want more electrons. • Try to take them away from others • High electronegativity. Trends in Electron Affinity • The energy change associated with adding an electron to a gaseous atom. • Easiest to add to group 7A. • Gets them to full energy level. • Increase from left to right: atoms become smaller, with greater nuclear charge. • Decrease as we go down a group.