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Periodic Table

Periodic Table. Modern Periodic Law. The properties of the elements repeat in a regular pattern when arranged by their atomic numbers. History. Johann Dobereiner – 1829 (friend of Goethe) He was the first to organize elements by their properties

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Periodic Table

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  1. Periodic Table

  2. Modern Periodic Law The properties of the elements repeat in a regular pattern when arranged by their atomic numbers.

  3. History • Johann Dobereiner – 1829 (friend of Goethe) • He was the first to organize elements by their properties • He grouped them in groups of three called triads

  4. triads • He noticed that the atomic mass of the middle member of the group was close to the arithmetic mean of the others. • Chlorine = 35.5, Bromine = 80, Iodine = 127 (average of Cl and I = 81) • Properties in common: • All react vigorously with first column metals to form soluble salts (compounds of a metal and nonmetal) • Hydrogen compounds are strong acids • All form -1 ions

  5. triads • Lithium = 7, Sodium = 23, Potassium = 39 (average of Li and K = 23) • Properties in common: • All salts are soluble • All give brightly colored flames • All react vigorously with water • All form +1 ions

  6. Other triads • Calcium = 40, Strontium = 88, Barium = 137 (average of Ca and Ba = 88.5) - All give +2 ions • S = 32, Se = 79, Te = 127.6 (average of S and Te = 79.8) - All give smelly compounds with hydrogen

  7. Failure of triads • Not all elements could be fit into triads: iron, manganese, nickel, cobalt, zinc and copper are similar elements but cannot be placed in the triads. • Newly discovered elements did not fit into triads • Very dissimilar elements could be fit into triads • Dobereiner’s triads were discarded

  8. Newlands’ octaves • John Newlands  1838 - 1898 • Law of Octaves (1863) • Elements can be arranged in “octaves” because certain properties repeated every 8th element when the elements are arranged in order of increasing atomic mass.

  9. Newlands’ Octaves

  10. Newlands’ octaves • Newlands’ Octaves also failed • It was not valid for elements that had atomic masses higher than Ca. • The octaves mixed metals and nonmetals – for example he put iron (metal) in the same group as oxygen and sulfur (non-metals) • When more elements were discovered, such as noble gases He, Ne, Ar, they could not be accommodated in his table.

  11. Newlands’ importance • Concept of groups of eight carried over to modern table • Reinforced concept of periodicity from Dobereiner’s table

  12. Mendeleev and Meyer First useable periodic table (1869) Dmitri Mendeleev 1834 – 1907 Lothar Meyer 1830 – 1895

  13. Modern Periodic table • The table was organized by atomic mass (not atomic number) and by properties. • When organized by atomic mass, both found that the chemical properties repeated on a regular basis – “Periodicity” • Both scientists noticed holes in the periodic table where elements seemed to be missing.

  14. Modern Periodic Table • However, Mendeleev…. ….published first (1869, Meyer in 1870) ….corrected the atomic mass of several elements ….classified anomalous elements by properties rather than atomic mass – he said that future measurements would correct anomalous masses

  15. Modern Periodic table Ar and K Co and Ni  Te and I  Th and Pa ….accurately predicted the properties of missing elements Sc, Ga, and Ge \Mendeleev is remembered as the inventor of the modern periodic table, not Meyer.

  16. Moseley and Seaborg • Henry Moseley discovered the proton and atomic number in 1913 • Arranging the periodic table by atomic number eliminated the problem of anomalous atomic weights. • Glenn Seaborg came up with the idea of the actinide series – last major modification

  17. Structure of the table • Rows = periods • All elements in a period have the same valence shell and the same number of occupied energy levels • Columns = groups or families • All elements in a group have the same dot structure • All elements in a group have similar properties

  18. Coloring time! • Label the representative elements (s and p blocks) • The number of valence electrons of these elements increases by one moving left to right • Label the transition elements (d block) • Label the inner transition elements (f block) • Transition elements all considered to have two valence electrons

  19. More coloring! • Label the dividing line between metals (on the left) and nonmetals (on the right) • Label the following groups: • Column 1: Alkali metals (Li to Fr) • Column 2: Alkaline earth metals (Be to Ra) • Representative column 6: Chalcogens (oxygen family) • Representative column 7: Halogens (fluorine family) • Representative column 8: Noble gases (include helium)

  20. Even more coloring! • First row of inner transition metals: Lanthanide Series • Second row of inner transition metals: Actinide Series • Label the metalloids (B, Si, Ge, As, Sb, Te, Po) • Label the “other metals” (Al, Ga, In, Sn, Tl, Pb, Bi)

  21. periodic trends • Atomic radius decreases across a period • Result of increasing nuclear charge • Radius increases down a column • Valence electrons are in higher energy levels

  22. Periodic trends • Ionic radius: ions are atoms that have gained or lost an electron • Ions have a charge equal to # protons - # electrons • “Isoelectronic species” are atoms or ions with the same number of electrons • Na+, F- and Ne are isoelectronic (10 e-)

  23. Periodic trends • Radius of isoelectronic ions decreases left to right • Metals lose electrons and make + ions • Nonmetals gain electrons to make - ions

  24. Ionization energy • Ionization energy is the energy needed to remove the highest energy electron from a gaseous atom (makes a +1 ion) • Increases across a row due to increased nuclear charge • Decreases down a column – electrons in higher energy levels are easier to remove, and are shielded by inner shell electrons • Alkaline earth metals and nitrogen family are slightly higher than expected due to breaking symmetry of half-filled and completely filled shells

  25. First ionization energy

  26. Electronegativity • Electronegativity is an atom’s attraction for electrons in a bond O H H • Metals have low electronegativity, nonmetals high

  27. Electron affinity • The energy gained or lost when a gaseous atom of an element gains an electron • Sometimes defined as the energy required to detach an electron from a -1 charged ion • Values are generally positive (endothermic process) • Values generally increase from left to right, with more exceptions than ionization energy • Values for noble gases are very small or negative

  28. Electron affinity

  29. Properties of metals • Physical properties: • Shiny (Luster) • Flexible (malleability – can be hammered into a sheet) • Ductility (can be drawn into wire) • Conductors of heat and electricity • Hardness – transition metals are the hardest (Ti, Cr) though they are less hard than C (diamond) or B. Alkalis are soft; Alkaline earths are hard.

  30. Physical properties of metals • Most are solids – only mercury is a liquid • Magnetism • Diamagnetism: no unpaired electrons, unaffected or repelled by magnet • Paramagnetism: Unpaired electrons, attracted to magnet • Ferromagnetism: Ability to form a permanent magnet (Fe, Co, Ni, some inner transitions, some alloys and compounds of these metals) • Curie temperature: temperature at which a material loses its ferromagnetic properties (1388K for Co, 88K for Dy, 1043K for Fe, 627K for Ni)

  31. Metals • Chemical properties: • Tend to lose electrons and form + ions • The further left on the table, the more readily the metal loses electrons • Left side of table are better conductors, more malleable, etc. • Charge of ions depends on column; transition metals vary • More reactive metals are at the bottom of the group because of shielding • Form salts with non-metals • Many react with acids to give hydrogen gas and a salt

  32. Alkali metals in water http://www.youtube.com/watch?v=QSZ-3wScePM

  33. Transition metals • All considered to have two valence electrons, though many different valence states (charges on ions) can exist • Most tend to be hard and dense • Tc and all metals past Bi are radioactive; many others have radioactive isotopes as well

  34. Nonmetals • Physical properties: • Can be solids, liquid (Br only) or gas • Solids are generally hard • Gases are the Noble Gases and the seven diatomic gases (BrINClHOF: Br2, I2, N2, Cl2, H2, O2, F2) • Br2 is a volatile liquid, and I2 an easily sublimed solid • Many are colored (S is yellow, Cl pale green, Br orange, I purple, O pale blue) • Most are diamagnetic, except oxygen

  35. Chemical properties of nonmetals • Nonmetals tend to gain electrons and form negative ions • Will react with metals to form salts – for example, Fe2O3 (rust) • When forming compounds with each other, electrons are shared rather than transferred • Noble gases are monatomic and don’t react with anything except fluorine (only Xe, Kr and Rn)

  36. Metalloids • Properties are intermediate between metals and nonmetals • Poor conductors, semi-shiny solids • Tend to share electrons rather than transfer • Used in semiconductors

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