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The Periodic Chart

The Periodic Chart. From then to Now. The History. 1669-Henning Brand discovered Phosphorus 1680-Robert Boyle rediscovered Phosphorus 1789-Lavosier wrote the 1st chemistry text 1809-There were 47 known elements 1862-Beguyer deChancourtois noticed periodicity

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The Periodic Chart

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  1. The Periodic Chart From then to Now . . .

  2. The History • 1669-Henning Brand discovered Phosphorus • 1680-Robert Boyle rediscovered Phosphorus • 1789-Lavosier wrote the 1st chemistry text • 1809-There were 47 known elements • 1862-Beguyer deChancourtois noticed periodicity • 1863-Newlands classified 56 elements into 11 groups, octaves • 1869-Mendeleev created a table and was able to predict the existence of 2 new elements • 1869-Meyer also created a table, but did not get the credit for it • 1900-Moseley developed the Periodic Law • 1944-Seaborg proposed the Actinide series

  3. History of the periodic table • In the 1700’s only 30 elements were identified • Dobereiner in the 1800’s noticed certain elements could be grouped into sets of 3 called triads • Dobereiner--triads

  4. TRIAD PROPERTIES • Properties similar – Group 1 are soft metals • Reactiveness similar – Group 1 are very reactive with water • Middle element value is average of one above and one below • Triad—3 elements with similar properties one value is an average of the other 2

  5. PROPERTIES OF TRIADS

  6. Newlands – mid 1800’s • Now 49 elements • Noticed that when arranged by increasing mass, every 8th element had similar properties • Called law of octaves Newlands -- octaves

  7. MENDELEEV VS MEYER • Both made discoveries at the same time but Mendeleev was the first to publish them • Wrote names and properties on cards and arranged them in various ways: In increasing mass In repetitive properties • Both couldn’t be done at the same time • Decided putting them in order of repetitive properties was more important

  8. DISCREPANCIES • In order to put the elements in similar groups according to properties, some of the masses were out of order • Thought that the atomic masses were wrong

  9. MENDELEEV’S PERIODIC TABLE • When he put elements in order according to their properties without regard to their masses, some elements seemed to be missing • He predicted the existence of these missing elements and when discovered, they fit perfectly into his pattern • But – Mendeleev was not entirely correct The atomic masses, when recalibrated, were not incorrect. • This left some atomic masses out of order on his periodic table

  10. Mendeleev’s notes

  11. Mendeleev’s 1869 Periodic Table

  12. LATE 1800’S MOSELY’S PERIODIC TABLE • Developed the idea of atomic #’s • Assigned one to each element based on the # of protons in their nucleus • Arranged elements according to the number of protons instead of mass • Now, elements are in a numerical repetitive order as well as grouped according to their properties • Since masses aren’t figured into arranging the periodic table, it’s ok for them to be out of order

  13. The Periodic Law • The periodic properties of the elements are functions of their atomic number. • In other words, the elements are arranged on the basis of their ground state electron configuration

  14. Periodic Table of 1944

  15. The Modern Periodic Table

  16. Vertical Columns • The vertical columns are arranged in groups or families. • They are numbered from left to right • Elements in a group have the same electron structure in their outer subshell (valence electrons)

  17. Electron Review • An electron shell, also known as a main energy level, is a group of atomic orbitals with the same value of the principal quantum number n. • Electron shells are made up of one or more subshell, which have orbitals with the same angular momentum quantum number l. (1 of s, 3 of p, 5 of d and 7 of f orbitals)

  18. States with the same value of n are related, and said to lie within the same electron shell. • Example: 1s22s22p6 • 1s2 and 2s22p6 are in the same electron shell • States with the same value of n and also l are said to lie within the same electron subshell. • Example: 1s22s22p6 • 1s2 are in the same electron subshell • 2s2 are in the same electron subshell • 2p6 are in the same electron subshell

  19. Electron shells make up the electron configuration. • It can be shown that the number of electrons that can reside in a shell is equal to 2n2. • Shells and subshells are defined by the quantum numbers. • In large atoms, shells above the second shell overlap (Aufbau principle)

  20. Valence Shell • The valence shell is the outermost shell of an atom, which contains the electrons most likely to participate in a chemical reaction with other atoms or to determine chemical properties. • Electrons in the valence shell are referred to as valence electrons.

  21. Let’s see ... Group 18 Ne =1s22s22p6 Ar = 1s22s22p63s23p6 Kr = [Ar]4s23d104p6 Xe = [Kr]5s24d105p6 Rn = [Xe]6s24f145d106p6

  22. Further Breakdown s-orbital elements d-orbital elements p-orbital elements f-orbital elements

  23. Horizontal Rows • The horizontal rows are the periods. • The periods are numbered from the top down. • Elements in the same period have the same principal energy level

  24. Let’s see ... Period 2 (Period n) Li = 1s22s1 Be = 1s22s2 B = 1s22s22p1 C = 1s22s22p2

  25. Group Names • Groups 1-2 and 13-18 (except Hydrogen) are the main group elements (also known as the representative elements). • Groups 3-12 are the transition metals

  26. Specific Group Names • Group 1: alkali metals • Group 2: alkaline earth metals • Group 11: coinage metals (not IUPAC approved) • Group 15: pnictogens (not IUPAC approved) • Group 16: chalcogens • Group 17: halogens • Group 18: noble gases

  27. Period Identifications • The elements in the 1st f-period are the Lanthanide series. • The elements in the 2nd f-period are the Actinide series

  28. Group 1: Alkali Metals • Hydrogen is NOT included in Group 1 • Metals that react with water to make an alkaline solution (basic) • Highly reactive, soft (less than 1 on the Mohs scale), and conductive

  29. Group 1 Electrons • Not found in their elemental form but in compounds • example: NaCl, KOH • There is only 1 valence electron. (ns1) • If the one electron is lost, it will be stable

  30. Mohs Hardness Scale • The scale used to describe the hardness of a material is the Mohs Hardness Scale • The scale is from 0-10 (softest to hardest) • example: Talc is 1 on the Mohs scale and the Diamond is 10

  31. Group 2: Alkaline Earth Metals • The alkaline earth metals are silvery colored, soft, low-density metals, which react readily with halogens to form ionic salts, and with water, to form strongly alkaline hydroxides. • Highly reactive, but not as reactive as alkali metals, usually found as compounds not in elemental form

  32. Alkaline Earth Electrons • There are 2 valence electrons. (ns2) • It takes more energy to lose 2 electrons than it does to lose only one (like the alkali metals)

  33. Valence Electrons of Groups 13-18 • Group 13 = ns2np1 • Group 14 = ns2np2 • Group 15 = ns2np3 • Group 16 = ns2np4 • Group 17 (halogens) = ns2np5 • Group 18 (noble gases) = ns2np6

  34. Group 17: The Halogens • Halogens are highly reactive non-metals. • Only 7 valence electrons (just one short of a full and stable valence shell) so they want to gain an electron • Reactive with most metals to form salts

  35. Group 18: Noble Gases • Have a full set of electrons (n2p6) • Low chemical reactivity and so they are very stable

  36. Hydrogen • Hydrogen is in a class by itself because it is the most common element in the Universe! • Hydrogen only has one proton and one electron and can react with almost anything

  37. Transition Metals • Groups 3-12 (d-block) • Do NOT have identical electron configurations in the outer shell. Why? • The Lanthanide and Actinide series are contained within the d-block and have f-orbitals

  38. Lanthanide & Actinide • Lanthanide are the rare earth series from atomic #58 to #71 • shiny metals with similar reactivity to alkaline • Actinide are from atomic #89 to #103 • nuclei are unstable, radioactive • As you move to the right, electrons are filled in the f-orbital

  39. Metallic Character • Approximately 2/3’s of the elements are metals. • See periodic chart • Metals have unique properties: • luster: mirror like shine that reflects light • conductivity: ability to conduct heat or electricity • malleable: ability to be rolled or hammered • ductile: ability to be drawn into wire

  40. Alloys • Metals that are mixed with other metals to form a stable compound are called alloys • example: Brass is Copper and Zinc • example: Steel is Iron, Tin, Nickel, Lead, etc.

  41. Nonmetals • Poor conductors of heat and electricity • Not malleable • Many are gasses One is liquid – Br Some are solids (brittle and dull) • More electrons in outer level • Form negatively charged ions

  42. METALLOIDS Metalloids – have properties of both metals and nonmetals • On the stairstep; exclude Aluminum and Polonium

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