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The Periodic Chart

Explore the fascinating history of the periodic table from its inception in 1669 to the modern version in 1944. Learn about key milestones, influential chemists, and the basis of the Periodic Law. Discover the structure, organization, and significance of elements in the periodic table.

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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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