Notes 11 14 08 To 11 21 08

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  • 1. The Periodic Table Unit 3 Chapter 4 pp 114-131
  • 2. Ok, so now what?
    • Recall: Element cannot be broken down any further.
    • In the late 1800’s, we had “discovered” about 60 elements.
    • We had no idea of the structure of atoms.
    • Many scientists attempted to put order to the rapidly expanding list of elements.
  • 3. Off to the Newlands…
    • Arranged known elements in a table by atomic mass in 1863.
    • Noticed a repeating pattern every 8 th element in 1865.
    • Law of Octaves – chemical properties repeat every 8 elements.
    • Was laughed at by peers.
    John Newlands (1837-1898)
  • 4. Newland’s Flaw
    • Knew nothing of subatomic particles
    • His table mixed some obviously different elements (like oxygen and iron)
  • 5. The Mad Russian
    • Produced a more orderly table independent of Newlands’ work in 1869 (also used atomic mass).
    • Left blanks for yet-undiscovered elements.
    • Predicted properties of Ga, Sc, and Ge (disc. 1875, 1877, & 1886).
    • Credited with the Periodic Table.
    Dmitri Mendeleev (1834-1907)
  • 6. Russian Roulette
    • Mendeleev’s table had a few problems.
    • Based on atomic mass, had to switch a few elements (e.g. Tellurium and Iodine) to keep reactivities in order.
    • Many believed he predicted too many elements (we had 63 already!!!).
    • Still, this is what we used for half a century.
  • 7. 45 Years Later…
    • Rearranged table according to electronic charge in 1914.
      • Became the # of protons after 1918.
    • Noticed his new table had spots for #’s 43, 61, 72, & 75.
    • Produced the modern periodic table we know today.
    • Enlisted in the army’s Royal Engineers when WWI broke out.
    Henry Moseley (1887 – 1915)
  • 8. Moseley’s New Order
    • Gave experimental meaning to atomic number.
    • Gave reason for Tellurium and Iodine being switched.
    • Moseley’s technique easily separated rare earth metals.
      • Plagued chemists for years and years.
    • Predicted how many elements remained between others.
    • (e.g. 13 elements between La and Lu)
  • 9. Moseley’s Lost Nobel
    • Many thought he should have won Nobel Prize.
    • It’s only given to the living…he was shot in the head by a sniper in Gallipoli.
    • Bohr (1962): "You see actually the Rutherford work [the nuclear atom] was not taken seriously. We cannot understand today, but it was not taken seriously at all. There was no mention of it any place. The great change came from Moseley."
    • British barred scientists from enlisting for combat.
  • 10. Elements Everywhere
    • Based on increasing number of protons, we now have a complete periodic table.
    • Will not find any lower elements, can only go up (118 so far).
    • Create new elements by smashing smaller atoms together:
    3 Neutrons 294 118 Uuo 48 20 Ca 249 98 Cf
  • 11. Periodicity
    • In order by atomic number (# of Protons)
      • H has 1 p + , U has 92
    • Arranged in Rows and Columns
    • Rows = Periods
      • Pd 3 = Na, Mg, Al, Si, P, S, Cl, Ar
    • Columns = Groups or Families
      • Group 1 = H, Li, Na, K, Rb, Cs, Fr
      • Have similar properties
      • (e.g. Form hydroxides: LiOH, NaOH, KOH, etc)
  • 12. Division of Labor
    • Different types of elements are found on different parts of the table:
      • Metals to the left (majority of the elements).
      • Nonmetals to the right (18 elements).
      • Metalloids found on a “staircase” dividing metals and nonmetals (7 elements).
      • Lanthanoids & Actinoids (metals) added to bottom to make table manageable.
  • 13. Metals
    • Lustrous (shiny)
    • Malleable (can be pounded into thin sheets)
    • Ductile (can be pulled into wires)
    • Conductive
      • Heat and electricity
    • Form solid oxides when burned.
    • Tend to react with acids to form Hydrogen gas.
  • 14. Nonmetals
    • Wide range of properties
    • Tend to:
      • Be Dull
      • Be Brittle (when solid)
      • Be Insulators
      • Form gaseous oxides
      • Not react with acids
      • Have lower melting & boiling points.
    Bromine
  • 15. Metalloids
    • Also called “semi-metals” or “staircase elements.”
    • Combination of properties of metals and nonmetals.
    • Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, & Polonium
    • Many exhibit semi-conducting behavior.
  • 16. Groups/Families
    • Alkali Metals
    • Alkali Earth Metals
    • Halogens
    • Noble Gases
    • Transition Metals
    • Inner Transition Metals
      • Lanthanoids (Rare Earths)
      • Actinoids
  • 17. s-Block Elements
    • At least 1 e - in s orbital (ns x )
    • Groups 1 & 2
      • Alkali Metals
      • Alkaline Earth Metals
    • Reactivity increases as you go down
    • All are metals, except H & He
    • Helium is technically an s-block, but placed with Noble Gases b/c of reactivity
  • 18. p-Block Elements
    • At least 1 e - in p orbital (np x )
    • Groups 13-18
    • Nonmetals at top, gradually transitioning into metals
    • All nonmetals and metalloids are p-block elements (excl H & He)
    • Some metals (Al, Ga, In, Sn, Tl, Pb, Bi)
  • 19. d-Block Elements
    • At least 1 e - in d orbital (nd x )
    • Groups 3-12
      • Transition Metals
    • Very little similarities w/in group
    • All are metals
    • Most form multiple ions (charged atoms)
  • 20. f-Block Elements
    • At least 1 e - in f orbital (nf x )
    • Oddballs
      • Lanthanoids start with #57, La
      • Actinoids start with #89, Ac
    • The groups are NOT similar up & down
    • All are metals
    • Lanthanoids (4f) are natural, most Actinoids (5f) are man-made
  • 21. Representative Elements
    • Exhibit nearly perfect periodicity.
      • All members of these groups behave as expected.
    • Groups on the outside of the table:
      • Alkali Metals (Group 1)
      • Alkaline Earth Metals (Group 2)
      • Halogens (Group 17)
      • Noble Gases (Group 18)
  • 22. Alkali Metals
    • Group 1 (excluding hydrogen) [ns 1 ]
    • Soft, lustrous, oxidize when exposed to air.
    • Difficult to isolate – never found in nature.
    • React (violently) with water to form a base.
    • React with chlorine to form a salt with a 1-to-1 ratio:
      • LiCl
      • NaCl
      • KCl
      • RbCl
      • CsCl (also FrCl)
  • 23. Alkaline Earth Metals
    • Group 2 [ns 2 ]
    • Harder & Denser than Alkali Metals.
    • Lustrous, oxidize slowly when exposed to air.
    • React with water or steam to form a base.
    • React with chlorine to form a salt with a 1-to-2 ratio:
      • BeCl 2
      • MgCl 2
      • CaCl 2
      • SrCl 2
      • BaCl 2
      • RaCl 2
  • 24. Halogens
    • Group 17 [np 5 ]
    • Nonmetals
    • Gases (F, Cl), liquid (Br), and solids (I, At)
    • Name means “salt former.”
    • React with sodium to form a salt with a 1-to-1 ratio:
      • NaF
      • NaCl
      • NaBr
      • NaI
      • NaAt
  • 25. Noble Gases
    • Group 18 [np 6 ]
    • Unreactive Gases – colorless, odorless.
    • Some of the last natural elements to be discovered.
    • Once called “Inert Gases.”
    • Monatomic in Nature
  • 26. Non-representatives
    • Other families have similarities, but do not behave exactly as expected
      • Groups 13-16, start with Boron – Oxygen
      • More differences than similarities
    • Others are lumped together for other reasons
      • Transition Metals
      • Lanthanoids
      • Actinoids
  • 27. Transition Metals
    • Groups 3 to 12 [nd x ]
    • Central portion of the PT.
    • Behavior and appearance vary.
    • Variable oxidation state (charge).
    • Different oxidation states can produce different colors.
    • Often used to make pigments.
    Co +2 Cr +6 Cr +6 Ni +2 Cu +2 Mn +7
  • 28. Lanthanoids
    • 1 st Row on Bottom of table [4f x ]
    • AKA Lanthanides & Rare Earths
    • Not so rare (Ce 25 th most abundant)
    • So similar, very difficult to separate – remember Moseley?
    • Most deflect UV – used in sunglasses
    • Shiny, silvery white, soft, react violently with most nonmetals, tarnish in air
  • 29. Actinoids
    • 2 nd Row on Bottom of table [5f x ]
    • AKA Actinides
    • All are radioactive
    • Not as similar as the Lanthanoids
    • Only Th and U are common in nature
    • Most are man-made
      • Nuclear fallout
      • Particle colliders
  • 30. State of the Union
    • Reacted State:
      • When elements are combined with other elements to form compounds
      • Most common state
    • Elemental State:
      • When elements are uncombined
      • Most elements are Monatomic (one atom)
      • Some are always Diatomic (two atoms)
      • A few are Polyatomic (>2 atoms)
  • 31. Diatomics
    • 7 elements always form diatomic molecules when they are isolated in their elemental state…ALWAYS!
    • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, & Iodine
    • These, you gotta memorize!
    • Luckily, Mr. Brinclhof is here to help!
    Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2
  • 32. Another Way
    • The rule of “7”
    • Diatomics form a “7” on the Periodic Table excluding H
    2 2 2 2 2 2 2 I Te Sb Sn Br Se As Ge Cl S P Si F O N C H
  • 33. The Oddballs
    • Sulfur is normally found as S 8
    • Selenium also forms Se 8
    • Phosphorus forms P 4
  • 34. Allotrope
    • When an element can be found in more than one form
    • Several elements have different allotropes, but most often cited is Carbon
    • Carbon has 3 common allotropes
      • Amorphous – Random arrangement of C atoms
      • Graphite – Hexagonal arrangement in sheets
        • Conducts electricity!
      • Diamond – 3-D network solid
  • 35. Allotropes of C Amorphous C Diamond (Network Solid) Graphite (Sheets)
  • 36. Trends in the Periodic Table
    • Several trends appear once we have the elements in order
      • Atomic Radius
      • Ionization Energy
      • Electronegativity
      • Reactivity
  • 37. Ray “D” Eye
    • Atomic Radii DECREASE from left to right
    • They INCREASE from top to bottom
    Na is bigger than Ar (223 pm) (88 pm) I is bigger than F (132 pm) (57 pm) Na Mg Al Si P S Cl Ar I Br Cl F
  • 38. Fluorine says, “Mine!”
    • Electronegativity is a measure of how badly an element wants to gain an electron
    • It INCREASES from left to right
    • It DECREASES from top to bottom
    Ne -- F 3.98 O 3.44 N 3.04 C 2.55 B 2.04 Be 1.57 Li 0.98 I 2.66 Br 2.96 Cl 3.16
  • 39. F has Codependency Issues
    • Ionization Energy is the amount of energy required to remove an electron.
      • It INCREASES from left to right
      • It DECREASES from top to bottom
    Na needs less NRG than Cl (496 kj/mol) (1256 kj/mol) F needs more NRG than I (1681 kj/mol) (1008 kj/mol)
  • 40. Major Trends in a Nutshell Atomic Radius Decreases Electronegativity Increases Ionization Energy Increases Fr F We usually ignore the Noble Gases
  • 41. Reactivity
    • Most reactive Metals are farther down and to the left
    • Most reactive Nonmetals are higher and to the right
  • 42. Tidbits
    • Hydrogen by far most abundant (4 out of every 5 atoms in universe)
    • Atoms in the Elemental state tend to be more dangerous/poisonous than those in the Reacted state – Exceptions: Cu & Pb
    • Oddo-Harkins Rule: even #’d elements more common than odd ones (protons apparently like to be paired up).