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The Periodic Law

The Periodic Law. History of the Periodic Table. * Antoine Lavoisier (France, 1789) Earned reputation as “father of chemistry” Established a common naming system of compounds and elements. First to organize elements Grouped them into four

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The Periodic Law

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  1. The Periodic Law

  2. History of the Periodic Table *Antoine Lavoisier (France, 1789) • Earned reputation as “father of chemistry” • Established a common naming system of compounds and elements. • First to organize elements • Grouped them into four categories:Gases, nonmetals, metals and “earths” (elements that could not be chemically separated at the time.)

  3. History of the Periodic Table Dmitri Mendeleev (Russia, 1869) 1.Placed elements in groups in which they shared similar properties which resulted in order of increasing atomic mass with a few exceptions. 2. Ex: If placed solely by atomic mass, iodine was not in group with chemically similar elements. 3. Left gaps for not-yet-discovered elements and predicted their properties: gallium, germaniun & scandium.

  4. History of the Periodic Table • Henry Mosely • Moseley (1911) modified the table by organizing elements in order of increasing atomic numbers. • Periodic Law: The phsical and chemical properties of the elements are the periodic functions of their atomic numbers. Glenn Seaborg (UC Berkeley, 1944) • Formed Actinide Series just like that of the lanthanides (#58-71)

  5. Basics of the Periodic Table periodic: a repeating pattern table: an organized collection of information • period: horizontal row on the P.T. • Designates e- energy levels groupor family: vertical column on the P.T. Periodic Table: an arrangement of elements in order of atomic number; elements with similar properties appear at regular intervals (are in the same group)

  6. e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- e- Electron Structures of Atoms nucleus 1st energy level 1st Period: Hydrogen (#1) Helium (#2) 2nd energy level 2nd Period: Lithium (#3) Neon (#10)

  7. S block elements: Group 1 & 2 • Chemically reactive metals, group 1 more reactive than group 2. Group Config: ns 1-2 • Alkali metals:silvery appearance, soft enough to cut with a knife, not found in nature as free elements. H shares e-config but not properties. • Alkaline-earth metals:harder, denser, stronger, and have a higher melting point than group 1. Too reactive to be found uncombined in nature. He shares e- config but not properties.

  8. p block elements: Group 13-18 Includes all the three types of elements: metals, non metals and metalloids. Group Config: ns2 np 1-6 Includes Halogens: most reactive of the nonmetals. React vigorously with most metals to form salts. P block metals are generally harder & denser than s block but softer & less dense than d block metals.Found in nature solely as compounds except for bismuth.

  9. d block elements: Group 3-12 • Transition Elements: metals with typical properties; good conductors, high luster. • Less reactive than s block, many existing in nature as free elements. • Electrons added to the d sublevel of the preceding energy level (n-1). Group configuration: (n-1)d1-10ns 0-2 • Some deviations from orderly d sublevel filling occur in group 4-11(s electrons jumping to d sublevel)

  10. f-block elements • F-block elements are wedged between groups 3 and 4 in the sixth and seventh period, consisting of lanthanides and actinides • Most elements are radioactive • Trans Uranium elements are all synthetic • Group Config: ns 0-2 (n-1) d 0-1 (n-2)f 1-14

  11. Covalent Radius for Covalently Bonded Atoms: half the distance between the nuclei of two covalently bonded atoms • F-F bond length is 144 pm, so F covalent radius is 72 pm. • H-F bond length is 109 pm, so H covalent radius is 37 nm. Atomic Radius for Elements like the Noble Gases • Ar atomic radius is 131 pm Metallic Radius for Metals • Al metallic radius is 143 pm. atomic radius:

  12. 6.3 Trends in Atomic Size • The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined.

  13. 6.3 Trends in Atomic Size

  14. 6.3

  15. Ionic Radii Cation=positive ion, Anion=negative ion • Forming a cation by losing electron(s) leads to a decrease in atomic radius, a smaller electron cloud. • Forming an anion by electron(s) leads to an increase in atomic radius, less pull from the nucleus & there is more repulsion between the greater number of electrons.

  16. 6.3 Cations

  17. 6.3 Anions

  18. Across a period atoms become smaller. Down a group atoms become larger.

  19. Ionization Energy • Amount of energy required to remove an e from a neutral atom in its gaseous state. • First Ionization Energy • A(g)  A+(g) + e- • Second Ionization Energy • A+(g)  A2+(g) + e- • Third Ionization Energy • A2+(g)  A3+(g) + e-

  20. Picture of IE trend I.E. increases across a period and decreases down a group.

  21. 6.3 Trends in Ionization Energy

  22. 6.3 Trends in Ionization Energy

  23. Electron Affinity • Amount of energy released when an e is added to a gaseous atom in its neutral state. • First Electron Affinity A(g) + e- A-(g) • Second Electron Affinity A-(g) + e- A2-(g) • Third Electron Affinity A2-(g) + e- A3-(g)

  24. Electronegativity • A measure of the ability of an atom in a chemical compound to attract electrons. • Fluorine, the most electronegative element, is arbitrarily assigned a value of 4.0. Values for other elements are calculated in relation to this. • Tend to increase across a period • Tend to decrease down a group or remain about the same. • If an element does not form a compound, some noble gases, will not have a value.

  25. Trends in Electron Affinity and Electronegativity • Both electron affinity and electronegativity increase from L to R across a period. • Both electron affinity and electronegativity decrease down a group.

  26. Two Factors Used to Explain Trends The principal energy level • All other factors being equal, increased n for the orbitals in which electrons are found means increased size of orbitals, which leads to decreased attraction for electrons from the nucleus.

  27. Effective Nuclear Charge • Effective charge is the approximate net nuclear charge felt by the highest energy electrons. • All other factors being equal, increased effective charge means increased attraction for electrons, which leads to decreased size of orbitals. • Effective charge depends upon two factors: • Total nuclear charge: # of protons (greater the total nuclear charge, higher the attraction felt by electrons) • # of shielding electrons (e present in between the nucleus and the valence shell electrons, the higher the number of shielding electrons, the lesser is the effective nuclear charge)

  28. SHIELDING: • The net nuclear charge felt by an outer electron is substantially lower than the actual nuclear charge. the outer electrons are shielded from the full charge of the nucleus by the inner electrons, which is called shielding effect.

  29. Explanation of Trends (1)

  30. Explanation of Trends

  31. Explanation of Trends

  32. Bibliography • http://www.homeoint.org/morrell/british/originidea.htm • http://www.chemsoc.org/viselements/pages/history.html • http://www.library.upenn.edu/etext/smith/d/dobereiner.html • http://www.library.upenn.edu/etext/smith/n/newlands.html • http://www.ulb.ac.be/sciences/cudec/ressources/Mendeleev.gif • http://intro.chem.okstate.edu/1314F00/Lecture/Chapter7/ATRADIID.DIR_PICT0003.gif • http://scidiv.bcc.ctc.edu/wv/4/0004-000-IE.GIF

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