1 / 33

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

kailey
Download Presentation

The Periodic Law

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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.(Periodicity) 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. • Modified Periodic Law: The physical 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: Reactive metals,Form +1 ions • Alkaline-earth Less reactive than group 1 but still reactive,Form +2 ions

  8. p block elements: Group 13-18 Includes all the three types of elements: metals, non metals and metalloids. Form +3 -1 ions. Group Config: ns2 np 1-6 Includes Halogens: most reactive of the nonmetals. React vigorously with most metals to form salts. Note: Groups can be named several ways. Also consider the group A 1-8 and B 1-8 instead of 1-18

  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 1-2 Most form +2 ions. • 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 • Atomic numbers greater than 92 (the atomic number of uranium). • Group Config: ns 0-2 (n-1) d 0-1 (n-2)f 1-14

  11. Factors Used to Explain Trends 1. Principal Energy Level • All other factors being equal, increased n for the orbitals in which e- are found means increased size of orbitals, which leads to decreased attraction for e- from the nucleus.

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

  13. 3. 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.

  14. 4. Atomic Radius • The atomic radius is one half of the distance between the nuclei of two atoms of the same element when the atoms are joined.

  15. Trends in Atomic Size

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

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

  18. Cations

  19. Anions

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

  21. I.E. increases across a period and decreases down a group.

  22. Trends in Ionization Energy

  23. 7. ELECTRON AFFINITY • Change in energy when an e is added to a gaseous atom in its neutral state. • The more negative the value, the greater the quantity of energy released. • 1st Electron Affinity A(g) + e-A-(g) • 2nd Electron Affinity A-(g) + e-A2-(g)

  24. Electron Affinity cont. • Noble gases have full valence shells, are very stable, and do not want to add more e- • Atoms with full subshells, or ½ filled also have more positive e- affinities (e- are less attracted) than the elements around them.

  25. 8. 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.

  26. Explanation of Trends

  27. Explanation of Trends

  28. Explanation of Trends

  29. 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 CHAPTER 5 ANIMATIONS more on radii periodic trends References Believe created by Travis Hambleton and modified at MVHS

More Related