1 / 46

Chapter 10

Chapter 10. The Periodic Table. Early Attempts at Classification. Johann Dobereiner. German chemist – 1817 He found the properties of metals Ca, Ba and Sr were very similar. He also noted the atomic mass of Sr was about midway between those of Ca and Ba. He grouped elements into triads.

rozene
Download Presentation

Chapter 10

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. Chapter 10 The Periodic Table

  2. Early AttemptsatClassification

  3. Johann Dobereiner • German chemist – 1817 • He found the properties of metals Ca, Ba and Sr were very similar. He also noted the atomic mass of Sr was about midway between those of Ca and Ba. • He grouped elements into triads

  4. John Newlands • English Chemist – 1863 • He arranged elements in order of increasing atomic masses. He noted there was a repetition of similar properties every 8th element. • He placed the 49 known elements in 7 rows of 7 elements. • Law of Octaves

  5. Dmitri Ivanovich Mendeleev In Search of: The Perfect Periodic Table

  6. Dmitri Ivanovich Mendeleev • Born: Feb. 7, 1834 Toboslk, Siberia • Died: Feb. 2, 1907 St. Petersburg, Russia • Attended St. Petersburg University • Graduated top of his class • Studied in France & Germany

  7. Dmitri Ivanovich Mendeleev • 1866 – named Professor of Chemistry at the University of St. Petersburg • Published: “The Principles of Chemistry” called one of the best Russian chemistry texts ever written

  8. Dmitri Ivanovich Mendeleev • 1869: Published his first periodic table • Contained 69 elements • Based on increasing atomic wt. • Not immediately accepted • Kept improving his table

  9. Dmitri Ivanovich Mendeleev • Jan. 1871 Published a new table which left gaps for elements not discovered • His genius – assumed all elements had not yet been discovered. • Used his table to predict other elements • Predicted chemical and physical properties • Gallium 1875 Scandium 1879

  10. Ekasilicon (Es) Predicted Properties Atomic mass 72 High melting point Density 5.5 g/cm3 Dark gray metal Slightly dissolved in HCl Will for EsO2 Germanium (Ge) Actual Properties Atomic mass 72.61 M. Pt. 945 C D = 5.323 g/cm3 Gray metal Not dissolved by HCl Forms GeO2 Mendeleev’s Predictions

  11. Mendeleev • Within a decade his work was hailed and widely accepted • Shortly before his death he lost the Nobel Prize by 1 vote. • 1955 an element, mendelevium, 101, was named for him

  12. A Problem Develops • Tellurium and Iodine appear to misplaced. • Elements 52 and 53 • They don’t match other elements in their column.

  13. Modern Periodic Law • If they were switched according to properties then they would be out of order according to atomic mass. • Mendeleev explained the problem by saying their masses had not been correctly determined • Soon, other pairs showed the same type of reversal

  14. Why the Problem? • Henry Mosely – English Physicist • Nov. 23, 1887 to Aug. 10, 1915 • Used X-rays to determine the atomic numbers of elements • His calculations showed certain elements were missing. • Died as a combat foot soldier in Gallipoli, Turkey during WWI

  15. Periodic Law • The properties of the elements are a periodic function of their atomic numbers.

  16. Element for each Block • Symbol • Name • Atomic Number • Atomic Mass • Electron Distribution • Much more

  17. Periodic Terms • Horizontal Row: period or series • Vertical Column: group or family

  18. Sublevel s p d f Electron Capacity 2 6 10 14 Electron Distribution

  19. METALS • Good conductor of heat and electricity • Malleable • Ductile • Lustrous

  20. NONMETALS • Poor conductors of heat and electricty • Nonmalleable • Nonductile • More variability: almost all metals are solids at room temperature. Many nonmetals are gases or liquids at room temperature.

  21. Metalloids - Semimetals • Show a mixture of metallic and nonmetallic properties • Lie close to the “stair-step” line • Ex.: silicon, germanium, arsenic, antimony and tellurium

  22. Octet Rule • 8 electrons in the outer level render an atom unreactive. • Larger elements bond (gain, lose or share electrons) in order to achieve a noble gas like electron configuration (s2p6) – an octet in the last energy level.

  23. Octet Rule • Smaller atoms attempt to achieve a helium like (1s2) electron structure in order to reach increased stability. They are too small to achieve an octet but they can fill their only energy level. • These elements do not become He or a noble gas but do achieve a lower energy content and thus increased stability.

  24. Periodicity & Periods 1, 2 & 3 • Period 1 • H, hydrogen, 1s1 gas, left and above on some tables, reactive, nonmetal • He, helium, 1s2 noble gas, inert-does not react

  25. Period 2 • Li, lithium, 1s2 2s1 soft, silvery, active metal • Be, beryllium, 1s2 2s2 silver metal, less active than Li because it has to lose 2 electrons • B, boron, 1s2 2s2 2p1 black solid • C, carbon, 1s2 2s2 2p2 black solid, able to form multiple bonds, organic compounds

  26. Period 2 • N, nitrogen, 1s2 2s2 2p3 colorless gas • O, oxygen, 1s2 2s2 2p4 colorless gas • F, fluorine, 1s2 2s2 2p5 pale-yellow gas, chemically active, most active nonmetal • Ne, neon, 1s2 2s2 2p6 colorless, inert, noble gas, does NOT react

  27. Period 3 • Na, sodium, 1s2 2s2 2p6 3s1 soft, silver, active metal • Mg, magnesium, 1s2 2s2 2p6 3s2 soft, active metal but less active than Na. Must lose 2 electrons • Al, aluminum, 3s2 3p1 soft, silvery • Si, silicon, 3s2 3p2 dark colored • P, phosphorus, 3s2 3p3 solid nonmetal

  28. Period 3 • S, sulfur, 3s2 3p4 yellow solid • Cl, chlorine, 3s2 3p5 yellow-green gas, distinctive odor, active nonmetal • Ar, argon, 3s2 3p6 colorless, inert, noble gas

  29. Groups • Group 1 Alkali Metals • 6 similar & very active metals • Outermost – s1 • Most active metal -francium

  30. Groups • Group 2 alkaline earth metals • 6 active metals but less active than Group 1 because they have to lose 2 electrons • All end in s2 • Most active: radium

  31. Groups • Group 7 halogen family • 5 nonmetals • Outermost electrons – s2 p5 • Most active - fluorine

  32. Groups • Group 8 noble gas family • All are inert gases • Except for He (1s2), all have octets s2p6

  33. Transition Elements • Begin at Period 4 • 1 or 2 electrons in outer shell • Filling occurs in the d sublevel

  34. Begins at period 6 Located at the bottom of the table Lanthanide and actinide series Filling occurs in the f sublevel Rare Earth Elements

  35. Periodic Properties • Radii of Atoms: • As the principal quantum number increases, the size of the electron cloud increases.

  36. Periodic Properties • Electron Cloud: Gives shape & size Excludes other atoms Boundary is fuzzy & indefinite

  37. Periodic Properties • Atomic Radius: radius of an atom without regard to surrounding atoms. Measurements vary depending upon the method used.

  38. Atomic Radius is a Periodic Property • Group or Family: radius increases from top to bottom. It does so because of increasing principal quantum no., energy level.

  39. Atomic Radius is a Periodic Property • Period or Series: atomic radius decreases going left to right in a series as the atomic number increases. Electrons are being added to the same energy level while positive charges are being added to the nucleus. As both positive and negative charges increase, they attract each other and reduce the size of the electron cloud.

  40. Ionic Radii • Metallic Ions: located left and center of the table are formed by the loss of electron(s). Positive ions are smaller than the atoms from which they came. • Nonmetallic Ions: located on the right side of the periodic table. Formed by the gain of electron(s) and are larger than the atoms from which they came.

  41. Ionization Energy • The energy required to remove an electron from an atom (Kj/mole) • First Ionization Energy: the energy required to remove the first electron from an atom. • I.E. tends to increase in a period L to R and decreases in a group T to B

  42. Metals, Nonmetals and IE • Metals: tend to have a low first IE • Nonmetals: tend to have a high first IE

  43. Factors Affecting IE • Nuclear Charge: in a period the larger the nuclear charge the greater the IE • Shielding: the greater the shielding effect the lower the IE • Radius: the greater the radius the lower the IE • Sublevel: atoms with full or half-full sublevels have a higher IE. More stable.

  44. IE is a Periodic Property • L to R in a period IE increases. Why? • T to B in a group IE decreases. Why?

  45. Multiple Ionization Energies • As more electrons are removed IE increases. • When an octet is broken there is a huge jump in IE.

  46. Electron Affinity • EA is the tendency of an atom to attract electrons. • Metals have low EA. Why? • Nonmetals have high EA. Why? • Periodic Property: • L to R in a series EA increases. Why? • T to B in a group EA decreases. Why?

More Related