1 / 53

AP Chapter 7

AP Chapter 7. Periodic Properties of the Elements HW: 25 31 33 35 36 37 39 41 47 53 63 67 69 73. 7.1 – Development of the PT. -Newlands = Law of Octaves – Noticed that when organized by atomic mass, similar properties appear every 8 th element

Download Presentation

AP Chapter 7

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. AP Chapter 7 Periodic Properties of the Elements HW: 25 31 33 35 36 37 39 41 47 53 63 67 69 73

  2. 7.1 – Development of the PT -Newlands = Law of Octaves – Noticed that when organized by atomic mass, similar properties appear every 8th element -Mendeleev – Father of the PT. Made the first accepted PT. Based on periodicity and atomic mass. Could predict missing elements. -Meyer – Did same work as Mendeleev -Rutherford – Gold Foil Experiment – Estimated # of positive charges in the nucleus -Moseley – Determined atomic #

  3. Development of Periodic Table Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

  4. Element Discoveries: Page 256

  5. Development of Periodic Table • Elements in the same group generally have similar chemical properties. • Properties are not identical, however.

  6. Periodic Classifications -Representative Elements – s and p block. AKA “main group” elements -IUPAC numbers groups as 1 – 18 -American style uses IA-IIA and IIIA – VIIIA -Noble Gases – Group 18 -Transition Metals – d block (Groups 3 – 12) - # System - We use IUPAC, not American

  7. 7.2 - Effective Nuclear ChargeIn 10th Grade we called this “Strength of the Nucleus” • In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. • The nuclear charge that an electron experiences depends on both factors. • Nuclear Charge = The amount of pull the electron is experiencing.

  8. Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z−S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons. (In same period, Z increases across the PT and # shielding remains the same, so Zeff INCREASES across the PT) (Down a group, the valence e-s are farther away, so Zeff decreases down a group) Mg = 12-10 = 2

  9. 7.3 - Atomic Radius The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

  10. Atomic Radius Bonding atomic radius tends to… …decrease from left to right across a row due to increasing Zeff. …increase from top to bottom of a column due to increasing value of n

  11. 7.3 - Ionic Radius • Ionic size depends upon: • Nuclear charge. • Number of electrons. • Orbitals in which electrons reside.

  12. Ionic Radius • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions are reduced.

  13. Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

  14. Sizes of Ions • Ions increase in size as you go down a column. • Due to increasing value of n.

  15. Ionic Radius • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing atomic #. =[He] =[Ne]

  16. 7.4 - Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. Always requires energy (+). • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc.

  17. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

  18. Trends in First Ionization Energies • Down a group, less energy is required to remove the first electron (lower IE). • For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

  19. Trends in First Ionization Energies • Generally, across a period, it gets harder to remove an electron (higher IE). • As you go from left to right, Zeff increases.

  20. Trends in First Ionization Energies However, there are two apparent exceptions to this trend.

  21. Trends in First Ionization Energies • The first occurs between Groups 2 and 13. • Group 13 – The electron removed from p-orbital. Group 2 it is removed from an s-orbital • Electron farther from nucleus, so easier to remove • s electrons are stable as a pair

  22. Trends in First Ionization Energies • The second occurs between Groups 15 and 16. • Group 15 has one e- in each orbital of the p subshell. • Group 16 has a pair and two singles in the p-subshell. Repulsion between the electrons in the same orbital makes it easier to remove.

  23. Electron Affinity Energy change accompanying addition of electron to gaseous atom: Cl + e− Cl− -If EA is + = the process is endothermic and energy must be added to make it occur -If EA is - = the process is exothermic and energy is released when the process occurs (this is different than what your book says!)

  24. Trends in Electron Affinity -In general, electron affinity becomes more exothermic (easier to add an e-) as you go from left to right across a period. -Becomes more endothermic as you go down a group.

  25. Trends in Electron Affinity There are two discontinuities in this trend.

  26. Trends in Electron Affinity • The first occurs between Groups 1 and 2. • For Group 2, the added electron must go into a p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons, so requires more energy than adding to Group 1 atom.

  27. Trends in Electron Affinity • The second occurs between Groups 14 and 15. • Group 15 has all half-filled p orbitals in the p-subshell. • Extra electron must go in to form a pair, which has repulsion, which takes more energy, so overall the process is less exothermic than it was for Group 14.

  28. Valence Electrons and Configurations Valence Electrons – Outer shell electrons. Involved in bonding. Configurations of ions – Gain (anions) or Lose (cations) to obtain the configuration of a noble gas. Ex: Na F (Isoelectronic = same number and location of the electrons)

  29. 7.6 - METALLIC PROPERTIES

  30. Metals versus Nonmetals

  31. Metals versus Nonmetals • Metals tend to form cations. • Nonmetals tend to form anions.

  32. Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

  33. Metals • Compounds formed between metals and nonmetals tend to be ionic. • Metal oxides tend to be basic.

  34. Nonmetals • Dull, brittle substances that are poor conductors of heat and electricity. • Tend to gain electrons in reactions with metals to acquire noble gas configuration.

  35. Nonmetals • Substances containing only nonmetals are molecular compounds. • Most nonmetal oxides are acidic.

  36. Metalloids • Have some characteristics of metals, some of nonmetals. • For instance, silicon looks shiny, but is brittle and fairly poor conductor.

  37. 7.7 – Group Trends and the Variations in Chemical Properties • (sorry about the grammar…) The jump from Period 2 to 3 is “more different” than the rest of the group trends because Period 1 and 2 elements do NOT have the option for d orbitals • Ex: Li more different than Na, than Na is to K • Diagonal relationships: There is a similarity between elements in different group and period (ex: Li and Mg are similar; Be and Al; B and Si, etc.) • Groups 13 – 16 have more variation in properties because of the “steps” separating metals and nonmetals

  38. Group 1 Hydrogen: Can be +1 or -1; Forms both molecular (water) and ionic substances (NaH) Alkali Metals: • Soft, metallic solids. • Name comes from Arabic word for ashes.

  39. Alkali Metals • Found only as compounds in nature. • Have low densities and melting points. • Also have low ionization energies.

  40. Alkali Metals Their reactions with water are famously exothermic. -Form the hydroxide and hydrogen gas

  41. Alkali Metals • Alkali metals (except Li) react with oxygen to form peroxides. • K, Rb, and Cs also form superoxides: K + O2 KO2 • Produce bright colors when placed in flame.

  42. Group 2: Alkaline Earth Metals • Less reactive than Group 1 • Have higher densities and melting points than alkali metals. • Have low ionization energies, but not as low as alkali metals.

  43. Alkaline Earth Metals • Bedoes not react with water, Mg reacts only with steam, but others react readily with water. • Reactivity tends to increase as go down group.

  44. Group 13 • B is a metalloid. Does not form binary ionic compounds. • Al, Ga, In, Tl are metals • Al forms +3 ions • Ga, In, Tl form +1 and +3 ions (lose p or lose s and p valence e-)

  45. Group 14 C = Nonmetal Si, Ge = Metalloid Sn, Pb = Metal -All can be +2 or +4 oxidation states -Sn and Pb more stable at +2; C and Si more stable at +4 -C can be -4 oxidation state

  46. Group 15 • N, P = nonmetals ex: N2,NO, N2O, NO2 P4, P4O6, P4O10 • As, Sb = metalloids • Bi = metal

  47. Group 16 • Oxygen, sulfur, and selenium are nonmetals. • Tellurium is a metalloid. • The radioactive polonium is a metal.

  48. Oxygen • Two allotropes: • O2 • O3, ozone • Three anions: • O2−, oxide • O22−, peroxide • O21−, superoxide • Tends to take electrons from other elements (oxidation)

  49. Sulfur • Weaker oxidizing agent than oxygen. • Most stable allotrope is S8, a ringed molecule.

  50. Group 17: Halogens • Prototypical nonmetals • Name comes from the Greek halos and gennao: “salt formers”

More Related