The periodic table & periodic law

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The periodic table & periodic law

  1. 1. The Periodic Table & PeriodicLaw
  2. 2. Objective: ▫ Trace the development of the periodic table ▫ Identify key features of the periodic table
  3. 3. • Sheldon & The Element Song
  4. 4. Development of the periodic table• - John Newlands• - Meyer and Meedeleev
  5. 5. John Newlands• 1864 English chemist• -Proposed a scheme for the elements• -He noticed that when elements where arranged based on atomic mass their properties repeated every 8th element.
  6. 6. John Newlands• -Newlands named the periodic relationship Law of octaves. (after a musical, a name which was frowned upon because it was unscientific)• -Although not accepted Newland’s correct in the aspect that elements do repeat each other.
  7. 7. Meyer & Mendeleev• Meyer-Russian chemist-1869• Dmitri Mendeleev- 1869 ▫ Demonstrated a connection between atomic mass and elemental properties.
  8. 8. Meyer ▫ Demonstrated a connection between atomic mass and elemental properties ▫ Arranged the elements in order of increasing atomic mass
  9. 9. Mendeleev ▫ Demonstrated a connection between atomic mass and elemental properties ▫ Arranged the elements in order of increasing atomic mass ▫ Predicted the existence of properties of undiscovered elements.
  10. 10. Mendeleev ▫ Mendeleev is given more credit. ▫ History of the Periodic Table
  11. 11. Periodic Table Activity ▫ Recognizing patterns
  12. 12. What did Mendeleev miss? ▫ New elements were discovered, it became evident that the periodic table was not in the correct order. ▫ Henry Moseley (1887-1915), arrange elements by atomic number.
  13. 13. Moseley ▫ Discovered that atoms contain a unique number of protons called the atomic number ▫ Arranged element in order of increasing atomic number, which resulted in a periodic pattern of properties.
  14. 14. The Periodic Law• Atoms with similar properties appear in groups or families (vertical columns) on the periodic table.• They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.
  15. 15. The Periodic Table & PeriodicLawElements & Modern periodic table
  16. 16. • Element Song
  17. 17. Modern Periodic Table• Groups= columns “Vertical”• Periods= rows “horizontal” ▫ Periodic properties  Periods have similar characteristics.
  18. 18. Metals, Nonmetals, Metalloids• How can you identify a metal?• What are its properties?• What about the less common nonmetals?• What are their properties?• And what the heck is a metalloid?
  19. 19. Alkali Metals• Li• Na• K• Rb• Cs• Fr
  20. 20. Alkali Metals ▫ 1 valance electron very reactive  Alkali metals with water
  21. 21. Alkali metals ▫ Very reactive, often exists as compounds with other elements ▫ Two familiar alkali metals are.. Na- used in table salt Li- Used in batteries
  22. 22. Alkaline Earth▫ 2 valance electron▫ Reactive-forms oxides▫ Important in living things  Found on planet in raw forms
  23. 23. Halogens ▫ Not “super” stable ▫ 7 valance electrons  Need 1 more
  24. 24. Halogens Cl- Halogen Used as a gas in WWI
  25. 25. Noble Gases ▫ Most un-reactive ▫ Why?  8 valance electrons
  26. 26. Inert gases… ▫ Is a gas which does not undergo chemical reactions under a set of given conditions  Mig welder- uses argon
  27. 27. CHNOPSS-nonmetals “important to life” ▫ Carbon- thats what we are “made of” ▫ Hydrogen-your body is mostly H2O ▫ Nitrogen – amino acids ▫ Oxygen-energy out of food ▫ Phosphorous-DNA ▫ Sulfur-Protein ▫ Selenium –micro amounts/ a deficiency is thought to cause cancer
  28. 28. Nonmetals • Nonmetals are the opposite. • They are dull, brittle, nonconductors (insulators). • Some are solid, but many are gases, and Bromine is a liquid.
  29. 29. Transition Metals ▫ “weird” number of electrons ▫ Cu-copper ▫ Au-Gold ▫ Ag-Silver
  30. 30. Titanium ▫ Strong and light it is often used to make frames for bicycles and eyeglasses.
  31. 31. “Poor” metals ▫ Al- Aluminum ▫ Ga- Gallium ▫ In – Indium ▫ Sn-Tin ▫ Ti- thallium ▫ Pb- lead ▫ Bi- Bismuth
  32. 32. “inner” transition metals ▫ Actinide series ▫ Lanthanide series  used extensively as phosphors, substances that emit light when struck by electrons
  33. 33. Metals ▫ Malleable & ductile, meaning they can be pounded into thin sheets & drawn into wire  Metals met the nonmetals, forming a stair step on the right hand side of the periodic table.
  34. 34. Metals • Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity. • They are mostly solids at room temp. • What is one exception?
  35. 35. “The disappearing Spoon” • Disappearing Spoon Gallium Low melting point Highly radioactive= good bye hand
  36. 36. Metalloids ▫ B-Boron ▫ Si-Silicon ▫ Ge-Germanium ▫ As-Arsenic ▫ Sb-Antimony ▫ Te-Tellurium ▫ Po-Polonium ▫ At-Astatine
  37. 37. Metalloids • Metalloids, aka semi- metals are just that. • They have characteristics of both metals and nonmetals. • They are shiny but brittle. • And they are semiconductors. • What is our most important semiconductor?
  38. 38. Metalloids▫ Semiconductors▫ Properties of both metals and nonmetals.
  39. 39. Metalloids▫ Two important metalloids  Ge-germanium  Si-Silicon Used in computers chips and solar cells
  40. 40. Metals, Nonmetals, Metalloids• There is a zig-zag or staircase line that divides the table.• Metals are on the left of the line, in blue.• Nonmetals are on the right of the line, in orange.
  41. 41. Metals, Nonmetals, Metalloids• Elements that border the stair case, shown in purple are the metalloids or semi- metals.• There is one important exception.• Aluminum is more metallic than not.
  42. 42. The Periodic Table & PeriodicLawClassification of Elements
  43. 43. Objective: ▫ Explain why elements in the same group have similar properties ▫ Identify the four block of the periodic table based in their electron configuration
  44. 44. Valence Electrons• Do you remember how to tell the number of valence electrons for elements in the s- and p- blocks?• How many valence electrons will the atoms in the d-block (transition metals) and the f-block (inner transition metals) have?• Most have 2 valence e-, some only have 1.
  45. 45. Valance electrons
  46. 46. Electron configuration
  47. 47. s, p, d, and f blocks
  48. 48. The Periodic Table & PeriodicLawPeriodic Trends
  49. 49. Atomic RadiusDefinition: Half of the distancebetween nuclei in covalentlybonded diatomic molecule Radius decreases across a period  Increased effective nuclear charge due to decreased shielding Radius increases down a group  Each row on the periodic table adds a “shell” or energy level to the atom
  50. 50. Table ofAtomicRadii
  51. 51. Explanation Video-Atomic Radii
  52. 52. Period Trend:Atomic Radius
  53. 53. Ionization EnergyDefinition: the energy required to removean electron from an atom  Tends to increase across a period  As radius decreases across a period, the electron you are removing is closer to the nucleus and harder to remove  Tends to decrease down a group  Outer electrons are farther from the nucleus and easier to remove
  54. 54. Periodic Trend:Ionization Energy
  55. 55. Electronegativity Definition: A measure of the ability of an atom in a chemical compound to attract electronso Electronegativity tends to increaseacross a period o As radius decreases, electrons get closer to the bonding atom’s nucleuso Electronegativity tends to decrease downa group or remain the same o As radius increases, electrons are farther from the bonding atom’s nucleus
  56. 56. Periodic Table of Electronegativities
  57. 57. Periodic Trend:Electronegativity
  58. 58. Summary ofPeriodic Trends
  59. 59. Ionic Radii  Positively charged ions formed when an atom of a metal loses one orCations more electrons  Smaller than the corresponding atom  Negatively charged ions formed when nonmetallic atoms gain one Anions or more electrons  Larger than the corresponding atom
  60. 60. The Octet Rule • The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level. • They may accomplish this by either giving electrons away or taking them. • Metals generally give electrons, nonmetals take them from other atoms. • Atoms that have gained or lost electrons are called ions.
  61. 61. Ions• When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion.• In the same way that nonmetal atoms can gain electrons, metal atoms can lose electrons.• They become positively charged cations.
  62. 62. Ions• Here is a simple way to remember which is the cation and which the anion: + + This is Ann Ion. This is a cat-ion. She’s unhappy He’s a “plussy” cat! and negative.
  63. 63. Ionic Radius• Cations are always smaller than the original atom.• The entire outer PEL is removed during ionization.• Conversely, anions are always larger than the original atom.• Electrons are added to the outer PEL.
  64. 64. Cation FormationNa atom Effective nuclear charge on remaining electrons1 valence electron increases. 11p+ Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Valence e- lost in ion formation Result: a smaller sodium cation, Na+
  65. 65. A chloride ion isAnion Formation produced. It is larger than the original Chlorine atom atom. with 7 valence e- 17p +One e- is added tothe outer shell. Effective nuclear charge is reduced and the e- cloud expands.
  66. 66. Graphic courtesy Wikimedia Commons user Popnose

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