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Bell work

Bell work. Pull out Electron Arrangements and the special periodic table that goes with it. On a scale of 1-10 (1-easy, 10-impossible), how difficult was this assignment? What parts were easy to understand? What parts were difficult?. Agenda. Bell work Notes

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Bell work

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  1. Bell work • Pull out Electron Arrangements and the special periodic table that goes with it. • On a scale of 1-10 (1-easy, 10-impossible), how difficult was this assignment? • What parts were easy to understand? What parts were difficult?

  2. Agenda • Bell work • Notes • Periodic table, periodic law, electron configurations • Practice electron configurations

  3. Development of the Periodic Table • Lavoisier (1743- 1794) • In the 1700s, Antoine Lavoisier organized a list of the known elements of his day into categories

  4. Development of Periodic Table • John Newlands (1837 – 1898) • Proposed an arrangement where elements were organized by increasing atomic mass • He was the first to organize the elements and show that properties repeated in a periodic way (every 8th element) • Law of octaves- after the musical octave

  5. Development of the Periodic Table • Mendeleev (1834- 1907) and Meyer (1830-1895) • Proposed periodic tables showing a relationship between atomic mass and elemental properties • Organized elements in order of increasing atomic mass into columns with similar properties • Mendeleev published his table first • He was able to predict properties of elements that had yet to be discovered (scandium, gallium, germanium)

  6. Mendeleev’s table

  7. Development of Periodic Table • Moseley • Organized elements by atomic number instead of atomic mass • Showed a clear, periodic pattern • Periodic law: periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number

  8. Check • Compare and contrast the ways in which Mendeleev and Moseley organized the elements • Mendeleev organized the elements by increasing atomic mass. Moseley organized them by increasing atomic number.

  9. Brief intro to the periodic table • Periods- Horizontal rows in the periodic table • Numbering goes down • Period 1 • Period 2 • Period 3 • Groups- Vertical columns in the periodic table • Numbering goes across (they’re already numbered on yours) … Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 16 Group 17 Group 18

  10. The Modern Periodic Table • Representative elements • Elements in groups 1, 2, 13, and 18 that possess a wide range of chemical and physical properties • Transition metals • Elements in groups 3 to 12 that are classified as metals, nonmetals, and metalloids

  11. Metals

  12. Metals • Characteristics • Shiny when smooth and clean • Malleable – easy to pound into sheets • Ductile – able to be drawn into wires • Alkali Metals • Group 1 elements • Extremely reactive, so they typically exist as compounds with other elements • Alkaline Earth Metals • Group 2 elements • Highly reactive • Transition metals • Group 3 – 12 elements • Inner transition metals • Lanthanide series and actinide series

  13. Metals – light blue

  14. Main group metals Transition metals Alkaline earth metals Alkali metals Inner transition metals

  15. Non-Metals

  16. Nonmetals • Nonmetals occupy the upper right side of the periodic table • Characteristics • Generally gases or brittle, dull-looking solids • Poor conductors of heat and electricity • Group 17 is comprised of highly reactive elements known as halogens • Group 18 is comprised of highly unreactive elements known as noble gases

  17. Nobel gases Non-metals – yellow Halogens Non-metals

  18. Metalloids

  19. Metalloids • Also known as semi-metals • Characteristics • Exhibit characteristics of both metals and nonmetals

  20. Metalloids – green Metalloids

  21. Valence Electrons • Electrons in the highest energy level of an atom • Only electrons available for bonding • An atom can have a MAXIMUM of 8 valence electrons (**Except: H & He – 2)

  22. Valence electrons • What is the maximum number of valence electrons any element can have? • How can you use the periodic table to determine the number of valence electrons?

  23. Electron arrangements • Pull out your electron arrangement homework and the special periodic table • Pass around the markers and erasers

  24. Which element am I describing? • 1s2 2s22p63s23p5 • What do the big numbers tell you? • What do you the letters tell you? • What do the little numbers tell you?

  25. Electron configurations • Electron configuration is a fancy way of saying electron arrangements • The large number tells you the period (energy level) • The letters tell you the block • The small (superscript) numbers tell you where in that block

  26. Classification of the Elements • Elements are organized into different blocks in the periodic table according to their electron configuration. • Just as there are four different atomic energy sublevels, there are four sections, or blocks, in the periodic table: s, p, d, f

  27. Four Blocks of the Periodic Table

  28. How many electrons in the sublevel? What is the electron configuration of sodium? 1s22s22p6 3s1 [Ne]3s1

  29. 1 2 2 Ne 3 Na

  30. Electron configurations • You can read electron configurations off of the periodic table but there are a few exceptions • The d block is always 1 big number (energy level) less than the others • The f block is always 2 big numbers less than the others • The order is based on the way you find them on the periodic table

  31. Noble gas configuration • You can write electron configurations in an abbreviated way • You use a noble gas to represent some of the electrons • 1s2 2s22p63s23p64s23d104p3 • What element is this? • What noble gas is within this configuration? • Ar : 1s2 2s22p63s23p6 • 1s2 2s22p63s23p64s23d104p3 • [Ar]4s23d104p3

  32. Practice • Write the electron configuration for the following elements: • Boron • Vanadium • Germanium • Technetium • Promethium

  33. Identify the following elements: • 1s2 2s22p63s23p5 • [Ne]3s1 • [Kr]5s24d2 • 1s2 2s22p63s23p64s23d104p1 • 1s2 2s22p63s23p64s23d104p65s24d5 • [Xe]6s24f10

  34. Bell work • Using the periodic table, write the electron configuration for the following elements: • As • Kr • F • C • Fe • Ni • I • Re • Zr • Cs

  35. Periodic Trends • Take out periodic trends study guide

  36. Practice • Which has the largest atomic radius: magnesium, silicon, sulfur, or sodium? • Which has the smallest atomic radius: helium, krypton, or radon?

  37. Atomic Radius

  38. Atomic Radius • Atomic radius is a periodic trend influenced by electron configuration • Atomic radius = how closely an atom lies to a neighboring atom (or half the distance between two nuclei)

  39. Atomic Radius • Trends within periods • What do you notice about the atomic radius as you move across a period? • There is a decrease in atomic radii as you move from left to right along a period • This is caused by increasing positive charge and filling orbitals of the same principal energy level • The valence electrons are not shielded from increasing nuclear charge, which pulls the outermost electrons closer to the nucleus

  40. Atomic Radius • Trends within groups • What do you notice about atomic radii as you move down a group? • There is an increase in atomic radii as you down a group • As the nuclear charge increases, electrons are added to successively larger principal energy levels • So, increased nuclear charge doesn’t pull the outer electrons toward the nucleus to make the atom smaller

  41. Atomic Radius Summary

  42. Practice • Which has the largest atomic radius: magnesium, silicon, sulfur, or sodium? • Which has the smallest atomic radius: helium, krypton, or radon?

  43. Ionic Radius • Ion = an atom or bonded group of atoms that has a positive or negative charge

  44. Ionic Radius What happens to the ionic radius as you go down a group? What do you notice about the charges in each group? How does charge relate to ionic radius?

  45. Ionic radius • When atoms lose electrons and form positively charged ions, they always become smaller. • The loss of a valence electron may leave an empty outer orbital • The decreased electrostatic repulsion between the now-fewer remaining electrons allows the electrons to be pulled closer to the positively charged nucleus

  46. Ionic Radius • When atoms gain electrons and form negatively charged ions, they become larger. • The addition of an electron increases the electrostatic repulsion between the atom’s outer electrons, forcing them to move further apart.

  47. Ionic Radius • Trends within periods • Elements on the left side of the periodic table form smaller positive ions and elements on the right form larger negative ions • As you move from left to right across periods 1 to 14, the size of positive ions decreases • As you move from left to right across periods 15 to 18, the size of the much larger negative ions also decreases

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