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Introduction to the Periodic Table

Introduction to the Periodic Table

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Introduction to the Periodic Table

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  1. Introduction to the Periodic Table

  2. alkali Alkaline earth Noble gas halogens <---transition--------------------> <-----------------Inner transition------------------>

  3. I. History of the Periodic Table DemitriMendeleev (1860’s Russia) • Arranged known elements: • by mass • in vertical columns by Phs/Chem Prop. • Left blank spaces (very accurate!) • Published his PT in 1871

  4. As time passed: *more elements were discovered *when placed by P/C properties, masses were out of order Henry Mosley (1913) • Found a way to determine the atomic number of elements using their emission frequency • Arranged the PT by atomic number • “Father” of the modern periodic table

  5. II. The Modern Periodic Table • Arises from Periodic Law: A. Includes: 1. periods/rows 2. groups/families • There are 4 categories of groups: • Representative (also Main Group Elements) • Transition elements • Inner transition (also Lanthanide and Actinide) • Noble Gases • Some groups have special names: • 1A (Alkali Metals) • 2A (Alkaline Earth Metals) • 7A (Halogens) • 8A I (Noble Gases) • The letters A and B in the group distinguish families • A = representative • B = transition

  6. Properties and Trends Electron Configurations

  7. L to R: metals to nonmetals • ii) Nonmetals are at the right of the Table. • They tend to be insulators and react easily with metals. • iii) Metalloids separate the metals and nonmetals and have intermediate properties • iv) Noble Gases exist at the extreme right, are chemically stable and have full valence shells Metallic Character

  8. Atomic Radius • Atomic radius – half distance between the nuclei of two atoms of same element • Shielding effect – inner energy levels ‘shield’ the outermost electrons from the positive charge pull of the nucleus • Atomic radius increasesas you move down the groups • Great distance (adding energy levels) from nucleus = less pull towards center • Decreasesas you move left to right • More pull from nucleus (more protons), but no new distance • EXCEPTION: Noble Gases – much bigger than group 17 – full outer shell

  9. Atomic Radii

  10. Ionization Energy • Ion – atom which has gained or lost electrons • Cation – (+) charged ion (lost e-) • Anion – (-) charged ion (gained e-) • Ionization energy – the energy that is required to remove an e- from an atom • Decreasesas you move down the periodic table • Outermost electron gets further from nucleus, easier to pull off • Increases as you move left to right • No more distance from nucleus, but higher charge = held more tightly

  11. Ionization Energy

  12. Electron Affinity • Electron Affinity – the energy change that occurs when an electron is ADDED to a neutral atom • Bigger negative number = easier to add e- • Harder to add as you go down a group • Further distance from nucleus & more inner e- = more repulsion felt from e- • Easier to add from left to right • Increased nuclear charge = more attraction to nucleus (+) • Noble Gases – don’t accept e- • Halogens gain most easily, they want to complete that ‘perfect eight’

  13. Electron Affinity

  14. Electron Affinity & Ionization Energy

  15. Ionization Energy vs. Electron Affinity

  16. Ionic Size (Radius) • Ionic Radius- ½ the distance between the nuclei of two ions • Cations(+) are always SMALLER than neutral atom • Nuclear charge the same, less e- = strong pull inwards • Anions (-) are always LARGER than neutral atom • Nuclear charge the same, MORE e- = less pull inwards • Increases as you go down a group • Decreases from left to right

  17. Ionic Radii

  18. Electronegativity • Electronegativity– tendency for an element to have a stronger pull on the shared e- in a covalent bond (values btwn 0-4) • Decreasesdown a group – less likely to keep the shared e- • Increasesfrom left to right – more likely to have the shared e-

  19. Electronegativity