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Chemistry Chapter 5. The Periodic Table. Sept 1860, group of chemists met in Germany to review scientific matters & coming to consensus about: measurement of atomic masses determining composition of compounds using atomic masses

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chemistry chapter 5

ChemistryChapter 5

The Periodic Table

slide3

Sept 1860, group of chemists met in Germany to review scientific matters & coming to consensus about:

    • measurement of atomic masses
    • determining composition of compounds using atomic masses
    • Cannizzaro presented method for measuring mass & scientists agreed upon the std values for atomic masses
slide4

in 1869, Dmitri Mendeleev used Cannizzaro’s method for measuring relative masses of atoms in textbook he wrote

  • published the arrangement of elements in – periodic table
slide5

only 60 elements known at this time

  • organized elements acc to properties & new atomic masses on cards
  • “game of patience”
  • Mendeleev grouped elements according to incr atomic mass & noticed certain properties appeared at regular intervals- periodic
slide7

in 1871, Mendeleev predicted properties of elements that weren’t even discovered at that time!

  • not all elements fit in according to increasing atomic mass:
    • I & Te Ar & K Co & Ni
    • Mendeleev couldn’t explain why
    • other scientist accepted the periodic table & considered him the Father of the Periodic Law
slide9

in 1913, Henry Moseley discovered patterns w/ x-ray tubes that led to atomic number (ch 3 notes)

  • he noticed that when he reordered elements on table acc to incr atomic number they fit into their patterns in better way
  • led to the Modern periodic table
  • periodic law- phy & chem properties of the elements are periodic functions of the atomic #
group properties
Group Properties
  • valence e- same for all elements in group
  • group 1: Alkali metals
    • e- conf: ns1
  • group 2: Alkaline Earth metals
    • e- conf: ns2
  • groups 3-12: Transition metals
    • e- conf: (n-1)d1ns2
slide12

groups 4-11 deviations occur

  • sum of outer s & d e- equal to group #
  • group 13: ns2np1
  • group 14: ns2np2
  • group 15: ns2np3
  • group 16: ns2np4
  • group 17: Halogens ns2np5
  • group 18: Noble gases ns2np6
slide13

noble gases have 8 valence e-

  • have stable octet very stable & unreactive
  • f-block elements
    • lanthanides- rare earth metals 1st row
    • actinides- all radioactive; most synthetic
periodic properties
Periodic Properties
  • phy & chem properties vary in periodic fashion
  • properties arise from e- configuration
  • 5 properties:
1 atomic radii
1. Atomic Radii
  • ½ distance betw nuclei of identical atoms joined in a molecule
  • e- occupy large region around nucleus & size atom varies
  • periodic trends- gradual decrease in radii across periods
    • due to increasing pos chrg of nucleus (pulled tighter by nucleus)
slide16

group trends: as go down group, atomic radii increases due to addition of e- to larger orbitals in higher energy levels

2 ionization energy
2. Ionization Energy
  • minimum amount of energy required to remove the most loosely bound e- from an isolated gaseous atom to form an ion w/ a +1 charge
  • if enough energy is supplied, e- can be removed from atoms
  • ex: 1st IE for Ca is 590 kJ/mol
  • Ca + 590kJ/mol  Ca+ + e-
slide19

ionization- process that results in the formation of ion

  • 2nd IE is 1145kJ
  • IE2 > IE1
  • ALWAYS more difficult to remove additional e- from positive ion
  • IE measures how tightly e- are bound to atoms
slide20

low IE indicates ease of e- removal & cation formation

  • group trends: as atomic radii increases in a group, 1st IE decreases
    • b/c the valence e- are further from nucleus “shielding effect”
    • period trends: IE incr from L to R due to increasing nuclear charge which holds e- tighter
3 e affinity
3. E- affinity
  • amount of energy involved in the process in which an e- is added to an isolated gaseous atom to produce an ion w/ a -1 charge
  • many atoms readily add e- & release energy
    • ex: Cl + e-  Cl- + energy (exothermic)
    • Why?
slide23

some have to be forced to gain e- by the addition of energy

    • Be + e- + energy  Be-
    • period trends: group 17 elements gain e- most easily ( large neg values) reason for the reactivity of these halogens
slide25

exceptions are seen betw groups 14 & 15 b/c ½ filled sublevels are a little more stable than ones not ½ full

  • group trends: generally more difficult to add e- to larger atoms than to smaller atoms
  • elements w/ very negative EA gain e- readily to form anions (ions w/ negchrg)
slide26

more difficult to add e- to an anion so 2nd EA are all positive

  • cation- positive ion
  • anion- negative ion
4 ionic radii
4. Ionic radii
  • ½ the diameter of an ion in a chemical compound
  • formation of a cation leads to a decrease in radius due to the e- cloud being drawn inward as valence e- are removed
  • formation of anion leads to an increase in radius as additional e- repel one another
slide29

periodic trends- metals form cations

    • nonmetals form anions
    • group trends- IR increases down group Why?
    • as you add higher energy levels, radius of ion incr
5 electronegativity en
5. Electronegativity EN
  • measure of the power of an atom in a chemical compound to attract e-
  • valence e- hold atoms in compound together & properties of compound are influenced by conc of negchrg closer to one atom than another
  • ex: NaCl
slide33

numerical values assigned to indicate the tendency of atom to attract e-

  • Fluorine – most EN element & assigned value of 4
  • periodic trends- gradual incr in EN from L to R across period
  • nonmetals tend to be more EN than metals
slide34

groups 1 & 2 least EN elements

  • halogens are most EN elements
  • group trends- EN either decreases down group or remains similar
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