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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

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ChemistryChapter 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

    • Cannizzaro presented method for measuring mass & scientists agreed upon the std values for atomic masses

  • 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

  • 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

  • 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 & TeAr & KCo & Ni

    • Mendeleev couldn’t explain why

    • other scientist accepted the periodic table & considered him the Father of the Periodic Law

  • 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 #

Getting Acquainted With the Periodic Table

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

  • 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

  • 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

  • phy & chem properties vary in periodic fashion

  • properties arise from e- configuration

  • 5 properties:

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)

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

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-

  • 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

  • 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

nonmetals tend to have higher IE than metals

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?

  • 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

  • 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)

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

  • cation- positive ion

  • anion- negative ion

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

  • 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

  • chemical compounds form b/c e- are lost, gained, or shared to bring an atom to a stable octet

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

  • 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

  • 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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