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Ch 5 Periodic Table

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					                           Chapter 5: Periodic Table
                                  Objectives




5-1 Development of Periodic Table

 State the periodic law
 Discuss the contributions of Dobereiner, Newlands, Mendeleev, and
   Moseley made to the periodic table.



5-2 Reading the Periodic Table

       Explain why the elements in a group have similar properties
       Identify the four blocks of the periodic table




5-3 Periodic Trends
       Define the term periodic trend
       Identify four important periodic trends and explain how each
         reflects the electron configuration of the elements.




The Element Song- Tom Lehrer
http://www.youtube.com/watch?v=GFIvXVMbII0


http://www.youtube.com/watch?v=d0zION8xjbM




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5-1 Development of the Periodic Table
   Periodic Table- was developed to help chemists organize & classify
     the elements, it also predicted the existence of elements yet to be
     discovered
Forerunners of the Periodic Table
   Late 1700’s scientists had discovered ~ 30 elements, IE> gold, silver
     (metals) H,O,N,C (nonmetals)
   Early 1800’s lab techniques improved. In less than 100 years the # of
     known elements doubled. Since there were more known elements it
     became more to organize & classify them.
   J.W. Dobereiner- Early 1800’s this German chemist observed that
     many elements could be classified by similar properties into sets of
     three called triads. It was also noted that the middle element of the
     triad had ~ average of properties of the first and third elements. IE>
     (Li, Na, K) (Ca, Sr, Ba) (Cl, Br, I). Using chart p.160 compare the
     atomic mass & density of the middle element (Sr) to the average of
     the first & third (Ca, Ba)
   J.A.R Newlands- in 1865 this English chemist arranged the 62 known
     elements according to increasing atomic mass & noticed that the
     properties repeated every 8 elements. Because it repeated every 8th
     element he called it the law of octaves, after the 8 notes on the
     musical scale. His ideas were not taken seriously for over 20 years.
   Dimitri Mendeleev- this Russian chemist wanted to organize elements
     for his students to learn more easily, he used notecards to arrange. He
     produced the 1st Periodic Table which included the repeating 8
     pattern & placed in order of  atomic weight.
   Not all elements were known. Mendeleev would break the 8 pattern in
     his table in order to keep elements w/ like properties together, he
                                                                               2
      suggested that some elements had had their atomic mass incorrectly
      measured and they needed to be re-massed.
    Also predicted the existence of 3 elements, leaving gaps in his table
      where they should fit in. IE> below Si he left a gap for element called
      Ekasilicon, 20 years later germanium was discovered
http://spiff.rit.edu/classes/phys314/lectures/period/period.html
http://www.vanderkrogt.net/elements/chemical_symbols.html
http://chemistry.about.com/library/weekly/aa030303a.htm
Look @ current periodic table. How are elements arranged?
    H. G. J. Moseley- in late 1800’s this English chemist did x-ray work w/
      nuclei. He noticed that metals produce x-rays when bombarded w/
      energetic e- & that the frequency was distinct to that metal. He
      predicted that this was due to the different amounts of positive charge
      in the nucleus (today we know this is true- protons). He assigned a
      series of whole #’s, called atomic numbers, to each element based on
      the frequency that was emitted. Today we know that the atomic # = #
      protons, this is what ID’s ther element
    Mendeleev was wrong after all, it wasn’t atomic mass but atomic #. He
      had used the wrong property but had found the right pattern, this is
      why he is given so much credit.
    Periodic Law- when elements are arranged in order of increasing
      atomic number, their physical and chemical properties show a
      periodic pattern.
   Homework 5-1 Review, 5-1 Apply




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5-2 Reading the Periodic Table
Organizing the squares
    See figure 5-7 p. 164. What info can we get from the square?
    Atomic #, Chemical symbol, Element name, Atomic mass, Electron
      configuration, and more.
    Organized by increasing atomic #
    Elements w/ similar properties arranged into columns, called groups
   or families (18)
    Horizontal rows called periods (7)
Labeling the groups- see 5-9 p.165.
    Different chemists use different labeling systems for groups of
      elements. This book uses the American system w/ Roman numerals.
    Some common family names
         o 1A- alkali metals
         o 2A- alkali earth metals
         o 7A- halogens
         o 8A- Noble gases
         o other groups sometimes named for first element in the group,
            IE> carbon group
    Hydrogen is in the alkali metal family but slightly separated. Even
      though it has similar properties it is not a metal. It is a unique element
      and therefore is slightly spaced from the rest.
Metals, Nonmetals, and Semimetals
    Metals- shaded light blue on left side. See figure 5-10 p. 167
      Properties include:
       Metallic luster or shine
       Good conductors of heat and electricity
       Usually solid @ room temp (except Hg)
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       Malleable (hammered into thin sheets, IE> aluminum)
       Ductile (drawn into fine wire, IE> copper)
 Nonmetals- shaded light red and located to the right (+ H). Properties
   include:
    Basically show none of characteristics of metals (review)
    Can be any phase @ room temp. Most = gas, some = solid, Br = liquid
    Wide range of physical properties (color/colorless, hard/soft solids)
 Semimetals/metalloids- show some properties of both, shaded light
   purple. Located between metal & nonmetals on either side of the stair
   step.
http://www.youtube.com/watch?v=LDHg7Vgzses
Electron configurations & the periodic table
    1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d
    s=1, p=3, d=5, f=7
    Periods are based on valence e-.
    Valence electrons- electrons that occupy the highest energy level; the
      outermost electrons in an atom. They are responsible for the chemical
      properties of the element.
    Periodic table is structured so that elements with similar valence e- are
      placed in the same group, see 5-14 p. 171
    Abbreviated e- configurations-
       Li = 1s2 2s1 or [He]2s1
       Na = [Ne]3s1
       K = [Ar]4s1
       Notice that all end in s and their period #




                                                                             5
The s-, p-, d-, f- block elements
    s-block- H, He, Alkali & Alkali Earth metals. These elements have
      valence e- in the s-orbital only, so it only contains 2 groups because s-
      orbital only holds 2 electrons
    p-block- from left to right, elements’ valence e- fill p orbitals.
2 Groups 13- 18
       6 groups wide because 6 valence e- in the p orbital
    d-block- takes up most of middle of table.
       Elements’ valence e- fill the d orbital, so this block is 10 elements
        wide because d orbital holds 10 electrons
    f-block- essentially the 28 elements below the periodic table. This
      block is 14 wide because the f orbital holds 14 e-
    s & p block also called representative or main group elements
    d block also called transition metals
    f block also called inner transition metals
Homework 5-2




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5-3 Periodic Trends
 Periodic trends- many properties of the elements change in predictable
   ways as you move through the periodic table


Atomic Radius (see figure 5-16 p. 175)
    Atomic radius = the distance from the center of atoms nucleus to
      outermost electron (this e- position is estimated, so how do we
      measure?) Take the distance between the 2 nuclei & divide by 2 in
      order to get atomic radius




 2 trends seen in atomic radius-
1. Atoms get larger moving down a group, IE> add a new energy level;
   principle # increases = higher orbitals = larger atomic radius
2 Atoms get smaller moving from left to right, IE> as go right = more
   protons = more pull on e- = smaller orbital = smaller atomic radius




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8
Ionic size (see figure 5-17 p.176)
 Ions- atoms that have gained or lost e-
 Cations (+ charge)- have lost e-, so are smaller =  attractive force
 Anions (- charge)- have gained e-, so are larger =  attractive force

Which is larger F or F- ?

 Groups 1A-2A form small + ions

 Groups 5A-8A form large – ions




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Ionization Energy

 Ionization energy – the amount of energy required to remove an e- from

   an atom, IE> Li(g)  Li + (g) + e- where ionization energy = 8.64 x 10-19

   J/atom

 Ionization energy (Ion E)- is a reflection of how strongly an atom holds

   onto its outermost e-. High Ion E means a tight hold.

 Ion E- usually reported for a large collection of atoms. Common unit for

   a large collection of atoms is the mole (mol). 1 mol = 6.023 x 1023 atoms.

   Li(g)  Li+ (g) + e-, Ion E = 521 kJ/mol

 Ion E- varies w/ atomic #




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1st Ionic Energy shows 2 trends:

1. 1st Ion E  as you move  a group

2. 1st Ion E  as you move  in a period




 These trends are exactly opposite of the trends for atomic radius due to

   the fact that both an atom’s size & Ion E depend on how strongly its e-

   are attracted to the nucleus

http://www.youtube.com/watch?v=ywqg9PorTAw

Work through sample problem 1, Practice problems 1-2, and alt practice 1.




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Successive Ionization Energies

 1st Ion E -is the energy to remove the 1st ion

 Successive Ion E- (table p. 180) energies required to remove the e-

   beyond the first electrons. Energy required to remove second e- is the 2nd

   Ion E, the energy required to remove the third e- is the 3rd Ion E, and ect.

    Li(g)  Li+ (g) + e-, 1st Ion E = 521 kJ/mol

    Li(g)  Li2+ (g) + e-, 2nd Ion E = 7304 kJ/mol

    Li(g)  Li3+ (g) + e-, 3rd Ion E = 11,752 kJ/mol

 Ion E  after each e- is removed due partially to reduced e-e- repulsion

 There is a large jump in Ion E’s for each element, but it does not proceed

   smoothly because:

    Atoms hold on to inner e- more than valence e-




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    Sodium = [Ne] 3s1 It has a small 1st Ion E but much larger 2nd Ion E,

      so it is likely to lose 1 valence electron and be found as Na +, but not

      found as Na+2

    Mg = [Ne] 3s2 The large jump is after 2nd e- is removed




Electron affinity (see figure5-21 p. 182)

    Electron affinity- Energy change that occurs when an atom gains an e-

    Ne(g) + e-  Ne-(g) Electron affinity = 29 kJ/mol




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Looking at your chart p.182, what do you notice about most electron

affinities? Most –

    F(g) + e-  F-(g) Electron affinity = -328 kJ/mol

    EA is the measure of an atom’s attraction (affinity) for an extra e-.

      Energy isn’t required; it is released.

    The more negative the EA, the more or stronger attraction it has for e-.

    EA is somewhat irregular across the table. It is related to valence e- and

      filling the outer orbital

    F = [He]2s22p5 + e- [He]2s22p6 this e- fills the shell (orbital). Noble

      gases don’t want an e-, their valence shell (outer orbital) is full

    Alkali metals, alkali earth metals & halogens commonly exist as ions.

      Why?

    Octet Rule- atoms tend to gain, lose, or share e- in order to acquire a full

      set of valence electrons.

    Electronegativity- an atoms ability to attract electrons in a chemical

      bond. It is an arbitrary scale with no units and can not be measured by its

      amount of energy. The electronegativity trend is increasing and

      increasing ,  F is the most electronegative

http://www.youtube.com/watch?v=XMLd-O6PgVs

Homework 5-3 Practice, 5-3 Apply & 5-3 Review



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Four Periodic Trends




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