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					                                   Unit VI: Periodicity (Chapter 3 & 13)

    1. Vocab: periodicity, period, group, family, first ionization energy, electronegativity, atomic radii,
       ionic radii, metallic radius, covalent radius, van der waals’ radius, valence electrons, alkali
       metals, alkaline earth metals, halogens, noble gases, transition metals, cations and anions,
       electron arrangement vs. electron configuration, shielding effect, effective nuclear charge,
       isoelectronic species, oxidation states, complex ions, ligands, dative bonds, coordination
       number, ligand replacement, oxides and chlorides, amphoteric

    2. Introduction
     For years, scientists have been searching for a way to arrange the elements that reflects the fact
       that certain physical and chemical properties occur ___________________. (read “History of
       Chemistry, p. 75)
     ______________________ was a Russian chemist who designed the modern periodic table. The
       only difference was that he arranged the elements by ___________________________, while
       now we arrange them by ________________________________.
     One of the brilliant aspects of Mendeleev’s periodic table is that not all the elements had been
       discovered yet; he left blanks on the table and predicted that those elements would eventually
       be discovered!
     _____________________– repeating patterns of chemical and physical properties



    3. Arrangement of the Periodic Table

The arrangement of elements into groups and periods is a reflection of their
_____________________________________ e.g. the alkali metals all have one valence electron and are
therefore very reactive.

The representative groups that you should know from previous Science classes are: alkali metals,
alkaline earth metals, halogens and noble gases.

_________________________________________: elements that separate Group 2 and 3. They are
unique in that there are similarities both down and across the block. Transition metals form more than
one stable cation with an incomplete d-subshell. E.g. Fe 2+ and Fe 3+ both exist in nature.

Note: based on this definition, zinc and scandium, though d-block elements, are not considered
transitional metals. The d-block elements are discussed later in more detail.

    4. Periodic Trends - Physical Properties

Effective Nuclear Charge (ENC): A force that keeps the _________ close to the nucleus. It is the
_________ attractive force felt by an e- after the _______________________________ of other e- in
the atom has been taken into consideration.
Visualizing the ENC helps us understand the trends in physical properties: atomic and ionic radius,
ionization energy, electronegativity and melting point.

    A. Atomic and Ionic radius

ACT: Graphing atomic radius

This table shows atomic radius measured in angstroms (1 angstrom = 1.0 × 10-10 meters)




Source: http://intro.chem.okstate.edu/1314f00/lecture/chapter7/lec111300.html

Atomic radius: measures the size of the e- probability cloud for a neutral element

       Across a period, atomic radius _________
            o Reason:
                ________________________________________________________________________
                ________________________________________________________________________
                ________________________________________________________________________
       Down a family/group: ___________
            o Reason:
                ________________________________________________________________________
                ________________________________________________________________________
                ________________________________________________________________________
Ionic Radius: measures the size of the e- probability cloud for a charged particle, or ion

   When atoms become ions, they lose or gain electrons.
   Cations (+ charged ions) are a result of losing electrons. This creates an ion that is
    ______________ than the neutral atom. Reason: greater ENC from nucleus
   Anions (- charged ions) are the result of gaining electrons. This creates an ion that is larger
    than the neutral atom. Reason: greater e- to e- repulsion.

For a good analogy of periodic trends explained by ENC, read:
http://laude.cm.utexas.edu/courses/ch301/lecture/ln5f07.pdf

Read “Isoelectronic Species” on p. 83


    B. Ionization energy

Energy required to _______________ an ________ from a gaseous atom or ion. The most weakly held e-
(valence e-) are ejected first. Usually measured in kJ/mol.

ACT: Graphing Ionization Energy

   Another way to explain IE is how tightly an electron is held by the atom. The more energy (IE) it
    takes to remove an electron, the greater the electron is held by the atom.

   The same factors that affect atomic size affect ionization energy.

   The further down a group, the larger the orbitals and the further the distance the outer electrons
    from the nucleus. Electrons further from the nucleus are bound less tightly – lower IE.

   As we go across a period, the number of protons increases and therefore the effective nuclear
    charge increases on each of the valence (outer) electrons. Thus a greater amount of energy is
    needed to remove an electron.

   We commonly use the terms first IE, second IE, third IE, etc. This just refers to the energy needed to
    remove the first electron, second electron, …

   The energy needed to remove additional electrons follow the same rules.

   The more stable the atom (or ion), the higher the IE. e.g. noble gases have high IE, and any ions with
    a full outer orbit have a high IE

e.g. Aluminum

First IE: 577 kJ/mol

Second IE: 1817 kJ/mol
Thrid IE: 2745 kJ/mol

Fourth IE: 11 577 kJ/mol

Why is there such a large jump between IE3 and IE4?
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________



    C. Electronegativity:

ACT: Graphing Electronegativity

Electronegativity is the ability of an atom in a covalent bond to _____________ electrons to itself.

Once again, the atomic radii will determine the electronegativity of an atom.

Going down a family:

The further the outer orbital from the nucleus, the smaller the attraction between the positive nucleus
and the negative electron.



Going across a period:

The greater the effective nuclear charge, the greater the attraction between the nucleus and an
electron.



    D. Melting Points

At left of periodic table

   elements are metallic (Li, Na, K)

   Going down a metal group, the metallic bonding increases. More on metallic bonding in the next
    unit, but it occurs because valence electrons are shed and become delocalized. As the distance
    between positive nuclei and delocalized e- increases, the attractive forces decrease, and therefore
    the melting points decrease
   Going across the period, increase in valence electrons = increase in melting point



At the centre of each period (Si & C)

   Giant covalently bonded structures with very strong bonds between particles – therefore very high
    melting points

At the right of each period

   Elements have molecular covalent structures and there are only weak van der Waal’s forces
    between the molecules (more on this in our Bonding unit)

Check your understanding: you should be able to explain the observed trends in physical properties for
Group 1, Group 7 and going across Period 3.



    5. Periodic Trends - Chemical Properties

    a) Reaction of alkali metals (Li, Na, K)
       i.     with H2O

                                         Li + H2O  LiOH + H2

        -makes a strong _________________ (basic) solution.

        -a highly exothermic reaction

        -this is called a redox reaction, where lithium acts as a reducing agent. This means it donates
        electrons to the water, and the lithium is said to be oxidized. More on this in Grade 12



        ii.     with Cl2 & Br2

                                           2 Li + Br2  2 LiBr

        -reacts readily because of _____________________________________________



Note – reactivity of alkali metals increases with decreasing ionization energies (down the group)
    b) Reaction of halogens (Cl2 , Br2, I2)

    i.      with H2O

   halogens are only slightly soluble in H2O (as halogens non-polar)

   dissociate to form slightly acidic solutions

                                      X2 + H2O  H+ + X- + HOX

   Compound HOX can readily donate O to other substances (acts as an oxidant)



Eg. It will oxidize colored dyes to colorless products



    ii.    with halide ions (X-)
   ACT – Halogen-halide lab

   more electronegative elements will displace a less electronegative element from its salts. These are
    called displacement reactions e.g.



                                           Cl2 + 2 I-  I2 + 2 Cl-

                                               I2 + F-  no rxn



    iii.    with metals

   form white salts which are all H2O soluble forming colorless solutions, EXCEPT Ag and Pb



                                          Ag+(aq) + X-(aq)  AgX(s)
       6. Trends across period 3 (Na  Ar)

You will learn to explain the physical and chemical trends in the chlorides and oxides of the third period

Oxides: Na2O, MgO, Al2O3, SiO2, P4O6 and P4O10, SO2and SO3, Cl2O and Cl2O7

Chlorides: NaCl, MgCl2, Al2Cl6 (AlCl3), SiCl4, PCl3 and PCl5, Cl2



       A. Chlorides of metals

      have high melting points

      when added to water, form ions – conduct electricity

Note: AlCl3 is actually Al2Cl6 in its gaseous state



       B. Chlorides of non-metals

      Have molecular covalent structure

      Weak forces between molecules – low melting and boiling points

      Don’t conduct

      When added to water (Hydrolysis) the bonds between the element and chlorine are replaced by the
       bonds between the element and oxygen

      Result is an acidic solution – containing chloride ions, & an oxyacid of the element



E.g.               PCl3 + 3 H2O  3 HCl + P(OH)3             [ P(OH)3 is H3PO3 , or phosphoric (III) acid]



      the oxyacid can also dissociate:

                                             H3PO3 = H+ + H2PO3-



Therefore, pure chlorides will not conduct electricity, but aqueous solutions of these chlorides will.

Note: Cl2 acts as a chlorine chloride, having the same properties of other non-metal chlorides
Note: Aluminum acts more as a non-metal chloride as a result of the small size and high charge of the
ion it forms

Sublimes at relatively low temperature of 178oC

Also undergoes vigorous hydrolysis when added to water.




    C. Oxides

Change from basic to acidic character as they move across period



Oxides on left

   are ionic solids

   behave as basic oxides and react with water to form an alkaline solution of the hydroxide (Mg(OH)2
    only slightly acidic due to low solubility)



                                 Na2O(s) + H2O(aq)  2 Na+(aq) + 2 OH-(aq)



   They will react with acids to form a salt and water



                                    MgO + 2 H+  Mg2+(aq) + H2O(l)



Center of periodic table

   Oxides are giant covalent lattices with very high melting and boiling points

   Different from chlorides as chlorine only forms single bond, making lattice impossible

   Al2O3 and SiO2 virtually insoluble due to large structures



Al2O3 is amphoteric (will dissolve in both acids and bases)
                                 Al2O3(s) + 6 H+(aq)  2 Al3+(aq) + 3 H2O(l)



                               Al2O3(s) + 2 OH-(aq) + 3 H2O(l)  2 Al(OH)4-(aq)

   Silicon dioxide has little acid-base activity, but does show weak acidic properties by dissolving in hot
    concentrated alkalis to form silicates (important in earth science)



Far Right

   Molecular bonding occurs

   Gases, liquids or low melting point solids

   Often a variety of oxides exist, containing element in different oxidation states

   Non-metal oxides react with water to form acids



                                 E.g. SO3(s) + H2O(l)  H+(aq) + HSO4-(aq)



    7. d – block elements

   occur in the central block of the periodic table (d subshell being filled)
   scandium to zinc


   transition metals are those in which the element has a partially filled d-electron shell in one
    of its common oxidation states



   due to partially filled subshell, transition metals have common characteristics not shared by
    other metals
    1. variety of oxidation states
    2. ability to form complex ions
    3. colored ions
    4. catalytic activity
   A. Variable Oxidation States

s-block metals (Na, Ca) have s-electrons that are easily lost, but ionization energies of inner
electrons very high so these electrons are never lost, therefore they always have same oxidation
states


Transition metals have higher effective nuclear charge – first ionization energies higher.
However, no sudden large increase in ionization energy.

The oxidation state (i.e. whether iron would form a +2 or +3 ion) therefore depends strength of
oxidizing environment



Oxidizing agent – an atom, molecule, or ion that causes an increase in the oxidation state of
another substance and is itself reduced.



Most common oxidation state of transition metals is +2 state – loss of the two 4s electrons
       very stable on the right of the d-block – becomes an increasingly powerful reductant
         going to the left.


Second most common oxidation state (as far as Mn) is loss of 3d and 4s valence electrons e.g.
Sc+3, Ti+4, V+5, Cr+6, Mn+7

Stable on left of d block but become increasingly powerful oxidants going across period

Highest oxidation states usually occur as oxanions, such as dichromate (Cr2O72-) and
permanganate (MnO4-)



Commonly encountered states that do not fit into this pattern are Cr +3 and Fe+3
B. Formation of Complex ions


   Ions of d-block metals and those in lower section of p-block (e.g. Pb) have low energy
    unfilled d- and p-orbitals.
   These orbitals can accept a pair of electrons from some species, known as ligands to form a
    dative bond between the ligand and the metal ion.


Ligands are species that can donate a pair of electrons to the central metal ion

       Most common examples are water, ammonia (NH3), and chloride ion.



Species which contain ligands bonded to a central metal ion are known as complex ions

e.g. [Cu(NH3)4]2+ [Ag(NH3)2]+



Most complex ions have either
   six ligands arranged octahedrally around the central atom (usually H2O or NH3 ligands)
   four ligands arranged tetrahedrally (Cl-)


Complex ions stabilize certain oxidation states and can affect solubility of compounds




C. Coloured Ions



   In an isolated atom, all d-orbitals have exact same energy
   However, if an atom or ion is surrounded by charged or polar ions, the electric field has
    different effect on different d-orbitals
   Because of symmetry of orbitals, orbitals split up into two different groups, usually three
    orbitals at low energy and two orbitals at higher energy *see p. 365
   When white light passes through a compound of a transition metal, light of a particular
    frequency is absorbed and an electron is promoted form a lower energy orbital to a higher
    energy orbital




   Light that passes through therefore appears colored because some of the frequencies have
    been absorbed.
        E.g. in most Copper (II) compounds, yellow and red light is absorbed, so they look a
        blue-green color

           Exact shade depends on species surrounding the copper


   If there aren’t any electrons in the d-orbitals, Sc+3 and Ti+4, there are no electrons to move
    and the compounds are colorless.
   Same is also seen if d-orbitals are completely filled, Zn+2




D. Catalytic activity



Catalytic activity caused by

1. Forming complex ions that can donate a lone pair of electrons
   This results in close contact between different ligands, as well as between the metal and the
   ligand
2. Transition metals have wide variety of possible oxidation states, so they can readily lose and
   gain electrons in redox reactions




2 types of d-block catalysts

    1. heterogeneous catalysts – catalysts are in a different phase than the reactants

           surface of the metal or compounds provides an active surface that provides a reduced
           activation energy
                                                  MnO2

                                       e.g. H2O2  H2O + O2



                                                   Ni


                                         CH2=CH2  CH3-CH3

*** Haber Process – see text pg. 368 (simplified) & pg. 198 ***

*** Contact process – see text pg. 203 ***



    2. homogeneous catalysts – catalysts in the same phase of reactants

               particular metal is oxidized in one stage and then reformed by being reduced in
                the second stage


       e.g. iron (II/III) catalysing the reaction between H2O2 and I-
        iron (II) oxidized by the peroxide to iron (III), which is then reduced by the iodide
           ions to reform iron (II)




    E. Chemistry of Selected d-block elements

Manganese
[Ar]3d54s2

common oxidation states are +2, +4, +6, +7

   most common manganese (VII) compound is MnO4- which has an intense purple color
   very powerful oxidant with other oxidation states formed by reduction under various
    conditions (pH)


              in strongly alkaline conditions, dark green manganate (VI) ion is the product
                   MnO4- + e-  MnO42-

              In neutral solution, brown precipitate of manganese (IV) oxide is formed
              MnO4- + 4 H+ + 3 e-  MnO2 + 2 H2O
          In acidic solution, almost colorless manganese (II) ion is the product
               MnO4- + 8 H+ + 5 e-  Mn2+ + 4 H2O




Iron

[Ar] 3d6 4s2

   common oxidation states are +2, +3
   iron(II) compounds are pale green in color
   iron (III) compounds vary from yellow to brown in color
   iron (II) acts as a mild reductant
   iron (III) acts as a mild oxidant


       Fe3+ + e-  Fe2+



   Oxidation state of iron in one of its compounds can be detected by adding aqueous alkali
    until a precipitate of hydroxide forms.


Fe2+(aq) + 2 OH-(aq)  Fe(OH)2 (s) green precipitate



Fe3+(aq) + 3 OH-(aq)  Fe(OH)3 (s) red-brown precipitate

				
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