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					                  Review - Element Properties
            Critical Atomic Properties
                Electron Configuration
                Atomic Size
                Ionization Energy
                Electronegativity




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                Review - Element Properties
            Electron Configuration
              (nl#) – The distribution of electrons in the energy
                levels and sublevels of an atom
                 n – principal quantum # indicating the energy and
                   distance of orbitals from the nucleus
                    Higher value of “n” means higher energy and
                      greater distance from nucleus
                 l – Angular Momentum
                    indicates shape of orbitals
                    Values: l = 1 → n -1
                 ml – Magnetic Quantum #
                    Values: -l …. 0 …..+l
                                               ml = 0                 (s orbital
                                               ml = -1 0 +1           (p orbitals)
                                               ml = -2 -1 0 1 2       (d orbitals)
                                               ml = -3 -2 -1 0 1 2 3 (f orbitals)
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              Review - Element Properties
        Electron Configuration
                           Quantum Numbers
      Name        Symbol      Permitted           Property
                              Values
      principal    n       positive integers   orbital energy
                             (1, 2, 3, …)          (size)

      angular      l       integers from       orbital shape
      momentum                 0 to n -1
                                               The l values
                                                  0, 1, 2, and 3
                                               correspond to
                                                  s, p, d, and f orbitals,
                                               respectively

      magnetic     ml      integers from       orbital (x,y,z)
                           -l to 0 to +l       orientation

      spin         ms      +1/2 or -1/2        e- spin orientation
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             Review - Element Properties




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                Review - Element Properties
            Electron Configuration
              Ground state – Lowest Energy Level
              “s” block elements (groups 1 & 2)
              “p” block elements (groups 3, 4, 5, 6, 7, 8)
              Outer electron configurations (valence electrons) are
               similar within a “Group”
              Outer electron configurations (valence electrons) are
               different within a “Period”
              Outer electrons occupy the “ns” and “np” sublevels
              Four valence-level orbitals (one “ns” and 3 “np”) occur
               among the “main-group” elements
              The “A” group number (1A – 8A) equals the number of
               valence electrons



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                 Review - Element Properties
            Electron Configuration
              Outer electrons are shielded from the full nuclear
                charge by electrons in the same level and even more
                so by the electrons in the lower levels
              Shielding reduces attraction of nuclear charge
                resulting in a “reduced effective nuclear charge, Zeff)
              Within a level, electrons that penetrate
                more (closer to nucleus) shield the
                other electrons more effectively
              Extent of penetration
                      S > p > d > f
              Inner (n=1) electrons shield outer
                (n=2) electrons very effectively
              2s electrons spend more time closer to nucleus than
                2p electrons, thus somewhat shielding the 2p
                electrons
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                Review - Element Properties
            Atomic Size
              Atomic size generally decreases left to right across a
               period of periodic chart
              Atomic size generally increases down a group
              Outer electrons in higher periods (n= 1 , 2, etc.) lie
               farther from the nucleus




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                     Review - Element Properties
            Ionization Energy (IE)
                Energy required to remove the highest energy electron from
                 1 mol of gaseous atoms
                Relative magnitude of IE influences the types of bonds
                    Element with a “low” IE is more
                     likely to lose electrons
                    Element with a “High” IE is
                     more likely to gain electrons
                    IE generally increases left to
                     right (Higher Zeff holds electrons
                     tighter)
                    IE generally decreases down a
                     group (greater distance from nucleus
                    Trends in IE are opposite those in atomic size
                    Easier to remove low IE electron that is farther from nuclues

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                 Review - Element Properties
            Electronegativity (EN)
              A number that refers to the relative ability of an atom in a
                covalent bond to attract shared electrons
                 EN generally increases left to right across a period
                    Higher Zeff and shorter distance from the nucleus
                      strengthen the attraction for the shared pair
                 EN generally decreases down a group
                    Greater distance from the nucleus weakens the attraction
                      for the shared pair
              Trends is Electronegativity are opposite those in Atomic size and
                the same as those in Ionization Energy
              The difference in electronegativity (EN) between atoms in a
                bond greatly influences physical & chemical behavior




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                Review - Element Properties
            Atomic Size, Ionization Energy, Electronegativity
              Trends in IE are opposite those in atomic size
              Trends is Electronegativity are opposite those in
               Atomic size and the same as those in Ionization
               Energy




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                Review - Element Properties
            Bonds – Forces that hold atoms together
            Types of Bonding
              The types of bonding, bond properties, nature of
               orbital overlap, and number of bond determine both
               physical and chemical behavior
              3 Types – Ionic, Covalent, Metallic
                Ionic
                  Results from attraction between positive and
                    negative ions
                  Ions arise through the transfer of electrons
                    between atoms with a large EN – metal –
                    metalloid – non-metal
                  Forms crystalline solids with ions packed tightly
                    in regular arrays


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                Review - Element Properties
            Types of Bonding
              Covalent
                Results from the attraction between two nuclei and
                 a “localized” electron pair
                Bond arises through electron sharing between
                 atoms with small EN, usually between 2 non-
                 metals
                Bond forces include:
                  “Strong” covalent bonding forces holding atoms
                    together forming a molecule
                  “Weak” intermolecular forces holding separate
                    molecules together, thus determining the
                    physical properties of covalent compounds
                Produces discrete molecules with specific shapes or
                 extended networks of molecules
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                Review - Element Properties
            Types of Bonding
              Metallic
                Results from the attraction between the cores of
                 metals atoms (metal cations) and their
                 “delocalized” valence electrons
                The bonding arises through the shared pooling of
                 valence electrons from many atoms and leads to
                 crystalline solids




                    Ionic             Covalent          Metallic
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                     Review - Element Properties
            Bond Overlap
                Actual bonding in real substances lies between the distinct ionic,
                 and decreasingly polar covalent models
                Electron Density Maps below showaa:
                    Small overlap region for the ionic bond in NaCl
                    Increase overlap for slightly polar covalent SiCL bond in SiCl4
                    Highest overlap for non-polar covalent bond in ClCl molecule




                      Ionic               Slightly Polar           Non-Polar


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               Review - Element Properties
         Continuum of Bond Types among Period
          3 Elements
           Left Side
             Chlorine compounds display a
              gradual change from ionic to covalent
              from top to bottom
             Decrease in ionic character (bond
              polarity) from bottom to top
           Right Side
             Elements themselves display a
              gradual change from covalent to
              metallic from top to botton
           Along the Base
             Compounds of each element display
              gradual change to metallic bonding
              from left to right
             Decrease in bond polarity (ionic
              character) from left to right
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                Review - Element Properties
            Bond Properties
              There are two (2) important properties of a covalent
               bond
                Bond Length – Distance between the nuclei of
                  bonded atoms
                Bond Energy (Bond Strength) – The Enthalpy
                  change (H) required to break a given bond in 1
                  mole of gaseous molecules
              As bond length increases, bond energy decreases,
               i.e., short bonds are the stronger bonds
              As bond energy decreases, reactivity increases




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              Review - Element Properties
        Orbital Overlap
          In a covalent bond, the shared electrons reside in the entire
           region composed of the overlapping orbitals of the two atoms
          Orbitals overlap in two (2) ways
            End-to-End
                s, p, and hybrid orbitals lead to sigma () bonds
                Electron density distributed symmetrically along bond
                  axis
                Single bond is a  bond
            Side-to-Side
                p with p (or p with d ) leads to a pi () bond
                Electron density distributed above and below bond axis
                A double bond consists of one  bond and one  bond
                A  bond restricts rotation around the bond axis,
                  allowing for different spatial arrangements of the atoms,
                  thus, different compounds

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                  Review - Element Properties
            Orbital overlap (con’t)
              Side-to-Side (con’t)
                Pi bonds are often sites of reactivity


                 CH2=CH2(g) + HCl(g) → CH3-CH2-Cl(g)




                Bond Order – ½ the number of electrons shared
                  Bond Order 1 – single bond
                  Bond Order 2 – double bond
                  Bond Order 3 – triple bond
                  Fractional bond Order – Occurs when a molecule has
                   resonance structures for species with adjacent single
                   and double bonds
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                 Review - Element Properties
            Number of Bonds and Molecular Shape
              The shape of a molecule is defined by the positions of the
               nuclei of the bonded atoms
              The VSEPR theory describes the number of bonding and
               non-bonding electron groups in the valence level of a
               “Central” atom
                Molecular Notation:
                 A       – The Central Atom (Least Electronegative atom)
                 X       – The Ligands (Bonding Pairs)
                 a       – The Number of Ligands          AXaEb
                 E       – Non-Bonding Electron Pairs
                 b       – The Number of Non-Bonding Electron Pairs
              Double & Triple Bonds count as a “single” electron pair
              The Geometric arrangement is determined by:
                                      sum (a + b)

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             Review - Element Properties
                       a=2
                                              AX6E0    a=6
              AX2E0    b=0       Linear        or      b=0
                                                               Octahedral
                                             AX5X1E0   a+b=6
                       a+b=2


                       a=3
                                Trigonal      AX5N0    a=5
              AX3E0    b=0                                       Trigonal
                                 Planar        or      b=0
                                                               Bipyramidal
                       a+b=3                 AX4X1E0   a+b=5



                       a=4
              AX4E0                           AX4E1    b=4
                                                                 Trigonal
               or      b=0     Tetrahedral     or      a=1
                                                               Bipyramidal
             AX2X2E0                         AX3X1E1   a+b=5
                       a+b=4



                       a=2                    AX6E0    a=6
              AX2E2    b=2     Tetrahedral     or      b=0
                                                               Octahedral
                                             AX4X2E0   a+b=6
                       a+b=4


                       a=5
                                                       a=2
                                 Trigonal                        Trigonal
              AX5E0    b=0                             b=3
                               Bipyramidal   AX2E3             Bipyramidal
                                                       a+b=5
                       a+b=5

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                 Review - Element Properties
            Metallic Behavior
              Elements are often classified as:
                Metals
                Metalloids
                Non-metals




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              Review - Element Properties
        Metals, Metalloids, Non-Metals
          Metals lie in the lower-left
           portion of the Period table
          Non-metals lie in the upper-right
           portion of the table
          Metalloids lie between the
           metals and non-metals
            Intermediate values of
              Atomic size
              Ionization Energy
              Electronegativity
              Shiny solids with low
                 conductivity
                                        -
              React cation-like (lose e )
                 with nonmetals
              React anion-like (gain e-)
                 with metals
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                 Review - Element Properties
            Metals, Metalloids, Non-Metals
              Metallic Behavior
                Metallic Behavior changes
                  gradually among elements
                Metallic behavior parallels
                  atomic size:
                   Larger members of group
                     (bottom) or period (left) are
                     more metallic
                   Smaller members are less
                     metallic
                Metals & non-metals typically
                  form crystalline compounds
                  when they react with each
                  other



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                 Review - Element Properties
            Metals, Metalloids, Non-Metals
              Metallic Character
                Ionic size and charge determine the packing in the solid
                Ionic size increases down a group
                Ionic size decreases left to right across period
                Cations are smaller than their parent atoms
                Anions are larger than their parent Atoms
                Anions are much large than cations




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                Review - Element Properties
            Acid-Base Behavior
              Oxides are known for almost all elements
              The metallic behavior of an element corresponds with
               the acid-base behavior of its oxide in water
              Acids:
                produce Hydrogen (H+) ions (Hydronium H3O- ions)
                  when dissolved in water
                React with bases to form a salt & water
              Bases:
                Produces OH- ion in water
                React with acids to form a salt & water




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                 Review - Element Properties
            Acid-Base Behavior (con’t)
              The chart below shows that the electronegativity and
                metallic behavior determine the type of bonding between
                the metal (E) and oxygen (O) in the metal oxide
              Elements with low Electronegativity (EN) (metals) form
                basic oxides
              Elements with high EN (non-metals) form acidic oxides
              Elements with intermediate EN (some metalloids and
                metals) form amphoteric oxides




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                 Review - Element Properties
            Acid-Base Behavior (con’t)
              Oxide Acidity
                 Increases left to right across a period
                 Decreases down a group
                 Acidity trends are opposite the trends in metallic behavior and
                   atomic size
              When an element forms two oxides, the element has a higher
                oxidation number in the more acidic oxide
                                                                               Increasing Acidity →




                                                        Increasing Acidity →
      Example:

              SO2 forms weak acid, H2SO3
                     (O.N. S = +4)
      whereas SO3 forms strong acid, H2SO4
                    (O.N. S = +6)



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                     Review - Element Properties
            Oxidation-Reduction - The relative ability of an element to lose or
             gain electrons when reacting with other elements
                Oxidation Number (O.N.) (also called Oxidation State)
                    O.N. for elements in native state = 0
                    O.N. = the number of electrons that have shifted away from
                     the atom (positive O.N.) or toward it (negative O.N.)
                    An oxidation-reduction (redox) reaction occurs when the O.N.
                     values of any atom in the reactants are different from those in
                     the products
                    All reactions that involve an elemental substance (native
                     element) involve an oxidation-reduction reaction
                                      2K + Cl2 → 2KCl
                    All combustion reactions (reactions with elements oxygen, i.e.,
                     burning)
                                CH4 + 2O2 → CO2 + 2H2O

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                 Review - Element Properties
                Oxidation is the loss of electrons
                Reduction is the gain of electrons
                Reducing Agent – loses electrons (is oxidized, attains more
                 positive O.N.)
                Oxidizing agent – gains electrons (is reduced, attains more
                 negative O.N.)
                Elements with low IE and low EN (groups 1A and 2A) are
                 strong reducing agent
                Elements with high IE and high EN (/groups 7A and oxygen in
                 Group 6A are strong oxidizing agents




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                 Review - Element Properties
            Oxidation State of the Main-Group Elements
              Oxidation State – A number equal to the magnitude of the
               “Charge” an atom would have if its shared electrons were
               held completely by the atom that attracts them
              The highest (most positive) state in a group equals the A
               group number after all its outer (valence) electrons shift
               toward a more electronegative atom
              Among non-metals, the lowest (most negative) state
               equals the A-group number minus 8
              Non-metals have more oxidation states than metals in the
               same group (oxygen and Fluorine are exceptions)
              Odd-numbered oxidation states are the most common ones
               in odd-numbered groups
              Even-numbered oxidation states are the most common
               ones in even-numbered groups
              Oxidation states differ by units of 2 because electrons are
               lost or gained in pairs
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              Review - Element Properties
        Oxidation State of the Main-Group
         Elements (con’t)
          For many metals and metalloids
            with more than one oxidation
            state (groups 3A – 5A), the
            lower state becomes more
            common down the group
            because the np electrons only
            are lost
          An element with more than one
            oxidation state exhibits greater
            metallic behavior in its lower
            state
            Ex. As3+ (lower state) oxide is
            more basic, more like a metal
            oxide, than is As4+ (higher        Most common state in Bold
            state)

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                Review - Element Properties
            Physical States of Elements
              Physical state – solid, liquid, gas – and heat of phase
               change – vaporization, melting point, etc. – reflect the
               relative strengths of the bonding and/or
               intermolecular forces between the atoms, ions, or
               molecules that make up an element or a compound
              Metals (lower left of periodic chart):
                Solids
                Strong metallic bonding holds atoms in crystalline
                  structures
              Metalloids:
                Along staircase line in table and carbon are solids
                Strong covalent bonding holds atoms together in
                  extensive networks


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                Review - Element Properties
            Physical States of Elements (con’t)
              Lighter non-metals and Group 8A (right side and top
               of periodic chart):
                Gases
                Dispersion forces are weak between molecules such
                  as H2, N2, O2, F2, Cl2 or atoms with smaller, less
                  polarizable electron clouds
              Heavier non-metals
                Liquid (Br2) or soft solids (P4, S8, I2)
                Dispersion forces are stronger between molecules
                  with larger, more polarizable electron clouds




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             Review - Element Properties




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                  Review - Element Properties
            Phase Changes of the Elements
                 Melting Point     Boiling Point    Hfus    Hvap
                Group 1(A)
                  These properties generally increase up the group
                  The smaller the atomic core, the stronger the
                   attraction of the delocalized electrons
                  More energy required to melt a solid, boil a liquid,
                   etc.
                Groups 7A and 8A
                  These properties generally decrease down a group
                  Dispersion forces become stronger with the larger,
                   more polarizable atoms



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                  Review - Element Properties
                Groups 3A to 6A
                  These properties reflect changes in interparticle
                   forces down the group
                    Lower values for molecular non-metals
                    Higher values for covalent networks of metalliods
                      (and carbon)
                    Intermediate values for metals




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                Review - Element Properties
            Physical Properties of Compounds
              Molecular Compounds
                Physical state depends on “intermolecular forces”
                Polar Compounds
                   Dipole-Dipole forces predominate
                Non-Polar Compounds
                   Dispersion forces dominate
                Most Molecular compounds are gases, liquids, or
                  low melting point solids at room temperature




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              Review - Element Properties
        Physical Properties of Compounds (con’t)
          Network Covalent Compounds
            Separate particles absent; strong
              covalent bonds link atoms together
              throughout a “network”
            Extremely high melting points, boiling
              points, Hfus, Hvap
            Ex. Silica – extended arrays of
              covalently bond silica & oxygen atoms
            Ex. Carbon – network of covalently
              bond carbon atoms
               Graphite (soft)
                   flat sheets of hexagonal carbon
                    rings consisting of strong  bonds
                    and delocalized  bonds
               Diamond (hardest known substance)
                   face-centered cubic cell units

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                Review - Element Properties
            Physical Properties of Compounds (con’t)
              Ionic Compounds
                Compounds composed of oppositely charged ions
                Very high Melting Point, Boiling Point, Hfus, Hvap




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              Review - Element Properties
        Physical Properties of Compounds (con’t)
          Hydrogen Bonding
            Hydrogen bonding to N, O, F
            Water (H2O) vs Methane (CH4)
                Water
                   Hydrogen bonding
                   Polar compound
                   Much higher
                       MP, BP, Hfus, Hvap
                   Higher specific Heat
                   Higher surface tension
                   Higher viscosity
                Methane
                   No Hydrogen Bonding
                   Non-polar compound
                   Dominated by dispersion
                     forces


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