periodicity by xiaohuicaicai

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									Periodic Relationships Among
        the Elements
                                   Chapter 8

  Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
When the Elements Were Discovered




                                    8.1
                                                                                                 ns2np6
ns1        Ground State Electron Configurations of the Elements




                                                   ns2np1

                                                            ns2np2
                                                                     ns2np3

                                                                              ns2np4
                                                                                       ns2np5
       ns2




                                             d10
             d1




                           d5




      4f
      5f

                                                                                                8.2
Classification of the Elements




                                 8.2
      Electron Configurations of Cations and Anions
                Of Representative Elements

 Na [Ne]3s1          Na+ [Ne]
                                  Atoms lose electrons so that
 Ca [Ar]4s2         Ca2+ [Ar]     cation has a noble-gas outer
                                  electron configuration.
 Al [Ne]3s23p1      Al3+ [Ne]


                          H 1s1         H- 1s2 or [He]
Atoms gain electrons
so that anion has a       F 1s22s22p5   F- 1s22s22p6 or [Ne]
noble-gas outer
                          O 1s22s22p4   O2- 1s22s22p6 or [Ne]
electron configuration.
                          N 1s22s22p3   N3- 1s22s22p6 or [Ne]
                                                             8.2
     Cations and Anions Of Representative Elements




                                     +3
+1
     +2




                                           -3
                                                -2
                                                     -1
                                                          8.2
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.

     Zeff = Z - s        0 < s < Z (s = shielding constant)

    Zeff  Z – number of inner or core electrons

                    Z       Core    Z
                                   → eff    Radius (pm)

        Na          11       10      1          186

        Mg          12       10      2          160

        Al          13       10      3          143

         Si         14       10      4          132
                                                                8.3
           Periodic Properties
• Elements show gradual changes in certain
  physical properties as one moves across a
  period or down a group in the periodic table.
  These properties repeat after certain
  intervals. In other words they are PERIODIC

                          Periodic properties
                          include:
                          -- Ionization Energy
                          -- Electronegativity
                          -- Electron Affinity
                          -- Atomic Radius
                          -- Ionic Radius
                                                  .8
              Atomic Size
• The electron cloud doesn’t have a definite
  edge.
• They get around this by measuring more
  than 1 atom at a time.
• Summary: it is the volume that an atom
  takes up
• http://www.mhhe.com/physsci/chemistry/e
  ssentialchemistry/flash/atomic4.swf
8.3
Atomic
Radius

• The radius increases on going down a group.
• Because electrons are added further from the
  nucleus, there is less attraction. This is due to
  additional energy levels and the shielding
  effect. Each additional energy level “shields”
  the electrons from being pulled in toward the
  nucleus.
• The radius decreases on going across a
  period.
                                                      .11
   The Electron Shielding Effect
• Electrons
  between the
  nucleus and
  the valence
  electrons repel
  each other
  making the
  atom larger.




                                   .12
  Group trends
                         H
• As we go down a        Li
  group (each atom
  has another energy     Na
  level) the atoms get
  bigger, because         K
  more protons and
  neutrons in the
  nucleus                 Rb
• The radius decreases across a period owing to
  increase in the positive charge from the protons.

• Why? Stronger attractive forces in atoms (as you go
  from left to right) between the opposite charges in the
  nucleus and electron cloud cause the atom to be
  'sucked' together a little tighter. Remember filling up
  same energy level, little shielding occurring.

• Each added electron feels a greater and greater +
  charge because the protons are pulling in the same
  direction, whereas the electrons are scattered.




     Large            All values are in        Small
                      nanometers
                                                            .14
Atomic Radius




                .15
Atomic Radius




                .16
8.3
Atomic Radii




               8.3
Trends in Ion Sizes
    Radius in pm




                      .19
Comparison of Atomic Radii with Ionic Radii




                                              8.3
               Ionic Size
• Cations form by losing electrons.
• Cations are smaller than the atom they
  come from. The electron/proton attraction
  has gone Up and so the radius Decreases.
• Cations like atoms increase as one
  moves from top to bottom in a group.
• Metals form cations.
• Cations of representative elements have
  noble gas configuration.
               Ionic size
• Anions form by gaining electrons.
• Anions are bigger than the atom they
  come from. The electron/proton attraction
  has gone Down and so size Increases.
• Trends in ion sizes are the same as atom
  sizes.
• Nonmetals form anions.
• Anions of representative elements have
  noble gas configuration.
                Periodic Trends
• Metals - losing from outer energy level, more
  protons than electrons so more pull, causing it to
  be a smaller species.
• Non metals gaining electrons in its outer energy
  level, but there are less protons than electrons in
  the nucleus, so there is less pull on the protons,
  so found further out making it larger.

             B+3         N-3        O-2      F-1
  Li+1

         Be+2    C+4
       Size of Isoelectronic ions
• Positive ions have more protons so they
  are smaller.


                                   N-3
                          O-2
               Ne   F-1
Al+3    Na+1

   Mg+2
Cation is always smaller than atom from
which it is formed.

Anion is always larger than atom from
which it is formed.
                                          8.3
     Trends in Ionization Energy
Ionization energy is the energy required to
remove an electron from an atom
• Metals lose electrons
  more easily than
  nonmetals.
• Nonmetals lose electrons
  with difficulty. (They like
  to GAIN electrons).
• Ionization energy
  increases across a period
  because the positive
  charge increases.
                                              .26
  Trends in Ionization Energy
• The ionization energy is
  highest at the top of a
  group. Ionization energy
  decreases as the atom
  size increases.
• This results from an
  effect known as the
  Shielding Effect




                                .27
Ionization Energies of the
 Representative Groups




                             .28
Ionization energy is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground
state.

  I1 + X (g)    X+(g) + e-         I1 first ionization energy

  I2 + X+(g)    X2+(g) + e-        I2 second ionization energy

  I3 + X2+(g)    X3+(g) + e-       I3 third ionization energy




                             I1 < I 2 < I3


                                                                 8.4
8.4
Variation of the First Ionization Energy with Atomic Number


         Filled n=1 shell
              Filled n=2 shell


                    Filled n=3 shell
                                       Filled n=4 shell
                                                    Filled n=5 shell




                                                                       8.4
              Electronegativity
  Electronegativity is a
  measure of the ability
  of an atom in a
  molecule to attract
  electrons to itself.
• How fair it shares.
• Big electronegativity
  means it pulls the
  electron toward it.
     This concept was first proposed by Linus
     Pauling (1901-1994). He later won the Nobel
     Prize for his efforts.
                                                   .32
Periodic Trends:
Electronegativity
      • In a group: Atoms with fewer
        energy levels can attract
        electrons better (less shielding).
        So, electronegativity increases
        UP a group of elements.

      • In a period: More protons, while
        the energy levels are the same,
        means atoms can better attract
        electrons. So, electronegativity
        increases RIGHT in a period of
        elements.
                                        .33
Electronegativity is the ability of an atom to attract
toward itself the electrons in a chemical bond. (Desire
to gain electrons)

                  X (g) + e-    X-(g)


       Electronegativity - relative, F is highest - or most
       electronegative element




                                                         9.5
The Electronegativities of Common Elements




                                             9.5
Variation of Electronegativity with Atomic Number




                                                    9.5
               Group Trend
• The further down a group the farther the electron
  is away and the more electrons an atom has.
• So as you go from fluorine to chlorine to bromine
  and so on down the periodic table, the
  electrons are further away from the nucleus
  and better shielded from the nuclear charge and
  thus not as attracted to the nucleus. For that
  reason the electronegativity decreases as you
  go down the periodic table.
             Period Trend
• Electronegativity increases from left to
  right across a period
• When the nuclear charge increases, so
  will the attraction that the atom has for
  electrons in its outermost energy level and
  that means the electronegativity will
  increase
Summary of Periodic Trends
Melting Point

-based upon types of intermolecular forces

-higher mp with metallic bonds (strong
intermolecular forces,

-network solid very high mp (very strong bonds
between atoms forming large molecules),

-covalent bonds lower mp (weak
intermolecular forces)
Melting Points of Group 1
  Element   Melting Point (K)

  Li        453
  Na        370
  K         336
  Rb        312
  Cs        301
  Fr        295
    Melting Points for halogens
Element          Melting Point (K)

Fluorine         85

Chlorine         238

Bromine          332

Iodine           457

Astatine         610
Melting Point
• the temperature at which a solid changes
  to a liquid

• Trends within
• a. alkali metal: MP DECREASES down the
  family/group
• b. halogens: MP INCREASES down the
  family/group
            Metallic bonding
• Collective bond, not a single bond
• Strong force of electromagnetic attraction
  between delocalized electrons (move freely).
• This is sometimes described as "an array of
  positive ions in a sea of electrons
Explanation
• Generally - MP depends upon the strength of forces
  holding atoms or molecules together. The stronger the
  IM (intermolecular) force the higher the MP (more energy
  is needed to separate molecules from the solid to the
  liquid phase)

• a. Alkali Metals have metallic bonds between atoms .
  As the size of the atom increases (down the family) the
  metallic bonding weakens and so the MP decreases.

• b. Halogens have Van der Waal’s forces between
  diatomic molecules . The molecules are NP (nonpolar)
  and have relatively weak IM forces. The strength of the
  forces INC with INC molecular mass, so MM increases
  down the family and therefore VdW forces increase, and
  therefore MP increases.
Why does the melting point decrease
  going down the alkali metals family?
• Atoms are larger and their outer electrons
  are held farther away from the positive
  nucleus.
• The force of attraction between the metal
  ions and the sea of electrons thus gets
  weaker down the group.
• Melting points decrease as less heat
  energy is needed to overcome this
  weakening force of attraction.
  Why does melting point increase
    going down the halogens?
• The halogens are diatomic molecules, so
  F2, Cl2, Br2, I2
• As the molecules get bigger there are
  more electrons that can cause more
  influential intermolecular attractions
  between molecules.
• The stronger the I.M. forces, the more
  difficult it will be to melt. (more energy
  needed to break the I.M. forces)
    3.3 Chemical properties
• Reactions of elements within the same
  family
• in general, if the electron arrangement
  determines the chemical reactivity of an
  element, then the members of the same
  family/group can be expected to have
  similar chemical reactivity.
Chemical properties
• Alkali Metals (Li and Na and K)
• most characteristic property is ability to lose an
  electron - they have low ionization energies and
  are very reactive and form ionic solids.

• Reactivity of alkali metals will increase down the
  family as reactive electrons are farther from the
  nucleus and easier to access and react with;
  these elements tend to lose electrons and
  become reducing agents (provide electrons for
  oxidation rxns to occur)
                  Group 1 Elements (ns1, n  2)

 M        M+1 + 1e-

2M(s) + 2H2O(l)            2MOH(aq) + H2(g)

4M(s) + O2(g)       2M2O(s)
                      Increasing reactivity




                                                  8.6
Group 1A Elements (ns1, n  2)




                                 8.6
Alkali Metals
• Physical properties of alkali metals -soft
  malleable metals with low mp and low
  densities.
• Very reactive chemically -- including
  exposure to air and water
Alkali Metals

• Reaction with water
Metal + water → H2(g) + metal hydroxide

• Due to the decrease in Ionization Energy
  of metals moving down the Periodic Table
  the reactivity of the metal INCREASES
  down the table (Li reacts less violently
  than does Na...)
Alkali Metals
• Reactions with halogens (ie. Cl2 and Br2)
• redox reactions to form ionic salts
  2 Na + Cl2 → 2NaCl

• Note: oxidation/reduction reactions (redox)
• 1. oxidation is the increase in oxidation
  number or the loss of electrons
• 2. reduction is the decrease in oxidation
  number or the gain of e-.
Chemical properties
• Halogens (Cl2,Br2 and I2) – Group 17 (VII)

General Properties –
• diatomics, colored, phase changes as one goes
  down the family. Cl2 is gas (green yellow), Br2 is
  liquid (brown/red) and Iodine is a purple solid
• not soluble in water (non polar substance)
  (hence use of oil in experiments-non polar to
  dissolve halogens).

General Reactivity-
• highly reactive due to need for a single electron
  to fill valence shell
Group 7A Elements (ns2np5, n  2)




                                    8.6
Halogens
• Reactivity decreases as one goes down the
  halogen family.

• Halogens will react by adding an electron to
  themselves (they behave as oxidizing agents -
  they are reduced - gain electrons). The smallest
  and most electronegative element F is the most
  reactive.

• Valence electrons that are farther from the
  nucleus will have less attraction and are
  therefore less reactive.
            Group 7A Elements (ns2np5, n  2)

X + 1e-         X-1

X2(g) + H2(g)         2HX(g)




                                                Increasing reactivity
                                                                        8.6
Halide Ions (F-, Cl-, Br- and I-)
• Reactivity
  oxidizing power of the ions decreases
  going down the table (size of atom
  increases and attraction for electrons
  decreases) so Cl will oxidize I but I will not
  oxidize Cl (higher halogen will displace a
  lower halogen from its salts.)
Halide Ions
• Reactions : assume that the halogen is the one
  reacting by removing electrons from the ion,
  therefore if the halogen (diatomic) is higher on
  the table than the ion , the reaction will take
  place, but if the ION is higher on the table than
  the HALOGEN the reaction will not take place.
• Cl2 + 2 I- → I2 + 2 Cl-
• Br2 + 2 I- → I2 + Br-
• I2 + 2 Br- → no rxn
  Properties of Oxides Across a Period


basic                             acidic




                                           8.6
Metal oxide + water → metal hydroxide (base)
     ie. Na2O(s) + H2O(l) → 2 NaOH(aq)

        MgO(s) + H2O(l) → Mg(OH)2(aq)




 Nonmetal oxide + water → acid
       ie. SO3(g) + H2O(l) → H2SO4(aq)

          P4O10(s) + 6 H2O(l) → 4 H3PO4(aq)
Properties of the Third Period
           Oxides
Properties of the Third Period
          Chlorides
The D Block Elements
          • The d block elements
            fall between the s
            block and the p block.
          • They share common
            characteristics since
            the orbitals of d
            sublevel of the atom
            are being filled.
    The D Block Elements
•   The D block elements include the transition
    metals. The transition metals are those d block
    elements with a partially filled d sublevel in one
    of its oxidation states.
•   Since the s and d sublevels are very close in
    energy, the d block elements show certain
    special characteristics including:
               1. Multiple oxidation states
               2. The ability to form complex ions
               3. Colored compounds
               4. Catalytic behavior
               5. Magnetic properties
   The D Block Elements
• The d electrons are close in energy to the s
  electrons.
• D block elements may lose 1 or more d
  electrons as well as s electrons. Hence they
  often have multiple oxidation states

     Some common D block oxidation states
  Multiple Oxidation States
• There is no sudden sharp increase in ionization
  energy as one proceed through the d electrons as
  there would be with the s block.
• D block elements can lose or share d electrons
  as well as s electrons, allowing for multiple
  oxidation states.
• Most d Block elements have a +2 oxidation State
  which corresponds to the loss of the two s
  electrons.
• This is especially true on the right side of the d
  block, but less true on the left.
  ---- For example Sc+2 does not exist, and
     Ti+2 is unstable, oxidizing
     in the presence of any
     water to the +4 state.
            Complex Ions
• The ions of the d block and the lower p block
  have unfilled d or p orbitals.
• These orbitals can accept electrons either an
  ion or polar molecule, to form a dative bond.
  This attraction results in the formation of a
  complex ion.
• A complex ion is made up of two or more ions
  or polar molecules joined together.
• The molecules or ions that surround the metal
  ion donating the electrons to form the complex
  ion are called ligands.
         Complex Ions
• Compounds that are formed with
  complex ions are called coordination
  compounds
• Common ligands



• Complex ions usually have either 4 or 6
  ligands.
          K3Fe(CN)6        Cu(NH3)42+
         Complex Ions
• The formation of complex ions stabilizes
  the oxidations states of the metal ion
  and they also affect the solubility of the
  complex ion.

           »         The formation of a
           »         complex ion often has
           »         a major effect on the
           »         color of the solution of
           »         a metal ion.
 The D Block Colored Compounds
• In an isolated atom all of the d sublevel electrons
  have the same energy.
• When an atom is surrounded by charged ions or polar
  molecules, the electric field from these ions or
  molecules has a unequal effect on the energies of the
  various d orbitals and d electrons.
• The colors of the ions and complex ions of d block
  elements depends on a variety of factors including:
         – The particular element
         – The oxidation state
         – The kind of ligands bound to the element


  Various oxidation
  states of Nickel (II)
         Colors in the D Block
• The presence of a partially filled d sublevels in a
  transition element results in colored compounds.
• Elements with completely full or completely empty
  subshells are colorless,
   – For example Zinc which has a full d subshell. Its
     compounds are white
• A transition metal ion is colored, if it absorbs light in
  the visible range (400-700
  nanometers).
• If the compound absorbs a
  particular wavelength of light its
  color will be the composite of those
  wavelengths that it does not absorb.
• In other words it shows its
   complimentary color.
Colors and d Electron Transitions
• When ligands are attached to transition metal ions, the
  d orbitals may split into two groups. Some of the
  orbitals are at a lower energy than the others
• The difference in energy of these orbitals varies slightly
  with the nature of the ligand or ion surrounding the
  metal ion




• The energy of the transition: ∆E =hn may occur in the
  visible region.
• When white light passes through a compound of a
  transition metal, light of a particular frequency is
  absorbed as an electron is promoted from a lower
  energy d orbital to a higher one.
• The result is a colored compound
Magnetic Properties
   • Paramagnetism --- Molecules with
     one or more unpaired electrons are
     attracted to a magnetic field. The
     more unpaired electrons in the
     molecule the stronger the attraction.
     This type of behavior is called
   • Diamagnetism --- Substances with
     no unpaired electrons are weakly
     repelled by a magnetic field.
   • Transition metal complexes with
     unpaired electrons exhibit simple
     paramagnetism.
   • The degree of paramagnetism
     depends on the number of unpaired
     electrons
Catalytic Behavior
   • Many D block elements are
     catalysts for various chemical
     reactions
   • Catalysts speed up the rate of a
     reaction with out being consumed.
   • The transition metals form complex
     ions with ligands that can donate
     lone pairs of electrons.
   • This results in close contact
     between the metal ion and the
     ligand.
   • Transition metals also have a wide
     variety of oxidation states so they
     gain and lose electrons in
     oxidation- reduction reactions
Some Common D Block Catalysts
  • Examples of D block elements that are
    used as catalysts:
                       1. Platnium or
                           rhodium in a
                           catalytic converter
                        2. MnO2 decomposition
                          of hydrogen peroxide
                        3. V2O5 in the contact
                           process
                        4. Fe in Haber process
                        5. Ni in conversion of
                            alkenes to alkanes
The Periodic Table--Summary

The periodic table is a classification
system. Although we are most
familiar with the periodic table that
Seaborg proposed more than 60
years ago, several alternate designs
have been proposed.
Alternate Periodic Tables
Alternate Periodic Tables II

								
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