Quantum Mechanics 7-05 Periodic Variation in Properties of Elements

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					      Quantum Mechanics 7-05:                                                           Groups or chemical families
      Periodic Variation in Properties of Elements                                        x the vertical columns into which elements are arranged.
                                                                                          x group number, for main block element, represents the
      Periodic table                                                                        number of electrons found in the valence shell of atoms
       x groups similar elements together                                                   within the group.
       x helps to focus on similarities and differences among known
         elements
       x helps to predict properties of elements not yet discovered.                    Some group names are:
         that there are periodic variations in physical and chemical                     x 1A    – alkali metals
         behaviors of elements.                                                          x 11A – alkaline earth metals
                                                                                         x VIB. – chalcogens
                                                                                         x VIIB – halogens
                                                                                         x VIII  – inert or noble gases.




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      Periods                                                                           Atomic Radius
       x the horizontal rows into which elements are arranged
       x period number represents the number of shells of electrons                         x effective atomic radius defined as distance from nucleus
         that atoms in a particular period possess.                                           within which 95% of all electron charge density is found
                                                                                            x is measured by determining distance between nuclei of
      Example                                                                                 adjacent atoms
      Oxygen is found in period 2 and group 6                                               x covalent radius is half the distance between nuclei of two
       x has two energy levels of electrons                                                   identical atoms joined by single covalent bond
       x has 6 electrons in the valence shell of each atom.                                 x ionic radius based on distance between nuclei of ions
                                                                                              joined by ionic bond




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      Different types of radii are illustrated below.                                   Examination of periodic table shows that atomic radii (covalent
                                                                                        radii for non metals and metallic radii for metals – see page
                                                                                        347) tend to increase from top to bottom down a group and
                                                                                        decrease from left to right across a period.




                                                                                              Atomic radii are measured in picometers (1 pm = 10-12m)



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      The radius of atom is influenced by attraction of outermost                       Example
      electrons towards nucleus or atomic core of atom.                                 Consider the element Li
                                                                                         x Li has electron configuration of 1s22s1
      Core Charge                                                                        x The two 1s electrons are in core (everything but valence
      When discussing trends, convenient to use concept of core                            shell)
      charge.                                                                            x One 2s electron is in valence shell
        x Core charge represents charge experienced by outermost                         x 2s electron sees nucleus containing 3 protons through
          electrons, assuming inner electrons shield the nucleus                           screen created by two core electrons
          from view.                                                                     x 2s electron sees a core charge of (3 – 2) =1. The two core
                                                                                           electrons effectively neutralize the charge on two of the
                                                                                           protons.

                                                                                        Core Charge = Z(atomic number) - # inner (core) electrons




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      Core charge helps explain some trends                                              Effective Nuclear Charge

                                                                                             x Core charge assumes inner electrons shield the nuclear
   From left to right across a period                                                          charge completely from valence shell electrons.
     x elements have the same number of electron shells                                      x In reality, the effective shielding is not 100%.
     x core charge increases                                                                 x In additions, electrons in the same shell (energy level)
     x predicted decrease in atomic radius does occur                                          repel one another which further reduces force of attraction
                                                                                               between valence shell electrons and the atomic core.

      From top to bottom down a group core
        x charge remains constant
        x number of energy levels of electrons increases
        x atomic radius increases.




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      Effective Nuclear Charge used rather than core charge when                             Chemical properties of any atom are determined by
      discussing valence shell attractions because                                           configuration of the atom’s valence electrons. Stability of
            x refers to nuclear charge experienced by an electron in                         these outermost electrons is reflected directly in the atom’s
              atom after shielding effect of other electrons is                              ionization energy.
              accounted for.
                                                                                             Ionization energy (kJ/mol) is quantity of energy a gaseous
                                                                                             atom must absorb to be able to expel an electron. Obtained
      From left to right across a period                                                     by bombarding gaseous atoms at low pressure with photons
                                                                                             of sufficient energy to eject an electron from atom.
                 x effective nuclear charge increases.                                         x Gaseous atoms specified because an atom in gas phase
                                                                                                 is virtually uninfluenced by its neighbours so no
                 x This explains why atomic radius decreases steadily                            intermolecular forces need to be taken into account.
                   from lithium to fluorine.                                                   x The higher the ionization energy, the more tightly the
                                                                                                 electron is held in the atom, the larger the force of
                                                                                                 attraction and the larger will be the amount of energy
                                                                                                 required to pull electron away.


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             For many electron atom,                                                         Example
              x First ionization energy (I1) is energy needed to remove                         From left to right across a period where the number of
                first electron from the atom in its ground state.                               shells of electrons remains constant

                                            Energy + X(g) Î X+(g) + e-                              x effective nuclear charge is increasing

                 x Second and third ionization energies (I2 and I3)are                              x valence electrons are more strongly attracted from left
                   energies needed to remove 2nd and 3rd electrons                                    to right across a period

                        Energy + X+(g) Î X2+(g) + e- second ionization                              x atomic sizes decrease
                         Energy + X2+(g) Î X3+(g) + e- third ionization
                                                                                                    x first ionization energies increase

          Ionization energy is directly related to effective nuclear                                Experimental data agree with these assertions.
          charge.



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                                                                                             First Ionization Energy

                                                                                             Consider the first ionization curve (I1) for period 2 (Li to Ne)
                                                                                              x from left to right across the period, I1 values increase
                                                                                              x this occurs because effective nuclear charge increasing
                                                                                                while number of shells remains constant
                                                                                              x dip between Be and B explained because Be atom has a
                                                                                                filled 2s orbital while in second shell of B, electron
                                                                                                configuration is 2s2p1. The p orbital is higher in energy
                                                                                                and electrons in this energy level easier to remove
                                                                                              x dip between N and O explained because second shell in
                                                                                                nitrogen is 2s22px1py1pz1 while that in O is 2s22px2py1pz1
                                                                                                In this case, transition occurring from exactly half-filled p-
                                                                                                orbitals to one in which one extra electron is present.
                                                                                                The half filled p orbitals are more stable and electrons
                                                                                                are more difficult to remove.

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             2nd Ionization Energies                                                                Why is 1st ionization energy of Be > 1st ionization energy
                                                                                                    of Li, while 2nd ionization energy of Be+ < 2nd ionization
             Again consider following graph. Blue line represents 2nd                               energy of Li+?
             ionization energies.                                                                     x Both Li and Be have two shells of electrons with Be
                                                                                                        having a greater effective nuclear charge. Therefore
                                                                                                        it is easier to remove a valence electron from Li. This
                                                                                                        explains its lower 1st ionization energy
                                                                                                      x When 2nd electron is removed from Li+ and Be+
                                                                                                        cations, electron is being removed from 1st shell of Li+
                                                                                                        but from 2nd shell of Be+.
                                                                                                      x There is a large jump in difficulty in removing electron
                 x 2nd ionization energies higher than 1st ionization                                   from 1st shell of Li+ as compared to energy required
                   energies for every element                                                           to remove the second 2nd shell electron from Be+
                 x makes sense because more difficult to remove an                                    x Because many factors involved in explaining these
                   electron from a 1+ cation than from neutral atom. Also,                              relative energy conversions, it is dangerous to carry
                   valence electrons of cations closer to nucleus than                                  this simplified argument too far
                   valence electrons of atoms

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                    Electron Affinity                                                           Electron affinity
                                                                                                      x Measures how strongly electron attracted to neutral
                    Electron affinity (EA) is the negative of energy change                             atom
                    that occurs when a gaseous atom gains an electron                                 x High electron affinity means that neutral atom has
                    (Different than textbook definition).                                               strong attraction for electron – electron difficult to
                                                                                                        remove from anion
                                   Cl(g) + e- Î Cl-(g)
                                                     ¨H = -349 kJ/mol                                 x Periodic trends largely parallel those in ionization
                                                    EA = 349 kJ/mol                                     energy, increasing across a period to become large
                    Therefore, electron affinity is 349 kJ/mol                                          and positive for halogen atom before falling
                                                                                                        abruptly to a negative value for noble gas atom.
                    Electron affinity is the same as the electron                                     x Negative value means electron must be forced to
                    detachment energy when an anion is converted to a                                   “stick” to atom and will spontaneously fly away.
                    gaseous atom.                                                                       This indicates that existing noble gas electron
                                                                                                        configuration is very stable without addition of extra
                                Cl-(g) Î Cl(g) + e- ¨H = EA = 349 kJ/mol                                electron



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      General Trends in Elemental Properties                                             Group 1A Elements – Alkali Metals (ns1, n• 2)

      Trends on the periodic table are summarized below.




                                                                                             x Have low ionization energies and great tendency to lose
                                                                                               the single valence electron
                                                                                             x So reactive that never found in pure state in nature
                                                                                             x Unipositive ions of these often found in salts (NaCl)
                                                                                             x Reactivity increases down group because valence electron
                                                                                               is further from nucleus and less tightly held



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          Group IIA Elements – Alkaline Earth Metals (ns2, n• 2)                         Group 111B (ns2np1, n•2)




          x Somewhat less reactive than alkali metals because
            valence electrons held more tightly because effective                            x Boron is a metalloid (an element with properties
            nuclear charge greater                                                             somewhere between those of metals and those of non
          x Tend to form M2+ ions                                                              metals) and does not form binary ionic compounds
          x Metallic characteristics increase (tend to lose valence shell                    x Other elements in this group form +3 ions and ionic
            electrons easier) from top to bottom down a group because                          compounds such as AlCl3
            electrons held less strongly as number of electron shells                        x Boron tends to act more like a non metal, not forming
            increases                                                                          cations easily, while aluminum has metallic characteristics

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          From top to bottom down a group, there is transition from                      Group 1VB (ns2np2, n•2)
          non metals to metalloids to metals in the p-orbital segment of
          the periodic table.




                                                                                             x   Carbon is a non metal
                                                                                             x   Both silicon and germanium are metalloids
                                                                                             x   Carbon, silicon and germanium do not form ionic bonds
                                                                                             x   Tin and lead exhibit metallic properties and can easily form
                                                                                                 cations




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      Group VB (ns2np3, n•2)                                                             Group VIB chalcogens (ns2np4, n•2)




                                                                                             x Oxygen, sulfur and selenium are non metal
             x   Nitrogen and phosphorus are non metals                                      x Tellurium and polonium are metalloids
             x   Arsenic and antimony are metalloids
             x   Bismuth is a metal but is far less reactive than most
             x   From top to bottom down this group there is a significant
                 change in the properties of the elements


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                                                                                         Group VIII –Inert or Noble Gases - (ns2np6, n•2)
                                                          2     5
      Group VIIB – halogens - (ns np , n•2)




                                                                                             x All noble gases exist as monatomic species
          x Are nonmetals with the general formula X2                                        x Have completely filled ns and np subshells which gives
          x Very reactive so never found in elemental form in nature                           elements great stability
          x Fluorine is the most reactive                                                    x Ionization energies are high and these elements have little
          x Halogens have high ionization energies and large positive                          tendency to gain or lose electrons
            electron affinities
          x Need one electron to fill outer shell so tend to form -1
            anions

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Suggested Homework
 x Read Petrucci 339 347 and 349 - 361 560
 x Look at Examples 9-1, 9-3, 9-4
 x Problems Chapter 9: 3, 9, 11, 13, 15, 21, 23, 29, 41, 43 (1, 2 &
   4)




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