11/13/2008




    Anomalous Electron Configurations

   • A few exceptions to the Aufbau
     principles exist. Stable co...
11/13/2008




  An application of Electronic Configuration of
                      Atom
       Locating elements in The ...
11/13/2008




 Electron Configurations can be
 Determined From the Position in
       the Periodic Table:

• Elements in ...
11/13/2008




                  Down the Periodic Table
               •Family: Are arranged vertically down the periodic...
11/13/2008




        Electron Configuration of
                 Cations
• cations are formed when an atom loses all its
...
11/13/2008




      The Periodic Table of
           Elements

                Week -7
             Lecture 13 &14




De...
11/13/2008




                  Arranging the Elements
The elements were first arranged in this way by Dmitri
Mendeleev, ...
11/13/2008




 Introduction of The Periodic Table

– The periodic table is made up of rows of elements
  and columns.
– A...
11/13/2008




                                Atoms & Elements
  Key Concepts:
                                      Elem...
11/13/2008




• Chemical Families
  – IA are called alkali metals because the react with
    water to from an alkaline so...
11/13/2008




Alkaline Earth Metals




  Transition Metals


                               11
11/13/2008




          These elements are also
            called the rare-earth
                        elements.




I...
11/13/2008




       Noble Gases


Reading the Periodic Table: Classification
        •   Nonmetals, Metals, Metalloids, ...
11/13/2008




Metals, Nonmetals, and Metalloids
Metals
•Metallic character refers to the properties of metals
    – Shiny...
11/13/2008




 Properties of Non-Metals
§ They are poor conductors of electrical
  energy




                           ...
11/13/2008




Tro, Chemistry: A Molecular
                              31
Approach




  The Groups of the Periodic Tabl...
11/13/2008




  The Groups of the Periodic Table
• Group 14: The Carbon Family
  – Contains elements that can form unusua...
11/13/2008




Electron Shells and the Sizes of
Atoms
Atomic Sizes
• As a consequence of the ordering in the
  periodic ta...
11/13/2008




                                                           Atomic Radius

Fig. 8.15 Atomic Radii for Main G...
11/13/2008




See Figure 8.16




              Trends in Ionic Radius
    • Ions in same group have same charge
    • Io...
11/13/2008




                              41




Tro, Chemistry: A Molecular
                              42
Approach
...
11/13/2008




             Ionization Energy
• minimum energy needed to remove an
  electron from an atom
   – gas state
...
11/13/2008




Ionization Energy
• The first ionization energy, I1, is the amount of
energy required to remove an electron...
11/13/2008




Tro, Chemistry: A Molecular
                                         47
Approach




     Ionization Energy...
11/13/2008




                               49




  Example – Choose the Atom in Each Pair
   with the Higher First Ion...
11/13/2008




              Irregularities in the Trend
• Ionization Energy generally increases from
  left to right acro...
11/13/2008




                     Irregularities in the
              First Ionization Energy Trends

         ↑↓       ...
11/13/2008




HIGHER IONIZATION ENERGIES




                       Chapter 8-55




            56




                 ...
11/13/2008




Electron Affinities
• Electron affinity is the opposite of ionization
energy.
• Electron affinity is the en...
11/13/2008




              Summary of Trend
          1. Ionization Energy: Largest toward
          3. Electron Affinit...
11/13/2008




Example 8.6 – Write the Electron Configuration
 and Determine whether the Fe atom and Fe3+
     ion are Par...
11/13/2008




  Melting Point Trends in Period

•In the second period, melting points
increase, going from left to right ...
11/13/2008




       Trends in Density

•Densities of elements increase in a
group as atomic number increases.
•In period...
11/13/2008




       34
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2 Slide Lecture 13

  1. 1. 11/13/2008 Anomalous Electron Configurations • A few exceptions to the Aufbau principles exist. Stable configuration: – half-filled d shell: • Cr has [Ar]4s13d5; • Mo has [Kr] 5s14d5 – filled d subshell: • Cu has [Ar]4s13d10 • Ag has [Kr]5s14d10. • Au has [Xe]6s14f145d10 • Exceptions occur with larger elements where orbital energies are similar. Anomalous electron configuration of some elements!!! Element Atomic Expected Actual number configuration Configuration Cr 24 3d4 4s2 3d5 4s1 Cu 29 3d9 4s2 3d10 4s1 Mo 42 4d4 5s2 4d5 5s1 *Pd 46 4d8 5s2 4d10 5s0 Ag 47 4d9 4s2 4d10 5s1 *Pt 78 5d8 6s2 5d9 6s1 Au 79 5d9 6s2 5d10 6s1 The explanation for this deviation lies in the superior stability of completely filled or all half-filled orbitals than half- nearly filled of nearly half-filled orbitals. half- orbitals. ****They are exception to this rule also. So, remember. Some other of this kinds are Nb(41), Ru(44), W(74), Nb(41), Ru(44), Sg(106)**** Sg(106)**** 1
  2. 2. 11/13/2008 An application of Electronic Configuration of Atom Locating elements in The Periodic Table • In The Periodic Table elements are organized on the basis of their electronic configuration • The highest ‘n’ value in electronic configuration determines the period of that element. • No of outer shell e determines the Group § If at s & p orbital of the outer shell have 1 to 7 e (s1 to s2 p5) then they are elements of Group I to VII Ex: Na (11) – 1s22s22p63s1 : Group I Cl (17) – 1s22s22p63s23p5 : Group VII § If at s & p orbital of outer shell have s2 p6 then the element is in group O. Ex: Ne (10) § Again at d & s orbital of outer shell determines the Group Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group III Cont’d § But at d & s orbital of outer shell, if there are 8/9/10 e then it is in the Group VIII Ex: Fe (26) - 1s22s22p63s23p63d64s2 : Group VIII § If you have more then 10 e at d & s of outer shell then only ‘s’ orbital’s e will give the Group Ex: Cu (29) - 1s22s22p63s23p63d104s1 : Group I • ‘A’ Sub Group: if at the outer shell you do not have ‘d’ orbital or ‘d’ is filled only then it in ‘A’ Sub Ex: Cl (17) – 1s22s22p63s23p5 : Group VIIA Ga (31) - 1s22s22p63s23p63d104s24p1 : Group IIIA • ‘B’ Sub Group: if you have e in ‘d’ orbital at outer shell (d1–d10) then it is in ‘B’ Sub Group Ex: Sc (21) - 1s22s22p63s23p63d14s2 : Group IIIB Zn (30) - 1s22s22p63s23p63d104s2 : Group IIB 2
  3. 3. 11/13/2008 Electron Configurations can be Determined From the Position in the Periodic Table: • Elements in group 1(1A) end in ns1 • Elements in group 2 (2A) end in ns2 • Elements in group 13 (3A) end in ns2np1 • Elements in group 14 (4A) end in ns2np2 • Elements in group 15 (5A) end in ns2np3 • Elements in group 16 (6A) end in ns2np4 • Elements in group 17 (7A) end in ns2np5 • Elements in group 18 (8A) end in ns2np6 Valence Electrons • The electrons in all the sub-shells with the highest principal energy shell are called the valence electrons • e in lower energy shells are called core electrons • Starting with one valence electron for the first element in a period, the number of electrons increases as you move from left to right across a period. • chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons 3
  4. 4. 11/13/2008 Down the Periodic Table •Family: Are arranged vertically down the periodic table (columns or group, 1- 18 or 1-8 A,B) 1- 1- •These elements have the same number electrons in the outer most shells, the valence shell. 1 18 IA VIIIA 2 Alkali Family: 13 14 15 16 17 1 IIA 1 e- in the valence shell IIIA IVA VA VIA VIIA 2 3 4 5 6 7 8 9 10 11 12 3 IIIB IVB VB VIB VIIB VIIIB IB IIB 4 Halogen Family: 7 e- in the valence shell 5 6 7 When an atom or molecule gain or loses an e- it becomes an ion • A cation has lost e- and therefore has positive charge • An anion has gained an e- and therefore has a negative charge. Electron Configuration of Anions • Anions are formed when atoms gain enough electrons to have 8 valence electrons – filling the s and p sublevels of the valence shell • The sulfur atom has 6 valence electrons S atom = 1s22s22p63s23p4 • In order to have 8 valence electrons, it must gain 2 more S2- anion = 1s22s22p63s23p6 4
  5. 5. 11/13/2008 Electron Configuration of Cations • cations are formed when an atom loses all its valence electrons – resulting in a new lower energy level valence shell – however the process is always endothermic • the magnesium atom has 2 valence electrons Mg atom = 1s22s22p63s2 • when it forms a cation, it loses its valence electrons Mg2+ cation = 1s22s22p6 Electron Configuration of Cations • for transition metals electrons, may be removed from the sublevel closest to the valence shell Al atom = 1s22s22p63s23p1 Al+3 ion = 1s22s22p6 Fe atom = 1s22s22p63s23p64s23d6 Fe+2 ion = 1s22s22p63s23p63d6 Fe+3 ion = 1s22s22p63s23p63d5 Cu atom = 1s22s22p63s23p64s13d10 Cu+1 ion = 1s22s22p63s23p63d10 5
  6. 6. 11/13/2008 The Periodic Table of Elements Week -7 Lecture 13 &14 Development of the Periodic Table • There were 114 elements known by 1999. • The majority of the elements were discovered between 1735 and 1843. • How do we organize 114 different elements in a meaningful way that will allow us to make predictions about undiscovered elements? • Arrange elements to reflect the trends in chemical and physical properties. • First attempt (Mendeleev and Meyer) arranged the elements in order of increasing atomic weight. 6
  7. 7. 11/13/2008 Arranging the Elements The elements were first arranged in this way by Dmitri Mendeleev, a professor at St. Petersburg University, in 1869. His arrangement was based on atomic mass. When Mendeleev was setting out the table, only 63 elements had been discovered. His big idea was to leave gaps for yet to be discovered elements. He was able to predict the properties of some of these elements, including silicon and boron. When his predictions were shown to be accurate his table became accepted, and it is the basis of the one we use today. The Father of the Periodic Table — Dimitri Mendeleev • Mendeleev was the first scientist to notice the relationship between the elements – Arranged his periodic table by atomic mass • Moseley later discovered that the periodic nature of the elements was associated with atomic number, not atomic mass • The Periodic Law – When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties. 7
  8. 8. 11/13/2008 Introduction of The Periodic Table – The periodic table is made up of rows of elements and columns. – An element is identified by its chemical symbol. – The number above the symbol is the atomic number – The number below the symbol is the atomic weight of the element. – A row is called a period – A column is called a family or group – Elements are arranged left to right and top to bottom in order of increasing atomic number – This order usually coincides with increasing atomic mass 8
  9. 9. 11/13/2008 Atoms & Elements Key Concepts: Elements are consist of Metals or Non metals Atoms that have that have Subatomic particles Chemical symbols Neutrons Electrons arranged in the Protons are in Periodic Table determine Make up the by Energy levels Nucleus with Groups Atomic Periods Number Outer shell electrons has a determine Mass Number Periodic law Group number • Periodic Patterns – The chemical behavior of elements is determined by its electron configuration – The first three periods contain just A families. Each period begins with a single electron in a new outer electron shell. – Each period ends with a completely filled outer shell that has the maximum number of electrons for that shell. – The outer shell electrons are responsible for chemical reactions. Elements in the same family have the same number of outer shell electrons; so they will have similar chemical properties. – Group A elements are called representative elements – Group B elements are called transition elements. 9
  10. 10. 11/13/2008 • Chemical Families – IA are called alkali metals because the react with water to from an alkaline solution, very soft metals. (except H2) – Group IIA are called the alkali earth metals because mostly we found them in soils as salts/minerals & they are also reactive, but not as reactive as Group IA. • They are also soft metals, though not as soft as alkali metals – Group VIIA are the halogens • These need only one electron to fill their outer shell • They are very reactive.( disinfectants, bleach) – Group VIIIA are the noble gases as they have completely filled outer shells • They are almost non reactive. Alkali Metals 10
  11. 11. 11/13/2008 Alkaline Earth Metals Transition Metals 11
  12. 12. 11/13/2008 These elements are also called the rare-earth elements. Inner-Transition Metals Halogens 12
  13. 13. 11/13/2008 Noble Gases Reading the Periodic Table: Classification • Nonmetals, Metals, Metalloids, Noble gases 13
  14. 14. 11/13/2008 Metals, Nonmetals, and Metalloids Metals •Metallic character refers to the properties of metals – Shiny or lustrous – Malleable (can be hammered into shape) – Ductile (can be drawn out into wires) – All except mercury are solids at room temperature – They are sonorous (make a ringing sound when hit) – In solution lose electrons in reactions - oxidized – Most oxides are basic and ionic Ex: Metal oxide + water → metal hydroxide Na2O(s) + H2O(l) → 2NaOH(aq) – Tends to form cation in aqueous solution • Metallic character increases down a group. • Metallic character decreases across a period. Only a few metals are magnetic. Magnetism is not a property of most metals! Metals, Nonmetals, and Metalloids Metals • When metals are oxidized they tend to form characteristics cations. • All group 1A metals form M+ ions. • All group 2A metals form M2+ ions. • Most transition metals have variable charges. 14
  15. 15. 11/13/2008 Properties of Non-Metals § They are poor conductors of electrical energy Both a diamond and a pencil ‘lead’ are § They are poor conductors of thermal made of the same element – carbon. energy § Many of them are gases § They are brittle if they are solid § Form anions § Most oxides are acidic Ex: nonmetal oxide + water → acid P4O10(s) + H2O(l) → 4H3PO4(aq) § Gain electrons in reactions – reduced § When nonmetals react with metals, nonmetals tend to gain electrons: metal + nonmetal → salt 2Al(s) + 3Br2(l) → 2AlBr3(s) Metals, Nonmetals, and Metalloids Metalloids Metalloids have properties that are intermediate between metals and nonmetals. Example: Si (shown here) has a metallic luster but it is brittle. Metalloids have found fame in the semiconductor industry. 15
  16. 16. 11/13/2008 Tro, Chemistry: A Molecular 31 Approach The Groups of the Periodic Table • Group 1: The Alkali Metals – Most reactive metals on the PT – Rarely found free in nature – Charge of +1 = 1 valence electron • Group 2: The Alkaline Earth Metals – Still quite reactive – Charge of +2 = 2 valence electrons • Groups 3-12: Transition Metals – Found freely and in compounds in nature – Charge is usually +2 but can vary = usually 2 valence electrons • Group 13: Boron Family – Charge is +3 = 3 valence electrons 16
  17. 17. 11/13/2008 The Groups of the Periodic Table • Group 14: The Carbon Family – Contains elements that can form unusual bonds (carbon and silicon) – Charge is +4 or -4 = contains 4 valence electrons • Group 15: The Nitrogen Family – Charge is -3 = contains 5 valence electrons • Group 16: The Oxygen Family – Also known as the chalcogens – Charge is -2 = 6 valence electrons • Group 17: The Halogens – Most reactive nonmetals – charge is -1 = 7 valence electrons • Group 18: The Noble Gases (The Inert Gases) – Nonreactive – Charge is 0 = 2 or 8 valence electrons Periodic Properties • Periodic law = elements arranged by atomic number gives physical and chemical properties varying periodically. • Various Elemental Properties change fairly smoothly going across a period or down a group. • We will study the following periodic trends: – Atomic radii – Ionization energy – Electron affinity – Melting Points and Boiling Points – Density 17
  18. 18. 11/13/2008 Electron Shells and the Sizes of Atoms Atomic Sizes • As a consequence of the ordering in the periodic table, properties of elements vary periodically. • Atomic size varies consistently through the periodic table. • As we move down a group, the atoms become larger. • As we move across a period, atoms become smaller. There are two factors at work: •principal quantum number, n, and •the effective nuclear charge, Zeff. Electron Shells and the Sizes of Atoms Atomic Sizes • As the principle quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases. • As we move across the periodic table, the number of core electrons remains constant. However, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease. 18
  19. 19. 11/13/2008 Atomic Radius Fig. 8.15 Atomic Radii for Main Group Elements Trends in Atomic Radius Transition Metals • increase in size down the Group • atomic radii of transition metals roughly the same size across the d block – valence shell ns2, not the d electrons – effective nuclear charge on the ns2 electrons approximately the same 19
  20. 20. 11/13/2008 See Figure 8.16 Trends in Ionic Radius • Ions in same group have same charge • Ion size increases down the group – higher valence shell, larger • Cations smaller than neutral atom; Anions bigger than neutral atom • Cations smaller than anions – except Rb+1 & Cs+1 bigger or same size as F-1 and O-2 • Larger positive charge = smaller cation – for isoelectronic species – isoelectronic = same electron configuration • Larger negative charge = larger anion – for isoelectronic series 20
  21. 21. 11/13/2008 41 Tro, Chemistry: A Molecular 42 Approach 21
  22. 22. 11/13/2008 Ionization Energy • minimum energy needed to remove an electron from an atom – gas state – endothermic process – valence electron easiest to remove – M(g) + IE1 → M1+(g) + 1 e- – M+1(g) + IE2 → M2+(g) + 1 e- • first ionization energy = energy to remove electron from neutral atom; 2nd IE = energy to remove from +1 ion; etc. • IE increases (irregularly) as you move from left to right across a period. • IE decreases (irregularly) as you move down a group. 44 22
  23. 23. 11/13/2008 Ionization Energy • The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom: Na(g) → Na+(g) + e-. •The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion: Na+(g) → Na2+(g) + e-. The larger ionization energy, the more difficult it is to remove the electron. There is a sharp increase in ionization energy when a core electron is removed. Chapter 7 45 General Trends in 1st Ionization Energy • larger the effective nuclear charge on the electron, the more energy it takes to remove it • the farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it • 1st IE decreases down the group – valence electron farther from nucleus • 1st IE generally increases across the period – effective nuclear charge increases 46 23
  24. 24. 11/13/2008 Tro, Chemistry: A Molecular 47 Approach Ionization Energy: Periodic table Fig. 8.18 Ionization Energy vs atomic # Chapter 8-48 24
  25. 25. 11/13/2008 49 Example – Choose the Atom in Each Pair with the Higher First Ionization Energy 1) Al or S 2) As or Sb 3) N or Si 4) O or Cl? opposing trends Tro, Chemistry: A Molecular 50 Approach 25
  26. 26. 11/13/2008 Irregularities in the Trend • Ionization Energy generally increases from left to right across a Period • except from 2A to 3A, 5A to 6A ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ Be N 1s 2s 2p 1s 2s 2p ↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑ ↑ B O 1s 2s 2p 1s 2s 2p Which is easier to remove an electron from N or O? Why? from B or Be? Why? Tro, Chemistry: A Molecular 51 Approach Irregularities in the First Ionization Energy Trends ↑↓ ↑↓ ↑↓ ↑ Be Be+ 1s 2s 2p 1s 2s 2p To ionize Be you must break up a full sublevel, cost extra energy ↑↓ ↑↓ ↑ ↑↓ ↑↓ B B+ 1s 2s 2p 1s 2s 2p When you ionize B you get a full sublevel, costs less energy B, Al, Ga, etc.: their ionization energies are slightly less than the ionization energy of the element preceding them in their period. • Before ionization ns2np1. • After ionization is ns2. Higher energy ⇒ smaller radius. Tro, Chemistry: A Molecular 52 Approach 26
  27. 27. 11/13/2008 Irregularities in the First Ionization Energy Trends ↑↓ ↑↓ ↑ ↑ ↑ ↑↓ ↑↓ ↑ ↑ N N+ 1s 2s 2p 1s 2s 2p To ionize N you must break up a half-full sublevel, cost extra energy ↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ ↑↓ ↑ ↑ ↑ O O+ 1s 2s 2p 1s 2s 2p When you ionize O you get a half-full sublevel, costs less energy Group 6A elements. • Before ionization ns2np4. • After ionization ns2np3 where each p electron in different orbital (Hund’s rule). Tro, Chemistry: A Molecular 53 Approach Trends in Successive Ionization Energies • removal of each successive electron costs more energy – shrinkage in size due to having more protons than electrons – outer electrons closer to the nucleus, therefore harder to remove • regular increase in energy for each successive valence electron • large increase in energy when start removing core electrons Tro, Chemistry: A Molecular 54 Approach 27
  28. 28. 11/13/2008 HIGHER IONIZATION ENERGIES Chapter 8-55 56 28
  29. 29. 11/13/2008 Electron Affinities • Electron affinity is the opposite of ionization energy. • Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: Cl(g) + e- → Cl-(g) • Electron affinity can either be exothermic (as the above example) or endothermic: • more energy released (more -); the larger the EA generally increases across period becomes more negative from left to right not absolute lowest EA in period = alkali earth metal or noble gas highest EA in period = halogen 57 The added electron in Cl is placed in the 3p orbital to form the stable 3p6 electron configuration. 29
  30. 30. 11/13/2008 Summary of Trend 1. Ionization Energy: Largest toward 3. Electron Affinity: Most favorable 1. Atomic Radius: Largest toward Magnetic Properties of Transition Metal Atoms & Ions • electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field – this is called paramagnetism – will be attracted to a magnetic field • electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field – this is called diamagnetism – slightly repelled by a magnetic field • both Zn atoms and Zn2+ ions are diamagnetic, showing that the two 4s electrons are lost before the 3d – Zn atoms [Ar]4s23d10 – Zn2+ ions [Ar]4s03d10 60 30
  31. 31. 11/13/2008 Example 8.6 – Write the Electron Configuration and Determine whether the Fe atom and Fe3+ ion are Paramagnetic or Diamagnetic • Fe Z = 26 • previous noble gas = Ar – 18 electrons • Fe atom = [Ar]4s23d6 • unpaired electrons 4s 3d • paramagnetic • Fe3+ ion = [Ar]4s03d5 • unpaired electrons • paramagnetic 61 Melting Points and Boiling Points •Trends in melting Points and boiling points can be used as a measure of the attractive forces between atoms or molecules. •Within the halogens (group 17 or VIIA) melting points and boiling points increase so that at room temperature fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid as you go down this periodic group. •This indicates that the intermolecular forces become stronger going down a group. 31
  32. 32. 11/13/2008 Melting Point Trends in Period •In the second period, melting points increase, going from left to right across the period for the first four elements. • Melting points then decrease drastically for nitrogen, oxygen, and fluorine, which are all diatomic molecules. • The lowest melting point is for neon, which is monatomic. Melting Points of Elements 32
  33. 33. 11/13/2008 Trends in Density •Densities of elements increase in a group as atomic number increases. •In periods, going from left to right, densities increase, then decrease. •Elements with the greatest densities are at the center of period 6. Densities of Elements 33
  34. 34. 11/13/2008 34

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