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The Modern Periodic Table Chapter 6

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Title: The Modern Periodic Table Chapter 6


1
The Modern Periodic TableChapter 6
2
Objectives for Friday
  • 4. Define period, group, and family.
  • 5. Explain the different systems for numbering
    the groups on the Periodic table.
  • 6. Name and locate on the periodic table metals,
    nonmetals, metalloids, noble gases, alkali
    metals, alkaline earth metals, halogens,
    chalcogens, Lanthanides, Actinides, Transition
    metals, and Inner transition metals.

3
Objectives
  • 7. Apply the octet rule in predicting the
    stability of an element.
  • 8. Predict the electron configuration of the
    outer energy level of any element from its
    position on the Periodic table.
  • 9. Recognize exceptions to the Aufbau order due
    to half-full or full d-sublevels.
  • 10. Define atomic radius, ionic radius,
    ionization energy, electron affinity, and
    electronegativity.
  • 11. Predict trends in atomic radius, ionic
    radius, oxidation number, ionization energy,
    electron affinity, and electronegativity.
  • 12. Explain trends in terms of nuclear charge and
    shielding effects.

4
Arrangement and Nomenclature
  • Rows are called periods
  • Columns are designated as groups
  • Each column in the main table and each row at the
    bottom is also designated an individual family
  • Groups 1A, 2A, and 3-8A are the main groups, or
    representative elements
  • Groups 1B-8B are called the transition elements

5
The Periodic Table With Atomic Symbols, Atomic
Numbers, and Partial Electron Configurations
6
Broad Periodic Table Classifications
  • Representative Elements (main group) filling s
    and p orbitals (Na, Al, Ne, O)
  • Transition Elements filling d orbitals (Fe,
    Co, Ni)
  • Lanthanide and Actinide Series (inner transition
    elements) filling 4f and 5f orbitals (Eu, Am,
    Es)

7
Information Contained in the Periodic Table
  1. Each group member has the same valence electron
    configuration (these electrons primarily
    determine an atoms chemistry).
  2. The electron configuration of any representative
    element.

8
Information Contained in the Periodic Table
  • Certain groups have special names (alkali metals,
    alkaline earth metals, chalcogens, halogens,
    etc).
  • Metals and nonmetals are characterized by their
    chemical and physical properties.

9
Special Names for Groups in the Periodic Table
10
Metals
  • Metals makeup more than 75 of the elements in
    the periodic table. Metals are characterized by
    the following physical properties
  • They have metallic shine or luster.
  • They are usually solids at room temperature.
  • They are malleable. Malleable means that metals
    can be hammered, pounded, or pressed into
    different shapes without breaking.
  • They are ductile meaning that they can be drawn
    into thin sheets or wires without breaking.
  • They are good conductors of heat and electricity.

11
Metals (cont)
  • All B and most A elements are metals.
  • The B ? At stairstep designates the border
    between metals and non-metals
  • 1A elements are alkali metals
  • They are soft shiny metals that usually combine
    with group VIIA nonmetals in chemical compounds
    in a 11 ratio.
  • 2A elements are the alkaline earth metals
  • Both alkali and alkaline earth metals are
    chemically reactive, but 2A metals are less
    reactive than 1As.
  • They combine with the group VIIA nonmetals in a
    12 ratio.

12
Transition Metals Metalloids
  • Transition metals
  • The remaining 1-8B elements are all transition
    elements
  • The transition elements also have valence
    electrons in two shells instead of one.
  • Inner transition metals
  • The lanthanide and actinide series comprise the
    inner transition metals

13
Metalloids
  • Metalloids have characteristics of both metals
    and nonmetals and so cant be classified as
    either, but something in between.
  • They are good conductors of heat and electricity
  • They are not good conductors or insulators.
  • The six metalloids are B, Si, Ge, As, Sb, and Te.

14
Nonmetals
  • There are 17 nonmetals in the periodic table, and
    they are characterized by four major physical
    properties.
  • They rarely have metallic luster.
  • They are usually gases at room temperature.
  • Nonmetallic solids are neither malleable nor
    ductile.
  • They are poor conductors of heat and electricity.
  • The elements above the B ? At stairstep are
    nonmetals

15
Nonmetals (cont)
  • Group 6A contains the chalcogen elements
  • Group 7A contains the highly reactive halogen
    elements
  • They are fluorine, chlorine, bromine, and iodine.
  • The halogens exist as diatomic molecules in
    nature.
  • Group 8A comprises the completely non-reactive
    noble gases
  • The noble gases are also called rare gas
    elements, and they all occur in nature as gases.
  • The noble gases fulfill the octet rule by having
    a full outer level with 8 valence electrons.
  • Therefore, they do not undergo chemical reactions
    because they do not accept any electrons.

16
Valence Electrons and the Periodic Table
  • Valence Electrons and Group
  • Atoms in the same group have the same chemical
    properties because they have the same number of
    valence electrons.
  • Moreover, they have the same outermost orbital
    structure
  • E.g. 1A elements all have s1 valence electrons
  • E.g. 2A elements all have s2 valence electrons
  • Valence Electrons and Period
  • The primary quantum number (n) for an elements
    valence electrons is the same its period.
  • E.g. Lithiums valence electron is n2 and Li is
    found in the 2nd period

17
The Octet Rule
  • Atoms tend to lose, gain, or share electrons
    until they are surrounded by 8 valence electrons

18
Exceptions to Aufbau Order
  • Subshell degeneracies occur in elements larger
    than Vanadium
  • i.e. different 4s and 3d orbitals have nearly the
    same energy
  • Also, it turns out that full and half-full
    sublevels have the most stability.

19
Exceptions Copper (Cu), Silver (Ag), Gold (Au)
  • Strict Aufbau ordering of Cu would be Ar4s23d9
  • experimental observation shows this to be an
    excited state
  • the ground state has a configuration of
    Ar4s13d10
  • The observed configuration for Cu creates a
    ½-full s and a full d, which is more stable than
    a full s and a partial d
  • Ag is NOT Kr5s24d9, but Kr5s14d10
  • Au is NOT Xe6s25d9, but Kr6s15d10

20
Exceptions Lanthanum and Actinium
  • Aufbau would place them in the inner transition
    series, but instead they are in the scandium
    family
  • i.e. La is Xe6s25d1
  • i.e. Ac is Rn7s26d1

21
Exceptions Chromium (Cr), Molybdenum (Mo), but
NOT Tungsten (W)
  • Cr is Ar4s13d5, NOT Ar4s23d4
  • Mo is Kr5s14d5, NOT Kr5s24d4
  • W IS Xe6s15d5

22
Ionization Energy
  • The quantity of energy required to remove an
    electron from the gaseous atom or ion.

23
For Aluminum
  • Al (g) Al (g) e- I1 580 kJ/mol
  • Al (g) Al2 (g) e- I2 1850 kJ/mol
  • Al2 (g) Al3 (g) e- I3 2740 kJ/mol
  • Al3 (g) Al4 (g) e- I4 11,600 kJ/mol

24
Periodic Trends
  • First ionization energy
  • increases from left to right across a period
    Why?
  • decreases going down a group. Why?

25
Trends in Ionization Energies for the
Representative Elements
26
The Values of First Ionization Energy for the
Elements in the First Six Periods
27
Question
  • The first ionization energy for the group IIA
    elements are significantly higher than those of
    the Group IA elements in the same periods. Why?

28
Question
  • The first ionization energy of the Group IIIA
    elements are lower than the IIA elements in the
    same period. Why?

29
Question
  • Group VIA elements have slightly lower first
    ionization energies than Group VA elements in the
    same period. Why?

30
Electron Affinity
  • The energy change associated with the addition
    of an electron to a gaseous atom.
  • X(g) e? ? X?(g)
  • Note the more negative the electron affinity,
    the more energy is released

31
The Electronic Affinity Values for Atoms Among
the First 20 Elements that Form Stable, Isolated
X- Ions
32
Questions
  • Helium and Beryllium do not form stable isolated
    negative ions. Why?
  • Nitrogen does not form a stable, isolated N- (g)
    ion, whereas carbon forms C-(g). Why?
  • In contrast to nitrogen, oxygen can add an
    electron to form the stable O- ion. Why?

33
Periodic Trends
  • Atomic Radii
  • decrease going from left to right across a
    period Why?
  • increase going down a group. Why?

34
The Radius of an Atom
35
Atomic Radii for Selected Atoms
36
Ionic Radii
  • What is the trend for ionic radii?
  • Which of the Period 3 ions would be the smallest?
  • Na, Mg2, Al3, S2-, Cl-

37
Sizes of Ions Related to Positions of the
Elements in the Periodic Table
38
Electronegativity Increases Up and To the Right
39
What is electronegativity?
  • How tightly an atom holds on to its valence
    electrons.
  • Essentially, this value depends on
  • the number of positively charged protons in the
    atoms nucleus
  • the radius of the outermost electron shell

40
What is electronegativity?
  • The more positive the nucleus
  • The smaller the valence electron shell around it
  • The greater the attraction between nucleus and
    electrons
  • Thus, the more electronegative the atom!
  • Thus, a high electronegativity value implies that
    the valence electrons are tightly held and
    require a large amount of energy to remove.

41
Oxidation Numbers
  • The Octet Rule states that atoms want to have
    their valence shell filled with electrons.
  • This means that, ideally, atoms are most stable
    with 8 valence electrons
  • N.B. This is not true for Period 1. Why?
  • Atoms will gain or lose electrons to form ions in
    order to fulfill the Octet Rule.
  • The charge they take on in this process is called
    the valence.
  • The oxidation state is, for ions, equal to the
    valence.
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