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.