Title: The Modern Periodic Table Chapter 6
1The Modern Periodic TableChapter 6
2Objectives 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.
3Objectives
- 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.
4Arrangement 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
5The Periodic Table With Atomic Symbols, Atomic
Numbers, and Partial Electron Configurations
6Broad 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)
7Information Contained in the Periodic Table
- Each group member has the same valence electron
configuration (these electrons primarily
determine an atoms chemistry). - The electron configuration of any representative
element.
8Information 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.
9Special Names for Groups in the Periodic Table
10Metals
- 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.
11Metals (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.
12Transition 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
13Metalloids
- 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.
14Nonmetals
- 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
15Nonmetals (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.
16Valence 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
17The Octet Rule
- Atoms tend to lose, gain, or share electrons
until they are surrounded by 8 valence electrons
18Exceptions 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.
19Exceptions 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
20Exceptions 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
21Exceptions Chromium (Cr), Molybdenum (Mo), but
NOT Tungsten (W)
- Cr is Ar4s13d5, NOT Ar4s23d4
- Mo is Kr5s14d5, NOT Kr5s24d4
- W IS Xe6s15d5
22Ionization Energy
- The quantity of energy required to remove an
electron from the gaseous atom or ion.
23For 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
24Periodic Trends
- First ionization energy
- increases from left to right across a period
Why? - decreases going down a group. Why?
25Trends in Ionization Energies for the
Representative Elements
26The Values of First Ionization Energy for the
Elements in the First Six Periods
27Question
- The first ionization energy for the group IIA
elements are significantly higher than those of
the Group IA elements in the same periods. Why?
28Question
- The first ionization energy of the Group IIIA
elements are lower than the IIA elements in the
same period. Why?
29Question
- Group VIA elements have slightly lower first
ionization energies than Group VA elements in the
same period. Why?
30Electron 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
31The Electronic Affinity Values for Atoms Among
the First 20 Elements that Form Stable, Isolated
X- Ions
32Questions
- 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?
33Periodic Trends
- Atomic Radii
- decrease going from left to right across a
period Why? - increase going down a group. Why?
34The Radius of an Atom
35Atomic Radii for Selected Atoms
36Ionic Radii
- What is the trend for ionic radii?
- Which of the Period 3 ions would be the smallest?
- Na, Mg2, Al3, S2-, Cl-
37Sizes of Ions Related to Positions of the
Elements in the Periodic Table
38Electronegativity Increases Up and To the Right
39What 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
40What 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.
41Oxidation 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.