Chapter 8: Periodic Properties of the Elements

I) Preliminary Questions

A) What is periodicity? Why is it so important to chemistry? B) How can periodicity be used to explain observed trends in properties of groups of elements?

II) History of Periodic Table

A) J. Dobreiner

- Law of Triads - groups of three elements share similar properties.

B) J. Newlands

- Law of Octaves - elements arranged in octaves (did not account for Group VIIIA). - works well for s & p-block elements only.

C) Dmitri Mendeleev & J. Lothar Meyer

- Meyer based his periodic table on physical properties of elements.

- Mendeleev based periodic table on chemical properties of elements and went so far as to make predictions concerning the properties of undiscovered elements. (see Figure 8.1)

- Both based their periodic tables on increasing atomic mass.

- Modern table based on increasing atomic number.

3) Concept of Periodicity

A) Properties of certain elements are similar & occur at fixed intervals.

B) Periodicity used to predict trends due to electron configurations. (Example: Alkali Metals) 65

III) Electronic Configurations

1) Aufbau Principle (Diagonal Rule)

A) accounts for how electrons are introduced into atomic orbitals.

1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p

2) Hund’s Rule

A) electrons fill degenerate orbitals halfway first (results in lowest energy configuration).

B) electrons listed as having parallel spins.

3) Orbital Filling Diagrams

A) An alternate method of expressing electron configurations, where arrows represent the electrons.

B) Within a given orbital, electrons are paired with opposing spins (one arrow points up, the other points down).

4) Valence Electrons

A) electrons found in outermost principal quantum level of atom. Other electrons are called core electrons.

B) electrons in same family (group) of periodic table have same outer electronic configuration and share similar chemical behavior.

Examples: Group IA ns1

Group IIA ns2

Group IIIA ns2np1 66

C) 4 Blocks of Elements in Periodic Table

1. s-block main group (representative) elements

2. p-block main group (representative) elements

3. d-block transition metals

4. f-block inner transition metals

Table 1. Distribution of Electrons Inside Electron Cloud

Principal Energy Level (n) 1 2 3 4 Maximum No. of Electrons (2n2) 2 8 18 32 Atomic Orbitals in Sublevel 1s 2s, 2p 3s, 3p, 3d 4s,4p,4d,4f Designation of Filled Orbitals 1s2 2s2, 2p6 3s2, 3p6, 3d10 4s2, 4p6, 4d10, 4f14 Maximum Electrons per Sublevel 2 2, 6 2, 6, 10 2, 6, 10, 14 Orbitals per Sublevel 1 1, 3 1, 3, 5 1, 3, 5, 7

5) Pauli Exclusion Principle

A) no two electrons in an atom can have the same set of 4 quantum numbers.

6) Polyelectronic Atoms

A) atoms with more than 1 electron. B) solutions to Schroedinger equation are much more complicated.

IV) Electronic Configurations & Periodic Properties: Atomic Radii

1) Atomic Radius (Size)

- atomic radius decreases across period. - atomic radius increases down a column.

2) Ionization Energy

General Equation: X (g)  X + (g) + e -

A) Ionization energy represents the energy required to remove an electron from a gaseous atom (ion).

B) Ionization energy increases across period.

C) Ionization energy decreases down a column. 67

3) Electron Affinity

General Equation X (g) + e -  X - (g)

A) Electron affinity refers to the energy change associated with addition of an electron to a gaseous atom.

B) Electron affinity increases across period.

C) Electron affinity decreases down a column.

V) Introduction to Chemical Bonding

1) The chemical forces that hold atoms together are referred to as bonds. 2) Two main categories of bonding

A) Ionic Bonding B) Covalent Bonding

VI) Ions & Electronic Configurations

1) Valence shell electrons are what participate in chemical bonding. 2) Metals lose electrons to form cations, while nonmetals gain electrons to form anions. 3) Atoms want to attain a noble gas electron configuration. Since the noble gases (except He) have a total of 8 valence electrons, this condition is often called the octet rule. 4) For main group elements, the group # gives the number of valence electrons directly. (Group IA has 1 valence electron, Group IIA has 2 valence electrons, etc.). 5) Predicting Formulas for Ionic Compounds

Ionic Compounds  dealing with solids

Example: Calcium oxide CaO

Ca [Ar] 4s2 Ca2+ [Ar] O [He] 2s2p4 O2- [Ne]

Mention Exceptions that are found along diagonal (metalloids)

Isoelectronic ions: ions with same number of electrons.

Size decreases with increasing Z for a set of isoelectronic ions. 68

VII) Ionic Bonds & Formation of Ionic Solids

1) Ionic bonding involves a transfer of electrons from a metal atom to a nonmetal atom. This transfer results in the formation of metal cations and nonmetal anions.

2) All salts are ionic compounds. These compounds consist of a three dimensional solid consisting of alternating cations and anions.

3) We will not concern ourselves with Born-Haber Cycle or lattice energy calculations.