Lecture notes: CHEM103 Fall 2008 – October 23

LECTURE OUTLINE:

0) Review: Periodic trends based on electronic structure

Atomic radius

Ionization energy

Electron affinity

1) Ionic radius

2) FINALLY – BRINGING IT ALL TOGETHER…

Periodic properties: metals vs. non-metals

Formation of ionic compounds & their properties

Ion pair energy (lattice energy, kinda) as driving force of ionic bond formation

Relative strengths of ion pair energies

3) Ionic compounds in solution

Some vocabulary

Energy of the solvation process – driving force of dissolution

(competition with ion pair energy)

SOLUBILITY RULES AS AN ALTERNATIVE

Writing dissolution & precipitation reactions

Strong and weak electrolytes and non-electrolytes

RESULT of IE and EA:

in both cases, NOBLE-GAS CONFIGURATION (FILLED ORBITAL STABLE)

(additional note: Group #s indicating likely charge on ions)

IONIC RADIUS (SIZE)

note: ionic radius given for “most stable” ion (noble gas config)

examples…

i) Comparing ions to ions

ii) Comparing ions to neutrals

iii) Isoelectronic examples

METALS/NON-METALS AND PROPERTIES; SEMI-METALS (METALLOIDS)

Properties of metals (i.e. sodium)

Tendency to lose electrons in reactions (to form cations)

• Low ionization energy

• Solids w/ moderately high melting points (exceptions)

• Conduct electricity well

• Conduct heat well

• Malleable

Properties of non-metals (i.e. chlorine)

Tendency to gain electrons in reactions (to form anions)

• Have much electron affinity

• A variety of physical forms, including gases at room temp

• Generally electrically insulating

• And thermally insulting (poor conductor)

NOTE: why does Zr make a 4+ ion and C NOT make a 4+ (or even a 2+)? More to do with metal/non-metal than with # of electrons.

EXCEPTIONS TO GENERAL PERIODIC TABLE STRUCTURE:

He & especially H!

(Ionic) bond formation (CHEMISTRY!)

by reaction of metal with non-metal to form an ionic compound!

Properties of ionic compounds (i.e. sodium chloride)

• Consist of cations & anions in a regular crystal lattice

• Solids w/ very high melting points

• Solid form is hard; breaks along lines within the crystal

• Solid form is often transparent (“clear as crystal”)

• (Sometimes) dissolve to form electrically conductive solutions

NOTE: “Na” is NOT “Na⁺”; “Cl” or (Cl₂) is NOT “Cl^–” ;

if you leave off the charge (or just get it wrong…)

THE ANSWER IS WRONG – ALWAYS!

(involves transfer of valence electrons between reactants)

P.S. How do ions combine? (IN WHOLE NUMBER RATIOS – remember Dalton!)

http://www.youtube.com/watch?v=Ftw7a5ccubs&feature=related

valence shell revisited; formation of ionic bonds as electron “transfer”

HOW? – valence electrons transferred (what direction?)

Ca + F₂ à CaF₂

(Fluorite)