I Chemical Bonding

Chemical Bond  attractive force between atoms or ions that binds them together as a unit  bonds form in order to…  decrease potential energy (PE)  increase stability

COMPOUND Ternary Compound Binary Compound 2 elements more than 2 elements NaNO 3 NaCl

ION Polyatomic Ion Monatomic Ion 1 atom 2 or more atoms NO 3 - Na +

IONIC COVALENT Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties e - are transferred from metal to nonmetal high yes (solution or liquid) yes e - are shared between two nonmetals low no usually not Melting Point crystal lattice true molecules TYPES OF BONDS Physical State solid liquid or gas odorous

“electron sea” METALLIC Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties Melting Point TYPES OF BONDS Physical State e - are delocalized among metal atoms very high yes (any form) no malleable, ductile, lustrous solid

IONIC BONDS

IONIC BONDING - CRYSTAL LATTICE

Covalent Bonding - True Molecules Diatomic Molecule

METALLIC BONDING - “ELECTRON SEA”

BOND POLARITY  Most bonds are a blend of ionic and covalent characteristics.  Difference in electronegativity determines bond type.

BOND POLARITY Electronegativity  Attraction an atom has for a shared pair of electrons.  higher e - neg atom   -  lower e - neg atom   +

BOND POLARITY Electronegativity Trend (p. 151) Increases up and to the right.

BOND POLARITY Nonpolar Covalent Bond e - are shared equally symmetrical e - density usually identical atoms

++ --  Polar Covalent Bond  e - are shared unequally  asymmetrical e - density  results in partial charges (dipole)

Nonpolar Polar Ionic

BOND POLARITY Examples: Cl 2 HCl NaCl =0.0 Nonpolar =0.9 Polar =2.1 Ionic

Chemical Bond  attractive force between atoms or ions that binds them together as a unit  bonds form in order to…  decrease potential energy (PE)  increase stability

I LEWIS DIAGRAMS Molecular Structure

RULE  Remember…  Most atoms form bonds in order to have 8 valence electrons.

 Hydrogen  2 valence e -  Groups 1,2,3 get 2,4,6 valence e -  Expanded octet  more than 8 valence e - (e.g. S, P, Xe)  Radicals  odd # of valence e - A. OCTET RULE  Exceptions: F B F F H O H N O Very unstable!! F F S F F

B. DRAWING LEWIS DIAGRAMS  Find total # of valence e -.  Arrange atoms - singular atom is usually in the middle.  Form bonds between atoms (2 e - ).  Distribute remaining e - to give each atom an octet (recall exceptions).  If there aren’t enough e - to go around, form double or triple bonds.

B. DRAWING LEWIS DIAGRAMS  CF 4 1 C × 4e - = 4e - 4 F × 7e - = 28e - 32e - F F C F F - 8e - 24e -

B. DRAWING LEWIS DIAGRAMS  BeCl 2 1 Be × 2e - = 2e - 2 Cl × 7e - = 14e - 16e - Cl Be Cl - 4e - 12e -

B. DRAWING LEWIS DIAGRAMS  CO 2 1 C × 4e - = 4e - 2 O × 6e - = 12e - 16e - O C O - 4e - 12e -

C. POLYATOMIC IONS  To find total # of valence e - :  Add 1e - for each negative charge.  Subtract 1e - for each positive charge.  Place brackets around the ion and label the charge.

C. POLYATOMIC IONS  ClO Cl × 7e - = 7e - 4 O × 6e - = 24e - 31e - O O Cl O O + 1e - 32e - - 8e - 24e -

C. POLYATOMIC IONS  NH N × 5e - = 5e - 4 H × 1e - = 4e - 9e - H H N H H - 1e - 8e - - 8e - 0e -

C. POLYATOMIC IONS  OH - 1 O × 6e - = 6e - 1 H × 1e - = 1e - 7e - O H + 1e - 8e - - 8e - 0e -

D. RESONANCE STRUCTURES  Molecules that can’t be correctly represented by a single Lewis diagram.  Actual structure is an average of all the possibilities.  Show possible structures separated by a double-headed arrow.

D. RESONANCE STRUCTURES O O S O O O S O O O S O n SO 3

I MOLECULAR GEOMETRY

VSEPR THEORY  Valence Shell Electron Pair Repulsion Theory  Electron pairs orient themselves in order to minimize repulsive forces.

VSEPR THEORY  Types of e - Pairs  Bonding pairs - form bonds  Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!

VSEPR THEORY  Lone pairs reduce the bond angle between atoms. Bond Angle

DETERMINING MOLECULAR SHAPE  Draw the Lewis Diagram.  Tally up e - pairs on central atom.  double/triple bonds = ONE pair  Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles!

COMMON MOLECULAR SHAPES 2 total 2 bond 0 lone LINEAR 180° BeH 2

COMMON MOLECULAR SHAPES 3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3

COMMON MOLECULAR SHAPES 3 total 2 bond 1 lone BENT <120° SO 2

COMMON MOLECULAR SHAPES 4 total 4 bond 0 lone TETRAHEDRAL 109.5° CH 4

COMMON MOLECULAR SHAPES 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NH 3

COMMON MOLECULAR SHAPES 4 total 2 bond 2 lone BENT 104.5° H2OH2O

COMMON MOLECULAR SHAPES 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90° PCl 5

COMMON MOLECULAR SHAPES 6 total 6 bond 0 lone OCTAHEDRAL 90° SF 6

EXAMPLES  PF 3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° F P F F

EXAMPLES  CO 2 O C O 2 total 2 bond 0 lone LINEAR 180°