Notes

Outline

Chemical Bonding Basic Concepts

Chemical Bond An electrical interaction between electrons of one atom and the positive nucleus of another atom that results in the binding together of atoms in a stable unit Types of Interaction Electron-Nuclear attractions (stabilizing) Electron-Electron Repulsion (destabilizing) Nuclear-Nuclear Repulsion (destabilizing)

Lewis Electron Dot Symbols For Main Groups(1-2 and 13-18) Symbol represents the core (nucleus and all but the valence electrons) Dots (or x’s) represent the valence electrons Unpaired electrons are potential source of chemical bonding Paired electrons are called “lone pairs” and usually not involved in bonding (there are exceptions) Display Lewis dot symbols for period 2 elements from groups 1-2 and 13-18

Ionic Radius Distance from the nucleus of an ion to its outermost occupied region

T-9.5 Determining Iodide Radii

Ionic Radii vs Atomic Radii Cation radii are smaller than the neutral atoms of the same element Na + < Na Higher charged cations are smaller than lower charged cations in the same period Be +2 < Li + Anion radii are larger than neutral atoms of the same element F- > F Higher charged anions are larger than smaller charged anions in the same period O-2 > F- Examples

T-71 Comparison of Atomic and ionic radii of Elements

Ionic Bond A bond created by the transfer of one or more  valence electrons from a metal atom to a non-metal atom and the electrostatic force of attraction of oppositely charged ions K•    --à  K + (isoelectronic to Ar)   +   e-    F  +  e- -à F – (isoelectronic to Ne)  K +  +  F -  -à K + F –

T-114 Interionic Forces of Attraction:Bond Length

Co-Valent Bond A bond created by the mutual sharing of one or more pairs of valence electrons in order for each atom to be iso-electronic to the nearest Noble Gas Single co-valent bond-sharing of one pair of valence electrons Each atom supplies half of the shared electrons each

T-9.9 Potential Energy Curve For H-H bond

Simple Molecular Compounds Using Lewis Dot Structures                                      H-Cl                    H-O                       |                      H              H              |            :N-H              |             H

Co-Ordinate Co-valentBond Co-valent Bond where one atom donates both electrons to be shared Called a donor-acceptor bond Examples H3N:      +   H+  -à   H3N:-àH+ -à NH4 + Hydronium Ion

Multiple Bonding Double co-valent bond-sharing of two pairs of valence electrons Examples O=O O=C=O Triple co-valent bond- sharing of three pairs of valence electrons Examples :NºN: :CºO: -CºC-

Electronegativity The ability of an atom to attract valence electrons to itself Pauling Electronegativity Scale (0-4)

T-9.12 Electronegativities of the Elements(Periodic Table)

Electronegativity and Periodicity Electronegativity increases from right to left in a period and from bottom to top in a group

Bond Polarity Polar Co-valent Bond- Covalent Bond where the atoms share the electrons unequally due to differences in electronegativity Bonding electrons are unevenly distributed lying closer to the more electronegative atom Results in an electrical force known as a “bond dipole moment” Greater the bond dipole moment the greater the polarity of the bond

Relationship Between Electronegativity and Bond Polarity The greater the electronegativity difference the greater the bond dipole moment and the more polar the bond C-H  2.5 – 2.1 = 0.4 difference = small dipole moment C-F 4.0 – 2.5 = 1.5 difference = large dipole moment Symbol for dipole moment =  +® Bond dipole are vectors Vector-a measurement having a magnitude and direction

Simple Vector Math Vectors directed in the same direction add and result in a larger vector in that same direction Opposing vectors result in a smaller resulting vector having a direction the same as the larger of the opposing vectors Opposing vectors of the same magnitude will cancel each other and result in no vector Vectors at an angle to one another result in a vector positioned between the two vectors at an angle Precisely determined using triginometry

Steps To Writing Lewis Structures Determine the total number of valence electrons for all atoms represented in the molecular formula  Decide which atom(s) will occupy the center(mono-valent ions can not be considered) Arrange all other atoms around the central atom(s) as symmetrical as possible

Steps to Writing Lewis Structures (cont)      4. Using the valence electrons determined in step 1 connect all atoms using dashes as bonding pairs into one cohesive unit     5. Distribute the remainder of the  valence electrons as lone pairs so as to bring each atom into compliance with the Octet Rule     6. If not enough valence electrons in the pool established in step 1 then consider converting a lone pair into a bonding pair Examples-

Determination of Formal Charge Of An Atom Write the Lewis Structure of the molecule Divide all bonding pairs distributing half to each atom bonded Count the lone pairs and total them with the bonding electrons of the atom Consider the number of valence electrons the atom should have to be neutral (Lewis Dot Symbol) Subtract electrons determine in step 3 from the valence electrons determined in step 4 to get formal charge Examples

Resonance and Lewis Structures Resonance-Equivalent hypothetical structures that differ by the position of one or more Pi bonding pairs within the structure The more Lewis Structures that can be written the more resonance stability Examples Carbonate ion Carboxylate ion Benzene

Exceptions To The Octet Rule Free Radicals Incomplete Octets BF3 AlCl3 Expanded Octets XeF6 SF6

Bond Formation Bond Formation Bond Energy- Energy released when a bond is formed Bond formation is an exothermic process Bond Breaking Dissociation Energy-minimum energy required to break a bond Bond Breaking is an endothermic process

T-85 Bond Breaking and Bond Formation Energy Exchanges

T-9.5 Bond Energies kJ/mol