14 Covalent Bonding Flashcards
What is a sigma bond?
-Direct head-on/end to-end overlap of atomic orbitals
-results in electron density concentrated between nuclei of bonding atoms
What is a pi bond?
- sideways overlap of atomic orbital
- results in electron density above and below the plane of the nuclei of the bonding atoms
Formal Charge
- decide which Lewis structure is preferred from several
- Lewis structure with the atoms having FC values closest to 0 is preferred
- Treats covalent bonds as if they were purely covalent with equal e- distribution
- ignores electronegativity
- counts how may e- belong to each atom in Lewis structure and compares this with the # of valence e- in non-bonded atom
Octet Rule exceptions
- incomplete octets
- expanded octets
Delocalization
- electrons that are shared between all atoms in a molecule/ion opposed to being localized between a pair of atoms
Resonance
- 2 or more Lewis structures to represent the same molecule
- structures that cannot be described by a single structure
Why do expanded octets exist?
- d orbital available in valence shell of these atoms have energy values relatively close to those of p orbitals
- Example: from 3p to 3d orbitals –> allows additional e- pairs to form
Species with 5 e- domains
- triangular bipyramidal shape (all 5 domains are bonding pairs)
- angles of 90, 180 and 120
- shape will change if there are non-bonding e- as they cause greater repulsion
- 1 non-bonding pair – > seesaw
- 2 non-bonding pairs –> T-shaped
- 3 non-bonding pairs –> linear
Species with 6 e- domains
- octahedral
- angles of 90
- 1 non-bonding pair –> square pyramidal
- 2 non-bonding pairs –> square planar
How to interpret Formal Charge
- low formal charges means less charge transfer has taken place in forming a structure from it’s atoms (stable structure)
- sum of formal charges in molecule for neutral molecules to equal charge of ion
- more stable Lewis structure has have negative values for more electro negative atoms+
Oxygen dissociation wavelength
<242nm
Ozone dissociation wavelength
<330nm
Does it take less energy to break ozone or oxygen and why?
- it takes less energy to break ozone because it has a resonance structure which counts its bonds as 1.5 opposed to the double (2) bond of oxygen
Why is ozone important?
- it absorbs radiation in the range of 200-315 which is UV-B and UV-C (veri scawy to living tissue)
- henceforth, ozone protec human being
Catalytic destruction of Ozone
- when ozone absorbs UV radiation, it become unstable and allows it to react with other compounds such as Nitrogen oxides (NOx) and chlorofluorocarbons (try saying that 5 times fast) (CFCs)
- both compounds produce highly reactive free radials that catalyze the decomposition of ozone to oxygen
Catalytic destruction of Ozone (Nitrogen Oxides)
- Nitrogen dioxide forms from the oxidation of Nitrogen Oxide (also a free radical)
Catalytic destruction of Ozone (chloroflurocarbons)
- higher energy UV radiation breaks CFCs down releasing free chlorine atoms (reactive free radicals)
- weaker C-Cl bond breaks in preference to the C-F bond, thus, Cl radicals catalyze the decomposition of ozone
Sigma Bond
- overlap along bond axis
- always forms in a single covalent bond
- electron density is concentrated between the nuclei of bonded atoms
Pi bond
- overlap sideways
- weaker than sigma bonds
- only form during the overlap of p orbitals –> pi bonds can only form in double or triple bonds
- electron is concentrated above and below bond axis
What forms a hybrid orbital?
- mixing of different types of atomic orbitals on the same atom
Formation of Covalent Bonds
- often starts with excitation of atoms
- excitation is where and e- is promoted within the atom from the 2s orbital to the vacant 2p orbital
Hybridization
- mixing atomic orbitals to form new hybrid atomic orbitals
- there is an unequal atomic orbitals as a result of excitation
- these orbits are mixed together resulting gin hybrid atomic orbitals
- this mix of atomic orbits will balance themselves out due to the imbalance (refer to paint analogy)
sp3 hybridization
Example: Carbon (CH4)
- when carbon forms 4 single bonds, it undergoes sp3 hybridization, producing 4 equal orbitals
- forms a tetrahedron (109.5 degrees)
- each hybrid orbital overlaps with the atomic orbital of another atom forming 4 sigma bonds
sp2 hybridization
Example: Carbon (C2H4)
- produces 3 equal orbits
- triangular planar (120 degrees)
- forms 3 sigma bonds
-p orbitals form pi bonds
sp hybridization
Example: Carbon (C2H2)
- 2 equal orbits
- linear (180 degrees)
- 2 sigma bonds
2 pi bonds form representing 4 lobes of e- density around the atoms
Lone pairs in Hybridization
Example: Amonia
- non-bonding pair allows for sp3 hybrid orbital even though it has only 3 bonding pairs
Hybridization to predict molecular shape
Tetrahedral = sp3 hybridized
Triangular planar = sp2 hybridized
Linear = sp hybridized
