ch9 molecular geometry and bonding theory Flashcards
molecular architecture
Lewis structures do not show the overall shapes of molecules (molecular architecture), they only show number and types of bonds between atoms
bond angles
The shape of a molecule is determined by its bond angles: angles made by the lines joining the nuclei of atoms in the molecules. Bond angles and bond lengths together accurately define the shape/size of a molecule
5 geometric structures
for ABn (A is bonded to n B atoms) linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral When A is a representative element (from s or p block) we can predict shape of molecule using valence-shell electron-pair repulsion (VSEPR model)
bonding pair
a region in which the electrons will most likely be found
nonbonding/lone pair
electron domain that is located principally on one atom
electron domain
each nonbonding pair, single bond, or multiple bond produces an electron domain around the central atom. The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them (basis of Valence Shell Eelectron Pair Repulsion model)
electron domain geometry2
- arrangement of electron domains about the central atom of ABn
- The different shapes of ABn molecules/ions depend on the number of electron domains surrounding central A atom.
molecular geometry
arrangement of only the atoms in the molecule/ion, does not describe nonbonding pairs of electrons
ideal bond angles
angles for specific electron geometry when surrounding electrons and domains are identical
factors for greater bond angle
Electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and thus tend to compress bond angles. Electron domains for multiple bonds exert a greater repulsive force on adjacent electron domains than do electron domains for single bonds (more nonbonding electrons or double/triple bonds=greater angle for that electron domain)
molecules with 5 domains
- electron domain geometry: trigonal bipyramidal (2 trigonal pyramids sharing a base)
- (0 lone pairs) 2 of 5 domains: axial positions, makes 90º with equatorial bonds
- (0lone pairs) Other 3 domains: equatorial positions, makes 90º with axial bonds, 120º with other equatorial bonds, nonbonding pairs always occupy this position
- 1 lone pair: see saw, <120, <90
- 2lone pair: tshaped, <90
- 3 lone pair: linear, 180
molecules with 3 domains
- 0 lone pair: trigonal planar, 120
- 1 lone pair: bent, <120
4 electron domains
- 0 lone pair: tetrahedral, 109.5
- 1 lone pair: trigonal pyramidal, <109.5
- 2 lone pair: bent, «109.5
2 electron domains
linear, 180
6 electron domains
- electron domain geometry: octahedron (2 square pyramids sharing a base)
- 0 lone pairs: octahedron, 90
- 1 lone pair: square pyramidal, <90
- 2 lone pair: square planar, 90
bond dipole
the dipole moment that is due only to the two atoms in that bond. Bond dipoles/dipole moments are vector quantities: they have magnitude and direction
polar vs nonpolar electron geometry
if bond dipoles dont cancel out (if lone pairs), then polar.
valence-bond theory
ORBITALS OVERLAP
- model of chemical bonding that combines lewis structures with orbitals. The buildup of electron density between two nuclei is visualized as occurring when a valence atomic orbital of one atom merges with that of another atom. The orbitals are said to share a region of space, aka to overlap.
- As the distance between the atoms decreases, the overlap region increases, so there is an increase in electron density between nuclei, and potential energy of the system decreases, and strength of the bond increases.
- However, if the nuclei become too close, energy increases due to electrostatic repulsion between nuclei
internuclear distance
The internuclear distance at minimum potential energy is the bond length. Bond length is the distance at which the attractive forces between unlike charges are balanced by the repulsion forces between like charges
hybrid orbitals
to explain geometries, assume that the atomic orbitals on an atom mix to form new hybrid orbitals. Hybridization: mixing atomic orbitals as atoms approach each other to form bonds. Number hybrid orbitals=number of atomic orbitals mixed
internuclear axis
The electron density in covalent bonds is concentrated symmetrically about the line connecting the nuclei (the internuclear axis)
sigma bonds
In single bonds, the line joining the two nuclei passes through the middle of the overlap region. These bonds are all called sigma () bonds. (the overlap region is sigma bond)
pi bonds
In multiple bonds, the sideways overlap of p orbitals (perpendicular to internuclear axis) produces as pi () bond. In pi bonds, the overlap regions lie above and below internuclear axis and there is no probability of finding the electron on the internuclear axis. Pi bonds are weaker than sigma bonds because the overlap is less in pi than sigma
sigma and pi bonds equal multiple bonds
Single bond=sigma bond, double bond=one sigma bond, one pi bond, triple bond=one sigma bond, two pi bonds
delocalized2
- In most molecules, Bonding electrons are localized: the sigma and pi electrons are associated totally with the two atoms that form the bond. In molecules with resonance structures with pi bonds: not entirely localized.
- Delocalized: cannot describe the pi bonds as individual electron-pair bonds between neighboring atoms (ex: benzene because it has two resonance structures)
molecular orbital theory5
- alternative model to explain bonding that describe the electrons in molecules by using specific wave functions called molecular orbitals (MO)
- An MO can hold max two electrons with opposite spins, has definite energy, and has electron density distribution. However, unlike atomic orbitals, MOs are associated with the entire molecule, not with a single atom
- Whenever two atomic orbitals overlap, two molecular orbitals form
- atomic orbs combine with other atomic orbs of similar energy
- as overlap increases, the energy of the bonding MO decreases and the energy of the antibonding MO increases
bonding MO
lower-energy MO concentrates electron density between two nuclei. Sausage shaped. Summed the two atomic orbitals into one. Electron is more stable (lower energy) because it is attracted to both nuclei. This is covalent bond
antibonding MO
higher-energy MO has very little electron density between the nuclei. Atom orbitals cancel each other in the region between the nuclei, greatest electron density on opposite sides of nuclei. An electron in this MO is repelled from boning region and is less stable (higher energy)
energy level diagram
shows interaction between two 1s atomic orbital and molecular orbitals that result. Bonding electrons=electrons occupying a bonding molecular orbital.
bond order
stability of covalent bond, =half the difference between the number of bonding electrons and the number of antibonding electrons. Bond order= ½ (bonding electrons – antibonding electrons). Bond order of 1=single bond, 2=double bond, 3=triple bond, 0=no bond exists.
first row diatomic molecules
only 1s atomic orbitals used and only sigma MOs produced
second row diatomic molecules5
- 2s, 2px, 2py, 2pz atomic orbitals used.
- 2s atomic orbs form sigma2s MO
- 2pz forms sigma2p MO
- 2px and 2py form pi2p MO
- The sigma2p MO has lower energy (more stable) than pi2p MO because there is greater overlap for sigma MOs
doubly degenerate
Both pi2p and pi*2p MO are doubly degenerate: two orbitals (same energy) MOs of each type.
second row MO using sp interactions2
- LARGE INTERACTION= HOPSCOTCH: B2, C2, and N2, the sigma2p MO is above the pi2p MOs in energy.
- SMALL INTERACTION= SIG, PI,PI, SIG: For O2, F2, and Ne2, the sigma2p MO is below the pi2p MOs in energy because of 2s-2p interactions
paramagnetism2
- molecules with one or more unpaired electrons are attracted into a magnetic field. The more unpaired electrons= stronger force of attraction
- gains weight if in magnetic field
diamagnetism2
- substances with no unpaired electrons are weakly repelled from a magnetic field. Weaker effect than paramagnetism
- loses weight if in a magnetic field
bond order/length/enthalpy trend
As bond orders increase=bond distance decrease and bond enthalpies increase
heteronuclear diatomic molecules2
- two atoms are not the same, (NO). if the atoms in heteronuclear diatomic molecule don’t differ too greatly in electronegativities, MO will be similar to homonuclear diatomics, except the atomic energies of the more electronegative atom will be lower in energy than those of the less electronegative element.
- An MO will have a greater contribution from the atomic orbital to which it is closer in energy. Not an equal mixture between the two atoms
