SECTION 4 - CHEMICAL BONDING II Flashcards
VSEPR
Valence shell electron repulsion theory that tells us the 3 dimensional shape of a molecule
Electron group geometry (egg)
Based on the number of electrons surrounding central atom (bonded atoms + lone pairs)
5 basic electron group geometry shapes
- Linear (2 e)
- Trigonal planar (3 e)
- Tetrahedral (4 e)
- Trigonal bypyramidal (5 e)
- Octahedral (6 e)
Dash lines
Extend BEHIND plane
Solid thick lines
Protrude OUT plane
Molecular shape (mg)
Only count atoms bonded to central atom
egg = mg
When all electron groups are atoms
egg doesnt equal mg when
Lone pairs present
5 molecular geometry shape
- Bent
- Seesaw
- T shape
- Square pyramidal
- Square planar
BLT
BP LP
2 1
2 2
BENT
2 3
LINEAR
3 1
TRIGONAL PYRAMIDAL
3 2
T SHAPE
4 1
SEESAW
4 2
SQUARE PLANAR
5 1
SQUARE PYRAMIDAL
Axial bonding position
Up/down
Equatorial bonding position
Horizontal/diagnol
- LP go on equatorial
Dipole moment
Measure of molecular polarity with unit of Debye (D). 1D = 3.34x10^-30cm
Dipole moment in diatomic molecule
There is only 1 bond so the dipole determined by bond
Dipole moment in polyatomic molecule
The total polarity determined by bond polarity and bond angle
Arrows facing out from one another
U = 10 (not polar)
Arrows facing towards each other
Polar
Valence bond theory
Tells us how electrons interact to hold atoms together into molecules and defines covalent bonds as electron density between 2 atoms as a result of constructive interference
Ideal distance of valence bond theory
Occurs with overlap between partially filed atomic bonds on each atom and the overlap creates new covalent bond
Inphase
Atoms that are in phase will overlap constructively and form bond
Outphase
Atoms that are out phase creates region of 0 electron density (node)
Node
N-1 (region of 0 electron density)
Sigma bond (3)
- Formed by overlap of atomic orbitals head to head
- All single bonds are sigma bonds
- Free rotation allowed
Pi bond (5)
- Formed by overlap of p orbitals side to side
- Multiple bonds consists of one sigma bond and some pi bonds
- Contains 2 regions of electron density
- Rotation not allowed
- Overlapping must occur in similar phase orbitals
Sp hybridization
2 electron groups
Sp2 hybridization
3 electron groups
Sp3 hybridization
4 electron groups
Sp3d hybridization
5 electron groups
Sp3d hybridization
6 electron groups
How to valence theory (4)
- Identify valence orbitals of bonding atoms
- Sketch valence orbital
- Bring in atoms to be bonded to central atom and sketch overlap
- Describe structure
Hybrid orbitals
The “mixing’ of orbitals
Relationship between the # of atomic orbitals and the # of hybrid orbitals produced
Must equal
Molecular orbital theory
States that electron in molecules are not assigned to individual chemical bonds, but rather moving under the influence of atomic nuclei in molecule. Electrons are assigned to molecules as a whole
Constructiveness interference
Forms a bonding molecular orbital with high region of electron density between nuclei
Destructiveness interference
Forms an anti bonding molecular orbital with a node between nuclei
Molecular orbital rules (6)
- # of atomic orbitals combined = # of molecular orbitals
- Bonding orbitals hold molecules together with low energy
- Antibonding orbitals do not help hold molecules together and are high energy
- Electrons fill lowest energy first
- Degenerate orbitals fill singly before pairing
- Maximum 2 electrons per orbital
*
High energy antibonding molecular orbital
Bonding molecular orbital
Force goes towards atom
Antibonding molecular orbital
Forces repel
Bond order
Bonded electrons-antibonded electorns/2
Bond order 1
Single bond
Bond order 2
Double bond
Bond order 1/2
Single bond with one electron
Bond order 0
Indicates that a bond is unlikely to form
Higher bond order
Indicates stronger bond, high stability, low energy, smaller bond length
Bond order 3
Triple bond
Paramagnetic
Contains unpaired electrons
Diamagnetic
Contains paired electrons
HOMO
Highest energy orbitals (has electrons)
LUMO
Lowest energy orbital (no electrons)
3 states of matter
Solid, liquid, gas
Condensed phases
Solid and liquids. This is because they are all packed closer together
How does temperature impact kinetic energy
Higher temperature = higher kinetic energy
Intermolecular forces
Attractive/repulsive force between molecules of a substance that impact physical properties of matter. Without this, all matter would exist as a gas
Strong intermolecular force
Requires more thermal energy to break the molecules apart = higher boiling/melting point
3 types of intermolecular force
- London dispersion
- Dipole-dipole
- Hydrogen bonding
Intramolecular force
Found within the molecule that keeps it together
London dispersion
Present in all neutral molecules/atoms and the weakest bond. Caused by when electrons are not evenly distributed inside orbital and results in small partial charges
Polarizability
Tendency of an atom/molecule to form a distorted electron cloud based on 3 conditions
Polarizability dependancy (3)
- Larger atoms = more polarizable
- Larger molecules = more polarizable
- Larger surface area = stronger dispersion force
Dipole-dipole (3)
- Occurs when a molecule has a permanent dipole and partial negative lines up with partial positive
- The more polar, the stronger the dipole
- Stronger the dipole = higher the boiling point (harder to pull)
Hydrogen bonding
Strongest intermolecular force that occurs when a partial positive hydrogen atom is covalently bonded to small electronegative partial negative atom N, O, F
Order in strength of intermolecular force
London dispersion –> dipole-dipole –> hydrogen bond
Water physical property
Each H2O molecule can form up to 4 hydrogen bonds at one time which makes water liquid at room temperature
Why is liquid water more dense than solid water
Because of the orientation of hydrogen bonds that cause the molecules to push further apart in solid state, lowering the density
