S2.2 Covalent Model - Hybridisation & Bonds Flashcards
Orbital Hybridisation
concept of mixing atomic orbitals to form new hybrid orbitals that can accommodate the bonding and lone pair electrons in a molecule
SP3
creates 4 equal length bonds (exmaple carbon)
* ground state electron configuration changes = excitation
* 2s orbitals are promoted to higher energy
* creates 4 new bonding orbitals of equal energy
* make uo tetradhedral shape = 109.5 degrees
SP2
when 3 bonds of equal length & shape want to be formed
* 2s orbital joins with p to make SP2 and 1 remaining 2p orbital
* sp3 orbitals form sigma bond
* unhybridised p orbital forms pi bond
* trigonal planar = 120 degrees
SP
- 2s orbitals joins with p orbital and leaves 2 unhybridised p orbitals and 2sp orbitals
- sigma bond between sp orbitals
- 2 pi bonds
Benzen hybridisation
- all 6 carbons in benzene are sp2 hybridised
- each C has sigma bond between adjacent C
- each C has delocalised pi bond with overlapping unhybridised p orbitals
How are electron pairs arranged
Electron pairs are arranged in a way that they will maximise seperation and minimize repulsion
Electron domain geometry
the shape of the arrangement of electron domains surrounding a central atom in a molecule or ion
electron domain
a region in which bonding & non bonding pairs of electrons are most likely to be found
Rules for bond angles
each lone pair reduces the bond angle by 2.5 degrees
Trend in decreasing repulsion strength
bonding pair & bonding pair< non bonding pair & bonding pair< non bonding pair & non bonding pair
How does repulsion affect bonding angle
the higher the repulsion, the smaller the bonding angle
expanded octets/hypervalent
when an atom can accept more than 8 electrons using their available d orbitals (3rd period atoms can)
linear
no lone pairs
180 degrees
CO2
3 electron pairs
bent/angular (one lone pair)
<120 degrees
trigonal planar
no lone pairs
120 degrees
BCl3
bent or angular (two lone pairs)
two lone pairs
<109.5 degrees
H2O
Octahedral
no lone pairs
90 degrees
SF6
Trigonal planar
no lone pairs
90 degrees and 120 degrees
PCl5
Seesaw/Bisphenoidal
<90 degrees and 117 degrees
5 electron domains
1 lone pair
SF4
T shape/ triangular bypyramidal
TBP= 90 degrees (full bonded)
5 electron domains
T shape = <90 degrees
2 lone pairs
BrF3
linear (3 lone pairs)
180 degrees
3 lone pairs
I3
square planar
90 degrees
2 lone pairs
XeF4
square pyramidal
90 degrees
1 lone pair
BF5
formal charge
useful to determine possible lewis structures and which minimizes the charges
based on finding how many e belong to the atom and the difference between how many it actually has
=> most likely structure will always have charge closest to zero
bond strength
measure of the energy required to break a bond
bond length
distance between the nuclei of two bonded atoms
How bond strength works
- double and triple bonds have greater strength
- more shared e- = higher electrostatic attraction between e and nuclei
- bond is harder to overcome
How bond length works
- double and triple bonds have shorter bond length
- the more shared electrons the stronger the electrostatic attraction
- two nuclei are pulled closer together
Evidence contradicting benzene
Bond lengths= if structure consisted of alternate double 6 single bonds = some shorter than others
=> X Ray Crystallygrophy showed that all bonds are euqal in length
Enthalpy of hydrogenation= should be 3x that of cyclyhexane but it isnt = difference is due to extra stability from delocalisation of electrons = less exothermic
Chemical evidence= test for double bond; YET benzene does not decolorise bromine water
=> resonance; electrons are delocalised in pi bond throughout ring= give greater stability= lack of reactivity
What causes the proof of benzene to be so extroardinary
- lack of expected nethalpy of combustion
- lack of reactivity towards electrophiles
How the Ozone layer protects from UV
Formation
* Oxygen absorbes UV light (240nm) shining on O2= 2 oxygen radicals
* radical reacts with 02 = o3
Decomposition
* Ozone absorbs Uv light (330nm) = Oxygen radical and 02
* Ozone reacts with =xygen radicla = 2O2
=> reduced UV concentration reation earth
Why O2 and O3 are dissociated by different wavelengths
- covalent molecules absord UV radiation when radiation has sufficient energy to break bonds within molecule = dissociates
- Ozone has weaker O-O bonds than oxygen (double bond)
- bonds are stronger than a single bond but weaker than a double bond
- require a longer wavelength of light to break their bonds than O=O double bond in oxygen molecules
- but a shorter wavelength than oxygen–oxygen single bond.
Resonance
occurs when there is more than one possible position for a double bond in a molecule by delocalising a pair of electrons
Delocalisation
occurs when an electron is not fixed to one position in a molecule but can move freely across bonds
How can you identify resonnace structures
- dashed line
- when multiple version of the lewis structure can be drawn
Bond order
the number of bonds between 2 atoms
Equation showing NO removing ozone
NO∙(g)+O 3 (g)→NO2∙(g)+O 2 (g) 1 mark
NO2∙(g)+O∙(g)→NO∙(g)+O2(g)
NO∙(g)+O∙(g)→NO∙(g)+O 2(g)
