Molecular Shape and Reactivity Flashcards
What does orbital hybridisation account for?
The bonding of carbon atoms into chains and rings.
Describe the bonding in ethane.
- Two carbon atoms - both are tetrahedral
- C-H bonds formed by overlap of carbon sp3 and hydrogen s orbitals
- C-C bond formed by overlap pf two carbon sp3 orbitals
- Increased bond length for C-C
- Bond angles are near 109.5
Describe the bonds in ethene.
- Each carbon forms four bonds but is bonded to only three atoms
- Hybridisation of an s atomic orbital with two 2p atomic - one p orbital is left non-hybridised
- Electron repulsion minimised by the three sp2 orbitals lying in a plane
- Bond angles close to 120
- Unhybridised p orbital is perpendicular to the plane
Describe how the double bond forms in ethene.
- One bond formed by overlap of one sp2 orbital from each carbon
- second bond formed by side-to-side overlap of the unhybridised p orbitals
How many electrons are involved in the formation of a C=C bond?
four
Compare a C=C bond to a C-C bond.
C=C bond is stronger and shorter than C-C
Why does the rotational barrier exist in ethene?
the p orbitals are well aligned for maximum overlap leading to formation of the pi bond.
Describe the bonds in ethyne.
- forms a triple bond
- hybridisation of only two orbitals - two p orbitals are left non-hybridised
- sp hybrid orbital derived from the combo of one s and one p atomic orbital
- the two sp hybrids are separated by an angle of 180
How many electrons are involved in the formation of a C C triple bond?
six
Compare a C triple bond to a C=C
C triple bond is stronger and shorter than a C=C bond
Describe the bonding in methyl cation.
- Carbon is positively charged and bonded to three hydrogen atoms
- Three orbitals are hybridised
- sp2 hybridisation
- The positively charged C and the three H atoms lie in a plane
- Non-hybridised p orbital remains empty and is perpendicular to the plane
Describe the bonding in methyl radical
- Carbon is bonded to three hydrogen atoms and is sp2 hybridised
- The C and three H atoms lie in a plane
- Radical has one more electron than the cation
- Resides in the non-hybridised p orbital
Describe the bonding in methyl anion
- Carbon is bonded to three hydrogen atoms
- Three pairs of bonding electrons and one lone pair of electrons
- Electron repulsion minimised by adopting tetrahedral geometry
- Negatively charged carbon is sp3 hybridised
- Three bonds formed by sp3 - s overlap
- Lone pair resides in the fourth sp3 orbital of the carbon
Describe the bonding in nitrogen
- Nitrogen hybridises to form four sp3 orbitals
- One of the four sp3 orbitals is occupied by two non-bonding electrons
- N-H bonds formed by sp3 - s overlap
Describe the bonding in amines
C-N bonds formed by sp3-sp3 overlap
- Hydrogens may be replaced by 1/2/3 alkyl or aryl groups
- Electron rich therefore nucleophilic
Describe the bonding in ammonium cations
- Possesses four identical N-H bonds and no lone pairs
- Nitrogen has four sp3 hybridised orbitals
- All bond angles = 109.5
Describe the bonding in oxygen
- O-H bonds formed by sp3-s orbital overlap (water, alcohols)
- C-O bond formed by sp3-sp3 orbital overlap (alcohols, Esthers)
- Lone pairs occupy the remaining two oxygen sp3 orbitals
- Bond angle smaller than in methane (109.5)
Describe the bonding in sulphur atoms
- S-H bonds formed by sp3-s orbital overlap (water, alcohols)
- S-C bond formed by sp3-sp3 orbital overlap (alcohols, Esthers)
- Lone pairs occupy the remaining two sulphur sp3 orbitals
Describe the bonding in hydrogen halides
- Bond formed between s-sp3
(HF = s-2sp3 and HCL = s-3sp3) - Electron density in region s-sp3 overlap decreases as the size of X increases
- HX bond becomes weaker and longer as the X increases
Describe the bonding in Alkyl Halides
- C-X bond formed by sp3-sp3 orbital overlap
- As size of halogen increases sp3 orbital from higher shell used in bond = decreased electron density and longer and weaker bond
Define electronegativity.
- The tendency of an atom to pull bonding electrons toward itself
Define polar covalent bond.
- A covalent bond in which the electron distribution between atoms is unsymmetrical
- Bond polarity is due to difference in electronegativity
Describe the inductive effect.
- An atom’s ability to polarise a bond
- Electron-attracting or electron-withdrawing effect transmitted through s bonds
- Electronegative elements have an electron withdrawing inductive effect
Define dipole moment.
- A measure of the net polarity of a molecule