Molecular shape and reactivity

describe the shape and bonding in ethane

two tetrahedral carbon atoms

C-H bonds formed by overlap of carbon sp3 and hydrogen s orbitals

C-C bond formed by overlap of two carbon sp3 orbitals

increased bond length for C-C compared to C-H

bond angles near 109.5

describe the shape and bonding in ethene - hybridisation of carbon’s orbitals

each C forms 3 bonds but is bonded to only 3 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 and pointing to the corners of an equilateral triangle (C nucleus at centre of this triangle)

bond angles between sp2 orbitals close to 120

unhybridised p orbital perpendicular to the plane

describe the C=C bond in ethene

carbon atoms form two non-identical bonds with each other

one bond formed by overlap of one sp2 orbital from each carbon (sigma bond)

second bond formed by side-to-side overlap of the unhybridised p orbital (pi bond)

four electrons involved

compare a C=C bond with a C-C bond and triple C-C bond

triple C-C is both strongest and shortest, then C=C, with C-C longest and weakest

why does a rotational barrier exist around a C=C bond?

p orbitals must be well aligned parallel to each other above and below the plane for maximum overlap and formation of the pi bond

describe the bonding in ethyne

triple bond between two carbons, each carbon connected to only two atoms (C and H)

hybridisation of only two orbitals - two p orbitals are left non-hybridised

sp hybrid orbital derived from the combination of one s and one p atomic orbital

two sp hybrids are separated by an angle of 180

compare the hybridisation of atomic orbitals in ethane, ethene and ethyne

ethane - four sp3 orbitals (all p orbitals involved in hybridisation)

ethene - three sp2 orbitals, one p orbital unhybridised

ethyne - two sp orbitals, two p orbitals unhybridised

describe the carbon-carbon triple bond in ethyne

unhybridised p orbitals are perpendicular to each other and to the sp orbitals

two side-to-side overlaps to give two pi bonds

region of high electron density around the C-C axis

describe the bonding in a methyl cation

carbon is positively charged and bonded to 3 hydrogen atoms

three orbitals sp2 hybridised

positively charged C and 3 H atoms lie in a plane

non-hybridised p orbital remains empty and is perpendicular to the plane

describe the bonding in a methyl radical

C bonded to 3 H atoms and is sp2 hybridised

C and H atoms lie in a plane

radical has one more electron than the cation, which resides in the non-hybridised p orbital

describe the bonding in a methyl anion

C bonded to 3 H atoms

3 pairs of bonding electrons and 1 lone pair of electrons

electron repulsion minimised by adopting tetrahedral geometry

negatively charged C is sp3 hybridised

three bonds formed by sp3-s overlap

lone pair resides in the fourth sp3 orbital of the carbon

describe hybridisation in nitrogen/ammonia

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

experimental bond angle 107.3 (close to tetrahedral)

describe the bonding in amines

N has same geometry as in ammonia

C-N bonds formed by sp3-sp3 overlap

electron rich so nucleophilic

describe hybridisation in oxygen

O-H bond formed by sp3-s overlap (water, alcohols)

C-O bond formed by sp3-sp3 overlap (alcohols, esters)

lone pairs occupy the remaining two oxygen sp3 orbitals

bond angle 109.5

what form of hybridisation can form sigma and pi bonds?

sp3 sigma only, sp2 and sp sigma and pi

why is the bonding in HF stronger than in HCl?

in HF bond formed between H 1s and F 2sp3 orbitals, whilst in HCl bond formed between H 1s and Cl 3sp3 orbital

average distance of electrons from nucleus increases, and average electron density in 3sp3 orbital is lower than in 2sp3

electron density in region of s-sp3 overlap decreases as size of X increases

H-X bond becomes weaker and longer as size of X increases

what is the C-X bond (X=halogen) formed by?

sp3-sp3 overlap

how does C-X bonding change as size of X increases?

weakens

sp3 orbital from higher shell is used in bond, which has decreased electron density (+ is longer)

define electronegativity

the tendency of an atom to pull bonding electrons towards itself

what differentiates between polar covalent and ionic bonds?

polar covalent requires difference in electronegativity of atoms of <2, ionic of >2

what is 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

what is a dipole moment?

a measure of the net polarity of a molecule

results from the vector summation of all individual bond polarities and lone-pair contributions in the molecule

list the 4 main types of interactions between discrete molecules that make up covalent compounds, in order of increasing strength

1. Van der Waals forces (aka London)

2. dipole-dipole interactions

3. hydrogen bonding (bond strength can vary widely)

4. ion-ion (electrostatic)

describe Van der Waals forces

weak interactions caused by momentary changes in electron density in a molecule

the only attractive forces present in nonpolar compounds

what determines the strength of the van der Waals interactions between molecules?

surface area of the molecules

define polarizability

a measure of how the electron cloud around an atom responds to changes in its environment

which atoms are more polarizable?

larger atoms with stronger forces of attraction

what are dipole-dipole interactions?

the attractive forces between the permanent dipoles of two polar molecules

describe hydrogen bonding

short range, directional, inter- or intramolecular non-bonded interaction

between a hydrogen atom bonded to an electronegative atom (D) and another electronegative atom (A) (ie D-H…A)

D/A must be O, N or F

what are the terms used for the electron deficient hydrogen and the electron rich heteroatom in a hydrogen bond?

hydrogen bond donor (HBD) and hydrogen bond acceptor (HBA)

state 2 reasons for the importance of hydrogen bonding

responsible for unusual properties of water (MP, BP, surface tension etc)

drug molecules - receptor interactions

what is the most electronegative atom on the periodic table?

F

in what 2 conformations can alkyl chains be, and why?

staggered and eclipsed

there is free rotation around the C-C bond

describe what a staggered alkyl chain conformation is

the most stable, lowest-energy, 3D conformation of an alkane chain

substituents on adjacent carbon atoms are oriented as far apart from each other as possible (rotated by 60 degrees from eclipsed)

describe what an eclipsed alkyl chain conformation is

high-energy spatial arrangement of an alkane

substituents on adjacent carbon atoms are aligned with each other when viewed down the C-C single bond

groups are in the closest proximity, causing torsional strain and lower stability

what makes cycloalkanes more complicated than alkyl chains?

angular, torsional and steric strains

what is angle strain?

the increased potential energy and instability in a molecule caused by bond angles deviating from their ideal values (usually in sp3 carbons)

what is torsional strain?

the increase in the potential energy and instability of a molecule due to repulsion between electrons in bonds that do not share an atom (ie eclipsed conformation)

what is steric strain?

the increase in the potential energy and instability of a molecule caused by repulsive forces between electrons from non-bonded atoms that are forced too close together