A - Structure

Question 1

What is the value of one atomic mass unit, amu?

Answer 1

1.6605 x 10^-27 kg

Question 2

What is the Aufbau principle?

Answer 2

Lowest energy orbitals are occupied first

Question 3

What is the Pauli principle?

Answer 3

Only 2 electrons may occupy each orbital

Question 4

What function describes an electron?

Answer 4

A wavefunction ie a sin wave. It is a mathematical description of the distribution of electrons in space and time.

Question 5

What is the uncertainty principle?

Answer 5

The position and energy of an electron cannot be determined. The energy of an electron can be known, but where it is exactly is unknown - only know where it is likely to be located, in an orbital.

Question 6

What are the 4 quantum numbers, and what do they determine?

Answer 6

Principal quantum number: n
How big the orbital is (the shell) -> 1, 2, 3, 4 etc

Angular momentum quantum number: l -> 0 - (n-1)
The shape of the orbital
l = 0 -> s l = 1 -> p l = 2 -> d

magnetic quantum number: ml = -l ….. +l
Gives the orientation of the orbital

Spin quantum number: ms = + or - 1/2

Question 7

What are wavefunctions used to calculate?

Answer 7

The radial distribution function of the orbital. This is the chance of finding an electron at a certain distance.

Question 8

What is the Pauli exclusion principle?

Answer 8

No 2 electrons in an atom can have the same 4 quantum numbers.

Question 9

What is Hund’s rule of maximum multiplicity?

Answer 9

A greater total spin state makes the resulting atom more stable. So, if two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs.
Electrons in separate orbitals will have less electrostatic repulsion, so will be in a lower energy arrangement. Additionally, they will have lower energy if their spins are parallel due to reduced repulsion from spin correlation - parallel spins mean they stay further away from each other

Question 10

What is the chemical reactivity of each block in the periodic table?

Answer 10

s block = forms stable cations
p block = forms multiple bonds
d block = range of oxidation states

Question 11

What is the radial distribution function of the 1s orbital?

Answer 11

Starts at 0
Rapidly increases until the maximum
Probability decreases and tends towards 0

Question 12

Why is the 2s orbital filled before the 2p orbital?

Answer 12

2p orbital has similar probability distribution to 1s. 2s probability distribution has a small peak close to the nucleus, then falls to zero again -> node.
Probability then rises again to form a larger peak than the first.
This means there is a greater probability of an electron being found near the nucleus in the 2s orbital than the 2p orbital, hence it is filled first.

Question 13

What is the Lewis model of bonding?

Answer 13

Covalent bonding occurs when valence electrons are shared between two atoms.
Maximum stability is achieved when each has a filled valence shell.

Question 14

What is the octet rule?

Answer 14

Each atom acquires shares in electrons until its valence shell has 8 electrons.

Question 15

When do atoms form a bond?

Answer 15

The potential energy varies with distance. A bond is formed at the distance with the lowest energy - the energy well.

Question 16

What is a resonance structure?

Answer 16

When a molecule has more than one satisfactory Lewis structure. The true structure is a resonance hybrid - it has equal contributions from all the resonance forms

Question 17

What are hypervalent compounds?

Answer 17

Compounds that require more than an octet of electrons in order to draw a satisfactory Lewis structure, eg PCl5, SF6. This is possible through the use of low-lying d orbitals

Question 18

What is VSEPR theory?

Answer 18

Valence Shell Electron Pair Repulsion.
Electrons in bonds and lone pairs can be considered as charge clouds which repel each other. The lowest energy arrangement is the one where they are as far apart as possible.

Question 19

In VSEPR theory, how do much do lone pairs and bonding pairs repel each other?

Answer 19

Most repulsion = lone pair to lone pair
Medium repulsion = bond pair to lone pair
Least repulsion = bond pair to bond pair

Question 20

Will a lone pair choose an equatorial position or an axial position, and why?

Answer 20

Equatorial. There is less repulsion from 2 axial bonding pairs in equatorial position than there is from 3 equatorial bonding pairs in the axial position.

Question 21

What are the problems with the Lewis model?

Answer 21

It does not explain hypervalency well, nor does it explain how/why O2 is paramagnetic.

Question 22

What is valence-bond theory?

Answer 22

Half-filled atomic orbitals on two atoms overlap to create a bond filled with paired electrons, ie a bonding orbital

Question 23

How does the valence-bond theory explain relative bond strengths?

Answer 23

More orbital overlap means a stronger bond is formed. Therefore in larger atoms, electrons are not as tightly held because the orbital is bigger and there is less overlap.

Question 24

What is the orbital hybridisation theory?

Answer 24

Orbitals can combine constructively to form hybrid orbitals which all have the same energy. This explains molecular geometry.

Question 25

How are orbitals hybridised in CH4?

Answer 25

The s and p valence orbitals combine to make 4 equivalent sp3 orbitals. These all have the same energy and are directed towards the corner of a tetrahedron.
Each sp3 orbitals then overlaps with a 1s H orbital. They overlap head on and so form a σ bond.

Question 26

How are orbitals hybridised in C2H4?

Answer 26

One s orbital and the px and py orbitals combine to form 3 equivalent sp2 orbitals. These form a trigonal planar shape. pz remains unhybridised.
An sp2 orbital from each C atom overlaps to form a σ bond, and the other sp2 orbitals overlap with a 1s H orbital to form σ bonds. The unhybridised pz orbital on each C atom then overlap sideways to form a π bond.

Question 27

What is molecular orbital theory?

Answer 27

Orbitals are wavefunctions and so can combine constructively or destructively. Combining constructively = electron density increases, forming a bonding molecular orbital. There is a high probability of finding an electron between the atoms, so it contributes towards bonding. Combining destructively = electron density decreases, forming an antibonding molecular orbital. There is zero probability of finding an electron between the atoms, so it negates from bonding.

Question 28

How is the energy of MOs different to that of AOs?

Answer 28

The bonding molecular orbital is lower in energy than the atomic orbitals. The antibonding orbital is higher in energy than the atomic orbitals. So if electrons occupy the MO, the energy has been lowered and bonding is favourable.

Question 29

What is the equation for bond order?

Answer 29

bond order = (number of bonding electrons - number of antibonding electrons) / 2

Question 30

What do different bond orders correspond to?

Answer 30

Bond order = 0 -> no bond
Bond order = 1 -> single bond
Bond order = 1/2 -> bond with half the strength of a single bond; the bond is transient
Bond order = 2 -> double bond

Question 31

What conditions must be kept when forming MOs?

Answer 31

The AOs must be close in energy in order to be able to overlap.
The degree of overlap is reflected in the separation in energy of bonding and antibonding orbitals. That is, the better the overlap, the lower the bonding MO is in energy, the stronger the bond

Question 32

How are π bonds formed in MO theory?

Answer 32

pz orbitals overlap head on. px and py orbitals can then only overlap side on, and hence form π bonds.

Question 33

What are non-bonding MOs?

Answer 33

Some molecules form non-bonding MOs. These do not change in energy level from the original AOs, and have no effect on the bonding of the molecule.

Question 34

Can the order of MOs change?

Answer 34

Yes, they change depending on the molecule. As the atoms involved get larger, the valence orbitals will be closer in energy and so will be able to overlap forming MOs.

Question 35

How does MO theory explain the paramagnetism of O2?

Answer 35

The MO diagram shows that the 1π* orbitals each have an unpaired electron, thus O2 is affected by an applied external magnetic field and is paramagnetic.

Question 36

What are the properties of a covalent bond?

Answer 36

Links atoms within a molecule
The bonds are strong
The bonds are short ranged and localised between the two atoms
They are highly directional as the geometry is well-defined
Generally long lasting

Question 37

What are the properties of non-covalent interactions?

Answer 37

Intermolecular and intramolecular
Weak interactions
Long ranged ie distance is longer than a bond
They are less directional and well-defined than covalent bonds
Generally short lived or fluctuating

Question 38

What non-covalent interactions make up Van de Waal’s forces?

Answer 38

London disperion + dipole-dipole + dipole-induced dipole

Question 39

How does electronegativity change across the Periodic Table?

Answer 39

It increases across a period: the nuclear charge increases due to more protons
It decreases down the group: outer electrons get further away from the nucleus, so the nuclear attraction felt by them decreases

Question 40

How can the polarity of a bond be calculated?

Answer 40

The difference between the electronegativity, χ, of the two atoms. If χa - χb >2, then the bond is so polarised that it can be considered ionic, ie there was total electron transfer

Question 41

What is the equation for the energy between freely rotating dipolar molecules?

Answer 41

E ≈ - 1 / r^6

E is also inversely dependent on T - as T increases, the interaction weakens.

Question 42

How are London dispersion interactions formed?

Answer 42

Fluctuating electron density causes an instantaneous dipole to be formed in a non-polar molecule. This can then induce a dipole in the neighbouring non-polar molecule. The induced dipoles attract each other.

Question 43

What is the polarisability of a molecule?

Answer 43

How easily it is polarised. The more electrons in outer shells, the less nuclear attraction they feel, so the more polarisable they are. Also, the size of the dipole-inducing the polarisation affects it.
The more polarisable, the stronger the dispersion interactions = more energy needed to break interactions = higher boiling point

Question 44

What is a dielectric constant?

Answer 44

How the polarity of a solvent is measured. It is how well the material stores electrical energy when exposed to an electric field. More polar molecules have a higher dielectric constant because they correctly align their positive and negative ends to the field.

Question 45

How do ionic substances dissolve in polar liquids?

Answer 45

The dipoles of the polar molecules align to the field of the ion, lowering the self-energy of the ion.
If the solvent is water, this is called hydration.

Question 46

What is a hydrogen bond?

Answer 46

A H bond forms between two electronegative atoms when one is bonded to a H and the other has a lone pair of electrons with which to attract the H atom.

Question 47

What are the properties of H bonds?

Answer 47

They are very short ranged because they involve the overlap of atomic orbitals of two electronegative atoms and H.
They are usually linear due to the head on atomic overlap.
H bonds are strong due to orbital overlap - weaker than covalent bonds but stronger than Van de Waal’s

Question 48

Why is the boiling point of water so high?

Answer 48

H2O can form up to 4 H bonds per molecule as it has 2 donors and 2 acceptors. The strong H bonding network significantly increases its boiling point, because for a molecule to enter the gas phase, it must break its H bonds.

Question 49

Why is ice less dense than water?

Answer 49

When H2O freezes, H bonds are longer and locked into a lattice, leaving lots of space between the molecules. As a liquid however, these bonds are formed and destroyed because of the kinetic energy of the molecules. So as a liquid the molecules are closer together.

Question 50

What is the Grotthuss mechanism?

Answer 50

It explains why the conduction of H+ through water is much faster than similarly charged ions.
H bonds can become covalent bonds and vice versa. All H bonds and covalent bonds switch, so charge can be transferred through the solution without the movement of any atoms.

Question 51

How does TLC chromatography work?

Answer 51

The solid stationary phase is an SiO2 layer on the plate surface, which displays free Si-O-H groups.
The mobile layer is a liquid solvent driven by capillary forces
The analyte (solute) will have different interactions with the stationary and mobile phase
The analyte will adsorb to the stationary phase sooner and more strongly if there are strong interactions -> H bonding, dipole-charge, dipole-dipole, dispersion
If the analyte interacts favourably with the mobile phase, it will travel further.

Question 52

What are the assumptions of VSEPR theory?

Answer 52

1- electrons in bonds and lone pairs can be seen as charge clouds that repel each other. The lowest energy arrangement is when these charge clouds are as far apart as possible, and this determines the equilibrium molecular shape.
2- electrons in lone pairs repel more than electrons in bonding pairs, and this gives rise to bent molecules, ie. molecules in which the angles between bonds are different than in molecules with tetrahedral geometry.

Question 53

What happens when a non-polar solvent is dissolved in water?

Answer 53

H bonding cannot occur between water and the hydrophobic molecule. There are fewer ways of arranging the H bonding network because H2O molecules near the non-polar molecules have fewer options of where they can H bond. This reduces the entropy but increases the free energy.
It is therefore unfavourable for hydrophobic molecules to be surrounded by water - they tend to separate to minimise the contact area.

Question 54

What is the hydrophobic effect?

Answer 54

The collecting together of non-polar molecules. It is driven by a reduction in their unfavourable interactions with water.

Question 55

What is an amphiphile?

Answer 55

A molecule that contains both hydrophobic and hydrophilic parts.

Question 56

What role does symmetry play in the formation of dipoles?

Answer 56

A dipole moment is a vectorial property. If a molecule is symmetrical, the dipoles cancel out and it has no overall permanent dipole moment.
H2O is non-linear due to the lone pairs of electrons on the O atom. The two dipoles do not cancel out so the molecule has an overall dipole moment.

Question 57

How does the polarity of molecules relate to their boiling point?

Answer 57

The more polarised a molecule is, the stronger and more numerous the intermolecular forces, so the more energy is required to break these interactions. The more energy that has to be put into the system in order to vaporise the molecules, the higher the boiling point.

Question 58

How can you determine the miscibility of two substances?

Answer 58

Like dissolves like:

Polar dissolves polar, non-polar dissolves non-polar

Question 59

What drives the self-assembly of amphiphiles in water?

Answer 59

The hydrophobic effect:

amphiphiles self-assemble to minimise entropically unfavourable interactions with water

Question 60

What does it mean when a bond has a dipole?

Answer 60

When two atoms with different electronegativities are bonded, the electronic charge is not evenly spread across the bond. Two poles are formed, one with negative electron density and one with positive electron density.

Question 61

What is the formula for the dipole moment of a bond?

Answer 61

ν = e x d
where e = charge on the atom (either delta positive or negative, as they have the same magnitude
and d = distance between the two charges

Question 62

What determines the overall dipole moment of a molecule?

Answer 62

Both the magnitude and the direction of the individual bond dipole moments: the vector sum