inorganic bonding - valence bond theory

Question 1

electronegativity definition

Answer 1

the ability of an atom to attract electron density towards itself

Question 2

covalent bond lewis definition

Answer 2

electron pair equally shared between 2 atoms

Question 3

polar covalent bond definition

Answer 3

electron pair unequally shared between 2 atoms

Question 4

ionic bonding definition

Answer 4

electrons are transferred between atoms

Question 5

bond length definition

Answer 5

the mean distance between 2 atoms in a molecule, found from many measurements of that bond

Question 6

how is bond length determined for homonuclear diatomics?

Answer 6

bond length = 2 x atomic radius

Question 7

atomic radius definition

Answer 7

half the distance between nuclei in a homonuclear bond

Question 8

bond dissociation enthalpy definition

Answer 8

a measure of bond strength, the standard enthalpy change for the reaction in which that specific bond is broken (endothermic, so +ve)

Question 9

why do atoms form bonds?

Answer 9

atoms achieve stability by sharing electrons to attain the electron configuration of the closest noble gas = the octet rule!!

Question 10

how does electron sharing work in the lewis model of covalent bonding?

Answer 10

electrons are always shared in pairs, to achieve octet configuration - lone pairs are non involved in bonds

Question 11

state the bond orders of single, double and triple bonds

Answer 11

single - bond order = 1
double - bond order = 2
triple - bond order = 3

Question 12

how are lewis bonding models shown?

Answer 12

with dot and cross diagrams

Question 13

give 3 limitations of the lewis model

Answer 13

• cannot easily explain the existence of radicals
• it is difficult to explain expansion of the octet
• it cannot explain why O2 is paramagnetic (is a consequence of having unpaired electrons)

Question 14

outline the concept of valence bond theory

Answer 14

more complicated version of lewis model - suggests bonding involved interactions between atomic orbitals, atoms are brought together and allowed to interact, with all interactions between and involving electron pairs

Question 15

outline how the lennard-jones potential describes bonding

Answer 15

steep repulsive term and smoother attractive term, describes how as atoms get closer together
repulsion between nuclei dominates (destabilising), and as they get further apart they do not interact, graph tends to 0
graph minima = equilibrium bond length, the point where attraction dominates most causing a stable bond to form (stabilising)
distance from minima to x-axis = bond dissociation energy
distance from minima to y-axis = bond length

Question 16

what are boundary surfaces of atomic orbitals and what do they show?

Answer 16

boundary surface = the surface of a 3D rendering of an atomic orbital
indicate the volume of space in which the electron will spend most of its time

Question 17

what does a cross section of an atomic orbital show?

Answer 17

radial nodes and nodal planes

Question 18

what does ϕA(1)ϕB(2) mean?

Answer 18

this is a wavefunction - the first quantum mechanical theory of bonding expresses ideas about covalent bonding in terms of wave functions
ϕA(1)ϕB(2) describes the wavefunction for 2 atoms sufficiently isolated so there is no bond/interactions - H atoms particularly, as must be only 2 bodies, (1) and (2) represents the 2 electrons present, from each atom

Question 19

how does the wavefunction change when atoms are brought closer together?

Answer 19

when atoms are closer it is no longer possible to tell the electrons apart, e-1 is equally likely to be associated with either H atom
this is shown by adding a second term:
Ψ+ = ϕA(1)ϕB(2) + ϕA(2)ϕB(1) - if spins are paired
Ψ- = ϕA(1)ϕB(2) - ϕA(2)ϕB(1) - if spins are parallel

Question 20

what do the Ψ+ and Ψ- functions look like when plotted?

Answer 20

Ψ+ curve resembled lennard-jones potential, this curve is dominated by nucleus-electron attraction
Ψ- curve is a downwards exponential curve, it does not cross the x-axis, this curve is dominated by internuclear repulsion

Question 21

what information can we get from the Ψ+ wavefunction?

Answer 21

the equation/graph for Ψ+ gives calculated bond dissociation enthalpies that are less endothermic than experimental values, and longer calculated bond lengths

Question 22

why does the Ψ+ function give data that deviates from experimental data?

Answer 22

this approach underestimates the strength of the bond, because it assumes that 1 of the 2 electrons are always associated with each atom

Question 23

give the improved Ψ wavefunction that removes the restriction that one electron is always associated with an atom

Answer 23

Ψ = ϕA(1)ϕB(2) + ϕA(2)ϕB(1) + λϕA(1)ϕA(2) + λϕB(1)ϕB(2)
- this gives closer calculations

Question 24

what do the terms λϕA(1)ϕA(2) + λϕB(1)ϕB(2) mean?

Answer 24

λϕA(1)ϕA(2) = when both electrons are associated with hydrogen/atom A
λϕB(1)ϕB(2) = when both electrons are associated with hydrogen/atom B

Question 25

what would λ equal for H2?

Answer 25

λ < 1, as the terms in which both electrons are associated with 1 atom involve greater electron electron repulsion

Question 26

why is resonance so important in valence bond theory?

Answer 26

bonding in molecules/diatomics is a weighted average of its different resonance forms, some of which are more covalent, or more ionic - resonance forms help to give a complete picture of the bonding, they are not in equilibrium with each other and do not have independent existence

Question 27

what is the effect of resonance?

Answer 27

stabilising

Question 28

hybridisation definition

Answer 28

a way of mixing together the atomic orbitals on an atom as it interacts with another atom

Question 29

how do the energies of hybrid orbitals compare to the original orbitals?

Answer 29

energy of hybrid orbitals is a fraction of the way between the 2 types of orbitals that made it - specific fraction depends on specific character of the hybridised orbital

Question 30

can hybrid orbitals form π-bonds?

Answer 30

no, only p-orbitals can form π-bonds

Question 31

how does the structure of N2 provide evidence for hybridisation + what kind of hybridisation is present?

Answer 31

N2 displays sp hybridisation, without this interaction the lone pairs would lie in 2s orbitals therefore would not be directional - lone pairs can be observed by x-ray diffraction experiments and chemically by interaction with H+ or a transition metal

Question 32

what are the 3 main types of hybridisation?

Answer 32

sp - 1 s and 1 p orbital hybridise to form 2 sp orbitals and leaves 2 p-orbitals sp2 - 1 s and 2 p orbitals hybridise to form 3 sp2 orbitals and leaves 1 p-orbital sp3 - 1 s and all 3 p orbitals hybridise to form 4 sp3 orbitals

Question 33

how are polar covalent bonds formed?

Answer 33

in heteronuclear diatomics, the more electronegative element attracts the shared electron pair more towards itself, so electron density is not equally shared between atoms, making the bond polar

Question 34

how does electronegativity change across the periodic table?

Answer 34

down a group - decreases across a period L -> R - increases

Question 35

rewrite the wavefunction for a heteronuclear diatomic molecule

Answer 35

for A-B, where electronegativity A >> B: Ψ = ϕA(1)ϕB(2) + ϕA(2)ϕB(1) + xϕA(1)ϕA(2) + yϕB(1)ϕB(2) where x and y = constants, and x >> y - necessary to account for the A-B+ form which contributes more significantly than the A+B- form

Question 36

describe the interactions between bonding atoms in valence bond theory

Answer 36

all interactions are localised between pairs of atoms and involving pairs of electrons