The Ionic Model Flashcards
What does the van Arkel triangle of bonding look like? What does this tell us about bonding?
Bonding is not always restrictive to ionic/covalent/metallic and instead there are contributions for other types within each other
What are the assumptions of the ionic model?
Solids are hard, incompressible, non-polarisable, charged spheres
Charges are integer multiples
The only interaction between ions is electrostatic
What evidence is there supporting the ionic model?
Appearance: ionic typically non-volatile, brittle, transparent- share similar properties
Electrical conductivity only when molten or dissolved in a polar solvent
Spectroscopy and magnetism: similar absorption spectra/ magnetic properties as free-gas phase ions
X-ray diffraction: shows electron density with larger anions and smaller cations in a lattice
How can you determine the size of an ion?
The lattice parameter contains the size of r+ + r-
If the radius of one ion is known, subtract from 1/2 lattice parameter (assumes ions touching)
Or an estimate from Cn/Zeff = R
where Cn is a constant dependent on quantum number
Given in pm, 10^-12
Consult database for ion size
What are the trends for ionic radius in the periodic table and why? What about between cations and anions?
Across the period, size of the cations/anions decreasing
This is because Zeff is increasing as protons are added whilst shielding does not increase at the same rate
The outer electrons are more greatly attracted by the stronger nuclear charge
Down the group, ionic radius increases, as although Zeff increases, the increase in quantum number has the larger effect, and so overall the higher energy, more diffuse orbitals results in a larger ionic radius
Cations<Anions as cations have fewer electrons for the same nuclear charge
How does ionic radius change with coordination number and oxidation state and why?
Higher coordination numbers increase ionic radius
This is because the ion is no longer fully incompressible and so larger
Higher oxidation states mean fewer electrons for the same nuclear charge
This means the valence electrons are more strongly attracted to the nucleus, and so the orbitals contract resulting in a smaller ionic radius
Define the lattice energy of a compound?
The amount of energy required to disassociate 1 mol of an ionic solid into a gas of its ions at infinite separation, so from 0K
The negative of this is lattice enthalpy
How can lattice enthalpy be calculated from a thermodynamics perspective?
This quantity cannot be directly measured
How has the Madelung Constant been derived? And the Coulombic interaction component of the Born-Lande equation?
What is the Born-Lande equation? What are the two components to it?
The coulombic electrostatic attraction between ions
And the repulsion between same charges
From the derivation of the coulombic interactions, derive the Born-Lande equation? What are the main steps for the whole derivation?
Find the coulombic attractions from Coulombs law and summing all the interactions
Derive the Madelung constant by considering the interactions next to, same plane repulsion, to centre, 6, 12, 8, add r as a parameter
Born repulsion as inversely proportional to r to the power of n
Add them, from graph, need to minimise energy, so differentiate with respect to r, equate to 0
Find B (constant for repulsion) in terms of the other
Substitute back into the equation and factorise
1- 1/n not 1-r
How does the Born-Mayer equation differ from Born Lande?
Differ in terms of repulsion, with Mayer using ae^ -r/constant, which is related to the compressibility of the crystal
What is the Kapustinskii equation and what assumptions have been used/
The madelung constant has been replaced with 0.87v, where v is the number of species in the unit cell
The interatomic radius r (1/2 A) has been replaced by r+ + r-
n=9 (rocksalt for all)
Derive the Kapustinskii equation from the Born-Lande equation?
Make sure to change m to pm, and J to Kj
What is the value for e (fundamental charge?
1.602 x 10^-19
What is the value for epsilon 0?
8.85 x 10 ^ -12
What should a Hess cycle for NaCl look like?
Includes Formation of NaCl(S) from standard states
Sublimation and atomisation energies for both
Ionisation energy
Electron affinity
Lattice enthalpy
What is the synthesis and general reactions of the alkali metal hydrides?
Formed via direct reaction with hydrogen
M + 1/2 H2 ⟶ MH
All hydrolyse to formed MOH and H2
Used a deprotonating agents to form H2 and Na+
And to make other hydrides e.g
4LiH + AlCl3 ⟶ LiAlH4 + 3LiCl
What happens to the thermal stability of the G1 metal hydrides down the group? Using a Hess cycle, explain why this occurs?
What type of lattice are they?
Thermal stability decreases, decomposes at lower temperatures
Using the Gibbs equation at change in G=0 at equilibrium, decomposition at
T=H/S
Entropy change is the same for both, so H must be smaller for formation of the hydride
This is due to the lattice enthalpy decreasing down the group as the cation ionic radius increases, resulting in smaller charge densities, meaning decreasingly energetic and thermal stability
All NaCl type structures
What is the bonding like for the group 2 hydrides?
BeH2: covalent, linear molecule in gas phase, tetrahedral in solid
MgH2: Ionic, rutile, evident from XRD electron density
- reversibly, so potential for hydrogen fuel storage
CaH2, SrH2, BaH2: ionic, PbCl2 structure with 9 fold coordination
Compare the stability of the G1 vs G2 hydrides?
G2 higher stability, as they have a higher lattice enthalpy, arising from higher charges and a greater number of ions in the formula unit
Use Kapustinskii for comparison (radi similar)
What compounds can form from combustion of a metal in air?
Oxides, O (2-)
Peroxides (O-O) (2-)
Superoxides (O-O) (1-)
How do formation enthalpies relate to stability?
The more negative the enthalpy of formation, the more stable the compound
This is because more energy is released from formation, and so a lower energy, more stable compound
Why are the oxides more stable for smaller cations whilst peroxides/superoxides more stable for larger cations?
To form the oxide, an addition electron must be added to the O- ion, which is an endothermic process
This is most favourable when the energy given out as the lattice enthalpy is large enough to compensate for the additional energy requirement
For small ion, the lattice enthalpy is more negative, so does compensate, and so the oxide is more favourable as there is a better option
For larger ions, the lattice is enthalpy is not as negative, so rather only the first electron affinity for superoxides is favourable (although all theoretically are thermodynamically favourable, but superoxides the most)
So the superoxides decompose into peroxides, and peroxides into oxides when the cation is small enough as their is a more thermodynamically favourable product available
What are the thermodynamically favourable compounds for G1 and G2? And their solid structure type?
Lithium Oxide- antifluorite
Sodium Peroxide, K/Rb/Cs superoxide- all with peroxide/superoxide subunit
BeO- covalent, Wurtzite
MgO,CaO,SrO- rocksalt
BaO2- peroxide, but decomposes upon heating
What is the thermochemical radius and how can it be calculated?
A way of finding the radius of ions from thermochemical experiments
e.g difficult to identify radius of CO3 2-
so instead calculate the lattice enthalpy using Hess Cycles and experimental data
Then rearrange for the anion/cation radius
Why is lattice enthalpy not exactly the same as internal energy change? What do our equations give and why are they valid?
Lattice enthalpy defined from absolute 0K
Equations derived to give change in internal energy
change in H= change in U + change in pv
However, as pv is small, almost negligible so we can assume enthalpy change is approximately internal energy change
What happens to the thermal stability of the carbonates down a group? Why?
Thermal stability increases
This is because cation size increases, meaning the stability of the oxide decreases, as the lattice enthalpy will fall
Therefore, the difference in energy between the carbonate and oxide is smaller down the group, meaning decomposition is favourable but becomes less favourable
Decomposition relies on the high lattice enthalpy of MO, but as this decreases, higher temperatures are needed for decomposition
Does CaCl exist and why?
Theoretically, there is an exothermic change for formation of CaCl, however, CaCl will disproportionate into CaCl2 so it does not exist
When will a reaction favour disproportionation and how can this be calculated?
When the enthalpy for disproportionation < 0, so more stable and exothermic
Start with the hypothetical disproportion reaction, and use Hess Cycles to calculate the enthalpy change
i.e ionisation energies, electron affinities, atomisation…
Out of the G2 halides, why is Iodide the species which would most likely stabilise MX the most?
To form MI2, needs energy for additional ionisation which is very endothermic
As the electron affinity for iodine is the lowest for the halides, and formation enthalpy of iodide the least exothermic, the addition of an extra iodide is the least likely to compensate for the additional energy of ionisation
Additionally, the lattice enthalpy of MI2 is the smallest of the halides due to the large ionic radius of I-
These factors would limit the ability to form MX2, however, the M+ cation is still too large to adequately stabilise so disproportionation still occurs
Why does NaF2 not form?
The second ionisation energy is far to endothermic to be compensated by the energy released from formation of NaF2 and additional F-
This is because the additional electron is being removed from a shell of a lower quantum number, experiencing a far larger Zeff
Why does the lattice enthalpy AgCl not agree with the Kapustinskii prediction? How can this be alleviated?
There is a degree of covalency within AgCl, not purely/perfectly ionic
This means there are additional covalent interactions to overcome which the Kapustinskii equation does not consider, so the prediction is too low
Instead, use the thermochemical radius for a more accurate estimate
Why does the addition of fluoride ions support higher oxidation states relative to the other halides?
The enthalpy of formation of F- is the most exothermic out of the halides, as it has the most negative electron affinity
Additionally the difference in lattice enthalpy between Mxn and Mxn+1 is largest for F-, as F- has the smallest anion radius
This results in F- having the most negative enthalpy of formation of a higher oxidation state of the halides
What happens to ions dissolved in a polar solvent?
The ions will be surrounded by molecules of the polar solvent, and coordinate to the ion
The lattice will break apart
Stabilises the ion, with ion-dipole interactions the main interaction
What happens to an ion in a structure-less medium?
Average ion-dipole interactions are considered
Polarisable, so described by the dielectric constant which tells us about the polarisability of a given medium in response to an electric field
What Hess Cycle should be considered for whether a substance dissolves?
Includes lattice enthalpy, hydration enthalpies of gaseous ions, and solution enthalpies for dissolving
if sol enthalpies <0, dissolves
What is the equation for the hydration enthalpy of a substance?
What simplifications can be made to known equations to enable comparisons of solubility? How can this be interpreted?
From the equations, the only time when the lattice enthalpy is small but hydration is large, is when one ion is small and the other large
There is the greatest solubility when there is a mismatch in ion size
What is the chelate effect? Why does it occur?
Larger, the more polydentate ligands will replace less polydentate ligands already coordinated
This is because there is an increase in entropy from the ligands liberated from the metal ion, replaced by fewer ligands
Therefore, the gibbs energy for the reaction is more negative, a driving force
The entropy increase is the dominant factor, will very low enthalpy values
What happens when an alkali metal is placed in ammonia?
Alkali metals dissolve in liquid ammonia
Na(s) –> Na+ (solv) + e- (solv)
Free electrons localised in the cavity of the solvent
Optical observation with electron excitation, colours
However solution metastable (will decay over time)
e.g NH3 + e- –> NH2- (sol) + 1/2 H2(g)
What are grignard reagents? When can Mg be found in an uncommon state?
Mg-R compounds, formed from reactions with RX, good nucleophiles
Mg(I) when coordinated by very bulky carbon ligands with nitrogen
What are carbides? And dizaenides?
[C2] 2- or C 4-
[N2] 2-, from thermal decomposition of azides
How would you justify your answers from a hess cycle from a stability perspect?
The enthalpy of the reaction > 0, and so the reaction is not thermodynamically feasible from an enthalpy perspective
The energy demand for ionisation e.g, is not compensated by the additional stabilisation of the (lattice enthalpy) so the reaction is unfavourable
What should you include when showing Kapustinskii from Born-Lande?
Why the approximations have been used
- 0.87v as for most solids a good replacement for their madelung constant
