Inorganic Acids and Bases

Question 1

What is a Brønsted Acid/Base? What is a Lewis Acid/Base? How about in general?

Answer 1

Brønsted Acid: Proton donor
Brønsted Base: Proton acceptor

i.e HNO₃ + H₂O → NO₃ - + H₃O +

 Acid      Base    Base   Acid

Lewis Acid: Electron pair acceptor
Lewis Base: Electron pair donor

Generally acids have positive character that is decreased upon reacting with a base
And a base has negative character decreased when reacted with an acid

Question 2

How are Ka and pKa defined?

Answer 2

Ka = [H₃O⁺] [A⁻] / [HA]

where we assuming the concentration of water is the same as activity, 1

pKa = -log (Ka)

Question 3

What is autoprotolysis? How does it give rise to Kw? And when do pH and pKa align?

Answer 3

2H₂O → H₃O⁺ + OH⁻

Water undergoes autoionisation

Using the Ka expression,
Kw= [H₃O⁺] [OH⁻] = 10^-14

pH=pKa when HA is 50% disassociated

Question 4

How does the pKa of polyprotic acids change with successive dissociations?

Answer 4

pKa increases, disassociation less favourable each time
As a negative ion is gaining a further negative charge, destabilising the conjugate base, as with the electrostatic model

Successive deprotonations occur when the pH is high enough

Question 5

How does the acidity of M(H20)6 complexes arise and change across the period?

Answer 5

The positive metal cations polarised the O-H bond, resulting in a partial positive charge on the hydrogen

Across the period, Zeff increases, as there is increased nuclear charge and decrease in ionic radius
Therefore, charge density increases across the period, resulting in greater polarity, and so increased acidity

Question 6

How does the acidity of M(H20)6 complexes change down the group?

Answer 6

Decrease in (Zeff/n)^2 down the group, and more diffuse orbitals so there is a larger ionic radius, resulting in smaller charge densities
Therefore less polar bonds and decreased acidity

(although this is an ionic model, and larger metals have some covalency)

Question 7

What can happen if the pH of aqua metal acids is increased?

Answer 7

Polymerisation, with bridging oxygens
Precipitation, with hydroxide ligands

Question 8

What are hydroxoacids? What are oxoacids?

Answer 8

Hydroxo: H-O acidic proton without adjacent oxygens
Oxo: H-O acidic proton with adjacent oxygens, bonded to the same central atom

Question 9

What are Paulings rules for Oxo acids? Why on the whole do these trends apply?

Answer 9

1) pKa1 = 8 - 5p
where p is the number of X=O, (not the O of O-H)
2) Successive pKas reduce by around 5 units

This is empirical, trends that fit the data rather than an explanation
Increased M=O results in greater resonance stabilisation, so more stable conjugate bases

Question 10

Why is boric acid less acidic than expected?

Answer 10

Boric acid does not deprotonate from the hydroxyl as expected
Instead, water coordinates to the empty p orbital , and a proton is lost from this water instead

Question 11

What is proton affinity a measure of?

Answer 11

H+ (g) + A (g) →HA +b(g)

The energy released from the species accepting a proton

In the hess cycle, as the process is exothermic, it is - this enthalpy for proton affinity
As it is for electron affinity

Question 12

How does proton affinity change across a period?

Answer 12

Across the period, the increase in electron affinity dominates the reaction, arising from increase Zeff
Therefore, proton affinity decreases

Question 13

How does proton affinity change down a group?

Answer 13

Down the groups, the decrease in BDE dominates, arising from larger atom sizes resulting in reduced electron-nuclear attraction, and poorer orbital overlap as more diffuse
Therefore, proton affinity decreases

Question 14

Why is NMe3 a weaker base than NH3 in water?

Answer 14

Using the Hess Cycle, the enthalpy changes accompanying protonation is dependent on hydration enthalpies and proton affinities
The more substituted compounds have more less exothermic hydration enthalpies of the reactants/products as a smaller charge density, so less easily solvated
Hydrogen bonding is less favourable in more substituted compounds, with less hydrogens present, so less stabilisation of the conjugate acid

Question 15

What is solvent levelling?

Answer 15

If an acid is dissolved in the solvent, if the acid is more acidic than the solvent and its conjugate acid, the acid itself will be fully dissociated
If comparing two acids, and both acids are more acidic than the conjugate acid of the solvent, they will both give the same pKa as fully dissociated
The same applies to bases with the conjugate base of the solvent

See the examples below, as an equilibria exists with acetic acid, it can be used to compare acid strength

Question 16

Why types of acids and bases can be compared in strength in ammonia?

Answer 16

Acids which are less acidic than NH4+, but more acidic than NH3

Bases which are less basic than NH2-, but more basic than NH3

Question 17

How do the orbitals interact when a lewis acid reacts with a lewis base? What affects the strength of this interaction?

Answer 17

The HOMO of the base overlaps with the LUMO of the acid
The smaller the energy gap, the stronger the bond formed, the greater the stabilisation

Question 18

What are examples of species which behave as lewis acids and how?

Answer 18

Metal cations, complexes e.g hexa-aqua cobalt, coordination of water

G3 compounds e.g B(Me)3, as it contains an empty p orbital, or hypervalent compounds such as SbF6- via use of the empty d orbital

BeCl2 acting as a base and acid, AlCl3 dimers

Carbon dioxide, via moving electron density to the oxygen

Question 19

How can the strength of a lewis acid/base pair be determined?

Answer 19

Using an equilibrium constant calculation

Question 20

How do formation constant values change for the halides with Mg vs Hg? Why?

Answer 20

Mg: F>Cl>Br>I
Hg: F<Cl<Br<I

Based on Hard-Hard vs Soft-Soft interactions

Question 21

What are hard species? What are soft species ? And examples ?

Answer 21

Hard: Small, high charge density, low polarisability e.g H+, Li+ , Ni 2+, H2O, F-

Soft: large, low charge density, high polarisability e.g CU+, Cd2+, Hg 2+, I- , CO

Question 22

How do hard/soft interactions occur?

Answer 22

Hard:Hard, electrostatic in nature, large differences between HOMO-LUMO and so ionic

Soft:Soft, covalent in nature, small HOMO-LUMO gap

Question 23

How can Drago-Wayland parameters be used to assess acid-base strength?

Answer 23

Calculation using electrostatic and covalent contributions to the acid and base
But does not take into account solvation effects

Question 24

What factors normally drive an equilibrium to the right?

Answer 24

Solvation effects: smaller ions are more easily solvated, so can be more easily stabilised

BDE, stronger bonds forming drive equilibriums

Question 25

What is the donor number of a solvent?

Answer 25

A measure of lewis basicity, defined as the negative reaction enthalpy of the bases reaction with SbF5 to form SbF5B

Question 26

What is the acceptor number of a solvent?

Answer 26

A measure of lewis acidity, quantified by the 31P NMR shift with triethylphosphine oxide as a reference base dissolved in pure solvent But does not take into account electrostatic/ covalent differences

Question 27

How does the methyl position in pyridine affect basicity? And frustrated systems?

Answer 27

If too close to the nitrogen, steric repulsion making adduct formation less exothermic If further away, increase adduct formation via increasing electron density via hyperconjugation and inductive effects Sterics too large to enable bond formation

Question 28

What is the trend in the silicon tetrahalides lewis acidity and why?

Answer 28

F>Cl>Br>I As greater electronegativity, more electron deficient centre, more electrophilic and susceptible to basic attack

Question 29

What is the trend in boron trihalide lewis acidity and why?

Answer 29

F

Question 30

What are superacids? How is magic acid formed?

Answer 30

Species which donate protons more efficiently than pure sulphuric acid Normally formed when a lewis acid is dissolved in a Bronsted acid 2FSO₃H + SbF₅ → [H₂FSO₃]⁺ + [SbF₅SO₃F]⁻ Here, the acid acts as a lewis acid and protonates the other acid further, forming the super acid or SbF₅ + 2HF →SbF₆ ⁻ + [H₂F]⁺

Question 31

How can the standard electrode potential of a Zn/Zn2+ system be calculated experimentally? What is a reductant and oxidant?

Answer 31

Connect the half cell to a standard hydrogen half cell, a Pt electrode in pH 0 acid, 1 bar H2 gas, and measure voltage with a half cell Reductant= Reducing agent = oxidised Oxidant= Oxidising agent= reduced

Question 32

What is the equation relating a reduction half reaction, Gibbs energy, and electrode potentials? And for a cell?

Answer 32

ΔG (red) = -nFE° Where n is the number of electrons transferred And F is faraday's constant, 96485 c mol-1 ΔG (rxn) = -nFEcell°

Question 33

What is the equation for E cell from the half cells?

Answer 33

Ecell° = E red - E ox But both values need to be the reduction potential voltages (Even for the reaction being oxidised)

Question 34

How can you determine if a redox reaction will occur from electrode potentials?

Answer 34

First compare the reduction potentials, the more negative one will undergo oxidation Calculate E cell from the reduction potentials If positive, the reaction is thermodynamically spontaneous, as ΔG (rxn) = -nFEcell°, and so negative, assuming kinetics are not too slow

Question 35

How can you calculate the equilibrium constant from electrode potentials? What's the derivation?

Answer 35

Careful with the reaction If a scalar multiple, the number of electrons differ, and so you get different K/delta G values, so use the equation given Standard conditions

Question 36

How does the electrode potentials for M/M+ vary for group 1?

Answer 36

Quite similar Two main factors dominate the enthalpy change for formation of the ion Down the group, larger ions so less exothermic hydration enthalpies, less negative formation But small ionisation energies, so more negative formations Opposing factors result in Na having the least negative enthalpy, so least negative electrode potential (linked by delta G, and as this is an oxidation, Goxd= nFE)

Question 37

What are the main steps for determining half equations for a reaction where atoms are exchanged?

Answer 37

Look at the change in oxidation states to reflect electrons Add protons/water to either side until charge and species balance

Question 38

What is the half equation for water? When is this important?

Answer 38

When comparing to another chemical species, and seeing if a reaction is possible with the water system i.e will the species be oxidised by water

Question 39

How can an electrode potential be calculated at non-standard conditions? How can this be derived? And a half cell?

Question 40

What equation can be used to show how electrode potential varies with pH for a reaction involving protons?

Answer 40

Valid where the number of protons is the same as number of electrons if not, multiply by m/n where m is the number of protons and n number of electrons

Question 41

How can complexation help extract metals from ores?

Answer 41

Sometimes formation of the metal ion is not spontaneous But forming a complex stabilises this oxidation state, making formation of the ion more favourable e.g Au with CN

Question 42

How do Latimer diagrams work?

Answer 42

Species placed in increasing oxidation state, with reduction potentials between them per electron as reduction potential is independent of electron number To calculate reduction potentials between species, multiply the reduction potential by the number of electrons and add sum, divide by the number of electrons Including multiplying by - when an oxidation state below 0

Question 43

How can a latimer diagram be used to plot a frost diagram?

Answer 43

Form a table with each species in order of oxidation state Run a cumulative tab of relative Gibbs free energy, by multiplying the reduction potential by number of electrons, and summing as you go down This is proportional Gibbs free energy If below 0 in oxidation state, we include multiplying by the negative sign

Question 44

What does a Frost diagram look like?

Answer 44

X axis: Oxidation state Y axis: NE, proportion to G Label the points with the species

Question 45

When do species disproportionate? And comproportionate?

Answer 45

Disproportionate when the species in between is above the connecting line, as on average, this species is higher in gibbs energy Comproportionate when the species in between is below the connecting line, as on average, this species is lower in gibbs energy

Question 46

How do you compare which species are the better oxidant from a frost diagram?

Answer 46

Join lines from the species being oxidised to the oxidant The steepest line has the species which is the greatest oxidant, most positive gradient, i.e the steepest line to being connected to it

Question 47

Qualitatively, how does pH affect reduction and oxidation?

Answer 47

Oxidation is supported at higher pH, easier to access higher oxidation states in alkaline conditions Reduction is supported lower pH Unless a redox couple has no proton dependence, in which case pH does not affect Gibbs energy

Question 48

Why does the X2/X- redox couple decrease in reduction potential down the group?

Answer 48

Although atomisation is highest for chlorine, the electron affinity is the most exothermic for chlorine, as highest (zeff/n), and most negative hydration enthalpy

Question 49

How do pourbaix diagrams work?

Answer 49

Electrode potential against pH Vertical line: couple has no pH dependence but redox e.g Fe2+/Fe3+ Horizontal: couple has pH dependence but not redox e.g Fe2+/Fe(OH)2 Diagonal: dependence on both, gradient dependent on m/n, from the Nernst pH equation

Question 50

How can you determine the pH at which a precipitate will form looking at a redox couple?

Answer 50

Use the solubility product, and then kw

Question 51

How can pourbaix diagrams be used to determine stability of a compound in water?

Answer 51

Shows the electrode potentials at different pH ranges Species, with electrode potentials, above the oxygen/water boundary , want to be reduced, and so oxidise water into oxygen Species below the hydrogen/H+ species will want to be oxidised more, and so reduce H+ Shows the stability of different substances in water based on electrode potentials, at different pH

Question 52

How do you calculate electrode potentials at different pHs?

Answer 52

E= E° - 0.059 pH x m/n where mass is the n of protons and n the n of electrons

Question 53

What does the slope of a frost diagram represent?

Answer 53

The reduction potential Negative, and the more negative, the more favourable, lower in energy, for the reduction

Question 54

How do you calculate the reduction potential between two species from a latimer diagram?

Answer 54

Multiply electrons transferred by reduction potential and sum Divide by the total number of electrons

Question 55

What is the gradient of a Pourbaix diagram?

Answer 55

Compare to a y=mx +C of the ph dependent Nernst equation Gradient= -0.059(m/n), with pH the x of the graph