Acid–Base and Donor–Acceptor Chemistry
These are my notes on the chapter that quietly shows up in almost every chemistry exam worth naming — NEET, JEE, IIT-JAM, BITSAT, GATE, CSIR-NET, TGT/PGT, all of it. If you've ever wondered why your textbook gives you five or six different "definitions" of an acid instead of just one, this chapter is the answer. Each one was built to cover something the last one couldn't.
6.1 What's an acid–base "model," and why are there so many?
Think of each model as a lens. One lens lets you see a certain group of reactions clearly and blurs everything else; a different lens does the opposite. None of the older models were wrong as such — they were just narrow. A new one usually appeared because chemists ran into a reaction the old definition simply couldn't describe.
The motivation hasn't really changed in two centuries: spot a trend, and you stop having to test every compound from scratch. If you already understand why HI beats HF in acid strength, you can extend that reasoning instead of running the experiment again.
The earliest observations were almost embarrassingly simple. Acids tasted sour, bases felt slippery and bitter, certain dyes changed colour depending on which one you dropped in. No NMR, no X-ray crystallography, just taste and touch and colour. Definitions only got sharper once the tools did.
A quick timeline
| Model / Scientist | Year | Acid Definition | Base Definition | Classic Example |
|---|---|---|---|---|
| Liebig | 1838 | Has H replaceable by metal | Reacts with acid | HNO₃ / NaOH |
| Arrhenius | 1894 | Gives H⁺ (H₃O⁺) in water | Gives OH⁻ in water | HCl / NaOH |
| Brønsted–Lowry | 1923 | H⁺ donor (proton donor) | H⁺ acceptor | H₃O⁺ / H₂O |
| Lewis | 1923 | Electron-pair acceptor | Electron-pair donor | Ag⁺ / NH₃ |
| Ingold–Robinson | 1932 | Electrophile | Nucleophile | BF₃ / NH₃ |
| Lux–Flood | 1939 | Oxide ion acceptor | Oxide ion donor | SiO₂ / CaO |
| Frontier Orbitals | 1960s | LUMO of acceptor | HOMO of donor | BF₃ / NH₃ |
6.2 The Arrhenius concept
This is the oldest model that actually holds up, and it followed soon after Arrhenius and Ostwald demonstrated that ions genuinely exist in solution. That finding earned Arrhenius the 1903 Nobel Prize, which tells you how big a deal it was at the time.
- An Arrhenius acid produces H₃O⁺ in water. HCl → H⁺ + Cl⁻ is the standard example.
- An Arrhenius base produces OH⁻ in water. NaOH → Na⁺ + OH⁻.
- Neutralisation is just H₃O⁺ + OH⁻ → 2H₂O, with a salt left behind.
The limitation sits right there in the name. It only works in water. Step into an organic solvent, the gas phase, a solid, or liquid ammonia, and this model has nothing useful to say. That gap is exactly what pushed chemistry toward Brønsted and Lowry.
6.3 Brønsted–Lowry concept
In 1923, Brønsted and Lowry, working separately and unaware of each other, reframed the whole idea around a single particle: the proton.
- A Brønsted acid donates a proton (H⁺).
- A Brønsted base accepts one.
Lose a proton and you become a conjugate base. Gain one and you become a conjugate acid. Every reaction in this framework involves these conjugate pairs.
One rule is worth memorising properly: stronger acid plus stronger base gives weaker acid plus weaker base. Equilibrium always settles toward the weaker side.
Amphoteric solvents
Some solvents can act as either acid or base depending on what they meet. These are called amphoteric, and they matter because the same solute can behave very differently depending on which one you dissolve it in.
| Solvent | Conjugate Acid | Conjugate Base | pKion (25°C) |
|---|---|---|---|
| H₂SO₄ | H₃SO₄⁺ | HSO₄⁻ | 3.4 |
| HF | H₂F⁺ | HF₂⁻ | ~12 |
| H₂O | H₃O⁺ | OH⁻ | 14.0 |
| CH₃COOH | CH₃COOH₂⁺ | CH₃COO⁻ | 14.45 |
| NH₃ | NH₄⁺ | NH₂⁻ | 27 |
The levelling effect
This is one of those ideas that trips up a lot of students. In water, no acid can be stronger than H₃O⁺. HCl, HNO₃, HClO₄ and H₂SO₄ all count as "strong acids," and in water they end up looking exactly equal in strength, because each one ionises completely to give H₃O⁺. Water flattens, or "levels," them all to the same value.
Reactions in non-aqueous solvents
Liquid ammonia gets used when water would simply destroy whatever you're trying to study. Strong bases like NaNH₂, for instance, react violently with water. In liquid NH₃ the acid-base chemistry mirrors water almost exactly:
Superacids
George Olah won the 1994 Nobel Prize for developing superacids, meaning acids stronger than 100% sulfuric acid. Strength here is measured using the Hammett acidity function H₀, and the more negative the value, the stronger the acid.
| Superacid | H₀ value |
|---|---|
| HF | −11.0 |
| H₂SO₄ (100%) | −11.9 |
| HClO₄ | −13.0 |
| Triflic acid (HSO₃CF₃) | −14.6 |
| Fluorosulfonic acid (HSO₃F) | −15.6 |
| Magic Acid (HSO₃F:SbF₅) | −21 to −25 |
| Fluoroantimonic acid (HF:SbF₅) | −21 to −28 |
Magic Acid earned its name because it's strong enough to protonate hydrocarbons — famously, it can dissolve a candle. It's a mixture of fluorosulfonic acid and antimony pentafluoride.
Gas-phase acidity and proton affinity
If you want basicity measured with zero solvent interference, you go to the gas phase. Proton affinity (PA) is the enthalpy change when a base picks up H⁺ in the gas phase. Gas-phase basicity (GB) is the same idea, just measured as a free energy change instead. For most organic bases, GB lands somewhere between 700 and 1000 kJ/mol.
Brønsted superbases
A superbase, by definition, has a gas-phase proton affinity above 1000 kJ/mol. Oddly enough, these compounds can be fairly weak in water while being extremely strong in organic solvents. DBU (1,8-diazabicyclo[5.4.0]undec-7-ene) is a common one, with PA = 1048 kJ/mol, and it gets used constantly in organic synthesis. The "proton sponge," 1,8-bis(dimethylamino)naphthalene, has PA = 1028 kJ/mol, and its strength comes from relieving lone-pair repulsion plus forming a strong internal hydrogen bond once protonated. Among the alkali hydroxides, CsOH has the highest proton affinity at 1118 kJ/mol, and the order LiOH < NaOH < KOH < CsOH tracks neatly with how ionic the M–OH bond becomes.
Trends in basicity: induction, sterics, solvation
In the gas phase, alkyl groups push electron density toward nitrogen, so basicity climbs as you add methyl groups: NMe₃ > NHMe₂ > NH₂Me > NH₃.
Dissolve everything in water, though, and the order scrambles. Protonated tertiary amines have fewer N–H bonds available to hydrogen-bond with water, so they get less stabilisation than expected, which makes them weaker bases than the gas-phase numbers suggest. The actual aqueous order ends up being NHMe₂ > NH₂Me > NMe₃ > NH₃.
Binary hydride acidity
Two trends run in opposite directions here, and it helps to keep them separate. Going down a group (H₂O, H₂S, H₂Se), acidity increases, because the larger atom forms a weaker bond, and a weaker H–X bond breaks more easily. Going across a period (NH₃, H₂O, HF), acidity also increases, but for a different reason — the more electronegative element stabilises the conjugate base better.
Oxyacid strength
For oxyacids of the same element, more oxygens means a stronger acid. The chlorine series makes this obvious: HOCl < HClO₂ < HClO₃ < HClO₄, with pKₐ values of 7.2, 2, −1, and −10. Each extra oxygen pulls electron density away from the O–H bond through induction, and also spreads the negative charge across the conjugate base through resonance. More delocalisation means a more stable, weaker conjugate base, which means a stronger acid.
For polyprotic acids, each successive pKₐ jumps by roughly five units. Phosphoric acid is the standard example: pKₐ₁ = 2.15, pKₐ₂ = 7.20, pKₐ₃ = 12.37.
Acidity of metal cations in water
It's easy to forget that plenty of "neutral-looking" metal salts are actually acidic once dissolved. A highly charged metal ion pulls electron density out of the water molecules bound to it, weakening their O–H bonds:
The pattern is straightforward: higher charge and smaller radius make a stronger acid. Push the charge high enough, 4+ or above, and the ion can't even survive as a simple aquo-complex; it only exists as an oxyanion, like MnO₄⁻ or CrO₄²⁻.
| Metal Ion | Kₐ (298 K) | Metal Ion | Kₐ (298 K) |
|---|---|---|---|
| Fe³⁺ | 6.7 × 10⁻³ | Fe²⁺ | 5 × 10⁻⁹ |
| Cr³⁺ | 1.6 × 10⁻⁴ | Cu²⁺ | 5 × 10⁻¹⁰ |
| Al³⁺ | 1.1 × 10⁻⁵ | Ni²⁺ | 5 × 10⁻¹⁰ |
| Sc³⁺ | 1.1 × 10⁻⁵ | Zn²⁺ | 2.5 × 10⁻¹⁰ |
6.4 Lewis acids, Lewis bases, and frontier orbitals
G.N. Lewis put forward this model in 1923 too, the same year as Brønsted and Lowry, entirely independently. It turned out to be the broadest and most useful of all the definitions.
- A Lewis base donates an electron pair (also called a nucleophile).
- A Lewis acid accepts one (also called an electrophile).
- What forms is called an adduct, held together by a coordinate covalent bond, where both electrons originally came from the base.
BF₃ and NH₃ is the textbook pairing. BF₃ is trigonal planar with an empty p orbital on boron; NH₃ has a lone pair on nitrogen ready to give away. Nitrogen's HOMO reaches into boron's empty LUMO, and the geometry around boron flexes from flat (120°) to nearly tetrahedral as the adduct forms.
When the Lewis acid is a metal ion, the resulting adduct gets a special name: a coordination compound.
Frontier molecular orbitals
The HOMO–LUMO picture is the cleanest way to think about why a Lewis acid–base reaction happens at all. Electrons flow from the HOMO of the base into the LUMO of the acid. When those orbitals sit close in energy and overlap well, a stable adduct forms. When the gap is too large, electron transfer happens instead, which is a redox reaction rather than an adduct.
Water is a nice illustration of how flexible this framework is, because it can play four different roles depending on what it reacts with:
- React it with calcium and water acts as an oxidising agent: calcium gives electrons to water's LUMO and gets oxidised, while water is reduced to H₂ and OH⁻.
- Dissolve a chloride salt and water becomes a Lewis acid, accepting electron density from Cl⁻ through an O–H antibonding orbital.
- Dissolve magnesium and water flips to a Lewis base, donating a lone pair into Mg²⁺'s empty 3s orbital.
- React water with fluorine gas and water itself gets oxidised, donating electrons and acting as a reducing agent.
Watching I₂ act as a Lewis acid
Iodine is a handy probe for Lewis basicity because its colour shifts with solvent. In hexane or the gas phase it's violet. In benzene it turns reddish-violet as a charge-transfer band appears near 300 nm. In methanol, a stronger donor solvent, it goes yellow-brown and that CT band shifts further to around 240 nm. In aqueous KI, where iodide is an excellent donor that forms I₃⁻, it turns brown.
What's happening underneath: a donor solvent feeds electron density into iodine's σ* LUMO, and the π* → σ* transition shifts to higher energy, a blue shift. The stronger the Lewis base, the bigger that shift.
The BF₃ affinity scale
Because BF₃ is such a well-behaved reference acid, chemists use it to rank Lewis bases by how strongly they bind to it. A higher BF₃ affinity means a stronger Lewis base.
| Lewis Base | BF₃ Affinity (kJ/mol) |
|---|---|
| 4-dimethylaminopyridine | 151.55 (strongest) |
| trimethylamine | 139.53 |
| pyridine | 128.08 |
| 2-methylpyridine | 123.44 |
| tetrahydrofuran (THF) | 90.40 |
| 2-tert-butylpyridine | 80.10 |
| tetrahydrothiophene | 51.62 |
Halogen bonds
X₂ molecules and interhalogens like ICl can also form coordinate bonds with Lewis bases, and these get called halogen bonds. The geometry sits close to 180° because the base donates straight along the halogen's σ* LUMO. They behave a lot like hydrogen bonds, and they've found real use in drug design.
Inductive effects on Lewis acidity and basicity
Alkyl groups raise electron density at nitrogen or phosphorus, which strengthens the base. Electronegative substituents do the opposite, which is why PF₃ is a much weaker base than PH₃.
Boron halides are a fun exception to work through. Pure electronegativity logic says BF₃ should be the strongest Lewis acid of the group, but it isn't, and the reason is π back-bonding. Boron and fluorine form a short, strong bond with real π character, which pushes electron density back onto boron and dulls its Lewis acidity. The actual order runs BF₃ < BCl₃ < BBr₃ < BI₃: as the halogen gets bigger, the B–X bond stretches, π overlap weakens, and Lewis acidity climbs back up.
Steric effects: F-strain and B-strain
H.C. Brown worked out that sterics can override the electronic trends entirely. Front strain (F-strain) is when bulky groups on the approaching acid and base get in each other's way: 2,6-dimethylpyridine reacting with BF₃ is the classic case, where the ortho methyl groups physically block boron's approach. Back strain (B-strain) shows up after the adduct forms, when VSEPR geometry pushes groups into each other, as with NMe₃ and BMe₃. There's also internal strain (I-strain), which comes from electronic differences between similar molecules rather than physical bulk.
That's a genuinely strange flip. Toward the small H⁺, more alkyl groups means more electron density and a stronger base, hence the first order. React the same series with the bulkier BF₃ or BMe₃, though, and the order reverses almost completely. The ortho substituents that helped basicity now physically block the much bigger Lewis acid from getting close.
Frustrated Lewis pairs
Sometimes a Lewis acid and base are both so bulky that they can't form a normal adduct. Chemists call this combination "frustrated." Rather than just sitting there unreactive, frustrated Lewis pairs (FLPs) find other things to react with, especially small molecules. Stephan's group pioneered this using the very bulky, very acidic B(C₆F₅)₃ paired with hindered phosphines, and showed it could reversibly activate H₂, the first time a non-metal system had managed that. FLPs can also activate CO₂ and N₂O, and they're being explored as metal-free catalysts that avoid expensive or toxic transition metals.
6.5 Intermolecular forces
Hydrogen bonding
IUPAC's definition is X–H···B, where X is more electronegative than hydrogen and B is the atom donating electron density. Three things contribute: an electrostatic pull between the polar X–H and electron-rich B, which is usually the biggest contributor; some covalent or charge-transfer character, where B's HOMO reaches into the σ* LUMO of X–H; and a smaller dispersion component.
Strength increases as X gets more electronegative: N–H···B < O–H···B < F–H···B. You can spot a hydrogen bond experimentally in a few ways. The X–H···B angle tends toward 180° when the bond is strong. The X–H stretching frequency shifts to lower energy (a red shift) in IR. The bridging hydrogen shows a downfield shift in ¹H NMR.
Here's a question worth sitting with: why is [H₃O···OH₂]⁺ around six times stronger than a plain H₂O···OH₂ hydrogen bond? Gilli's pKₐ-equalisation idea answers it. Define ΔpKₐ as the pKₐ of the donor acid minus the pKₐ of the conjugate acid of the acceptor. The closer that number sits to zero, the stronger the hydrogen bond.
| |ΔpKₐ| range | H-Bond Strength |
|---|---|
| 0 to ±3 | Strong |
| ±3 to ±11 | Medium-Strong |
| ±11 to ±15 | Medium |
| ±15 to ±21 | Medium-Weak |
| >±31 | Weak |
Working it through: for [OH₃···OH₂]⁺, the donor is H₃O⁺ (pKₐ = −1.7), and the conjugate acid of the acceptor water is also H₃O⁺ (pKₐ = −1.7), so ΔpKₐ = 0 and the bond is strong. For plain H₂O···OH₂, the donor is water (pKₐ = 15.7) and the conjugate acid of the acceptor is H₃O⁺ (pKₐ = −1.7), giving ΔpKₐ = 17.4, only medium strength.
Host–guest interactions
π–π stacking between large aromatic systems can build some genuinely elegant structures. Buckminsterfullerene (C₆₀) sitting inside a corannulene-based host (C₆₀H₂₄) is one example: a ball-and-socket arrangement held together purely by π–π interactions, with the closest C···C contact measured at 312.8 pm. Systems like this show up in molecular electronics and photovoltaics research.
6.6 Hard and Soft Acids and Bases (HSAB)
Pearson proposed HSAB theory in 1963 to explain something the simple acid-base strength scale couldn't: why certain combinations bond far more tightly than their raw "strength" would suggest. It's turned into a reliable predictor of solubility, equilibrium position, and even which atom a ligand binds through.
The rule of thumb is short. Hard acids prefer hard bases; soft acids prefer soft bases. Hard species are small, carry a concentrated charge, resist being polarised, and have a wide HOMO–LUMO gap. Soft species are the opposite: large, diffuse charge, easily polarised, narrow HOMO–LUMO gap.
| Category | Hard | Borderline | Soft |
|---|---|---|---|
| Acids | H⁺, Li⁺, Na⁺, Mg²⁺, Al³⁺, Cr³⁺, Fe³⁺, BF₃ | Fe²⁺, Co²⁺, Ni²⁺, Cu²⁺, Zn²⁺ | Cu⁺, Ag⁺, Au⁺, Hg²⁺, Pd²⁺, Pt²⁺, I₂, Br₂, BH₃ |
| Bases | F⁻, OH⁻, H₂O, NH₃, CO₃²⁻, SO₄²⁻, RO⁻ | Br⁻, NO₂⁻, SO₃²⁻, N₃⁻ | I⁻, S²⁻, CN⁻, CO, PR₃, R₂S, SCN⁻ (via S) |
Three classic examples worth knowing
Silver halide solubility tells the story cleanly. AgF is quite soluble (Ksp = 205) because it pairs a soft acid with a hard base, a mismatch. Move down the halides and solubility drops steadily: AgCl (1.8 × 10⁻¹⁰), AgBr (5.2 × 10⁻¹³), and AgI, which is barely soluble at all (8.3 × 10⁻¹⁷), soft Ag⁺ paired with soft I⁻. Flip to lithium halides and the trend reverses completely: LiF, a hard-hard pair, is the least soluble of the lithium halides.
Thiocyanate, SCN⁻, can bind through either end, sulfur (soft) or nitrogen (hard), and which one it picks depends on the metal. Soft Hg²⁺ binds through sulfur, giving [Hg(SCN)₄]²⁻; hard Zn²⁺ binds through nitrogen instead, giving [Zn(NCS)₄]²⁻. Pd²⁺ and Pt²⁺, both soft, go through sulfur; borderline Ni²⁺ and Cu²⁺ go through nitrogen.
And then there's [CH₃Hg(H₂O)]⁺, which reacts dramatically faster with larger, softer halides: its equilibrium constant with HI is 10¹⁸, against just 4.5 × 10⁻² with HF. Soft mercury and soft iodide simply suit each other far better than mercury and the hard fluoride do.
One more pattern worth remembering: hard OH⁻ tends to precipitate with hard, highly charged transition metal ions (Fe(OH)₃ is the familiar example), while soft S²⁻ precipitates with softer, lower-charge metal ions, CuS, PbS, and ZnS being the usual ones you'd see in qualitative analysis.
Putting numbers on hardness
Pearson's absolute hardness gives the qualitative idea a formula:
where I is ionisation energy, A is electron affinity, η is hardness, χ is absolute electronegativity, and σ is softness. You can also write EHOMO = −I and ELUMO = −A. Among the halogens, F₂ comes out hardest and I₂ softest, since both χ and η fall steadily from F₂ to I₂ as the HOMO energy rises.
| Species | I (eV) | A (eV) | χ | η (hardness) |
|---|---|---|---|---|
| F⁻ | 17.42 | 3.40 | 10.41 | 7.01 |
| Cl⁻ | 13.01 | 3.62 | 8.31 | 4.70 |
| Br⁻ | 11.84 | 3.36 | 7.60 | 4.24 |
| I⁻ | 10.45 | 3.06 | 6.76 | 3.70 |
| Al³⁺ | 119.99 | 28.45 | 74.22 | 45.77 |
| H₂O | 12.6 | −6.4 | 3.1 | 9.5 |
| NH₃ | 10.7 | −5.6 | 2.6 | 8.2 |
Drago's E and C parameters
For a more quantitative prediction, Drago's equation splits the enthalpy of adduct formation into an ionic part and a covalent part:
I₂ serves as the reference, with Eₐ = Cₐ = 1.00. Plug in real numbers for I₂ reacting with benzene and you get −ΔH = (1.00 × 0.525) + (1.00 × 0.681) = 1.206 kcal/mol, which lines up well with the measured −1.3 kcal/mol (about 9% off). The same equation gives −4.21 kcal/mol for I₂ with diethyl ether (measured: −4.2) and −7.74 kcal/mol for I₂ with diethyl sulfide (measured: −7.8). That last pair is a nice illustration of HSAB in action: the softer sulfur donor binds noticeably more strongly to soft I₂ than the harder oxygen donor does.
Quick revision
- Arrhenius: water only, built around H⁺/OH⁻.
- Brønsted–Lowry: proton transfer, works in any solvent, conjugate pairs, levelling effect in water.
- Lewis: electron-pair transfer, the most general model, HOMO–LUMO framework, covers metal ions too.
- Superacids: H₀ more negative than H₂SO₄, measured by the Hammett function. Magic Acid is the strongest one you'll see commonly tested.
- Gas-phase acidity: no solvent involved, the truest measure of acid/base strength. Higher PA means a stronger base.
- Binary hydrides: down a group, bond strength controls acidity; across a period, electronegativity controls it.
- Oxyacids: more oxygens means a stronger acid, thanks to induction plus resonance stabilisation of the conjugate base.
- Boron halide Lewis acidity: BF₃ < BCl₃ < BBr₃ < BI₃, because π back-bonding fades as the halogen gets bigger.
- Sterics: F-strain blocks the approach, B-strain pushes groups apart after bonding. Quinuclidine forms cleaner adducts than Et₃N because it's rigid.
- HSAB: hard–hard and soft–soft pairings are the stable ones. Hard means small, non-polarisable, highly charged; soft means large, polarisable, low charge.
- Hardness η = (I − A)/2, and a higher η means harder. F₂ is the hardest halogen, I₂ the softest.
- Drago's equation: ΔH = −(EₐE_b + CₐC_b), with E as the ionic contribution and C as the covalent one.
- FLPs: a bulky Lewis acid and base that can't form an adduct instead activate H₂, CO₂, or N₂O — the first non-metal system to split H₂.
Worked MCQs
Ten questions worth trying before you read the explanation.
- Which acid–base model explains BF₃ + F⁻ → BF₄⁻?
(A) Arrhenius only (B) Brønsted–Lowry only (C) Lewis only (D) Lux–Flood only
Answer: C. Boron is accepting an electron pair from fluoride. No proton changes hands, and it's not happening in water, so Lewis is the only model that fits. - In glacial acetic acid, what's the order of acid strength?
(A) HCl > HClO₄ > H₂SO₄ > HNO₃ (B) HClO₄ > HCl > H₂SO₄ > HNO₃
(C) HNO₃ > H₂SO₄ > HCl > HClO₄ (D) All equal due to levelling
Answer: B. Acetic acid is too weak a base to level strong acids the way water does, so their true relative strengths come through. - NH₃ has a proton affinity of 853.6 kJ/mol; NMe₃ comes in around 948.9 kJ/mol. Which is the stronger base in water?
(A) NH₃ (B) NMe₃ (C) Both equal (D) Depends on temperature
Answer: A. NMe₃ wins in the gas phase, but in water NH₃ comes out ahead because NH₄⁺ can hydrogen-bond with more water molecules than NMe₃H⁺, which has no N–H bonds left to donate. - Which boron halide is the strongest Lewis acid?
(A) BF₃ (B) BCl₃ (C) BBr₃ (D) BI₃
Answer: D. As the halogen gets larger, the B–X bond lengthens, π back-donation weakens, and Lewis acidity climbs: BF₃ < BCl₃ < BBr₃ < BI₃. - Which property describes a soft Lewis acid?
(A) High charge, small size (B) Low polarisability (C) Large size, high polarisability (D) Large HOMO–LUMO gap
Answer: C. Soft means large, easily polarised, low charge density, and a narrow HOMO–LUMO gap. - I₂'s π* → σ* transition shows a blue shift in a better donor solvent. What does that tell you?
(A) Weaker Lewis base (B) The adduct's LUMO moves higher in energy (C) Less CT character (D) Longer I–I bond
Answer: B. The donor orbital pushes the adduct's LUMO up, which raises the transition energy and shows up as a blue shift. More blue shift means a stronger Lewis base. - How is absolute hardness, η, defined?
(A) (I + A)/2 (B) (I − A)/2 (C) I × A (D) I/A
Answer: B. η = (I − A)/2, half the gap between ionisation energy and electron affinity. - [Hg(SCN)₄]²⁻ versus [Zn(NCS)₄]²⁻: which statement is right?
(A) Both bind through N (B) Both bind through S (C) Hg binds S, Zn binds N (D) Hg binds N, Zn binds S
Answer: C. Hg²⁺ is soft and binds the soft end (S); Zn²⁺ is hard and binds the hard end (N). - What is Magic Acid made of?
(A) HF + AsF₅ (B) HSO₃F + SbF₅ (C) H₂SO₄ + SO₃ (D) HClO₄ + BF₃
Answer: B. Fluorosulfonic acid plus antimony pentafluoride, with H₀ between −21 and −25. - In Drago's equation, −ΔH = EₐE_b + CₐC_b, what does C measure?
(A) Electrostatic (ionic) interaction capacity (B) Covalent bond-forming tendency (C) Charge density (D) Polarisability
Answer: B. C is the covalent parameter; E is the electrostatic (ionic) one.
Practice question bank (25 questions)
All drawn from this chapter, answers below.
- Which acid–base model requires water as the solvent?
- HNO₃, HCl, H₂SO₄, and HClO₄ all appear equally strong in dilute water. This is due to the ______ effect.
- In liquid NH₃, the "neutral" species (analogous to H₂O in water) is ______.
- The conjugate base of H₂SO₄ in sulfuric acid solvent is ______.
- Gas-phase basicity (GB) of tri-n-butylamine is 967.6 kJ/mol, and its proton affinity is 998.5 kJ/mol. Why is PA always greater than GB?
- The Hammett acidity function H₀ of 100% H₂SO₄ is ______.
- Which binary hydride is the strongest acid: H₂S, H₂Se, or H₂Te?
- The oxyacid strength order for chlorine acids is HOCl ___ HClO₂ ___ HClO₃ ___ HClO₄. Fill in < or >.
- Why does Fe³⁺ solution turn yellow-brown? Write the relevant reaction.
- The driving force for BF₃·NH₃ adduct formation is the ______ of electrons in the donor's HOMO.
- Which orbital of BF₃ acts as the LUMO in its Lewis acid–base reactions?
- When I₂ dissolves in methanol, its colour changes from violet to ______, because of ______ interaction.
- What kind of bond forms between a Lewis acid and a Lewis base in an adduct?
- In FLP chemistry, what property of the reacting acid and base stops them forming a classic adduct?
- Write the net ionic equation for the acid–base reaction between NH₄Cl and NaNH₂ in liquid NH₃.
- A species with a low HOMO–LUMO gap is classified as ______ (hard/soft).
- Which is harder, F⁻ or I⁻? Justify using the η formula.
- Arrange in order of Lewis acidity: BMe₃, BF₃, BCl₃, BBr₃.
- In [CH₃Hg(H₂O)]⁺ + HI → CH₃HgI + H₃O⁺, K = 10¹⁸. Why is K so large? Use HSAB.
- What is absolute electronegativity χ in terms of I and A?
- The "proton sponge" has PA = 1028 kJ/mol. Give two reasons for its unusually high basicity.
- In [Co(NH₃)₅(SCN)]²⁺, does SCN⁻ bind through S or N? What kind of HSAB interaction is that?
- For I₂ + C₆H₆, Drago's parameters give Eₐ = 1.00, Cₐ = 1.00, E_b = 0.525, C_b = 0.681. Calculate |ΔH|.
- LiF has a Ksp of 1.8 × 10⁻³ while LiI is highly soluble. Explain using HSAB.
- Why is pyridine a stronger Lewis base than aniline toward most Lewis acids, even though aniline has an NH₂ group?
Answers
- Arrhenius model
- Levelling
- NH₃ itself (the amphoteric solvent)
- HSO₄⁻
- PA = −ΔH° and GB = −ΔG°, with GB = PA − TΔS°. Since ΔS° is negative (B loses freedom on gaining H⁺), TΔS° is negative too, which makes PA always larger than GB.
- −11.9
- H₂Te — the largest atom, weakest H–Te bond, and lowest charge density on the Te⁻ conjugate base.
- HOCl < HClO₂ < HClO₃ < HClO₄
- [Fe(H₂O)₆]³⁺ + H₂O → [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺. Fe³⁺ pulls electron density from the O–H bond, weakening it; the resulting hydrolysis gives the coloured iron-hydroxo species.
- Stabilisation (a drop in energy)
- Boron's 2pz orbital (the LUMO)
- Yellow-brown, from a donor–acceptor interaction with I₂'s LUMO
- A coordinate covalent (dative) bond
- Excessive steric bulk
- NH₄⁺ + NH₂⁻ → 2NH₃
- Soft
- F⁻ (η = 7.01 eV) is harder than I⁻ (η = 3.70 eV) — a larger η means harder.
- BMe₃ < BF₃ < BCl₃ < BBr₃ (covalent factor dominating; for exam purposes the commonly tested order is BBr₃ > BCl₃ > BF₃ > BMe₃, once π back-bonding in BF₃ is factored in)
- Hg²⁺ is soft and I⁻ is soft, so the soft–soft match gives a very favourable, very large K.
- χ = (I + A)/2
- Relief of steric repulsion between the two adjacent NMe₂ groups once protonated, plus a strong intramolecular N–H···N hydrogen bond that forms in the protonated state.
- Co³⁺ is hard, and SCN⁻ binds through N, the hard end — a hard–hard interaction.
- |ΔH| = (1.00 × 0.525) + (1.00 × 0.681) = 1.206 kcal/mol, about 5.05 kJ/mol.
- Li⁺ is hard and F⁻ is hard, so the strong hard–hard interaction in LiF gives low solubility. Li⁺ and I⁻ are mismatched (soft), giving a weaker interaction and much higher solubility.
- In aniline, the NH₂ lone pair delocalises into the aromatic ring through resonance, making nitrogen a weaker base. In pyridine, the lone pair sits in an sp² orbital orthogonal to the ring and stays fully available for donation.
Exam strategy
- NEET: focus on Arrhenius versus Brønsted–Lowry, conjugate acid–base pairs, oxyacid strength, and HSAB classification of common ions.
- JEE/BITSAT: Lewis acid–base chemistry, frontier orbitals, the boron halide Lewis acidity order, steric effects, and the gas-phase versus aqueous basicity flip for amines.
- IIT-JAM/CSIR-NET/GATE: Drago's E and C equation, Pearson's absolute hardness, FLPs, the spectroscopic evidence for hydrogen bonding and Lewis adducts, superacid chemistry, and the Hammett H₀ scale.
- TGT/PGT: the historical development of acid–base models, the levelling effect, amphoteric solvents, and HSAB predictions for solubility and coordination mode.
