Equilibrium · Lesson 7 of 14
Ionic Equilibrium and Models of Acids and Bases
“Connect ions in water with three useful ways of recognising acid–base reactions.”
• Distinguish strong, weak and non-electrolytes from the particles present in solution. • Explain hydration, dissociation and molecular ionization. • Identify acids, bases and conjugate pairs using proton transfer. • Recognise electron-pair donation and acceptance in the Lewis model.
Ions allow a solution to conduct
Dissolving sugar and dissolving common salt can both give clear solutions, but the particles present are different. Sugar remains mainly as neutral molecules. Sodium chloride supplies mobile sodium and chloride ions. In an electric field, oppositely charged ions move in opposite directions and carry current. An electrolyte produces mobile ions in its liquid or dissolved state; a nonelectrolyte does not supply an appreciable concentration of such ions in the solution considered.
A strong electrolyte forms ions almost completely in water at suitable dilute concentrations. A weak electrolyte forms only some ions and retains a substantial molecular population. Weak electrolytes therefore have a reversible ionization equilibrium. These terms describe the extent of ion formation, not how much solute was added. A dilute strong acid and a concentrated weak acid are both possible.
| Solute in water | Main particles | Classification |
|---|---|---|
| Sugar | Neutral sugar molecules | Nonelectrolyte |
| Sodium chloride | Hydrated Na+ and Cl− | Strong electrolyte |
| Hydrochloric acid | Hydronium and chloride ions | Strong acid / strong electrolyte |
| Acetic acid | Mostly CH3COOH with some hydronium and acetate | Weak acid / weak electrolyte |
In solid sodium chloride the ions already exist in an ionic lattice. Water separates and hydrates them: this is dissociation. In hydrogen chloride or acetic acid, a neutral molecule reacts with water to generate ions: this is ionization. The chapter sometimes uses the two words broadly, but the particle distinction remains useful. Water’s polarity and high relative permittivity reduce electrostatic attraction between separated ions, and hydration stabilises them. Whether dissolution is favourable also depends on the complete free-energy balance.
Faraday’s studies connected solution conductivity with ions and electrolysis. Arrhenius later developed a theory of electrolytic dissociation. Their work helped connect observable electrical behaviour with the microscopic species in a solution.
Arrhenius acids and bases in water
An Arrhenius acid increases the concentration of hydrogen ions in water, while an Arrhenius base increases hydroxide-ion concentration. Hydrogen ions do not remain as bare protons in water. They bind to water and form hydrated proton structures. H3O+, the hydronium ion, is the simplest useful representation; H+(aq) is a shorthand for this hydrated proton.
The direct-dissociation picture is useful for hydroxides such as NaOH, but ammonia contains no OH group and still produces hydroxide by reacting with water. A proton-transfer model explains this explicitly. The water-specific Arrhenius description is also less general than models that apply in other solvents or to reactions without proton transfer. Familiar sourness, litmus changes, hydrogen release with some metals, and neutralisation are observations associated with many acids and bases, not instructions to identify chemicals by tasting them.
The Brønsted–Lowry model follows a proton
A Brønsted–Lowry acid donates a proton and a Brønsted–Lowry base accepts one. In , ammonia accepts a proton and water donates it. In the reverse direction ammonium donates a proton to hydroxide. Each side therefore contains an acid and a base.
Two species differing by one proton. Removing H+ from an acid produces its conjugate base; adding H+ to a base produces its conjugate acid.
Adding one proton raises the charge by one; removing one lowers it by one. A conjugate pair need not consist of neutral molecules. Water accepts a proton from HCl but donates one to ammonia, so its acid or base role depends on its reaction partner. A species that can both donate and accept a proton is amphiprotic.
Problem
Give the conjugate bases of HF, H2SO4 and HCO3−.
- 1.Remove exactly one H+ from each species; do not remove every hydrogen at once.
- 2.HF gives F−. H2SO4 gives HSO4−.
- 3.HCO3− gives . Each product has one lower charge than its parent acid.
Problem
Give the conjugate acids of NH2−, NH3 and HCOO−.
- 1.Add one H+ to each species and raise its charge by one.
- 2.NH2− gives NH3, while NH3 gives NH4+.
- 3.HCOO− gives HCOOH. The acid–base pair differs by a proton, not by an electron.
| Amphiprotic species | Conjugate acid | Conjugate base |
|---|---|---|
| H2O | H3O+ | OH− |
| HCO3− | H2CO3 | |
| HSO4− | H2SO4 | |
| NH3 | NH4+ | NH2− |
In a proton-transfer equilibrium, the stronger acid tends to react with the stronger base to form the weaker acid and weaker base. The stronger an acid is, the weaker its conjugate base. This is a relative relationship: the conjugate base of an ordinary weak acid is not automatically a very strong base in water. For example, acetate is a weak base, while amide, NH2−, accepts protons from water extremely readily.
The Lewis model follows an electron pair
A Lewis acid accepts an electron pair and a Lewis base donates one to form a bond. Ammonia has an available lone pair on nitrogen. Electron-deficient BF3 can accept that pair to form an adduct. No proton has to be transferred, so the Lewis model includes reactions outside the Brønsted–Lowry description.
Problem
Classify OH−, F−, H+ and BCl3 by their electron-pair behaviour.
- 1.OH− and F− have available lone pairs and can donate a pair: they are Lewis bases.
- 2.H+ can accept a pair to make a bond and is a Lewis acid.
- 3.Electron-deficient boron in BCl3 can accept a pair, so BCl3 is a Lewis acid. The classification follows the participating site and reaction, not simply the sign of the overall charge.
An acid and its conjugate base differ by one proton, not one electron. Also, describing NH4+ as a Lewis acid in a proton-transfer reaction refers to electron-pair acceptance at a transferable proton; it does not require nitrogen to expand its octet.
Quiz
Which solution mainly contains neutral solute molecules?
In NH3 + H2O forming NH4+ + OH−, water acts as
The conjugate base of HSO4− is
In the BF3–NH3 adduct, NH3 is a Lewis base because it
A weak acid is necessarily
Which statement about H+(aq) is correct?
Practice Problems
- Compare the particles in aqueous NaCl, sugar and acetic acid.
- Explain the difference between dissociation of an ionic solid and ionization of a neutral acid.
- Identify both conjugate pairs in HCl + H2O forming H3O+ + Cl−.
- Give the conjugate acid and base of HCO3−, including their charges.
- Explain BF3 + NH3 using electron-pair transfer and explain NH3 + H2O using proton transfer.
- Distinguish a strong acid from a concentrated acid.
Key Takeaways
Electrolyte behaviour depends on mobile ions, while strength describes the extent of ion formation. Brønsted–Lowry reactions transfer protons and create conjugate pairs. Lewis reactions involve electron-pair donation and acceptance; the appropriate model depends on the reaction.