Acid: An acid dissociates and releases H+ ions in aqueous solution. A Bronsted- Lowry acid is a
proton donor.
A strong acid completely dissociates AND a weak acid partially dissociates.
Can measure current through an acid to determine if strong and weak acid. Strong conducts
more electricity as more ionised.
Alkali: An alkali dissociates and releases OH- ions in aqueous solution. An alkali is a soluble base.
A Bronsted- Lowry base is a proton acceptor.
Neutralisation: This is neutralisation. H+ (aq) + OH- (aq) -> H2O (l)
Conjugate Acid- Base Pairs: A conjugate acid- base pair contains two species that can be
interconverted by gain or loss of proton. For example in this dissociation HCl (aq) -> H+ (aq) + Cl-
(aq). Combining dissociation and neutralisation ionic equation above get and acid- base
equilibria…
HCl is acid as donates H+ and OH- is base as
accepts H+ in forward reaction.
In reverse reaction H2O an acid as donates H+ and Cl- base as accepts H+.
Equilibrium Sign: Single arrow in dissociations indicates forward reaction complete if strong
acid. Equilibrium sign for weaker acids.
Hydronium Ion: Dissociation only takes place when water present as requires proton transfer
from acid to base. Water can act as base and acid. H2O below accepts a proton to form H3O+.
Hydronium ion present in all aqueous acids.
H+ in equation is actually the simplified version H3O+ so can interchange. H3O
+ + OH-
-> 2H2O.
When constructing acid- base equilibrium, instead of forming OH- or H+, form H3O+.
Two acids can react together. The stronger acid will act as a proton donor (acid) the weaker acid
will act as a proton acceptor (base).
Reactants were given. To find products take a H and then add a H. No water involved here in
reactant so none in product. Must have negative charge for base formed and positive charge
for acid formed. Unless one of them already has a charge, may form neutral or double charge,
always +/- 1.
Monobasic: Monobasic, dibasic and tribasic aka monoprotic refer to number of hydrogen ions in
acid that can be replaced per molecule in an acid- base reaction.
In organic compounds, do not replace H on carbon chain e.g. CH3COOH is monobasic. H3BO3 is
tribasic.
Neutralisation Equations: Decide whether acid is mono, di or tribasic. Then write equation using
as many Na in NaOH needed e.g. H2SO4 + 2NaOH -> Na2SO4 + 2H2O. Then balance as normal. Now
can this equation to work out the volume of NaOH requires to neutralise 25cm3
of acid.
Ionic Equation: Split all ionic compounds (solid/ aqueous) into ion. Leave solids, liquids and
gases. Cancel ions that do not change- spectator ions. Rewrite.
Redox Reaction: Acids- H2SO4, HCl, H3PO4 and HNO2. These form sulphate, chloride, phosphate
and nitrate salts.
Acid (aq) + metal (s) -> salt + hydrogen. Solid dissolves and fizzing.
Ionic equation- Mg + 2H+ -> Mg2+ + H2.
Neutralisation Reactions between Acids and Bases…
Acid + metal carbonate -> salt + water + carbon dioxide. Effervescence and if soluble carbonate,
the solid dissolves.
2H+ + CO3
2- -> H2O + CO2.
Acid + base (metal oxide) -> salt + water. Solid dissolves.
2H+ + O2- -> H2O.
Acid + alkali -> salt + water. No observations. Alkali can be NaOH or KOH.
Acid + ammonia -> Ammonium salt.
H2SO4 + 2NH3 -> (NH4)2SO4. No observations.
Be careful using sodium carbonate as formula is Na2CO3 so needs two moles of acid and salt.
Always check formula of alkali by working out.
If any kind of acid reacts with base (question may say they are acid/ base), the ionic equation will
always be H+ + OH- -> H2O. And in normal equation, ONLY two products are salt and water if
acid + alkali.