M6 Acids and Bases Flashcards
Indicators
Phenolphthalein - white at 8, pink at 10.
Bromothymol blue - yellow at 6, blue at 8.
Methyl Orange - red at 3, yellow-orange at 5.
Methyl red - red at 4, yellow at 7.
Antoine Lavoisier (1743-94)
Said an acid contained oxygen
Humphrey Davy (1778-1829)
series of experiments in which he demonstrated HCl, H2S and H2Te were acids, showing that oxygen could not be responsible for acidity. Proposed that acids contained hydrogen that could be displaced by a metal.
Svante Arrhenius (1859 - 1927)
proposed that acids ionised in solution to produce H+ and that bases ionised in solution to produce OH-.
Bronsted-Lowry Theory
An acid is a substance that donates one or more protons or hydrogen ions. A base is a substance that accepts one or more protons.
Arrhenius Theory Limitations
ammonia dilemma - could not explain why ammonia was a base, nature and role of the solvent was not considered, not all bases are soluble, the H+ ion does not exist as such in aqueous solutions, it bonds to H2O to produce hydronium.
Bronsted-Lowry Theory Limitations
requires the present of a solvent which has a hydrogen attached to an oxygen or nitrogen. It does not explain the acid-base behaviour of substances in non-aqueous solvents where a proton is not involved. It cannot explain reactions between acidic and basic oxides, which take place even where there is no solvent. There are also substances such as BF3 and AlCl3 that act as acids but do not have a hydrogen present, so they cannot donate a proton.
Gilbert Lewis Theory
an acid is an electron pair acceptor, a base is an electron pair donor. Does not require a solvent. Explains why BF3 reacts with ammonia. All Bronsted-Lowry acids and bases are Lewis acids and bases, but the reverse is not true.
Amphoteric
substance capable of acting as both a base and an acid. e.g. water, hydrogen sulfate ion.
Conjugate pair theory
every acid has a conjugate base, a substance that has one proton less than the acid. If the acid is strong, then the base is weak and vice versa.
Monoprotic
one proton or hydrogen ion available to donate. diprotic, triprotic, polyprotic…
Burette
Large thin glass apparatus used to deliver solution slowly in titrations.
Half-equivalence point
where the amount of titrant is half the amount of the chemical in our initial solution.
Henderson-Hasselbalch Equation
where a weak acid is titrated with a strong base (or vice versa), we use a buffer solution. pH = pKa + log10([A-] / [HA]) [A-] is concentration of acid. [HA] is concentration of base. At the half-equivalence point, [A-] = [HA], therefore pH = pKa
Primary Standard
A solution of known concentration that is the starting point of a titration. Must be: available in a highly pure form, have large molar mass, stable in air, not absorb moisture or carbon dioxide from atmosphere, readily soluble in distilled water, reacts readily with the solution of unknown concentration. E.g. anhydrous sodium carbonate Na2CO3
Secondary standard
a solution whose concentration is determined by titration against a primary standard.
End/Equivalence point
Point where the exact stoichiometric ratios of moles of reactants exist.
Titration curves
y-axis: pH of the analyte solution. X-axis: volume of titrant added
Conductivity graphs
Strong acids/bases have higher conductivity (greater dissociation, ions). Y-axis: conductance. X-axis: mL of strong acid/base.
Back Titration
method where the concentration of an analyte is determined by reacting it with a known amount of excess reagent. The remaining excess reagent is then titrated with another, second reagent. Used when: it is difficult to determine a definite end point because the reaction occurs too slowly, the sample is not soluble in water but will react with an acid/base, the sample is toxic or volatile or in gas form or fairly unreactive or low concentration.
Buffers
maintain a pH within a certain range despite the addition of an acid or base up to a point. Occur in conjugate acid-base pairs. Acidic buffer: weak acid, strong conjugate base. Basic buffer: weak base, strong conjugate acid. Uses equilibrium reactions, Le Chatelier’s principle. E.g. bicarbonate buffer system: maintains blood pH between 7.35 and 7.45, ocean acidification carbonic acid buffer system.
Optimal Buffer Conditions
work best when the acid and base have the same concentrations. Solution is buffered if pH is maintained with 1 unit of the optimal pH.
Buffer use
Standards for precision instruments such as pH probes. Household products such as soaps, cosmetics, washing powder, fizzy drinks, paints, beer and wine. Industry such as fermentation, dyes and pharmaceuticals.
