Studied by 0 people
Define a Bronsted-Lowry acid.
Substance that can donate a proton (hydrogen ion).
Define a Bronsted-Lowry base.
Substance that can accept a proton (hydrogen ion).
Define conjugate acid.
Species formed after base has accepted a proton.
Define conjugate base.
Species remaining after acid has lost a proton.
Define a Lewis acid.
Substance which can ACCEPT a PAIR OF ELECTRONS.
Define a Lewis base.
Substance which can DONATE a PAIR OF ELECTRONS.
What are the relative strengths of the conjugates of strong and weak acids.
Strong acids→Weak conjugate base Weak acids→Strong conjugate base
Deduce whether or not a species could act as a Brønsted-Lowry and/or a Lewis acid or base.
•All bronsted-lowry acids are lewis acids. •Not all lewis acids are bronsted lowry acids as they do not need to release H⁺.
Deduce the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid).
HA + B ⇌ A⁻ + BH⁺ A⁻ conjugate base pair of HA BH⁺ conjugate acid pair of B
Describe Lewis acid base reactions.
Result in the formation of a covalent bond which is will always be a dative bond as both electrons come from the base. There is NO transfer of H⁺.
Describe alkalis.
Bases that dissolve in water. They all release hydroxide ion OH⁻ when dissolved in water. e.g. NH₃ +H₂O ⇌ NH₄⁺ + OH⁻
Strength of an acid.
Measure of how readily acid dissociates in aqueous solution.
What are monoprotic acids?
1 mole of acid produces 1 mole of H⁺ions.
What are diprotic acids?
1 mole of acid produces 2 moles of H⁺ions.
Describe the pH scale.
pH = -log₁₀[H⁺] [H⁺]=10⁻(pH) Each change of one pH unit represents a 10-fold change in the hydrogen ion concentration.
Concentration of strong acids and bases.
Assuming full dissociation occurs, [H⁺] or [OH⁻] will be equal to the initial concentration of the acid or base respectively.
State the equation for the reaction of weak acid with water, and hence deduce the expressions for Ka.
Use equilibrium expressions as they do not fully dissociate. For weak acid HA: HA(aq)+H₂O(l)⇌H₃O⁺(aq)+A⁻(aq) Kc=[H₃O⁺][A⁻]/[HA][H₂O] Given [H₂O] is constant, combine with Kc to create Ka, acid dissociation constant. Ka=[H₃O⁺][A⁻]/[HA]
State the equation for the reaction of weak base with water, and hence deduce the expressions for Kb.
Use equilibrium expressions as they do not fully dissociate. For weak base B: B(aq)+H₂O(l)⇌BH⁺(aq)+OH⁻(aq) Kc=[BH⁺][OH⁻]/[B][H₂O] Given [H₂O] is constant, combine with Kc to create Kb, base dissociation constant. Kb=[BH⁺][OH⁻]/[B]
Identify the relative strengths of acids and bases using values of Ka, Kb, pKa and pKb.
• Strong acid Ka increase↑ ↓ pKa increases Weak acid • Strong base Kb increase↑ ↓ pKb increases Weak Base •The larger the pKw the weaker the acid. •Larger the pKw, the weaker the base.
Describe a buffer solution.
•Aqueous solution resistant to changes in pH when small amounts of acid or alkali are added. •Blood is a natural buffer. •Pure water is NOT a good buffer as its pH fluctuates significantly with addition of acid or alkali.
Define buffer capacity.
Amount a buffer can be diluted without a change in pH when acid or base is absorbed.
Describe the composition of acidic buffers.
Mixing equal amounts of moles of WEAK ACID with the ITS SALT OF A STRONG BASE (often a sodium salt) in aqueous solution.
Response of acid buffers to added acid and base.
ADDITION OF ACID (H⁺): CH₃COO⁻+H⁺⇌CH₃COOH H⁺ combines with the base CH₃COO⁻ to form CH₃COOH, removing added H⁺. ADDITION OF BASE (OH⁻): CH₃COOH+OH⁻⇌CH₃COO⁻+H₂O OH⁻ combines with acid CH₃COOH to form CH₃COO⁻ and H₂O, removing added OH⁻.
Describe the composition of basic buffers.
Mixing equal amounts of moles of WEAK BASE with a solution of ITS SALT OF A STRONG ACID.
Response of basic buffers to added acid and base.
ADDITION OF ACID (H⁺): NH₃+H⁺⇌NH₄⁺ H⁺ combines with the base NH₃ to form NH₄⁺, removing added H⁺. ADDITION OF BASE (OH⁻): NH₄⁺+OH⁻⇌NH₃+H₂O OH⁻ combines with acid NH₄⁺ to form NH₃ and H₂O, removing added OH⁻.
Amounts of acid and base needed to create acid and basic buffers.
ACIDIC BUFFER:>50% conjugate base salt solution than weak acid (need large reservoir) BASIC BUFFER: >50% conjugate acid salt solution than weak base (need large reservoir)
What happens when buffer solution contains equal amounts in moles of acid and salt (or base and salt)?
When [acid]=[salt], pH=pKa When [base]=[salt], pOH=pKb
Affect of temperature on a buffer.
Temperature affects Ka and Kb so temperature affects pH
Describe salt hydrolysis.
•Many salts dissolve in water to produce solutions which are not neutral. •pH depends on whether and to what extent ions react with water and hydrolyse it releasing H⁺ and OH⁻ ions.
Deduce whether salts form acidic, alkaline or neutral aqueous solutions.
•STRONG ACID+STRONG BASE: neither ion hydrolyses, neutral salt. •WEAK ACID+STRONG BASE: anion hydrolyses, basic salt. •STRONG ACID+WEAK BASE: cation hydrolysed, acidic salt. •WEAK ACID+WEAK BASE: anion and cation hydrolysed, type of salt depends on relative strengths of conjugates.
Describe acid base titrations.
Neutralization reactions are carried out by titration sequential addition of one reactant from a burette to a fixed volume of the other reactant. Usually alkali 0.1mol is added to 25cm³ of acid 0.1mol
Explain equivalence point.
[Acid]=[Alkali]
Describe titration of strong acid and strong base.
e.g. HCl+NaOH→NaCl+H₂O •Initial pH 1 (strong acid) •After equivalence curve flattens out at pH 13 •pH equivalence point = 7
Describe titration of weak acid and strong base.
e.g. CH₃COOH+NaOH→CH₃COO⁻Na⁺+H₂O •Initial pH 3 (weak acid) •After equivalence curve flattens out at pH 13 •pH equivalence point >7 •at half equivalence pH=pKa
Describe titration of strong acid and weak base.
e.g. HCl+NH₃→NH₄Cl •Initial pH 1 (strong acid) •pH relatively constant through buffer region until equivalence •jump in pH at equivalence from pH 3 to pH 7 •pH equivalence point < 7
Describe titration of weak acid and weak base.
e.g. CH₃COOH+NH₃→CH₃COO⁻NH₄⁺ •Initial pH fairly high 3 (weak acid) •addition of base causes pH to rise steadily •change in pH at equivalence point much less sharp than other titrations •After equivalence curve flattens out at fairly low pH •pH equivalence point difficult to identify
Describe indicators.
Weak acids or weak bases whose undissociated and dissociated forms have different colours. Indicators are substances that change colour reversibly according to pH of solution.
Describe qualitatively the action of an acid-base indicator.
For indicator HIn that is a weak acid: HIn ⇌ H⁺ + In⁻ Colour A Colour B Increase in [H⁺]: equilibrium shift to left in favour of HIn (colour A) Decrease in [H⁺]: equilibrium shift to right in favour of In⁻ (colour B) Increase [OH⁻]: OH⁻ react with H⁺ to make H₂O, position of equilibrium shifts to products.
What is the end point in an acid base titration?
pH at which an acid alkali indicator changes colour. Choose an indicator with end point in range of equivalence point.
State and explain how the pH range of an acid-base indicator relates to its pKa value.
Ka=[H⁺][In⁻]/[HIn] At equilibrium, half equivalence point [HIn]=[In⁻] ∴at equilibrium Ka=[H⁺] or pKa=pH
Note
Flashcard