Acids & Bases: Buffers
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What are buffer solutions
Maintain an approximately constant pH when small amounts of acid, base or water are added
When are buffer solutions utilised
Standardisation of pH meters Biological systems eg. fermentation Maintain pH in shampoos & cosmetics
Properties of an acidic buffer
pH below 7 Weak acid Strong conjugate base salt
Properties of a basic buffer
pH above 7 Weak base Strong conjugate acid salt
How do buffers maintain pH?
A proton accepter that can react with added acid: Weak base = Basic buffer Strong conjugate base salt = Acidic buffer A proton donor that can react with added base: Weak acid = Acidic buffer Strong conjugate acid salt = Basic buffer
Strong salt definition
Fully dissociates into it’s ions with a high concentration
In an acidic buffer containing CH3COOH: What is the equilibrium reaction? Reaction if HCl was added? Reaction if NaOH was added? How is equilibrium maintained?
CH3COOH ⇌ CH3COO- + H+ H+ (from HCl) + CH3COO- → CH3COOH OH- (from NaOH) + H+ → H2O Equilibrium will shift to maintain concentration of a weak acid dissociation reactions
Explain a buffer solution section on a graph
Only weak acid to start with so salt and water forms. Swiftly becomes more alkali Buffer solution containing unreacted weak acid and salt of weak acid Reactions cause the pH to become more gradual All acid and salt run out - equivalence point and rapid pH occurs
In a basic buffer containing NH3: What is the equilibrium reaction? Reaction if HCl was added? Reaction if NaOH was added? How is equilibrium maintained?
NH3 + H2O ⇌ NH4+ + OH- H+ (from HCl) + OH- → H2O OH- (from NaOH) + NH4+ → NH3+ + OH- Equilibrium will shift to maintain concentration of a weak basic dissociation reactions
Why cannot the equation [H+]^2 / [HA] be used in weak acid buffers
Not a pure weak acid so are not approximate concentrations to eachother
Ka solutions for weak acid buffers: rearranged for [H+] (pH) Implication of this
[H+] = Ka x ( [HA]/[A-] ) Ratio of [HA] : [A-] = [H+]
Why does adding water not cause an effect
Diluting a buffer affects both [HA-] and [A-] equally so ratio and pH are unchanged
Reaction process to find pH when a weak acid reacts with a strong base: CH3COOH + NaOH → CH3COONa + H2O Ka given
Moles NaOH used = Moles A- made ICE table (H2O not relevant) [H] = Ka x (moles of HA / moles of A-) -log(H+)
Moles of HA and A- in a buffer (with 3 moles initially each) after X moles of NaOH added OH- + HA → H2O + A-
HA: 3 - X A-: 3 + X
Moles of HA and A- in a buffer (with 3 moles initially each) after X moles of HCl added H+ + A- → HA
HA: 3 + X A-: 3 - X
Ethanedioic acid is a weak acid. Ethanedioic acid acts, initially, as a monoprotic acid. COOHCOOH ⇌ COOHCOO- + H+ Use the concept of electronegativity to justify why the acid strengths of ethanedioic acid and ethanoic acid are different. [6]
Stage 1: Difference in structure The acids are of the form RCOOH but in ethanoic acid R = CH3 whilst in ethanedioic acid R = COOH Stage 2: Inductive effect The unionised COOH group contains two very electronegative oxygen atoms therefore has a negative inductive (electron withdrawing) effect The CH3 group has a positive inductive (electron pushing) effect Stage 3: Polarity of OH The O–H bond in the ethanedioic acid is more polarised More dissociation into H+ ions Ethanedioic acid is stronger than ethanoic acid
Two solutions, one with a pH of 4.00 and the other with a pH of 9.00, were left open to the air. The pH of the pH 9.00 solution changed more than that of the other solution. Suggest what substance might be present in the air to cause the pH to change. Explain how and why the pH of the pH 9.00 solution changes. [3]
[1] CO2 [2] pH falls / decreases [3] reacts with alkali
HCO3- ⇌ H+ + CO32- Explain how a solution containing sodium hydrogencarbonate and sodium carbonate can act as a buffer when small amounts of acid or small amounts of alkali are added. [3]
[1] Acid: Increase in concentration of H+ ions, equilibrium moves to the left [2] Alkali: OH- reacts with H+ ions, equilibrium moves to the right [3] Concentration of H+ remains (almost) constant