What is the acid dissociation constant , K a ?

An order of reaction is the power to which the concentration of a reactant is raised in the experimentally determined rate equation. It reflects how the rate changes with concentration of that reactant.

True or False? A weak acid with a higher K a value is more dissociated and therefore stronger than one with a lower K a .

True. The K a value reflects the extent of dissociation. A larger K a means the equilibrium lies further to the right, so a greater proportion of acid molecules have dissociated.

Given the K a expression for a weak acid HA: K a = [H + ][A - ] / [HA] State the two assumptions made when simplifying this expression.

The two simplifying assumptions are: The concentration of H + from water ionisation is negligible The equilibrium concentration of HA approximates its initial concentration, [HA] eq ≈ [HA] initial , because the degree of dissociation is small

What formula is used to calculate [H + ] for a weak acid from its K a and concentration?

[H + ] = √( K a x [HA]) This follows from [H + ] = [A - ] when one H + and one A - form per molecule of HA that dissociates.

Calculate the pH of 0.100 mol dm -3 ethanoic acid where K a = 1.74 x 10 -5 mol dm -3 . [H + ] = √( .......... x .......... ) = .......... mol dm -3 pH = ..........

[H + ] = √( 1.74 x 10 -5 x 0.100 ) = 1.32 x 10 -3 mol dm -3 pH = 2.88

True or False? In a 0.10 mol dm -3 solution of ethanoic acid, the equilibrium concentration of CH 3 COOH is approximately equal to its initial concentration.

True. Ethanoic acid is a weak acid and only partially dissociates. The change in [CH 3 COOH] at equilibrium is negligible compared to 0.10 mol dm -3 , so [HA] eq ≈ [HA] initial .

Why does the K a expression for a weak acid not include the concentration of water?

Water is the solvent and its concentration is so large it remains effectively constant throughout the reaction. It is therefore incorporated into the equilibrium constant and omitted from the expression.

What is p K a and how is it calculated?

p K a is the negative logarithm (base 10) of the acid dissociation constant: p K a = -log 10 K a . Most weak acids have p K a values between 3 and 7.

Ethanoic acid has K a = 1.74 x 10 -5 mol dm -3 . Complete the calculation of p K a . p K a = -log 10 ( .......... ) = ..........

p K a = -log 10 ( 1.74 x 10 -5 ) = 4.76

True or False? A weaker acid has a larger p K a value.

True. A weaker acid has a smaller K a . Taking the negative log of a smaller number gives a larger p K a . So the stronger the acid, the lower the p K a .

At the half-neutralisation point of a weak acid titration, what is the relationship between p K a and pH?

At the half-neutralisation point, [HA] = [A - ], so the K a expression simplifies to K a = [H + ]. Therefore p K* a = pH at this point.

Describe how to determine K a from a pH titration curve. Plot .......... against volume of base added Identify the volume at .......... neutralisation Read off the .......... at this volume At this point, p K a = ..........

Plot pH against volume of base added Identify the volume at half neutralisation Read off the pH at this volume At this point, p K a = pH

Why is p K a more convenient to use than K a for comparing weak acids?

K a values for weak acids are very small numbers (e.g. 10 -5 ) spread over a wide range. The p K a scale compresses these into manageable values typically between 3 and 7, making comparison easier.

True or False? p K w = pH + pOH = 14.00 at 298 K.

True. Since K w = [H + ][OH - ] = 1.00 x 10 -14 , taking -log of both sides gives p K w = pH + pOH = 14.00 at 298 K.

Acid dissociation constant ( K a )

The acid dissociation constant ( K a ) is the equilibrium constant for the partial ionisation of a weak acid: HA ⇌ H + + A − , giving K a = [H + ][A − ] / [HA]. A larger K a indicates a stronger weak acid .

What is the equivalence point (stoichiometric point) in a titration?

Equivalence point is the point at which the two solutions have completely reacted in their stoichiometric ratio. It is detected by an indicator changing colour (the end point) or by a sharp pH change on a pH curve.

In a titration, 25.00 cm 3 of 0.0500 mol dm -3 HCl was neutralised by 8.50 cm 3 of NaOH. Complete the calculation for [NaOH]. moles HCl = .......... mol moles NaOH = .......... mol [NaOH] = .......... mol dm -3

moles HCl = 1.25 x 10 -3 mol moles NaOH = 1.25 x 10 -3 mol [NaOH] = 0.147 mol dm -3

True or False? A burette is read to ±0.05 cm 3 because its smallest markings are 0.10 cm 3 and the uncertainty is half the smallest division.

True. Burettes are analogue instruments marked to 0.10 cm 3 . The uncertainty is half the smallest division, giving ±0.05 cm 3 .

What is the difference between the end point and the equivalence point in a titration?

The equivalence point is where the acid and base have exactly reacted. The end point is where the indicator changes colour. These may not coincide exactly, which is why choosing the right indicator matters.

50.0 cm 3 of 0.10 mol dm -3 NaOH is added to 25.0 cm 3 of 0.15 mol dm -3 HCl. After adding 45 cm 3 of NaOH: moles acid = .......... mol; moles NaOH added = .......... mol; NaOH in excess = .......... mol [OH - ] = .......... mol dm -3 ; pH = ..........

moles acid = 3.75 x 10 -3 mol; moles NaOH = 4.50 x 10 -3 mol; NaOH excess = 7.5 x 10 -4 mol [OH - ] = 0.0107 mol dm -3 ; pH = 12.03

True or False? Titration calculations involving diprotic acids such as H 2 SO 4 use a 1:2 acid-to-base mole ratio.

True. H 2 SO 4 has two acidic protons, so it reacts with NaOH in a 1:2 ratio. pH-curve questions at A Level are limited to monoprotic acids, but concentration calculations can involve diprotic acids.

What are concordant titration results?

Concordant results are titration readings that are within 0.10 cm 3 of each other. Multiple trials are carried out until concordant results are obtained before an average titre is calculated.

The titre is the volume of solution delivered from a burette during a titration, recorded in cm 3 . Multiple titrations are carried out until concordant results (within 0.10 cm 3 ) are obtained, and an average titre is calculated.

What is a pH curve in the context of a titration?

A conjugate acid–base pair is two species that differ by a single proton (H + ). The acid donates H + to form its conjugate base; the base accepts H + to form its conjugate acid.

An indicator is suitable if its pH range falls .......... the steep/vertical section of the pH curve. For a weak acid-strong base titration, the pH changes from .......... to .......... at the equivalence point, so .......... is the best indicator.

An indicator is suitable if its pH range falls entirely within the steep/vertical section of the pH curve. For a weak acid-strong base titration, the pH changes from 7 to 10 at the equivalence point, so phenolphthalein is the best indicator.

True or False? No indicator is suitable for a weak acid-weak base titration.

True. In a weak acid-weak base titration there is no sharp pH change at the equivalence point. The pH changes too gradually for any indicator to give a distinct colour change at a specific volume.

How does the starting pH of a pH curve reveal whether the acid used is strong or weak?

A strong acid starts at a low pH (e.g. approximately 1 for 0.1 mol dm -3 ). A weak acid of the same concentration starts at a higher pH (e.g. approximately 3) because it only partially dissociates.

What is the pKa of an indicator and how does it relate to its end point?

The p K a of an indicator equals the pH at its end point, because at the end point [HIn] = [In - ], which simplifies K a = [H + ]. So p K a = pH at the colour-change mid-point.

Titration Suitable indicator Strong acid + strong base .......... Weak acid + strong base .......... Strong acid + weak base .......... Weak acid + weak base ..........

Titration Suitable indicator Strong acid + strong base methyl red or phenolphthalein Weak acid + strong base phenolphthalein Strong acid + weak base methyl red (or methyl orange) Weak acid + weak base none suitable

True or False? A strong acid-weak base pH curve has its equivalence point below pH 7.

True. At the equivalence point the salt formed (e.g. ammonium chloride) undergoes hydrolysis to give a slightly acidic solution, so the pH at equivalence is below 7.

What is a buffer solution ?

A strong acid is an acid that fully dissociates in aqueous solution, releasing all its protons to form hydrogen ions and the corresponding anion.

When OH - ions are added to an ethanoic acid / sodium ethanoate buffer: OH - reacts with .......... to form .......... [H + ] .......... Equilibrium shifts .......... and more .......... dissociates pH remains ..........

OH - reacts with H + to form water [H + ] decreases Equilibrium shifts right and more CH 3 COOH dissociates pH remains approximately constant

True or False? Diluting a buffer solution changes its pH because the ratio of [salt] to [acid] changes.

False. On dilution both the acid and salt concentrations decrease by the same factor, so the [salt]/[acid] ratio is unchanged. The Henderson-Hasselbalch equation shows that pH therefore stays the same.

Write the Henderson-Hasselbalch equation for calculating buffer pH.

pH = p K a + log 10 ([salt]/[acid]) This is derived from K a = [H + ][A - ]/[HA], rearranged to find [H + ] and then converted to the logarithmic form.

How do you calculate the pH of a buffer solution given K a and the concentrations of the weak acid and its conjugate base?

Substitute into [H + ] = K a × ([acid] / [salt]), then calculate pH = −log[H + ]. Example: K a = 1.74 × 10 -5 mol dm -3 , [acid] = 0.305 mol dm -3 , [salt] = 0.520 mol dm -3 → [H + ] = 1.02 × 10 -5 mol dm -3 , pH = 4.99.

How does the carbonate buffer system in blood prevent acidosis?

CO 2 from respiration dissolves in blood to form H + and HCO 3 - . If [H + ] rises, the equilibrium shifts left, converting H + and HCO 3 - back to CO 2 and H 2 O, keeping blood pH between 7.35 and 7.45.

True or False? A basic buffer can be made from ammonia and ammonium chloride, where NH 3 neutralises added acid and NH 4 + neutralises added alkali.

True. Added H + reacts with NH 3 to form NH 4 + ; added OH - reacts with NH 4 + to reform NH 3 and water. Both reactions prevent significant pH change.

A LevelChemistryAQAFlashcardsAdvanced Physical ChemistryFurther Acids & Bases Calculations

Further Acids & Bases Calculations (AQA A Level Chemistry): Flashcards

Exam code: 7405

1/37

0Still learning

Know0

FrontAcid Dissociation Constant
What is the acid dissociation constant, K_a?
FrontAcid Dissociation Constant
Given the K_a expression for a weak acid HA:
K_a = [H⁺][A^-] / [HA]
State the two assumptions made when simplifying this expression.
BackAcid Dissociation Constant
The two simplifying assumptions are:
The concentration of H⁺ from water ionisation is negligible
The equilibrium concentration of HA approximates its initial concentration, [HA]_eq ≈ [HA]_initial, because the degree of dissociation is small

Was this flashcard helpful?

Cards in this collection (37)

What is the acid dissociation constant, K_a?
An order of reaction is the power to which the concentration of a reactant is raised in the experimentally determined rate equation. It reflects how the rate changes with concentration of that reactant.
True or False?
A weak acid with a higher K_a value is more dissociated and therefore stronger than one with a lower K_a.
True.
The K_a value reflects the extent of dissociation. A larger K_a means the equilibrium lies further to the right, so a greater proportion of acid molecules have dissociated.
Given the K_a expression for a weak acid HA:
K_a = [H⁺][A^-] / [HA]
State the two assumptions made when simplifying this expression.
The two simplifying assumptions are:
1. The concentration of H⁺ from water ionisation is negligible
2. The equilibrium concentration of HA approximates its initial concentration, [HA]_eq ≈ [HA]_initial, because the degree of dissociation is small
What formula is used to calculate [H⁺] for a weak acid from its K_a and concentration?
[H⁺] = √(K_a x [HA])
This follows from [H⁺] = [A^-] when one H⁺ and one A^- form per molecule of HA that dissociates.
Calculate the pH of 0.100 mol dm^-3 ethanoic acid where K_a = 1.74 x 10^-5 mol dm^-3.
[H⁺] = √(.......... x .......... ) = .......... mol dm^-3
pH = ..........
[H⁺] = √(1.74 x 10^-5 x 0.100) = 1.32 x 10^-3 mol dm^-3
pH = 2.88
True or False?
In a 0.10 mol dm^-3 solution of ethanoic acid, the equilibrium concentration of CH₃COOH is approximately equal to its initial concentration.
True.
Ethanoic acid is a weak acid and only partially dissociates. The change in [CH₃COOH] at equilibrium is negligible compared to 0.10 mol dm^-3, so [HA]_eq ≈ [HA]_initial.
Why does the K_a expression for a weak acid not include the concentration of water?
Water is the solvent and its concentration is so large it remains effectively constant throughout the reaction. It is therefore incorporated into the equilibrium constant and omitted from the expression.
What is pK_a and how is it calculated?
pK_a is the negative logarithm (base 10) of the acid dissociation constant: pK_a = -log₁₀ K_a. Most weak acids have pK_a values between 3 and 7.
Ethanoic acid has K_a = 1.74 x 10^-5 mol dm^-3. Complete the calculation of pK_a.
pK_a = -log₁₀(.......... ) = ..........
pK_a = -log₁₀(1.74 x 10^-5) = 4.76
True or False?
A weaker acid has a larger pK_a value.
True.
A weaker acid has a smaller K_a. Taking the negative log of a smaller number gives a larger pK_a. So the stronger the acid, the lower the pK_a.
At the half-neutralisation point of a weak acid titration, what is the relationship between pK_a and pH?
At the half-neutralisation point, [HA] = [A^-], so the K_a expression simplifies to K_a = [H⁺]. Therefore pK*_a = pH at this point.
Describe how to determine K_a from a pH titration curve.
1. Plot .......... against volume of base added
2. Identify the volume at .......... neutralisation
3. Read off the .......... at this volume
4. At this point, pK_a = ..........
1. Plot pH against volume of base added
2. Identify the volume at half neutralisation
3. Read off the pH at this volume
4. At this point, pK_a = pH
Why is pK_a more convenient to use than K_a for comparing weak acids?
K_a values for weak acids are very small numbers (e.g. 10^-5) spread over a wide range. The pK_a scale compresses these into manageable values typically between 3 and 7, making comparison easier.
True or False?
pK_w = pH + pOH = 14.00 at 298 K.
True.
Since K_w = [H⁺][OH^-] = 1.00 x 10^-14, taking -log of both sides gives pK_w = pH + pOH = 14.00 at 298 K.
Acid dissociation constant (K_a)
The acid dissociation constant (K_a) is the equilibrium constant for the partial ionisation of a weak acid: HA ⇌ H⁺ + A⁻, giving K_a = [H⁺][A⁻] / [HA]. A larger K_a indicates a stronger weak acid.
What is the equivalence point (stoichiometric point) in a titration?
Equivalence point is the point at which the two solutions have completely reacted in their stoichiometric ratio. It is detected by an indicator changing colour (the end point) or by a sharp pH change on a pH curve.
In a titration, 25.00 cm³ of 0.0500 mol dm^-3 HCl was neutralised by 8.50 cm³ of NaOH. Complete the calculation for [NaOH].
moles HCl = .......... mol
moles NaOH = .......... mol
[NaOH] = .......... mol dm^-3
moles HCl = 1.25 x 10^-3 mol
moles NaOH = 1.25 x 10^-3 mol
[NaOH] = 0.147 mol dm^-3
True or False?
A burette is read to ±0.05 cm³ because its smallest markings are 0.10 cm³ and the uncertainty is half the smallest division.
True.
Burettes are analogue instruments marked to 0.10 cm³. The uncertainty is half the smallest division, giving ±0.05 cm³.
What is the difference between the end point and the equivalence point in a titration?
The equivalence point is where the acid and base have exactly reacted. The end point is where the indicator changes colour. These may not coincide exactly, which is why choosing the right indicator matters.
50.0 cm³ of 0.10 mol dm^-3 NaOH is added to 25.0 cm³ of 0.15 mol dm^-3 HCl. After adding 45 cm³ of NaOH:
moles acid = .......... mol; moles NaOH added = .......... mol; NaOH in excess = .......... mol
[OH^-] = .......... mol dm^-3; pH = ..........
moles acid = 3.75 x 10^-3 mol; moles NaOH = 4.50 x 10^-3 mol; NaOH excess = 7.5 x 10^-4 mol
[OH^-] = 0.0107 mol dm^-3; pH = 12.03
True or False?
Titration calculations involving diprotic acids such as H₂SO₄ use a 1:2 acid-to-base mole ratio.
True.
H₂SO₄ has two acidic protons, so it reacts with NaOH in a 1:2 ratio. pH-curve questions at A Level are limited to monoprotic acids, but concentration calculations can involve diprotic acids.
What are concordant titration results?
Concordant results are titration readings that are within 0.10 cm³ of each other. Multiple trials are carried out until concordant results are obtained before an average titre is calculated.
Titre
The titre is the volume of solution delivered from a burette during a titration, recorded in cm³. Multiple titrations are carried out until concordant results (within 0.10 cm³) are obtained, and an average titre is calculated.
What is a pH curve in the context of a titration?
A conjugate acid–base pair is two species that differ by a single proton (H⁺). The acid donates H⁺ to form its conjugate base; the base accepts H⁺ to form its conjugate acid.
An indicator is suitable if its pH range falls .......... the steep/vertical section of the pH curve.
For a weak acid-strong base titration, the pH changes from .......... to .......... at the equivalence point, so .......... is the best indicator.
An indicator is suitable if its pH range falls entirely within the steep/vertical section of the pH curve.
For a weak acid-strong base titration, the pH changes from 7 to 10 at the equivalence point, so phenolphthalein is the best indicator.
True or False?
No indicator is suitable for a weak acid-weak base titration.
True.
In a weak acid-weak base titration there is no sharp pH change at the equivalence point. The pH changes too gradually for any indicator to give a distinct colour change at a specific volume.
How does the starting pH of a pH curve reveal whether the acid used is strong or weak?
A strong acid starts at a low pH (e.g. approximately 1 for 0.1 mol dm^-3). A weak acid of the same concentration starts at a higher pH (e.g. approximately 3) because it only partially dissociates.
What is the pKa of an indicator and how does it relate to its end point?
The pK_a of an indicator equals the pH at its end point, because at the end point [HIn] = [In^-], which simplifies K_a = [H⁺]. So pK_a = pH at the colour-change mid-point.

Titration	Suitable indicator
Strong acid + strong base	..........
Weak acid + strong base	..........
Strong acid + weak base	..........
Weak acid + weak base	..........

Titration	Suitable indicator
Strong acid + strong base	methyl red or phenolphthalein
Weak acid + strong base	phenolphthalein
Strong acid + weak base	methyl red (or methyl orange)
Weak acid + weak base	none suitable

True or False?
A strong acid-weak base pH curve has its equivalence point below pH 7.
True.
At the equivalence point the salt formed (e.g. ammonium chloride) undergoes hydrolysis to give a slightly acidic solution, so the pH at equivalence is below 7.
What is a buffer solution?
A strong acid is an acid that fully dissociates in aqueous solution, releasing all its protons to form hydrogen ions and the corresponding anion.
When OH^- ions are added to an ethanoic acid / sodium ethanoate buffer:
1. OH^- reacts with .......... to form ..........
2. [H⁺] ..........
3. Equilibrium shifts .......... and more .......... dissociates
4. pH remains ..........
1. OH^- reacts with H⁺ to form water
2. [H⁺] decreases
3. Equilibrium shifts right and more CH₃COOH dissociates
4. pH remains approximately constant
True or False?
Diluting a buffer solution changes its pH because the ratio of [salt] to [acid] changes.
False.
On dilution both the acid and salt concentrations decrease by the same factor, so the [salt]/[acid] ratio is unchanged. The Henderson-Hasselbalch equation shows that pH therefore stays the same.
Write the Henderson-Hasselbalch equation for calculating buffer pH.
pH = pK_a + log₁₀([salt]/[acid])
This is derived from K_a = [H⁺][A^-]/[HA], rearranged to find [H⁺] and then converted to the logarithmic form.
How do you calculate the pH of a buffer solution given K_a and the concentrations of the weak acid and its conjugate base?
Substitute into [H⁺] = K_a × ([acid] / [salt]), then calculate pH = −log[H⁺].
Example: K_a = 1.74 × 10^-5 mol dm^-3, [acid] = 0.305 mol dm^-3, [salt] = 0.520 mol dm^-3 → [H⁺] = 1.02 × 10^-5 mol dm^-3, pH = 4.99.
How does the carbonate buffer system in blood prevent acidosis?
CO₂ from respiration dissolves in blood to form H⁺ and HCO₃^-. If [H⁺] rises, the equilibrium shifts left, converting H⁺ and HCO₃^- back to CO₂ and H₂O, keeping blood pH between 7.35 and 7.45.
True or False?
A basic buffer can be made from ammonia and ammonium chloride, where NH₃ neutralises added acid and NH₄⁺ neutralises added alkali.
True.
Added H⁺ reacts with NH₃ to form NH₄⁺; added OH^- reacts with NH₄⁺ to reform NH₃ and water. Both reactions prevent significant pH change.

Sign up to unlock flashcards

By signing up you agree to our Terms and Privacy Policy

Previous:Fundamentals of Acids & BasesNext:Properties of Period 3 Elements & their Oxides