Acids, Bases, Buffers and pH
Acids, Bases, Buffers and pH
Lesduur is: 50 minOnderdelen in deze les
Acids, Bases, Buffers and pH
Open by linking the topic to familiar lab experiences such as universal indicator, titrations and pH values, then signal that this lesson moves from definitions to equilibrium reasoning and buffer behaviour.
Learning objective
I can explain acids, bases, buffers and pH using Brønsted-Lowry theory and equilibrium ideas, and apply this to strong and weak acids, titrations and buffer systems.
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What words do you link with acids, bases and pH?
Use this to activate prior knowledge from GCSE and practical work. Likely responses include pH, indicator, neutralisation, H+, alkali, strong, weak and titration.
Lesson overview
This KS5 chemistry lesson explores Acids, Bases, Buffers and pH. Students apply Brønsted-Lowry theory and equilibrium calculations to strong and weak acids, titrations and buffer systems. It is suitable for A level classroom teaching, independent study or exam revision.
This slide is included verbatim as requested. Keep it separate near the beginning of the lesson.
Acids and bases in particles
At A level, acids and bases are best defined by proton transfer.
Acid: proton donor
Base: proton acceptor
A reaction between them is an acid–base reaction
This model explains more reactions than the simple idea of acids only being substances with pH below 7
Stress particle-level reasoning: what matters is H+ transfer between species, not just whether something is labelled acidic or alkaline in everyday language.
Brønsted–Lowry theory
In the Brønsted–Lowry model:
an acid loses H+
a base gains H+
proton transfer always creates related pairs
the equation shows which species change by one proton
Use a simple example such as HCl + H2O → H3O+ + Cl− or NH3 + H2O ⇌ NH4+ + OH− to model how to track proton movement.
Conjugate acid–base pairs
A conjugate acid–base pair differs by one proton.
HCl / Cl−
NH4+ / NH3
H2O / OH−
H3O+ / H2O
When an acid donates H+, it forms its conjugate base. When a base accepts H+, it forms its conjugate acid.
A common misconception is that conjugate pairs must appear on opposite sides only; the key rule is that the two species differ by exactly one proton.
Amphoteric species
Some species can act as both an acid and a base.
Water can accept H+ to form H3O+
Water can donate H+ to form OH−
Hydrogencarbonate, HCO3−, is another common example
Whether a species behaves as an acid or base depends on what it reacts with
Link amphoteric behaviour to why water is central in many acid–base equilibria.
Writing acid–base equations
When writing equations:
identify the proton donor and proton acceptor
move one H+ from acid to base
check charge balance as well as atoms
use ⇌ for reversible weak acid or weak base systems
use state symbols where appropriate in full equations
You can contrast a complete ionic equation with a simpler molecular equation if useful, but keep the focus on proton transfer.
Which statement matches Brønsted–Lowry theory?
Acid gains H+
Base donates H+
Acid donates H+
Base loses OH−
Checks the particle model definition. Distractors target confusion with Arrhenius ideas and reversed proton transfer.
Complete the summary
In Brønsted–Lowry theory, an acid is a proton donor and a base is a proton acceptor. Conjugate acid–base pairs differ by one proton. Water is amphoteric because it can act as either an acid or a base.
Useful retrieval of the key vocabulary from slides 3 to 7.
What is your first check when writing an acid–base equation?
Expected elements: identify donor and acceptor; move one H+ only; check atoms and charge balance; use reversible arrow for weak systems where appropriate.
Strong and weak acids
Strong and weak describe the extent of ionisation, not the concentration.
Strong acid: nearly complete ionisation in water
Weak acid: partial ionisation in water
A concentrated weak acid can still have a lower pH than a dilute strong acid only depending on amount present
Strength is about equilibrium position
Students often confuse strength with concentration. Make the contrast explicit: strength = degree of ionisation; concentration = amount per unit volume.
Ionisation and equilibrium
For a weak acid, only some molecules donate a proton:
HA(aq) ⇌ H+(aq) + A−(aq)
both reactants and products are present
the equilibrium lies mainly to the left
changing concentration can shift the position
If preferred, write H3O+ instead of H+ and keep notation consistent across your delivery. The slide uses standard A level shorthand.
The acid dissociation constant, Ka
For a weak acid:
Ka = [H+][A−] / [HA]
larger Ka means greater dissociation
stronger weak acids have larger Ka values
pKa = −log Ka
smaller pKa means a stronger acid
Make clear that Ka is used for weak acids because strong acids ionise almost completely, so an equilibrium constant is not usually useful in the same way.
Calculating pH
A level pH work usually follows these patterns:
strong monoprotic acid: find [H+] directly from concentration
strong alkali: find [OH−], then use pOH or Kw
weak acid: use Ka with equilibrium reasoning
always finish with pH = −log[H+]
Depending on exam board, students may use approximation methods for weak acids. Mention that assumptions must be justified by the size of dissociation.
What makes an acid weak?
Low concentration
Partial ionisation
High pH only
Few hydrogen atoms
Targets the common misconception that weak means dilute.
Complete the pH and Ka summary
A strong monoprotic acid gives [H+] directly from its concentration. For a weak acid, Ka describes the equilibrium. A larger Ka means greater dissociation, and pH is found at the end using −log[H+].
This checks the calculation route rather than a full numerical calculation.
The pH scale
pH measures hydrogen ion concentration on a logarithmic scale.
pH = −log[H+]
each 1 pH unit change means a tenfold change in [H+]
lower pH means higher [H+]
at 25 °C, neutral water has pH 7 because [H+] = [OH−]
Remind students that neutral does not always mean pH 7 unless temperature is specified as 25 °C.
Water and Kw
Water self-ionises slightly:
H2O(l) ⇌ H+(aq) + OH−(aq)
Kw = [H+][OH−]
at 25 °C, Kw = 1.0 × 10−14 mol2 dm−6
if you know one ion concentration, you can find the other
Some boards write H3O+ explicitly; if so, adapt consistently. The key idea is the fixed relationship between [H+] and [OH−] at a given temperature.
Comparing acidic and alkaline solutions
At 25 °C:
acidic: [H+] > [OH−]
neutral: [H+] = [OH−]
alkaline: [OH−] > [H+]
pH below 7 is acidic
pH above 7 is alkaline
Keep the particle-level comparison central so students do not reduce everything to memorising the number 7.
Indicators and pH measurement
Methods of measuring acidity differ in precision.
Indicator paper gives a rough pH range
Universal indicator shows approximate pH by colour
A pH meter gives a more precise reading
In quantitative work, a pH meter is preferred
You can link this to titration curves and why endpoint detection may depend on the indicator chosen.
At 25 °C, if [H+] = 1.0 × 10−9 mol dm−3, the solution is...
acidic
neutral
alkaline
buffered
Students must link pH 9 or compare [H+] with [OH−] using Kw, rather than relying only on labels.
Complete the measurement summary
The pH scale is logarithmic, so a change of one pH unit means a tenfold change in [H+]. At 25 °C, Kw equals 1.0 × 10−14, and for quantitative work a pH meter gives the most precise reading.
Echoes slides 12 to 15 and reinforces precision in measurement.
Titration basics
A titration measures how much acid reacts with a base, or vice versa.
the reaction must have a known stoichiometry
one solution is added gradually from a burette
the equivalence point is when the reacting amounts are exact
pH changes most rapidly near this point
Differentiate equivalence point from endpoint: equivalence is theoretical exact reaction, endpoint is the observed indicator change.
Titration curves
A titration curve shows pH against volume added.
strong acid–strong base curves have a very steep vertical section
weak acid–strong base curves start at a higher pH
the curve shape depends on acid or base strength
the graph helps identify a suitable indicator
If useful, sketch how the buffer region appears before equivalence in a weak acid–strong base titration.
Choosing an indicator
An indicator works when its colour change range falls within the steep pH change of the titration.
strong acid–strong base: several indicators may work
weak acid–strong base: an indicator with a higher transition range is often needed
a poor choice gives an inaccurate endpoint
Phenolphthalein and methyl orange are useful examples in UK A level courses, but keep the principle general unless you want board-specific detail.
Which titration usually starts at the higher pH?
Strong acid–strong base
Weak acid–strong base
Strong acid–weak base
Any titration curve
Checks understanding of curve shape from the subtopic, not just vocabulary.
Which indicator would you expect to suit a weak acid–strong base titration best?
Higher range
Lower range
Any indicator
No indicator
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Why must the indicator match the steep pH change?
Expected elements: indicator transition range should lie within the steep vertical part of the curve; a poor match gives an inaccurate endpoint; higher transition range often suits weak acid–strong base titrations.
Buffers: what they do
A buffer solution resists large changes in pH when small amounts of acid or alkali are added.
it does not keep pH perfectly constant
it works best over a limited pH range
many buffer solutions contain a weak acid and its conjugate base
others contain a weak base and its conjugate acid
Students may think buffers stop pH changing completely; emphasise that they reduce the change, rather than prevent it entirely.
How an acidic buffer works
An acidic buffer commonly contains a weak acid, HA, and its conjugate base, A−.
Added H+ → removed by A−
Added OH− → removed by HA
Because both species are present in significant amounts, the equilibrium shifts to oppose the change.
A good example is ethanoic acid with sodium ethanoate. Tie the explanation to Le Châtelier’s principle without drifting away from acid–base equilibrium.
Buffer calculations
For an acidic buffer:
Ka = [H+][A−] / [HA]
Rearranging gives:
[H+] = Ka × [HA] / [A−]
This lets you estimate pH from the ratio of weak acid to conjugate base.
If students know the Henderson–Hasselbalch equation from your course, you may mention it, but this slide stays aligned to the equilibrium expression itself.
Why buffers matter
Buffers are important wherever pH must stay within narrow limits.
biological systems rely on stable pH
many industrial and analytical processes depend on controlled pH
buffer action also explains part of the shape of some titration curves
Keep examples chemistry-centred. You may briefly mention blood buffering, but avoid turning the slide into a biology detour.
Putting the ideas together
This topic connects through equilibrium and proton transfer.
Brønsted–Lowry theory explains who donates or accepts H+
strength depends on extent of ionisation
pH depends on [H+]
titration curves show how pH changes during neutralisation
buffers work because equilibria oppose small changes
Use this as a synthesis slide before moving to glossary terms or later interactive elements in another pass.
What is an acidic buffer made from?
Strong acid and salt
Weak acid and base
Weak acid and conjugate base
Strong base and water
Directly checks the buffer composition taught on slides 19 to 21.
Complete the buffer explanation
A buffer resists large changes in pH when small amounts of acid or alkali are added. In an acidic buffer, added H+ is removed by A−, while added OH− is removed by HA. This works because the equilibrium shifts to oppose the change.
A concise application of buffer action using the exact species from the content deck.
If [A−] falls in a buffer, what happens to pH?
Expected elements: use [H+] = Ka × [HA] / [A−]; if [A−] decreases or [HA] increases, [H+] rises and pH falls; if [A−] increases relative to [HA], [H+] falls and pH rises. This serves as the applied task and allows differentiation through the depth of explanation.
What is one idea you understand better now?
Reflection prompt with no right answer. Invite students to name one idea, method or misconception they clarified.
Which part of this lesson felt most challenging?
Reflection prompt with no right answer. Useful for planning reteach or revision focus.
What helped you learn this topic best today?
Reflection prompt with no right answer. Encourage students to mention a strategy, example or representation that helped them learn.
Glossary
acid - a proton donor in the Brønsted–Lowry model
base - a proton acceptor in the Brønsted–Lowry model
conjugate acid–base pair - two species that differ by one proton
amphoteric - able to act as either an acid or a base
strong acid - an acid that ionises almost completely in water
weak acid - an acid that ionises only partially in water
ionisation - formation of ions when a substance dissociates in solution
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Glossary
Ka - the acid dissociation constant for a weak acid equilibrium
pKa - the negative logarithm of Ka
pH - the negative logarithm of hydrogen ion concentration
Kw - the ionic product of water, equal to [H+][OH−]
equivalence point - the point in a titration where reactants are present in exact reacting amounts
indicator - a substance that changes colour over a particular pH range
buffer solution - a solution that resists large changes in pH when small amounts of acid or alkali are added
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