HSC Chemistry Module 6 Acids and Bases Topic Summary

A bottle is labelled “0.10 mol L\(^{-1}\) acid”. Can you tell its pH from that information alone?

Not necessarily. If the acid is strong, almost every acid particle releases a proton into water. If it is weak, only a fraction does. The concentration on the bottle tells you how much acid was added, but acid-base chemistry asks a deeper question: what actually exists in the solution after the particles interact with water?

That idea connects almost everything in HSC Chemistry Module 6.

This summary covers the core Module 6 map: acid and base models, characteristic reactions, Brønsted-Lowry theory, conjugate pairs, strong and weak acids, pH calculations, \(K_a\) and \(pK_a\), titrations, titration curves, conductivity, and buffers. Experimental technique is included where it changes how you interpret results, but detailed practical-method evaluation belongs in your practical notes.

The equilibrium reasoning underneath weak acids and buffers comes from Module 5. If equilibrium expressions or Le Chatelier’s principle feel shaky, revise the HSC Chemistry Module 5 Equilibrium summary alongside this page.

01What actually makes something an acid or a base?

Suppose you dissolve hydrogen chloride, \(\ce{HCl}\), in water and ammonia, \(\ce{NH3}\), in another beaker.

Both solutions behave very differently. The hydrochloric acid solution contains a high concentration of hydronium ions. The ammonia solution is basic, even though an \(\ce{NH3}\) molecule contains no \(\ce{OH^-}\).

That second observation is the problem that makes acid-base models useful.

Arrhenius gives us a useful first model

An Arrhenius acid increases the concentration of \(\mathrm{H}^{+}\) in aqueous solution. An Arrhenius base increases the concentration of \(\ce{OH^-}\).

For example:

\[
\ce{HCl(aq) -> H+(aq) + Cl-(aq)}
\]

and

\[
\ce{NaOH(aq) -> Na+(aq) + OH-(aq)}
\]

This model works nicely for familiar acids and hydroxide bases.

Its limitation is that it focuses heavily on aqueous solutions and struggles to explain bases such as ammonia without adding extra steps.

Brønsted-Lowry follows the proton instead

The more useful HSC model is based on proton transfer.

A Brønsted-Lowry acid donates a proton, \(\mathrm{H}^{+}\).

A Brønsted-Lowry base accepts a proton.

Consider ammonia in water:

\[
\ce{NH3(aq) + H2O(l) <=> NH4+(aq) + OH-(aq)}
\]

Ammonia accepts a proton from water, so \(\ce{NH3}\) is acting as a base. Water donates that proton, so water is acting as an acid.

This is much more powerful than simply hunting for \(\mathrm{H}^{+}\) and \(\ce{OH^-}\) in formulas.

Quick check: In

\[
\ce{HCl + H2O -> H3O+ + Cl-}
\]

which species is the Brønsted-Lowry base?

02Acid reactions follow patterns you should recognise

Before calculating anything, you should be able to predict what an acid will do when it meets common reactants.

The point isn’t to memorise twenty equations. Learn the reaction pattern, then apply it.

Acid plus base

An acid and a base undergo neutralisation.

For a strong acid reacting with a strong hydroxide base, the essential ionic process is:

\[
\ce{H+(aq) + OH-(aq) -> H2O(l)}
\]

For example:

\[
\ce{HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)}
\]

Neutralisation is generally exothermic because forming water releases energy. In practical work, the enthalpy of neutralisation can be estimated using calorimetry.

Acid plus carbonate

The pattern is:

acid + carbonate \(\rightarrow\) salt + water + carbon dioxide

For example:

\[
\ce{2HCl(aq) + Na2CO3(aq) -> 2NaCl(aq) + H2O(l) + CO2(g)}
\]

The bubbling is carbon dioxide.

Acid plus reactive metal

The pattern is:

acid + reactive metal \(\rightarrow\) salt + hydrogen gas

For example:

\[
\ce{2HCl(aq) + Mg(s) -> MgCl2(aq) + H2(g)}
\]

The exact reaction depends on the metal and acid, so don’t turn the pattern into an absolute rule for every possible metal.

Indicators don’t cause acidity

An acid-base indicator is itself a weak acid-base system whose forms have different colours.

A simplified representation is:

\[
\ce{HIn(aq) <=> H+(aq) + In-(aq)}
\]

If \(\ce{HIn}\) and \(\mathrm{In}^{-}\) have different colours, changing the acidity shifts their relative amounts and changes the observed colour.

Think of the indicator as a chemical weather vane. It responds to the conditions. It doesn’t create them.

Common trap: “The indicator changes colour exactly at pH 7.”

Not generally. Different indicators change colour over different pH ranges. This becomes important when choosing an indicator for a titration.

03Conjugate pairs, water, and amphiprotic substances

Look again at:

\[
\ce{NH3 + H2O <=> NH4+ + OH-}
\]

Something useful has happened. Each reactant has turned into a closely related product.

\(\ce{NH3}\) accepts \(\mathrm{H}^{+}\) to become \(\ce{NH4+}\).

\(\ce{H2O}\) donates \(\mathrm{H}^{+}\) to become \(\ce{OH^-}\).

So the conjugate acid-base pairs are:

• \(\ce{NH4+ / NH3}\)
• \(\ce{H2O / OH^-}\)

A conjugate acid-base pair differs by exactly one proton.

That is the fastest reliable test.

Water can play either role

Water behaves a bit like the friend who changes personality depending on who they’re dating.

With a stronger proton donor, water accepts a proton:

\[
\ce{HCl + H2O -> H3O+ + Cl-}
\]

Here water is a base.

With ammonia, water donates a proton:

\[
\ce{NH3 + H2O <=> NH4+ + OH-}
\]

Here water is an acid.

The analogy breaks because water isn’t making emotional decisions. Proton transfer depends on the relative acid-base tendencies of the species involved.

A substance able to act as either a proton donor or proton acceptor is amphiprotic.

For example, hydrogen carbonate can accept a proton:

\[
\ce{HCO3- + H+ -> H2CO3}
\]

or donate one:

\[
\ce{HCO3- + OH- -> CO3^2- + H2O}
\]

Quick check: Is \(\ce{H2PO4^-}\) amphiprotic?

04Strong is not the same thing as concentrated

This is probably the most important language distinction in the topic.

Imagine two rooms.

Room A contains ten people, and every single person starts dancing.

Room B contains one thousand people, but only ten start dancing.

“How many people are in the room?” and “what fraction are dancing?” are different questions.

Acid concentration and acid strength work similarly.

Concentration tells you how much acid is present per unit volume.

Strength tells you the extent to which the acid ionises in water.

A strong monoprotic acid such as \(\ce{HCl}\) is treated as essentially completely ionised:

\[
\ce{HCl(aq) + H2O(l) -> H3O+(aq) + Cl-(aq)}
\]

A weak acid establishes an equilibrium:

\[
\ce{CH3COOH(aq) + H2O(l) <=> H3O+(aq) + CH3COO-(aq)}
\]

Most weak-acid particles remain un-ionised.

So a dilute strong acid can have a higher pH than a concentrated weak acid. Strength alone does not tell you the final pH.

\(K_a\) measures weak-acid ionisation

For a weak acid represented as:

\[
\ce{HA(aq) + H2O(l) <=> H3O+(aq) + A-(aq)}
\]

the acid dissociation constant is:

\[
K_a = \frac{[\ce{H3O+}][\mathrm{A}^{-}]}{[\ce{HA}]}
\]

A larger \(K_a\) means the equilibrium lies further towards ionised products, so the acid is stronger.

We also use:

\[
pK_a = -\log_{10}(K_a)
\]

Therefore:

• larger \(K_a\) = stronger acid
• smaller \(pK_a\) = stronger acid

That reversal catches people constantly.

If equilibrium expressions feel mysterious here, this is exactly where Module 5 equilibrium comes back into the course.

05pH calculations are really concentration calculations

pH compresses a very large range of hydronium concentrations into a manageable scale.

At HSC level, the key relationships are:

Concentrations are in mol L\(^{-1}\).

Because pH is logarithmic, a change of one pH unit corresponds to a tenfold change in \([\ce{H3O+}]\).

So pH 3 is not “a little more acidic” than pH 4. It has ten times the hydronium ion concentration.

Mixing acids and bases

When an acid and base are mixed, don’t average their pH values.

pH is logarithmic, and more importantly, the acid and base chemically react.

The reliable process is:

1. calculate moles of acid and base,
2. use the balanced equation to determine the limiting reagent,
3. find any excess \(\ce{H3O+}\) or \(\ce{OH^-}\),
4. divide excess moles by the new total volume,
5. calculate pH or pOH.

That mole-first method becomes essential in titrations.

06Titration turns stoichiometry into a measurement

Suppose you have an acid of unknown concentration.

You could measure its pH, but that doesn’t automatically tell you its analytical concentration, especially if the acid is weak.

Instead, react it with a solution whose concentration is accurately known.

That is the logic of an acid-base titration.

A measured volume of the unknown solution is placed in a conical flask. A standard solution is delivered gradually from a burette. The volume required to reach the reaction’s stoichiometric point tells you how many moles reacted.

Equivalence point versus endpoint

These terms are related but not identical.

The equivalence point is the theoretical point at which the reactants have been added in their exact stoichiometric ratio.

The endpoint is the experimentally observed point, usually indicated by a colour change.

A good indicator changes colour close to the equivalence point.

Choosing an indicator simply because “it changes near pH 7” is not a reliable rule. The equivalence-point pH depends on the acid and base involved.

What titration curves tell you

A titration curve plots pH against volume of titrant added.

The shape tells you more than the final titre.

For a strong acid-strong base titration, the pH changes very sharply around an equivalence point near pH 7.

For a weak acid-strong base titration, the initial pH is higher, a buffer region appears before equivalence, and the equivalence point is above pH 7.

For a strong acid-weak base titration, the equivalence point is below pH 7.

Why?

Because the salt left at equivalence may react with water. The equivalence point means stoichiometric neutralisation has occurred. It does not automatically mean the resulting solution is neutral.

Conductivity gives a second view of the same reaction

Conductivity depends on the concentration and mobility of ions.

During a titration, ions disappear, appear, or are replaced by ions with different mobilities. The conductivity graph therefore changes as the reaction proceeds.

This is useful because it lets you analyse a neutralisation without relying only on colour.

07Buffers are controlled resistance to pH change

Imagine adding a small amount of acid to pure water. The extra \(\ce{H3O+}\) remains available, so the pH can change substantially.

Now suppose the solution contains both a weak acid, \(\ce{HA}\), and its conjugate base, \(\mathrm{A}^{-}\).

If acid is added:

\[
\ce{A^- + H3O+ -> HA + H2O}
\]

The conjugate base removes much of the added hydronium.

If base is added:

\[
\ce{HA + OH^- -> A^- + H2O}
\]

The weak acid removes much of the added hydroxide.

That mixture is a buffer.

A buffer therefore resists changes in pH when small amounts of acid or base are added.

It does not lock the pH permanently.

Think of a buffer as two friends standing near a badly behaved mate at a party. One deals with trouble coming from one direction, the other deals with trouble coming from the other. This works until the amount of trouble exceeds what they can handle.

Chemically, that limit is the buffer capacity. Once too much acid or base has been added and one buffer component is substantially depleted, the pH changes much more rapidly.

Buffers matter in natural systems because many chemical and biological processes operate properly only over restricted pH ranges.

The key conceptual link is equilibrium. A buffer works because added acid or base is consumed by reactions involving a weak acid-conjugate base pair, shifting equilibria rather than allowing the added ions simply to accumulate.

08Common misconceptions worth fixing before an exam

“Strong acid means low pH.”

Not by itself. Strength describes ionisation. pH also depends on concentration.

“A weak acid barely reacts.”

Wrong idea. “Weak” means partial ionisation in water. A weak acid can still react substantially with an appropriate base.

“Equivalence point always has pH 7.”

Only in cases such as a strong acid-strong base titration under standard assumptions. Weak conjugate species can make the equivalence solution acidic or basic.

“A buffer stops pH changing.”

It resists change. Add enough acid or base, and the buffer will be overwhelmed.

“When acid is diluted, it becomes weaker.”

Its concentration decreases. Its intrinsic acid strength, represented by quantities such as \(K_a\), does not change just because water was added.

“\(K_a\) and concentration are the same kind of quantity.”

No. Concentration tells you how much acid is present. \(K_a\) describes the position of the acid-ionisation equilibrium at a given temperature.

“An indicator tells you the exact pH.”

Usually not. An indicator gives a colour over a transition range. A calibrated pH probe provides a numerical measurement.

The most useful way to revise Module 6 is as one connected chain:

proton transfer \(\rightarrow\) equilibrium \(\rightarrow\) acid strength \(\rightarrow\) ion concentration \(\rightarrow\) pH \(\rightarrow\) neutralisation \(\rightarrow\) titration and buffers.

If one calculation starts feeling like a random formula exercise, move backwards along that chain. Usually the missing step is not algebra. It is deciding which particles are actually present, which reaction occurs, and whether that reaction is effectively complete or an equilibrium.