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Inquiry Question 1: What is an acid and a base?

Investigate the properties of acids and bases and the historical development of the Arrhenius model of acids and bases

A focused answer to the HSC Chemistry Module 6 dot point on the properties of acids and bases. Observed properties, indicator colours, the Arrhenius model, limitations of Arrhenius, and the historical development that led to Bronsted-Lowry.

Generated by Claude Opus 4.89 min answerUpdated 2026-06-22

Reviewed by: AI editorial process; not yet individually human-reviewed

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In plain English

Think of an acid and a base as two teams that only recognise each other by the "jersey" they hand out when they dive into a pool of water. Arrhenius noticed that every acid, no matter what it looks like on dry land (a bottle of vinegar, a chunk of citric acid, a tank of sulfuric acid), hands out the exact same jersey the moment it hits water: a hydrogen ion, $H^+$ . Every base hands out a different, matching jersey: a hydroxide ion, $OH^-$ . Because every acid wears the same jersey, they all behave the same way in the pool, sour taste, red litmus, fizzing with a carbonate, and every base behaves the same way too, bitter, slippery, blue litmus. The catch is that Arrhenius could only referee the game inside the pool (water). The moment you take the game outside, onto dry land (a non-aqueous solvent) or bring in a player who behaves like a base but never actually hands out an $OH^-$ jersey (ammonia), Arrhenius's rulebook has nothing to say. That gap is exactly why the Bronsted-Lowry rulebook, built on passing a single ball (a proton) rather than handing out jerseys, took over as the more complete set of rules.

What this dot point is asking

NESA wants you to describe the physical and chemical properties shared by acids and bases, recall the Arrhenius model and the reasoning behind it, and explain why later models (Bronsted-Lowry, Lewis) had to extend it. This is the entry point to Module 6 and the foundation for every later calculation, including reactions of acids, strong vs weak ionisation, and Bronsted-Lowry conjugate pairs.

The answer

Observed properties of acids

Taste sour (citric acid in lemons, ethanoic acid in vinegar). Never taste laboratory chemicals.
Turn blue litmus red.
React with active metals (Mg, Zn, Fe) to produce hydrogen gas.
React with metal carbonates and hydrogencarbonates to produce $CO_2$ .
React with bases to form a salt and water (neutralisation).
Aqueous solutions conduct electricity (they are electrolytes).
Have pH less than 7 at 25 degrees C.

Observed properties of bases

Taste bitter and feel soapy or slippery (do not test by taste or touch).
Turn red litmus blue.
React with acids to form a salt and water.
React with ammonium salts to release ammonia gas.
Aqueous solutions conduct electricity.
Have pH greater than 7 at 25 degrees C.

Indicators

An indicator is a weak acid or weak base whose protonated and deprotonated forms have different colours. The colour change occurs across a narrow pH range, usually about 2 pH units wide.

Indicator	Colour in acid	Colour in base	pH range
Methyl orange	red	yellow	3.1 to 4.4
Bromothymol blue	yellow	blue	6.0 to 7.6
Phenolphthalein	colourless	pink	8.3 to 10.0
Litmus	red	blue	4.7 to 8.3

Universal indicator is a mixture of indicators that gives a continuous colour scale from red (very acidic) to violet (very basic).

Must-know constants and the neutralisation equation

Neutralisation (ionic equation): $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}$ , true for every strong acid-strong base pair.
Ionic product of water at 25 degrees C: $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ . This lets you convert between $[H^+]$ and $[OH^-]$ in any aqueous solution.
pH scale definition: $pH = -\log_{10}[H^+]$ and $pOH = -\log_{10}[OH^-]$ , with $pH + pOH = 14.00$ at 25 degrees C.
Indicator colour range: an indicator changes colour over roughly 2 pH units centred near its own $pK_a$ ; choose one whose range overlaps the steep part of the titration curve, not necessarily pH 7.

An owned illustrative titration curve shows how pH changes as a base is added to a strong acid, and is the graph you should expect to read and calculate from in an exam:

The Arrhenius model

Svante Arrhenius (1887) proposed that acids and bases are substances that ionise in water.

Arrhenius acid: a substance that releases $H^+$ in aqueous solution.
Arrhenius base: a substance that releases $OH^-$ in aqueous solution.
Neutralisation: $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}$ .

The model elegantly explained why all aqueous acids share the same chemistry (because they all produce the same ion, $H^+$ ) and why all aqueous bases share the same chemistry (because they all produce $OH^-$ ).

Historical development

The Arrhenius model built on earlier ideas:

Lavoisier (1780s). Believed all acids contained oxygen (the name "oxygen" means "acid former"). Disproved when Humphry Davy showed that hydrochloric acid contains no oxygen.
Davy (1810). Proposed that hydrogen is the essential element in acids.
Liebig (1838). Refined Davy's idea: an acid is a hydrogen-containing compound whose hydrogen can be replaced by a metal.
Arrhenius (1887). Provided the ionic explanation: acids dissociate in water to give $H^+$ .

This progression is a classic example of how a scientific model is refined as new evidence (electrolysis, conductivity, ionic theory) becomes available.

Limitations of the Arrhenius model

Arrhenius works well for simple aqueous acid-base reactions, but it cannot explain:

Basic species without hydroxide. Ammonia ( $NH_3$ ) turns litmus blue and reacts with acids, yet contains no $OH^-$ . Arrhenius cannot account for its basicity.
Non-aqueous acid-base chemistry. $HCl$ reacts with $NH_3$ in the gas phase to form $NH_4Cl$ with no water involved.
The hydrated proton. The bare $H^+$ ion never exists in water; it always attaches to a water molecule to form $H_3O^+$ . Arrhenius treats $H^+$ as a free species.
Acid-base behaviour of salts. Solutions of $NH_4Cl$ are slightly acidic and solutions of $CH_3COONa$ are slightly basic, yet Arrhenius offers no mechanism for these effects.

These limitations motivated the Bronsted-Lowry model (1923), which defines acids as proton donors and bases as proton acceptors.

Common traps

Calling ammonia an Arrhenius base: It is not. Ammonia is a Bronsted-Lowry base because it accepts a proton from water: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$ . The $OH^-$ comes from the water, not from $NH_3$ itself.
Forgetting state symbols: Acid-base equations are expected to carry $(aq)$ or $(l)$ on every species.
Treating the model as wrong, rather than limited: Arrhenius is correct for aqueous, ionic acids and bases. It is a special case of the broader Bronsted-Lowry model. Markers want "limited" not "incorrect".
Confusing strength and concentration: A property like "turns blue litmus red" depends on $[H^+]$ , which is set by both the strength (degree of ionisation) and the concentration. Cover this carefully on the strong vs weak page.
Listing taste or feel as a test: Never test laboratory chemicals by taste or touch. State the property but not the procedure.

Exam technique for acid-base property and titration questions

When asked to IDENTIFY or DESCRIBE a solution's acidity, always pair the observation with the reagent AND the correct colour change (for example, "blue litmus turns red" rather than just "litmus changes colour"). For any concentration or titration calculation, work in this fixed order: write the balanced equation first, convert every volume to litres before substituting, find moles of the KNOWN species, apply the mole ratio from the equation, then divide by volume for concentration, rounding your final answer to match the significant figures given in the question. When reading a titration curve, the equivalence point is the volume at the steepest part of the curve, not necessarily where pH crosses exactly 7 (that is true only for strong acid-strong base titrations).

Examples in context

Example 1. Acid sulfate soils in the Tweed-Richmond floodplain. Coastal floodplains in northern NSW contain pyrite ( $FeS_2$ ) buried in waterlogged sediment. When agricultural drainage exposes the pyrite to oxygen it forms sulfuric acid: $FeS_2 + \frac{7}{2}O_2 + H_2O \rightarrow Fe^{2+} + 2SO_4^{2-} + 2H^+$ . The released acid drops creek pH below 3, killing fish at the Yamba prawn fishery and corroding concrete drains. NSW DPI monitors with blue litmus paper for quick screening, with pH probes for quantitative work. The properties students learn (acids release $H^+$ , lower pH, conduct, react with carbonates) explain every observation in the field, including the white limestone treatment beds installed to neutralise drainage downstream.

Example 2. Drain cleaner manufactured at Bunnings supplier near Smithfield. Domestic drain cleaners sold across NSW contain 30 to 50 percent sodium hydroxide pellets. When dissolved the pellets give $Na^+_{(aq)} + OH^-_{(aq)}$ , a textbook Arrhenius base. The solution turns red litmus blue, conducts strongly, feels slippery on skin (saponification of skin fats), and neutralises acidic blockages by hydrolysing fats and proteins. Safety labelling references the corrosive nature explicitly. Arrhenius theory explains the bulk behaviour cleanly but cannot explain why aqueous ammonia bottles on the same shelf are also basic, which is precisely the reason Bronsted-Lowry was developed to extend the framework.

Try this

Q1. List four observable properties of acids and four of bases, and state how the Arrhenius model accounts for each. [4 marks]

Cue. Acids: sour, conduct, redden blue litmus, react with metals/carbonates; release $H^+$ . Bases: bitter, slippery, turn red litmus blue, react with acids; release $OH^-$ .

Q2. Calculate the pH of a 0.0250 mol L $^{-1}$ solution of $Ca(OH)_2$ , a strong diprotic base. [3 marks]

Cue. $[OH^-] = 2 \times 0.0250 = 0.0500$ , $pOH = 1.30$ , $pH = 12.70$ .

Q3. Discuss the limitations of the Arrhenius model. (a) State one species that is basic but contains no hydroxide. (b) Identify one solvent in which Arrhenius theory does not apply. (c) Explain why Bronsted-Lowry theory was a necessary extension. [1+1+2 marks]

Cue. (a) Ammonia. (b) Liquid ammonia or any non-aqueous solvent. (c) Bronsted-Lowry defines acids and bases by proton transfer, accommodating non-aqueous solvents and species without hydroxide.

Exam-style practice questions

Practice questions written in the style of NESA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

2020 HSC4 marksOutline the key features of the Arrhenius model of acids and bases and identify two limitations of the model.

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A 4 mark answer needs the two definitions, an example, and at least two clear limitations.

Arrhenius model (1887). Svante Arrhenius proposed that acids and bases ionise in water.

An acid is a substance that ionises in water to produce hydrogen ions ( $H^+$ ). Example: $HCl_{(aq)} \rightarrow H^+_{(aq)} + Cl^-_{(aq)}$ .
A base is a substance that ionises in water to produce hydroxide ions ( $OH^-$ ). Example: $NaOH_{(aq)} \rightarrow Na^+_{(aq)} + OH^-_{(aq)}$ .

Neutralisation is the reaction between $H^+$ and $OH^-$ to form water.

Limitations.

The model only describes reactions in aqueous solution. It cannot account for acid-base behaviour in non-aqueous solvents (for example, $HCl$ dissolved in liquid ammonia).
It cannot explain basic species that contain no hydroxide ion, such as $NH_3$ , $CO_3^{2-}$ , or $HCO_3^-$ , which clearly behave as bases by accepting protons.
The notion of a "bare" $H^+$ ion in water is unrealistic. Protons in water exist as hydronium, $H_3O^+$ , so the model is descriptively incomplete.

Markers reward (1) a precise definition of acid and base, (2) a representative equation, (3) two distinct limitations.

2017 HSC3 marksA student is given an unknown colourless solution and asked to determine whether it is acidic, neutral or basic using everyday equipment. Describe a procedure the student could follow and the expected observations.

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A 3 mark answer needs a method, at least two indicators or tests, and the matching observations.

Procedure. Place small samples of the unknown solution in three test tubes. Test each with a different indicator or test.

Universal indicator (or red and blue litmus). Add a few drops. Acidic solutions turn universal indicator red, orange or yellow; neutral solutions give green; basic solutions give blue or purple. Litmus is simpler: red litmus turns blue in base, blue litmus turns red in acid, and neither colour changes in a neutral solution.
Reactivity with a small piece of magnesium ribbon. In an acid, $Mg$ reacts to produce hydrogen gas (visible bubbling, "pop" test with a flame). In neutral or basic solutions there is no reaction.
Conductivity. All ionic solutions conduct, but a strong acid or strong base gives a noticeably brighter bulb than a neutral solution at the same concentration. (This test is supporting, not definitive.)

Markers reward (1) a safe and feasible procedure, (2) at least two distinct tests with the expected observations, (3) a correct interpretation of each observation.

Practice questions

Original practice questions graded from foundation to exam level, each with a full worked solution. Try them before revealing the solution.

foundation3 marksA student is given three unlabelled colourless solutions and told one is acidic, one is neutral and one is basic. Describe how universal indicator and litmus paper could be used together to identify each, and state the expected observation for each solution.

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A 3-mark identify needs a stated method, the two indicators used, and the expected colour for each solution.

Method: Add a few drops of universal indicator to a fresh sample of each solution, and separately dip both red and blue litmus paper into a fresh sample of each solution.
Acidic solution: Universal indicator turns red, orange or yellow. Blue litmus turns red; red litmus stays red.
Neutral solution: Universal indicator stays green. Neither litmus paper changes colour.
Basic solution: Universal indicator turns blue or violet. Red litmus turns blue; blue litmus stays blue.

Marking criteria: 1 mark for a safe, workable method using fresh samples, 1 mark for correctly matching universal indicator colours to each solution, 1 mark for correctly matching litmus colour changes (or lack of change) to each solution.

foundation3 marksState the Arrhenius definition of an acid and of a base, and use them to classify

H_2SO_4

and

Mg(OH)_2

, writing an ionisation equation for each.

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Definitions. An Arrhenius acid is a substance that ionises in water to produce $H^+_{(aq)}$ . An Arrhenius base is a substance that ionises in water to produce $OH^-_{(aq)}$ .

Classification and equations.

H_2SO_{4(aq)} \rightarrow 2H^+_{(aq)} + SO_4^{2-}{}_{(aq)}

$H_2SO_4$ is a diprotic Arrhenius acid, releasing two moles of $H^+$ per mole of acid.

Mg(OH)_{2(aq)} \rightarrow Mg^{2+}_{(aq)} + 2OH^-_{(aq)}

$Mg(OH)_2$ is a diacidic Arrhenius base, releasing two moles of $OH^-$ per mole of base.

Marking criteria: 1 mark for both correct definitions, 1 mark for the correct classification of each substance, 1 mark for both balanced ionisation equations with state symbols.

foundation4 marksCompare Lavoisier's, Davy's and Liebig's models of acids with Arrhenius's model, and explain why each earlier model was replaced.

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Lavoisier (1780s): Proposed all acids contain oxygen (hence the name 'oxygen', 'acid former'). Replaced because Davy showed hydrochloric acid contains no oxygen at all, yet is clearly acidic.
Davy (1810): Proposed hydrogen, not oxygen, is the essential element in acids. This was a step forward but did not explain WHY hydrogen-containing compounds behave as acids, or distinguish acids from other hydrogen-containing compounds like alkanes.
Liebig (1838): Refined Davy's idea: an acid is a compound whose hydrogen can be replaced by a metal (forming a salt). This described behaviour but gave no mechanism for how or why the replacement occurs.
Arrhenius (1887): Provided the missing mechanism: acids ionise in water to release $H^+$ , explaining Liebig's metal-replacement observation (the metal displaces the ionisable $H^+$ ) and unifying all aqueous acid behaviour under one ionic model.

Marking criteria: 1 mark per correct model attributed to the correct scientist and year (Lavoisier, Davy, Liebig), 1 mark for explaining that Arrhenius supplied the ionic mechanism that explained the earlier observations.

core4 marksA 25.00 mL sample of a

Ca(OH)_2

solution is exactly neutralised by 22.40 mL of 0.100 mol L

^{-1}

hydrochloric acid. Calculate the concentration of the

Ca(OH)_2

solution, to 3 significant figures.

Show worked solution →

Step 1: write the balanced equation.

Ca(OH)_{2(aq)} + 2HCl_{(aq)} \rightarrow CaCl_{2(aq)} + 2H_2O_{(l)}

Step 2: moles of HCl used.

n(HCl) = c \times V = 0.100\ \text{mol L}^{-1} \times 0.02240\ \text{L} = 2.240 \times 10^{-3}\ \text{mol}

Step 3: moles of $Ca(OH)_2$ , using the 1:2 mole ratio from the equation.

n(Ca(OH)_2) = \frac{1}{2} \times n(HCl) = \frac{1}{2} \times 2.240 \times 10^{-3} = 1.120 \times 10^{-3}\ \text{mol}

Step 4: concentration of $Ca(OH)_2$ .

c(Ca(OH)_2) = \frac{n}{V} = \frac{1.120 \times 10^{-3}\ \text{mol}}{0.02500\ \text{L}} = 0.04480\ \text{mol L}^{-1}

Step 5: round to 3 significant figures (matching the data).

c(Ca(OH)_2) = 0.0448\ \text{mol L}^{-1}

Marking criteria: 1 mark for the correctly balanced equation, 1 mark for correct moles of HCl, 1 mark for correctly applying the 1:2 mole ratio, 1 mark for the final concentration to 3 significant figures with correct units.

core5 marksThe titration curve below is an owned illustrative graph of pH versus volume of 0.100 mol L

^{-1}

NaOH added to 25.00 mL of a strong monoprotic acid of unknown concentration. (a) Read off the equivalence point volume from the curve. (b) Calculate the original concentration of the acid. (c) State, with reasoning, a suitable indicator for this titration.

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(a) Reading the curve. The steep vertical inflection (the equivalence point) is marked on the curve at $V = 20.00$ mL of NaOH added, where pH rises sharply through pH 7.

(b) Concentration calculation.

At equivalence, moles of $OH^-$ added equal moles of $H^+$ originally present (1:1 stoichiometry for a monoprotic acid with NaOH).

n(OH^-) = c \times V = 0.100\ \text{mol L}^{-1} \times 0.02000\ \text{L} = 2.00 \times 10^{-3}\ \text{mol}

n(H^+)_{\text{original}} = n(OH^-) = 2.00 \times 10^{-3}\ \text{mol}

c(\text{acid}) = \frac{n}{V} = \frac{2.00 \times 10^{-3}\ \text{mol}}{0.02500\ \text{L}} = 0.0800\ \text{mol L}^{-1}

(c) Suitable indicator. Because both the acid and the base are strong, the equivalence point sits at pH 7 with a very steep, near-vertical jump in the curve. Either phenolphthalein (colour change 8.3 to 10.0) or bromothymol blue (colour change 6.0 to 7.6) is suitable, since the steep jump spans several pH units around neutral and both ranges fall within that vertical section. Methyl orange (3.1 to 4.4) would change colour too early, before the true equivalence point, and is not suitable here.

Marking criteria: 1 mark for correctly reading the equivalence volume from the curve, 1 mark for correct moles of $OH^-$ at equivalence, 1 mark for correctly equating moles of acid to moles of base at equivalence (1:1 ratio), 1 mark for the correct final concentration with units and significant figures, 1 mark for naming a suitable indicator with reasoning tied to the steep region of the curve.

exam7 marksAssess the claim that 'the Arrhenius model of acids and bases is simply wrong and should be discarded' by evaluating its explanatory power and its limitations, using ammonia and the acid sulfate soils of a coastal NSW floodplain as evidence.

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This is a 7-mark ASSESS: markers reward a judgement supported by specific evidence, not a one-sided list.

Band 6 PLAN.

Thesis: the claim overstates the case. Arrhenius is not wrong, it is a correct but limited special case that works well within its domain (aqueous, ionisable acids and bases) and should be understood as superseded rather than discarded, because it still correctly predicts and explains the great majority of everyday aqueous acid-base behaviour.
Evidence FOR Arrhenius's continuing usefulness: acid sulfate soils on a coastal NSW floodplain release $H^+$ when pyrite oxidises ( $FeS_2 + \frac{7}{2}O_2 + H_2O \rightarrow Fe^{2+} + 2SO_4^{2-} + 2H^+$ ), and this single ionic species, $H^+_{(aq)}$ , correctly predicts every observed property: lowered pH, litmus colour change, reaction with limestone treatment beds. The model's core insight, that acidity is due to a common ion, is exactly right here.
Evidence AGAINST treating Arrhenius as complete: ammonia, $NH_3$ , dissolved in the same floodplain waterways, behaves as a base (turns red litmus blue, neutralises acid, forms ammonium salts) yet contains no hydroxide ion to ionise. Arrhenius has no mechanism for this and would wrongly classify ammonia solutions as "neither acid nor base" if the definition were applied literally.
Judgement: the model should be described as SUPERSEDED for general use (replaced by Bronsted-Lowry, which correctly classifies ammonia as a proton acceptor) rather than "wrong", because within its aqueous, ionic domain it remains accurate and pedagogically essential as the historical foundation for later models.

Model paragraph (excerpt). The Arrhenius model retains real explanatory power: in the acid sulfate soils of a northern NSW floodplain, the oxidation of pyrite releases $H^+_{(aq)}$ directly into drainage water, and every downstream observation, from the drop in pH to the corrosion of concrete drains, is fully accounted for by Arrhenius's single defining ion. However, the model cannot classify ammonia, a substance found in the same catchments from agricultural runoff, as a base at all, since $NH_3$ contains no hydroxide to release; only the Bronsted-Lowry extension, which defines a base as a proton acceptor, captures this behaviour correctly. The claim that Arrhenius should be "discarded" therefore overstates the case: the model is not false within its domain, it is incomplete outside it, and the historically accurate description is that it was superseded, not disproven.

Marker's note: top-band answers (1) state an explicit thesis that directly engages with the word "wrong" in the prompt, (2) use BOTH pieces of evidence given (acid sulfate soils AND ammonia) rather than only one, (3) reach a judgement using the word "superseded" or equivalent rather than a flat "agree/disagree", and (4) tie the judgement back to the aqueous/ionic domain restriction explicitly.

exam6 marksExplain how the historical progression from Lavoisier through to Arrhenius illustrates the nature and development of scientific models, and justify why this progression is considered scientifically valid rather than a series of failures.

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This is a 6-mark EXPLAIN and JUSTIFY: markers reward a clear causal chain plus an explicit justification of scientific validity.

Band 6 PLAN.

Structure the explanation as a chain: each model was proposed to fit the evidence available at the time, then REFINED (not discarded outright) when new evidence emerged that it could not explain.
Lavoisier (1780s): all acids contain oxygen. Refuted by Davy's 1810 demonstration that hydrochloric acid, clearly acidic, contains no oxygen at all.
Davy (1810): hydrogen is the essential element in acids. An improvement, but did not explain the mechanism of acidity or distinguish acids from other hydrogen-containing compounds.
Liebig (1838): an acid's hydrogen can be replaced by a metal to form a salt. Descriptive and testable (e.g. $2HCl + Mg \rightarrow MgCl_2 + H_2$ ), but still gave no reason WHY replacement happens.
Arrhenius (1887): supplied the mechanism, ionisation in water releasing $H^+$ , which explains Liebig's metal-replacement observation directly (the metal reduces the ionisable $H^+$ ) and unifies all prior observations under one testable, quantitative model.
Justification of validity: this is exactly how science is meant to progress. Each model was not "wrong" in an absolute sense, it was the best available explanation of the evidence THEN known, and was legitimately revised as new experimental evidence (electrolysis, conductivity measurements, ionic theory) became available. This is falsifiability and refinement in action, not failure.

Model paragraph (excerpt). The progression from Lavoisier to Arrhenius shows scientific models being refined, not discarded as failures, in response to new evidence. Lavoisier's oxygen-based model was a reasonable inference from the acids known in the 1780s, but Davy's demonstration that oxygen-free hydrochloric acid is still a strong acid falsified it, prompting Davy's own hydrogen-based replacement. Liebig then made this testable by defining acids through the metal-replacement reaction, and Arrhenius finally supplied the causal mechanism, ionisation, that explained why replacement occurs at all. Each step retained what worked from its predecessor while fixing a specific, evidenced failure, which is precisely the iterative, evidence-driven process that gives scientific models their validity.

Marker's note: top-band answers (1) name all four scientists with correct years, (2) explicitly state what evidence falsified or motivated each transition (not just "it was replaced"), (3) use the word "refined" or equivalent rather than treating earlier models as simply "wrong", and (4) explicitly connect the case study to the general nature of scientific model development (falsifiability, evidence-driven revision).