In 1947, industrial chemist Wendell Latimer and William Jolly at the University of California systematically catalogued how naming conventions predict reaction products — showing that knowing whether HCl is a binary acid (hydro-prefix) or H₂SO₄ is an oxoacid (suffix from ion name) lets you predict every neutralisation product before mixing a drop. NESA adopted the same IUPAC two-pattern framework in its current syllabus.
Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
A winemaker adds three drops of a pale orange solution to a sample of wine. It turns yellow. She records "too acidic" in her logbook and adjusts the fermentation. A pharmacist dips a paper strip into dissolved aspirin tablets and it turns pink — he confirms the solution is appropriately basic. A soil scientist tests paddock soil and gets orange — he prescribes a lime treatment. All three are using indicators. All three solutions looked colourless or pale before the indicator was added. The indicator is the same type of substance in every case — a weak acid.
Question 1: How does a weak acid change colour based on the acidity of its surroundings? Write your best molecular-level explanation before reading on.
Question 2: You see two bottles on a shelf labelled "HBr" and "H₂SO₄." Without testing the solutions, can you name each compound from the formula alone? What rule or pattern do you use?
No calculation formulas this lesson — nomenclature, indicators, and reaction patterns are conceptual and classification-based.
Core Content
Two patterns: binary (hydro-/ic) vs oxoacid (from ion name) — never mix them
Acid names follow two distinct patterns depending on whether or not the acid contains oxygen — once you identify which pattern applies, you can name or decode any common inorganic acid systematically.
Binary acids contain hydrogen bonded to a single non-metal with no oxygen present. They are named: hydro + [non-metal root] + ic acid. Examples: HCl = hydrochloric acid; HF = hydrofluoric acid; HBr = hydrobromic acid; HI = hydroiodic acid; H₂S = hydrosulfuric acid.
Oxoacids contain hydrogen, a non-metal, and oxygen. They are named directly from the polyatomic ion they contain — using the ion's suffix to determine the acid's suffix:
| Formula | Name | Ion it contains | Binary or oxoacid | Strong or weak |
|---|---|---|---|---|
| HCl | Hydrochloric acid | Cl⁻ | Binary | Strong ✓ |
| HF | Hydrofluoric acid | F⁻ | Binary | Weak ⚠️ |
| HBr | Hydrobromic acid | Br⁻ | Binary | Strong ✓ |
| HI | Hydroiodic acid | I⁻ | Binary | Strong ✓ |
| H₂S | Hydrosulfuric acid | S²⁻ | Binary | Weak |
| H₂SO₄ | Sulfuric acid | SO₄²⁻ (sulfate) | Oxoacid | Strong (1st ionisation) |
| H₂SO₃ | Sulfurous acid | SO₃²⁻ (sulfite) | Oxoacid | Weak |
| HNO₃ | Nitric acid | NO₃⁻ (nitrate) | Oxoacid | Strong ✓ |
| HNO₂ | Nitrous acid | NO₂⁻ (nitrite) | Oxoacid | Weak |
| H₃PO₄ | Phosphoric acid | PO₄³⁻ (phosphate) | Oxoacid | Weak |
| H₂CO₃ | Carbonic acid | CO₃²⁻ (carbonate) | Oxoacid | Weak |
Binary acids (H + non-metal, no O) use the hydro-/ic pattern (e.g. HCl = hydrochloric acid); oxoacids use the ion suffix (-ate → -ic acid; -ite → -ous acid; never use "hydro-") — exception: HF is a weak acid despite binary naming because the short, strong H–F bond resists proton donation.
Pause — copy the highlighted definition into your book before moving on.
+5 XP — Which name and pattern is correct for H₂SO₄?
Four strong hydroxide bases + NH₃ + salt-type bases — memorise the strong ones
Most common inorganic bases follow standard ionic compound naming with one important extension — some bases contain no hydroxide at all, and their classification as bases requires the Brønsted-Lowry model rather than the Arrhenius model.
Hydroxide bases are named as standard ionic compounds: [metal name] + hydroxide. The four strong bases you must memorise are:
| Formula | Name | Classification | Strong or weak |
|---|---|---|---|
| NaOH | Sodium hydroxide | Hydroxide base | Strong ✓ |
| KOH | Potassium hydroxide | Hydroxide base | Strong ✓ |
| Ca(OH)₂ | Calcium hydroxide | Hydroxide base | Strong ✓ |
| Ba(OH)₂ | Barium hydroxide | Hydroxide base | Strong ✓ |
| NH₃ | Ammonia | Brønsted-Lowry base (proton acceptor) | Weak |
| Na₂CO₃ | Sodium carbonate | Salt with basic anion (CO₃²⁻ accepts H⁺) | Weak (by hydrolysis) |
| NaHCO₃ | Sodium hydrogen carbonate | Salt with amphiprotic anion | Weakly basic |
| Mg(OH)₂ | Magnesium hydroxide | Hydroxide base | Weak (sparingly soluble) |
NH₃ is the most important weak base in this module. Older texts sometimes write its aqueous solution as "ammonium hydroxide (NH₄OH)" — this name is chemically misleading because free NH₄OH molecules are not significantly present in solution. The correct description is: NH₃ dissolves in water and partially accepts protons from water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, producing a basic solution. The OH⁻ comes from water, not from NH₃ itself.
Na₂CO₃ is basic because the carbonate ion (CO₃²⁻) is the conjugate base of the weak acid H₂CO₃ — it accepts protons from water, making the solution basic. This is salt hydrolysis, developed further in L06.
The four strong bases are NaOH, KOH, Ca(OH)₂, Ba(OH)₂ (fully dissociate, use →); all others (NH₃, Na₂CO₃, Mg(OH)₂) are weak — NH₃ is a base because it accepts H⁺ from water (NH₃ + H₂O ⇌ NH₄⁺ + OH⁻), not because it contains OH⁻; never write NH₄OH for dissolved ammonia.
Add the highlighted point to your notes before the check below.
+5 XP — Which of the following is a STRONG base?
HIn ⇌ H⁺ + In⁻ · Le Chatelier drives the colour change · Three indicators to memorise
We just saw that bases like NH₃ and CO₃²⁻ accept H⁺ from water to produce OH⁻. That raises a question: If we can't see H⁺ or OH⁻ directly, how do we detect whether a solution is acidic or basic? This card answers it → indicators are weak acids whose two forms have different colours, so they visibly shift colour when [H⁺] changes.
An indicator is not a passive dye that detects pH from the outside — it is a weak acid actively participating in an equilibrium, and its colour change is a direct, predictable consequence of Le Chatelier's Principle responding to changes in [H⁺].
Indicators are weak acids represented as HIn, where the acid form (HIn) and its conjugate base (In⁻) have distinctly different colours. The equilibrium is:
HIn(aq) ⇌ H⁺(aq) + In⁻(aq) [acid colour] → [base colour]
In acidic solution: [H⁺] is high. The additional H⁺ shifts the indicator equilibrium to the LEFT (Le Chatelier — system opposes the increase in [H⁺] by consuming H⁺ and shifting toward HIn). The HIn form predominates → acid colour is seen.
In basic solution: [H⁺] is low (OH⁻ reacts with H⁺ → H₂O, removing H⁺). The equilibrium shifts to the RIGHT to replace H⁺ being consumed. The In⁻ form predominates → base colour is seen.
The colour transition occurs over a range of approximately pKa(In) ± 1 — spanning roughly two pH units. Inside this range, both forms are present and an intermediate (mixed) colour is observed.
Acid colour (HIn): Red
Base colour (In⁻): Yellow
Transition range: pH 3.1–4.4
pKa (approx): 3.5
Acid colour (HIn): Yellow
Base colour (In⁻): Blue
Transition range: pH 6.0–7.6
pKa (approx): 7.0
Acid colour (HIn): Colourless
Base colour (In⁻): Pink
Transition range: pH 8.3–10.0
pKa (approx): 9.2
Indicator equilibrium — direction of shift and resulting colour determined by [H⁺] in solution
Indicator choice depends on where the colour-change range sits on the pH scale.
An indicator is a weak acid (HIn ⇌ H⁺ + In⁻) whose two forms have different colours — high [H⁺] shifts the equilibrium left (acid colour); low [H⁺] shifts it right (base colour); memorise: methyl orange red/yellow (pH 3.1–4.4), bromothymol blue yellow/blue (6.0–7.6), phenolphthalein colourless/pink (8.3–10.0).
Pause — copy the highlighted definition into your book before moving on.
+5 XP — Phenolphthalein is added to a solution of pH 6. What colour is observed?
Identify the reactant type → apply the pattern → determine salt formula → balance
We just saw that indicators detect pH by responding to [H⁺] via Le Chatelier's Principle. That raises a question: When an acid actually reacts with something, how do we predict what products form? This card answers it → three reaction patterns (acid + base, acid + carbonate, acid + metal) each produce predictable products once you identify the reactant type.
Acids react with three classes of substances in predictable, reproducible patterns — learn the pattern once, identify the reactant type, and you can write and balance any of these equations without memorising individual reactions.
Pattern 1 — Acid + base (neutralisation). Products are always a salt and water. The salt is formed from the metal cation of the base and the anion of the acid. The net ionic equation for strong acid + strong base is always H⁺ + OH⁻ → H₂O (spectator ions don't participate).
Pattern 2 — Acid + carbonate or hydrogen carbonate. Products are always a salt, water, and carbon dioxide gas. CO₂ is produced because carbonic acid (H₂CO₃), the intermediate product, is unstable and decomposes immediately: H₂CO₃ → H₂O + CO₂. The bubbling observed is CO₂ escaping as a gas. All three products must appear in the balanced equation.
Pattern 3 — Acid + reactive metal. Products are always a salt and hydrogen gas (H₂). Only metals more reactive than hydrogen react: Mg, Al, Zn, Fe, Sn — yes; Cu, Ag, Au — no (below hydrogen in activity series). The H₂ gas bubbles are the observable sign.
Reactants: H⁺ source + OH⁻ source
Products: Salt + H₂O
Key observable: Temperature rise (exothermic)
Example: HCl + NaOH → NaCl + H₂O
Reactants: H⁺ source + CO₃²⁻
Products: Salt + H₂O + CO₂(g)
Key observable: Bubbling; solid dissolves
Example: H₂SO₄ + CaCO₃ → CaSO₄ + H₂O + CO₂
Reactants: H⁺ source + HCO₃⁻
Products: Salt + H₂O + CO₂(g)
Key observable: Bubbling; fizzing
Example: HCl + NaHCO₃ → NaCl + H₂O + CO₂
Reactants: H⁺ source + metal above H
Products: Salt + H₂(g)
Key observable: Bubbling; metal dissolves
Example: Zn + 2HCl → ZnCl₂ + H₂
How to find the salt formula: Identify the metal cation from the base (or metal) and the anion from the acid. Balance the charges to get the correct formula before you start balancing the equation.
The reaction pattern determines the products: acid + base gives salt and water, while carbonates and metals introduce gas products.
Three acid reaction patterns: (1) acid + base → salt + water; (2) acid + carbonate/H-carbonate → salt + water + CO₂ (all three products always); (3) acid + reactive metal (above H in activity series) → salt + H₂(g) — determine the salt formula by pairing the metal cation with the acid anion and balancing charges first.
Add the highlighted point to your notes before the check below.
+5 XP — Which set of products is correct for: HCl + Na₂CO₃ →
Ammonium sulfate [(NH₄)₂SO₄] is one of the world's most widely used nitrogen fertilisers, providing both nitrogen (for plant protein synthesis) and sulfur (for amino acid production). It is produced industrially by the Brønsted-Lowry acid-base reaction: 2NH₃(g) + H₂SO₄(aq) → (NH₄)₂SO₄(aq). This is an acid + base reaction following Pattern 1 — the products are the salt (NH₄)₂SO₄ and water. Sulfuric acid (the acid) donates protons to ammonia (the base), forming ammonium ions (NH₄⁺) and sulfate ions (SO₄²⁻) which combine as the ionic salt.
This reaction is also a direct industrial application of nomenclature: the salt name is ammonium sulfate — ammonium (from NH₄⁺) + sulfate (from SO₄²⁻). The formula (NH₄)₂SO₄ is determined by balancing the 1+ charge of NH₄⁺ against the 2− charge of SO₄²⁻ (need two NH₄⁺ per SO₄²⁻). Notice that the Haber process (Module 5) and this acid-base reaction together form the industrial chain from atmospheric N₂ → NH₃ → fertiliser.
"H₂SO₄ should be named hydrosulfuric acid." — Never use "hydro-" for an oxoacid. H₂SO₄ contains oxygen, so it is an oxoacid named from its ion: sulfate (SO₄²⁻) → sulfuric acid. Hydrosulfuric acid is the name for H₂S (binary acid, no oxygen).
"HF is a strong acid because it's a binary acid like HCl." — HF is a weak acid (Ka = 6.8 × 10⁻⁴). The H–F bond is exceptionally short and strong due to fluorine's small atomic radius, making it hard to donate the proton despite fluorine's high electronegativity. HBr and HI are strong because the H–Br and H–I bonds are much weaker (larger atoms, longer bonds).
"An indicator shows the exact pH of a solution." — Indicators show only whether the pH is below, within, or above their transition range. Within the range, an intermediate colour is seen but no exact pH is determined. For precise pH measurement, a calibrated digital pH meter is required.
"Phenolphthalein turns pink at pH 7." — Phenolphthalein's transition range is pH 8.3–10.0. It is colourless at pH 7 (and at all pH values below ~8.3). It only begins turning pink above pH 8.3 and is fully pink above pH 10.
"Acid + carbonate produces only salt and water." — The products are always salt + water + CO₂(g). All three must appear in a balanced equation. Omitting CO₂ is a formula error — it is the reason for the observed bubbling in these reactions.
✏️ Worked Examples
HNO₂ — contains H, N, and O → oxoacid. The ion is NO₂⁻ = nitrite (ends in -ite) → acid name ends in -ous → nitrous acid. Classification: weak acid.
H₃PO₄ — contains H, P, and O → oxoacid. The ion is PO₄³⁻ = phosphate (ends in -ate) → acid name ends in -ic → phosphoric acid. Classification: weak acid (not fully ionised).
Ba(OH)₂ — barium cation + hydroxide anion → standard ionic naming: barium hydroxide. Classification: strong base (fully dissociates: Ba(OH)₂ → Ba²⁺ + 2OH⁻).
H₂S — contains H and S, no oxygen → binary acid → hydro + sulf + ic → hydrosulfuric acid. Classification: weak acid (Ka₁ = 9.5 × 10⁻⁸; the H–S bond is weak but the acid only partially ionises).
Answers: (a) Nitrous acid — weak acid. (b) Phosphoric acid — weak acid. (c) Barium hydroxide — strong base. (d) Hydrosulfuric acid — weak acid.
Acid + base (Pattern 1). Salt formula: Mg²⁺ from Mg(OH)₂, SO₄²⁻ from H₂SO₄ → MgSO₄. Products: MgSO₄ + H₂O. Unbalanced: H₂SO₄ + Mg(OH)₂ → MgSO₄ + H₂O. Left has 4H; right has 2H → need 2H₂O.
H₂SO₄ + Mg(OH)₂ → MgSO₄ + 2H₂O
Check: S=1✓, Mg=1✓, H=4✓, O=4+2=6 left; 4+2=6 right ✓
Acid + carbonate (Pattern 2). Salt formula: Ca²⁺ from CaCO₃, Cl⁻ from HCl → CaCl₂. Products: CaCl₂ + H₂O + CO₂. Need 2HCl to give 2Cl⁻.
2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂
Check: Ca=1✓, Cl=2✓, H=2✓, C=1✓, O=3 left; 1+2=3 right ✓
Acid + reactive metal (Pattern 3). Al is above H in the activity series ✓. Salt formula: Al³⁺ + NO₃⁻ → Al(NO₃)₃ (need 3 NO₃⁻ for Al³⁺). Products: Al(NO₃)₃ + H₂. Unbalanced: HNO₃ + Al → Al(NO₃)₃ + H₂. Need 3HNO₃ → 3H → 1.5H₂; multiply ×2:
6HNO₃ + 2Al → 2Al(NO₃)₃ + 3H₂
Check: Al=2✓, N=6✓, H=6✓, O=18 left; 18 right ✓
Answers: (a) H₂SO₄ + Mg(OH)₂ → MgSO₄ + 2H₂O (b) 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂ (c) 6HNO₃ + 2Al → 2Al(NO₃)₃ + 3H₂
Orange at pH 4.0: Methyl orange is a weak acid: HIn (red) ⇌ H⁺ + In⁻ (yellow), transition range pH 3.1–4.4, pKa ≈ 3.5. At pH 4.0, the solution is within the transition range. Since pH 4.0 > pKa (3.5), the equilibrium lies slightly to the right — more In⁻ (yellow) than HIn (red) — but both forms are present in significant amounts. The solution appears orange because it is a visual mixture of the red (HIn) and yellow (In⁻) forms. Neither form so strongly dominates that a pure colour is visible.
Colour shifts to red after HCl added: Adding HCl increases [H⁺] significantly. By Le Chatelier's Principle, the indicator equilibrium HIn ⇌ H⁺ + In⁻ is disturbed by the increased [H⁺]. To oppose this increase, the equilibrium shifts LEFT, consuming H⁺ and converting In⁻ back to HIn. The HIn (red) concentration increases and the In⁻ (yellow) concentration decreases. The solution shifts toward red because the acid form (HIn) now dominates after the leftward shift.
Answer: (a) At pH 4.0, both HIn (red) and In⁻ (yellow) forms are present in significant amounts — pH is within the transition range (3.1–4.4). Orange = visual mixture of both forms. (b) Adding HCl increases [H⁺]; Le Chatelier shifts HIn ⇌ H⁺ + In⁻ LEFT; HIn (red) dominates; colour shifts toward red.
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Activities
Each statement or equation below contains at least one error. Identify the error and write the corrected version with an explanation.
Error 1 — Naming
Error 2 — Indicator prediction
Error 3 — Balanced equation
Error 4 — Acid + metal
Error 5 — Indicator equilibrium explanation
❓ Multiple Choice
1. A student prepares an aqueous solution of HBr and wants to name it correctly. Which name and classification is correct?
2. Phenolphthalein (transition range pH 8.3–10.0) is added to three solutions: A (pH 5), B (pH 9), C (pH 12). What colours are observed?
3. Excess zinc is added to dilute sulfuric acid. Which equation is correctly balanced and identifies all products correctly?
4. Which of the following best explains why HF is a weak acid but HCl, HBr, and HI are all strong acids?
5. A student adds bromothymol blue to a solution and observes a green colour (a mixture of yellow and blue). Which of the following pH values is most consistent with this observation?
✍️ Short Answer
6. (a) Explain how indicators function as weak acids. Use the equilibrium expression HIn ⇌ H⁺ + In⁻ and Le Chatelier's Principle to explain why the indicator appears the acid colour in acidic solution and the base colour in basic solution. (b) A student dissolves aspirin tablets and tests the solution with bromothymol blue — the indicator turns yellow. What can the student conclude about the pH of the solution? (5 marks)
7. Write balanced molecular equations for the following reactions. State the reaction type (Pattern 1, 2, or 3) for each.
(a) Phosphoric acid (H₃PO₄) reacts with potassium hydroxide (KOH) to form tripotassium phosphate (2 marks)
(b) Hydrochloric acid reacts with marble chips (CaCO₃) (2 marks)
(c) Dilute sulfuric acid reacts with iron filings (2 marks) (6 marks)
8. Real-World Application: A fertiliser manufacturer produces ammonium sulfate [(NH₄)₂SO₄] by reacting ammonia with sulfuric acid. The process uses 98% concentrated H₂SO₄ and requires careful pH monitoring using bromothymol blue indicator.
(a) Write the balanced equation for the production of ammonium sulfate. Identify the Brønsted-Lowry acid and base. (2 marks)
(b) Explain why bromothymol blue is a suitable indicator for monitoring whether the reaction has reached neutralisation. What colour change would indicate the endpoint has been reached? (3 marks) (5 marks)
Error 1: H₂SO₄ contains oxygen — it is an oxoacid, not a binary acid. The "hydro-" prefix is never used for oxoacids. Correct name: sulfuric acid (from sulfate SO₄²⁻ → -ate → -ic acid). "Hydrosulfuric acid" is the name of H₂S.
Error 2: Two errors. First, phenolphthalein's transition range is pH 8.3–10.0; at pH 6 the solution is below this range and phenolphthalein is colourless (HIn form dominates). Second, phenolphthalein is colourless in the acid form — not pink. It is pink in the base (In⁻) form only above pH 8.3.
Error 3: Missing two products and incorrect balancing. Acid + carbonate → salt + water + CO₂. Correct: 2HCl + Na₂CO₃ → 2NaCl + H₂O + CO₂. (2HCl for 2Na⁺ and 1CO₃²⁻; water and CO₂ always appear together.)
Error 4: Two errors. Copper (Cu) is below hydrogen in the activity series — dilute H₂SO₄ does NOT react with copper by Pattern 3. The reaction as written does not occur under these conditions. Additionally, the gas product of acid + metal is H₂ (diatomic), not H. If a metal above hydrogen were used (e.g. Fe), the product would be H₂, not H.
Error 5: The explanation ignores the indicator equilibrium and Le Chatelier entirely. Colour change is not a light-absorption/reflection property of the static molecules — it is a direct result of the equilibrium shift. In basic solution, [H⁺] decreases (OH⁻ reacts with H⁺ → H₂O), shifting HIn ⇌ H⁺ + In⁻ to the RIGHT. The In⁻ form (yellow) increases in concentration and dominates → yellow colour observed.
1. C — HBr contains H and Br, no oxygen → binary acid → hydro + brom + ic = hydrobromic acid. Strong acid (complete ionisation). Option B is the name of the pure gas HBr(g), not the aqueous acid solution. Option A applies the oxoacid pattern incorrectly. Option D invents a non-existent "-ous" form for a binary acid.
2. A — pH 5 is below 8.3 → HIn dominates → colourless. pH 9 is within range (8.3–10.0) → In⁻ dominates → pink. pH 12 is above 10.0 → In⁻ still dominates → remains pink. Common error: thinking the indicator returns to colourless above the range — it stays pink because In⁻ still dominates at very high pH.
3. A — Acid + reactive metal → salt + H₂. Zn²⁺ + SO₄²⁻ → ZnSO₄. One Zn, one H₂SO₄, one ZnSO₄, one H₂ — balanced. Option B incorrectly gives ZnS. Option C uses Zn₂SO₄ (Zn is 2+, not 1+). Option D describes concentrated hot H₂SO₄ — not tested at this level.
4. B — The H–F bond is unusually short and strong because F has the smallest atomic radius of the halogens, creating a very compact bond that requires substantial energy to break. As atomic radius increases (Cl < Br < I), the H–X bond becomes longer and weaker, allowing progressively easier proton donation (HCl, HBr, HI all strong). Option C is incorrect — F is MORE electronegative than Cl, Br, I; it holds electrons tightly, but the bond strength is the key factor.
5. C — Bromothymol blue transition range: pH 6.0–7.6. Green = mixture of yellow (HIn, acid form) and blue (In⁻, base form) = both forms present in similar concentrations = within the transition range. pH 6.8 is within 6.0–7.6 ✓. pH 4.0 is below the range (pure yellow). pH 5.5 is below the range (yellow only). pH 9.0 is above the range (pure blue).
Q6 (5 marks): (a) Indicators are weak acids that exist in equilibrium: HIn ⇌ H⁺ + In⁻ [1]. In acidic solution, [H⁺] is high — Le Chatelier's Principle predicts the equilibrium shifts LEFT (to oppose the increased [H⁺]), increasing [HIn] so the acid colour dominates [1]. In basic solution, OH⁻ reacts with H⁺ (H⁺ + OH⁻ → H₂O), reducing [H⁺] — Le Chatelier shifts the equilibrium RIGHT to replace H⁺, increasing [In⁻] so the base colour dominates [1]. (b) Bromothymol blue's transition range is pH 6.0–7.6. Yellow = acid form (HIn) dominates [1]. Therefore the aspirin solution has pH < 6.0 — it is acidic, consistent with the presence of acetylsalicylic acid [1].
Q7 (6 marks): (a) Type: acid + base (Pattern 1) [1]. Salt: K⁺ + PO₄³⁻ → K₃PO₄ (need 3KOH for one H₃PO₄ and one PO₄³⁻). Equation: H₃PO₄ + 3KOH → K₃PO₄ + 3H₂O [1]. (b) Type: acid + carbonate (Pattern 2) [1]. 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂ [1]. (c) Type: acid + reactive metal (Pattern 3) [1]. Fe is above H in the activity series; Fe²⁺ in dilute H₂SO₄. Fe + H₂SO₄ → FeSO₄ + H₂ [1].
Q8 (5 marks): (a) 2NH₃(g) + H₂SO₄(aq) → (NH₄)₂SO₄(aq) [1]. BL acid = H₂SO₄ (donates H⁺ to NH₃); BL base = NH₃ (accepts H⁺ from H₂SO₄) [1]. (b) Bromothymol blue (range 6.0–7.6) is suitable because this is a weak base (NH₃) + strong acid (H₂SO₄) reaction, giving an equivalence point at pH < 7 — the acidic NH₄⁺ salt forms [1]. The bromothymol blue range spans the expected equivalence point pH [1]. Colour change at endpoint: from blue (excess NH₃, basic solution) through green (near endpoint) to yellow (excess H₂SO₄ or at EP < 7) — the change from green/blue to yellow signals the endpoint has been reached [1].
Go back to your Think First response at the top of this lesson. The red colour of universal indicator in an acidic mystery solution is explained by the same equilibrium Runge described in 1834: HIn ⇌ H⁺ + In⁻ shifts left in acid, favouring the HIn colour. Now check your answers:
Review — Flashcard Drill
Flashcard 1: State the naming rule for a binary acid (no oxygen). Give an example.
Flashcard 2: State the naming rule for an oxoacid. Give an example showing both -ate and -ite cases.
Flashcard 3: Write the indicator equilibrium equation and explain the direction of shift in (a) acid and (b) base.
Flashcard 4: List the three acid reaction patterns and their products.
Flashcard 5: Give the transition range and colour change (acid→base) for all three HSC indicators.
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