What is the difference between Arrhenius, Brønsted–Lowry and Lewis acids and bases in H2 A-Level Chemistry? Many JC2 students can repeat that an acid donates a proton and a base accepts a proton, yet stumble when a question introduces NH₃ in water, a BF₃–NH₃ coordinate bond or an amphiprotic ion. The definitions are not three competing slogans. Each is a model that describes a particular kind of acid–base interaction, with different scope and evidence.
For students looking for a JC2 Chemistry tutor in Hougang, understanding these three theories builds stronger explanations of buffers, conjugate pairs, hydrated metal ions and organic mechanisms. The revised Singapore-Cambridge H2 Chemistry 9476 syllabus includes Arrhenius, Brønsted–Lowry and Lewis theories, explicitly including a non-aqueous Lewis example involving BF₃ and NH₃. Rather than memorise a paragraph for each, I recommend starting with the chemical species and asking what actually moves: a proton, or an electron pair.
This guide uses balanced equations, original examples and a practical correction method. eduKateSG’s premium three-student small-group model is at 8 Fourth Avenue near Sixth Avenue MRT, not in Hougang, so families should consider current lesson placement and travel before making arrangements.
Begin with the correct H2 Chemistry scope
The SEAB 2027 H2 Chemistry 9476 syllabus requires understanding of all three acid–base theories and their application, including conjugate acid–base pairs and the non-aqueous reaction between BF₃ and NH₃. It also addresses equilibrium chemistry, weak acids and bases, buffers and relevant pH calculations elsewhere in the course. Students should see these as connected ideas rather than isolated chapters copied from separate revision books.
If an older H1 or O-Level note gives a narrower definition, do not automatically use it as the complete H2 explanation. The Arrhenius model is particularly useful for aqueous hydrogen and hydroxide ions; Brønsted–Lowry broadens the discussion to proton transfer; Lewis reasoning extends it to electron-pair donation and acceptance. Each model can be scientifically useful when its assumptions fit.
Arrhenius acids in aqueous solution
An Arrhenius acid increases the concentration of hydrogen ions, commonly represented through hydronium H₃O⁺, when dissolved in water. Hydrochloric acid can be represented as HCl + H₂O → H₃O⁺ + Cl⁻ in the appropriate aqueous setting. The important restriction is that the theory concerns behaviour in water. A molecule is not classified only from the letters H appearing in its formula.
Students sometimes write H⁺ as though a naked proton exists independently in bulk aqueous solution without solvation. Many school equations use H⁺ as a useful shorthand, but the hydronium representation makes the interaction with water clearer. A student should know the convention used in the question and avoid mistaking notation for a different physical process.
Arrhenius bases in aqueous solution
An Arrhenius base increases hydroxide-ion concentration in water. Sodium hydroxide dissolves and dissociates as NaOH(aq) → Na⁺(aq) + OH⁻(aq) in the straightforward ionic model. This helps explain why a sodium hydroxide solution is strongly alkaline, but it does not fully describe every base that can accept a proton without containing an OH group in its initial molecular formula.
Ammonia exposes the limitation of a too-narrow memory rule. NH₃ does not simply consist of preformed hydroxide ions, yet aqueous ammonia produces OH⁻ through reaction with water. A more useful explanation uses proton transfer and identifies ammonia as a base. That is the reason to learn another theoretical model, rather than insist that every base must display OH in its formula.
Brønsted–Lowry acids donate protons
In the Brønsted–Lowry model, an acid is a proton donor. When ethanoic acid reacts with water, CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺, the ethanoic acid donates a proton and becomes ethanoate. The water molecule accepts that proton and becomes hydronium. This definition explains the reaction through actual proton transfer rather than the overall pH change alone.
A student should identify the particle before and after losing one proton. Do not confuse losing H⁺ with losing a neutral hydrogen atom or an electron. Charge and atom balance still apply. The acid’s conjugate base differs from it by one proton, which provides a quick way to check whether the proposed pair makes chemical sense.
Brønsted–Lowry bases accept protons
A Brønsted–Lowry base accepts a proton. For NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, ammonia accepts H⁺ from water to produce ammonium, while water loses the proton and becomes hydroxide. The same equation therefore illustrates both an acid and a base in a single reaction. The roles are determined by the actual direction of the stated proton transfer.
Students often say water is always neutral and therefore cannot be an acid. In the Brønsted–Lowry model, water can donate a proton to a suitable base as well as accept a proton from a suitable acid. A neutral molecule can participate in acid–base chemistry; its net charge does not alone fix its role.
Conjugate acid–base pairs differ by one proton
CH₃COOH/CH₃COO⁻ is one conjugate acid–base pair, while H₃O⁺/H₂O is another in the ethanoic acid equation. The acid becomes its conjugate base by donating one proton; the base becomes its conjugate acid by accepting one. The partners are directly related by H⁺ transfer, not by an arbitrary chemical reaction between two similar-looking formulas.
A useful annotation method is to draw one arrow between each pair and write ‘+H⁺’ or ‘−H⁺’ above it. This forces the student to conserve atoms and charges. If a proposed conjugate pair differs by oxygen, chlorine or an electron instead of one proton, the classification is wrong. That quick check protects marks in acid–base theory and later buffer calculations.
Water is amphiprotic
Water can act as a Brønsted–Lowry acid or base depending on its partner. It donates a proton to NH₃, forming NH₄⁺ and OH⁻, and accepts a proton from CH₃COOH, forming CH₃COO⁻ and H₃O⁺. A substance that can both donate and accept protons is described as amphiprotic. The term is not a claim that water always performs both roles simultaneously in every reaction.
Students should distinguish amphiprotic from amphoteric, which has a broader meaning involving acid- and base-like behaviour in specified reactions. A substance may be amphoteric without simply transferring an H⁺ as the full explanation. In examination answers, use the term demanded by the reaction and support it with the actual species involved.
Hydrogencarbonate can donate or accept a proton
The ion HCO₃⁻ can accept a proton to form H₂CO₃ in an appropriate acid–base description, and it can donate a proton to become CO₃²⁻. It therefore provides another accessible example of amphiprotic behaviour. The ion’s charge changes correctly in each process: adding a positive proton makes it neutral, while losing a positive proton makes its charge more negative.
Write the two related transformations before assigning the role. If HCO₃⁻ encounters a stronger proton donor, it may act as a base; if it encounters a suitable proton acceptor, it may act as an acid. Context decides the role. Memorising the name of the ion without writing its conjugate species is an easy way to make a charge error.
Why the strength of an acid is not its concentration
Acid strength concerns the extent to which an acid donates protons or ionises under relevant conditions; concentration states how much acid is present per unit volume. A dilute strong acid can have a lower hydronium concentration than a much more concentrated weak acid, even though the weak acid is less extensively dissociated. Calling a solution ‘strong’ merely because a bottle is concentrated confuses two different properties.
This distinction affects pH, electrical conductivity and titration behaviour. Always establish whether the question asks about intrinsic acidic strength, the amount of acid or the resulting solution pH. A scientifically accurate answer may require equilibrium reasoning rather than choosing the strongest-sounding label.
Ka is an equilibrium measure, not a pH value
For a weak acid HA in water, the equilibrium HA + H₂O ⇌ H₃O⁺ + A⁻ can be described by an acid dissociation constant with the appropriate concentration expression under the syllabus conventions. A larger Ka at the same temperature generally indicates greater dissociation and hence a stronger acid among the directly comparable systems. The numerical value of Ka is not itself a pH reading.
A student should not write Ka as though the full equilibrium mixture were a strong-acid solution with complete dissociation unless the approximation is justified. When calculations appear, identify initial and equilibrium species carefully. This is one place where the Brønsted–Lowry model helps explain the chemistry behind the algebra rather than letting symbols substitute for understanding.
Kb and the role of conjugate acids
A weak base B can react with water to produce BH⁺ and OH⁻ in the appropriate equilibrium model. Kb describes that base reaction under defined conditions. When comparing a conjugate acid and base pair, the relative tendencies are connected, but the particular numerical relationship must be used with care and under a consistent temperature convention. The student should understand which species is accepting the proton.
For ammonia in water, NH₃ is the base and NH₄⁺ its conjugate acid. Calling NH₄⁺ a base because it contains nitrogen misses the actual proton transfer. Ask students to draw the lone pair on ammonia and then show how it becomes a bond to the donated proton; the structure clarifies the change.
The Lewis model follows electron pairs
A Lewis base donates an electron pair to make a coordinate bond or take part in an appropriate electron-pair interaction, while a Lewis acid accepts that pair. This definition is broader than proton transfer. A Lewis acid does not have to contain a hydrogen atom, and a Lewis base does not have to release OH⁻ in water. BF₃ reacting with NH₃ is a central H2 example because the process can be understood without any proton movement.
At this level, identify which species possesses an available lone pair and which can accept it. Do not conclude that the donor becomes an Arrhenius acid merely because it contributes electrons. The Lewis model asks a different question from an aqueous pH model: where do the electrons flow when a new bond is formed?
Why BF₃ accepts an electron pair
In an isolated BF₃ molecule, boron has an electron-deficient valence environment and can accept an electron pair from a suitable Lewis base. The molecule has a trigonal planar arrangement in its simple VSEPR description, and an available acceptor orbital allows formation of an adduct. The Lewis acid is BF₃ because it accepts a pair, not because it must first release H⁺ into water.
Students sometimes identify fluorine as the acceptor simply because it is electronegative. But the relevant acceptor centre in the familiar BF₃–NH₃ adduct is boron. An accurate electron-pair arrow begins at the nitrogen lone pair in NH₃ and points toward boron. The arrow is a representation of electron-pair donation, not the physical movement of an entire nitrogen atom.
NH₃ is a Lewis base as well as a Brønsted base
Nitrogen in ammonia has a lone pair available for donation. In water, NH₃ can accept a proton and act as a Brønsted–Lowry base. With BF₃, the nitrogen lone pair can be donated directly to boron to form a coordinate bond. Both reactions are compatible with Lewis basicity, but only the first involves the specific proton-transfer process that defines a Brønsted base.
Students should compare these two settings explicitly. Ask what new bond forms, where its electron pair originated and whether any proton was transferred. The same reagent can be described through different theories depending on the reaction being explained. This is scientific flexibility, not a contradiction.
Drawing the BF₃–NH₃ adduct
A suitable shorthand is H₃N→BF₃, with the arrow indicating donation of nitrogen’s electron pair to boron. After coordination, the environment around boron can be represented as approximately tetrahedral in the simple valence picture. The arrow is not a reaction-rate arrow or a mechanism arrow showing independent proton transfer; it identifies the electron-pair source and destination in the coordinate bond.
The drawing should respect valence and any formal charges implied by the particular representation. Some diagrams mark formal positive and negative centres in the adduct; others emphasise the coordinate bond. Learn the chosen notation and why it reflects electron sharing rather than memorising one image without structural explanation.
Coordinate bonds and covalent bonds
A coordinate covalent bond begins with both shared electrons provided by one partner, according to the electron-pair donation description. Once formed, it is still a covalent interaction, not a permanently separate magical category of bond. The donor and acceptor labels explain the process by which the pair became shared, while the resulting structure determines its properties.
Students who draw an ordinary bond may correctly represent connectivity but fail to answer a question specifically asking for the direction of donation. Conversely, a donated-electron arrow with the wrong atom or lone pair is chemically incorrect. Read the instruction: is the examiner asking for bond formation, a Lewis acid, a Lewis base or a complete structure?
How Brønsted transfer also fits Lewis logic
A Brønsted base accepting H⁺ uses an electron pair to form a bond to the proton, so many proton-transfer reactions can also be interpreted in Lewis terms. In NH₃ + H⁺ → NH₄⁺, ammonia donates its nitrogen lone pair toward the proton. The proton acts as a Lewis acid because it accepts that pair. The proton-transfer description and the electron-pair description are different perspectives on the same chemical change.
That does not mean every Lewis acid–base reaction is Brønsted–Lowry. BF₃ accepting ammonia’s electron pair does not require proton donation. This one-way relationship is an excellent comparative exam question. Students should explain the extra scope of Lewis theory rather than recite that all three definitions are identical.
Metal ions and Lewis acidity
A hydrated metal cation can act as a Lewis acid because electron pairs from ligands such as water coordinate to the metal centre. Highly charged small ions can polarise bound water, connecting Lewis interactions with aqueous acidity. This helps explain why aluminium(III) salts may give acidic aqueous solutions, while a sodium chloride solution behaves differently in the straightforward school model.
The relation is chemically valuable, but a correct answer still has to follow the supplied reaction. If the problem asks about proton release from coordinated water, use the relevant acid–base transfer. If it asks why water binds to the metal, discuss electron-pair donation from water toward the metal centre. Clear questions require specific answers.
The Lewis idea in organic reaction mechanisms
Organic mechanisms often describe nucleophiles as electron-pair donors and electrophiles as electron-pair acceptors. These are closely related to Lewis bases and Lewis acids. When a nucleophile attacks an electrophilic carbon, a curly arrow begins at a lone pair or bond containing the donating electrons, not at an arbitrary atom label. This connection gives students a transferable principle across acid–base and organic topics.
Do not force every organic reaction into a Brønsted proton-transfer explanation. A nucleophilic substitution or carbonyl addition may involve bond formation at an electron-poor carbon without that step being merely an H⁺ donation. Recognising the Lewis model’s wider scope makes mechanism diagrams more coherent and reduces memorised arrow mistakes.
Acid–base models and buffers
A buffer based on a weak acid and its conjugate base involves proton-transfer equilibria. For example, a CH₃COOH/CH₃COO⁻ mixture can consume small additions of H⁺ through the ethanoate component and OH⁻ through the ethanoic acid component. The equilibrium and amounts determine the pH response; the simple phrase ‘resists change’ is an outcome, not a complete mechanistic explanation.
Ask which conjugate species reacts when a little strong acid is added, and which reacts when alkali is added. The Brønsted–Lowry model tells us who gives and receives a proton; a quantitative buffer task adds a mole ledger. Students who understand both can handle changed mixture ratios instead of memorising one numerical formula.
Amphoteric oxides are not necessarily amphiprotic species
Al₂O₃ and Al(OH)₃ display amphoteric acid–base behaviour in suitable reactions. That does not mean the oxide itself is simply a molecule that can donate and accept protons in the same direct way as water. ‘Amphoteric’ is the broader functional classification; ‘amphiprotic’ specifically describes ability to donate and accept protons. The distinction protects conceptual accuracy when Period 3 chemistry and acid–base theories appear together.
In an exam response, name the oxide or hydroxide, describe the reagent and show the resulting equation or relevant species. Do not substitute a label for evidence. When students can produce both acid and alkali responses, the classification becomes understandable and transferable.
How to identify acid and base in an unfamiliar reaction
Start with the actual particles, not their usual names. Look for an H⁺ transfer and track the conjugate pairs; if no proton is transferred, investigate whether one species supplies an electron pair to another. A molecule generally described as a base may act as a ligand or nucleophile in a different setting. If a given reaction does not clearly fit a model, do not force it into the wrong definition.
Check the charge and atom balance after marking donor and acceptor. This often reveals whether a supposed proton-transfer step has actually changed more than one chemical feature. A correct identification should explain the chemical event rather than merely attach the words acid and base to two formulas.
A three-column theory comparison
Create a table with headings theory, key event and relevant example. Arrhenius focuses on aqueous H₃O⁺ or OH⁻ behaviour; Brønsted–Lowry focuses on proton transfer and conjugate pairs; Lewis focuses on electron-pair donation and acceptance. Examples can be HCl in water, NH₃ reacting with water and BF₃ with NH₃. The student should be able to tell what extra range the next definition covers.
Do not let the table replace explanations. Hide the theory names and ask learners to classify three reactions by the actual event. Then change one reagent so that a proton is absent but electron-pair coordination still occurs. The purpose is to teach conceptual boundaries, not to win a vocabulary quiz.
A tutor diagnoses the first wrong model
Imagine Aisha thinks NH₃ cannot be a base because its formula lacks OH. Ben knows ammonia is a base but cannot identify NH₄⁺ as its conjugate acid. Mira draws BF₃ accepting a proton from NH₃ even though the problem concerns a Lewis adduct. These three learners may score similarly yet need different repairs: model scope, proton bookkeeping and electron-pair donation.
In a three-student tutorial, each should first solve a short changed example independently. The tutor then checks the first incorrect reasoning step and sets a different follow-up. Progress is visible when the student can identify donor, acceptor and product after a delay without another spoken hint.
Worked proton transfer: ethanoic acid and ammonia
Consider CH₃COOH + NH₃ ⇌ CH₃COO⁻ + NH₄⁺ in a suitable acid–base description. Ethanoic acid donates a proton to ammonia, becoming its conjugate base, ethanoate. Ammonia accepts that proton, becoming ammonium. The equation conserves atoms and total charge: both sides are overall electrically neutral. It is a clean example because proton transfer can be recognised without first constructing a complicated pH calculation.
Ask students to label two conjugate pairs, not merely the acid and base on the left side. CH₃COOH/CH₃COO⁻ is one pair and NH₄⁺/NH₃ the other. If a learner labels NH₄⁺ as the conjugate base of NH₃, ask whether adding a proton produces an acid or a base relative to that pair. This simple check prevents a repeated exam error.
Worked comparison: ammonia and methylamine
NH₃ and CH₃NH₂ both possess a nitrogen lone pair and can accept a proton, making them Brønsted bases under suitable conditions and Lewis bases when donating electron pairs. Their basic strengths in water differ for reasons involving electron-donating effects and solvation; the methyl substituent affects electron density at nitrogen. However, the comparison must be made with the relevant medium and conditions rather than claiming one universal numerical ordering.
Teach the learner to distinguish ability to accept a proton from the strength of that acceptance in an aqueous equilibrium. Structural features explain tendencies; Ka or Kb data can quantify them. A student who sees an electron-donating substituent should be ready to discuss lone-pair availability, but should not discard solvent and stabilisation effects when the question requests a more nuanced comparison.
Worked comparison: amine and amide
An amine’s nitrogen lone pair can be comparatively available for protonation, whereas an amide’s nitrogen lone pair is delocalised into the neighbouring carbonyl group. This delocalisation reduces the availability of the pair for accepting a proton at nitrogen and helps explain why ordinary amides are much less basic than comparable amines. Simply saying both contain nitrogen is not a sufficient comparison.
The Lewis perspective makes the structural difference explicit: basicity depends in part on how available an electron pair is. Draw the amide’s carbonyl and nitrogen together, then explain the role of resonance rather than pretending the amide nitrogen has the same isolated pair as in an alkylamine. This is a useful bridge between acid–base theory and H2 organic chemistry.
The conjugate base is not always a strong base
The conjugate base of an acid has accepted or retained the electron pair after that acid loses a proton, but its strength as a proton acceptor depends on chemical stability and the relevant solution environment. Cl⁻ is the conjugate base of the strong acid HCl and is extremely weak as a base in ordinary water. Ethanoate, the conjugate base of a weak acid, has more appreciable proton-accepting behaviour in suitable equilibria.
Students may wrongly infer that losing a proton automatically produces a strong base. The conjugate-pair relationship defines identities; equilibrium strength is an additional question. Use the actual acid dissociation tendency and relevant data instead of equating the word ‘conjugate’ with ‘strong’. This distinction becomes important when classifying aqueous salts.
Why ammonium chloride solution can be acidic
Ammonium ions, NH₄⁺, can donate a proton to water through the equilibrium NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺. In a solution of NH₄Cl, chloride ions are not generally the main source of the acidity in the standard aqueous explanation. The ammonium ion is the conjugate acid of the weak base ammonia, and its proton-transfer behaviour matters.
This example tests whether a student can transfer Brønsted theory to a dissolved salt rather than treat every acid as a molecular compound whose formula begins with H. A useful follow-up is sodium ethanoate solution, whose ethanoate ions can accept protons from water and generate some OH⁻. The ions present determine the chemical response.
Why sodium ethanoate solution can be alkaline
Ethanoate ions act as weak bases in water: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻. Sodium is ordinarily a spectator ion for this acid–base explanation. Therefore an aqueous sodium ethanoate solution can be alkaline under the usual conditions. The relevant species is the conjugate base of a weak acid, not a pre-existing NaOH molecule mixed into the salt.
This makes a valuable paired exercise with ammonium chloride. Ask which dissolved ion participates in proton transfer, what conjugate species forms and how the hydronium or hydroxide concentration is affected. Once the student can explain both without a memorised pH list, they have a transferable model for unfamiliar weak-acid and weak-base salts.
A buffer’s first step is sometimes a chemical reaction
Suppose a buffer contains CH₃COOH and CH₃COO⁻, and a small amount of strong acid is added. The ethanoate ions consume the added H⁺ to make more ethanoic acid. If a small amount of strong alkali is added, ethanoic acid can consume OH⁻ and produce ethanoate and water. Only after accounting for that reaction should a quantitative problem use the resulting weak-acid/conjugate-base ratio where the appropriate approximation is valid.
A common wrong move is to insert the initial amounts into a buffer pH relationship after adding a strong reagent, as though no neutralisation occurred. The Brønsted model shows which species donates or accepts the proton; a mole ledger then updates the amount. Chemical meaning must come before mathematical substitution.
pH at a different temperature
The water equilibrium and ion-product constant depend on temperature. Neutrality means equal hydronium and hydroxide concentrations in the relevant aqueous model; it does not guarantee pH is exactly 7 at every possible temperature. At approximately 25 °C, the familiar pure-water value is pH 7, but temperature changes the underlying equilibrium and the numerical neutral pH.
Students who remember only ‘pH 7 equals neutral’ may overgeneralise in a data-based question that explicitly changes temperature. If the question supplies a new Kw or relevant data, use it. This is another example of why definitions and model conditions should govern an answer rather than a single numerical fact learned earlier in school.
Lewis acids do not require an aqueous solution
The BF₃–NH₃ coordinate adduct illustrates a Lewis acid–base reaction that does not depend on producing aqueous H₃O⁺ or OH⁻. This broadens the acid–base idea beyond the Arrhenius frame. A student should recognise that BF₃ accepts a lone pair from nitrogen despite containing no proton, while ammonia donates the pair despite not having an OH group in its molecular formula.
Try a short response of exactly two linked sentences: ‘NH₃ is the Lewis base because nitrogen donates its lone pair. BF₃ is the Lewis acid because boron accepts that pair to form a coordinate bond.’ Then redraw the adduct with a donation arrow and ask whether any proton was transferred. The contrast makes the theory boundaries clear.
A VSEPR check on the Lewis adduct
In simple VSEPR terms, the isolated BF₃ molecule is trigonal planar around boron, with three bonding regions. On accepting ammonia’s lone pair, the boron centre gains an additional coordination interaction and can be described with four bonding regions in the adduct, giving an approximately tetrahedral arrangement. The geometry change is evidence of a new coordination environment rather than a statement that fluorine atoms were lost.
Students should not call the nitrogen the Lewis acid merely because the final structure contains a new N–B bond. The source of the pair matters. A shape question and an acid–base question can use the same molecule but ask different outputs; the response should be shaped by the command word.
An exam comparison of theories
Suppose a prompt gives HCl in water, NH₃ in water and BF₃ with NH₃. A strong response identifies aqueous hydronium formation for the first, proton transfer and conjugate pairs for the second, and electron-pair acceptance/donation for the third. It is unnecessary and potentially misleading to force all three examples into the narrowest definition. The models can overlap while differing in scope.
Train with an unlabeled set of equations. Ask learners to state what event is visible in each and which theory most directly explains it. Then ask why the same reaction might also fit a broader model. This produces flexible scientific understanding instead of three memorised sentences recited on command.
What parents can ask after a lesson
Ask your JC2 student to draw one proton-transfer equation and one electron-pair donation diagram, then point out what physically changed in each. If they can identify conjugate pairs and name the source of a lone pair without tutor hints, the lesson produced an observable skill. A large sheet of definitions is less convincing unless that knowledge transfers into unfamiliar reaction examples.
Learning evidence should be checked again after a delay, with different reagents. Parents need not assess every equilibrium constant or provide chemistry corrections themselves. A good teaching plan will show the student’s specific mistakes, the repaired principle and a changed task solved independently. That is a stronger measure than a tutor’s claim that acid–base theory has been fully ‘covered’.
Frequently asked questions about acid–base theories
Does every acid have to contain H? A Brønsted acid donates a proton, but a Lewis acid such as BF₃ can accept an electron pair without containing hydrogen. Is NH₃ a base without OH? Yes; it accepts a proton from water. What is a conjugate pair? Two species related by the gain or loss of one proton. What does amphiprotic mean? Able to donate and accept protons.
Why is BF₃ a Lewis acid? Boron accepts a donated electron pair. Why is NH₃ a Lewis base? Nitrogen has an available lone pair. Are strong and concentrated the same? No. Can a metal ion be a Lewis acid? Yes, through coordination with electron-pair-donating ligands. Should every H2 acid–base answer use all three models? No; use the model relevant to the reaction and command word.
Continue from definitions to actual mastery
Work through BF₃ and NH₃ electron-pair donation, JC2 acid–base equilibrium and buffers, Period 3 acidity and hydrolysis and the SEAB 2027 H2 Chemistry syllabus. These show how three definitions become chemical explanations.
For JC2 Chemistry tuition, contact eduKate Singapore or enquire about current classes. Confirm the correct syllabus level, timetable, fees and travel. The examples here are teaching illustrations; actual examination answers depend on the stated chemical situation.