Untitled Flashcard Set

Study Guide: Acid-Base Concepts

1. Fundamental Definitions and Terminology

• Proton (\text{H}^+): A positively charged hydrogen atom. A neutral hydrogen atom consists of one proton and one electron; when it loses its electron, only the proton remains.

• Hydronium Ion (\text{H}_3\text{O}^+): An ion formed when a water molecule accepts a proton from an acid.

• Amphiprotic Substance: A substance capable of both accepting and donating a proton depending on the reaction environment. Water (\text{H}_2\text{O}) is a primary example of an amphiprotic substance.

• Ionisation Reaction: A reaction in which a neutral molecule reacts with water to form ions.

• \text{pH}: A relative measure of the acidity or basicity (alkalinity) of a solution.

2. Theoretical Models of Acids and Bases

Arrhenius Theory

Proposed by Swedish chemist Svante Arrhenius, who observed that acids and bases increase the electrical conductivity of water.

• Acids: Donate protons to water to form hydronium ions (\text{H}_3\text{O}^+). \text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^-

• Bases: Accept protons from water to form hydroxide ions (\text{OH}^-). \text{NH}_3 + \text{H}_2\text{O} \rightarrow \text{NH}_4^+ + \text{OH}^-

Brønsted-Lowry Theory (Proton Transfer Theory)

Expanded by English chemist Thomas Lowry and Danish chemist Johannes Brønsted to include proton transfer reactions between any compounds, regardless of whether water is present.

• Acids: Proton donors.

• Bases: Proton acceptors. \text{HCl} + \text{NH}_3 \rightarrow \text{NH}_4^+ + \text{Cl}^-

3. Conjugate Acid-Base Pairs

Many proton-transfer reactions are reversible, meaning the products can react in reverse to regenerate the starting reactants. Products in these reactions function as conjugate acids and conjugate bases.

• Conjugate Base: Formed when an acid donates a proton. It contains one less proton (\text{H}^+) than its corresponding acid.

• Conjugate Acid: Formed when a base accepts a proton. It contains one more proton (\text{H}^+) than its corresponding base.

• Key Principle: A conjugate pair always differs by exactly one proton (\text{H}^+).

Reaction Examples

Example 1: Ethanoic Acid in Water

\text{CH}_3\text{COOH}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightleftharpoons \text{CH}_3\text{COO}^-_{(\text{aq})} + \text{H}_3\text{O}^+_{(\text{aq})}

• Acid: \text{CH}_3\text{COOH} (Ethanoic acid) \rightarrow Conjugate Base: \text{CH}_3\text{COO}^- (Ethanoate ion)

• Base: \text{H}_2\text{O} (Water) \rightarrow Conjugate Acid: \text{H}_3\text{O}^+ (Hydronium ion)

Example 2: Ammonia in Water

\text{NH}_{3(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightleftharpoons \text{NH}_{4(\text{aq})}^+ + \text{OH}^-_{(\text{aq})}

• Base: \text{NH}_3 (Ammonia) \rightarrow Conjugate Acid: \text{NH}_4^+ (Ammonium ion)

• Acid: \text{H}_2\text{O} (Water) \rightarrow Conjugate Base: \text{OH}^- (Hydroxide ion)

4. Acid and Base Strength vs. Concentration

In chemistry, strength and concentration represent two distinct properties:

• Strength: Refers to the degree of ionisation a substance undergoes when reacted with water.

• Concentration: Refers to the number of moles of acid or base dissolved per litre of solution (\text{mol L}^{-1}).

Classifications of Strength

Case Study: Concentrated Weak Acid

Hydrofluoric acid (\text{HF}) serves as an example of a concentrated, weak acid. It can exist at a very high molar concentration (28\text{ mol L}^{-1}), yet it undergoes only partial ionisation in solution.

5. Summary of Common Acids and Bases

6. Monoprotic and Polyprotic Acids

Acids are categorized by the number of protons available for donation per molecule during a chemical reaction.

Monoprotic Acids

Acids that transfer one proton per molecule to a base.

• Examples: Hydrochloric acid (\text{HCl}), Ethanoic acid (\text{CH}_3\text{COOH}).

Polyprotic Acids

Acids that can transfer more than one proton per molecule to a base. Polyprotic ionisation occurs in step-by-step stages.

Diprotic Acids

Can transfer up to two protons. Ionisation occurs in two stages.

• Examples: Sulfuric acid (\text{H}_2\text{SO}_4), Carbonic acid (\text{H}_2\text{CO}_3).

Ionisation stages of Sulfuric Acid (\text{H}_2\text{SO}_4):

1. Stage 1: \text{H}_2\text{SO}_{4(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightarrow \text{HSO}^-_{4(\text{aq})} + \text{H}_3\text{O}^+_{(\text{aq})}

2. Stage 2: \text{HSO}^-_{4(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightleftharpoons \text{SO}^{2-}_{4(\text{aq})} + \text{H}_3\text{O}^+_{(\text{aq})}

• Single Summarised Balanced Equation: \text{H}_2\text{SO}_{4(\text{aq})} + 2\text{H}_2\text{O}_{(\text{l})} \rightarrow \text{SO}^{2-}_{4(\text{aq})} + 2\text{H}_3\text{O}^+_{(\text{aq})}

Triprotic Acids

Can transfer up to three protons. Ionisation occurs in three stages.

• Examples: Phosphoric acid (\text{H}_3\text{PO}_4), Citric acid (\text{C}_6\text{H}_8\text{O}_7).

7. Acid-Base Indicators

Acid-base indicators are weak acids or bases composed of conjugate acid-base pairs. The acidic form (\text{HIn}) and basic form (\text{In}^-) exhibit different colours because they reflect different wavelengths of visible light depending on the solution's \text{pH} (concentration of \text{H}_3\text{O}^+ ions).

Phenolphthalein Mechanics

The molecular formula of phenolphthalein is simplified as \text{HIn}.

• In Acidic Solutions: Phenolphthalein exists predominantly as \text{HIn}, which is colourless because the molecule does not emit visible wavelengths of light. \text{In}^- + \text{H}^+ \rightarrow \text{HIn} \quad (\text{deep pink} \rightarrow \text{colourless})

• In Basic Solutions: A proton is transferred from the indicator (\text{HIn}) to the base, forming an anion (\text{In}^-) that reflects red and violet wavelengths of light, giving the solution a pink/magenta colour. \text{HIn} + \text{OH}^- \rightarrow \text{In}^- + \text{H}_2\text{O} \quad (\text{colourless} \rightarrow \text{deep pink})

• Titration Observations:

  1. The acidic solution is initially clear.

  2. As base is added, the solution turns pink momentarily, disappearing with swirling.

  3. The first permanent pink colour indicates the titration endpoint.

  4. Beyond the equivalence point, excess base causes a vivid dark pink/magenta colour.

Topic 5.2: Reactions of Acids and Bases — Comprehensive Study Guide

1. Fundamental Acid-Base Concepts & Theories

Core Definitions

• Proton (\text{H}^+): A positively charged hydrogen atom. A hydrogen atom consists of one proton and one electron; upon losing its electron, only the proton remains.

• Hydronium Ion (\text{H}_3\text{O}^+): Formed when a water molecule accepts a proton (\text{H}^+) from an acid.

Historical and Theoretical Models

1. Arrhenius Theory (Ionisation in Water):

   • Swedish chemist Svante Arrhenius identified that acids and bases increase the electrical conductivity of water.

   • Acids donate protons to water to form hydronium ions (\text{H}_3\text{O}^+). \text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^-

   • Bases accept protons from water to form hydroxide ions (\text{OH}^-). \text{NH}_3 + \text{H}_2\text{O} \rightarrow \text{NH}_4^+ + \text{OH}^-

2. Brønsted-Lowry Theory (Proton Transfer Theory):

   • English chemist Thomas Lowry and Danish chemist Johannes Brønsted expanded acid-base definitions to include proton transfer between any compounds, not limited to aqueous solutions.

   • Acids are proton donors.

   • Bases are proton acceptors. \text{HCl} + \text{NH}_3 \rightarrow \text{NH}_4^+ + \text{Cl}^-

3. Amphiprotic Nature of Water:

   • Water (\text{H}_2\text{O}) is an amphiprotic substance, meaning it can act as both a proton acceptor (base) and a proton donor (acid).

   • Reactions where a molecule reacts with water to produce ions are termed ionisation reactions.

Conjugate Acid-Base Pairs

• Many proton-transfer reactions are reversible, allowing products to revert back to reactants.

• Conjugate pairs differ from one another by one proton (\text{H}^+):

  • Acids donate a proton to form a conjugate base (which has 1 fewer \text{H}^+).

  • Bases accept a proton to form a conjugate acid (which has 1 more \text{H}^+).

Examples of Conjugate Pairs:

• Example 1: \text{CH}_3\text{COOH}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightleftharpoons \text{CH}_3\text{COO}^-_{(\text{aq})} + \text{H}_3\text{O}^+_{(\text{aq})}

  • Acid: \text{CH}_3\text{COOH} \rightarrow Conjugate Base: \text{CH}_3\text{COO}^-

  • Base: \text{H}_2\text{O} \rightarrow Conjugate Acid: \text{H}_3\text{O}^+

• Example 2: \text{NH}_{3(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightleftharpoons \text{NH}_{4(\text{aq})}^+ + \text{OH}^-_{(\text{aq})}

  • Base: \text{NH}_3 \rightarrow Conjugate Acid: \text{NH}_4^+

  • Acid: \text{H}_2\text{O} \rightarrow Conjugate Base: \text{OH}^-

2. Acid and Base Strength vs Concentration

Distinguishing Strength from Concentration

• Strength: Refers to the degree of ionisation a substance undergoes when reacted with water.

• Concentration: Refers to the number of moles of acid or base present per litre of solution (\text{mol L}^{-1}).

• Note: A weak acid can exist in high concentration (e.g. hydrofluoric acid at 28\text{ mol L}^{-1} is concentrated, yet remains a weak acid due to partial ionisation).

Classification of Acid and Base Strength

• Strong Acids: Ionise completely in water. One mole of a strong monoprotic acid yields one mole of hydronium ions. The conjugate base of a strong acid is weak (e.g. \text{Cl}^- is a weak conjugate base of \text{HCl}).

• Weak Acids: Ionise partially in water. One mole of a weak acid yields less than one mole of hydronium ions.

• Strong Bases: Ionise/dissociate completely in water to produce one mole of hydroxide ions per mole of base.

• Weak Bases: Ionise partially in water to produce less than one mole of hydroxide ions per mole of base.

Polyprotic Acids

Acids are categorized based on the number of protons available for donation:

1. Monoprotic Acids: Transfer one proton per molecule (e.g. \text{HCl}, \text{CH}_3\text{COOH}).

2. Polyprotic Acids: Transfer more than one proton per molecule.

   • Diprotic Acids: Transfer up to two protons in two sequential stages (e.g. \text{H}_2\text{SO}_4, \text{H}_2\text{CO}_3).

     • Stage 1: \text{H}_2\text{SO}_{4(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} \rightarrow \text{HSO}_{4(\text{aq})}^- + \text{H}_3\text{O}^+_{(\text{aq})}

     • Stage 2: \text{HSO}_{4(\text{aq})}^- + \text{H}_2\text{O}_{(\text{l})} \rightleftharpoons \text{SO}_{4(\text{aq})}^{2-} + \text{H}_3\text{O}^+_{(\text{aq})}

     • Overall Summary Equation: \text{H}_2\text{SO}_{4(\text{aq})} + 2\text{H}_2\text{O}_{(\text{l})} \rightarrow \text{SO}_{4(\text{aq})}^{2-} + 2\text{H}_3\text{O}^+_{(\text{aq})}

   • Triprotic Acids: Transfer up to three protons in three sequential stages (e.g. phosphoric acid \text{H}_3\text{PO}_4, citric acid \text{C}_6\text{H}_8\text{O}_7).

3. Acid-Base Indicators

Principles of Indicators

• Acid-base indicators are weak acids or bases composed of molecules of a weak acid and its conjugate base pair.

• They change colour based on the \text{pH} (relative measure of acidity or basicity) of the solution.

• The acidic form (\text{HIn}) and basic form (\text{In}^-) display different colours because they reflect different wavelengths of visible light.

Phenolphthalein Mechanism

• Formula simplified to \text{HIn}.

• In Acidic Solution: Colourless (\text{HIn} does not emit visible wavelengths of light). \text{In}^- + \text{H}^+ \rightarrow \text{HIn} \quad (\text{deep pink} \rightarrow \text{colourless})

• In Basic Solution: Proton is transferred from the indicator to the base, forming an anion (\text{In}^-) that reflects red and violet wavelengths, producing a pink/magenta colour. \text{HIn} + \text{OH}^- \rightarrow \text{In}^- + \text{H}_2\text{O} \quad (\text{colourless} \rightarrow \text{deep pink})

Indicator Colour Summary

4. Classification and Reactions of Oxides

Oxides are categorized as acidic, basic, or amphoteric based on their chemical behaviour with water, acids, and bases.

1. Acidic Oxides (Non-Metal Oxides)

• Non-metal oxides react chemically with water to form oxyacids.

• They react with bases to form a salt (oxyanion) and water.

Reaction Examples (\text{SO}_2):

• Reaction with water: \text{SO}_{2(\text{g})} + \text{H}_2\text{O}_{(\text{l})} \rightarrow \text{H}_2\text{SO}_{3(\text{aq})} \quad (\text{sulphurous acid})

• Reaction with base (full equation): \text{SO}_{2(\text{g})} + 2\text{NaOH}_{(\text{aq})} \rightarrow \text{Na}_2\text{SO}_{3(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

• Complete ionic equation: \text{SO}_{2(\text{g})} + 2\text{Na}^+_{(\text{aq})} + 2\text{OH}^-_{(\text{aq})} \rightarrow 2\text{Na}^+_{(\text{aq})} + \text{SO}_{3(\text{aq})}^{2-} + \text{H}_2\text{O}_{(\text{l})}

• Net ionic equation: \text{SO}_{2(\text{g})} + 2\text{OH}^-_{(\text{aq})} \rightarrow \text{SO}_{3(\text{aq})}^{2-} + \text{H}_2\text{O}_{(\text{l})}

Common Non-Metal Oxides, Oxyanions, and Oxyacids:

2. Basic Oxides (Metal Oxides)

• Metal oxides react with water to form hydroxide ions (\text{OH}^-).

• They react with acidic substances to produce a salt and water.

Reaction Examples (\text{MgO}):

• Reaction with water: \text{MgO}_{(\text{s})} + \text{H}_2\text{O}_{(\text{l})} \rightarrow \text{Mg}^{2+}_{(\text{aq})} + 2\text{OH}^-_{(\text{aq})}

• Reaction with acid (full equation): \text{MgO}_{(\text{s})} + 2\text{HCl}_{(\text{aq})} \rightarrow \text{MgCl}_{2(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

• Complete ionic equation: \text{MgO}_{(\text{s})} + 2\text{H}^+_{(\text{aq})} + 2\text{Cl}^-_{(\text{aq})} \rightarrow \text{Mg}^{2+}_{(\text{aq})} + 2\text{Cl}^-_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

• Net ionic equation: \text{MgO}_{(\text{s})} + 2\text{H}^+_{(\text{aq})} \rightarrow \text{Mg}^{2+}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

• Other common basic oxides: \text{Na}_2\text{O}, \text{K}_2\text{O}, \text{CaO}.

3. Amphoteric Oxides

Amphoteric oxides exhibit both acidic and basic properties and can neutralise both acids and bases. Examples include lead oxide, aluminium oxide (\text{Al}_2\text{O}_3), and zinc oxide (\text{ZnO}).

Behaviour of Aluminium Oxide (\text{Al}_2\text{O}_3):

1. Acting as a Base (Reacting with Acid):

   • Full equation: \text{Al}_2\text{O}_{3(\text{s})} + 6\text{HCl}_{(\text{aq})} \rightarrow 2\text{AlCl}_{3(\text{aq})} + 3\text{H}_2\text{O}_{(\text{l})}

   • Complete ionic equation: \text{Al}_2\text{O}_{3(\text{s})} + 6\text{H}^+_{(\text{aq})} + 6\text{Cl}^-_{(\text{aq})} \rightarrow 2\text{Al}^{3+}_{(\text{aq})} + 6\text{Cl}^-_{(\text{aq})} + 3\text{H}_2\text{O}_{(\text{l})}

   • Net ionic equation: \text{Al}_2\text{O}_{3(\text{s})} + 6\text{H}^+_{(\text{aq})} \rightarrow 2\text{Al}^{3+}_{(\text{aq})} + 3\text{H}_2\text{O}_{(\text{l})}

2. Acting as an Acid (Reacting with Base):

   • Form sodium aluminate (\text{NaAlO}_2) and water.

   • Full equation: \text{Al}_2\text{O}_{3(\text{s})} + 2\text{NaOH}_{(\text{aq})} \rightarrow 2\text{NaAlO}_{2(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

   • Complete ionic equation: \text{Al}_2\text{O}_{3(\text{s})} + 2\text{Na}^+_{(\text{aq})} + 2\text{OH}^-_{(\text{aq})} \rightarrow 2\text{Na}^+_{(\text{aq})} + 2\text{AlO}_{2(\text{aq})}^- + \text{H}_2\text{O}_{(\text{l})}

   • Net ionic equation: \text{Al}_2\text{O}_{3(\text{s})} + 2\text{OH}^-_{(\text{aq})} \rightarrow 2\text{AlO}_{2(\text{aq})}^- + \text{H}_2\text{O}_{(\text{l})}

Behaviour of Zinc Oxide (\text{ZnO}):

• When acting as an acid, zinc oxide forms the zincate oxyanion (\text{ZnO}_2^{2-}).

5. Thermochemistry of Neutralisation Reactions

Definition & Enthalpy

• An acid-base reaction is classified as a neutralisation reaction.

• If 1 mole of a strong acid reacts with 1 mole of a strong base, the resulting solution is neutral.

• Protons (\text{H}^+) combine with hydroxide ions (\text{OH}^-) to form liquid water: \text{H}^+_{(\text{aq})} + \text{OH}^-_{(\text{aq})} \rightarrow \text{H}_2\text{O}_{(\text{l})}

• Thermochemical Behaviour: Neutralisation is exothermic because chemical bonds are being formed. It releases 55.9\text{ kJ} of energy per mole of water produced, carrying a negative enthalpy value (\Delta H).

\text{H}^+_{(\text{aq})} + \text{OH}^-_{(\text{aq})} \rightarrow \text{H}_2\text{O}_{(\text{l})} \quad \Delta H = -55.9\text{ kJ mol}^{-1}

6. Acid Reaction Types & Equation Writing Guide

When writing complete and net ionic equations:

• Soluble strong electrolytes (strong acids, soluble salts, strong soluble bases) are written as dissociated free ions.

• Solids, liquids, gases, weak acids, and weak bases are not split into ions.

Reaction Type 1: Acid + Metal Oxide

\text{Acid} + \text{Metal Oxide} \rightarrow \text{Salt} + \text{Water}

• Example: Sodium oxide + Sulfuric acid

  • Full Equation: \text{Na}_2\text{O}_{(\text{s})} + \text{H}_2\text{SO}_{4(\text{aq})} \rightarrow \text{Na}_2\text{SO}_{4(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

  • Full Ionic Equation: \text{Na}_2\text{O}_{(\text{s})} + 2\text{H}^+_{(\text{aq})} + \text{SO}_{4(\text{aq})}^{2-} \rightarrow 2\text{Na}^+_{(\text{aq})} + \text{SO}_{4(\text{aq})}^{2-} + \text{H}_2\text{O}_{(\text{l})}

  • Net Ionic Equation: \text{Na}_2\text{O}_{(\text{s})} + 2\text{H}^+_{(\text{aq})} \rightarrow 2\text{Na}^+_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

Reaction Type 2: Acid + Metal Hydroxide

\text{Acid} + \text{Metal Hydroxide} \rightarrow \text{Salt} + \text{Water}

A. Soluble Hydroxide + Strong Acid

• Example: Potassium hydroxide + Hydrochloric acid

  • Full Equation: \text{HCl}_{(\text{aq})} + \text{KOH}_{(\text{aq})} \rightarrow \text{KCl}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

  • Full Ionic Equation: \text{H}^+_{(\text{aq})} + \text{Cl}^-_{(\text{aq})} + \text{K}^+_{(\text{aq})} + \text{OH}^-_{(\text{aq})} \rightarrow \text{K}^+_{(\text{aq})} + \text{Cl}^-_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

  • Net Ionic Equation: \text{H}^+_{(\text{aq})} + \text{OH}^-_{(\text{aq})} \rightarrow \text{H}_2\text{O}_{(\text{l})}

  • Spectator Ions: \text{K}^+, \text{Cl}^-

B. Insoluble Hydroxide + Strong Acid

• Example: Calcium hydroxide + Hydrochloric acid

  • Full Equation: 2\text{HCl}_{(\text{aq})} + \text{Ca(OH)}_{2(\text{s})} \rightarrow \text{CaCl}_{2(\text{aq})} + 2\text{H}_2\text{O}_{(\text{l})}

  • Full Ionic Equation: 2\text{H}^+_{(\text{aq})} + 2\text{Cl}^-_{(\text{aq})} + \text{Ca(OH)}_{2(\text{s})} \rightarrow \text{Ca}^{2+}_{(\text{aq})} + 2\text{Cl}^-_{(\text{aq})} + 2\text{H}_2\text{O}_{(\text{l})}

  • Net Ionic Equation: 2\text{H}^+_{(\text{aq})} + \text{Ca(OH)}_{2(\text{s})} \rightarrow \text{Ca}^{2+}_{(\text{aq})} + 2\text{H}_2\text{O}_{(\text{l})}

C. Soluble Strong Base + Weak Acid

• Example: Sodium hydroxide + Ethanoic acid (Note: Ethanoic acid is weak and only partially ionises, so it cannot be separated into ions in the ionic equation).

  • Full Equation: \text{CH}_3\text{COOH}_{(\text{aq})} + \text{NaOH}_{(\text{aq})} \rightarrow \text{CH}_3\text{COONa}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

  • Full Ionic Equation: \text{CH}_3\text{COOH}_{(\text{aq})} + \text{Na}^+_{(\text{aq})} + \text{OH}^-_{(\text{aq})} \rightarrow \text{CH}_3\text{COO}^-_{(\text{aq})} + \text{Na}^+_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

  • Net Ionic Equation: \text{CH}_3\text{COOH}_{(\text{aq})} + \text{OH}^-_{(\text{aq})} \rightarrow \text{CH}_3\text{COO}^-_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

Reaction Type 3: Acid + Metal Carbonate

\text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}

• Example: Calcium carbonate + Nitric acid

  • Full Equation: 2\text{HNO}_{3(\text{aq})} + \text{CaCO}_{3(\text{s})} \rightarrow \text{Ca(NO}_3)_{2(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} + \text{CO}_{2(\text{g})}

  • Full Ionic Equation: 2\text{H}^+_{(\text{aq})} + 2\text{NO}_{3(\text{aq})}^- + \text{CaCO}_{3(\text{s})} \rightarrow \text{Ca}^{2+}_{(\text{aq})} + 2\text{NO}_{3(\text{aq})}^- + \text{H}_2\text{O}_{(\text{l})} + \text{CO}_{2(\text{g})}

  • Net Ionic Equation: 2\text{H}^+_{(\text{aq})} + \text{CaCO}_{3(\text{s})} \rightarrow \text{Ca}^{2+}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} + \text{CO}_{2(\text{g})}

Reaction Type 4: Acid + Metal Hydrogencarbonate

\text{Acid} + \text{Metal Hydrogencarbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}

• Example: Sodium hydrogencarbonate + Hydrochloric acid

  • Full Equation: \text{HCl}_{(\text{aq})} + \text{NaHCO}_{3(\text{s})} \rightarrow \text{NaCl}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} + \text{CO}_{2(\text{g})}

  • Full Ionic Equation: \text{H}^+_{(\text{aq})} + \text{Cl}^-_{(\text{aq})} + \text{NaHCO}_{3(\text{s})} \rightarrow \text{Na}^+_{(\text{aq})} + \text{Cl}^-_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} + \text{CO}_{2(\text{g})}

  • Net Ionic Equation: \text{H}^+_{(\text{aq})} + \text{NaHCO}_{3(\text{s})} \rightarrow \text{Na}^+_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})} + \text{CO}_{2(\text{g})}

Reaction Type 5: Acid + Reactive Metal (Redox Reaction)

\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen Gas}

• Example: Magnesium + Nitric acid

  • Full Equation: 2\text{HNO}_{3(\text{aq})} + \text{Mg}_{(\text{s})} \rightarrow \text{Mg(NO}_3)_{2(\text{aq})} + \text{H}_{2(\text{g})}

  • Full Ionic Equation: 2\text{H}^+_{(\text{aq})} + 2\text{NO}_{3(\text{aq})}^- + \text{Mg}_{(\text{s})} \rightarrow \text{Mg}^{2+}_{(\text{aq})} + 2\text{NO}_{3(\text{aq})}^- + \text{H}_{2(\text{g})}

  • Net Ionic Equation: 2\text{H}^+_{(\text{aq})} + \text{Mg}_{(\text{s})} \rightarrow \text{Mg}^{2+}_{(\text{aq})} + \text{H}_{2(\text{g})}

7. Stoichiometric Calculations Involving Neutralisation Reactions

Quantitative Laboratory Volumetric Analysis

• In chemical laboratories, neutralisation reactions are utilized to determine the unknown concentration of an acidic or basic solution via volumetric titration.

1. A standard solution of known concentration is prepared.

2. The standard solution is reacted with a second solution of unknown concentration.

3. The exact volume required for complete neutralisation is recorded.

4. Stoichiometric ratios from the balanced chemical equation are applied to compute the unknown concentration.

Core Stoichiometric Equation Model:

\text{NaOH}_{(\text{aq})} + \text{HCl}_{(\text{aq})} \rightarrow \text{NaCl}_{(\text{aq})} + \text{H}_2\text{O}_{(\text{l})}

• Molar Ratio (\text{NaOH} : \text{HCl}) = 1 : 1.