Deduce the symbol equation with state symbols for a chemical reaction, given relevant information

Writing Symbol Equations with State Symbols (IGCSE Chemistry 0620)

Learning Objectives (Assessment Objectives)

• AO1 – Knowledge & Understanding: Identify reactants and products, write correct chemical formulas (including ionic formulas), assign appropriate state symbols and recognise reaction types (neutralisation, redox, precipitation, combustion, etc.).
• AO2 – Application: Balance symbol equations, use the mole concept to relate masses, moles, particles, gas volumes and to perform limiting‑reactant, percentage‑yield and stoichiometric calculations.
• AO3 – Analysis & Evaluation: Interpret experimental observations (colour change, gas evolution, precipitate), assign oxidation numbers, discuss energy changes (ΔH, activation energy) and evaluate environmental relevance.

1. States of Matter & Diffusion

In the IGCSE syllabus the four physical states are distinguished by state symbols. The symbols are written in brackets immediately after the formula.

Why gases diffuse rapidly: In a gas the particles are far apart and move freely at high speeds (kinetic‑particle‑theory). This explains the characteristic “bubbling” or “effervescence” observed when a gas is produced in a reaction.

2. Core Concepts for Symbol Equations

• Chemical formulas represent the composition of each substance. Ionic formulas are written using the charges of the constituent ions (e.g. Ca²⁺ + CO₃²⁻ → CaCO₃).
• State symbols show the physical state under the experimental conditions.
• Balancing is achieved by adjusting coefficients (never subscripts) so that the number of atoms of each element (and the total charge for ionic equations) is the same on both sides.
• Arrow notation:
  • → : irreversible reaction (most IGCSE reactions)
  • ⇌ or ↔ : reversible reaction (equilibrium)
• Oxidation numbers help identify redox processes. Remember the common rules (O = –2, H = +1 except in metal hydrides, halogens = –1 unless combined with a more electronegative element, etc.).
• Acid–base definitions (AO1):
  • Acid = proton (H⁺) donor (Bronsted‑Lowry) or H⁺ source (Arrhenius).
  • Base = proton acceptor or OH⁻ source.
  • Strong acids/bases dissociate completely in water; weak acids/bases only partially dissociate.
• Solubility rules (quick reference) – essential for predicting precipitates in ionic equations:
  • All nitrates (NO₃⁻), acetates (CH₃COO⁻) and alkali‑metal salts are soluble.
  • Most chlorides, bromides, iodides are soluble, except those of Ag⁺, Pb²⁺, Hg₂²⁺.
  • Sulfates are soluble except those of Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺, Hg₂²⁺.
  • Carbonates, phosphates, sulfides, hydroxides are generally insoluble except those of alkali metals and NH₄⁺.

3. Step‑by‑Step Procedure to Write a Balanced Symbol Equation

1. Read the experimental description and list every reactant and product, noting observable changes (colour change, gas evolution, precipitate, temperature change).
2. Write formulas for each substance. Use ionic notation for strong electrolytes (e.g. Na⁺, Cl⁻) when a full ionic equation is required.
3. Assign state symbols based on the description or typical behaviour (solid metal, liquid water, gas released, aqueous solution).
4. Balance the equation:
   • Start with the most complex species.
   • Adjust coefficients, never subscripts.
   • For ionic equations first balance the total charge (charge‑balance method), then balance atoms.
5. Check that atoms of every element and, for ionic equations, the total charge are equal on both sides.
6. Write the final equation in a boxed format, including state symbols and the correct arrow.

4. Stoichiometry & the Mole Concept (Supplementary)

Once the symbol equation is balanced you can use it for quantitative work.

Typical calculations:

• Limiting‑reactant: Use the balanced coefficients to compare the mole ratios of the reactants actually present.
• Percentage yield: \(\%\; \text{yield}= \dfrac{\text{actual mass of product}}{\text{theoretical mass}} × 100\).
• Gas volume: Convert moles of a gaseous product to volume using the 24 dm³ mol⁻¹ value (unless the question specifies other conditions).

5. Energy Changes & Bond‑Energy Calculations (Supplementary)

• Enthalpy change (ΔH):
  • ΔH < 0 – exothermic (heat released, temperature rise).
  • ΔH > 0 – endothermic (heat absorbed, temperature fall).
• Activation energy (E_a): Minimum energy required for reactant particles to collide effectively. Catalysts lower E_a but do not appear in the symbol equation.
• Bond‑energy method (optional for the supplement):
  1. Write the balanced equation.
  2. Sum the bond energies of all bonds broken (positive sign).
  3. Sum the bond energies of all bonds formed (negative sign).
  4. ΔH ≈ Σ(bonds broken) – Σ(bonds formed).

6. Redox, Half‑Equations & Electrolysis (Supplementary)

Redox reactions involve transfer of electrons. For IGCSE you should be able to:

• Assign oxidation numbers to all atoms.
• Identify the oxidising agent (gains electrons) and the reducing agent (loses electrons).
• Write the overall balanced redox equation.
• Write separate half‑equations (useful for electrolysis questions).

Electrolysis reminder (AO1):

• Anode – positive electrode – oxidation occurs.
• Cathode – negative electrode – reduction occurs.
• Inert electrodes (Pt, graphite) are used when neither reactant is a metal.
• Overall cell reaction = sum of the two half‑equations, electrons cancel.

7. Reversible Reactions & Equilibrium (Supplementary)

• Write reversible reactions with a double arrow (⇌) and include state symbols for all species.
• Equilibrium concepts (Le Chatelier’s principle) are not required for writing the equation, but you may be asked to predict the direction of shift when conditions change.

8. Worked Examples

Example 1 – Metal + Solid Carbon Dioxide (Redox)

Experimental description: Magnesium ribbon is heated in a crucible containing dry ice (solid CO₂). A colourless gas is released and a white solid remains.

1. Reactants: Mg(s) and CO₂(s).
   Products: MgO(s) and CO(g).
2. Tentative equation:
   \(\text{Mg(s)} + \text{CO}_2\text{(s)} \rightarrow \text{MgO(s)} + \text{CO(g)}\)
3. Balance: atoms already balanced.
4. Final boxed equation:
   
   \(\boxed{\text{Mg(s)} + \text{CO}_2\text{(s)} \rightarrow \text{MgO(s)} + \text{CO(g)}}\)

Example 2 – Acid + Carbonate (Neutralisation, gas evolution)

Experimental description: Excess HCl(aq) is added to solid CaCO₃. Effervescence occurs and the solution becomes clear.

1. Reactants: CaCO₃(s) and HCl(aq).
   Products: CaCl₂(aq), H₂O(l) and CO₂(g).
2. Initial equation:
   \(\text{CaCO}_3\text{(s)} + \text{HCl(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}\)
3. Balance H and Cl (coefficient 2 before HCl).
   \(\text{CaCO}_3\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}\)
4. Final boxed equation:
   
   \(\boxed{\text{CaCO}_3\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}}\)

Example 3 – Metal + Dilute Acid (Redox)

Experimental description: Zinc metal is placed in dilute HCl(aq). Bubbles of colourless gas are observed.

1. Formulas with state symbols:
   \(\text{Zn(s)} + \text{HCl(aq)} \rightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)}\)
2. Oxidation numbers: Zn 0 → Zn²⁺ (oxidation); H⁺ +1 → H₂ 0 (reduction).
3. Balance charge: 2 HCl needed.
   \(\text{Zn(s)} + 2\text{HCl(aq)} \rightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)}\)
4. Final boxed equation:
   
   \(\boxed{\text{Zn(s)} + 2\text{HCl(aq)} \rightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)}}\)

Example 4 – Acid‑Base Neutralisation (Strong acid & strong base)

Experimental description: Sodium hydroxide solution is added to sulphuric acid solution. The mixture becomes warm.

1. Formulas:
   \(\text{NaOH(aq)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{Na}_2\text{SO}_4\text{(aq)} + \text{H}_2\text{O(l)}\)
2. Balance Na (coefficient 2 before NaOH).
   \(2\text{NaOH(aq)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{Na}_2\text{SO}_4\text{(aq)} + 2\text{H}_2\text{O(l)}\)
3. Final boxed equation:
   
   \(\boxed{2\text{NaOH(aq)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{Na}_2\text{SO}_4\text{(aq)} + 2\text{H}_2\text{O(l)}}\)

Example 5 – Precipitation (Ionic & Molecular equations)

Experimental description: Solutions of copper(II) sulfate and sodium hydroxide are mixed. A blue precipitate forms.

1. Molecular equation:
   \(\text{CuSO}_4\text{(aq)} + 2\text{NaOH(aq)} \rightarrow \text{Cu(OH)}_2\text{(s)} + \text{Na}_2\text{SO}_4\text{(aq)}\)
2. Full ionic equation (show all strong electrolytes dissociated):
   \(\text{Cu}^{2+}(aq) + \text{SO}_4^{2-}(aq) + 2\text{Na}^+(aq) + 2\text{OH}^-(aq) \rightarrow \text{Cu(OH)}_2\text{(s)} + 2\text{Na}^+(aq) + \text{SO}_4^{2-}(aq)\)
3. Net ionic equation (spectator ions removed):
   \(\text{Cu}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow \text{Cu(OH)}_2\text{(s)}\)

9. Links to Other Syllabus Topics

• Atomic structure & ions: Oxidation‑number rules rely on knowing the typical charges of common ions (e.g., Fe²⁺, Fe³⁺, NH₄⁺).
• Electrochemistry: Redox equations underpin half‑equation writing for electrolysis and fuel‑cell questions.
• Environmental chemistry: Reactions that produce CO₂(g), SO₂(g) or NH₃(g) are examined in the context of air‑pollution; always write the gas state symbol.
• Organic chemistry: Simple combustion reactions (e.g., C₃H₈ + 5 O₂ → 3 CO₂ + 4 H₂O) follow the same balancing rules and require (g) for the gaseous products.
• Physical chemistry: Enthalpy changes (ΔH) and activation energy are linked to the observations of temperature rise or fall during a reaction.

10. Practice Questions

1. Write the balanced symbol equation with state symbols for the reaction of sodium metal with chlorine gas to form sodium chloride.
2. Calcium reacts with water to produce calcium hydroxide and hydrogen gas. Include state symbols and balance the equation.
3. When copper(II) sulfate solution is mixed with sodium hydroxide solution, a blue precipitate forms and sodium sulfate remains in solution. Write the balanced molecular, full ionic and net ionic equations with state symbols.
4. Balance the redox equation and assign oxidation numbers: Fe(s) + H₂SO₄(aq) → Fe₂(SO₄)₃(aq) + H₂(g). Show the half‑equations.
5. Calculate the mass of carbon dioxide formed when 5.0 g of calcium carbonate reacts with excess hydrochloric acid (use Example 2). Show all steps, including limiting‑reactant check.
6. For the reaction in Example 3, calculate the volume of H₂(g) produced at r.t.p. when 0.50 mol of Zn reacts completely.
7. Write the overall cell reaction for the electrolysis of aqueous NaCl (use half‑equations). Include state symbols.

11. Quick‑Check Checklist

• ✔︎ Reactants and products identified from the description.
• ✔︎ Correct chemical formulas (including ionic charges where required).
• ✔︎ Appropriate state symbols attached.
• ✔︎ Equation balanced – atoms (and total charge for ionic equations) are equal.
• ✔︎ Arrow type chosen correctly (→ for irreversible, ⇌ for reversible).
• ✔︎ Oxidation numbers assigned for redox questions.
• ✔︎ Acid‑base definitions and solubility rules applied where needed.
• ✔︎ If required, half‑equations written for electrolysis or redox.

12. Suggested Diagram