Comprehensive Notes on Symbols, Formulae, Ionic & Covalent Compounds, and Chemical Equations

Symbols for Elements and Their Molecular Forms

Metallic and Most Non-Metallic Elements

• Metals and common non-metals (B, C, Si, P, S, all Group 0) are represented by their single/two-letter symbol in equations (e.g., $\text{Na},\; \text{Al},\; \text{C}$).

• A-level note: While P and S exist as $\text{P}<em>4$ and $\text{S}</em>8$, writing $\text{P}$ and $\text{S}$ is acceptable at this level.

Diatomic Elements

• Seven elements are diatomic and use a subscript 2: $\text{H}<em>2,\; \text{N}</em>2,\; \text{O}<em>2,\; \text{F}</em>2,\; \text{Cl}<em>2,\; \text{Br}</em>2,\; \text{I}_2$ .

• Exam tip: Use a mnemonic to memorize these 7.

Ionic Formulae: Main-Group Ions

Variable-Charge (Transition-Metal & Post-Transition) Ions

• Transition metals (between Groups 2 and 3) have charges indicated by Roman numerals in brackets:

  • $\text{Fe}^{3+}$ (iron(III) oxide)

  • $\text{Cu}^{2+}$ (copper(II) chloride)

  • $\text{Ag}^+$ (silver(I) nitrate, often just 'silver nitrate').

• Some p-block metals (e.g., Pb, Sn) also show multiple charges, e.g., $\text{Pb}^{2+}$ in lead(II).

• Default: If no Roman numeral is given, assume $2+$ (except Ag, which is $1+$).

Molecular (Polyatomic) Ions You Must Know

Chemical Formulae Rules

1. Subscripts indicate ion quantities; omit '1'.

2. Net charge must be zero (positive charge = negative charge).

3. For multiple polyatomic ions, use brackets: e.g., $\text{Ca(OH)}_2$ .

4. Covalent formulae are memorized, not charge-balanced.

Worked Examples (Building Formulae)

| # | Compound Name | Ions Involved | Charge Balance | Final Formula | Notes |

| : | :------------ | :------------ | :------------- | :------------ | :---- |

| | | | | | |

| 1 | sodium chloride | $\text{Na}^+, \text{Cl}^-$ | $1+$ vs $1-$ | $\text{NaCl}$ | |

| 2 | magnesium oxide | $\text{Mg}^{2+}, \text{O}^{2-}$ | $2+$ vs $2-$ | $\text{MgO}$ | |

| 3 | calcium fluoride | $\text{Ca}^{2+}, \text{F}^-$ | need 2 F$-$ | $\text{CaF}<em>2$ | | | 4 | aluminium oxide | $\text{Al}^{3+}, \text{O}^{2-}$ | LCM(3,2)=6 | $\text{Al}</em>2\text{O}<em>3$ | total $6+$ and $6-$ | | 5 | lithium hydroxide | $\text{Li}^+, \text{OH}^-$ | $1+$ vs $1-$ | $\text{LiOH}$ | no brackets needed | | 6 | copper(II) hydroxide | $\text{Cu}^{2+}, \text{OH}^-$ | need 2 OH$-$ | $\text{Cu(OH)}</em>2$ | brackets required |

| 7 | silver nitrate | $\text{Ag}^+, \text{NO}<em>3^-$ | $1+$ vs $1-$ | $\text{AgNO}</em>3$ | |

| 8 | chromium(III) sulfate | $\text{Cr}^{3+}, \text{SO}<em>4^{2-}$ | need 2 Cr$^{3+}$ & 3 SO$_4^{2-}$ | $\text{Cr}</em>2(\text{SO}<em>4)</em>3$ | $6+$ vs $6-$ |

| 9 | ammonium carbonate | $\text{NH}<em>4^+, \text{CO}</em>3^{2-}$ | need 2 NH$_4^+$ | $(\text{NH}<em>4)</em>2\text{CO}<em>3$ | | | 10 | copper(II) sulfate | $\text{Cu}^{2+}, \text{SO}</em>4^{2-}$ | $2+$ vs $2-$ | $\text{CuSO}_4$ | |

Essential Covalent Compounds to Memorise

Exam tip: HCl(g) and HCl(aq) have identical formulae but different names and properties.

Counting Atoms in Formulae

• Example: $\text{LiOH}$ has 1 Li, 1 O, 1 H (total 3 atoms).

• Example: $\text{Fe(NO}<em>3)</em>3$ has 1 Fe, 3 N, 9 O (total 13 atoms).

Balanced Symbol Equations

• Mass is conserved: no atoms created or destroyed.

• Layout: $\text{Reactants} \;\longrightarrow\; \text{Products}$ .

Exam mark rules:

1. Never change chemical formulae to balance equations.

2. Use only whole-number coefficients before formulae.

3. A coefficient multiplies all atoms in its formula.

4. Equations needing coefficients earn 3 marks; those naturally balanced by formulae earn 2 marks.

5. Adding state symbols (s, l, g, aq) adds 1 mark; a 2-mark equation becomes 3 marks with states.

State Symbols

• $(s)$ – solid

• $(l)$ – liquid

• $(g)$ – gas

• $(aq)$ – aqueous (dissolved in water)

Dilute mineral acids and soluble salts are typically $(aq)$.

Example – dynamite explosion (trinitrotoluene approximation):

$\text{2}\,\text{C}<em>7\text{H}</em>5\text{N}<em>3\text{O}</em>6(s) \;\rightarrow\;\; \text{7}\,\text{CO}(g) + \text{7}\,\text{C}(s) + \text{5}\,\text{H}<em>2\text{O}(g) + \text{3}\,\text{N}</em>2(g)$

Rapid gas formation and heat release create blast force.

Ionic Equations & Half-Equations

Spectator Ions & Net Ionic Equations

• Ionic equations focus on changing species only.

• Universal acid–base neutralisation:

  $\text{H}^+ + \text{OH}^- \;\rightarrow\; \text{H}_2\text{O}$

  All ions except $\text{H}^+$ and $\text{OH}^-$ are spectator ions and omitted.

Displacement Example

Full reaction: $\text{Mg} + \text{CuSO}_4 \;\rightarrow\; \text{MgSO}_4 + \text{Cu}$ .

Species present:

• Reactants: $\text{Mg}(s),\; \text{Cu}^{2+}(aq),\; \text{SO}_4^{2-}(aq)$

• Products: $\text{Mg}^{2+}(aq),\; \text{SO}_4^{2-}(aq),\; \text{Cu}(s)$

Spectator ion: $\text{SO}_4^{2-}$ (unchanged). Net ionic equation:

$\text{Mg}(s) + \text{Cu}^{2+}(aq) \;\rightarrow\; \text{Mg}^{2+}(aq) + \text{Cu}(s)$

Half-Equations (Electron Accounting)

1. Magnesium oxidation:

   $\text{Mg} \;\rightarrow\; \text{Mg}^{2+} + 2e^-$

2. Copper(II) reduction:

   $\text{Cu}^{2+} + 2e^- \;\rightarrow\; \text{Cu}$

Lost electrons in one half-equation are gained in the other, balancing charge and showing redox.

Ethical, Practical & Exam-Technique Notes

• Memorizing correct formulae prevents errors in balancing, percentage-yield, titration, and redox calculations.

• Mnemonics/flashcards help memorize diatomic molecules/polyatomic ions.

• Always include state symbols in past-paper questions for full marks.

• Relate ionic equations to real-world processes: e