Chapter 6 Notes: Chemical Formulae & Equations

6.1 Formulae of Ionic Compounds

Ionic compounds are formed from a cation (positive ion) and an anion (negative ion).
Cations (positive) and Anions (negative): common ions breaking down by groups
- Cations (Group 1, Group 2, Group 13):
- Group 1: Li, Na, K → ions: $ext{Li}^{+}, ext{Na}^{+}, ext{K}^{+}$
- Group 2: Be, Mg, Ca → ions: $ext{Be}^{2+}, ext{Mg}^{2+}, ext{Ca}^{2+}$
- Group 13: B, Al → ions: $ext{Al}^{3+}$
- Anions (Group 16 and 17):
- Group 16: O, S → oxide, sulfide ions: $ext{O}^{2-} ext{(oxide)}, ext{S}^{2-} ext{(sulfide)}$
- Group 17: F, Cl, Br, I → halide ions: $ext{F}^{-}, ext{Cl}^{-}, ext{Br}^{-}, ext{I}^{-}$
- Other common cations: $ext{Cu}^{+}, ext{Cu}^{2+}, ext{Sn}^{2+}, ext{Pb}^{2+}, ext{Pb}^{4+}, ext{Fe}^{2+}, ext{Fe}^{3+}, ext{H}^{+}, ext{Ag}^{+}, ext{Zn}^{2+}$
Transition metals can have varying charges; Roman numerals in brackets indicate the charge (except Zn and Ag which typically have fixed charges).
Polyatomic ions (ions composed of multiple atoms) include:
- Hydroxide: $ext{OH}^{-} \, (1^{-})$
- Nitrate: $ext{NO}_3^{-} \, (1^{-})$
- Sulfate: $ext{SO}_4^{2-} \, (2^{-})$
- Carbonate: $ext{CO}_3^{2-} \, (2^{-})$
- Ammonium: $ext{NH}_4^{+} \, (1^{+})$
Memorisation aid (shaded portions in original notes) for polyatomic ions:
- Hydroxide: OH−, 1− (Oh)
- Nitrate: NO3−, 1− (No)
- Sulfate: SO4^{2−}, 2− (So)
- Carbonate: CO3^{2−}, 2− (Co)
- Ammonium: NH4^{+}, 1+ (Naaaa..)
How to construct the chemical formula of an ionic compound
- An ionic compound is formed from a cation and an anion (often metal cation with a non-metal anion or polyatomic ions).
- The total charges must sum to zero: charges balance to neutral.
- Examples:
- Sodium chloride: $ext{Na}^{+} + ext{Cl}^{-} ightarrow ext{NaCl}$
 - Net charge: $(+1) + (-1) = 0$
- Aluminium carbonate: $2 ext{Al}^{3+} + 3 ext{CO}3^{2-} ightarrow ext{Al}2( ext{CO}3)3$
 - Net charge: $2(+3) + 3(-2) = 0$
Example practice interpretations:
- There are 3 carbonate ions for Al2(CO3)3 to balance the charge of 2 Al3+ ions.
- The overall formula must reflect the smallest whole-number ratio that makes the total charge zero.

6.1 Formulae of Ionic Compounds (Examples)

Example 1: NaCl
- Cation: Na⁺; Anion: Cl⁻ → Formula: $ext{NaCl}$
- Net charge: $+1 + (-1) = 0$
Example 2: Al2(CO3)3
- Cation: Al³⁺ (2 ions) → total positive charge: $2 imes (+3) = +6$
- Anion: CO₃²⁻ (3 ions) → total negative charge: $3 imes (-2) = -6$
- Formula: $ext{Al}2( ext{CO}3)_3$
Important reminder:
- For covalent and ionic compounds, different rules apply; ionic compounds balance charges, covalent compounds use prefixes and do not involve charged ions to balance.

6.2 Counting of Atoms

Counting atoms and ions in formulas and formula units
Representation rules:
- H = 1 hydrogen atom (1 H)
- 2H = 2 hydrogen atoms (2 H)
- H2 = 1 molecule of hydrogen (2 H atoms)
- 2H2 = 2 molecules of hydrogen (4 H atoms)
- NH3 = 1 formula unit of ammonia; 1 N atom and 3 H atoms
- 2NH3 = 2 formula units; 2 N atoms and 6 H atoms
- NaNO3 = 1 formula unit of sodium nitrate; 1 Na, 1 N, 3 O
- Ca(NO3)2 = 1 formula unit of calcium nitrate; 1 Ca, 2 N, 6 O
- 2Ca(NO3)2 = 2 formula units; 2 Ca, 4 N, 12 O
Practice questions (from WS 1):
- Determine the number of atoms/ions in given formula units.

6.2 Chemical Equations

Word equations can be written as chemical equations but must be balanced.
A chemical equation tells:
- Which reactants and products are involved.
- The relative amounts of reactants and products.
- The physical states of reactants and products (state symbols).
Common elements can be monatomic or diatomic in chemical equations:
- Noble gases (Group 18) are typically monatomic.
- Many common elements exist as diatomic molecules: H₂, N₂, O₂, Cl₂, etc.
Gases of elements example list (monatomic/diatomic):
- Neon (Ne) – Monatomic: $ext{Ne}$
- Helium (He) – Monatomic: $ext{He}$
- Argon (Ar) – Monatomic: $ext{Ar}$
- Hydrogen (H₂) – Diatomic: $ext{H}_2$
- Nitrogen (N₂) – Diatomic: $ext{N}_2$
- Oxygen (O₂) – Diatomic: $ext{O}_2$
- Chlorine (Cl₂) – Diatomic: $ext{Cl}_2$

6.3 State Symbols

State symbols indicate physical state in reactions:
- Solid: (s)
- Liquid: (l)
- Gas: (g)
- Aqueous: (aq)
Notes:
- Liquid state examples: melted sugar, etc.
- Aqueous state: substances dissolved in water (e.g., sugar solution).
Common examples in reactions:
- Soluble salts in water exist as ions in solution (aqueous).
- Insoluble salts or precipitates form as solids (s).
Practical reminder:
- State symbols are often indicated in ionic equations and balanced chemical equations.

6.3 Writing Balanced Equations with State Symbols

Steps to write a balanced chemical equation with state symbols
- Step 1: Write the word equation for the reaction.
- Step 2: Write the chemical equation using formulas.
- Step 3: Count atoms on both sides to check balance.
- Step 4: Balance by placing coefficients in front of chemical formulas (do not change subscripts).
- Step 5: Include state symbols for reactants and products.
Example 3: Nitric acid and sodium hydroxide produce sodium nitrate and water
- Word equation: sodium hydroxide + nitric acid → sodium nitrate + water
- Chemical equation: $ext{NaOH} + ext{HNO}3 ightarrow ext{NaNO}3 + ext{H}_2 ext{O}$
- Balance check: add coefficients to balance Na, H, N, O
- Balanced equation with state symbols: $ext{NaOH (aq)} + ext{HNO}3 ext{ (aq)} ightarrow ext{NaNO}3 ext{ (aq)} + ext{H}_2 ext{O (l)}$

6.4 Ionic Equations

Soluble ionic compounds exist as ions in aqueous solution.
An ionic equation is a simplified equation showing only the ions that participate in the reaction, excluding spectator ions.
Spectator ions are ions that appear unchanged on both sides of the equation and do not participate in the reaction.
State symbols must be included in any ionic equation.
How to identify and cancel spectator ions:
- Step 1: Write the balanced molecular equation with state symbols.
- Step 2: Write the full ionic equation by splitting soluble ionic reactants and products into their ions.
- Step 3: Cancel spectator ions on both sides.
- Step 4: Write the net ionic equation with state symbols.
Example: Silver nitrate and sodium chloride
- Molecular equation (aq): $ext{AgNO}3 ext{ (aq)} + ext{NaCl (aq)} ightarrow ext{AgCl (s)} + ext{NaNO}3 ext{ (aq)}$
- Complete ionic equation:
 $ext{Ag}^+ ext{(aq)} + ext{NO}3^- ext{(aq)} + ext{Na}^+ ext{(aq)} + ext{Cl}^- ext{(aq)} ightarrow ext{AgCl (s)} + ext{Na}^+ ext{(aq)} + ext{NO}3^- ext{(aq)}$
- Spectator ions cancel: $ext{Ag}^+ ext{(aq)} + ext{Cl}^- ext{(aq)} ightarrow ext{AgCl (s)}$
Practice question 1:
- Reaction: aqueous NaOH + aqueous Pb(NO₃)₂ → solid Pb(OH)₂ + aqueous NaNO₃
- Molecular equation: $2 ext{NaOH (aq)} + ext{Pb(NO}3)2 ext{ (aq)} ightarrow ext{Pb(OH)}2 ext{ (s)} + 2 ext{NaNO}3 ext{ (aq)}$
- Ionic particles: $2 ext{Na}^+ ext{(aq)} + 2 ext{OH}^- ext{(aq)} + ext{Pb}^{2+} ext{(aq)} + 2 ext{NO}3^- ext{(aq)} ightarrow ext{Pb(OH)}2 ext{(s)} + 2 ext{Na}^+ ext{(aq)} + 2 ext{NO}_3^- ext{(aq)}$
- Ionic equation after canceling spectator ions: $2 ext{OH}^- ext{(aq)} + ext{Pb}^{2+} ext{(aq)} ightarrow ext{Pb(OH)}_2 ext{(s)}$
Practice question 2: Aqueous NaOH and sulfuric acid react to form sodium sulfate and water
- Ionic equation: $2 ext{OH}^- ext{(aq)} + 2 ext{H}^+ ext{(aq)} ightarrow 2 ext{H}2 ext{O (l)}$ or simplified: $ext{OH}^- ext{(aq)} + ext{H}^+ ext{(aq)} ightarrow ext{H}2 ext{O (l)}$

Notes and Connections

Key objective across sections:
- Write and balance chemical equations, including covalent and ionic species.
- Determine and use correct state symbols in reactions.
- Understand how to derive ionic equations by separating ions in aqueous solutions and canceling spectator ions.
Practical implications:
- Balancing reactions ensures conservation of atoms and charge.
- Ionic equations help identify actual chemical species that participate in reactions in solution, particularly in precipitation and acid-base reactions.
Real-world relevance:
- Predicting products of reactions in solutions (e.g., precipitation, neutralisation).
- Understanding electrochemistry and catalysis where ionic species are central.
Mathematical and notational conventions:
- Charges written with superscripts: $1^{+}, 2^{+}, 3^{+}, 3^{-}, 2^{-}, 1^{-}$
- Subscripts in formulas denote fixed ratios (cannot be changed during balancing of equations): e.g., $ext{CO}3^{2-}$ , $ext{H}2 ext{O}$ , $ext{NaCl}$ .
- Coefficients balance atoms and charge; subscripts reflect the chemical formula, not altered during balancing.

Quick Reference: Common Ion Symbols and Formulas

Cations
- $ext{Li}^+, ext{Na}^+, ext{K}^+, ext{Be}^{2+}, ext{Mg}^{2+}, ext{Ca}^{2+}, ext{Al}^{3+}, ext{Cu}^+ , ext{Cu}^{2+}, ext{Fe}^{2+}, ext{Fe}^{3+}, ext{Zn}^{2+}, ext{H}^+ , ext{Ag}^+$
Anions
- $ext{O}^{2-} \, ( ext{oxide}), ext{F}^- , ext{Cl}^- , ext{Br}^- , ext{I}^- , ext{S}^{2-} \, ( ext{sulfide}), ext{NO}3^{-}, ext{SO}4^{2-}, ext{CO}_3^{2-}, ext{OH}^-$
Polyatomic ions
- $ext{OH}^- (1^-), ext{NO}3^- (1^-), ext{SO}4^{2-} (2^-), ext{CO}3^{2-} (2^-), ext{NH}4^+ (1^+)$

End of Chapter 06 Notes