chem

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6.1 What Are Chemical Formulae?
Learning Outcome
• State chemical symbols of elements and formulae of compounds.

Key ideas

  • A chemical formula is like a recipe that shows the ratio of elements in one unit of a substance.
  • It is made up of:
    • chemical symbol(s) for the elements present; and
    • subscript(s) indicating the number of atoms of each element in the molecule or formula unit.
  • Examples:
    • Carbon dioxide: one C and two O atoms in each molecule: CO2CO_2
    • Water: two H atoms and one O atom: H2OH_2O
  • Similar-looking molecules can be different substances because their compositions differ.
    • One molecule of carbon dioxide is made of 1 C and 2 O: CO2CO_2
    • One molecule of water is made of 2 H and 1 O: H2OH_2O

Formulae of Elements

  • Elements can exist as monoatomic, diatomic, or polyatomic species.
  • Monoatomic elements: exist as single, uncombined atoms. Example: noble gases in Group 18 often exist as monoatomic species (full valence shell).
  • Group 18 elements (noble gases) are commonly monoatomic (e.g., Ne, Ar).
  • Word Alert: A formula is a list of constituents.
  • Prefixes mono- and di- indicate one and two atoms, respectively.
    • In CO₂, the prefix di- indicates there are two oxygen atoms.

Diatomic Molecules

  • Diatomic molecules are composed of two atoms of the same element chemically bonded together.
  • Examples of diatomic molecules include elements like H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. The text notes these as common diatomic molecules (Table 6.1).
  • Some elements in Group 17 (the halogens) also exist as diatomic molecules (e.g., Cl₂, Br₂, F₂).

Polyatomic Molecules

  • Polyatomic molecules contain three or more atoms chemically bonded together (e.g., S₈, P₄, O₃).
  • Example in the text: sulfur (S), phosphorus (P), ozone (O₃).

Chemical Formulae of Compounds

  • Compounds are formed from atoms of elements that are chemically combined and have fixed formulae. Mixtures do not have fixed formulae.
  • The modern chemical formulae can be traced to Hill system ideas; the Hill system was proposed in 1900 by Edwin A. Hill.
  • Hill system vs modern system:
    • Hill system originally listed elements in alphabetical order.
    • Modern system writes formulae to reflect composition and bonding (e.g., H₂SO₄ for sulfuric acid).
  • Examples comparing Hill vs Modern system (Table 6.3 ideas):
    • Ammonia: NH₃ (Hill) vs NH₃ (Modern) – unchanged.
    • Copper(II) sulfate: CuSO₄ (both systems in common use).
    • Sodium hydroxide: NaOH (both systems).
    • Hydrochloric acid: HCl (both systems).
    • Ethanoic acid: CH₃COOH (Hill) vs CH₃COOH (Modern).
    • Ethene (ethene): CH₂=CH₂ (Modern) vs CH₂CH₂ (Hill).
    • Trichloromethane (chloroform): CHCl₃ (both).
    • Propanoic acid: CH₃CH₂COOH (Modern) vs CH₃CH₂COOH (Hill).

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6.2 How Are Chemical Formulae Constructed?
Learning Outcomes
• Deduce the formulae of simple compounds from the relative numbers of atoms present and vice versa.
• Deduce the formulae of ionic compounds from the charges of ions and vice versa.

Analogy: Fixed studs in plastic bricks

  • Atoms have fixed numbers of electrons involved in bonding; this determines the formula of an ionic compound.
  • The “studs” analogy: atoms have fixed bonding capacities (valences) that determine which other ions/atoms can attach to form a stable compound.

Valences from Group Numbers

  • Valence is the number of electrons that must be lost, gained, or shared to achieve a noble-gas electronic configuration.
  • A quick guide from the periodic table: valence patterns follow group numbers (as shown in Table 6.4 in the text).
  • Example mappings (from Table 6.4 in the text):
    • Valency 1: Group 1 (Li, Na, K)
    • Valency 2: Group 2 (Be, Mg, Ca)
    • Valency 3: Group 13 (B, Al)
    • Valency 4: Group 14 (C, Si)
    • Valency 3: Group 15 (N, P)
    • Valency 2: Group 16 (O, S)
    • Valency 1: Group 17 (F, Cl, Br, I)
  • Note: The valence concept is used to determine how ions combine to form compounds.

Valences from Roman Numerals

  • Transition metals can form more than one stable ion; their valences are indicated with Roman numerals in brackets (Table 6.5).
  • Example ions/valences:
    • Iron(III) ion: Fe³⁺
    • Iron(II) ion: Fe²⁺
    • Copper(II) ion: Cu²⁺
    • Copper(I) ion: Cu⁺
  • Silver (Ag) forms only one stable ion (Ag⁺); its valency is fixed (Table 6.6).

Valences of Polyatomic Ions

  • Polyatomic ions are composed of more than one atom covalently bonded and carry an overall charge.
  • Common polyatomic ions and their valences (Table 6.7):
    • Ammonium: NH4+NH_4^+, valence 1
    • Hydroxide: OH−OH^-, valence 1
    • Nitrate: NO3−NO_3^-, valence 1
    • Sulfate: SO42−SO_4^{2-}, valence 2
    • Carbonate: CO32−CO_3^{2-}, valence 2
    • Phosphate: PO43−PO_4^{3-}, valence 3

Deducing Chemical Formulae of Compounds

  • Crossing (cross-multiplication) method: balance ions by crossing their valences to form neutral compounds.
  • This method balances electrons gained/lost/shared in bonding.
  • Worked Example 6A: Deduce the chemical formula of sodium phosphate.
    • Sodium (Na) has valency 1; phosphate (PO₄³⁻) has valency 3.
    • To neutralize, combine 3 Na⁺ with 1 PO₄³⁻ → formula: Na<em>3PO</em>4Na<em>3PO</em>4
  • Worked Example 6B: Aluminium sulfate.
    • Aluminium (Al³⁺) with sulfate (SO₄^{2-}). To balance, take 2 Al³⁺ and 3 SO₄^{2-} → formula: Al<em>2(SO</em>4)3Al<em>2(SO</em>4)_3

Deducing Charges