Chemical Formulae, Naming Conventions, and Representation of Molecules

Types of Chemical Formulae

When asked to write a chemical formula, it generally refers to the empirical formula unless dealing with a molecular substance, in which case the molecular formula is required.

Empirical Formula

• Shows the relative number of each type of atom present in a substance.
  • Examples: H2O, NaCl, CH3
  • A subscripted number refers only to the atom immediately preceding it.

Molecular Formula

• Refers to a discrete molecule.
  • Example: C2H6

Structural Formula

• Aims to show how atoms in a molecule are bonded together.
  • Example: H-O-H
• Chemical symbols represent each element.
• Atoms are placed in the order in which they are bonded.
• Bonds between neighboring atoms are represented as lines.
• Does not accurately represent the geometry inside the molecule; it only shows bonding.

Chemical Formula Rules

• There is a standardized way of writing a formula.

Binary Compounds (2 Different Elements)

• The element further to the left on the periodic table appears first.
  • Examples: KCl, Al2S3
• Exception: Hydrogen often appears second.
  • Examples: NH_3, HCl
• If both elements are from the same group, the element with the highest period is first.
  • Examples: SiC, BrF_3
• In ionic compounds, the cation is written first, followed by the anion.
  • Examples: NaBr, MgCl_2
• Formulae for compounds with greater than 2 elements require some knowledge of the bonding involved.

Formula Rules II

Ionic Compounds

• Cation followed by anion, with the total charge equaling zero.
  • Example: Calcium nitrate \Rightarrow Ca(NO3)2(s) \rightarrow Ca^{2+}(aq) + 2NO_3^-(aq)
  • A subscripted number refers only to the atom immediately preceding it.

Covalent Compounds

• Generally, the chemical formula is written with the elements on the left-hand side (LHS) before those on the right-hand side (RHS), and the highest period first.
• For carbon compounds, carbon (C) is first, followed by hydrogen (H), and then the remaining elements in alphabetical order.
  • Examples: SO2, AsF3, IF5, C2H6O, C4H_9BrO
  • Exceptions exist with hydrogen.
    • Examples: H2CO3, H3PO4

Representing Molecules

• Different ways of representing molecules:
  • Chemical formula shows only the relative number of atoms.
  • Lewis dot formula (or electron dot) shows bonds between atoms as pairs of dots.
  • Structural formulas (bond-line formulas) show connections between atoms.
  • Ball-and-stick models show atoms as spheres and bonds as sticks, with accurate angles and relative sizes but exaggerated distances.
  • Space-filling models are accurately scaled-up versions of molecules but do not show bonds.
  • Electron-density models show the ball-and-stick model within the space-filling shape and color regions of high (red) and low (blue) electron charge.
• Why different representations?
  • Understanding shape, polarity, and potential reactivity.

Inorganic Naming

Naming of Standard Compounds

• Most areas of science require familiarity with the names of common chemicals.
• Basic rules and memory work are required.
• Divided into 4 classes:
  • Inorganic ionic compounds
  • Inorganic covalent compounds
  • Inorganic acids
  • Organic compounds (starts in week 6)

Nomenclature

IUPAC Nomenclature

• The system for naming developed by the International Union of Pure and Applied Chemistry (IUPAC).
• Some compounds are better known by their common unsystematic names than their systematic IUPAC names.
  • Water's systematic name is oxidane.
  • Acetone's systematic name is CH3C(=O)CH3
• For most compounds, systematic names are essential.

Common Names

• IUPAC names are systematic and unambiguous.
  • Name ⇔ formula
• Prior to IUPAC, common names were memorable and convenient.
  • Example: nitrous oxide = N_2O (laughing gas).
• Common names are nearly always used in catalogs for simple compounds.
• We will use common names for common chemicals.

IUPAC Naming

• The International Union of Pure and Applied Chemistry sets the rules.
  • IUPAC Organic naming 2013e 11MB or 1306 pages!
• First, understand classes of compounds in Inorganic main group chemistry.

1) Covalent Compounds

• Generally form from elements on the RHS of the periodic table combining with elements from the RHS (non-metals with non-metals).
• Covalent bonding involves atoms sharing electrons in pairs.
  • Example: H-O-H, O=C=O

2) Ionic Compounds

• Generally, form from elements on the LHS of the periodic table combining with elements from the RHS (metals with non-metals).
• Ionic bonding involves the attraction of cations to anions.
  • Example: MgO

3) Inorganic Acids

• Possess a labile proton (easily lost) and a stable component (anion) that is left behind.
  • Examples: H…Cl  H+ + Cl-, H..OSO3 -  H+ + SO4 2- ( H2SO4  H+ + HSO4 -)

Oxidation Number Rules

• Many metallic cations in ionic compounds have variable oxidation states.
• Rules to determine oxidation states/numbers.
• These rules are used throughout the subject.

Oxidation Number Rules

• Rule 1: For anions & cations, the oxidation number is equal to the charge on the ion.
  • Examples: Na^+ oxid. no. +1, O^{2-} oxid. no. 2-
• Rule 2: The oxidation number of a pure element is zero.
  • Examples: O2, P4, S_8 oxid. no. = 0
• Rule 3: Fluorine (F) always has an oxidation state of -I
• Rule 4: Oxygen (O) normally has an oxidation state of -II
• Rule 5: Hydrogen (H) has an oxidation state of +I
• Rule 6: The sum of oxidation numbers must equal the charge written on the formula.
  • Example: H2O – charge written is 0 (H2O = H_2O^0) sum = 0 = 1 x O + 2 x H = 1 x -II + 2 x +I

Worked Problem

• Common anions: hydroxide OH^- and perchlorate ClO_4^-
• Determine the oxidation state of each of the atoms in the hydroxide ion.
  • e.g. OH^-
    • O -II
    • H +I
    • sum = -I
    • -1 = 1 x O + 1 x H = 1 x -II + 1 x +I
    • RHS = LHS
• For perchlorate you should get +VII for Cl.
• BURSTING STAR
  • KClO_4 - 50
  • K3[Fe(CN)6] - 20
  • C dust - 25
  • Al powder - 5

Type 1 Ionic Compounds I

• Inorganic Ionic Compounds
  • Two Types:
    • (i) Type I Metal Cations + Non-metal anion e.g. Na Cl, MgO
      • The metal cation is either an alkali or alkaline earth metal and these metals have fixed oxidation states
        • e.g. Na^+ = Na (I)
      • The name is composed as follows:
        • Name: cation (first word) anion (second word)
          • e.g. sodium chloride

Type 1 Ionic Compounds II

• Name: cation (first word) anion (second word)
  • Where the cation is monatomic (1 atom): monatomic cation: uses the element name
  • Where the anion is a monatomic non-metal: monatomic anion: uses the root of the element name and adds -ide
    • chlor- (root of chlorine), add –ide \Longrightarrow chloride
    • e.g. magnesium iodide MgI_2
• N.B. PREFIXES ARE NOT USED!

Polyatomic Ions / Naming

• Polyatomic ions use common names that must be memorised.
  • e.g. NH4^+ ammonium, ClO4^- perchlorate
    • NH4ClO4 ammonium perchlorate
• The common polyatomic cations are mercurous Hg2^{2+}, ammonium NH4^+, and hydronium H_3O^+, see Table 2.4 BLB.
• There are many common polyatomic anions, see Table 2.5 BLB.

Common Ions & Charges

• Polyatomic ion names you need to know:
  • Cations
    • NH_4^+ ammonium
    • H_3O^+ hydronium
  • Anions
    • (fewer O) (more O)
      • NO2^- nitrite NO3^- nitrate
      • SO3^{2-} sulfite SO4^{2-} sulfate
    • OH^- hydroxide
    • CN^- cyanide
    • PO_4^{3-} phosphate
    • CO_3^{2-} carbonate
    • HCO_3^- bicarbonate
    • ClO_4^- perchlorate
    • CH_3COO^- acetate
    • MnO_4^- permanganate
    • Cr2O7^{2-} dichromate
    • O_2^- peroxide
    • CaCO_3

Type 2 Ionic Compounds

• (ii) Type II Metal Cations + Non-metal anions e.g. FeCl_3, MnO
  • The metal cation is a transition metal, and these metals have multiple oxidation states
    • e.g. Fe^{3+} = Fe (III)
    • Mn^{2+} = Mn(II)
  • The name is composed as follows:
    • cation(ox. no.) first word anion (second word)
      • use the element name for the cation & add the ox. state in Roman numerals.
        • e.g: FeCl_3 iron (III) chloride
        • N.B. PREFIXES ARE NOT USED!

Type 2 Ionic Compounds Common Names

• Common names: use the