Topic 3 – Molecules (SACE Stage 1 Chemistry)

Topic 3.1 – Molecular Polarity

Molecular Shapes and the VSEPR Model

Shape is dictated by the spatial arrangement of valence‐electron pairs (bonding & lone pairs) around the central atom.
Valence Shell Electron Pair Repulsion (VSEPR) theory: electron groups repel and adopt a geometry that maximises separation and minimises repulsion.

Applying VSEPR – 3-Step Method

Draw a full Lewis structure.
Count electron groups (bonding + lone). Any multiple bond counts as one electron group.
Match the count to an ideal geometry, then adjust for the number of lone pairs.

Ideal Geometries (no lone pairs)

Electron groups	Geometry	Generic formula
2	Linear	$AX_2$
3	Trigonal planar	$AX_3$
4	Tetrahedral	$AX_4$
5	Trigonal bipyramidal	$AX_5$
6	Octahedral	$AX_6$

Lone-Pair Variations (common examples)

Bent/V-shaped: $AX<em>2E$ or $AX</em>2E_2$
Trigonal pyramidal: $AX_3E$
Seesaw: $AX_4E$
T-shaped: $AX<em>3E</em>2$
Square pyramidal: $AX_5E$
Square planar: $AX<em>4E</em>2$

Worked Examples (workbook pp. 112-113)

CS₂ → linear (no lone pairs, 2 groups).
SO₂ → bent $AX_2E$ .
NF₃ → trigonal pyramidal $AX_3E$ .
CCl₄ → tetrahedral $AX_4$ .
CH₂O → trigonal planar $AX_3$ .

Expansion of the Octet (Hypervalency)

Period 3 and beyond can accommodate more than 8 valence e⁻ due to access to 3d, 4d … subshells.
Atoms containing >8 e⁻ in their valence shell are called hypervalent; they can form 5+ bonds (e.g. $PF<em>5$ , $SF</em>6$ ).

Electronegativity and Bond Polarity Recap

Polarity arises when electronegativity (EN) differences create partial charges (δ⁺, δ⁻).
Negligible EN difference ⇒ non-polar bond/molecule.

Molecular Polarity in 3-D

Overall dipole depends on both bond polarities and 3-D geometry.
CCl₄ example: although each C-Cl bond is polar, the tetrahedral symmetry places the centre of negative charge at the same point as the positive one → net dipole = 0 (non-polar).
Water (bent) & ammonia (trigonal pyramidal) have non-coincident centres → polar.

Topic 3.2 – Interactions Between Molecules

Intra- vs Intermolecular Forces

Intramolecular = covalent/ionic/metallic bonds within a molecule.
Intermolecular = forces between molecules; substantially weaker → called secondary interactions.

Dispersion (London) Forces

Caused by instantaneous, temporary dipoles from electron movement.
Exist in all molecules/atoms; strength ∝ number of electrons & surface area.
Larger atoms/molecules (e.g. I₂, 53 e⁻) have stronger dispersion forces than smaller (e.g. F₂, 17 e⁻) → higher melting/boiling points.

Dipole–Dipole Interactions

Occur between permanent molecular dipoles of polar molecules.
Stronger than dispersion; magnitude still scales with size/electron count.

Hydrogen Bonding (special dipole–dipole)

Requirements: H bonded within a molecule to N, O, or F and a lone pair on N/O/F of a neighbouring molecule.
Represented by dashed lines (···).
Considerably stronger than other dipole–dipole forces; key to water’s high b.p. and biological macromolecular structure.

Topic 3.3 – Hydrocarbons

Definition & General Features

Molecules composed solely of C & H.
Store chemical potential energy → excellent fuels; highly flammable.
Backbone of organic chemistry.

Combustion Reactions

Oxidiser (usually $O_2$ ) gains electrons, removing them from C-H bonds, liberating energy as heat/light.

Complete Combustion

$2C<em>8H</em>{18}(l) + 25O<em>2(g) \rightarrow 16CO</em>2(g) + 18H_2O(l)$

All C converted to $CO_2$ .

Incomplete Combustion

$2C<em>8H</em>{18}(l) + 17O<em>2(g) \rightarrow 16CO(g) + 18H</em>2O(l)$
$2C<em>8H</em>{18}(l) + 9O<em>2(g) \rightarrow 16C(s) + 18H</em>2O(l)$

Produces $CO$ or elemental C (soot) when $O_2$ limited.

Physical Properties vs Chain Length

Longer C-chains ⇒ more dispersion forces ⇒ higher m.p./b.p., viscosity.

Homologous Series & Functional Groups

Functional group: atom/group dictating reactivity.
Homologous series: molecules with same functional group, differing by one $CH_2$ per successive member.
• Share a general formula.
• Gradual physical property changes; chemical reactivity nearly identical.

Three Core Hydrocarbon Series

Series	General formula	Key bond type	Name ending
Alkanes	$C<em>nH</em>{2n+2}$	Single	–ane
Alkenes	$C<em>nH</em>{2n}$	≥1 C=C	–ene
Alkynes	$C<em>nH</em>{2n-2}$	≥1 C≡C	–yne

Saturation

Saturated: only single bonds (alkanes).
Unsaturated: contains C=C or C≡C (alkenes/alkynes) → addition reactions possible.

Molecular Representations

Molecular formula (e.g. $C<em>6H</em>{12}$ ).
Empirical formula → simplest ratio.
Structural / expanded formula → shows all bonds.
Condensed formula → e.g. $CH<em>3CH</em>2CH=CHCH<em>2CH</em>3$ .
Skeletal formula → zig-zag lines; H on C implied.

Structural Isomers

Same molecular formula, different connectivity. Example: C₅H₁₂ → pentane vs 2-methylbutane.

IUPAC Nomenclature (Alkanes/Enes/Ynes)

Identify principal functional group ⇒ fixes suffix.
Select longest continuous C-chain containing that group ⇒ parent chain.
Number chain giving:
• Lowest locant to double/triple bond (priority)
• Then lowest to first substituent.
Name substituents (methyl, ethyl, chloro, etc.), list alphabetical.
Use di-, tri-, tetra- for multiples; separate numbers with commas, numbers/letters with hyphens.

Multiple Double/Triple Bonds

Indicate all positions: e.g. penta-1,3-diene.

Example with Two Different Branches

Correct: 3-ethyl-2-methylhexane (alphabetical order: e < m).

Naming from Skeletal Structures

Each line terminus/intersection = C.
Locate functional groups, count longest path, assign substituents, follow naming rules.

Alcohol Functional Group (–OH)

General depiction $R–OH$ .
Naming: drop final –e, add –ol; place locant number before –ol (except sometimes 1).
• Example: 2-ethylpentan-1-ol.

Primary, Secondary, Tertiary Alcohols

Determined by how many C atoms the carbinol carbon (α-carbon) is attached to.
• Primary: 1 C (e.g. ethanol).
• Secondary: 2 C (e.g. 2-propanol).
• Tertiary: 3 C (e.g. 2-methyl-2-propanol).

Topic 3.4 – Polymers

Fundamentals

Polymers = very large molecules (high $M_r$ ) composed of repeating units called monomers.
Occur naturally (cotton, rubber) & synthetically (plastics).
Notation: enclose repeating unit in brackets, append $n$ : $[\;\;\;]_n$ .

Addition Polymerisation

Monomers must contain C=C (or C≡C). The π-bond breaks; new σ-bonds link monomers head-to-tail.
No atoms lost (contrast condensation polymerisation in later courses).
Example: Polyethylene from ethene.

Identifying the Monomer

Locate the smallest repeating sequence in the polymer chain, re-insert double bond → original alkene monomer.

Properties Determined By…

Monomer identity (functional groups, polarity, side chains).
Chain length – longer chains → higher tensile strength, higher m.p.
Branching / secondary interactions:
• Extensive branching ↓ packing, ↓ density, ↑ flexibility.
• Strong intermolecular attractions ↑ rigidity.

Thermoplastics vs Thermosets

Thermoplastics: weak secondary forces; soften/melt with heat; recyclable.
Thermosetting polymers: covalent cross-links/double bonds → strong networks; do not melt; hard, heat-resistant, non-recyclable.

Benzene Interlude (Slide 54)

Benzene = C₆H₆, planar hexagon with 3 alternating C=C.
6 π‐electrons become delocalised across ring → unusual stability (aromaticity).
Basis of many aromatic monomers and substituent phenyl groups (Ph–).

Workbook References & Practice

Topic 3.1: Q 1-7 (p.112-117)
Topic 3.2: Q 9-12 (p.120-123)
Topic 3.3: Q 13-14 (p.125-126)
Topic 3.4: Q 21-24 (p.149-154)

Ethical / Practical Relevance

Understanding polarity & intermolecular forces underpins solvent selection, drug formulation, material design.
Combustion chemistry connects to environmental issues: CO₂ (greenhouse gas) & CO (toxic) emissions.
Polymer science informs recycling strategies (thermoplastics) vs landfill concerns (thermosets).