Topic 2 - Combining of Atoms (SACE Stage 1 Chemistry)

Topic 2.1 – Combining Atoms

Materials (solids, liquids, gases) are produced when individual atoms interact through primary, intramolecular forces that we call bonds. The bond type that forms is dictated entirely by the nature of the atoms involved.

Primary Bond Types

• Metallic – metal ↔ metal.
• Ionic – metal ↔ non-metal.
• Covalent – non-metal ↔ non-metal.

The patterns common to the three classes are summarised below.

1. Metallic: a lattice of fixed, positive metal cations is immersed in a mobile “sea” of delocalised valence electrons. Strong, non-directional electrostatic attraction operates between nuclei (+) and the mobile electrons (–).

2. Ionic: extended 3-D arrays of alternating cations and anions are stabilised by the directional electrostatic attraction between opposite charges. Charge neutrality (total + = total –) is compulsory.

3. Covalent: atoms share one, two or three pairs of electrons. The sharing can be localised inside individual molecules (covalent-molecular) or extend throughout a giant network (covalent-network).

Metallic Bonding – Structure and Origin

• Example elements: Mg, Zn, Cu, Ag, Fe.
• Each metal atom donates its valence electrons to the common pool, producing $\text{M}^{n+}$ ions in a rigid lattice.
• Although neighbouring cations repel, the attraction to the electron sea dominates, yielding a cohesive, malleable solid.

Ionic Bonding – Structure and Origin

• Typical solids: $\text{NaCl}$, $\text{MgO}$.
• A regular lattice packs positive and negative ions in fixed whole-number ratios (e.g. $1:1$ for NaCl, $1:1$ for MgO). Charge balance follows $\sum q<em>{+}=\sum q</em>{-}$.

Covalent Bonding – Two Sub-classes

1. Covalent Molecular (e.g. sucrose, $\text{I}_2$): discrete molecules are held together by weak intermolecular forces (dispersion, dipole-dipole, H-bonding).

2. Covalent Network (e.g. diamond, graphite, SiO2): an extended lattice of strong, directional covalent bonds produces hardness, rigidity and very high melting points.

Melting Points – Particle View

Heating a solid increases vibrational KE. Fusion occurs when vibrations overpower the attractive forces (inter- or intramolecular) that maintain order. Therefore network solids > ionic solids > metals > molecular solids in $T_m$.

Electrical Conductivity – Bond Type Fingerprint

• Metals: delocalised electrons move freely → excellent conductors in any phase.
• Ionic: conduct only when molten or aqueous (ions mobile). Solids are insulators.
• Covalent molecular & network: generally insulators; exception = graphite whose layered network contains one delocalised $\pi$ electron per C atom.

Topic 2.2 – Bond Formation & Energy

Chemical reactions rearrange bonds. Forming bonds always releases energy (exothermic at the bond level), whereas breaking bonds requires energy.

Endothermic vs Exothermic

• Endothermic: E{\text{break}}>E{\text{make}} → net absorption, surroundings cool.
• Exothermic: E{\text{make}}>E{\text{break}} → net release, surroundings warm.

Activation Energy ( $E_a$ )

The minimum energy required for reactants to reach the transition state. Think of reactants climbing an energy hill then descending to products.

Questions to test understanding (lecture prompts): definition of $E<em>a$, temperature effect (higher T gives more particles $E\ge E</em>a$), role of catalysts (lower $E_a$), collision orientation, etc.

Metallic Bonding – Detailed Properties

• High electrical/thermal conductivity – free electrons transport charge/energy.
• Malleable & ductile – layers of cations slide while electrons maintain cohesion.
• Lustre – electrons absorb & re-emit photons across broad frequencies.
• Density – high due to close packing; varies with ionic mass & lattice type.
• High $T<em>m, T</em>b$ – breaking strong electrostatic metal-electron attractions is energy intensive.

Heat Treatments of Metals

Tempering, quenching, annealing alter microstructure → modify hardness, ductility. Rapid quench traps defects (hard, brittle); slow cool allows recrystallisation (soft, ductile).

Metal Alloys

When different metals mix, size mismatch distorts the lattice. Extra strain hampers layer slippage → alloy is harder than either parent. Example: brass (Cu+Zn).

Ionic Bonding – Forming Ions to Obey Octet Rule

• Atoms in groups 1–3 lose electrons → cations. Energy input = first (or successive) ionisation energies.
• Atoms in groups 5–7 gain electrons → anions. Energy input must overcome e^––e^- repulsion; driving force is large electron affinity + lattice energy released on salt formation.

Example Equations

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

Charge vs Group

Group V gains 3e^- → $3^-$, group VI gains 2e^- → $2^-$, group VII gains 1e^- → $1^-$, noble gases typically inert.

Writing Ionic Formulae – Balancing Charges

1. Write symbols with oxidation states.

2. If |charge| equal → subscripts 1:1 (e.g. $\text{Na}^+\, \text{Cl}^-$ → NaCl).

3. Otherwise cross-multiply to smallest whole-number ratio.
   • $\text{Ca}^{2+}$ & $\text{Cl}^-$ → $\text{CaCl}<em>2$ (1:2). • $\text{Al}^{3+}$ & $\text{O}^{2-}$ → $\text{Al}</em>2\text{O}_3$ (2:3).

Transition-Metal Oxidation States

E.g. $\text{Fe}^{2+}$ = iron(II), $\text{Fe}^{3+}$ = iron(III).

Polyatomic Ions

E.g. $\text{SO}<em>4^{2-},\;\text{NO}</em>3^-,\;\text{NH}_4^+$. Treat entire ion as a single charged unit when balancing.

Naming Ionic Compounds

Rule 1 (two elements, metal+non-metal): metal name + non-metal stem + “ide” (MgO → magnesium oxide). Hydrogen precedes other non-metals if present.
Rule 2 (compound contains oxygen + another non-metal): metal name + stem + “ate” (Na₂CO₃ → sodium carbonate).
Rule 3 (two non-metals only): use Greek prefixes (CO₂ = carbon di-oxide; CO = carbon mon-oxide).

Ionic Lattice Properties

• Hard & brittle – compression misaligns layers; like charges juxtapose and crystal cleaves along planes.
• High $T<em>m, T</em>b$ proportional to |ionic charge| and ionic radii.
• Non-conducting as solids; conductive when molten or aqueous (electrolytes).

Covalent Bonding – Fundamentals

Shared electron pair attracts both atomic nuclei.
Single bond (1 pair), double (2 pairs), triple (3 pairs) with increasing bond strength and decreasing bond length.

Molecules

• Homonuclear (same element) vs heteronuclear (different). Seven elemental diatomics: $\text{H}<em>2, \text{N}</em>2, \text{O}<em>2, \text{F}</em>2, \text{Cl}<em>2, \text{Br}</em>2, \text{I}_2$.

Electron-Dot (Lewis) Diagrams

• Dots = valence electrons. Lines can replace paired dots to indicate bonding. Central atom is usually least electronegative.
• Construction algorithm: write symbols, sum valence e^- count, place single bonds, satisfy outer atoms’ octets, place leftovers on centre, form multiple bonds as required.

Bond Polarity

Difference in electronegativity $\Delta EN$:
• If $\Delta EN\approx0$ → electrons shared equally, non-polar.
• If 0<\Delta EN<\sim2 (for covalent range) → unequal sharing → polar with partial charges $\delta^- / \delta^+$.
Dipole moment magnitude increases with $\Delta EN$ and bond length.

Properties of Covalent Substances

• Molecular: weak intermolecular forces → low $T<em>m, T</em>b$, softness, poor electrical & thermal conduction.
• Network: extensive σ-bonds → extreme hardness, very high $T<em>m, T</em>b$, generally insulating (graphite is the noted conductive exception due to delocalised $\pi$ electrons between layers).

Naming Covalent Molecules & Networks

Use Greek prefixes to show atom count; drop mono- on first element. For networks the empirical formula (formula unit) is quoted, sometimes written $(AX<em>n)</em>n$ to emphasise repetition.

Topic 2.3 – Quantitative Chemistry Recap

The Mole

One mole contains $6.022\times10^{23}$ entities (Avogadro’s constant $N_A$).

Number of particles $n<em>p$ is found via $n</em>p = n \times N_A$
where $n$ is amount of substance in moles.

Percentage Composition

1. Identify molecular formula.

2. Compute molar mass $M$.

3. For each element $X$ use
   $\%\,X = \frac{(\text{atoms of }X)\times M<em>X}{M</em>{\text{compound}}}\times100$.

Empirical Formula from % Composition

Assume 100 g → % converts directly to mass. Then
$n = \frac{m}{M}$ for each element → divide all $n$ by smallest to obtain simplest ratio.
Example: 40.0 % C, 6.71 % H, 53.28 % O gave $\text{CH}_2\text{O}$.

Empirical Formula from Mass Data

Same procedure but start from measured masses. If subscripts include fractions, multiply all by a whole number factor to clear them (e.g. 3.5 → ×2).

Example: 2.66 g Cl & 4.20 g O → initial ratio $\text{Cl}<em>{0.075}\text{O}</em>{0.262}$. Divide by 0.075 → $\text{ClO}<em>{3.5}$; multiply by 2 → $\text{Cl}</em>2\text{O}_7$ (final empirical formula).

These notes capture every principle, definition, example, formula and practical implication discussed in the transcript, providing a self-contained study resource for Topic 2: Combining of Atoms and Quantities in Stage 1 Chemistry.