Topic 1 — Materials & Their Atoms (SACE Stage 1 Chemistry)

Properties and Uses of Materials

Materials underpin every aspect of daily life and their selection is governed by the relationship between atomic‐scale structure and bulk properties.

• Key bulk properties used to classify or select a material
• Solubility – ability of a substance to dissolve (typically in water). When a material dissolves a solute disperses in a solvent to give a homogeneous solution. Insoluble microscopic dispersions are termed colloids.
• Electrical conductivity – ease with which electrons flow through a lattice. Metallic lattices (delocalised electrons) give high conductivity; ionic or covalent lattices generally do not conduct when solid.
• Thermal conductivity – capacity to transport heat. Good thermal conductors transfer energy efficiently, whereas thermal insulators hinder heat flow.
• Melting point (m.p.) – temperature where solid → liquid (e.g. water $0\,^{\circ}\mathrm{C}$).
• Boiling point (b.p.) – temperature where liquid → gas (e.g. water $100\,^{\circ}\mathrm{C}$).

• Classical states of matter recap:
– Solids: rigid, fixed shape & volume.
– Liquids: not rigid, no fixed shape, fixed volume.
– Gases: not rigid, no fixed shape, no fixed volume.

Size Regimes: Macro → Nano

• Matter ranges from macroscopic (visible) to microscopic (invisible) and further to the nanoscale ($1\,\mathrm{nm}=10^{-9}\,\mathrm{m}$).
• A glucose molecule (≈$1\,\mathrm{nm}$) is $10^{6}$ times smaller than a 1 mm printed full stop.

Nanomaterials

• At $1-100\,\mathrm{nm}$ the discrete nature of atoms/molecules dominates; materials cannot be treated as continuous.
• Nanoparticles exhibit new properties relative to bulk forms because:
– Surface atoms dominate – very high surface‐area-to-volume ratio (SA:V).
– Surface forces that appear weak macroscopically (e.g. van der Waals) become significant.
– Greater proportion of reactive atoms gives enhanced reactivity or catalytic activity.

Example: Dividing a $1\,\mathrm{cm}$ cube into $1\,\mathrm{nm}$ cubes increases total surface area from $60\,\mathrm{cm^{2}}$ to $6.0\times10^{7}\,\mathrm{cm^{2}}$.

Historical Perspective & Vision

• Richard Feynman’s 1959 talk “There’s Plenty of Room at the Bottom” posed questions about manipulating atoms, building molecular machines and ultra‐dense data storage.
• Modern nanotechnology realises these ambitions by engineering matter atom-by-atom to deliver fundamentally new properties and functions.

Everyday & Emerging Applications

• Invisible sunscreen – ZnO nanoparticles (< visible wavelength) do not scatter white light yet still absorb UV.
• Lotus effect – nanoscale waxy bumps on lotus leaves create extreme hydrophobicity; mimicked to make self-cleaning, stain-resistant fabrics and surfaces (e.g. carbon nanotube-coated cloth).
• Self-cleaning glass – $\text{TiO}<em>2$ nanocoating (~40 nm) becomes super-hydrophilic under UV; water sheets off carrying altered dirt particles. • Aerospace concepts – carbon nanotube composites for light, strong shells; nanoelectronics; H$2$ storage; smart MEMS/NEMS control surfaces.

Classification of Substances

• Elements – single type of atom (pure).
• Compounds – different atoms chemically bonded in fixed ratios (pure).
• Mixtures – two or more elements/compounds physically combined without reaction.

• Homogeneous mixtures – uniform at molecular level (e.g. alloy, salt solution, air).
• Heterogeneous mixtures – visibly distinguishable phases (e.g. soil, vinaigrette, crude oil).

Separating Mixtures

Differences in physical properties enable separation:

1. Filtration – separates solids from liquids in heterogeneous suspensions via a porous barrier; both filtrate (liquid) and residue (solid) recovered.

2. Evaporation – heats mixture so lower-b.p. solvent vaporises; high-b.p. solute remains.

3. Simple distillation – heats mixture, the component with the lowest b.p. vaporises, condenses and is collected (distillate).

4. Fractional distillation – for liquids with close b.p.’s; fractionating column provides surfaces for repeated vaporisation-condensation cycles, effecting finer separation.

Atomic Structure

• Atom = central nucleus (protons $p^{+}$, neutrons $n^{0}$) + orbiting electrons $e^{-}$.
• Sub-atomic summary:
– Proton: charge $+1\,e$, mass $1\,\text{amu}$, nucleus.
– Neutron: charge $0$, mass $1\,\text{amu}$, nucleus.
– Electron: charge $-1\,e$, mass $\approx1/1800\,\text{amu}$, outside nucleus.

• Electron shells hold $2,\ 8,\ 18,\ 32$ e⁻ (for principal quantum numbers $n=1$ to 4).
• On the periodic table:
– Atomic number (Z) = protons = electrons (neutral atom).
– Mass number (A) ≈ protons + neutrons.
– Neutrons $N = A - Z$.

Electron Energy & Spectroscopy

• Electrons possess electrical potential energy due to Coulomb attraction; energy increases with distance (higher shells).

Atomic Absorption & Emission

• Absorption: photon energy promotes an electron to a higher shell (excited state).
• Emission: when the electron relaxes it emits a photon with energy equal to the gap.
• Energy levels are quantised → atoms absorb/emit only specific photon ‘chunks’.
• Low-energy / long-λ photons appear red; high-energy / short-λ photons appear violet.
• Each element therefore has a characteristic line spectrum usable for identification (atomic absorption spectroscopy or emission spectroscopy).

Nuclear Symbols & Isotopes

• General notation: $\displaystyle ^{A}_{Z}!X$ where $X$ = element symbol, $A$ = mass number, $Z$ = atomic number.
• Isotopes – atoms of the same element (same $Z$) but different $A$ (different neutrons).
– Chemical behaviour identical (same e⁻ configuration).
– Physical properties (mass, density, radioactivity) differ.
– Example: $^{12}\text{C},\ ^{13}\text{C},\ ^{14}\text{C}$.

Quantum Model: Subshells & Orbitals

• Each shell $n$ contains $n$ subshells labelled $s,p,d,f$.
– $s$: 1 orbital, 2 e⁻; $p$: 3 orbitals, 6 e⁻; $d$: 5 orbitals, 10 e⁻; $f$: 7 orbitals, 14 e⁻.
• Orbitals are 3-D probability regions for locating an electron.

Electron Configuration Rules

1. Aufbau principle – fill lowest available energy orbitals first.

2. Pauli exclusion – max 2 e⁻ per orbital with opposite spins.

3. Hund’s rule – degenerate orbitals singly occupied before pairing.

• Exceptions (energy stabilisation via half/full $d$ sublevel):
– Chromium: $[\text{Ar}]\,3d^{5}\,4s^{1}$ (not $3d^{4}4s^{2}$).
– Copper: $[\text{Ar}]\,3d^{10}\,4s^{1}$ (not $3d^{9}4s^{2}$).

Ions

• Main-group ions: add/remove electrons according to charge and then write configuration.
• Transition metals lose $s$ electrons before $d$ when forming cations (e.g. Fe: $[\text{Ar}]3d^{6}4s^{2}$ → Fe²⁺: $[\text{Ar}]3d^{6}$).

Quantifying Matter: The Mole

• Counting individual atoms is impractical; chemists use the mole.
• Avogadro’s number: $N_{A}=6.022\,140\,76\times10^{23}\,\text{particles·mol}^{-1}$.

Molar Mass

• Molar mass (M) = mass of 1 mol in $\mathrm{g\,mol^{-1}}$. For elements, $M$ equals the relative atomic mass (periodic table).
• For a compound $\displaystyle M = \sum n<em>{i}M</em>{i}$ where $n<em>{i}$ = number of atoms of element $i$. Example: $M</em>{\mathrm{SO_{2}}} = 1\times32.07 + 2\times15.99 = 64.06\,\mathrm{g\,mol^{-1}}$.

Mole Calculations

Relationship: $\boxed{n = \dfrac{m}{M}}$
Where $n$ = moles, $m$ = mass (g), $M$ = molar mass.
Example: Mass $0.214\,\mathrm{g}$ Fe ($M=55.85$) ⇒ $n = 0.214/55.85 = 3.83\times10^{-3}\,\text{mol}$.

Periodic Table Architecture

• Groups (columns) – same valence e⁻ count; 18 groups total.
• Periods (rows) – same number of electron shells; 7 periods.
• Block notation:
– $s$-block (Groups 1–2 + He), $p$-block (13–18), $d$-block (transition metals), $f$-block (lanthanides & actinides).

• Electron configuration reveals position:
– Silicon $1s^{2}2s^{2}2p^{6}3s^{2}3p^{2}$ → ends in $3p^{2}$ → period 3, group 14 (4 valence e⁻).
– Potassium $[\text{Ar}]4s^{1}$ → period 4, group 1, $s$-block.

Metals, Non-Metals & Metalloids

• Metals: good conductors, malleable, lustrous; located left & centre.
• Non-metals: poor conductors, diverse states; upper right.
• Metalloids: intermediate properties (B, Si, Ge, As, Sb, Te, etc.).

Periodic Trends

Atomic Radius

• Increases down a group – additional shells increase distance from nucleus.
• Decreases across a period – nuclear charge rises while shielding constant ⇒ stronger pull on same‐shell electrons.

Electronegativity (Pauling Scale 0–4)

• Decreases down a group – larger atoms hold valence e⁻ less tightly.
• Increases across a period – higher effective nuclear charge and smaller radius enhance electron-attracting power.
• Extrema: Fluorine $\chi=4.0$ (highest); Caesium/Francium $\chi=0.7$ (lowest).

Practical implication: bond polarity and type (ionic vs covalent) can be predicted from electronegativity differences.

Ethical & Practical Considerations of Nanotechnology

• Enhanced reactivity raises concerns over toxicity, environmental persistence, and unforeseen biological interactions.
• Regulation and safe design ("safer‐by‐design" nanoparticles, recycling strategies) are crucial for responsible adoption.

Connections & Relevance

• Fundamental atomic structure governs spectra, guiding analytical tools (AAS, AES) used in mining, environmental monitoring, and astrophysics.
• Nanomaterial surface phenomena connect directly to catalytic industrial processes and biomedical drug delivery.
• Periodic trends enable rational design of alloys, semiconductors, and battery materials.