Atoms, Elements, and Compounds: Properties of Matter and Atomic Theory

The Nature and Classification of Matter

• Definition of Matter: Matter is defined as anything that possesses mass and occupies physical space.

• Definition of Chemistry: Chemistry is the branch of science dedicated to the study of matter, its properties, its constitution, and the transformations it undergoes.

• The Particle Theory of Matter: Summarizes the fundamental composition and dynamical behaviour of matter through five key postulates:

  • All matter is composed of extremely small particles that have empty spaces between them.

  • Different pure substances are composed of fundamentally different kinds of particles.

  • Particles are in perpetual, random motion.

  • The velocity and kinetic energy of the particles increase as the temperature of the substance rises.

  • Particles exert electrostatic attractive forces on one another.

• States of Matter Explained by Particle Theory:

  • Solid: Interparticle attractive forces are strong enough to constrain particles in a fixed, rigid arrangement where they vibrate in place, resulting in a definite shape and volume.

  • Liquid: Gained thermal energy allows particles to move faster and slide past one another; attractive forces hold them close together, maintaining a fixed volume while allowing fluid flow to match the shape of a container.

  • Gas: Significant thermal energy enables particles to completely overcome attractive forces; particles move rapidly and independently in straight lines, separated by large distances, occupying the entire volume and shape of their container.

• Broad Classification of Matter:

  • Pure Substance: A class of matter composed of only one type of particle with uniform, invariant chemical and physical properties throughout (e.g., pure distilled water, pure elemental gold).

  • Mixture: A physical combination of two or more distinct types of particles that retain their respective chemical identities. Mixtures occur in solid, liquid, or gaseous phases.

• Subdivisions of Mixtures:

  • Mechanical Mixture (Heterogeneous Mixture): A mixture containing multiple types of matter where the individual components remain visually distinguishable to the unaided eye or under an optical microscope (e.g., granola bars, breakfast cereal, pizza, concrete, milk, fog, freshly squeezed orange juice with pulp).

  • Solution (Homogeneous Mixture): A uniform mixture of two or more substances that appears visually identical to a pure substance because the distinct particles are dispersed on a molecular level and cannot be visually differentiated. All liquid and gaseous solutions are optically clear and transparent; any cloudiness or opacity indicates a heterogeneous suspension or mechanical mixture (e.g., clear apple juice, filtered seawater, air).

  • Alloy: A solid solution composed of two or more metals (or a metal combined with non-metals) melted together and cooled to form a single solid phase:

  • Solder: An alloy traditionally composed of tin and lead; utilized to join electrical wires and copper pipes due to its fluidity at elevated temperatures and rapid solidification.

  • Lead-Free Solder: Modern non-toxic alternatives composed of tin, copper, silver, or bismuth, developed to eliminate lead poisoning hazards.

  • Brass: A solid solution of approximately $75\%$ copper and $25\%$ zinc.

  • Sterling Silver: A metallic alloy consisting of $92.5\%$ silver and $7.5\%$ copper, which improves hardness compared to pure silver.

  • $14\text{ kt}$ Gold: An alloy composed of $14$ parts gold and $10$ parts copper (approximately $58.3\%$ gold by mass).

Physical Properties and Physical Changes

• Physical Property: A characteristic of a substance that can be determined, observed, or measured without altering the chemical composition of the substance.

• Qualitative Physical Properties: Characteristics observed directly through the senses without numerical measurements:

  • Colour: Observed wavelength appearance (e.g., colorless, red, golden).

  • Odour: Detectable scent (e.g., odorless, sharp, sweet).

  • Taste: Flavor profile (e.g., sweet, salty, sour).

  • Texture: Tactile quality (e.g., powdery, smooth, rough).

  • Lustre: The degree of surface reflectiveness, shininess, or dullness (e.g., high lustre in polished silver; low lustre in a rusted iron nail).

  • Optical Clarity: The capacity of a material to transmit light:

  • Transparent: Transmits light cleanly without scattering (e.g., clear window glass, clear apple juice).

  • Translucent: Transmits light diffusely; images cannot be clearly resolved (e.g., frosted glass).

  • Opaque: Completely blocks and absorbs or reflects light (e.g., brick wall, cardboard).

  • Brittleness: The susceptibility of a material to crack, fracture, or shatter under mechanical stress rather than deforming (e.g., glass is brittle; modeling clay is flexible).

  • Viscosity: The quantitative or qualitative resistance of a fluid to flow or pour (e.g., molasses and ketchup possess high viscosity; water possesses low viscosity).

  • Hardness: The relative resistance of a substance to scratching by another material, frequently evaluated on the Mohs hardness scale (e.g., candle wax is soft; diamond is extremely hard and scratches almost all other materials).

  • Malleability: The capacity of a substance to be hammered, rolled into thin sheets, or molded without breaking (e.g., gold and aluminum foil are malleable; glass is non-malleable).

  • Ductility: The capacity of a material to be drawn or stretched into fine wires or strands under tensile force (e.g., copper wire, solder wire).

  • Electrical Conductivity: The ability of a substance to permit the flow of electric charge through it (e.g., copper is an excellent electrical conductor; plastics and ceramics act as non-conductive insulators).

• Quantitative Physical Properties: Characteristics that are measured with instruments and expressed using numerical values and standard scientific units:

  • Mass ($m$): Measure of the amount of matter in a sample ($g$, $kg$).

  • Volume ($V$): Measure of the three-dimensional space occupied by a substance ($cm^3$, $mL$, $L$, $m^3$).

  • Temperature: Measure of the average kinetic energy of the constituent particles ($^\circ\text{C}$, $\text{K}$).

  • Dimensions: Measured physical parameters including length, width, height, and diameter ($cm$, $m$).

• Physical Change: A transformation that alters the physical form, shape, size, or state of matter without modifying its chemical identity or creating new substances:

  • Changes of State: Reversible physical transitions between solid, liquid, and gas phases (e.g., melting ice, freezing water, boiling, condensation, sublimation of dry ice from solid carbon dioxide to gas).

  • Mechanical Deformation: Reshaping matter by cutting, grinding, sawing, tearing, or folding (e.g., folding a paper crane, cutting cloth, sawing timber into sawdust, grinding wheat into flour).

  • Dissolving (Dissolution): Dispersion of solute particles among solvent molecules without breaking intramolecular chemical bonds (e.g., dissolving table sugar or salt into water; solute can be recovered completely upon evaporating the solvent).

Chemical Properties and Chemical Changes

• Chemical Property: A characteristic of a substance that describes its potential to undergo alterations in chemical composition and react with other materials to produce one or more entirely new substances:

  • Combustibility / Flammability: The ability of a substance to react rapidly with oxygen, releasing thermal energy and light (e.g., burning charcoal, combustion of methane gas).

  • Reactivity with Water: Tendency of elements to undergo vigorous oxidation-reduction reactions upon contact with liquid water (e.g., alkali metals generating aqueous hydroxides and flammable hydrogen gas).

  • Reactivity with Acids: Generation of salt and gaseous products upon acidic contact (e.g., carbonates releasing carbon dioxide gas; metals displacing hydrogen).

  • Corrosion Resistance: Resistance to environmental oxidation (e.g., gold's non-reactivity versus iron's rapid rusting).

• Chemical Change: A transformation in which starting substances (reactants) undergo the breaking and reforming of chemical bonds, rearranging atoms to form one or more new substances (products) with distinct chemical compositions and physical properties.

• Primary Clues and Evidence of Chemical Change:

  • Change of Colour: Appearance of a distinct hue indicative of new chromophores (e.g., oxidation of yellow dye to indigo blue in denim fibres; red-orange copper oxidizing to green copper carbonate patina; clear egg whites denaturing to opaque white upon cooking).

  • Change of Odour: Evolution of detectable volatile compounds (e.g., composting organic waste, baking cake batter giving off aroma, souring milk).

  • Formation of Gas (Bubbles): Production of gaseous products not attributable to boiling or physical heating (e.g., effervescence observed when vinegar reacts with baking soda to liberate carbon dioxide gas).

  • Formation of a Precipitate: Separation of an insoluble powdery or crystalline solid phase out of a liquid solution during a reaction (e.g., limewater turning milky when exposed to carbon dioxide gas due to solid calcium carbonate precipitation).

  • Release or Absorption of Energy: Substantial thermal energy, light emission, or sound generation (e.g., pyrotechnic firework explosions, the brilliant glow of ignited magnesium ribbon, chemiluminescence in a glowstick).

• Reversibility Principle: While most physical changes are easily reversed by physical means (e.g., refreezing melted wax), most chemical changes cannot be reversed without subsequent chemical reactions because chemical bonds have been broken and rearranged.

Laboratory Safety and WHMIS Standards

• Individual and Collective Responsibility: Laboratory safety requires adherence to rigorous protocols to protect the individual experimenter, peers, and instructors.

• Essential Laboratory Safety Equipment:

  • Emergency exit routes and floor plans.

  • Eyewash stations and emergency deluge showers (flushing for a continuous $15\,\text{min}$ period in the event of ocular or cutaneous exposure).

  • Chemical splash goggles and protective laboratory aprons.

  • Class ABC dry chemical fire extinguishers, fire blankets, and sand buckets.

  • Dedicated broken glass disposal containers.

  • Material Safety Data Sheet (MSDS) / Safety Data Sheet (SDS) documentation binders.

• Standard Laboratory Protocols:

  • Chemical splash goggles and lab aprons must remain on at all times during investigations; goggles must never be removed due to lens fogging.

  • Long hair must be tied securely back, loose garments and dangling accessories tucked in, and open-toed shoes avoided.

  • Eating, drinking, and chewing gum are prohibited in laboratories to prevent chemical ingestion.

  • Never taste any substance in a laboratory.

  • Chemical odours must be evaluated exclusively by wafting (gently sweeping vapours toward the nose with a hand) rather than inhaling directly.

  • Point open test tubes away from all individuals when heating liquids over a Bunsen burner flame.

  • Never return unused reagents to original stock bottles to prevent cross-contamination.

  • Inspect glassware for cracks or star fractures prior to heating; cracked beakers must be disposed of in glass disposal bins.

  • Disconnect electrical devices by pulling directly on the insulated plug casing rather than pulling the power cord.

• WHMIS (Workplace Hazardous Materials Information System):

  • Comprehensive Canadian national standard for communicating information regarding hazardous workplace chemicals.

  • Core classifications include:

  • Compressed Gas (cylinder pictogram).

  • Flammable and Combustible Material (flame pictogram).

  • Oxidizing Material (flame over circle).

  • Poisonous and Infectious Material Causing Immediate and Serious Toxic Effects (skull and crossbones).

  • Toxic and Infectious Material Causing Other Toxic Effects (exclamation point / stylized silhouette).

  • Biohazardous Infectious Material (trefoil biological hazard emblem).

  • Corrosive Material (liquids spilling onto a hand and metal plate).

  • Dangerously Reactive Material (exploding canister/bomb emblem).

Characteristic Physical Properties and Density

• Concept of Characteristic Physical Property: A physical property that is unique to a pure substance and remains invariant regardless of sample size, mass, or shape, allowing it to serve as a reliable diagnostic identifier.

• The Three Primary Characteristic Physical Properties:

  • Density

  • Melting / Freezing Point

  • Boiling Point

• Density ($D$ or $\rho$):

  • Quantitative measure defined as the ratio of a substance's mass to its volume:     $D = \frac{m}{V}$

  • Mass represents the quantity of matter contained within the object ($g$ or $kg$).

  • Volume represents the amount of physical space occupied ($cm^3$, $mL$, or $m^3$).

  • Equivalence of volumetric units: $1\,cm^3 = 1\,mL$.

  • Standard units: $g/cm^3$ for solids; $g/mL$ for liquids; $kg/m^3$ for bulk substances.

  • Influence of Temperature and Pressure:

  • Solids and liquids undergo minor density alterations with temperature due to subtle thermal expansion.

  • Gas density is highly sensitive to external conditions: increasing pressure compresses gas particles closer together (increasing density); increasing temperature drives gas particles farther apart (decreasing density).

• Published Densities of Metals and Common Substances (at Room Temperature and Atmospheric Pressure):

  • Aluminum: $2.70\,g/cm^3$

  • Zinc: $7.13\,g/cm^3$

  • Iron: $7.87\,g/cm^3$ (or $7.9\,g/cm^3$)

  • Copper: $8.96\,g/cm^3$ (or $8.90\,g/cm^3$)

  • Silver: $10.49\,g/cm^3$ (or $10.5\,g/cm^3$)

  • Lead: $11.36\,g/cm^3$ (or $11.3\,g/cm^3$)

  • Mercury: $13.55\,g/cm^3$ (or $13.53\,g/cm^3$, $13.5\,g/cm^3$)

  • Gold: $19.32\,g/cm^3$ (or $19.30\,g/cm^3$, $19.3\,g/cm^3$)

  • Platinum: $21.4\,g/cm^3$

  • Maple Wood: $0.65\,g/cm^3$

  • Gasoline: $0.89\,g/cm^3$

  • Pure Liquid Water: $1.00\,g/cm^3$ ($1.0\,g/mL$)

  • Ethanol (Pure Alcohol): $0.79\,g/mL$

  • Solid Ice: $0.92\,g/cm^3$

  • Iron Pyrite (Fool's Gold): $5.02\,g/cm^3$

• Analysis of Mass versus Volume Graphs:

  • Slope Representation:     $\text{Slope} = \frac{\Delta \text{Mass}}{\Delta \text{Volume}} = \text{Density}$

  • A steeper slope corresponds directly to a greater density. On the coordinate graph, the line for gold ($19.3\,g/cm^3$) is steeper than mercury ($13.5\,g/cm^3$), which is steeper than silver ($10.5\,g/cm^3$).

  • All linear mass-volume curves pass through the origin $(0, 0)$ because a volume of zero contains zero mass.

• Buoyancy and Immiscible Fluid Layering:

  • An object sinks if its density exceeds that of the surrounding fluid ($D_{\text{object}} > D_{\text{fluid}}$) and floats if its density is lower ($D_{\text{object}} < D_{\text{fluid}}$).

  • In a column containing gasoline ($0.89\,g/cm^3$), water ($1.0\,g/cm^3$), and liquid mercury ($13.5\,g/cm^3$):

  • Top layer: Gasoline ($0.89\,g/cm^3$).

  • Middle layer: Water ($1.0\,g/cm^3$).

  • Bottom layer: Mercury ($13.5\,g/cm^3$).

  • Silver Dollar ($10.5\,g/cm^3$): Sinks through gasoline and water, but floats on top of the liquid mercury layer ($10.5 < 13.5$).

  • Solid Platinum ($21.4\,g/cm^3$): Sinks through all three liquids to the container floor ($21.4 > 13.5$).

  • Solid Copper ($8.96\,g/cm^3$): Floats on mercury ($8.96 < 13.55$).

• The G.R.A.S.P. Method for Quantitative Problem Solving:

  • Given: State known numerical quantities with explicit units.

  • Required: Identify the unknown target variable.

  • Analysis: Write down the algebraic formula rearranged for the target variable.

  • Solution: Substitute numerical values into the formula and solve.

  • Paraphrase / Statement: Write a final sentence stating the computed answer with proper units.

• Worked Density Problem Examples:

  • Excavated Crown Identification:

  • Given: $m = 2.00\,kg = 2000\,g$, $V = 225\,cm^3$.

  • Required: Density ($D$), comparison to gold ($19.30\,g/cm^3$) or copper ($8.90\,g/cm^3$).

  • Analysis: $D = \frac{m}{V}$.

  • Solution: $D = \frac{2000\,g}{225\,cm^3} = 8.89\,g/cm^3$.

  • Statement: The calculated density of $8.89\,g/cm^3$ matches copper ($8.90\,g/cm^3$); the crown is constructed of copper, not gold.

  • Aluminum Cube:

  • Given: $m = 2.0\,kg = 2000\,g$, $V = 741\,cm^3$.

  • Required: Density ($D$) in $g/cm^3$.

  • Analysis: $D = \frac{m}{V}$.

  • Solution: $D = \frac{2000\,g}{741\,cm^3} = 2.70\,g/cm^3$.

  • Statement: The density of the aluminum cube is $2.70\,g/cm^3$.

  • Bulk Iron Block:

  • Given: $D = 7.9\,g/cm^3 = 7900\,kg/m^3$, $V = 1.0\,m^3$.

  • Required: Mass ($m$) in kilograms.

  • Analysis: $m = D \times V$.

  • Solution: $m = 7900\,kg/m^3 \times 1.0\,m^3 = 7900\,kg$.

  • Statement: The mass of the $1.0\,m^3$ iron block is $7900\,kg$.

  • Water Displacement of Lead Weights:

  • Given: $V_{\text{initial}} = 20.0\,mL$, $V_{\text{final}} = 42.0\,mL$, $m = 250\,g$.

  • Required: Density of lead ($D$).

  • Analysis: $V = V_{\text{final}} - V_{\text{initial}} = 42.0\,mL - 20.0\,mL = 22.0\,mL = 22.0\,cm^3$, $D = \frac{m}{V}$.

  • Solution: $D = \frac{250\,g}{22.0\,cm^3} = 11.36\,g/cm^3$.

  • Statement: The density of the lead sample is $11.36\,g/cm^3$.

• Phase Change Temperatures:

  • Freezing Point: The invariant temperature at which a substance changes from liquid to solid state.

  • Melting Point: The temperature at which a substance changes from solid to liquid state. For any pure substance, melting point and freezing point are the exact same temperature (e.g., pure water melts and freezes at $0\,^\circ\text{C}$).

  • Boiling Point: The temperature at which a substance rapidly transitions from liquid to gas throughout the liquid volume (e.g., pure water boils at $100\,^\circ\text{C}$ at standard atmospheric pressure).

  • Melting Points of Common Metals:

  • Mercury: $-39\,^\circ\text{C}$ (liquid at room temperature, conducts electricity, used in mechanical thermostats).

  • Tin: $232\,^\circ\text{C}$

  • Lead: $328\,^\circ\text{C}$

  • Zinc: $420\,^\circ\text{C}$

  • Aluminum: $660\,^\circ\text{C}$

  • Silver: $961\,^\circ\text{C}$

  • Gold: $1065\,^\circ\text{C}$

  • Copper: $1085\,^\circ\text{C}$

  • Nickel: $1453\,^\circ\text{C}$

  • Iron: $1535\,^\circ\text{C}$

  • Platinum: $1772\,^\circ\text{C}$

  • Tungsten: $3407\,^\circ\text{C}$ (highest melting point; ideal for incandescent lamp filaments).

• Freezing Point Depression and Applications:

  • The introduction of dissolved solutes (such as sodium chloride, $\text{NaCl}$) disrupts regular crystal lattice formation in solidifying solvents.

  • A $20\%$ aqueous salt solution lowers the freezing point of water to $-16\,^\circ\text{C}$, allowing road surfaces to remain liquid at sub-zero temperatures.

  • Traditional ice cream freezing employs rock salt mixed with crushed ice to depress the temperature of the bath below the freezing point of cream.

• The Anomalous Behaviour of Water:

  • Unlike almost all other materials, solid water (ice) is less dense than its liquid form near its freezing point ($0.92\,g/cm^3$ versus $1.00\,g/cm^3$), causing ice to float.

  • Water achieves maximum density at $4\,^\circ\text{C}$ ($1.000\,g/cm^3$). As liquid water cools from $50\,^\circ\text{C}$ down to $4\,^\circ\text{C}$, it contracts and densifies, sinking to lake bottoms.

  • Cooling below $4\,^\circ\text{C}$ down to $0\,^\circ\text{C}$ causes water molecules to expand into an open hexagonal crystalline framework held by directional hydrogen bonds, increasing volume while maintaining constant mass.

  • Ecological and Practical Significance:

  • In winter, lakes freeze from the top down; floating surface ice acts as an insulating barrier, while bottom waters remain liquid at $4\,^\circ\text{C}$, allowing aquatic ecosystems to survive.

  • Destructive expansion: Confined water expanding upon freezing generates substantial hydraulic force, bursting plumbing pipes, cracking glass containers, and degrading roadways.

Historical Development of Atomic Theory

• Scientific Theory Definition: An evidence-based, scientifically tested explanation of natural phenomena that evolves and refines as new observational technologies emerge.

• Democritus (c. 400 BCE):

  • Proposed that matter cannot be divided indefinitely and consists of ultimate, indivisible particles termed atomos ("cannot be cut").

  • Postulated that atoms exist in various sizes and geometric shapes, remain in continuous motion, and are separated by empty voids.

• Aristotle (c. 450 BCE):

  • Rejected atomism, proposing instead that all matter comprises four primordial elements: Earth, Water, Air, and Fire.

  • Assigned four elemental properties: dry (Earth), wet (Water), cold (Earth/Water), and hot (Air/Fire). Held dominance for nearly two millennia.

• John Dalton (1807) – The Billiard Ball Model:

  • Formulated the first modern atomic theory based on mass ratios and chemical observations:

  • All matter is composed of tiny, indivisible spheres called atoms.

  • All atoms of a given element are identical in mass and properties.

  • Atoms of different elements possess different properties and masses.

  • Atoms are rearranged in chemical reactions to yield new substances, but are never created or destroyed.

  • Limitation: Failed to account for electrical phenomena and charge attraction/repulsion.

• J. J. Thomson (1897) – The Plum Pudding Model and the Electron:

  • Experimented with evacuated cathode ray tubes, discovering streams of negatively charged subatomic particles attracted toward positive electrodes.

  • Discovered the electron ($e^-$), proving that atoms are divisible.

  • Postulated that because atoms are electrically neutral, the atom consists of a diffuse, positively charged spherical mass with negatively charged electrons embedded evenly throughout, analogous to raisins in plum pudding or chocolate chips in a muffin.

• Ernest Rutherford (1909–1920) – The Gold Foil Experiment and the Proton:

  • The Gold Foil Experiment (1909): Directed a beam of fast-moving, positively charged alpha particles through a slit toward an ultra-thin sheet of gold foil, monitoring scattering on a zinc sulfide screen.

  • Prediction: Based on Thomson's uniform charge distribution, all alpha particles should pass directly through the foil with little to no deflection.

  • Observation: Most alpha particles passed directly through undisturbed, but a small fraction were deflected at very wide angles, and some bounced almost directly backward.

  • Conclusions:

  • The atom is composed predominantly of empty space.

  • The atom's positive charge and nearly all of its mass are concentrated in an extremely small, dense central core named the nucleus.

  • The nucleus is surrounded by a diffuse cloud of orbiting electrons.

  • Discovery of the Proton (1920): Identified the proton ($p^+$) as a fundamental positively charged nuclear particle.

• James Chadwick (1932) – The Planetary Model and the Neutron:

  • Rutherford's student who discovered a penetrating, electrically neutral nuclear particle ejected during atomic disintegration experiments.

  • Identified the neutron ($n^0$), proving that the nucleus contains both positively charged protons and neutral neutrons of approximately equal mass ($\approx 1\,\text{u}$).

  • Planetary Model: Formulated the model wherein a tiny, dense central nucleus containing protons and neutrons is circled rapidly by electrons traversing the surrounding empty space.

• Niels Bohr (1913) – Quantized Electron Orbits:

  • Analyzed the discrete line-emission spectra of excited hydrogen gas, noting that excited atoms emit specific wavelengths/colours of light rather than a continuous rainbow spectrum.

  • Developed the Bohr-Rutherford Model with the following principles:

  • Electrons travel around the nucleus only in specific, fixed circular pathways called orbits or energy levels.

  • Each electron orbit corresponds to a definite, quantized amount of energy.

  • Energy increases with radial distance from the nucleus.

  • Electrons cannot exist between orbits; they absorb energy to jump to higher orbits and emit specific wavelengths of light when falling from higher to lower orbits.

  • Electron shell capacities for the first 20 elements follow a strict maximum occupancy: $2$ electrons in the first orbit, $8$ electrons in the second orbit, and $8$ electrons in the third orbit.

• Summary Properties of Fundamental Subatomic Particles:

  • Proton ($p^+$): Electric charge $+1$; Relative mass $1$; Location: Nucleus.

  • Neutron ($n^0$): Electric charge $0$; Relative mass $1$; Location: Nucleus.

  • Electron ($e^-$): Electric charge $-1$; Relative mass $\frac{1}{2000}$ (approx. $0.0005$); Location: Electron orbits surrounding nucleus.

Structure of the Periodic Table and Chemical Families

• Organization of the Modern Periodic Table:

  • Contains all 118 known chemical elements arranged in order of increasing Atomic Number ($Z$).

  • Dmitri Mendeleev (1869): Organized the original periodic table of 63 elements by increasing atomic mass, arranging elements with similar physical and chemical properties into vertical columns. Mendeleev left vacant spaces and predicted the existence and properties of undiscovered elements (e.g., Gallium, Germanium), which were later verified experimentally.

  • Groups (Chemical Families): The 18 vertical columns; elements within the same group possess the same number of valence (outermost) electrons and exhibit similar chemical and physical properties.

  • Periods: The 7 horizontal rows; each period indicates the number of occupied electron shells/orbits in an atom.

• Major Structural Divisions:

  • Metals: Located to the left and throughout the center of the periodic table; characterized by metallic lustre, electrical and thermal conductivity, malleability, and ductility. All are solids at room temperature except mercury ($\text{Hg}$, liquid).

  • Non-Metals: Positioned in the upper right quadrant (plus hydrogen); generally dull, brittle solids or gases, and poor conductors of heat and electricity. Bromine ($\text{Br}_2$) is the only liquid non-metal.

  • Metalloids (Semimetals): Form a staircase border starting at Boron ($\text{B}, \text{Si}, \text{Ge}, \text{As}, \text{Sb}, \text{Te}, \text{Po}$); exhibit properties intermediate between metals and non-metals (e.g., semiconductors).

• Profiles of Key Chemical Families:

  • Group 1: Alkali Metals ($\text{Li}, \text{Na}, \text{K}, \text{Rb}, \text{Cs}, \text{Fr}$):

  • Possess $1$ valence electron.

  • Shiny, silvery-white, soft enough to cut with a butter knife, low density (lithium, sodium, and potassium float on water).

  • Highly reactive with moisture and oxygen; violently displace hydrogen gas from water:       $2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)$

  • Reactivity increases down the group ($\text{Cs} > \text{Rb} > \text{K} > \text{Na} > \text{Li}$) as the valence electron is positioned further from the nucleus, shielded by inner shells, and lost more readily.

  • Stored under mineral oil; found in nature only within stable compounds (e.g., $\text{NaCl}$, $\text{NaHCO}_3$).

  • Hydrogen ($\text{H}$): Positioned above Group 1 due to possessing $1$ valence electron, but behaves as a non-metallic, reactive diatomic gas ($\text{H}_2$).

  • Group 2: Alkaline Earth Metals ($\text{Be}, \text{Mg}, \text{Ca}, \text{Sr}, \text{Ba}, \text{Ra}$):

  • Possess $2$ valence electrons.

  • Silvery-white, harder, denser, and less reactive than Group 1 metals.

  • Burn with brilliant, distinctive flame colours (magnesium burns white and is used in flares/fireworks).

  • Calcium ($\text{Ca}$) is an essential biological element for building bones and teeth; strontium ($\text{Sr}$) strengthens coral skeletons.

  • Group 17: Halogens ($\text{F}, \text{Cl}, \text{Br}, \text{I}, \text{At}$):

  • Possess $7$ valence electrons.

  • Highly reactive non-metals seeking $1$ additional electron to attain a stable octet.

  • Exhibit diverse physical states at room temperature: fluorine ($\text{F}_2$) and chlorine ($\text{Cl}_2$) are gases; bromine ($\text{Br}_2$) is a dense liquid; iodine ($\text{I}_2$) and astatine ($\text{At}$) are solid crystals.

  • Reactivity increases up the group ($\text{F} > \text{Cl} > \text{Br} > \text{I}$) due to smaller atomic radii exerting stronger electrostatic attraction for incoming electrons.

  • Highly toxic and corrosive (chlorine was deployed as a chemical weapon in WWI); used at low concentrations as disinfectants (chlorine in pools, iodine tinctures for wounds).

  • Group 18: Noble Gases ($\text{He}, \text{Ne}, \text{Ar}, \text{Kr}, \text{Xe}, \text{Rn}$):

  • Possess completely filled valence shells ($2$ valence electrons for Helium; $8$ for Ne, Ar, Kr, Xe, Rn).

  • Completely unreactive, chemically inert, non-toxic (except radioactive radon), odorless, and colorless gases.

  • Emit characteristic wavelengths when subjected to electrical discharges:

    • Neon ($\text{Ne}$): Bright red-orange

    • Argon ($\text{Ar}$): Blue

    • Krypton ($\text{Kr}$): Pink

    • Xenon ($\text{Xe}$): Purple

  • Helium ($\text{He}$) is non-flammable and low density, making it a safe alternative to hydrogen in airships and balloons.

Atomic Calculations and Bohr-Rutherford Diagrams

• Fundamental Atomic Calculations:

  • Atomic Number ($Z$): Represents the number of protons in an atomic nucleus. Uniquely defines the elemental identity. In any neutral atom:     $\text{Number of protons} = \text{Number of electrons} = Z$

  • Mass Number ($A$): Total sum of protons and neutrons within the nucleus:     $A = \text{Number of protons} + \text{Number of neutrons}$

  • Calculating Neutrons ($N$):     $N = A - Z$

  • Isotopes: Atoms of the same element containing identical proton counts ($Z$) but differing neutron counts ($N$) and mass numbers ($A$) (e.g., Lithium-6 with $3\,p^+, 3\,n^0$ versus Lithium-7 with $3\,p^+, 4\,n^0$).

  • Atomic Mass: The weighted average mass of all naturally occurring isotopes of an element, expressed in atomic mass units ($\text{u}$). Rounding atomic mass to the nearest integer provides the mass number of the predominant natural isotope (e.g., Boron $10.81\,\text{u} \rightarrow \text{B-11}$ with $5\,p^+, 6\,n^0$).

• Comprehensive Table of the First 20 Elements:

  • Hydrogen ($\text{H}$): Atomic Number $= 1$, Mass Number $= 1$, Protons $= 1$, Electrons $= 1$, Neutrons $= 0$

  • Helium ($\text{He}$): Atomic Number $= 2$, Mass Number $= 4$, Protons $= 2$, Electrons $= 2$, Neutrons $= 2$

  • Lithium ($\text{Li}$): Atomic Number $= 3$, Mass Number $= 7$, Protons $= 3$, Electrons $= 3$, Neutrons $= 4$

  • Beryllium ($\text{Be}$): Atomic Number $= 4$, Mass Number $= 9$, Protons $= 4$, Electrons $= 4$, Neutrons $= 5$

  • Boron ($\text{B}$): Atomic Number $= 5$, Mass Number $= 11$, Protons $= 5$, Electrons $= 5$, Neutrons $= 6$

  • Carbon ($\text{C}$): Atomic Number $= 6$, Mass Number $= 12$, Protons $= 6$, Electrons $= 6$, Neutrons $= 6$

  • Nitrogen ($\text{N}$): Atomic Number $= 7$, Mass Number $= 14$, Protons $= 7$, Electrons $= 7$, Neutrons $= 7$

  • Oxygen ($\text{O}$): Atomic Number $= 8$, Mass Number $= 16$, Protons $= 8$, Electrons $= 8$, Neutrons $= 8$

  • Fluorine ($\text{F}$): Atomic Number $= 9$, Mass Number $= 19$, Protons $= 9$, Electrons $= 9$, Neutrons $= 10$

  • Neon ($\text{Ne}$): Atomic Number $= 10$, Mass Number $= 20$, Protons $= 10$, Electrons $= 10$, Neutrons $= 10$

  • Sodium ($\text{Na}$): Atomic Number $= 11$, Mass Number $= 23$, Protons $= 11$, Electrons $= 11$, Neutrons $= 12$

  • Magnesium ($\text{Mg}$): Atomic Number $= 12$, Mass Number $= 24$, Protons $= 12$, Electrons $= 12$, Neutrons $= 12$

  • Aluminum ($\text{Al}$): Atomic Number $= 13$, Mass Number $= 27$, Protons $= 13$, Electrons $= 13$, Neutrons $= 14$

  • Silicon ($\text{Si}$): Atomic Number $= 14$, Mass Number $= 28$, Protons $= 14$, Electrons $= 14$, Neutrons $= 14$

  • Phosphorus ($\text{P}$): Atomic Number $= 15$, Mass Number $= 31$, Protons $= 15$, Electrons $= 15$, Neutrons $= 16$

  • Sulfur ($\text{S}$): Atomic Number $= 16$, Mass Number $= 32$, Protons $= 16$, Electrons $= 16$, Neutrons $= 16$

  • Chlorine ($\text{Cl}$): Atomic Number $= 17$, Mass Number $= 35$, Protons $= 17$, Electrons $= 17$, Neutrons $= 18$

  • Argon ($\text{Ar}$): Atomic Number $= 18$, Mass Number $= 40$, Protons $= 18$, Electrons $= 18$, Neutrons $= 22$

  • Potassium ($\text{K}$): Atomic Number $= 19$, Mass Number $= 39$, Protons $= 19$, Electrons $= 19$, Neutrons $= 20$

  • Calcium ($\text{Ca}$): Atomic Number $= 20$, Mass Number $= 40$, Protons $= 20$, Electrons $= 20$, Neutrons $= 20$

• Rules for Constructing Bohr-Rutherford Diagrams:

  • Draw a central circle representing the nucleus and write the exact number of protons ($p^+$) and neutrons ($n^0$) inside.

  • Draw concentric circles around the nucleus representing electron orbits/shells.

  • Populate electron orbits sequentially from innermost to outermost, following the $2, 8, 8$ maximum rule for the first 20 elements:

  • Orbit 1: Maximum $2\,e^-$

  • Orbit 2: Maximum $8\,e^-$

  • Orbit 3: Maximum $8\,e^-$

  • Orbit 4: Holds the remaining $1$ or $2$ electrons for Potassium and Calcium.

  • Place electrons individually at four equidistant compass points (North, East, South, West) before pairing them.

Allotropes of Carbon and Materials Science

• Allotrope Definition: Structurally distinct physical forms of the same pure element, arising from differences in how atoms bond and arrange in space.

• Charcoal:

  • An amorphous, non-crystalline, disorganized arrangement of carbon atoms mixed with ash and other compounds (up to $98\%$ carbon).

  • Black, porous, opaque, crumbly solid that burns easily, serving primarily as fuel.

• Graphite:

  • Organized two-dimensional structure where each carbon atom is covalently bonded to three neighboring carbon atoms, forming planar sheets of interconnected hexagons.

  • Layers are held together by weak intermolecular attractions, allowing them to slide over one another under shear stress (basis of pencil markings and dry lubricants).

  • Possesses delocalized valence electrons that move between sheets, making graphite an electrical conductor (unusual for non-metals; used in battery electrodes and electroplating).

• Diamond:

  • Formed under high pressure and temperature; each carbon atom is covalently bonded to four adjacent carbon atoms in a continuous, three-dimensional tetrahedral lattice.

  • Rigid structure confers extreme hardness (10 on the Mohs scale).

  • High optical refraction creates significant light dispersion and sparkle; lack of mobile electrons makes diamond an electrical insulator.

• Canadian Diamond Industry and Environmental-Societal Context:

  • Diamonds were discovered in the Northwest Territories during the 1980s; Canada is now a top-three global producer (e.g., Diavik Mine on Lac de Gras, NWT), generating over $15\%$ of the global diamond supply.

  • Canadian diamonds are certified conflict-free and laser-inscribed with a microscopic polar bear trademark, distinguishing them from African "blood diamonds."

  • Economic impacts: High-paying jobs and northern economic growth alongside housing shortages and high living costs in Yellowknife and Iqaluit.

  • Environmental concerns: Lake drainage, stream diversions, permafrost degradation, and aquatic habitat loss.

Chemical Compounds: Ionic versus Molecular Bonding

• Ionic Compounds:

  • Formed via the transfer of valence electrons from metal atoms to non-metal atoms.

  • Metals lose electrons to form positively charged cations ($M \rightarrow M^{n+} + n\,e^-$).

  • Non-metals gain electrons to form negatively charged anions ($X + m\,e^- \rightarrow X^{m-}$).

  • Ionic Bond: The electrostatic attraction holding oppositely charged cations and anions together in an alternating, three-dimensional crystal lattice.

  • Properties: High melting and boiling points, brittle crystalline solids, non-conductive when dry, but conduct electricity when dissolved in water or molten (acting as electrolytes).

• Common Ionic Compounds:

  • Sodium Chloride ($\text{NaCl}$): Table salt, highway de-icing agent.

  • Potassium Chloride ($\text{KCl}$): Potash, agricultural fertilizer.

  • Calcium Oxide ($\text{CaO}$): Quicklime, masonry and mortar.

  • Sodium Hydroxide ($\text{NaOH}$): Lye, industrial drain cleaner.

  • Calcium Carbonate ($\text{CaCO}_3$): Limestone, marble, chalk, antacids.

  • Sodium Hydrogen Carbonate ($\text{NaHCO}_3$): Baking soda, leavening agent.

  • Magnesium Hydroxide ($\text{Mg(OH)}_2$): Milk of magnesia, antacid suspension.

  • Copper(II) Sulfate ($\text{CuSO}_4$): Bluestone, algicide and fungicide.

• Molecular Elements and Diatomic Molecules:

  • Molecular Element: A molecule composed entirely of identical atoms of a single chemical element.

  • Diatomic Molecules: Molecules composed of exactly two chemically bonded atoms. The seven diatomic elements can be remembered by the mnemonic HOFBrINCl:

  • Hydrogen ($\text{H}_2$)

  • Oxygen ($\text{O}_2$)

  • Fluorine ($\text{F}_2$)

  • Bromine ($\text{Br}_2$)

  • Iodine ($\text{I}_2$)

  • Nitrogen ($\text{N}_2$)

  • Chlorine ($\text{Cl}_2$)

• Molecular Compounds and Covalent Bonding:

  • Formed when two or more non-metal atoms share pairs of valence electrons.

  • Covalent Bond: A chemical bond resulting from the mutual electrostatic attraction of two atomic nuclei for a shared pair of electrons.

  • Combining Capacities of Non-metals:

  • Hydrogen ($\text{H}$): 1 covalent bond

  • Chlorine ($\text{Cl}$): 1 covalent bond

  • Oxygen ($\text{O}$): 2 covalent bonds

  • Sulfur ($\text{S}$): 2 covalent bonds

  • Nitrogen ($\text{N}$): 3 covalent bonds

  • Carbon ($\text{C}$): 4 covalent bonds

  • Common Molecular Compounds:

  • Water ($\text{H}_2\text{O}$): Dihydrogen monoxide.

  • Ozone ($\text{O}_3$): Trioxygen; stratospheric UV shield and ground-level pollutant.

  • Carbon Dioxide ($\text{CO}_2$): Cellular respiration byproduct; solid phase is dry ice.

  • Carbon Monoxide ($\text{CO}$): Toxic, odorless gas formed during incomplete combustion in low-oxygen environments.

  • Hydrogen Chloride ($\text{HCl}$): Acid gas; aqueous solution is muriatic acid.

  • Methane ($\text{CH}_4$): Primary component of natural gas.

  • Ammonia ($\text{NH}_3$): Nitrogen trihydride; common household cleaning agent.

  • Propane ($\text{C}_3\text{H}_8$): Barbecue and heating fuel.

  • Butane ($\text{C}_4\text{H}_{10}$): Cigarette lighter fluid.

  • Acetic Acid ($\text{C}_2\text{H}_4\text{O}_2$ or $\text{CH}_3\text{COOH}$): Active acid component of vinegar.

  • Ethanol ($\text{C}_2\text{H}_6\text{O}$ or $\text{C}_2\text{H}_5\text{OH}$): Grain alcohol.

  • Acetylsalicylic Acid / ASA ($\text{C}_9\text{H}_8\text{O}_4$): Aspirin.

  • Sucrose ($\text{C}_{12}\text{H}_{22}\text{O}_{11}$): Table sugar.

  • Glucose ($\text{C}_6\text{H}_{12}\text{O}_6$): Monosaccharide carbohydrate.

  • Saturated Hydrocarbons (Alkanes): Follow the homologous general formula $\text{C}_n\text{H}_{2n+2}$.

Chemical Formulas and Atom Counting

• Rules for Counting Atoms in Chemical Formulas:

  • Element symbols designate the identity of each element present.

  • Subscripts immediately following an element symbol indicate the number of atoms of that element; if no subscript is written, it indicates exactly $1$ atom.

  • Subscripts positioned outside brackets multiply all atom counts inside those brackets.

  • Leading coefficients multiply all constituent atoms throughout the entire following formula unit.

• Comprehensive Formula Analysis Examples:

  • $\text{Na}_2\text{CO}_3$ (Sodium Carbonate):

  • $\text{Na} = 2$

  • $\text{C} = 1$

  • $\text{O} = 3$

  • Total atoms: $2 + 1 + 3 = 6\,\text{atoms}$

  • $\text{Ca}_3(\text{PO}_4)_2$ (Calcium Phosphate):

  • $\text{Ca} = 3$

  • $\text{P} = 1 \times 2 = 2$

  • $\text{O} = 4 \times 2 = 8$

  • Total atoms: $3 + 2 + 8 = 13\,\text{atoms}$

  • $\text{NH}_4\text{C}_2\text{H}_3\text{O}_2$ (Ammonium Acetate):

  • $\text{N} = 1$

  • $\text{H} = 4 + 3 = 7$

  • $\text{C} = 2$

  • $\text{O} = 2$

  • Total atoms: $1 + 7 + 2 + 2 = 12\,\text{atoms}$

  • $3\text{BaCl}_2$ (Barium Chloride):

  • $\text{Ba} = 3 \times 1 = 3$

  • $\text{Cl} = 3 \times 2 = 6$

  • Total atoms: $3 + 6 = 9\,\text{atoms}$

  • $\text{C}_6\text{H}_{12}\text{O}_3$:

  • $\text{C} = 6$

  • $\text{H} = 12$

  • $\text{O} = 3$

  • Total atoms: $6 + 12 + 3 = 21\,\text{atoms}$

  • $4\text{Al}_2(\text{CO}_3)_3$ (Aluminum Carbonate):

  • $\text{Al} = 4 \times 2 = 8$

  • $\text{C} = 4 \times (1 \times 3) = 12$

  • $\text{O} = 4 \times (3 \times 3) = 36$

  • Total atoms: $8 + 12 + 36 = 56\,\text{atoms}$

  • $\text{Pb}(\text{NO}_3)_2$ (Lead(II) Nitrate):

  • $\text{Pb} = 1$

  • $\text{N} = 1 \times 2 = 2$

  • $\text{O} = 3 \times 2 = 6$

  • Total atoms: $1 + 2 + 6 = 9\,\text{atoms}$

  • $4(\text{NH}_4)_2\text{Cr}_2\text{O}_7$ (Ammonium Dichromate):

  • $\text{N} = 4 \times (1 \times 2) = 8$

  • $\text{H} = 4 \times (4 \times 2) = 32$

  • $\text{Cr} = 4 \times 2 = 8$

  • $\text{O} = 4 \times 7 = 28$

  • Total atoms: $8 + 32 + 8 + 28 = 76\,\text{atoms}$

Gas Identification Tests, Chemical Reactions, and Catalysis

• Diagnostic Laboratory Tests for Common Gases:

  • Oxygen ($\text{O}_2$): Rekindles a glowing wooden splint.

  • Hydrogen ($\text{H}_2$): Explodes with a distinctive "pop" sound when exposed to a burning wooden splint, producing water vapour:     $2H_2(g) + O_2(g) \rightarrow 2H_2O(g)$

  • Carbon Dioxide ($\text{CO}_2$): Extinguishes a flame; turns clear limewater (aqueous calcium hydroxide, $\text{Ca(OH)}_2$) cloudy through the formation of a fine white precipitate of calcium carbonate:     $CO_2(g) + Ca(OH)_2(aq) \rightarrow CaCO_3(s) + H_2O(l)$

• Decomposition of Hydrogen Peroxide:

  • Chemical equation:     $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$

  • The extra oxygen atom in hydrogen peroxide makes the molecule unstable relative to water and oxygen.

  • Bleaching Action: Released oxygen reacts with organic pigments, disrupting conjugated double bonds and lightening hair or dental stains.

  • Antiseptic Action: Effervescing oxygen bubbles destroy anaerobic bacteria in cuts.

  • Catalysis: A catalyst accelerates a chemical reaction without being consumed or altered chemically. Catalase enzymes in yeast or living tissues, or manganese dioxide ($\text{MnO}_2$), rapidly speed up the decomposition of $\text{H}_2\text{O}_2$.

  • Storage: Packaged in opaque, expandable, vented plastic containers to prevent light-induced decomposition and explosive pressure accumulation.

• Corrosion Chemistry:

  • Rusting of Iron:     $4Fe(s) + 3O_2(g) + 6H_2O(l) \rightarrow 4Fe(OH)_3(s) \rightarrow 2Fe_2O_3 \cdot 3H_2O(s)$

  • Requires both oxygen and liquid water; accelerated by dissolved salts acting as electrolytes.

  • Unlike passivating oxides, iron rust is porous and flakes off, continuously exposing underlying metal to progressive degradation.

  • The Eiffel Tower requires $60\,\text{tonnes}$ of protective paint across $220,000\,m^2$ every seven years at a cost of roughly five million dollars to prevent structural failure.

  • Passivation of Copper:

  • Clean copper oxidizes to brown-black copper oxide ($\text{CuO}$), which slowly reacts with atmospheric moisture and carbon dioxide to form a protective green patina of basic copper carbonate ($\text{CuCO}_3 \cdot \text{Cu(OH)}_2$).

  • Patina adheres firmly to the underlying metal, shielding it from further decay (as seen on Canada's Parliament building roofs and the Statue of Liberty).

  • During the Statue of Liberty's 1980s restoration, galvanic and environmental corrosion of its internal iron armature was found to be accelerated by marine salt spray.

Environmental and Societal Impacts of Chemicals

• Heavy Metal Toxicity:

  • Mercury ($\text{Hg}$):

  • Neurological poison that bioaccumulates in aquatic food webs.

  • Historically caused "Mad Hatter's Disease" among felt hat manufacturers using mercury nitrate solutions (symptoms included tremors, ataxia, memory loss, anxiety, and slurred speech).

  • Industrial effluent discharged into the English-Wabigoon River system poisoned the Grassy Narrows First Nation community in northwestern Ontario, leading to the closure of local commercial fisheries and chronic neurological illness.

  • Found in older thermometers, fluorescent light bulbs, and dental amalgams; major global emissions stem from coal-fired power stations.

  • Lead ($\text{Pb}$):

  • Toxic heavy metal that causes irreversible cognitive and organ damage, particularly in children.

  • Phased out from plumbing pipes, automobile gasoline, and food-can solder.

• Tantalum Mining and Habitat Degradation:

  • Tantalum is an indispensable metal for capacitors in cellphones, laptops, and consumer electronics due to its ability to store large electrical charges in small volumes.

  • Mined from coltan ores in central African rainforests, leading to habitat fragmentation and a $70\%$ decline in eastern lowland gorilla populations due to forest clearing and bushmeat hunting.

  • Electronics recycling programs (e.g., the EcoCell program at the Toronto Zoo) recover tantalum and reduce mining pressure.

• Synthetic Compounds and Environmental Persistence:

  • Plastics: Petroleum-derived synthetic polymers that are durable, flexible, and non-biodegradable; accumulate in landfills and marine environments (such as the Great Pacific Garbage Patch), washing up on remote coastlines like Funafuti Atoll in Tuvalu.

  • Persistent Organic Pollutants (POPs): Synthetic pesticides (such as DDT) and industrial chemicals (such as PCBs) that resist environmental degradation, biomagnifying through food webs into top carnivores and northern Arctic human populations.