Year 9 Combined Chemistry: Atomic Structure, Periodic Table, and Bonding Notes

1.1 Atoms, Elements and Compounds

• Atom: Defined as the smallest part of an element that can exist and still retain the chemical properties of that element. Atoms are composed of three subatomic particles: protons, neutrons, and electrons.
• Element: A pure substance consisting of only one type of atom. Elements are fundamental in that they cannot be broken down into simpler substances through chemical reactions. There are currently $118$ known elements.
• Compound: A substance formed through the chemical combination of two or more elements in fixed proportions. Compounds exhibit properties distinct from the elements they comprise and can only be separated back into those elements via chemical reactions.
• Mixture: Consists of two or more substances that are not chemically bonded. In a mixture, substances retain their individual properties and can be separated using physical methods.
• Comparison of Compounds vs. Mixtures:     * Bonding: Elements in a compound are chemically bonded in fixed ratios (e.g., Water is always $H_2O$, meaning $2$ hydrogen atoms to $1$ oxygen atom). Mixtures contain substances in variable proportions without chemical bonds.     * Properties: Compounds have a fixed set of properties, including a single melting/boiling point. Mixtures exhibit a range of properties based on their composition.     * Separation: Mixtures are separable by physical methods; compounds require chemical reactions.
• Chemical Symbols and Formulae:     * Every element has a unique chemical symbol consisting of one or two letters (the first is always capitalized).     * Formulae indicate the specific elements in a compound and their ratios. For example, $H_2O$ denotes $2$ hydrogen atoms bonded to $1$ oxygen atom.     * Word Equations: Follow the format $\text{reactants} \rightarrow \text{products}$ (e.g., $\text{sodium} + \text{water} \rightarrow \text{sodium hydroxide} + \text{hydrogen}$).     * Symbol Equations: Use formulae and must be balanced, ensuring the same number of atoms exist on both sides of the equation.

1.2 Separating Mixtures

• Key Vocabulary:     * Solvent: The liquid in which a solute dissolves (e.g., water).     * Solute: The substance that dissolves within the solvent.     * Solution: The mixture formed by the dissolved solute and solvent.     * Soluble: The ability of a substance to dissolve in a solvent.     * Insoluble: A substance that cannot dissolve and remains solid.
• Filtration:     * Purpose: Separates an insoluble solid from a liquid (e.g., sand from water).     * Mechanism: The mixture is poured through filter paper in a funnel. The liquid (filtrate) passes through, while the solid (residue) is trapped. The solid must be insoluble for this to work.
• Crystallisation:     * Purpose: Separates a soluble salt from a solution (e.g., copper sulfate from water).     * Mechanism: The solution is heated to evaporate water and increase concentration, then left to cool. As solubility decreases with temperature, the solute crystallizes. Slow cooling yields larger, purer crystals.
• Simple Distillation:     * Purpose: Separates a solvent from a solution or liquids with significantly different boiling points (e.g., water from salt water).     * Mechanism: Heating evaporates the liquid with the lower boiling point. The vapour enters a condenser, is cooled, and turns back into a liquid (distillate) for collection.
• Fractional Distillation:     * Purpose: Separates mixtures of liquids with similar boiling points (e.g., ethanol and water, crude oil fractions).     * Mechanism: Uses a fractionating column with a temperature gradient (hotter at base, cooler at top). Components condense at the height corresponding to their specific boiling point.
• Chromatography:     * Purpose: Separates dissolved substances, such as dyes or inks.     * Mechanism: A spot of mixture is placed on paper above a solvent. Capillary action carries the solvent and substances up the paper. Solubility and attraction to the paper determine how far a substance travels.     * Rf Value Calculation:         * $Rf = \frac{\text{distance travelled by substance}}{\text{distance travelled by solvent front}}$         * Values are always between $0$ and $1$ and have no units.         * Example: If a substance travels $4.5\text{ cm}$ and the solvent front travels $9.0\text{ cm}$, then $Rf = \frac{4.5}{9.0} = 0.5$.

1.3 Development of the Atomic Model

• Solid Sphere Model (Dalton, early 1800s): John Dalton proposed atoms as tiny, indivisible solid spheres. He suggested different elements were made of different spheres. He could not explain electrical behavior.
• Plum Pudding Model (Thomson, 1897): J.J. Thomson discovered the electron. He modeled the atom as a positively charged sphere with negatively charged electrons embedded like plums in a pudding.
• Nuclear Model (Rutherford, 1911): Ernest Rutherford performed the alpha particle scattering experiment.     * The Experiment: Positively charged alpha particles were fired at thin gold foil in a vacuum.     * Observation 1: Most passed through (Conclusion: Atoms are mostly empty space).     * Observation 2: Some deflected at large angles (Conclusion: A concentrated positive charge/nucleus exists).     * Observation 3: A few bounced back (Conclusion: The nucleus is tiny and very dense).
• Bohr Model (Bohr, 1913): Niels Bohr used line emission spectra to propose that electrons orbit the nucleus in fixed energy levels (shells). Electrons emit or absorb energy when moving between these shells.
• Modern Nuclear Model (Chadwick, 1932): James Chadwick discovered the neutron (an uncharged particle with the same relative mass as a proton). This explained isotopes (atoms of same element with different masses).

1.4 Structure of the Atom

• Subatomic Particles:     * Proton: Located in the nucleus; Relative charge = $+1$; Relative mass = $1$.     * Neutron: Located in the nucleus; Relative charge = $0$; Relative mass = $1$.     * Electron: Located in shells; Relative charge = $-1$; Relative mass = $\frac{1}{1836}$ (approx. $0$).
• Calculations and Definitions:     * Atomic number (Z): The number of protons. Defines the element. In neutral atoms, this equals the number of electrons.     * Mass number (A): The total of protons + neutrons.     * $\text{Number of neutrons} = A - Z$     * Isotopes: Atoms of the same element (same protons) but different neutrons (different mass number). They share chemical properties but have different physical properties (e.g., density).         * Carbon-12: $6$ p, $6$ n, $6$ e.         * Carbon-13: $6$ p, $7$ n, $6$ e.         * Carbon-14: $6$ p, $8$ n, $6$ e.     * Relative atomic mass (Ar): Weighted average mass of all isotopes of an element compared to $1/12$ of a carbon-12 atom.         * $Ar = \frac{\text{sum of (mass of isotope} \times \text{percentage abundance)}}{100}$         * Example (Chlorine): $Cl-35$ (75\text{%}) and $Cl-37$ (25\text{%}). $Ar = \frac{(35 \times 75 + 37 \times 25)}{100} = 35.5$.
• Ions: Atoms that have lost or gained electrons. Metals lose electrons to form positive ions (cations); non-metals gain electrons to form negative ions (anions).
• Scale: Atomic sizes are approximately $1 \times 10^{-10}\text{ m}$.

1.5 Electronic Structure

• Filling Rules (First 20 Elements):     * Shell 1 (Innermost): Max $2$ electrons.     * Shell 2: Max $8$ electrons.     * Shell 3: Max $8$ electrons (for first 20 elements).     * Shell 4: Starts filling at Potassium (element $19$).
• Examples:     * Lithium (Li) (Atomic No. $3$): $2, 1$     * Neon (Ne) (Atomic No. $10$): $2, 8$     * Sodium (Na) (Atomic No. $11$): $2, 8, 1$     * Argon (Ar) (Atomic No. $18$): $2, 8, 8$
• Relationship to Periodic Table:     * Group Number: Indicates the number of electrons in the outer shell.     * Period Number: Indicates the total number of electron shells.

1.6 The Periodic Table

• Arrangement: Elements are ordered by increasing atomic number.
• Metals (Left and Center): Conductors of heat/electricity, malleable, ductile, lustrous, high melting points, form positive ions, mostly solid (except Mercury).
• Non-metals (Right): Poor conductors (except graphite), brittle (solid state), dull, low melting points, form negative ions or share electrons.
• Metalloids/Semi-metals: Elements near the stepped diagonal line separating metals from non-metals (e.g., silicon, germanium). They possess mixed properties.

1.7 Development of the Periodic Table

• Early Tables (1860s): Arranged by weight; properties didn't always fit; many undiscovered elements.
• Mendeleev (1869): Arranged by atomic weight but adjusted order for chemical fit. Crucially, he left gaps for undiscovered elements like eka-silicon (germanium) and predicted their properties. His model was supported when those elements were later discovered.
• Modern Table: Ordered by atomic number; structure explained by electronic configuration.

1.8 Group 0: Noble Gases

• Elements: Helium ($He$), Neon ($Ne$), Argon ($Ar$), Krypton ($Kr$), Xenon ($Xe$).
• Properties: Extremely stable and unreactive (inert) due to full outer shells. Exist as monatomic atoms (single atoms).
• Trends Down the Group: Boiling point increases as atoms get larger because intermolecular forces become stronger.     * $He$: $-269\text{ }^\text{o}C$     * $Xe$: $-108\text{ }^\text{o}C$
• Uses: $He$ for balloons (low density, non-flammable); $Ne$ for signs; $Ar$ for light bulbs and welding; $Xe$ for flashlamps.

1.9 Group 1: Alkali Metals

• Elements: Lithium ($Li$), Sodium ($Na$), Potassium ($K$).
• Properties: Soft (cut with a knife), low density (float on water), low melting points. React by losing $1$ electron to form $+1$ ions.
• Reactions:     * With Water: $\text{metal} + \text{water} \rightarrow \text{metal hydroxide} + \text{hydrogen}$. Observations: fizzing, metal movement, alkalinity (indicator turns purple). $K$ ignites with a lilac flame.     * With Oxygen: Forms metal oxide; surface tarnishes to a white/grey coating.     * With Chlorine: Violent reaction forming white metal chloride salts.
• Reactivity Trend: Reactivity increases down the group. This is because the outer electron is further from the nucleus with more shielding, making it easier to lose.

1.10 Group 7: Halogens

• Elements: Fluorine ($F_2$ - pale yellow gas), Chlorine ($Cl_2$ - yellow-green gas), Bromine ($Br_2$ - orange-red liquid), Iodine ($I_2$ - grey-black solid).
• Properties: Non-metals, exist as diatomic molecules. Melting/boiling points increase down the group.
• Reactivity Trend: Reactivity decreases down the group. It is harder for the nucleus to attract an incoming electron because the outer shell is further away and more shielded.
• Displacement Reactions: A more reactive halogen displaces a less reactive halogen from its salt solution.     * $Cl_2 + 2KBr \rightarrow 2KCl + Br_2$ (Chlorine displaces bromide, turning solution orange-brown).     * $I_2 + KCl \rightarrow \text{no reaction}$.

2.1 - 2.3 Ionic Bonding and Compounds

• Mechanism: Transfer of electrons from a metal (loses to become positive) to a non-metal (gains to become negative). Bonded by strong electrostatic attraction.
• Giant Ionic Lattice: A regular repeating 3D arrangement of ions. Forces of attraction act in all directions.
• Properties:     * Melting/Boiling Points: Very high ($NaCl$ melts at $801\text{ }^\text{o}C$) due to strong forces requiring massive energy to break.     * Conductivity: Do not conduct when solid (ions fixed); conduct when molten or dissolved (ions free to move).
• Math Skill Balance: The overall charge must be zero. For Aluminium oxide: $2 \times Al^{3+}$ and $3 \times O^{2-}$ gives $Al_2O_3$.

2.4 - 2.5 Covalent Bonding and Simple Molecules

• Mechanism: Non-metal atoms share pairs of electrons. A single bond ($-$) is one pair; a double bond ($=$) is two pairs.
• Examples: $H_2$, $Cl_2$, $O_2$ ($O=O$), $H_2O$, $CH_4$, $CO_2$ ($O=C=O$), $NH_3$.
• Simple Molecular Substances: Small molecules with strong covalent bonds but weak intermolecular forces.     * Properties: Low melting/boiling points (only weak intermolecular forces are broken); non-conductive (no free charged particles).

2.6 Giant Covalent Structures

• Diamond: Each $C$ bonded to $4$ others (tetrahedral). Extremely hard, non-conductive, very high melting point.
• Graphite: Each $C$ bonded to $3$ others in hexagonal layers. Contains delocalised electrons between layers.     * Properties: Soft/slippery (layers slide); conducts electricity; high melting point.
• Silicon Dioxide ($SiO_2$): Similar structure to diamond ($Si$ bonded to $4$ $O$ atoms). Melting point: $1,713\text{ }^\text{o}C$. Found in sand/quartz.

2.7 Metallic Bonding

• Structure: A lattice of positive metal ions in a "sea" of delocalised electrons.
• Properties:     * Conductivity: Delocalised electrons carry charge and heat through the structure.     * Malleability/Ductility: Layers of ions can slide over each other without breaking the metallic bond because the electron sea moves with them.     * Strength: Higher number of delocalised electrons (e.g., $Al$ vs $Na$) results in stronger bonding.

2.8 States of Matter

• Solid: Regular lattice, vibrating fixed positions, high density.
• Liquid: Random close arrangement, can flow, fixed volume.
• Gas: Random spread, rapid movement, fills container, low density.
• Transitions:     * Endothermic (Energy absorbed): Melting, Boiling, Sublimation.     * Exothermic (Energy released): Freezing, Condensation.     * Sublimation: Solid to gas directly (e.g., Iodine, $\text{solid } CO_2$).
• Physical vs. Chemical Change: Physical changes (e.g., freezing) are reversible with no new substances; chemical changes create new substances and are usually irreversible.

2.9 Nanoparticles

• Size: $1$ to $100\text{ nm}$ ($1 \times 10^{-9}\text{ m}$ to $1 \times 10^{-7}\text{ m}$). Much smaller than fine particles ($PM_{2.5}$).
• Surface Area to Volume Ratio (SA:V): As particles shrink, SA:V increases dramatically. This makes nanoparticles highly reactive since more atoms are at the surface.     * Cube Example: Cube side $2\text{ nm}$. $SA = 6 \times (2^2) = 24\text{ nm}^2$. $V = 2^3 = 8\text{ nm}^3$. $SA:V = \frac{24}{8} = 3\text{ nm}^{-1}$.
• Applications: Medicine (targeted drug delivery); Sunscreens ($ZnO$ or $TiO_2$ protecting from UV without white residue); Catalysts (high SA:V ratio); Antibacterial silver coatings.
• Risks: Can penetrate skin or lungs; unknown long-term toxicity; environmental impact on ecosystems.