Atomic Structure and the Periodic table C1

Elements, Compounds & Mixtures definitions

• Element: A substance made of only one type of atom.

• Compound: Two or more elements chemically combined in fixed proportions.

• Mixture: Two or more elements or compounds not chemically combined together, whose chemical properties remain completely unchanged.

Physical vs Chemical changes in mixtures

• Separating a mixture involves physical processes rather than chemical reactions. • This means no new chemical substances are formed, and the individual components retain their original chemical properties.

Subatomic Particles mass and charge

• Proton: Relative mass of 1, relative charge of +1, located in the nucleus.

• Neutron: Relative mass of 1, relative charge of 0 (neutral), located in the nucleus.

• Electron: Relative mass of very small (1/2000 or 0.0005), relative charge of -1, located orbiting the nucleus in energy levels (shells).

Atomic size and scale constants

• Atom radius: Approximately 1 x 10⁻¹⁰ metres (or 0.1 nanometres).

• Nucleus radius: Approximately 1 x 10⁻¹⁴ metres, which is 10,000 times smaller than the overall atom.

• Mass concentration: Virtually all the mass of an atom is concentrated inside its tiny nucleus.

Atomic Number vs Mass Number rules

• Mass Number (Top): Total number of protons and neutrons in the nucleus.

• Atomic Number (Bottom): Number of protons in the nucleus, which also dictates the chemical identity of the element.

• Total Electrons: In a neutral atom, the number of negative electrons always equals the number of positive protons.

Ions formation and electron shells

• Definition: A charged particle formed when an atom gains or loses electrons.

• Metal atoms: Lose electrons from their outer shell to form positive ions (cations) with a full outer shell.

• Non-metal atoms: Gain electrons into their outer shell to form negative ions (anions) with a full outer shell.

Isotopes and Relative Atomic Mass (Ar)

• Isotopes: Atoms of the same element with the same atomic number (protons) but a different mass number (neutrons).

• Ar Definition: The average mass of an atom of an element compared to 1/12th the mass of a carbon-12 atom, taking into account the abundance of each isotope.

• Ar Formula: Sum of (isotope abundance × isotope mass number) ÷ 100.

Electronic Structure 2,8,8 rule

• First shell: Holds a maximum of 2 electrons.

• Second shell: Holds a maximum of 8 electrons.

• Third shell: Holds a maximum of 8 electrons.

• Periodic Table link: The group number of an element tells you exactly how many electrons are in its outer shell.

History of the Atom timeline

• Dalton (1803): Atoms are solid, indivisible spheres.

• Thomson (1897): Discovered electrons, creating the Plum Pudding Model (a ball of positive charge with negative electrons embedded).

• Rutherford (1909): Nuclear Model shows a central positive nucleus surrounded by empty space.

• Bohr (1913): Proved electrons orbit the nucleus at specific energy levels (shells).

• Chadwick (1932): Discovered neutrons inside the nucleus.

Rutherford's Gold Foil Exam evidence

• Most alpha particles passed straight through the foil: Proves the atom is mostly empty space.

• Some alpha particles were deflected at small angles: Proves the center of the atom has a concentrated positive charge that repelled them.

• A few alpha particles deflected straight back: Proves the nucleus is tiny, dense, and contains most of the atom's mass.

Why the Nuclear Model replaced the Plum Pudding Model

• The Plum Pudding model predicted all alpha particles would pass straight through the positive cloud with minimal deflection.

• The observation of large-angle deflections completely contradicted this, proving positive charge must be concentrated in a tiny nucleus.

Filtration method and application

• Function: Separates an insoluble solid from a liquid solvent.

• Method: Place filter paper in a funnel over a flask. Pour the mixture in; the liquid passes through the pores as filtrate, while the solid particles are trapped as residue.

Crystallisation method and application

• Function: Separates a soluble solid from a liquid solvent to collect dry crystals.

• Method: Heat the solution gently in an evaporating dish until it becomes saturated (crystals start to form). Remove from heat and leave to cool slowly so large crystals grow, then filter and dry with filter paper.

Simple Distillation method and application

• Function: Separates a liquid solvent from a dissolved solid solution, or two liquids with wildly different boiling points.

• Method: Heat the solution in a flask. The liquid with the lowest boiling point evaporates into vapour, passes into a water-cooled Liebig condenser, condenses back into a liquid, and is collected.

Fractional Distillation method and application

• Function: Separates a mixture of multiple miscible liquids with close boiling points (e.g. ethanol and water, or crude oil).

• Method: Heat the mixture below a fractionating column packed with glass beads. The liquid with the lowest boiling point reaches the top first and enters the condenser, while higher boiling point vapours condense early on the beads and drip back down.

Paper Chromatography method and phases

• Function: Separates mixtures of soluble substances (like inks or food colourings) based on their solubilities.

• Stationary phase: The uniform paper which does not move. • Mobile phase: The liquid solvent that moves up the paper via capillary action.

Rf Value definition and calculations

• Rf (Retention Factor): A ratio showing how far a substance moved compared to the solvent front.

• Formula: Rf = Distance moved by substance ÷ Distance moved by solvent front.

• Properties: Has no units, is always a decimal less than 1, and varies depending on the specific solvent used.

Pure vs Impure substances in chromatography

• Pure substance: Produces exactly one single spot on the chromatogram regardless of the solvent used.

• Impure mixture: Separates into two or more distinct spots at different heights on the paper.

Early Periodic Table development

• Early 1800s: Elements were arranged strictly in order of increasing atomic weight because subatomic particles had not yet been discovered.

• Flaw: This primitive system placed completely dissimilar elements into the same columns because it ignored their chemical behaviors.

Mendeleev's Periodic Table breakthroughs

• Improvements: Kept elements arranged by atomic weight but swapped the order if their chemical properties did not fit the pattern.

• Gaps left: Left empty gaps in the table for undiscovered elements, allowing him to accurately predict the physical and chemical properties of these missing elements before they were found.

Modern Periodic Table arrangement

• Setup: Elements are arranged in order of increasing atomic (proton) number.

• Layout: Vertical columns are called Groups (elements have the same number of outer shell electrons and similar properties). Horizontal rows are called Periods.

Metals vs Non-Metals: Physical properties

• Metals: Good conductors of heat and electricity, malleable, shiny, high melting and boiling points, and high density.

• Non-Metals: Poor conductors (insulators), brittle, dull appearance, low melting and boiling points, and lower density.

Metals vs Non-Metals: Chemical properties

• Metals: Found on the left and center of the periodic table, reacting to lose electrons and form positive ions (+).

• Non-Metals: Found on the top right side, reacting to gain or share electrons to form negative ions (-) or covalent bonds.

Group 1 Alkali Metals: General properties

• Features: Very soft to cut, low density (the first three float on water), and have low melting and boiling points that decrease down the group.

• Ion formed: Lose their single outer electron to form a +1 ion.

Group 1 Alkali Metals: Reactivity trend

• Trend: Reactivity increases going down the group.

• Explanation: As you go down, atoms gain more electron shells, making the outer electron further away from the positive nucleus. The electron is more shielded from the nuclear charge, reducing electrostatic attraction and making it much easier to lose.

Group 1 Alkali Metals: Chemical reactions

• With Water: React vigorously to produce a metal hydroxide solution (alkaline) and hydrogen gas (e.g. 2Na + 2H₂O → 2NaOH + H₂).

• With Chlorine: React vigorously to form white solid metal chloride salts (e.g. 2Li + Cl₂ → 2LiCl).

• With Oxygen: React vigorously to form solid metal oxides, causing the metal to tarnish instantly (e.g. 4K + O₂ → 2K₂O).

Group 7 Halogens: General properties

• Features: Non-metals that exist as diatomic molecules (paired molecules like F₂, Cl₂, Br₂, I₂).

• Trends: Melting and boiling points increase going down the group. • Ion formed: Gain one electron to form a -1 halide ion.

Group 7 Halogens: Reactivity trend

• Trend: Reactivity decreases going down the group. • Explanation: As you go down, the atom gains more electron shells, so the outer shell is further away from the positive nucleus and more shielded. This reduces the electrostatic attraction, making it harder to attract and gain an incoming electron.

Group 7 Halogens: Displacement Reactions

• Rule: A more reactive halogen will physically displace a less reactive halogen from an aqueous solution of its salt.

• Example: Chlorine + Sodium bromide → Sodium chloride + Bromine (Cl₂ + 2NaBr → 2NaCl + Br₂). Chlorine is more reactive than bromine, so it takes its place.

Group 0 Noble Gases: Properties and trends

• Features: Colourless, non-flammable gases that exist as single monoatomic atoms. • Reactivity: Completely unreactive (inert) because they possess a stable, full outer electron shell, meaning they have no tendency to lose or gain electrons.

• Trends: Boiling point and density increase going down the group due to an increase in relative atomic mass.