Atomic Structure & the Periodic Table

• An atom is the smallest part of an element that can exist; about 100 elements are represented by one- or two-letter symbols.
• A compound contains two or more elements chemically bonded in fixed proportions and can be split into elements only by chemical reactions.
• A balanced symbol equation has equal numbers of each atom on both sides; change coefficients in front, never a substance's formula.
• Higher Tier: write balanced half equations and ionic equations when the reaction requires them.

The punchline: Keep identity in the formula; balance the count with coefficients.

• Filtration traps an insoluble solid as residue while the liquid passes through as filtrate.
• Crystallisation recovers a dissolved solid by concentrating a solution, then allowing crystals to form.
• Simple distillation recovers a solvent; fractional distillation separates liquids with different boiling points.
• Chromatography separates dissolved substances by their different attractions to the stationary paper and moving solvent.

The punchline: Name the differing physical property, then choose the method that exploits it.

• Dalton's indivisible solid spheres gave way to the plum pudding model after electrons were discovered.
• Most alpha particles passed through foil, showing that atoms are mostly empty space; a few deflected strongly near concentrated positive charge and mass.
• The nuclear model placed a tiny positive nucleus at the centre, with electrons around it.
• Bohr's fixed energy levels fitted evidence; later work identified protons, and Chadwick's experiments supplied evidence for neutrons.

The punchline: In an evidence question, link each observation to the feature it revealed.

• Protons have relative charge +1 and mass 1; neutrons have charge 0 and mass 1; electrons have charge −1 and negligible relative mass.
• Atomic number = protons. Mass number = protons + neutrons, so neutrons = mass number − atomic number.
• A neutral atom has equal numbers of protons and electrons; an ion has gained or lost electrons, not protons.
• In nuclear notation, the mass number sits at the upper left of the symbol and the atomic number at the lower left.

The punchline: Use atomic number for protons, subtract for neutrons, then account for charge when counting electrons.

• An atom's radius is about 0.1 nm, or 1 × 10⁻¹⁰ m. Its nucleus has less than 1/10,000 of that radius.
• Almost all atomic mass is in the nucleus, while most atomic volume is empty space occupied by electron shells.
• Isotopes are atoms of one element with the same proton number but different neutron numbers, so their mass numbers differ.
• Isotopes have the same electron arrangement and therefore the same chemical reactions.

The punchline: Same protons means same element; different neutrons means a different isotope.

• Relative atomic mass compares the average mass of an element's atoms with 1/12 of the mass of a carbon-12 atom and has no unit.
• To calculate it, multiply each isotope mass by its percentage abundance, add the products, then divide by 100.
• The answer often lies between whole-number isotope masses; chlorine is about 35.5 because chlorine-35 is more abundant than chlorine-37.

The punchline: Weight each isotope by how common it is; do not take an unweighted mean.

• For the first 20 elements, fill 2 electrons in the first shell, then up to 8 in the second and up to 8 in the third.
• Write sodium as 2,8,1 or draw electrons on shells; both representations must total its atomic number of 11.
• For Groups 1–7, the group number matches the outer-shell electron count; the period matches the number of occupied shells.
• Outer electrons control how readily atoms lose, gain or share electrons, linking electronic structure to reactions and ion formation.

The punchline: Check the total, then use outer electrons for group and occupied shells for period.

• Modern elements are arranged by increasing atomic number; vertical columns are groups and horizontal rows are periods.
• Members of a group have the same number of outer-shell electrons and therefore similar chemical properties.
• Metals occupy the left and lower regions; non-metals lie mainly to the right.
• Use an element's position to predict its electron arrangement, likely ion and broad chemical behaviour.

The punchline: Position is compressed evidence: read across for shells and down a group for shared outer structure.

• Early tables used atomic weight because protons and electrons were not yet known; Newlands forced every eighth element together and produced mismatches.
• Mendeleev mainly followed atomic weight but left gaps and changed some orders to keep elements with similar properties together.
• Later discoveries matched properties he had predicted for missing elements, strengthening acceptance of his table.
• Atomic number later supplied the correct order and explained mass-order exceptions caused by isotope averages.

The punchline: Credit Mendeleev for gaps, justified swaps and successful predictions—not for using atomic number.

• Metals are typically shiny, strong conductors, malleable and high-melting, though these are broad patterns rather than definitions.
• Solid non-metals are typically dull, brittle and poor conductors, and many non-metals have lower melting and boiling points.
• An element's position and outer electrons help predict whether it will lose electrons, gain them or share them in bonding.

The punchline: Use electron behaviour for the chemical distinction; use physical properties as typical evidence.

• Noble gases are very unreactive, monatomic gases; helium has 2 outer electrons and the others have 8.
• Their boiling points increase down the group as relative atomic mass increases and attractions between atoms become stronger.
• Use the trend to predict that a noble gas below another will generally have a higher boiling point.

The punchline: Full outer shell explains low reactivity; increasing mass tracks the boiling-point trend.

• Group 1 metals are soft, low-density metals that lose one electron to form +1 ions; reactivity increases down the group.
• Greater distance and shielding weaken attraction between the nucleus and outer electron, so that electron is lost more easily.
• Lithium, sodium and potassium react with water to make hydrogen and an alkaline metal-hydroxide solution; potassium reacts most vigorously.
• They react with oxygen to form oxides and with chlorine to form white metal-chloride salts.

The punchline: Explain the trend with distance, shielding and easier loss—not simply with a larger atom.

• Halogens are diatomic molecules with 7 outer electrons; they gain one electron to form −1 ions or share electrons in covalent bonds.
• Reactivity decreases down the group, while relative molecular mass and melting and boiling points increase.
• At room temperature chlorine is green gas, bromine orange-brown liquid and iodine grey-black solid; these observations support trend predictions.
• A more reactive halogen displaces a less reactive halogen from its salt; halogens form ionic salts with metals and covalent molecules with non-metals.

The punchline: For displacement, the free halogen must sit higher in Group 7 than the halogen in the salt.

• Separate Chemistry: transition metals occupy the central block between Groups 2 and 3.
• Compared with Group 1, they are less reactive, harder, stronger, denser and usually higher-melting; mercury is the melting-point exception.
• Many form ions with different charges, such as Fe²⁺ and Fe³⁺, and many compounds are coloured.
• Transition metals and their compounds often act as catalysts; iron is used in the Haber process.

The punchline: Label this separate-only content and compare transition metals directly with Group 1.