1. Alchemy

John Dalton (1803)

Proposed the Solid Sphere Model, stating that atoms are tiny, solid, indivisible spheres that cannot be broken apart and combine in fixed ratios to form compounds

J. J. Thomson (1904)

Proposed the Plum Pudding Model, where the atom is a sphere of positive charge with negatively charged electrons scattered throughout, following his discovery of the electron in 1897

Ernest Rutherford (1911)

Proposed the Nuclear Model after the gold foil experiment, showing that atoms are mostly empty space with a small, dense, positively charged nucleus that deflects some alpha particles

Niels Bohr (1913)

Proposed the Planetary Model, where electrons orbit the nucleus in fixed energy levels (shells) with specific energies and cannot exist between these levels (quantised energy)

Erwin Schrödinger (1926)

Proposed the Quantum Model, stating that electrons do not move in fixed paths but exist as waves in regions of probability called orbitals, meaning their exact position cannot be known

Proton

Positively charged subatomic particle located in the nucleus with a relative mass of 1

Neutron

Neutral subatomic particle located in the nucleus with a relative mass of 1

Electron

Negatively charged subatomic particle located in shells around the nucleus with a very small mass (1/1837)

Atomic Number

The number of protons in the nucleus of an atom, which determines the element

Mass Number (Nucleon Number)

The total number of protons and neutrons in the nucleus

Neutral Atom

An atom with equal numbers of protons and electrons, resulting in no overall charge

Ion

An atom or group of atoms that has gained or lost electrons and therefore carries a charge

Cation

A positively charged ion formed when an atom loses one or more electrons

Anion

A negatively charged ion formed when an atom gains one or more electrons

Electron Shell

An energy level around the nucleus where electrons are found

Electron Configuration

The arrangement of electrons in shells or energy levels around the nucleus

Isotope

Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers

Chemical Properties of Isotopes

Similar because isotopes have the same number of electrons and therefore the same electron configuration

Physical Properties of Isotopes

Different because they have different masses, which affects properties such as melting and boiling points

Döbereiner (1829)

Groups of three elements (triads) with similar chemical properties

John Newlands (1800s)

Arrangement of elements by increasing atomic mass where every eighth element showed similar properties, though some elements were forced into incorrect groups. Law of Octaves, where each element was similar to the element eight places further on.

Scientists rejection of Newland's Octaves

Some elements had to be put into groups that didn't match their chemical properties.

Dmitri Mendeleev (1869)

Arrangement of elements by increasing atomic mass with gaps left for undiscovered elements so that elements with similar properties aligned in vertical columns

Mendeleev different from Newland's

Put elements into groups based on similar chemical properties

Henry Moseley (1913)

Arrangement of elements by increasing atomic number instead of atomic mass, correcting inconsistencies such as the placement of iodine and tellurium

Noble Gases

Elements in Group 18 that are non-metals and very unreactive because they have full valence electron shells, making them stable

Inert Gases

Another name for noble gases due to their lack of reactivity

Monoatomic Gas

A gas that exists as single atoms rather than molecules because it does not need to bond

Boiling Point Trend (Noble Gases)

Boiling point increases down the group because larger atoms have more electrons, leading to stronger intermolecular forces

Density Trend (Noble Gases)

Density increases down the group due to increasing atomic mass

Helium Use

Used in balloons

Neon Use

Used in lights

Argon Use

Used in light bulbs

Krypton Use

Used in double glazing

Xenon Use

Used in satellites

Radon Use

Used in cancer treatment

Alkali Metals

Group 1 elements that have one electron in their outer shell and are highly reactive

Reactivity Trend (Group 1)

Reactivity increases down the group because outer electrons are further from the nucleus and more shielded, so they are lost more easily

Shielding Effect

The reduction in attraction between the nucleus and outer electrons due to inner electron shells blocking the nuclear charge

Halogens

Group 7 non-metal elements that need to gain one electron to achieve a full outer shell

Diatomic Molecule

A molecule consisting of two atoms bonded together, as seen in halogens

Reactivity Trend (Halogens)

Reactivity decreases down the group because outer electrons are further from the nucleus and more shielded, making it harder to gain an electron

Displacement Reaction

A reaction in which a more reactive halogen displaces a less reactive halide ion from a compound

Ionic Bond

A strong electrostatic attraction between oppositely charged ions formed when electrons are transferred from a metal to a non-metal

Ionic Lattice

A giant repeating structure of alternating positive and negative ions held together by strong electrostatic forces

Dot and Cross Diagram

A diagram used to show the transfer of valence electrons between atoms in ionic bonding

Group 1 Ions

Always Soluble (aq)

Ammonium

Always Soluble (aq)

Nitrates

Always Soluble (aq)

Acetates

Always Soluble (aq)

Perchlorates

Always Soluble (aq)

Chlorides, Bromides, Iodides

Soluble (aq)

Chlorides, Bromides, Iodides EXCEPTIONS (s)

Ag+, Pb2+, or Hg2+ (s)

Sulfates

Soluble (aq)

Sulfates EXCEPTION (s)

Ba2+, Sr2+, Pb2+, Ca2+, or Hg2+ (s)

Carbonates & Phosphates

Insoluble (s)

Carbonates & Phosphates EXCEPTION (aq)

Group 1 ions or NH4+ (aq)

Sulfides

Insoluble (s)

Sulfides EXCEPTION (aq)

Group 1, Group 2: Ca2+, Sr2+, Ba2+, or NH42+ (aq)

Hydroxides

Insoluble (s)

Hydroxides EXCEPTION (aq)

Group 1 and Ca2+, Sr2+, Ba2+ (aq)

Spectator Ion

An ion that remains unchanged on both sides of a chemical equation and does not participate in the reaction

Net Ionic Equation

An equation showing only the ions and substances that actually take part in the chemical reaction

HCl vs H2SO4 reactivity

sulfuric acid more reactive

Metal + acid

Salt + hydrogen

Iron + oxygen

iron turns from light grey to dark grey

Magnesium + oxygen

bright white flash, magnesium metal turns to white powder

Copper + oxygen

copper colour turns from bronze to chrome

Metal carbonate + acid

Salt + water + carbon dioxide