GCSE Chemistry Revision Booklet
1. Atomic Structure
Changing Models of the Atom
● John Dalton: Proposed that atoms were solid spheres.
● J.J. Thomson: Discovered electrons; proposed the "plum pudding" model
where electrons were embedded in a positively charged "pudding."
● Ernest Rutherford: Conducted the gold foil experiment, leading to the
discovery of a dense, positively charged nucleus.
● Niels Bohr: Introduced electron shells where electrons orbit the nucleus at
fixed distances.
● James Chadwick: Discovered neutrons in the nucleus.
Atomic Structure
● Protons: Positively charged particles found in the nucleus.
● Neutrons: Neutral particles found in the nucleus.
● Electrons: Negatively charged particles orbiting the nucleus in shells.
● Atomic Number: Number of protons in the nucleus.
● Mass Number: Total number of protons and neutrons in the nucleus.
● Electron Configuration: Arrangement of electrons in an atom’s shells.
Isotopes
● Definition: Atoms of the same element with different numbers of neutrons.
● Examples: Carbon-12 and Carbon-14.
● Importance: Used in dating, medical imaging, and as tracers in biological
systems.
Calculating Relative Atomic Mass
● Formula: (Isotope mass × Isotope abundance) / 100
● Example: For chlorine with isotopes Cl-35 and Cl-37:
2. The Periodic Table
Structure of the Periodic Table
● Groups: Vertical columns; elements with similar chemical properties.
● Periods: Horizontal rows; elements with increasing proton number.
● Blocks: s-block, p-block, d-block, and f-block.
Trends in the Periodic Table
, ● Atomic Radius: Decreases across a period; increases down a group.
● Ionization Energy: Increases across a period; decreases down a group.
● Electronegativity: Increases across a period; decreases down a group.
● Reactivity: Varies across periods and groups, e.g., Group 1 metals become
more reactive down the group.
Group 1 (Alkali Metals)
● Properties: Soft, highly reactive, low density.
● Reactivity: Increases down the group due to easier loss of the outer
electron.
● Examples: Lithium (Li), Sodium (Na), Potassium (K).
Group 7 (Halogens)
● Properties: Non-metals, diatomic molecules, high electronegativity.
● Reactivity: Decreases down the group due to increasing atomic size.
● Examples: Fluorine (F), Chlorine (Cl), Bromine (Br).
Group 0 (Noble Gases)
● Properties: Inert, colorless gasses, full outer electron shells.
● Uses: Helium in balloons, Neon in signs, Argon in light bulbs.
3. Bonding
Ionic Bonding
● Formation: Transfer of electrons from a metal to a non-metal.
● Properties: High melting and boiling points, conduct electricity when
molten or in solution.
● Examples: Sodium chloride (NaCl), Magnesium oxide (MgO).
Covalent Bonding
● Formation: Sharing of electrons between nonmetals.
● Properties: Low melting and boiling points (simple molecules), do not
conduct electricity.
● Examples: Water (H₂O), Methane (CH₄).
Metallic Bonding
● Formation: Positive metal ions in a sea of delocalized electrons.
● Properties: Conduct electricity and heat, malleable, ductile.
● Examples: Copper (Cu), Iron (Fe).
4. States of Matter & Mixtures
States of Matter
● Solids: Fixed shape and volume, particles in a regular arrangement.
1. Atomic Structure
Changing Models of the Atom
● John Dalton: Proposed that atoms were solid spheres.
● J.J. Thomson: Discovered electrons; proposed the "plum pudding" model
where electrons were embedded in a positively charged "pudding."
● Ernest Rutherford: Conducted the gold foil experiment, leading to the
discovery of a dense, positively charged nucleus.
● Niels Bohr: Introduced electron shells where electrons orbit the nucleus at
fixed distances.
● James Chadwick: Discovered neutrons in the nucleus.
Atomic Structure
● Protons: Positively charged particles found in the nucleus.
● Neutrons: Neutral particles found in the nucleus.
● Electrons: Negatively charged particles orbiting the nucleus in shells.
● Atomic Number: Number of protons in the nucleus.
● Mass Number: Total number of protons and neutrons in the nucleus.
● Electron Configuration: Arrangement of electrons in an atom’s shells.
Isotopes
● Definition: Atoms of the same element with different numbers of neutrons.
● Examples: Carbon-12 and Carbon-14.
● Importance: Used in dating, medical imaging, and as tracers in biological
systems.
Calculating Relative Atomic Mass
● Formula: (Isotope mass × Isotope abundance) / 100
● Example: For chlorine with isotopes Cl-35 and Cl-37:
2. The Periodic Table
Structure of the Periodic Table
● Groups: Vertical columns; elements with similar chemical properties.
● Periods: Horizontal rows; elements with increasing proton number.
● Blocks: s-block, p-block, d-block, and f-block.
Trends in the Periodic Table
, ● Atomic Radius: Decreases across a period; increases down a group.
● Ionization Energy: Increases across a period; decreases down a group.
● Electronegativity: Increases across a period; decreases down a group.
● Reactivity: Varies across periods and groups, e.g., Group 1 metals become
more reactive down the group.
Group 1 (Alkali Metals)
● Properties: Soft, highly reactive, low density.
● Reactivity: Increases down the group due to easier loss of the outer
electron.
● Examples: Lithium (Li), Sodium (Na), Potassium (K).
Group 7 (Halogens)
● Properties: Non-metals, diatomic molecules, high electronegativity.
● Reactivity: Decreases down the group due to increasing atomic size.
● Examples: Fluorine (F), Chlorine (Cl), Bromine (Br).
Group 0 (Noble Gases)
● Properties: Inert, colorless gasses, full outer electron shells.
● Uses: Helium in balloons, Neon in signs, Argon in light bulbs.
3. Bonding
Ionic Bonding
● Formation: Transfer of electrons from a metal to a non-metal.
● Properties: High melting and boiling points, conduct electricity when
molten or in solution.
● Examples: Sodium chloride (NaCl), Magnesium oxide (MgO).
Covalent Bonding
● Formation: Sharing of electrons between nonmetals.
● Properties: Low melting and boiling points (simple molecules), do not
conduct electricity.
● Examples: Water (H₂O), Methane (CH₄).
Metallic Bonding
● Formation: Positive metal ions in a sea of delocalized electrons.
● Properties: Conduct electricity and heat, malleable, ductile.
● Examples: Copper (Cu), Iron (Fe).
4. States of Matter & Mixtures
States of Matter
● Solids: Fixed shape and volume, particles in a regular arrangement.
