Year 9 Chemistry: Atoms, Elements, and Structure

• Fully detailed revision notes cover all the topics in the revision list with explanations and examples for Year 9 Chemistry.

Atoms, Elements, Compounds, Molecules, Isotopes, and Ions

• An atom is the smallest unit of matter, comprised of protons, neutrons, and electrons.
• An element is a pure substance consisting of only one type of atom, such as Oxygen (O) or Hydrogen (H).
• A compound is a substance formed when two or more elements chemically bond, such as Water (H2O).
• A molecule consists of two or more atoms covalently bonded, like O2 or CO2.
• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, for example, Carbon-12 and Carbon-14.
• An ion is a charged particle formed when an atom gains or loses electrons, such as Na+ (sodium ion) or Cl- (chloride ion).

Atomic Structure

• Proton: Subatomic particle with a charge of +1, a relative mass of 1, located in the nucleus.
• Neutron: Subatomic particle with a charge of 0, a relative mass of 1, located in the nucleus.
• Electron: Subatomic particle with a charge of -1, a relative mass of ~0, located in shells.
• Oxygen Example: Oxygen (O, Atomic Number = 8) contains 8 protons, 8 neutrons (if the mass number is 16), and 8 electrons.

Ionic Charges

• Elements in Groups 1, 2, and 3 lose electrons to form positive ions (cations), e.g., Na → Na⁺ + e⁻.
• Elements in Groups 5, 6, and 7 gain electrons to form negative ions (anions), e.g., Cl + e⁻ → Cl⁻.

Relative Atomic Mass (Isotopes)

• Formula: Relative Atomic Mass=∑(isotope mass×abundance)100
• Chlorine has two isotopes: 75% of Cl-35 and 25% of CI-37.
• Relative Atomic Mass=(35×75)+(37×25)100=35.5

Diatomic vs Monoatomic Molecules

• Diatomic molecules consist of two atoms chemically bonded, e.g., O2, N2, H2.
• Monoatomic molecules consist of single atoms not bonded to others, e.g., noble gases like He, Ne.

Electron Configuration

• Electrons fill shells using the 2,8,8 rule.
• Example: Sodium (Na, Atomic Number = 11) has an electron configuration of 2,8,1.

Dot and Cross Diagrams

• Ionic Bonding: Involves electron transfer between metals & non-metals, resulting in ions like Na⁺ and Cl in NaCl.
• Covalent Bonding: Involves electrons shared between non-metals, such as in H2O where Oxygen shares electrons with two Hydrogen atoms.

Bonding Types

• Ionic bonding occurs between a metal and a non-metal, resulting in a high melting point and electricity conduction when molten.
• Covalent bonding occurs between non-metals and non-metals, resulting in a low melting point and no electricity conduction.
• Metallic bonding occurs between metals only and conducts electricity, is malleable, and is ductile.

Carbon Allotropes

• Diamond: Each carbon bonds to 4 others, is very hard, and does not conduct electricity.
• Graphite: Each carbon bonds to 3 others and conducts electricity due to delocalized electrons.
• Fullerenes: Carbon is arranged in spheres/tubes (nanotubes used in medicine and electronics).

Ionic Lattices

• Ionic compounds form giant lattices of alternating positive and negative ions.
• Example: NaCl lattice is held together by strong electrostatic forces.

Metallic Bonding & Conductivity

• Delocalized electrons allow metals to conduct electricity and heat.
• Alloys (e.g., steel, brass) are stronger than pure metals because different atom sizes disrupt the layers.

Periodic Table

• Groups: Columns that show valence electrons.
• Periods: Rows that show electron shells.
• Metals: Located on the left side.
• Non-metals: Located on the right side.
• Alkali Metals (Group 1): Very reactive.
• Halogens (Group 7): Form salts with metals.
• Noble Gases (Group 8/0): Inert and monoatomic.

Group 1 – Alkali Metals

• Reactivity Increases Down the Group due to more shielding, making it easier to lose electrons.
• Reacts with Water: Example: Na + H2O → NaOH + H2

Group 7 – Halogens

• Reactivity Decreases Down the Group due to more shielding, making it harder to gain electrons.
• Displacement Reactions: A more reactive halogen displaces a less reactive halide, e.g., Cl2 + 2KBr → 2KCl + Br2 (chlorine displaces bromine).

Redox Reactions

• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.
• Example: Mg + O2 → MgO; Mg is oxidized (loses electrons), and O is reduced (gains electrons).

Noble Gases

• Monoatomic: Exist as single atoms.
• Unreactive: Exhibit a full outer shell.
• Trends Down the Group: Boiling Point Increases due to more intermolecular forces.
• Example Uses: Helium in balloons (lighter than air, non-flammable) and Argon in lightbulbs (prevents filament burning).