Chemistry Chapter C2: Atoms, Elements and Compounds

Hello future chemist! This chapter is absolutely fundamental. Think of it as the A, B, C of all chemistry. We are going to break down how everything around you—from the air you breathe to the metals in your phone—is built. Understanding atoms and how they stick together (bonding) is the key to unlocking the rest of the syllabus. Don't worry, we'll take it one step at a time!

C2.1 Elements, Compounds and Mixtures

Imagine building a giant LEGO castle. The basic LEGO bricks are like atoms. How you put them together determines whether you have an element, a compound, or a mixture.

1. Elements

An element is a pure substance consisting of only one type of atom.
It cannot be broken down into simpler substances by chemical means.

  • Examples: Oxygen (O₂), Gold (Au), Carbon (C).
2. Compounds

A compound is a pure substance formed when two or more different elements are chemically combined in a fixed ratio.

  • Compounds have properties totally different from the elements they are made of.
  • They can only be separated by chemical reactions.
  • Example: Water (H₂O). It is made of Hydrogen (explosive gas) and Oxygen (supports combustion), but water itself is a liquid used to put out fires!
3. Mixtures

A mixture consists of two or more substances (elements or compounds) that are not chemically combined.

  • They retain the properties of the original substances.
  • They can be separated using physical methods (like filtration, distillation, etc.).
  • Example: Saltwater (a mixture of salt (a compound) and water (a compound)).

Quick Review: Element (one type of atom), Compound (chemically joined), Mixture (physically mixed).

C2.2 Atomic Structure and the Periodic Table

The atom is the smallest unit of an element that retains its properties. Let’s look inside!

1. Structure of the Atom (Core)

Atoms consist of a central, dense nucleus surrounded by electrons moving in shells (or energy levels).

  • The Nucleus: Contains two types of particles: protons and neutrons.
  • The Shells: Contain electrons.

Key Sub-atomic Particles:

Particle Relative Mass Relative Charge
Proton 1 +1 (Positive)
Neutron 1 0 (Neutral)
Electron 1/1840 (almost zero) -1 (Negative)

Memory Aid: PROtons are POSITIVE. NEUtrons are NEUTRAL.

2. Key Atomic Numbers (Core)

The identity of an atom is defined by the number of protons it has.

  • Proton number / Atomic number (Z): The number of protons in the nucleus.
    For a neutral atom, Z also equals the number of electrons.
  • Mass number / Nucleon number (A): The total number of protons and neutrons in the nucleus.

How to calculate the number of neutrons:
Neutrons = Mass Number (A) − Proton Number (Z)

Did you know? Protons and neutrons are also called nucleons because they live in the nucleus!

3. Electronic Configuration (Core)

Electrons arrange themselves in specific shells around the nucleus. We determine the electronic configuration for elements with proton number 1 to 20.

  • Shell capacities (for the first 20 elements):
    1st shell: maximum 2 electrons
    2nd shell: maximum 8 electrons
    3rd shell: maximum 8 electrons (until you reach element 20)

Example: Sodium (Na) has 11 protons (Z=11). Configuration is 2, 8, 1.

4. Linking Electronic Configuration to the Periodic Table (Core)

The electron arrangement tells us exactly where an element sits in the Periodic Table:

  • Group Number (I to VII): Equals the number of outer-shell electrons. (Group VIII Noble Gases are special—they have a full outer shell, usually 8, except Helium which has 2.)
  • Period Number: Equals the number of occupied electron shells.

Example: Chlorine (2, 8, 7) is in Group VII (7 outer electrons) and Period 3 (3 occupied shells).

Key Takeaway: The number of protons determines the element (Z), and the number of outer electrons determines its chemical behaviour (Group).

C2.3 Isotopes

What are Isotopes? (Core)

Isotopes are different atoms of the same element that have the same number of protons (same Z) but different numbers of neutrons (different A).

Example: Carbon-12 and Carbon-14. Both have 6 protons. C-12 has 6 neutrons, C-14 has 8 neutrons.

Nuclide Notation (Core)

We use a standard symbol to represent atoms and ions:

\( \frac{A}{Z}X \)

  • A = Mass Number (Protons + Neutrons)
  • Z = Proton Number (Protons)
  • X = Element Symbol

Example: Carbon-12 is written as \( \frac{12}{6}C \). Chlorine ion is written as \( \frac{35}{17}Cl^- \).

Chemical Properties of Isotopes (Supplement)

Isotopes of the same element have identical chemical properties. This is because chemical reactions only involve the outer electrons.

  • Since isotopes have the same number of protons, they must have the same number of electrons, and therefore the same electronic configuration.
  • The difference in the number of neutrons does not affect how the atom reacts chemically.

Key Takeaway: Isotopes are chemically identical but differ in mass (due to different neutron counts).

C2.4 Ions and Ionic Bonds

Atoms are usually happiest when their outer shell is full (like the Noble Gases). They achieve this stability by gaining or losing electrons, forming ions.

1. Forming Ions (Core)

An ion is an atom (or group of atoms) that has lost or gained electrons, giving it an overall electrical charge.

  • Cations: Positive ions formed when atoms (usually metals from Groups I, II, III) lose electrons. (They lose negative charge, so they become positive.)
  • Anions: Negative ions formed when atoms (usually non-metals from Groups V, VI, VII) gain electrons. (They gain negative charge, so they become negative.)
2. The Ionic Bond (Core)

An ionic bond is the strong electrostatic attraction between oppositely charged ions.

Formation (Core/Supplement)

  • This bonding occurs between metals (which form positive ions/cations) and non-metals (which form negative ions/anions).
  • The metal atom transfers its outer electrons completely to the non-metal atom.

Example: Sodium Chloride (NaCl) formation (Group I and Group VII).
Na (2, 8, 1) loses 1 electron $\rightarrow$ Na$^+$ (2, 8)
Cl (2, 8, 7) gains 1 electron $\rightarrow$ Cl$^-$ (2, 8, 8)

You must be able to draw dot-and-cross diagrams to show this electron transfer and the resulting electronic configurations (Core/Supplement).

3. Structure and Properties of Ionic Compounds (Core/Supplement)

Ionic compounds form a giant lattice structure (exemplified by sodium chloride). This is a regular arrangement of alternating positive and negative ions held together by strong electrostatic forces.

Properties:

  1. High Melting Points and Boiling Points: (Explanation in Supplement) Large amounts of energy are required to break the very strong electrostatic forces of attraction throughout the giant lattice structure.
  2. Electrical Conductivity:
    • Poor when solid: The ions are held firmly in fixed positions and cannot move to carry charge.
    • Good when aqueous or molten: When melted or dissolved, the ions are free to move and carry the electric current.
  3. Solubility: Generally soluble in water.

Key Takeaway: Ionic bonds involve transfer of electrons between metals and non-metals, creating strong lattices with high melting points and only conduct electricity when liquid or dissolved.

C2.5 Simple Molecules and Covalent Bonds

Covalent bonds form when atoms share electrons to achieve a stable noble gas configuration.

1. The Covalent Bond (Core)

A covalent bond is formed when a pair of electrons is shared between two atoms, usually non-metals.

  • Sharing allows both atoms to count the shared electrons towards their outer shell, achieving stability.

Dot-and-Cross Diagrams (Core/Supplement):

You must be able to draw dot-and-cross diagrams showing the electronic configurations in simple covalent molecules. The syllabus specifies:

  • Core: H₂, Cl₂, H₂O, CH₄ (methane), NH₃ (ammonia), HCl.
  • Supplement adds: CH₃OH (methanol), C₂H₄ (ethene), O₂, CO₂, N₂.

Hint for Diagrams: Only show the outer electron shells in your diagrams!

2. Structure and Properties of Simple Molecular Compounds (Core/Supplement)

Compounds formed by covalent bonding exist as simple molecules (like H₂O or CO₂), where strong bonds exist within the molecule, but weak forces exist between the molecules.

Properties:

  1. Low Melting Points and Boiling Points: (Explanation in Supplement)
    • To melt or boil the substance, you only need to overcome the weak intermolecular forces (forces between molecules).
    • Since these forces are weak, little energy is needed, resulting in low MP/BP.
  2. Poor Electrical Conductivity: They have no free charged particles (no ions or delocalised electrons) to carry a current.

Key Takeaway: Covalent bonds form simple molecules held together by weak forces, leading to low melting points and poor conductivity.

C2.6 & C2.7 Giant Structures

Not all covalent or metallic substances exist as simple molecules; many form huge, continuous lattice structures.

C2.6 Giant Covalent Structures (Core & Supplement)

These substances contain many atoms held together by strong covalent bonds in a large, repeating structure.

1. Diamond (Core/Supplement)

  • Structure: Each carbon atom is bonded covalently to four other carbon atoms in a rigid tetrahedral structure.
  • Properties & Uses:
    • Extremely hard (used for cutting tools).
    • High melting point (due to many strong covalent bonds needing to be broken).
    • Does not conduct electricity (no free electrons).

2. Graphite (Core/Supplement)

  • Structure: Each carbon atom is bonded covalently to three other carbon atoms, forming layers of hexagons. The layers are held together by weak forces.
  • Properties & Uses:
    • Soft and slippery (used as a lubricant) because the weak forces between layers allow the layers to slide over each other easily.
    • Conducts electricity (used as an electrode) because there is one spare electron per carbon atom that becomes delocalised (free to move) within the layers.
C2.7 Metallic Bonding (Supplement Only)

Metals are held together by metallic bonding.

1. Metallic Bonding Description (Supplement)

Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and a 'sea' of delocalised electrons.

  • The metal atoms lose their outer electrons, which are then free to move throughout the structure.

2. Properties of Metals (Supplement)

These properties are explained by the structure:

  1. Good Electrical Conductivity: The delocalised electrons are mobile and can move through the lattice, carrying electrical charge.
  2. Malleability (can be hammered into sheets) & Ductility (can be drawn into wires): The layers of positive ions can slide over one another without shattering the structure, because the mobile electron sea holds the structure together regardless of the position of the ions.

Key Takeaway: Giant covalent structures like diamond are extremely hard insulators, while graphite is slippery and conductive. Metallic bonds use a 'sea' of electrons for high conductivity and malleability.