VCE Vault

Year 9 Science · Chemical Sciences · 18m read

Atomic Structure and the Periodic Table

Ever wonder why some elements are metals and others are gases? It all comes down to how atoms are structured—the tiny particles that make up everything around you. You'll explore how electrons arrange themselves and how this arrangement explains why the periodic table is organized the way it is. Mastering this foundation helps you predict how chemicals will react and is essential for VCE chemistry.

Learning objective: Describe atomic structure and the arrangement of the periodic table.

By the end of this lesson you can:

  • I identify protons, neutrons, electrons.
  • I use atomic number and mass number.
  • I describe trends in the periodic table.

The periodic table is an apartment directory. Each element lives in a specific apartment: its floor (period) tells you how many electron shells, and its column (group) tells you how many outer electrons — and so how it behaves.

Atoms, elements and the periodic table

ATOMIC STRUCTURE — the building block of matter.

Every atom has a NUCLEUS at the centre and ELECTRONS orbiting around it.

### Subatomic Particles

| Particle | Symbol | Charge | Location | Relative mass |
|----------|--------|--------|----------|---------------|
| Proton | p⁺ | +1 | Nucleus | 1 |
| Neutron | n | 0 | Nucleus | 1 |
| Electron | e⁻ | −1 | Shells (orbits) | ~0 |

Atoms are ELECTRICALLY NEUTRAL: number of protons = number of electrons.

### Key Numbers
• ATOMIC NUMBER (Z) = number of protons. Unique to each element.
• MASS NUMBER (A) = protons + neutrons.
• Number of neutrons = A − Z.

Example: Sodium (Na): Z = 11, A = 23. Neutrons = 23 − 11 = 12.

### Electron Shells (First 20 elements)
Electrons fill shells from the inside out:
• 1st shell: maximum 2 electrons.
• 2nd shell: maximum 8 electrons.
• 3rd shell: maximum 8 electrons (for elements up to Ca).

Shell configurations:
• Hydrogen (Z=1): 1
• Carbon (Z=6): 2, 4
• Oxygen (Z=8): 2, 6
• Sodium (Z=11): 2, 8, 1
• Chlorine (Z=17): 2, 8, 7

### Isotopes
ISOTOPES are atoms of the SAME element with DIFFERENT numbers of neutrons.
• Same Z (same element), different A.
• Examples: Carbon-12 (6p, 6n) and Carbon-14 (6p, 8n).
• Carbon-14 is radioactive — used in radiocarbon dating of archaeological remains (Australian Indigenous artifacts have been dated this way).

### THE PERIODIC TABLE
Elements arranged by INCREASING ATOMIC NUMBER.

PERIODS (rows) — elements in the same period have the same number of electron shells.
• Period 1: H, He (1 shell).
• Period 2: Li to Ne (2 shells).
• Period 3: Na to Ar (3 shells).

GROUPS (columns) — elements in the same group have the same number of OUTER SHELL electrons.
• Group 1: 1 outer electron (very reactive metals).
• Group 18: full outer shell (unreactive noble gases).

### Key Groups

| Group | Name | Properties | Examples |
|-------|------|-----------|---------|
| 1 | Alkali metals | Soft, react vigorously with water, 1 outer e⁻ | Li, Na, K |
| 2 | Alkaline earth metals | Reactive, 2 outer e⁻ | Mg, Ca |
| 17 | Halogens | Reactive non-metals, 7 outer e⁻ | F, Cl, Br |
| 18 | Noble gases | Unreactive, full outer shell | He, Ne, Ar |

### Periodic Trends
• Moving RIGHT across a period: increases in outer electrons, often less metallic.
• Moving DOWN a group: more electron shells, atoms are LARGER, outer electrons are less tightly held.
• Reactivity of metals INCREASES down Group 1 (potassium more reactive than lithium).
• Reactivity of non-metals DECREASES down Group 17 (fluorine most reactive halogen).

### Australian connection
The Olympic Dam mine in South Australia contains vast deposits of uranium ore. Uranium (Z = 92) is the heaviest naturally occurring element, found at the bottom of the periodic table.

Atomic Calculations

  • Atomic number: Z = \text{number of protons}
  • Mass number: A = \text{protons} + \text{neutrons}
  • Neutrons: n = A - Z
  • Neutral atom: \text{electrons} = \text{protons} = Z

Remember

  • Z = number of protons: This uniquely defines the element. Changing Z changes the element.
  • Neutral atom: Electrons = Protons. Gain electrons → anion (negative). Lose electrons → cation (positive).
  • Noble gases (Group 18): Full outer shell → chemically stable → very unreactive.

Worked example (Medium)

For oxygen (Z = 8, A = 16), state protons, neutrons, electrons and electron shell configuration.

  1. Protons. Z = 8 → 8 protons.
  2. Electrons. Neutral atom → 8 electrons.
  3. Neutrons. A − Z = 16 − 8 = 8 neutrons.
  4. Shell config. 1st shell: 2. Remaining: 8 − 2 = 6 in 2nd shell. Configuration: 2, 6.

Answer: Oxygen: 8p⁺, 8n, 8e⁻, configuration 2,6.

Worked example (Hard)

Element X has 11 protons, 12 neutrons, and loses 1 electron to form an ion. (a) Identify element X. (b) Write the mass number. (c) Describe the ion formed.

  1. Identify. Z = 11 → Sodium (Na).
  2. Mass number. A = 11 + 12 = 23.
  3. Ion. Losing 1 electron: electrons = 11 − 1 = 10. More protons than electrons → positive ion (cation). Ion: Na⁺ (sodium ion).
  4. Electron configuration. Na⁺ has 10 electrons: 2, 8 — same as neon (noble gas configuration).

Answer: Sodium-23 (²³Na) forms the Na⁺ cation (2, 8 electron configuration).

Common mistakes

  • Confusing mass number with atomic mass (relative atomic mass) Mass number (A) is the count of protons + neutrons for ONE isotope. Relative atomic mass is the weighted average across all naturally occurring isotopes. Fix: Use A = p + n for specific isotopes. The decimal number on the periodic table (e.g., Cl = 35.5) is the relative atomic mass.
  • Thinking isotopes are different elements Isotopes have the SAME number of protons (same Z = same element) — only neutron count differs. Fix: Carbon-12 and Carbon-14 are both carbon (Z = 6). They differ only in neutron number (6 vs 8).
  • Placing electrons in the wrong shells (e.g., putting 10 electrons in the second shell) Shells have maximum capacities. Shell 1: 2 max. Shell 2: 8 max. Fix: Fill shell 1 first (max 2), then shell 2 (max 8), then shell 3. For Z = 11 (Na): 2, 8, 1.

Practice questions

  1. An element has atomic number 17 and mass number 35. How many neutrons does it have?

    1. 17
    2. 18
    3. 35
    4. 52

    (1 marks)

  2. Explain why elements in the same group of the periodic table have similar chemical properties.

    (2 marks)

  3. Write the electron shell configuration for potassium (K, Z = 19).

    (2 marks)

  4. Carbon-12 and carbon-14 are isotopes. What is the same and what is different?

    (2 marks)

  5. Explain why reactivity increases down Group 1 (alkali metals) but decreases down Group 17 (halogens).

    (3 marks)

Going Deeper — Ions and Ionic Bonding

When atoms gain or lose electrons they become IONS:
• CATIONS (positive): metal atoms that LOSE electrons.
Na → Na⁺ + e⁻ (sodium loses 1 electron)
Mg → Mg²⁺ + 2e⁻ (magnesium loses 2 electrons)
• ANIONS (negative): non-metal atoms that GAIN electrons.
Cl + e⁻ → Cl⁻ (chlorine gains 1 electron)
O + 2e⁻ → O²⁻ (oxygen gains 2 electrons)

IONIC BONDING occurs between a cation and anion:
The electrostatic attraction between Na⁺ and Cl⁻ forms sodium chloride (table salt).

The driving force: atoms become more stable by achieving the electron configuration of the nearest noble gas.

Common ions to know:
• Na⁺, K⁺, Li⁺ (Group 1 form +1 ions)
• Mg²⁺, Ca²⁺ (Group 2 form +2 ions)
• Al³⁺ (Group 13 forms +3 ions)
• Cl⁻, Br⁻, F⁻ (Group 17 form −1 ions)
• O²⁻, S²⁻ (Group 16 form −2 ions)

Summary

  • Atoms: nucleus (protons + neutrons) + electrons in shells.
  • Z = atomic number = number of protons (unique to element).
  • A = mass number = protons + neutrons. Neutrons = A − Z.
  • Isotopes: same Z, different A (different neutron count).
  • Periodic table: periods = rows (same number of shells); groups = columns (same outer electrons → similar properties).
  • Noble gases (Group 18): full outer shell → chemically unreactive.

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