State the relative charges and relative masses of a proton, a neutron and an electron

Atoms, Elements and Compounds – Atomic Structure and the Periodic Table

Learning Objective

State the relative charges and relative masses of a proton, a neutron and an electron. Define the proton number (Z) and mass number (A), use these ideas to calculate relative atomic mass, and apply them to the electronic configuration, isotopes, ions and the main types of bonding of the first 20 elements.

1. Structure of the Atom

• Nucleus: contains positively‑charged protons and electrically neutral neutrons.
• Electron shells: electrons occupy discrete energy levels (K, L, M …) surrounding the nucleus.
• Proton number (Z): the number of protons in the nucleus; it uniquely identifies an element.
• Mass number (A): total number of protons + neutrons (A = Z + N).
• Isotopes: atoms of the same element (same Z) with different numbers of neutrons (different A). Isotopes have identical chemical behaviour because they have the same electron configuration.

2. Relative Charge and Relative Mass of Sub‑Atomic Particles

Why the electron mass is often ignored: 1 u ≈ 1836 × electron mass, so the electron contributes less than 0.1 % to the mass of a neutral atom. For most stoichiometric calculations the electron mass can be omitted without affecting the result.

3. Using Z and A – Calculating Relative Atomic Mass

Relative atomic mass (Ar) of an element is the weighted average of the masses of its naturally occurring isotopes.

Example (Nickel):

• ⁵⁸Ni: 68.08 % abundance, A = 58 u
• ⁶⁰Ni: 26.22 % abundance, A = 60 u
• ⁶¹Ni: 1.14 % abundance, A = 61 u
• ⁶²Ni: 4.56 % abundance, A = 62 u

Ar(Ni) = (0.6808 × 58) + (0.2622 × 60) + (0.0114 × 61) + (0.0456 × 62) ≈ 58.69 u

4. Ions

• Cation: atom that has lost one or more electrons; charge is positive (e.g., Na⁺, Mg²⁺).
• Anion: atom that has gained one or more electrons; charge is negative (e.g., Cl⁻, O²⁻).
• In a neutral atom, number of electrons = Z. When electrons are added or removed, the ion’s charge equals Δe⁻ (positive if electrons are lost, negative if gained).

5. Electronic Configuration

Two notations are useful:

1. K‑L‑M shell notation – counts electrons in each principal shell (e.g., 2, 8, 1 for Na).
2. Spectroscopic (sub‑shell) notation – uses the order 1s, 2s, 2p, 3s, 3p, 4s … (e.g., Na: 1s² 2s² 2p⁶ 3s¹).

Quick‑Reference Electronic Configurations (Z = 1 – 20)

6. Periodic‑Table Relationships

• Group I–VII (main‑group elements): the number of valence electrons equals the group number (e.g., Group II elements have two valence electrons, giving similar reactivity such as forming +2 cations).
• Group VIII (noble gases): full outer shell (2 e⁻ for He, 8 e⁻ for Ne, Ar, …) → very low chemical reactivity.
• Period number = number of occupied electron shells.  Period 1 → K shell only; Period 2 → K + L; Period 3 → K + L + M, etc.
• Trends linked to Z and electron configuration:
  • Metallic character increases down a group (larger atomic radius, easier loss of valence electrons).
  • Non‑metallic character increases across a period (higher effective nuclear charge, stronger attraction for electrons).

7. Types of Bonding Relevant to the First 20 Elements

7.1 Ionic Bonding

• Occurs between a metal that readily loses electrons (forms a cation) and a non‑metal that readily gains electrons (forms an anion).
• Resulting electrostatic attraction holds the ions together in a crystal lattice.
• Example: Na (Z = 11) loses one electron → Na⁺; Cl (Z = 17) gains one electron → Cl⁻; Na⁺ Cl⁻ forms sodium chloride.

7.2 Simple Covalent Molecules

• Atoms share pairs of electrons to achieve a full outer shell (usually the noble‑gas configuration).
• Typical of non‑metals with similar electronegativities.
• Examples: H₂, O₂, CO₂, NH₃.

7.3 Giant Covalent (Network) Structures

• All atoms are covalently bonded in a continuous three‑dimensional network.
• Very high melting points, hardness, and poor electrical conductivity (unless the network contains delocalised electrons).
• Examples in the first 20 elements: Carbon (diamond, graphite) and silicon (Si).

7.4 Metallic Bonding

• Metal atoms release some of their valence electrons to form a “sea of delocalised electrons” that moves freely throughout the lattice of positive metal ions.
• Explains metallic properties: conductivity, malleability, ductility, and luster.
• Examples: Na, Mg, Al, Ca (all in the first 20).

8. Summary Checklist (Cambridge IGCSE 0620 Core)

• Relative charge and mass of p, n, e – known.
• Define Z and A – known.
• Calculate relative atomic mass from isotopic data – example provided.
• Write electronic configurations (both shell and spectroscopic) for Z = 1‑20 – table included.
• Explain group‑valence‑electron relationship, noble‑gas rule and period‑shell rule – covered.
• Describe isotopes, ions, and the four main types of bonding (ionic, simple covalent, giant covalent, metallic) with examples – added.