Atomic Structure: Protons, Neutrons, Electrons and Electronic Configuration

What you'll learn

This topic forms the foundation of chemistry concepts tested in the CXC CSEC Integrated Science examination. You will master the structure of atoms, identify subatomic particles and their properties, calculate atomic numbers and mass numbers, and write electronic configurations using both notation systems. Questions on atomic structure appear consistently in Section A (multiple choice) and Section B (structured questions) of past papers.

Key terms and definitions

Atom — the smallest particle of an element that retains the chemical properties of that element and cannot be broken down by chemical means.

Proton — a positively charged subatomic particle found in the nucleus of an atom, with a relative mass of 1 and a relative charge of +1.

Neutron — a neutral subatomic particle (no charge) found in the nucleus with a relative mass of 1.

Electron — a negatively charged subatomic particle orbiting the nucleus in shells or energy levels, with negligible mass (1/1840) and a relative charge of -1.

Atomic number (Z) — the number of protons in the nucleus of an atom; this defines the element and equals the number of electrons in a neutral atom.

Mass number (A) — the total number of protons and neutrons in the nucleus of an atom.

Electronic configuration — the arrangement of electrons in shells (energy levels) around the nucleus of an atom.

Isotope — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.

Core concepts

Structure of the atom

The atom consists of a central nucleus containing protons and neutrons, surrounded by electrons arranged in shells or energy levels. The nucleus is extremely small compared to the overall size of the atom but contains nearly all the atom's mass.

Key characteristics of subatomic particles:

The atom as a whole is electrically neutral because the number of positively charged protons equals the number of negatively charged electrons. The strong nuclear force holds protons and neutrons together in the nucleus, overcoming the electrostatic repulsion between positive protons.

Atomic number and mass number

Every element is defined by its atomic number (symbol Z), which represents the number of protons. For example, all carbon atoms have 6 protons, all oxygen atoms have 8 protons, and all aluminium atoms (used in the bauxite industry in Jamaica and Guyana) have 13 protons.

The mass number (symbol A) equals protons plus neutrons. For a neutral atom:

• Number of protons = atomic number (Z)
• Number of electrons = atomic number (Z)
• Number of neutrons = mass number (A) - atomic number (Z)

Standard notation for representing atoms:

$$\text{Mass number (A)} \rightarrow \text{ }^{A}_{Z}\text{X}$$

Where X is the chemical symbol. Example: $^{23}_{11}\text{Na}$ represents a sodium atom with 11 protons, 11 electrons, and 12 neutrons (23 - 11 = 12).

Isotopes

Isotopes are atoms of the same element with identical atomic numbers but different mass numbers. They have the same chemical properties because they have the same number of electrons and therefore the same electronic configuration, which determines chemical behaviour.

Carbon has three naturally occurring isotopes:

• Carbon-12: $^{12}_{6}\text{C}$ (6 protons, 6 neutrons)
• Carbon-13: $^{13}_{6}\text{C}$ (6 protons, 7 neutrons)
• Carbon-14: $^{14}_{6}\text{C}$ (6 protons, 8 neutrons) — radioactive, used in dating

Chlorine exists as two main isotopes: chlorine-35 ($^{35}{17}\text{Cl}$) and chlorine-37 ($^{37}{17}\text{Cl}$), which explains why chlorine's relative atomic mass is 35.5 rather than a whole number.

The petroleum industry in Trinidad processes hydrocarbons containing carbon and hydrogen isotopes, and understanding isotope behaviour is crucial for refining processes.

Electronic configuration

Electrons occupy shells or energy levels around the nucleus. Each shell can hold a maximum number of electrons:

• 1st shell (closest to nucleus): maximum 2 electrons
• 2nd shell: maximum 8 electrons
• 3rd shell: maximum 8 electrons (for elements 1-20)
• 4th shell: maximum 2 electrons (for calcium, atomic number 20)

Electrons fill shells starting from the innermost shell, following the aufbau principle (building-up principle). Shells are filled in order of increasing energy, and each shell must be filled before electrons enter the next shell.

Two notation systems are used in CXC CSEC examinations:

1. Shell notation (comma-separated)

Write the number of electrons in each shell, separated by commas. Examples:

• Hydrogen (H, Z=1): 1
• Carbon (C, Z=6): 2,4
• Oxygen (O, Z=8): 2,6
• Sodium (Na, Z=11): 2,8,1
• Calcium (Ca, Z=20): 2,8,8,2

2. Orbital notation (using electron diagrams)

Draw circles representing shells with dots or crosses representing electrons. The nucleus can be shown as a central circle with the charge (e.g., +6 for carbon).

Electronic configuration and the Periodic Table

The group number in the Periodic Table equals the number of electrons in the outermost shell (valence electrons) for Groups 1, 2, and 13-18. The period number equals the number of shells containing electrons.

Examples relevant to Caribbean industries:

• Aluminium (Al, Z=13): 2,8,3 — Group 13, Period 3 — used in bauxite extraction in Jamaica
• Calcium (Ca, Z=20): 2,8,8,2 — Group 2, Period 4 — found in limestone used in cement production
• Iron (Fe, Z=26): 2,8,14,2 — transition metal used in steel manufacturing

Elements with the same number of valence electrons show similar chemical properties because chemical reactions involve only the outermost electrons. For example:

• Group 1 elements (lithium, sodium, potassium): all have 1 valence electron, all react vigorously with water
• Group 17 elements (chlorine, bromine, iodine): all have 7 valence electrons, all form salts with metals

Relationship between electronic structure and reactivity

Atoms react to achieve a stable electronic configuration, typically 8 electrons in the outermost shell (the octet rule), or 2 electrons for hydrogen and helium. This stability is characteristic of the noble gases (Group 18), which are unreactive.

Atoms achieve stability by:

• Losing electrons — metals form positive ions (cations)
• Gaining electrons — non-metals form negative ions (anions)
• Sharing electrons — forming covalent bonds

Example: Sodium (2,8,1) loses 1 electron to form Na⁺ ion (2,8), achieving the stable configuration of neon. Chlorine (2,8,7) gains 1 electron to form Cl⁻ ion (2,8,8), also achieving a stable configuration.

Worked examples

Example 1: Calculating subatomic particles

Question: An atom of potassium is represented as $^{39}_{19}\text{K}$. Calculate: (a) the number of protons [1 mark] (b) the number of electrons [1 mark] (c) the number of neutrons [2 marks] (d) Write the electronic configuration of potassium. [1 mark]

Solution:

(a) Number of protons = atomic number (Z) = 19 ✓

(b) Number of electrons = atomic number (for neutral atom) = 19 ✓

(c) Number of neutrons = mass number - atomic number = A - Z = 39 - 19 = 20 ✓✓

(d) Electronic configuration: 2,8,8,1 ✓ (2 electrons in first shell, 8 in second, 8 in third, 1 in fourth)

Example 2: Isotopes

Question: Copper has two isotopes: copper-63 and copper-65. Both isotopes have an atomic number of 29. (a) State what is meant by the term isotope. [2 marks] (b) Complete the table below: [4 marks]

(c) Explain why both isotopes have identical chemical properties. [2 marks]

Solution:

(a) Isotopes are atoms of the same element ✓ that have the same number of protons but different numbers of neutrons / different mass numbers. ✓

(b)

(c) Both isotopes have the same number of electrons ✓ / same electronic configuration, and chemical properties depend on the number and arrangement of electrons ✓ / valence electrons.

Example 3: Electronic configuration and the Periodic Table

Question: Chlorine (Cl) is used to purify water at the Water and Sewerage Authority facilities across Trinidad and Tobago. A chlorine atom has 17 protons. (a) Write the electronic configuration of chlorine. [1 mark] (b) State the group number and period number for chlorine in the Periodic Table. [2 marks] (c) Explain why chlorine is very reactive. [2 marks]

Solution:

(a) Electronic configuration: 2,8,7 ✓

(b) Group number: 17 or VII ✓ Period number: 3 ✓

(c) Chlorine has 7 electrons in its outer shell ✓ and needs only 1 more electron to achieve a stable configuration of 8 / to complete its outer shell, so it readily gains an electron from other atoms. ✓

Common mistakes and how to avoid them

• Mistake: Confusing atomic number with mass number. Correction: Atomic number (Z) = protons only; mass number (A) = protons + neutrons. The atomic number is always the smaller number and defines the element.

• Mistake: Incorrectly calculating neutrons by subtracting the wrong values (e.g., Z - A instead of A - Z). Correction: Always use the formula: Neutrons = Mass number - Atomic number (A - Z). Check that your answer makes sense; neutrons should never be negative.

• Mistake: Writing electronic configurations that violate the maximum shell capacity (e.g., writing 2,10 instead of 2,8,2 for magnesium). Correction: Remember the maximum electrons per shell: 1st shell = 2, 2nd shell = 8, 3rd shell = 8 (for elements 1-20). Fill shells in order without exceeding capacity.

• Mistake: Stating that isotopes have different chemical properties. Correction: Isotopes have identical chemical properties because they have the same number of electrons and same electronic configuration. They differ only in physical properties like mass and density.

• Mistake: Forgetting that the number of electrons equals the number of protons only in neutral atoms. Correction: In neutral atoms, electrons = protons. In ions, this differs: positive ions (cations) have fewer electrons than protons; negative ions (anions) have more electrons than protons.

• Mistake: Confusing group number with the total number of electrons instead of valence electrons. Correction: The group number (for main groups) equals the number of valence electrons (outermost shell only), not the total number of electrons in the atom.

Exam technique for Atomic Structure: Protons, Neutrons, Electrons and Electronic Configuration

• Command words matter: "State" requires a simple answer without explanation (1 mark). "Explain" or "Suggest why" requires reasoning with scientific terminology (usually 2-3 marks). "Calculate" requires working shown with correct units.

• Show all working in calculations: Even if you make an arithmetic error, you can still earn method marks. Always write the formula first: Neutrons = A - Z. For 3-mark calculation questions, typically 1 mark is for the correct method and 2 marks for the correct answer.

• Use standard notation correctly: When writing isotope notation $^{A}_{Z}\text{X}$, ensure the mass number (A) is top left and atomic number (Z) is bottom left of the element symbol. This appears regularly in structured questions worth 1-2 marks.

• Draw neat electron diagrams: Use a ruler for circles representing shells. Space electrons evenly around each shell. Label the nucleus clearly with the number of protons (e.g., +6) or write "nucleus" beside it. Such diagrams typically earn 2-3 marks in Paper 2.

Quick revision summary

Atoms contain protons (+1 charge, mass 1) and neutrons (neutral, mass 1) in the nucleus, with electrons (-1 charge, negligible mass) in shells around it. Atomic number (Z) = protons = electrons in neutral atoms. Mass number (A) = protons + neutrons. Electronic configuration shows electron arrangement in shells (maximum 2,8,8,2 for elements 1-20). Isotopes have same protons, different neutrons, identical chemical properties. Group number = valence electrons; period number = number of shells. Atoms react to achieve stable outer shell configurations, typically 8 electrons.