# Atoms and subatomic particles

> IB Chemistry HL · S1: Models of the particulate nature of matter
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u1-atoms-and-subatomic-particles/

This sub-topic introduces the fundamental structure of atoms, the key properties of protons, neutrons and electrons, and core notation for nuclides, ions and isotopes, forming the foundation for all further study in IB Chemistry HL.

**Prerequisites:** Basic understanding that matter is made of atoms

## Learning objectives

- Identify the three main subatomic particles and their relative properties
- Calculate the number of protons, neutrons and electrons for neutral atoms and ions
- Interpret standard nuclide notation correctly
- Distinguish between isotopes and compare their properties

## Properties of Subatomic Particles

**Atom** — The smallest unit of an element that retains all the chemical properties of that element. Atoms consist of a dense, positively charged nucleus surrounded by a cloud of negatively charged electrons.

| Particle | Relative Mass | Relative Charge | Location |
| --- | --- | --- | --- |
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/1836 | -1 | Electron cloud outside nucleus |

> **tip**
>
> IB exams almost always ask for relative mass and charge, not absolute values in kg or coulombs. You only need to memorize the relative values above.

**Worked example:** Compare the relative mass and charge of a proton and an electron.

1. Recall the relative properties of each particle from the table:
2. $$\text{Proton: mass} = 1, \text{charge} = +1$$
3. Recall the properties of an electron:
4. $$\text{Electron: mass} = \frac{1}{1836} \approx 0, \text{charge} = -1$$
5. Conclusion: A proton has ~1836 times the mass of an electron, and equal magnitude opposite charge.

## Nuclide Notation and Particle Counting

**Nuclide Notation** — Standard notation for a specific atomic species, where $Z$ = atomic number (number of protons), $A$ = mass number (total protons + neutrons), and $X$ = the element's chemical symbol.

*Notation:* _Z^A X

For neutral atoms, the number of protons equals the number of electrons, so net charge is zero. For charged ions, the number of electrons differs from the proton number: positive cations lose electrons, negative anions gain electrons. The number of protons and neutrons is never affected by ion formation.

**Worked example:** Find the number of protons, neutrons and electrons in the $_{17}^{35}\text{Cl}^-$ ion.

1. Extract values from the notation: $Z = 17$ (proton number), $A = 35$ (mass number), charge = -1
2. Number of protons = $Z$ = 17 (proton number never changes for an element)
3. Calculate number of neutrons as $A - Z$:
4. $$n = 35 - 17 = 18$$
5. For a -1 charge, add 1 electron to the proton number: electrons = 17 + 1 = 18
6. Final result: 17 protons, 18 neutrons, 18 electrons

**Check your understanding**

Test your understanding:

1. How many electrons are in $_{12}^{24}\text{Mg}^{2+}$?

   - 10
   - 12
   - 14
   - 24

   *Why:* Correct: $Z=12$ protons, +2 charge means 12 - 2 = 10 electrons. Charge only changes electron count.

## Isotopes and Their Properties

**Isotopes** — Atoms of the same element that share the same atomic number (number of protons) but have different mass numbers (different number of neutrons).

Isotopes of the same element have identical chemical properties because chemical reactivity depends on electron number, which is the same for all isotopes. They differ in physical properties like mass, density, boiling point and melting point, due to their different mass. Some isotopes are radioactive (radioisotopes) and have uses in medicine, archaeology and industry.

**Worked example:** Compare the subatomic composition and properties of carbon-12 and carbon-14, two common isotopes of carbon.

1. Carbon always has an atomic number $Z=6$, so both isotopes have 6 protons.
2. Calculate neutron count for each: carbon-12 (A=12) has 12 - 6 = 6 neutrons; carbon-14 (A=14) has 14 - 6 = 8 neutrons.
3. Both are neutral, so both have 6 electrons.
4. Properties comparison: chemical properties are identical. Physical properties differ: carbon-14 is heavier and radioactive, while carbon-12 is stable.

## Common pitfalls

- **Wrong:** Using atomic number Z as the electron count for all species, regardless of charge.
  - Why it fails: Only neutral atoms have equal numbers of protons and electrons. Ions gain or lose electrons, changing the electron count.
  - Correct: For +n cations: electrons = Z - n. For -n anions: electrons = Z + n. Proton count always equals Z.
- **Wrong:** Calculating neutron count as Z - A instead of A - Z.
  - Why it fails: Mass number A is the sum of protons and neutrons, so it is always larger than atomic number Z. This mistake gives a negative neutron count.
  - Correct: Neutron number = mass number (A) minus atomic number (Z): $n = A - Z$.
- **Wrong:** Claiming isotopes of the same element have different chemical properties.
  - Why it fails: Chemical reactivity is controlled by electron arrangement, which is identical for isotopes of the same element.
  - Correct: Isotopes have identical chemical properties, and only differ in physical properties like mass and density.
- **Wrong:** Using the average relative atomic mass from the periodic table as the mass number A for a specific nuclide.
  - Why it fails: Relative atomic mass is a weighted average of all isotopes of an element, often not a whole number. Mass number is a count of nucleons for a single nuclide, always a whole number.
  - Correct: Only use the mass number given for the specific nuclide in the question, not the average from the periodic table.

## Cheatsheet

| Quantity | Calculation | Rule for Ions |
| --- | --- | --- |
| Proton number | Equal to Z (atomic number) | Unchanged by charge |
| Neutron number | A - Z | Unchanged by charge |
| Electrons (neutral atom) | Equal to Z | N/A |
| Electrons (+n cation) | Z - n | Subtract charge value |
| Electrons (-n anion) | Z + n | Add charge magnitude |

## What's next

This sub-topic forms the foundation of all atomic structure and chemical bonding concepts in IB Chemistry HL. You will build on this core understanding to explore the historical development of atomic models, followed by learning how mass spectrometry separates isotopes and allows calculation of average relative atomic mass. Mastery of subatomic particle counting and nuclide notation is critical for all subsequent IB Chemistry topics, including stoichiometry, chemical bonding, and nuclear chemistry, and this foundational content is regularly tested in both Paper 1 and Paper 2 exams.

- [The nuclear atom and mass spectrometry](https://www.owlsprep.com/study/ib-chemistry-hl-u1-the-nuclear-atom-and-mass/)
- [Electron configurations](https://www.owlsprep.com/study/ib-chemistry-hl-u1-electron-configurations/)
- [The mole concept](https://www.owlsprep.com/study/ib-chemistry-hl-u1-the-mole-concept/)

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