All matter is composed of atoms, and every atom is built from three fundamental subatomic particles. Understanding these particles and how they define an element is the foundation of all chemistry at A-Level.
🔑 Key Principle
The chemical properties of an element are determined by its electron configuration, specifically the arrangement of outermost (valence) electrons. The physical properties such as mass and density depend primarily on the nucleus.
The Three Subatomic Particles
You need to know the relative charge and relative mass of each particle. Absolute values are not required at this level.
| Particle | Relative Charge | Relative Mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | -1 | 1/1836 (negligible) | Electron cloud / shells |
In calculations, the mass of electrons is treated as negligible. This is why the relative atomic mass of an atom is essentially the sum of its protons and neutrons (nucleons). You will never need to account for electron mass in mole or mass calculations.
Atomic Number and Mass Number
Atomic Number (Z)
The number of protons in the nucleus. This defines which element an atom is. Change Z and you change the element entirely.
Mass Number (A)
The total number of protons + neutrons (collectively called nucleons) in the nucleus. Used to calculate neutrons: N = A - Z
Standard Notation
\( ^A_Z X \)
For example, \( ^{23}_{11}\text{Na} \) has 11 protons, 12 neutrons, and 11 electrons (in a neutral atom).
Isotopes
Atoms of the same element (same number of protons) with different numbers of neutrons. They have the same atomic number but different mass numbers.
Because isotopes have the same electron configuration, they have identical chemical properties. However, they differ in mass, which means they can have different:
- Physical properties (density, rate of diffusion, melting point)
- Nuclear stability (some isotopes are radioactive)
\(^{35}\text{Cl}\): 17 protons, 18 neutrons (35 - 17), 17 electrons
\(^{37}\text{Cl}\): 17 protons, 20 neutrons (37 - 17), 17 electrons
Both isotopes have 17 electrons and therefore identical electron configurations (2, 8, 7). This is why they react identically in chemical reactions.
Ions
When atoms gain or lose electrons, they form ions. The number of protons remains unchanged.
Cations (+)
Formed when atoms lose electrons. Fewer electrons than protons. Metals typically form cations.
Example: Na → Na⁺ + e⁻
Anions (-)
Formed when atoms gain electrons. More electrons than protons. Non-metals typically form anions.
Example: Cl + e⁻ → Cl⁻
A common mistake is to say that an ion has "more protons" or "fewer protons". The number of protons never changes during ion formation. Only the number of electrons changes. If you write "Na⁺ has 11 protons and 10 electrons", you must make clear that it has lost one electron, not gained a proton.
Calculating Particles in Ions
Protons: 26 (given by atomic number)
Neutrons: 56 - 26 = 30
Electrons: 26 - 3 = 23 (lost 3 electrons to form the 3+ charge)
Relative Atomic Mass and Isotopic Abundance
The weighted mean mass of an atom of an element compared to one twelfth of the mass of a carbon-12 atom.
Because most elements exist as a mixture of isotopes, the Ar value shown on the periodic table is a weighted average. This is why chlorine has an Ar of 35.5 rather than a whole number.
\( A_r = \frac{(35 \times 75.0) + (37 \times 25.0)}{100} = \frac{2625 + 925}{100} = \frac{3550}{100} = 35.5 \)
Always show your working clearly in Ar calculations. Even if the answer seems obvious, marks are awarded for the method. Make sure you divide by the sum of the abundances (usually 100, but check the question as sometimes percentages don't add to exactly 100).
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