Subatomic Particles and the Nuclear Atom
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You should be comfortable with basic arithmetic, the concept of charge, and the idea that matter is made of atoms, as introduced in the previous lecture.
Why this matters
Understanding the internal structure of the atom is the foundation for all subsequent chemistry.
Before the early twentieth century, atoms were considered indivisible. The discovery of subatomic particles revealed that atoms are complex systems with a dense, positively charged nucleus surrounded by a cloud of electrons. This nuclear model explains why matter has mass, why it carries charge, and why elements exhibit periodic chemical behaviour. Without this structural understanding, concepts like bonding, reactivity, and spectroscopy remain opaque.
The Nuclear Model in Plain Language
Imagine a football stadium. If the nucleus were a pea placed in the centre of the pitch, the electrons would be like gnats flying randomly around the stands. This analogy highlights two key features: the nucleus is incredibly small compared to the total size of the atom, yet it contains almost all of the atom's mass. The electrons occupy the vast empty space surrounding the nucleus. This arrangement explains why atoms can be compressed or deformed (because they are mostly empty space) and why they interact primarily through their outer electrons, which are far from the dense core.
Discovery of the Electron
The electron was the first subatomic particle to be identified, discovered by J.J. Thomson in 1897 through cathode ray tube experiments. Thomson observed that cathode rays were deflected by electric and magnetic fields, proving they consisted of negatively charged particles. Crucially, the charge-to-mass ratio (\(e/m\)) of these particles was found to be independent of the gas used in the tube, indicating that electrons are universal constituents of all atoms. This discovery shattered the idea of the atom as an indivisible unit and established the existence of subatomic particles.
The electron is much lighter than the atom itself, which implies that the positive charge and most of the mass must reside elsewhere in the atom.
The Gold Foil Experiment and the Nucleus
From 1909, Hans Geiger and Ernest Marsden, working under Ernest Rutherford, carried out the alpha-particle scattering experiments. They fired alpha particles (positively charged helium nuclei) at a thin sheet of gold foil. According to the prevailing "plum pudding" model, where positive charge was spread evenly throughout the atom, the alpha particles should have passed through with only slight deflections. Instead, most passed straight through, but a small fraction were deflected at large angles, and some even bounced back. This indicated that the positive charge and mass were concentrated in a tiny, dense region at the centre, which Rutherford described in 1911 and which came to be called the nucleus. The large deflections occurred when alpha particles came close to this dense positive core and experienced strong electrostatic repulsion.
The inverse-square nature of the electrostatic force explains why the deflection is so sensitive to the distance of closest approach to the nucleus.
Properties of Subatomic Particles
Atoms consist of three main subatomic particles: protons, neutrons, and electrons. Protons carry a positive charge of \(+1e\), neutrons are electrically neutral, and electrons carry a negative charge of \(-1e\). The proton and neutron have nearly identical masses, approximately \(1.67 \times 10^{-27}\ \mathrm{kg}\), while the electron is significantly lighter, at approximately \(9.11 \times 10^{-31}\ \mathrm{kg}\). The mass of an atom is therefore concentrated in the nucleus, which contains the protons and neutrons. The number of protons defines the element (atomic number, \(Z\)), while the number of neutrons determines the isotope. The electrons determine the chemical behaviour of the atom.
This mass ratio is crucial for understanding atomic mass calculations and nuclear stability.
A common misconception is that the electron orbits the nucleus like a planet around the sun. In reality, electrons do not follow fixed paths; they exist in probability distributions called orbitals. The "empty space" in an atom is not truly empty but is filled with the electron's wavefunction, which determines where the electron is likely to be found. Treating electrons as classical particles leads to incorrect predictions about atomic stability and spectral lines. C
Worked example
Problem. Calculate the total charge of the nucleus of a carbon-12 atom and the total charge of the entire neutral atom. Use the elementary charge \(e = 1.602 \times 10^{-19}\ \mathrm{C}\).
- Determine the number of protons in carbon. Carbon has an atomic number \(Z = 6\), so it has 6 protons.
- Calculate the nuclear charge. Each proton has a charge of \(+e\). Total nuclear charge \(Q_{\text{nuc}} = 6 \times (+1.602 \times 10^{-19}\ \mathrm{C})\).
- Determine the number of electrons in a neutral atom. A neutral atom has equal numbers of protons and electrons, so there are 6 electrons.
- Calculate the total charge of the atom. Total charge \(Q_{\text{atom}} = Q_{\text{nuc}} + Q_{\text{electrons}}\). Each electron has a charge of \(-e\). So \(Q_{\text{atom}} = 6e + 6(-e) = 0\).
Units check. The charge is expressed in coulombs (\(\mathrm{C}\)), which is the SI unit for electric charge. The sum of positive and negative charges cancels out exactly for a neutral atom.
Three quick tests
- What is the primary evidence that led Rutherford to propose the nuclear model of the atom?
Answer
The observation that a small fraction of alpha particles were deflected at large angles or bounced back from the gold foil, indicating a dense, positively charged centre. - Which subatomic particle has a mass approximately 1836 times greater than that of an electron?
Answer
The proton (and the neutron, which has a very similar mass). - If an atom has 11 protons and 12 neutrons, what is its atomic number and mass number?
Answer
The atomic number is 11 (number of protons), and the mass number is 23 (sum of protons and neutrons).
Back to CU-101 →
Every lecture and every result in this unit, with prerequisites and threads.
Lecture notes drafted with a local model and checked line by line by the editor; every number in the worked example was recomputed.