Subatomic Particles

Evidence for Subatomic Particles

Dalton's model explained fixed ratios in chemical reactions, but it still pictured atoms as solid spheres. Electrical experiments later showed that atoms contain smaller charged particles.

In the late 1800s, scientists began using low-pressure glass tubes and high voltage. A beam called a cathode ray appeared inside the tube. The beam could be bent by electric and magnetic fields. From this evidence, atoms started to look less like solid balls and more like structures with smaller particles inside.

The evidence sequence runs from Thomson's cathode ray experiment through Rutherford's gold foil experiment to Chadwick's discovery of the neutron.

J. J. Thomson discovered the electron in $1897$ and showed that this particle was far lighter than an atom.

Evidence from Three Experiments

Switch the mode below and compare the direction of motion, which path bends, and which path stays straight. Use those observations to identify the particle.

Cathode Rays Identified Electrons

Thomson used a cathode ray tube. Inside the tube, most air was removed, then a high voltage was applied. The beam that appeared was always attracted toward positive charge and pushed away from negative charge.

Opposite charges attract, while like charges repel. Because the cathode ray was attracted to positive charge $+$, its particles had negative charge $-$. We now call those particles electrons.

Thomson also found that cathode ray particles were far lighter than atoms and had the same properties even when the electrode metal changed. The unchanged properties showed that electrons were not fragments of one particular metal. They were common components of atoms.

In Thomson's model, positive charge is spread throughout the atom and electrons are embedded in that distribution. The model accounts for electrons, but it does not contain a small, dense nucleus. Rutherford's experiment later revealed that limitation.

A Tiny Dense Nucleus

Rutherford's Manchester team tested Thomson's model. At Rutherford's direction, Hans Geiger and Ernest Marsden fired positively charged alpha particles at very thin metal foils, including gold. If positive charge were spread out evenly as Thomson's model suggested, almost all alpha particles should pass through with only small bends.

A small fraction of the observations contradicted Thomson's prediction:

• most alpha particles passed through the gold foil with little change in direction
• a small fraction bent
• a very tiny fraction bounced backward

Now connect those observations to the location of positive charge inside the atom. Because almost all alpha particles passed straight through, most of the atom must be empty space. The few particles that bent sharply or bounced back showed that the positive charge and most of the atom's mass are packed into a very small region. That region is called the atomic nucleus.

To picture the scale, imagine enlarging an atom until it is as wide as a $100\ \text{m}$ stadium. Its nucleus would be only about $1\ \text{mm}$ across, roughly the size of a grain of sand at the center. A real atom is about $10^{-10}\ \text{m}$ across, while its nucleus is about $10^{-15}\ \text{m}$ across. The nuclear diameter is therefore about $10^5$ times smaller than the atomic diameter.

Hydrogen Nuclei Identified the Proton

The gold-foil evidence located the atom's positive charge, but it did not yet identify the particles inside the nucleus. In $1919$, Rutherford bombarded nitrogen with alpha particles and detected fast hydrogen nuclei coming out. A hydrogen nucleus carries one positive elementary charge and is now called a proton.

Rutherford interpreted the nuclear atom in $1911$ and demonstrated free protons in $1919$. The two steps are related but distinct events.

Neutron Mass and Charge

Protons explained the nucleus's positive charge, but the mass of the protons alone could not explain the full atomic mass. Scientists still needed to find another particle that added mass without adding charge.

In $1932$, James Chadwick showed evidence for a neutral particle with almost the same mass as a proton. That particle is called the neutron. Neutral means its charge is $0$, so neutrons can add mass to a nucleus without changing the atom's electric charge.

Chadwick showed that radiation from beryllium bombarded with alpha particles consisted of neutral particles with about the same mass as protons.

Isotopes are atoms of the same element with different numbers of neutrons. The element stays the same because the proton number stays fixed, while the mass changes with the neutron number.

Comparing the Three Subatomic Particles

This table compares each particle's charge, main location, and contribution to atomic mass.

A neutral atom has the same number of protons and electrons. For example, if an atom has $6$ protons and is still neutral, it also has $6$ electrons.

Remember these three relationships: protons determine element identity, neutrons help determine mass, and electrons determine many chemical behaviors. Use them together when comparing atomic charge, mass, and identity.