High School Physical Science Tutorial: The Atom

What Is an Atom?

Atoms are the fundamental building blocks of all matter in the universe. Everything you see, touch, and interact with is composed of atoms. Despite their tiny size, atoms hold the secrets of the physical world, connecting chemistry, physics, and biology.

This tutorial will dive into the history, structure, and significance of atoms, helping you understand why they are so crucial to science and everyday life.

The History of Atomic Theory

The concept of the atom dates back to ancient Greece, around 400 BCE, when the philosopher Democritus proposed that all matter is composed of tiny, indivisible particles called “atomos,” meaning “uncuttable.” While his idea was purely philosophical and lacked experimental evidence, it was revolutionary for its time.

Another ancient Greek philosopher, Aristotle, disputed Democritus’s theory and proposed that matter is uniform throughout and is not made of smaller particles. Though Aristotle’s theory was incorrect, it was accepted as scientific dogma for nearly 2,000 years.

John Dalton and the Modern Atomic Theory

In the early 19th century, an English scientist named John Dalton revived the idea of atoms with experimental backing. He proposed that:

• Matter is made up of small, indivisible particles called atoms.
• Atoms of a given element are identical in mass and properties.
• Atoms combine in simple whole-number ratios to form compounds. This marked the beginning of modern chemistry, bridging ancient philosophy with experimental science.

J.J. Thomson and the Electron

In 1897 the English physicist J.J. Thomson discovered the electron, a negatively charged subatomic particle, through his experiments with cathode rays. He proposed the “plum pudding model,” where electrons were embedded in a positively charged sphere, resembling raisins in a pudding.

Ernest Rutherford and the Nucleus

Ernest Rutherford’s famous gold foil experiment in 1911 revealed that an atom has a dense, positively charged nucleus at its center, with electrons orbiting around it. This discovery overturned the plum pudding model. Rutherford proposed that nearly all of the mass of an atom was tied up in the central nucleus that had a positive charge, and the electrons were less massive and encircled the nucleus.

Niels Bohr and Electron Orbits

The Danish physicist Niels Bohr refined Rutherford’s model by suggesting that electrons orbit the nucleus in specific energy levels or shells, much like planets around the sun. This model explained why atoms emit light in discrete spectra. James Chadwick, a student of Rutherford, concluded that the nucleus contained positive protons and electrically neutral neutrons.

Quantum Mechanics and the Modern Model

By the 20th century, the advent of quantum mechanics further refined atomic theory. The modern atomic model, developed by scientists like Erwin Schrödinger and Werner Heisenberg, depicts electrons as existing in “clouds” or regions of probability around the nucleus, rather than fixed orbits. The electron clouds are 100,000 times larger than the diameter of protons and neutrons in the nucleus. Since the electrons move so quickly about the nucleus, it is impossible to describe its exact location in space.

Modern View of the Atom

At the center of every atom is the nucleus, composed of particles called nucleons, which consists of positively charged protons and negatively charged neutrons. Neutrons and protons have nearly the same mass and are much more massive than the surrounding electrons. The positively charged protons hold the negatively charged electrons in their orbits through the interaction of their electric charges. To make an atom electrically neutral, the number of electrons and protons must be the same.

The neutrons in the nucleus have no direct effect on the structure of the electrons and do not affect the chemistry of the atom. The neutrons’ primary role is to act as a form of nuclear cement to hold the nucleus together. Nucleons are bound by an attractive nuclear force known as the strong nuclear force, which hold the nucleus together.

The strong nuclear force that binds together the nucleus is very short ranged compared to the electrical force that binds the electron to the protons. The electrical force falls off as the inverse square of the distance between the electrons and protons, but the strong nuclear force decreases much more rapidly. When two nucleons are just a few nucleon diameters distant, the strong unclear force they exert on each other is nearly zero. The presence of neutrons adds to the nuclear attraction that allows the protons to remain together in the nucleus without flying apart.

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Atoms are made up of three primary subatomic particles:

• Protons: Positively charged particles found in the nucleus. They determine the atomic number of an element. The mass of a proton is approximately 1.67x10^-24 g.
• Neutrons: Neutral particles also located in the nucleus. Together with protons, they contribute to the atom’s mass. The mass of a neutron is approximately that of a proton.
• Electrons: Small negatively charged particles that orbit the nucleus in energy levels. They are responsible for chemical bonding and interactions. The mass of the electron is 1,836 times smaller than that of the proton and neutron.

Nuclear scientists believe that the electron is not composed of smaller particles; however, the proton and neutron are made of smaller particles called quarks. Scientists study quarks using particle accelerators, such as the one at Fermi National Accelerator Laboratory at Batavia, Illinois. Giant magnets and strong electric fields are combined, and these accelerate small particles, such as protons and neutrons, to extremely high velocities. When they collide, they produce products that give scientists the inner structure of the smallest particles.

Electron Energy Levels

Electrons are not scattered randomly around the nucleus; they are organized in specific regions called energy levels or shells. These shells are like layers of an onion, surrounding the nucleus at increasing distances. The arrangement of electrons in these shells follows specific rules and plays a critical role in determining an atom’s chemical properties and how it interacts with other atoms.

Each shell has a maximum capacity for electrons, determined by the formula 2n², where n = 1,2,3,4… is the shell number. For instance:

• The first shell, closest to the nucleus, is the smallest and can hold only up to two electrons.
• The second shell is larger, accommodating up to eight electrons.
• The third shell can hold up to 18 electrons.
• The fourth shell and those beyond can hold even more, up to 32 electrons, but these capacities are only filled in larger, more complex atoms.

The way electrons fill these shells follows a specific order. Electrons occupy the lowest available energy level before filling higher levels. This is why the first shell fills up before the second, and so on.

The number and arrangement of electrons in the outermost shell, known as the valence shell, are especially important. These valence electrons are the ones involved in chemical reactions and bonding. For example, elements with full outer shells, like helium and neon, are chemically stable and unreactive, while elements with incomplete outer shells, like sodium or chlorine, tend to form bonds to achieve a stable configuration.

The concept of energy levels also helps explain why atoms emit or absorb light. When an atom absorbs energy, its electrons can move to higher energy levels, a process called excitation. When they return to their original levels, they release energy in the form of light. This phenomenon is the basis for technologies like neon lights and lasers.

By understanding electron energy levels, students can better grasp how elements behave in the periodic table and why certain combinations of atoms form molecules while others do not.

Atomic Number and Mass Number

Two important numbers that help us describe the identity and composition of an atom are the atomic number and the mass number. These values provide a snapshot of what makes each atom unique and how it contributes to the elements we encounter in our daily lives.

The atomic number of an atom is the number of protons in its nucleus. It is the defining characteristic of an element, determining its place on the periodic table and its fundamental properties. For example, the element hydrogen has an atomic number of 1, meaning every hydrogen atom has one proton in its nucleus. Similarly, oxygen, a crucial element for life, has an atomic number of 8, which means its nucleus contains eight protons.

The atomic number also determines the number of electrons in a neutral atom, as the number of electrons matches the number of protons. This balance of protons and electrons shapes an element’s chemical behavior, enabling scientists to predict how it will interact with other elements.

While the atomic number identifies an atom, the mass number provides a fuller picture of its internal structure. The mass number is the sum of protons and neutrons in the nucleus, as these subatomic particles make up nearly all the atom’s mass. Electrons are so light that their contribution to the mass is negligible.

For example, sodium (Na) has 11 protons and 12 neutrons, giving it a mass number of 23. While all atoms of a given element share the same number of protons, they can differ in their number of neutrons. This leads to variations called isotopes, where atoms of the same element have different mass numbers.

Finding the Number of Neutrons

Calculating the number of neutrons in an atom is straightforward. Subtract the atomic number from the mass number:

Number of Neutrons = Mass Number − Atomic Number

For example, consider the sodium atom:

Neutrons = 23(mass number) − 11(atomic number) =12 neutrons

Why Atomic and Mass Numbers are Important

Atomic and mass numbers are more than just numerical values—they are essential for understanding how elements behave in chemical reactions and physical processes. For instance, knowing these numbers allows scientists to:

• Identify isotopes and study their applications, such as in medical imaging or carbon dating.
• Predict the arrangement of elements on the periodic table and their chemical properties.
• Analyze the stability of an atom’s nucleus, which is important in nuclear reactions.

Isotopes

When the atoms of an element have different numbers of neutrons in the nucleus, they are called isotopes. The nucleus of most hydrogen atoms has one proton and no neutron. The isotope of hydrogen called deuterium is formed when the single proton is accompanied by the neutron. When there are two neutrons in the hydrogen nucleus, it forms the isotope called tritium. Every element has isotopes and generally, the lighter nuclei have roughly equal numbers of protons and neutrons, whereas the more massive the nucleus the more neutrons are required relative to the number of protons to bind the nucleus together.

Radioactivity

The first conclusive evidence that atoms were complex entities, rather than “indivisible” objects as stated in Dalton’s atomic theory of matter, came with the discovery of radioactivity by Henri Becquerel in 1896. Radioactivity is the term applied to the spontaneous decomposition of atoms of certain elements, such as radium and uranium, into other elements through the simultaneous production of rays of particles. There are three basic types of rays associated with radioactive substances:

1. Alpha (α) particles consisting of two protons and two neutrons. They carry a net +2 electric charge.
2. Beta (β) particles, which are high-speed, high-energy electrons emitted by the radioactive decay of an atomic nucleus. Beta particles carry a -1 electric charge.
3. Gamma (γ) radiation, which is a penetrating form of electromagnetic radiation that is produced in the process of radioactive decay of a nucleus. Gamma radiation is much more penetrative in materials than either alpha or beta radiation and carries no electric charge.