3.1 Atomic Models

This topic traces the historical development of atomic theory — how our understanding of the atom changed over time, from Dalton's simple model to the modern quantum mechanical model.

Main Idea

The concept of the atomic model evolved over time as our understanding of atomic structure deepened through experimental observations and theoretical advances. Several important models of the atom had been proposed throughout history, each contributing to the understanding of atomic behaviour and properties. The most important atomic models are:

• Dalton's model
• Rutherford's model
• Bohr's model
• Quantum Mechanical Model

Simple explanation: An atomic model is a tool for understanding the structure and behaviour of atoms. It is not a physical model — it is a conceptual imagination that helps explain experimental observations. As research and technology progress, scientists continue to improve our knowledge and atomic models.

Core Concept

Each model improved on the previous one:

Detailed Explanation

Dalton's Model (1803)

In 1803, the British chemist John Dalton presented a scientific theory on the existence and nature of matter. This theory is called Dalton's atomic theory.

Main postulates:

• All elements are composed of tiny indivisible particles called atoms.
• Atoms of a particular element are identical. They have same mass and same volume.
• During chemical reactions atoms combine or separate or re-arrange. They combine in simple ratios.
• Atoms can neither be created nor destroyed.

Dalton was able to explain quantitative results that scientists of his time had obtained in their experiments. He nicely explained the laws of chemical combinations. His brilliant work became the main stimulus for the rapid progress of chemistry during the nineteenth century.

However, series of experiments performed in 1850s and beginning of twentieth century clearly demonstrated that atom is divisible and consists of subatomic particles — electrons, protons and neutrons.

Rutherford's Experiment (1911)

In 1911 Rutherford performed an experiment in order to know the arrangement of electrons and protons in atoms.

The experiment: Rutherford bombarded a very thin gold foil about 0.00004 cm thickness with α-particles. He used α-particles obtained from the disintegration of polonium. α-particles are helium nuclei that are doubly positively charged (He²⁺).

• Most of these particles passed straight through the foil.
• Only few particles were slightly deflected.
• But one in 1 million was deflected through an angle greater than 90° from their straight paths.

Rutherford performed a series of experiments using thin foils of other elements. He observed similar results from these experiments.

Rutherford's conclusions:

• Since majority of the α-particles passed through the foil undeflected, most of the space occupied by an atom must be empty.
• The deflection of a few α-particles through angles greater than 90° shows that these particles are deflected by electrostatic repulsion between the positively charged α-particles and the positively charged part of atom.
• Massive α-particles are not deflected by electrons.

Rutherford's model: On the basis of conclusions drawn from these experiments, Rutherford proposed a planetary model (similar to the solar system) for an atom.

• An atom is a neutral particle.
• The mass of an atom is concentrated in a very small dense positively charged region. He named this region as nucleus.
• The electrons are revolving around the nucleus in circles. These circles are called orbits.
• The centrifugal force due to the revolution of electrons balances the electrostatic force of attraction between the nucleus and the electrons.

Defects in Rutherford's Atomic Model:

Rutherford's model of an atom resembles our solar system. It has following defects:

• Classical physics suggests that electron being charged particle will emit energy continuously while revolving around the nucleus. Thus the orbit of the revolving electron becomes smaller and smaller until it would fall into the nucleus. This would collapse the atomic structure.
• If revolving electron emits energy continuously it should form a continuous spectrum.

Bohr's Atomic Theory (1913)

In 1913 Neil Bohr proposed a model for an atom that was consistent with Rutherford's model. But:

Bohr's postulates:

• The electron in an atom revolves around the nucleus in one of the circular orbits. Each orbit has a fixed energy. So each orbit is also called energy level.
• The energy of the electron in an orbit is proportional to its distance from the nucleus. The farther the electron is from the nucleus, the more energy it has.
• The electron revolves only in those orbits for which the angular momentum of the electron is an integral multiple of h/2π where h is Planck's constant (its value is 6.626 × 10⁻³⁴ J·s).
• Light is absorbed when an electron jumps to a higher energy orbit and emitted when an electron falls into a lower energy orbit. Electron present in a particular orbit does not radiate energy.
• The energy of the light emitted is exactly equal to the difference between the energies of the orbits.

Equation:

Δ E = E₂ - E₁

Where ΔE is the energy difference between any two orbits with energies E₁ and E₂.

Limitation: Bohr model does not depict the three-dimensional aspect of an atom.

Quantum Mechanical Model (Current Model)

This is the current model used by modern science to describe the structure of the atom. It incorporates the principles of quantum mechanics and treats electrons as wave-particle entities.

Instead of exact orbits, it defines probability regions, called orbitals, where electrons are likely to be found.

Heisenberg Uncertainty Principle:

Heisenberg uncertainty principle is one of the fundamental concepts of quantum mechanics and is named after the German physicist Werner Heisenberg, who formulated it in 1927.

This principle states that it is impossible to simultaneously determine the exact location and future trajectory of an electron. As a result, plotting the electron orbit around the nucleus becomes an irresistible challenge.

Explanation: Imagine that you have a single hydrogen atom and you decide to observe the position of that single electron at a given moment. Shortly after you repeated this process, the electron moved to another position. This means that from the original location to the next one is completely unknown to you. Continuous repetition of this process allows the gradual construction of a three-dimensional map representing the likely locations where the electron is expected to exist. You cannot know for sure where an electron is and where it goes next. This makes it impossible to draw the orbit of the electron around the nucleus.

In hydrogen, the electron has the potential to exist anywhere in the spherical region surrounding the nucleus. 95% of the time, the electron will be in a relatively simple region of space close to the nucleus, called an orbital. An orbital is the region of space where the electron lives.

de Broglie's Contribution:

Louis de Broglie, a French physicist, in 1924 proposed dual nature of electrons. He suggested that sub-atomic particles like electrons can exhibit both particle-like and wave-like behaviour. His idea opened the door for new possibilities in understanding behaviour of sub-atomic particles. This concept made a significant contribution to the development of quantum mechanics.

Davisson and Germer (1927):

In 1927, Davisson and Germer, experimentally confirmed the de Broglie hypothesis that electron has wave like behaviour. This discovery laid the foundation for the Modern Quantum Mechanics.

Understanding Atomic Models

An atomic model is a tool for understanding the structure and behavior of atoms and their interactions in chemical reactions. Any atomic model helps us understand the structure of an atom. An atomic model is not a physical model, but represents a conceptual imagination. This helps to explain experimental observations of atomic behavior. The atomic model gives us a simplified representation of complex reality. As research and technology progress, scientists continue to improve our knowledge and atomic models.

A Simple View of the Structure of an Atom

The nucleus of an atom is in the center. It contains protons and neutrons. Protons and neutrons are collectively called nucleons. The nucleus is surrounded by electrons in shells. Protons and neutrons are massive particles. The mass of electrons is so small that in practice, the mass of an atom is concentrated in the nucleus.

Nuclear Force

The nucleus contains protons and neutrons. Protons are positively charged and neutrons are neutral. The nucleus has no negative charge. The positively charged protons must cancel each other out and the nucleus must break apart. But atoms are stable and have existed for billions of years. Therefore, there must be some kind of attraction that connects them.

No electrostatic or magnetic forces occur within the core. This is because these forces involve both attraction and repulsion. Therefore, the force that binds protons and neutrons together is a strong force. This force is called strong nuclear force.

This is defined as the strong attractive force that binds protons and neutrons together. This force is stronger than electrostatic or magnetic forces. This force exists between neutrons and neutrons, protons and protons, and neutrons and protons.

Important Points

• - Dalton's model (1803): atoms are indivisible; atoms of same element are identical; atoms combine in simple ratios; atoms cannot be created or destroyed.
• - Rutherford's experiment (1911): gold foil bombarded with α-particles; most passed straight through; few deflected; led to discovery of nucleus.
• - Rutherford's model: planetary model; nucleus is small, dense, positively charged; electrons revolve in orbits.
• - Defects of Rutherford's model: electron should emit energy continuously and fall into nucleus; should form continuous spectrum.
• - Bohr's model (1913): electrons revolve in fixed energy orbits; angular momentum is integral multiple of h/2π; light absorbed/emitted when electron jumps; ΔE = E₂ − E₁.
• - Quantum Mechanical Model: current model; electrons as wave-particle entities; probability regions called orbitals; Heisenberg uncertainty principle.
• - Heisenberg uncertainty principle (1927): impossible to simultaneously know exact location and future trajectory of an electron.
• - de Broglie (1924): dual nature of electrons — particle-like and wave-like.
• - Davisson and Germer (1927): experimentally confirmed de Broglie hypothesis.
• - Nuclear force: strong attractive force that binds protons and neutrons together; stronger than electrostatic or magnetic forces.

Remember

• - Dalton: atom indivisible.
• - Rutherford: nucleus discovered; planetary model; α-particles deflected.
• - Bohr: fixed energy orbits; ΔE = E₂ − E₁.
• - Quantum Mechanical: orbitals (probability regions); Heisenberg uncertainty principle.
• - de Broglie: dual nature of electrons.
• - Davisson and Germer: confirmed wave nature of electrons.
• - Strong nuclear force: binds protons and neutrons; stronger than electrostatic/magnetic forces.

Common Mistakes

• 1. Thinking Dalton's model is still accepted. It was replaced after the discovery of subatomic particles.
• 2. Thinking Rutherford's model had no defects. It could not explain why electrons don't fall into the nucleus.
• 3. Thinking Bohr's model is the current model. It was replaced by the quantum mechanical model.
• 4. Confusing orbits with orbitals. Orbits are fixed paths (Bohr); orbitals are probability regions (quantum mechanical).
• 5. Thinking the uncertainty principle means we know nothing. It means we cannot know both position and momentum exactly at the same time.
• 6. Forgetting the strong nuclear force. It is what holds the nucleus together despite proton–proton repulsion.

Quick Revision

• Dalton's model: atoms indivisible, identical for same element, combine in simple ratios.
• Rutherford's experiment: α-particles on gold foil; most passed through; few deflected >90°; led to nuclear model.
• Rutherford's defects: electrons should lose energy and fall into nucleus; should give continuous spectrum.
• Bohr's model: fixed energy orbits; angular momentum quantized; ΔE = E₂ − E₁; light absorbed/emitted on jumps.
• Quantum Mechanical Model: electrons as wave-particle entities; probability regions (orbitals); current model.
• Heisenberg uncertainty principle: impossible to know exact position and future trajectory of electron simultaneously.
• de Broglie: dual nature of electrons (1924).
• Davisson and Germer: confirmed wave nature (1927).
• Strong nuclear force: binds protons and neutrons; stronger than electrostatic/magnetic forces.

60-Second Revision

• Dalton: atom indivisible.
• Rutherford: nucleus discovered; planetary model.
• Bohr: fixed energy orbits; ΔE = E₂ − E₁.
• Quantum Mechanical: orbitals; probability regions.
• Heisenberg: uncertainty principle.
• de Broglie: dual nature of electrons.
• Strong nuclear force: binds nucleons.
• Warning: Rutherford's model had defects — electrons would fall into the nucleus.