Atomic Theory

Atomic Structure

Now that we understand how atomic theory developed, we can examine what we currently know about the structure of the atom itself. Modern atomic structure is built from three fundamental subatomic particles, each with a specific location, charge, and mass. Together, these particles account for the mass, charge, and chemical behavior of every atom in the periodic table.

The modern atom consists of a dense central nucleus containing protons and neutrons, surrounded by electrons occupying the space outside the nucleus. Although the atom appears solid to us, it is actually mostly empty space, with nearly all of its mass concentrated within the tiny nucleus.

The Three Subatomic Particles

Every atom consists of three fundamental subatomic particles:

• Protons

• Neutrons

• Electrons

Each particle possesses unique properties that determine its role within the atom.

Properties of the Subatomic Particles

Table displaying the fundamental particles of an atom, including their symbols, charges, relative masses, and locations within the atom.

Mass is often expressed in:

Atomic Mass Units (amu)

where:

1 amu ≈ 1.66 × 10⁻²⁷ kg

Figure 1: The Three Subatomic Particles

Explore the fundamental subatomic particles—protons, neutrons, and electrons—that comprise atoms, highlighting their charges, relative masses, and locations within or outside the nucleus.

Relative Masses

One of the most important observations from the table is the enormous difference in mass between electrons and the particles found in the nucleus.

Protons and neutrons each possess a mass of approximately:

1 amu

Electrons possess a mass of only about:

1/1836 amu

As a result:

• Nearly all atomic mass resides in the nucleus.

• Electrons contribute negligibly to atomic mass.

Even though electrons contribute very little mass, they are extraordinarily important because they determine:

• Chemical bonding

• Chemical reactivity

• Periodic trends

• Physical properties

Figure 2: Relative Mass Comparison

Protons and neutrons have nearly the same mass, approximately 1 atomic mass unit (amu), while electrons are much lighter, with a mass about 1/1836 that of a proton or neutron, emphasizing their minimal contribution to atomic mass.

Reading Atomic Notation

Each element's atomic notation contains important information about its structure.

For example:

²⁴₁₂Mg²⁺

contains three important pieces of information:

• Mass number (upper left)

• Atomic number (lower left)

• Charge (upper right)

Understanding how to interpret these numbers allows us to determine the numbers of protons, neutrons, and electrons present in any atom or ion.

Figure 3: Atomic Notation Breakdown

Atomic notation illustrates key information about elements, including mass number, atomic number, element symbol, and ionic charge, providing a concise representation of atomic structure.

Atomic Number (Z)

The atomic number, symbolized by Z, is the number of protons in the nucleus.

For magnesium:

Z = 12

Therefore:

Magnesium contains 12 protons.

The atomic number is the defining characteristic of an element.

Every magnesium atom contains exactly:

12 protons

No other element contains exactly 12 protons.

If the number of protons changes, the element itself changes.

For example:

• 11 protons → sodium

• 12 protons → magnesium

• 13 protons → aluminum

The periodic table is organized according to increasing atomic number.

Figure 4: Atomic Number Defines the Element

The image explains how the atomic number, representing the number of protons, uniquely defines each element in the periodic table, highlighting magnesium as an example.

Mass Number (A)

The mass number, symbolized by A, equals the total number of protons and neutrons in the nucleus.

Mass Number Equation

A=p+nA = p + n

where:

• A = mass number

• p = protons

• n = neutrons

Because both protons and neutrons have masses close to 1 amu, the mass number serves as a whole-number approximation of atomic mass.

For magnesium-24:

A = 24

Since magnesium has:

12 protons

the number of neutrons is:

24 − 12 = 12 neutrons

Neutron Calculation

Number of neutrons = Mass number − Atomic number

Figure 5: Calculating Neutrons

Learn how to calculate the number of neutrons in an atom using magnesium-24 as an example, by subtracting the atomic number from the mass number.

Charge and the Number of Electrons

In a neutral atom:

Number of protons = Number of electrons

The positive charge of the protons exactly balances the negative charge of the electrons.

As a result:

Net charge = 0

Ions

When atoms gain or lose electrons, they become Ions.

Cations

Positive ions are called: Cations

They form when electrons are lost.

Examples:

• Na⁺

• Mg²⁺

• Al³⁺

Anions

Negative ions are called: Anions

They form when electrons are gained.

Examples:

• F⁻

• O²⁻

• N³⁻

Figure 6: Formation of Ions

Atoms become ions through the loss or gain of electrons, forming positively charged cations or negatively charged anions, as illustrated with sodium and chlorine atoms.

Calculating Electrons from Charge

For ions:

Number of electrons = Atomic number − Charge

Remember:

• Positive charges indicate electrons have been lost.

• Negative charges indicate electrons have been gained.

Example: Magnesium Ion

Mg²⁺

Atomic number: 12

Charge: +2

Electrons: 12 − 2 = 10 electrons

Magnesium still contains: 12 protons

but now contains only: 10 electrons

resulting in a net charge of: +2

Figure 7: Magnesium Ion Example

Illustration of how a magnesium atom becomes a magnesium ion by losing two electrons, resulting in a Mg²⁺ ion with a stable, full outer shell.

Worked Examples

The following examples illustrate how to determine the numbers of protons, neutrons, and electrons from atomic notation.

Table displaying the atomic and subatomic properties of various chemical species, including species names, atomic numbers, mass numbers, and counts of protons, neutrons, and electrons.

Figure 8: Determining Atomic Structure

Explore the relationships between protons, neutrons, and electrons in atoms and ions, with examples of how to calculate these subatomic particles for different elements and their charges.

Element Symbols

Each element is represented by a unique chemical symbol.

Symbols consist of:

• One or two letters

• First letter always capitalized

• Second letter, if present, always lowercase

Many symbols come directly from English names.

Examples:

• C = Carbon

• O = Oxygen

• N = Nitrogen

Others originate from Latin names.

Common Latin-Derived Symbols

This table lists chemical symbols, their corresponding elements, and their Latin origin names for quick reference.

Figure 9: Origins of Element Symbols

Explore the Latin origins of element symbols, highlighting the historical roots and meanings behind familiar chemical abbreviations.

How Elements Exist in Nature

Elements do not always exist as isolated individual atoms.

Different elements occur naturally in several structural forms.

Monatomic Elements

Some elements exist as individual atoms under normal conditions.

These are the:

Noble Gases

Examples:

• Helium (He)

• Neon (Ne)

• Argon (Ar)

• Krypton (Kr)

• Xenon (Xe)

• Radon (Rn)

Historically these were called:

"Inert gases"

because they were believed to form no compounds.

Today we know that the heavier noble gases can form compounds under special conditions.

Figure 10: Noble Gas Atoms

Illustrating the unique stability of noble gases, this chart shows how their full outer electron shells render them inert and unreactive, existing independently without forming chemical bonds.

Diatomic Elements

Seven elements naturally exist as molecules containing two atoms bonded together.

These are:

• H₂

• N₂

• O₂

• F₂

• Cl₂

• Br₂

• I₂

A useful mnemonic is: HONClBrIF (pronounced "hon-kl-brif")

When these elements appear in their elemental form, they must be written as diatomic molecules rather than individual atoms.

For example:

Correct:

O₂

Incorrect:

O

Figure 11: The Seven Diatomic Elements

Allotropes

Some elements exist in multiple structural forms.

These different forms are called: Allotropes

Allotropes contain the same element but differ in how the atoms are bonded and arranged.

Carbon Allotropes

Carbon provides one of the most dramatic examples.

Diamond

• Three-dimensional network structure

• Hardest naturally occurring substance

Graphite

• Layered structure

• Soft enough to write with

Buckminsterfullerene (C₆₀)

• Hollow spherical structure

• Often called a "buckyball"

All three substances are composed entirely of carbon atoms.

Figure 12: Carbon Allotropes

Explore the diverse forms of carbon allotropes: diamond's tetrahedral hardness, graphite's conductive layers, and the spherical structure of buckminsterfullerene, each with unique properties and applications.

Oxygen Allotropes

Oxygen also exists in multiple forms.

Oxygen Gas

O₂

The form required for respiration.

Ozone

O₃

A pale blue gas that absorbs harmful ultraviolet radiation in the upper atmosphere.

Figure 12 : Oxygen and Ozone

Explore the differences between oxygen (O₂) and ozone (O₃), both composed of oxygen atoms but with distinct structures and roles: O₂ is essential for life, while O₃ protects by absorbing UV radiation.

Additional Examples of Allotropes

Phosphorus

• White phosphorus

• Red phosphorus

Sulfur

• Rhombic sulfur

• Monoclinic sulfur

These allotropes possess different physical and chemical properties despite being composed of the same element.

Allotropes versus Isotopes

Students often confuse allotropes and isotopes.

The distinction is important.

Isotopes

Differ in:

• Number of neutrons

Examples:

• Carbon-12

• Carbon-13

• Carbon-14

Allotropes

Differ in:

• Atomic arrangement

• Bonding pattern

Examples:

• Diamond

• Graphite

• Buckminsterfullerene

The atoms themselves are identical in allotropes; only the structure changes.

Figure 14: Isotopes versus Allotropes

Explore the differences between isotopes and allotropes, focusing on carbon. Isotopes vary by neutron count, while allotropes have distinct atomic arrangements, affecting properties like hardness and conductivity.

Summary

Modern atomic structure consists of protons, neutrons, and electrons. Protons and neutrons occupy the nucleus and account for nearly all atomic mass, while electrons occupy the surrounding space and determine chemical behavior.

The atomic number identifies an element by specifying its number of protons, while the mass number gives the total number of protons and neutrons. Electrons may be gained or lost to form ions, producing charged species called cations and anions.

Elements exist in nature in several forms, including monatomic noble gases, diatomic molecules, and allotropes. Understanding these concepts provides the foundation for interpreting atomic notation, periodic trends, bonding, and chemical reactions.

Key Points

• Atoms contain protons, neutrons, and electrons.

• Protons have a +1 charge and reside in the nucleus.

• Neutrons are neutral and reside in the nucleus.

• Electrons have a −1 charge and occupy the space outside the nucleus.

• The atomic number equals the number of protons.

• The mass number equals protons plus neutrons.

• Neutral atoms contain equal numbers of protons and electrons.

• Cations form by losing electrons.

• Anions form by gaining electrons.

• Noble gases exist naturally as monatomic species.

• Seven elements exist naturally as diatomic molecules.

• Allotropes are different structural forms of the same element.

• Isotopes differ in neutron number, while allotropes differ in atomic arrangement.