Atomic Structure

Table of Contents

Overview: What Is Atomic Structure About?

Atomic structure describes how matter is built from extremely small particles called atoms, and how these atoms themselves are structured internally. In this chapter, we focus on:

The internal components of atoms
How these components are arranged
How this structure explains key observable properties (like atomic mass and charge)
How atomic structure provides the basis for later models and the periodic table

Detailed historical models and modern quantum ideas will be treated in later chapters; here you need just enough structure to understand what an atom is and how its parts are organized.

The Atom as the Basic Unit of Matter

Atoms are the smallest units of an element that still retain the chemical properties of that element.

Key points:

Atoms are extremely small: typical diameters are about $10^{-10}\,\text{m}$.
Atoms are mostly empty space: almost all mass is concentrated in a tiny central nucleus.
Atoms of different elements differ mainly in the number of protons in their nuclei.

In everyday matter, atoms combine to form molecules, ions, and solids. The way they combine is strongly influenced by their internal structure.

Subatomic Particles: Building Blocks of the Atom

Atoms consist of three main types of subatomic particles:

Protons: positively charged particles in the nucleus
Neutrons: neutral particles in the nucleus
Electrons: negatively charged particles surrounding the nucleus

Properties of Protons, Neutrons, and Electrons

Approximate relative properties:

Charge

Proton: $+1$ (elementary charge $+e$)
Neutron: $0$
Electron: $-1$ (elementary charge $-e$)

Relative mass (in units of electron mass)

Proton: about $1836$ times heavier than an electron
Neutron: about $1839$ times heavier than an electron
Electron: lightest; taken as reference

Location

Protons and neutrons: in the nucleus
Electrons: in the electron cloud around the nucleus

Because proton and electron charges are equal in magnitude but opposite in sign, an atom with the same number of protons and electrons is electrically neutral.

The Atomic Nucleus

The nucleus is a tiny, dense, positively charged region at the center of the atom.

Contains protons and neutrons (together called nucleons).
Occupies an extremely small volume compared to the whole atom (roughly $10^{-15}\,\text{m}$ across).
Contains almost all the mass of the atom.

The proton number determines the chemical identity of the element, while the combination of protons and neutrons influences nuclear properties (e.g. stability). Detailed nuclear behavior is discussed in nuclear chemistry, so here we focus on how the nucleus determines basic atomic characteristics.

Atomic Number, Mass Number, and Isotopes

The composition of an atomic nucleus can be described with two key numbers.

Atomic Number $Z$

Symbol: $Z$
Definition: Number of protons in the nucleus
Determines:

The element (hydrogen, carbon, oxygen, etc.)
The nuclear positive charge: $+Ze$
The number of electrons in a neutral atom

Examples:

Hydrogen: $Z = 1$ (1 proton)
Carbon: $Z = 6$ (6 protons)
Oxygen: $Z = 8$ (8 protons)

If the number of protons changes, you have a different element.

Mass Number $A$

Symbol: $A$
Definition: Total number of protons and neutrons in the nucleus
$$A = Z + N$$
where $N$ is the number of neutrons.
Determines (approximately) the mass of a single atom.

Mass number is always an integer. It is not the same as the atomic mass listed in the periodic table (which is an average over isotopes and usually not an integer).

Neutron Number $N$

Symbol: $N$
Definition: Number of neutrons in the nucleus
$$N = A - Z$$

For a given element (fixed $Z$), $N$ can vary, leading to isotopes.

Isotopes: Same Element, Different Mass

Isotopes are atoms of the same element (same $Z$) but different numbers of neutrons (different $N$ and therefore different $A$).

Same:

Number of protons
Chemical behavior (to a very good approximation)

Different:

Mass number $A$
Nuclear properties (e.g. stability, radioactivity)

Notation:

General: $^{A}_{Z}\text{X}$ or simply $^{A}\text{X}$ when $Z$ is implied.
Example (carbon):

$^{12}\text{C}$: $Z = 6$, $N = 6$
$^{13}\text{C}$: $Z = 6$, $N = 7$
$^{14}\text{C}$: $Z = 6$, $N = 8$

Isotopes explain why the atomic masses in the periodic table are often not integers: they are weighted averages of the isotopic masses, depending on natural abundance.

Atomic Mass and the Atomic Mass Unit

Defining the Atomic Mass Unit (u)

Chemists use a relative mass scale based on the isotope carbon-12.

1 atomic mass unit (1 u) is defined as:
$\,\text{u} = \frac{1}{12} \text{ of the mass of a } ^{12}\text{C atom}$$
Numerical value:
$\,\text{u} \approx 1.6605 \times 10^{-27}\,\text{kg}$$

Approximate masses:

Proton: $\approx 1.0073\,\text{u}$
Neutron: $\approx 1.0087\,\text{u}$
Electron: $\approx 0.00055\,\text{u}$ (so small it is often neglected in atomic mass calculations)

Atomic Mass of a Single Atom

The mass number $A$ gives a good approximation to the atomic mass in u, but not exactly, because:

Proton and neutron masses are not exactly 1 u.
The nucleus has binding energy, which slightly reduces the total mass (mass defect).

Thus, atomic masses are not exactly integers, even for an individual isotope.

Example:

$^{12}\text{C}$: defined as exactly $12\,\text{u}$
$^{16}\text{O}$: about $15.9949\,\text{u}$ (not exactly $16\,\text{u}$)

Relative Atomic Mass $A_\text{r}$

The relative atomic mass of an element is a weighted average of the masses of all naturally occurring isotopes.

Definition:

Relative atomic mass $A_\text{r}(\text{X})$ is the average mass of atoms of element X, compared to $1/12$ of the mass of a carbon-12 atom (dimensionless number).

This is the value you typically see in the periodic table.

Example (simplified for illustration):

Suppose element X has two isotopes:

$^{10}\text{X}$: mass $= 10.0\,\text{u}$, abundance $= 20\%$
$^{11}\text{X}$: mass $= 11.0\,\text{u}$, abundance $= 80\%$

Then:
$$A_\text{r}(\text{X}) = 0.20 \cdot 10.0 + 0.80 \cdot 11.0 = 10.8$$

The atomic weight of an element in chemical tables is effectively this average value, used to calculate molar masses in stoichiometry.

The Electron Cloud: Electrons in Atoms (Qualitative View)

Electrons occupy the space around the nucleus. In this chapter, we limit ourselves to a qualitative description; more precise models appear in later sections.

Key ideas:

Electrons are attracted to the positively charged nucleus by electrostatic (Coulomb) forces.
Electrons do not simply spiral into the nucleus; they occupy specific energy levels.
These energy levels are often grouped into shells roughly labeled K, L, M, … or with principal quantum numbers $n = 1, 2, 3, \dots$ in more advanced treatments.
Electrons behave with both particle-like and wave-like properties. Quantum-mechanical details (orbitals, quantum numbers) are discussed later.

The arrangement of electrons in these shells (and subshells) is crucial:

It determines the outer electron configuration.
The outer (valence) electrons govern chemical bonding and reactivity.
Many periodic trends (reactivity, atomic radius, ionization energy) follow from this arrangement.

Neutral Atoms, Ions, and Charge Balance

The relationship between protons and electrons determines the net charge of an atom or ion.

Neutral Atoms

Number of electrons equals number of protons:
$$\text{Neutral atom: } \text{electrons} = Z$$
Total positive charge of protons is balanced by total negative charge of electrons.
Overall charge: $0$.

Example:

Carbon: $Z = 6$

6 protons ($+6e$)
6 electrons ($-6e$)
Net charge $= 0$

Ions: Charged Atoms or Groups of Atoms

When atoms gain or lose electrons, they become ions.

Cations: positively charged ions (fewer electrons than protons)

Formed by loss of one or more electrons.
Charge $= +(\text{protons} - \text{electrons})$.

Anions: negatively charged ions (more electrons than protons)

Formed by gain of one or more electrons.
Charge $= -(\text{electrons} - \text{protons})$.

Examples:

Sodium atom: $^{23}\text{Na}$, $Z = 11$

Neutral atom: 11 protons, 11 electrons
Sodium cation: $\text{Na}^+$ has 11 protons, 10 electrons

Chlorine atom: $^{35}\text{Cl}$, $Z = 17$

Neutral atom: 17 protons, 17 electrons
Chloride anion: $\text{Cl}^-$ has 17 protons, 18 electrons

The nucleus remains unchanged in ordinary chemical processes (protons and neutrons do not change); only the electron count changes. Changes in the nucleus belong to nuclear chemistry.