2.3 Atomic Structure and Symbolism | Chemistry: Atoms First 2e

Learning Objectives

Write and interpret symbols that depict the atomic number, mass number, and charge of an atom or ion
Define the atomic mass unit and average atomic mass
Calculate average atomic mass and isotopic abundance

Core Concepts & Principles

Atoms consist of a tiny, dense nucleus containing positively charged protons and neutral neutrons, surrounded by a much larger cloud of negatively charged electrons. While the nucleus contains nearly all of the atom's mass, electrons occupy almost all of its volume.

To measure these subatomic scales, chemists use specialized units:

Unified atomic mass unit (u) or dalton (Da): Defined as exactly 1/12th the mass of a carbon-12 atom ( $1.6605 \times 10^{-24}\text{ g}$ ).
Fundamental unit of charge (e): Equal to the magnitude of an electron's charge ( $1.602 \times 10^{-19}\text{ C}$ ).

Key Atomic Identifiers

Atomic Number ( $Z$ ): The number of protons in the nucleus; it uniquely identifies the element. In a neutral atom, $Z$ also equals the number of electrons.
Mass Number ( $A$ ): The total count of protons and neutrons in an atom ( $A = \text{protons} + \text{neutrons}$ ).
Isotopes: Atoms of the same element (same $Z$ ) with different numbers of neutrons (different $A$ ).

Ions and Charges

When atoms gain or lose electrons, they become electrically charged particles called ions.

Cations: Positively charged ions formed by losing electrons ( $\text{protons} > \text{electrons}$ ).
Anions: Negatively charged ions formed by gaining electrons ( $\text{electrons} > \text{protons}$ ).

Average Atomic Mass

Because elements exist as mixtures of isotopes in nature, the atomic mass listed on the periodic table is a weighted average of all natural isotopes based on their isotopic abundances.

Problem-Solving Routines & Methods

Calculating Average Atomic Mass

1
Convert each isotope's percent abundance into a decimal (fractional abundance) by dividing by 100.
2
Multiply each isotope's fractional abundance by its respective isotopic mass.
3
Sum the products of all isotopes to find the average atomic mass.

Pro-Tip: Check that your fractional abundances add up to exactly 1.00 before multiplying.

Average Atomic Mass

\text{Average Atomic Mass} = \sum (f_i \times m_i)

Calculates the weighted average mass of an element's naturally occurring isotopes.

Variables & Constants

f_i

=fractional abundance of isotope

i

;

m_i

=isotopic mass of isotope

i

Neutron Calculation

A - Z = \text{number of neutrons}

Finds the number of neutrons in a specific nucleus.

Variables & Constants

A

=mass number (protons + neutrons);

Z

=atomic number (protons)

Practice & Concept Checks

Concept Check

An ion of platinum has a mass number of 195 and contains 74 electrons with a 4+ charge. How many protons and neutrons does it contain?

Concept Check

Why are the atomic masses on the periodic table not whole numbers if protons and neutrons each weigh approximately 1 amu?

Key Terms & Vocabulary

unified atomic mass unit (u)Measurement

A unit of mass defined as exactly 1/12 of the mass of a single carbon-12 atom, equal to 1.6605 × 10⁻²⁴ g.

Example: 1 u ≈ mass of one proton or neutron

fundamental unit of charge (e)Subatomic Properties

The elementary unit equal to the magnitude of the charge of a single electron, equal to 1.602 × 10⁻¹⁹ C.

Example: Electron charge is 1-, proton charge is 1+

dalton (Da)Measurement

An alternative mass unit that is numerically equivalent to the unified atomic mass unit (u).

Example: 1 Da = 1 u

atomic number (Z)Atomic Structure

The number of protons in the nucleus of an atom, which determines the chemical identity of the element.

Example: Carbon always has Z = 6

mass number (A)Atomic Structure

The total number of protons and neutrons in the nucleus of an atom.

Example: Carbon-12 has a mass number of 12

ionCharged Species

An electrically charged atom or molecule formed by gaining or losing one or more electrons.

Example: Na⁺ or Cl⁻

anionCharged Species

A negatively charged ion formed when an atom gains one or more electrons.

Example: O²⁻ or Cl⁻

cationsCharged Species

Positively charged ions formed when an atom loses one or more electrons.

Example: Na⁺ or Mg²⁺

chemical symbolNotation

A one-, two-, or three-letter abbreviation used to represent an element or an atom of an element.

Example: C for carbon, Hg for mercury

atomic massMeasurement

The weighted average mass of all naturally occurring isotopes of an element, expressed in atomic mass units (amu).

Example: Boron has an atomic mass of 10.81 amu