Atoms and Elements: Foundations of Matter (Chapter 4 Study Notes)

Atoms and Elements

Introduction to Atoms and Matter

Atoms are the fundamental building blocks of matter. The properties of atoms determine the properties of all substances, from rocks to air to living organisms.

• Atom: The smallest identifiable unit of an element, retaining the chemical properties of that element.

• Element: A substance that cannot be broken down into simpler substances by chemical means. Each element is defined by its number of protons.

• There are about 91 naturally occurring elements, with additional synthetic elements created in laboratories.

• Examples: Seaside rocks are composed of silicates (compounds of silicon and oxygen atoms); air contains nitrogen and oxygen molecules; decaying fish emit amines such as ethylamine.

Atomic Theory: Historical Development

The concept of atoms has evolved over centuries, from philosophical ideas to scientific theory.

• Democritus (460–370 BCE): Proposed that matter is composed of tiny, indivisible particles called atomos (meaning "indivisible").

• John Dalton (1808): Developed the first modern atomic theory, which states:

  • Each element is composed of tiny, indestructible particles called atoms.

  • All atoms of a given element have the same mass and properties.

  • Atoms combine in simple, whole-number ratios to form compounds.

Discovery of Subatomic Particles

Atoms are not indivisible; they are composed of smaller subatomic particles: electrons, protons, and neutrons.

• J. J. Thomson (1897): Discovered the electron, a negatively charged particle much smaller than the atom. Proposed the "plum pudding" model, where electrons are embedded in a sphere of positive charge.

• Ernest Rutherford (1909): Gold foil experiment showed that atoms have a small, dense, positively charged nucleus. Proposed the nuclear model of the atom:

  • Most of the atom's mass and all positive charge are in the nucleus.

  • Electrons are dispersed in the empty space around the nucleus.

  • The number of protons equals the number of electrons in a neutral atom.

Subatomic Particles: Properties and Comparison

Atoms are composed of three main subatomic particles, each with distinct properties.

Particle	Mass (kg)	Mass (amu)	Charge
Proton	1.67262 × 10−27	1.0073	+1
Neutron	1.67493 × 10−27	1.0087	0
Electron	0.00091 × 10−27	0.00055	−1

• Protons and neutrons have similar masses (~1 amu), while electrons are much lighter.

• Electrical charge is a fundamental property: like charges repel, opposite charges attract.

• Atoms are normally charge-neutral, with equal numbers of protons and electrons.

Atomic Number and the Periodic Table

The identity of an element is determined by its atomic number (Z), which is the number of protons in its nucleus.

• Atomic number (Z): Number of protons in the nucleus of an atom.

• The periodic table arranges elements in order of increasing atomic number.

• Each element has a unique name, symbol, and atomic number.

• Some symbols are derived from Latin or Greek names (e.g., Na for sodium from natrium).

Classification of Elements: Metals, Nonmetals, and Metalloids

Elements are broadly classified based on their properties and position in the periodic table.

• Metals: Left side of the table; good conductors, malleable, ductile, lustrous, tend to lose electrons.

• Nonmetals: Upper right side; varied properties, poor conductors, tend to gain electrons.

• Metalloids (Semimetals): Along the zigzag line; intermediate properties, semiconductors.

Groups and Periodic Trends

The periodic table is organized into groups (columns) and periods (rows), with elements in the same group sharing similar properties.

• Main group elements: Properties are predictable based on position.

• Transition elements: Properties are less predictable.

• Special groups:

  • Group 1A: Alkali metals

  • Group 2A: Alkaline earth metals

  • Group 7A: Halogens

  • Group 8A: Noble gases

Ions: Formation and Charge

Atoms can gain or lose electrons to form ions, which are charged particles.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• The charge of an ion is given by:

$\text{Ion charge} = \text{number of protons} - \text{number of electrons}$

• Main-group elements tend to form ions with the same number of valence electrons as the nearest noble gas.

• Example: $\text{Li} \rightarrow \text{Li}^+ + e^-$ (lithium loses one electron to form a cation)

• Example: $\text{F} + e^- \rightarrow \text{F}^-$ (fluorine gains one electron to form an anion)

Isotopes: Variations in Neutron Number

Atoms of the same element can have different numbers of neutrons, resulting in isotopes.

• Isotope: Atoms of the same element (same number of protons) with different numbers of neutrons.

• Mass number (A): Total number of protons and neutrons in the nucleus.

• Isotope notation: $^{A}_{Z}\text{X}$, where X is the element symbol, A is the mass number, and Z is the atomic number.

• Alternative notation: X-A (e.g., Ne-20, Ne-21, Ne-22 for neon isotopes).

Atomic Mass: Weighted Average of Isotopes

The atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

• Atomic mass calculation:

$\text{Atomic mass} = (\text{fraction of isotope 1} \times \text{mass of isotope 1}) + (\text{fraction of isotope 2} \times \text{mass of isotope 2}) + \ldots$

• Example: Chlorine has two isotopes, Cl-35 and Cl-37. The atomic mass is calculated using their masses and percent abundances.

• Sample calculation for gallium:

$\text{Atomic mass} = (0.6011 \times 68.9256\ \text{amu}) + (0.3989 \times 70.9247\ \text{amu}) = 69.72\ \text{amu}$

Radioactive Isotopes and Nuclear Radiation

Some isotopes are unstable and emit nuclear radiation, changing into different isotopes or elements.

• Radioactive isotope: An isotope whose nucleus is unstable and emits subatomic particles (nuclear radiation).

• Radioactive decay can be harmful to living organisms but also has beneficial uses (e.g., medical imaging with technetium-99).

Summary Table: Properties of Subatomic Particles

Particle	Symbol	Relative Mass (amu)	Charge
Proton	p+	1.0073	+1
Neutron	n0	1.0087	0
Electron	e−	0.00055	−1

Chemical Skills Learning Objectives

• Recognize that all matter is composed of atoms.

• Explain how the experiments of Thomson and Rutherford led to the nuclear theory of the atom.

• Describe the properties and charges of electrons, neutrons, and protons.

• Determine atomic symbols and numbers using the periodic table.

• Classify elements by group and type (metal, nonmetal, metalloid).

• Determine ion charge from numbers of protons and electrons.

• Determine the number of protons and electrons in an ion.

• Determine atomic numbers, mass numbers, and isotope symbols for isotopes.

• Calculate atomic mass from percent natural abundances and isotopic masses.