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

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Atoms and Elements

Experiencing Atoms in Nature

Atoms are the fundamental building blocks of matter and are present in everything around us, from rocks to air and living organisms. Understanding atoms helps explain the composition and properties of substances in our environment.

Atoms are the foundation of our sensations and experiences.
Seaside rocks are typically composed of silicates, which are compounds containing silicon and oxygen atoms.
Seaside air contains molecules such as nitrogen (N2) and oxygen (O2).
Substances called amines (e.g., triethylamine) may be present in the air, contributing to characteristic smells such as the fishy odor near the seaside.
Example: Triethylamine is emitted by decaying fish and is responsible for the fishy smell at the seaside.

The Scale and Number of Atoms

Atoms are incredibly small, and even a tiny object contains an immense number of them. This concept illustrates the vastness of atomic scale compared to everyday objects.

A single pebble from the shoreline contains more atoms than can be counted.
The number of atoms in a pebble far exceeds the number of pebbles on the bottom of San Francisco Bay.
Example: If every atom in a pebble were the size of the pebble itself, the resulting object would be larger than Mount Everest.

Atoms: The Foundation of Matter

Atoms compose all matter, and their properties determine the characteristics of substances. The study of atoms is essential for understanding chemical behavior and the diversity of elements.

An atom is the smallest identifiable unit of an element.
An element is a substance that cannot be broken down into simpler substances by chemical means.
There are about 91 naturally occurring elements, each with unique types of atoms.
Scientists have created about 20 synthetic elements not found in nature.

Atomic Theory: Historical Perspectives

The concept of atoms has evolved over centuries, from philosophical ideas to scientific theories supported by experimental evidence.

Democritus (460–370 B.C.E.) proposed that matter is composed of tiny, indestructible particles called atomos (atoms).
John Dalton (1808) formalized atomic theory, stating:
- 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.

Modern Evidence for Atomic Theory

Advancements in technology have allowed scientists to observe and manipulate individual atoms, confirming their existence and properties.

Scanning tunneling microscopes (STM) can move and image atoms, such as arranging xenon atoms to spell out letters.
Atoms are generally spherical in shape.

Subatomic Particles: Electrons, Protons, and Neutrons

Atoms are composed of smaller particles: electrons, protons, and neutrons, each with distinct properties.

Electrons are negatively charged, much smaller and lighter than atoms, and are present in all substances.
Protons are positively charged and have a mass nearly 2000 times that of electrons.
Neutrons have no electrical charge and a mass similar to protons.

Atomic Models: Thomson and Rutherford

Scientific experiments led to the development of atomic models that describe the structure of atoms.

Thomson's Plum Pudding Model: Electrons are embedded in a sphere of positive charge.
Rutherford's Gold Foil Experiment: Most alpha particles passed through gold foil, but some were deflected, leading to the nuclear model.
Nuclear Model: Most of the atom's mass and all positive charge are concentrated in a small nucleus; electrons are dispersed in empty space around the nucleus.

Distribution of Mass and Charge in the Atom

The nucleus contains most of the atom's mass, while electrons occupy most of its volume.

The nucleus is extremely dense and small compared to the overall size of the atom.
Electrons are distributed throughout a much larger region but contribute little to the atom's mass.
Electrical charge is a fundamental property: opposite charges attract, like charges repel, and equal numbers of protons and electrons result in a neutral atom.

Properties of Subatomic Particles

Particle	Symbol	Charge	Mass (amu)
Proton	p+	+1	1
Neutron	n	0	1
Electron	e-	-1	~0.0005

Elements and Atomic Number

Each element is defined by its number of protons, known as the atomic number (Z).

If the number of protons changes, the atom becomes a different element.
The periodic table lists elements in order of increasing atomic number.

Origins of Element Names and Symbols

Element symbols are often derived from English, Latin, or Greek names, and some elements are named after countries or scientists.

Examples: Potassium (K, from Latin kalium), Sodium (Na, from Latin natrium).
Polonium (after Poland), Curium (after Marie Curie).

The Periodic Table: Organization and Classification

The periodic table organizes elements by atomic number and groups elements with similar properties into columns called groups or families.

Metals are on the left side, tend to lose electrons, and have properties such as conductivity, malleability, ductility, and luster.
Nonmetals are on the upper right, tend to gain electrons, and have varied properties (some are gases, some solids).
Metalloids lie along the zigzag line and have mixed properties, often acting as semiconductors.

Special Groups in the Periodic Table

Alkali metals (Group 1A): Highly reactive metals such as lithium, sodium, potassium.
Alkaline earth metals (Group 2A): Reactive metals such as magnesium, calcium.
Halogens (Group 7A): Reactive nonmetals such as fluorine, chlorine, bromine, iodine.
Noble gases (Group 8A): Inert gases such as helium, neon, argon.

Ions: Formation and Charge

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

Cations: Positively charged ions formed by losing electrons.
Anions: Negatively charged ions formed by gaining electrons.
The charge of an ion is calculated as:

Example (Cation):
Example (Anion):

Isotopes: Variations in Neutron Number

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

Mass number (A):
Isotope notation: , where X is the chemical symbol.
Example: Neon has three isotopes: Ne-20, Ne-21, Ne-22.

Atomic Mass: Weighted Average

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

Example (Gallium):
- Ga-69: 60.11% abundance, mass 68.9256 amu
- Ga-71: 39.89% abundance, mass 70.9247 amu
- Atomic mass: amu

Radioactive Isotopes and Nuclear Radiation

Some isotopes are unstable and emit subatomic particles, a process known as nuclear radiation. These isotopes are called radioactive.

Radioactive isotopes can be harmful due to their interaction with biological molecules.
Some isotopes emit radiation for a short time, while others remain radioactive for millions of years.
Beneficial uses: Technetium-99 is used in medical imaging to diagnose diseases.

Summary Table: Properties of Subatomic Particles

Particle	Charge	Mass (amu)
Proton	+1	1
Neutron	0	1
Electron	-1	~0.0005

Chemical Skills Learning Objectives

Recognize that all matter is composed of atoms.
Explain how experiments 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 predict ion charges.
Calculate atomic mass from percent natural abundances and isotopic masses.

Additional info: These notes are based on textbook slides and introductory chemistry content, suitable for college-level General Chemistry students.