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

Atoms and Elements

Introduction to Atoms

Atoms are the fundamental building blocks of matter. Every substance, whether solid, liquid, or gas, is composed of atoms, which combine in various ways to form compounds and mixtures.

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

• Element: A pure substance that cannot be broken down into simpler substances by chemical means.

• Example: Typical seaside rocks are composed of silicates (compounds of silicon and oxygen atoms). Seaside air contains nitrogen and oxygen molecules, and may also contain amines such as trimethylamine, which is emitted by decaying fish and contributes to the fishy smell of the seaside.

The Scale of Atoms

Atoms are incredibly small, and their sheer number in even tiny objects is immense.

• A single pebble contains more atoms than can be counted, far exceeding the number of pebbles on the bottom of San Francisco Bay.

• If every atom in a pebble were the size of the pebble itself, the pebble would be larger than Mount Everest.

Atomic Theory: Historical Development

Democritus and Dalton

The concept of atoms dates back to ancient Greece, but modern atomic theory was developed in the 19th century.

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

• John Dalton (1808): Formulated atomic theory with three main points:

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

  2. All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.

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

Experimental Evidence for Atomic Theory

Modern technology allows scientists to observe and manipulate individual atoms.

• Scanning tunneling microscopes (STM) can move atoms to form patterns, such as spelling out letters.

Structure of the Atom

Discovery of Subatomic Particles

Atoms are not indivisible; they are composed of smaller particles.

• Electron: Discovered by J.J. Thomson; negatively charged and much smaller than atoms.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle found in the nucleus.

Models of the Atom

• Plum Pudding Model: Proposed by Thomson; electrons are embedded in a sphere of positive charge.

• Nuclear Model (Rutherford): Most of the atom's mass and positive charge are concentrated in a small nucleus, with electrons dispersed in the surrounding space.

Distribution of Mass and Charge

The nucleus contains nearly all the mass of the atom, while electrons occupy most of the volume.

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

• Atoms are electrically neutral when the number of protons equals the number of electrons.

Properties of Subatomic Particles

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-

The Periodic Table and Classification of Elements

Atomic Number and Element Identity

The number of protons in the nucleus (atomic number, Z) defines the element.

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

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

Origins of Element Names and Symbols

• Most symbols are based on English names; some use Latin or Greek roots (e.g., K for potassium from kalium, Na for sodium from natrium).

• Elements may be named after properties, countries, or scientists.

Periodic Law and Patterns

When elements are arranged by increasing atomic mass, certain properties recur periodically.

• Elements with similar properties align in vertical columns (groups or families).

Classification: Metals, Nonmetals, Metalloids

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

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

• Metalloids: Along the zigzag line; mixed properties, semiconductors (e.g., silicon, arsenic, germanium).

Main Group and Transition Elements

• Main Group Elements: Properties are predictable based on position.

• Transition Elements: Properties are less predictable.

Special Groups in the Periodic Table

• Alkali Metals (Group 1A): Lithium, sodium, potassium, etc.

• Alkaline Earth Metals (Group 2A): Beryllium, magnesium, calcium, etc.

• Halogens (Group 7A): Fluorine, chlorine, bromine, iodine, astatine.

• Noble Gases (Group 8A): Helium, neon, argon, krypton, xenon.

Ions and Isotopes

Formation of Ions

Atoms can lose or gain electrons to form ions during chemical reactions.

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

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

• Ion Charge Formula: $\text{Charge} = \text{Number of protons} - \text{Number of electrons}$

• Example: Lithium loses one electron: $\text{Li} \rightarrow \text{Li}^+ + e^-$ For Li+ (3 protons, 2 electrons): $3 - 2 = 1^+$

• Fluorine gains one electron: $\text{F} + e^- \rightarrow \text{F}^-$ For F- (9 protons, 10 electrons): $9 - 10 = 1^-$

Predicting Ion Charges

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

• The group number (1A–8A) indicates the number of valence electrons.

Isotopes

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

• Isotope: Atoms with the same number of protons but different numbers of neutrons.

• Mass Number (A): Sum of protons and neutrons: $A = \text{Number of protons} + \text{Number of neutrons}$

• Isotope Symbol Notation: $^{A}_{Z}\text{X}$, where X is the chemical symbol, A is the mass number, Z is the atomic number.

• Example: Neon isotopes: Ne-20, Ne-21, Ne-22.

Atomic Mass and Natural Abundance

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 (Gallium): Ga-69: 68.9256 amu, 60.11% abundance (0.6011) Ga-71: 70.9247 amu, 39.89% abundance (0.3989) $\text{Atomic mass} = (0.6011 \times 68.9256) + (0.3989 \times 70.9247) = 41.4312 + 28.2919 = 69.7231 \approx 69.72$ amu

Radioactive Isotopes

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

• Radioactive Isotope: An isotope that emits subatomic particles from its nucleus.

• Nuclear radiation can be harmful, but some radioactive isotopes have beneficial uses (e.g., Tc-99 in medical imaging).

Chapter Review and Learning Objectives

• All matter is composed of atoms.

• Experiments by Thomson and Rutherford led to the nuclear theory of the atom.

• Electrons, protons, and neutrons have distinct properties and charges.

• The periodic table organizes elements by atomic number and groups with similar properties.

• Atoms can form ions by gaining or losing electrons; ions are classified as cations or anions.

• Isotopes are atoms of the same element with different numbers of neutrons; atomic mass is a weighted average of isotopic masses.