Atoms and Elements: Foundations of Matter and the Periodic Table

Atoms and Elements

Atoms as the Foundation of Matter

Atoms are the basic building blocks of all matter. The properties of atoms determine the properties of the substances they compose. An atom is the smallest identifiable unit of an element, and an element is a substance that cannot be broken down into simpler substances.

• Atoms compose all matter, from rocks to air.

• There are about 91 naturally occurring elements, each with unique atoms.

• Scientists have created about 20 synthetic elements.

• The exact number of natural elements is debated due to rare occurrences.

• Example: Silicates in rocks are compounds of silicon and oxygen atoms; air contains nitrogen and oxygen molecules.

Atomic Theory: Historical Development

The concept of atoms dates back to ancient Greece, where Democritus proposed that matter is composed of tiny, indivisible particles called "atomos." In 1808, John Dalton formalized atomic theory, which became widely accepted.

• Democritus: Matter consists of indestructible particles (atoms).

• Dalton's Atomic Theory:

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

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

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

Modern Evidence for Atomic Theory

Advanced technology, such as the scanning tunneling microscope (STM), allows scientists to manipulate individual atoms, providing direct evidence for the existence of atoms.

• IBM scientists arranged xenon atoms to spell "IBM." Atoms are generally spherical, but appear cone-shaped due to instrumentation.

Structure of the Atom

Discovery of Subatomic Particles

Atoms are not indivisible; they contain smaller particles. J. J. Thomson discovered the electron, a negatively charged particle much smaller than the atom. He proposed the "plum pudding" model, where electrons are embedded in a sphere of positive charge.

• Electrons: Negatively charged, much lighter than atoms, present in all substances.

• Atoms must contain positive charge to balance electrons.

Rutherford's Gold Foil Experiment and Nuclear Model

Ernest Rutherford's gold foil experiment showed that most alpha particles passed through gold foil, but some were deflected, indicating a dense, positively charged nucleus.

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

• Electrons occupy most of the atom's volume, which is mostly empty space.

• The number of electrons equals the number of protons, making atoms electrically neutral.

Distribution of Mass and Charge in the Atom

The nucleus contains nearly all the atom's mass but occupies a tiny fraction of its volume. Electrons are distributed in a much larger region and have negligible mass.

• If matter were composed only of nuclei, it would be extremely dense (e.g., neutron stars).

• Protons are nearly 2000 times as massive as electrons.

Electrical Charge in Atoms

Protons and electrons possess electrical charge. Opposite charges attract, like charges repel, and positive and negative charges cancel each other, resulting in charge-neutral atoms.

• Proton: 1+ charge

• Electron: 1- charge

• Neutron: No charge

Properties of Subatomic Particles

Protons and neutrons have similar masses, while electrons are much lighter. The following table summarizes their 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-

The Periodic Table and Elements

Atomic Number and Element Identity

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

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

• Each element has a unique atomic number.

Origins of Element Names and Symbols

Element symbols are often derived from English or Latin names. Some elements are named after properties, countries, or scientists.

• Examples: Potassium (K, kalium), Sodium (Na, natrium), Lead (Pb, plumbum), Curium (named after Marie Curie).

• Bromine (from Greek bromos, meaning "stench") is a red-brown gas with a strong odor.

Periodic Law and Classification of Elements

Dmitri Mendeleev observed that when elements are arranged by increasing mass, their properties recur periodically. The periodic table organizes elements by atomic number, grouping those with similar properties.

• Rows (periods) and columns (groups/families) align elements with similar properties.

• Periodic law summarizes recurring properties but does not explain them (theories do).

Classification: Metals, Nonmetals, Metalloids

Elements are classified as metals, nonmetals, or metalloids based on their properties and position in the periodic table.

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

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

• Metalloids: Along zigzag line; intermediate conductivity, useful in electronics.

Main Group and Transition Elements

Main group elements have predictable properties based on their position. Transition elements have less predictable properties.

• Main group elements: Groups 1A–8A.

• Transition elements: Middle section of the periodic table.

Groups of Elements: Alkali Metals, Alkaline Earth Metals, Halogens, Noble Gases

Specific groups in the periodic table have characteristic properties.

• Alkali metals (1A): Highly reactive, lose one electron.

• Alkaline earth metals (2A): Reactive, lose two electrons.

• Halogens (7A): Highly reactive, gain one electron.

• Noble gases (8A): Inert, stable electron configuration.

Ions and Isotopes

Formation of Ions

Atoms can lose or gain electrons to form ions. Positive ions are called cations, and negative ions are called anions. The charge of an ion is determined by the difference between the number of protons and electrons.

• Ion charge formula:

$\text{Ion charge} = \text{number of protons} - \text{number of electrons}$ $= \#p^+ - \#e^-$

Predicting Ion Charges

The number of valence electrons (outermost electrons) determines the charge an atom will acquire when forming ions. Main-group elements tend to form ions with the same number of valence electrons as the nearest noble gas.

• Group number indicates typical ion charge.

Isotopes: Variations in Neutron Number

Atoms of the same element can have different numbers of neutrons, resulting in isotopes. Isotopes are characterized by their mass number (A), which is the sum of protons and neutrons.

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

• Percent natural abundance describes the proportion of each isotope in nature.

Isotope Symbol Notation

Isotopes are represented by their chemical symbol, atomic number, and mass number. Two notations are commonly used:

• Notation 1: $^A_ZX$ (A = mass number, Z = atomic number, X = chemical symbol)

• Notation 2: X-A (chemical symbol or name followed by mass number)

Atomic Mass: Weighted Average of Isotopes

The atomic mass listed in the periodic table is the weighted average of the masses of all naturally occurring isotopes of an element.

• Formula:

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

• Fractions are percent abundances converted to decimals.

• Example: Gallium atomic mass calculation using Ga-69 and Ga-71.

Radioactive Isotopes and Nuclear Radiation

Some isotopes are unstable and emit nuclear radiation, changing into different isotopes or elements. These radioactive isotopes can be harmful or beneficial, such as in medical imaging.

• Example: Technetium-99 (Tc-99) is used in disease diagnosis.

Summary 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.

• Ions form by gaining or losing electrons; isotopes differ by neutron number.

• Atomic mass is a weighted average of isotopic masses.

Learning Objectives:

• Recognize that all matter is composed of atoms.

• Explain how Thomson and Rutherford's experiments led to atomic theory.

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

• Determine atomic symbol and number using the periodic table.

• Classify elements by group.

• Determine ion charge from protons and electrons.

• Calculate atomic mass from percent abundances and isotopic masses.