Atoms and Elements: Foundations of Modern Chemistry

Atoms and Elements

Imaging and Moving Individual Atoms

Atoms are the fundamental building blocks of matter, connecting the macroscopic and microscopic worlds. An atom is the smallest identifiable unit of an element, and there are about 91 naturally occurring elements, with over 20 synthetic elements created by scientists.

Early Ideas About the Building Blocks of Matter

Historical Perspectives

Leucippus and Democritus first proposed that matter was composed of small, indestructible particles called atoms. They suggested that different kinds of atoms existed, each with unique shapes and sizes, moving randomly through empty space. Plato and Aristotle, however, believed matter had no smallest parts and was composed of fire, air, earth, and water. The scientific approach later prevailed, with John Dalton providing evidence for the atomic theory.

Modern Atomic Theory and the Laws That Led to It

Key Laws

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction. The total mass of substances remains unchanged.

• Law of Definite Proportions: All samples of a given compound have the same proportions of their constituent elements.

• Law of Multiple Proportions: When two elements form different compounds, the masses of one element that combine with 1 gram of the other can be expressed as a ratio of small whole numbers.

Example: Law of Definite Proportions

Two samples of carbon dioxide decomposed yield the same oxygen-to-carbon mass ratio, demonstrating the law of definite proportions:

• Sample 1: $\frac{25.6}{9.60} = 2.67$

• Sample 2: $\frac{21.6}{8.10} = 2.67$

Example: Law of Multiple Proportions

Carbon dioxide and carbon monoxide show different mass ratios of oxygen to carbon:

• CO2: 2.67 g O per 1 g C

• CO: 1.33 g O per 1 g C

Atomic Theory and Subatomic Particles

Dalton’s Atomic Theory

• 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.

• Atoms of one element cannot change into atoms of another element.

The Discovery of the Electron

J. J. Thomson’s cathode ray experiments led to the discovery of the electron, a negatively charged, low-mass particle present in all atoms.

Millikan’s Oil Drop Experiment

Robert Millikan determined the charge of a single electron using the oil drop experiment, finding it to be $-1.60 \times 10^{-19}$ C.

Plum-Pudding Model

Thomson proposed that electrons were embedded in a positively charged sphere, known as the plum-pudding model.

Rutherford’s Gold Foil Experiment

Ernest Rutherford’s experiment showed that most of the atom is empty space, with a dense nucleus containing protons. This led to the nuclear model of the atom.

Subatomic Particles

Atoms are composed of protons, neutrons, and electrons. Protons and neutrons have nearly identical masses, while electrons are much lighter. The charge of the proton and electron are equal in magnitude but opposite in sign; neutrons have no charge.

Particle	Mass (kg)	Mass (amu)	Charge (relative)	Charge (C)
Proton	1.67262 × 10-27	1.00727	+1	+1.60218 × 10-19
Neutron	1.67493 × 10-27	1.00866	0	0
Electron	0.00091 × 10-27	0.00055	-1	-1.60218 × 10-19

Elements and the Periodic Table

Defining Elements by Number of Protons

The number of protons in an atom’s nucleus (atomic number, Z) defines the element. Each element has a unique atomic number and chemical symbol.

The Periodic Table

The periodic table organizes elements by increasing atomic number. Each element is represented by its atomic number, chemical symbol, and name.

Isotopes and Atomic Mass

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. The mass number (A) is the sum of protons and neutrons. Isotopes are represented as $^A_ZX$ or as X-A.

Symbol	Number of Protons	Number of Neutrons	Mass Number (A)	Natural Abundance (%)
Ne-20	10	10	20	90.48
Ne-21	10	11	21	0.27
Ne-22	10	12	22	9.25

Atomic Mass

Atomic mass is the weighted average mass of an element’s isotopes, based on their natural abundance. It is listed beneath the element’s symbol in the periodic table.

For example, chlorine’s atomic mass is calculated as:

• Cl-35: $0.7577 \times 34.97 = 26.4968$ amu

• Cl-37: $0.2423 \times 36.97 = 8.9578$ amu

• Atomic mass Cl = $26.4968 + 8.9578 = 35.45$ amu

Mass Spectrometry

Mass spectrometry is used to determine the masses and abundances of isotopes, producing characteristic patterns for each element.

The Periodic Law and Classification of Elements

The Periodic Law

When elements are arranged in order of increasing mass, certain sets of properties recur periodically. Mendeleev organized elements so that those with similar properties fall in the same columns.

Classification of Elements

Elements are classified as metals, nonmetals, or metalloids. Metals are good conductors, malleable, ductile, and tend to lose electrons. Nonmetals are poor conductors, not malleable or ductile, and tend to gain electrons. Metalloids have mixed properties and are often semiconductors.

Periodic Table Groups and Periods

The periodic table is divided into vertical columns (groups) and horizontal rows (periods). Main-group elements have predictable properties, while transition elements are less predictable.

Special Groups

• Noble Gases (Group 8A): Mostly unreactive, e.g., helium, neon, argon.

• Alkali Metals (Group 1A): Highly reactive metals, e.g., sodium, lithium.

• Alkaline Earth Metals (Group 2A): Fairly reactive, e.g., calcium, magnesium.

• Halogens (Group 7A): Very reactive nonmetals, e.g., fluorine, chlorine.

Ions and the Periodic Table

Formation of Ions

Atoms can lose or gain electrons to become ions. Cations are positively charged (e.g., Na+), and anions are negatively charged (e.g., F-). Main-group metals tend to lose electrons, forming cations; main-group nonmetals tend to gain electrons, forming anions.

Predicting Ion Charges

• Alkali metals (1A): Lose one electron, form 1+ ions.

• Alkaline earth metals (2A): Lose two electrons, form 2+ ions.

• Halogens (7A): Gain one electron, form 1- ions.

• Oxygen family (6A): Gain two electrons, form 2- ions.

Molar Mass and Counting Atoms

The Mole Concept

The mole (mol) is a unit for counting atoms, defined as $6.022 \times 10^{23}$ particles (Avogadro’s number). The mass of 1 mole of atoms (molar mass) in grams is numerically equal to the atomic mass in amu.

Converting Between Mass, Moles, and Number of Atoms

To count atoms by weighing, use the molar mass as a conversion factor between mass and moles, and Avogadro’s number to convert between moles and atoms.

• Number of moles = $\frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of atoms = number of moles $\times 6.022 \times 10^{23}$

Example Calculation

Calculate the number of copper atoms in 2.45 mol of copper:

• $2.45 \text{ mol} \times 6.022 \times 10^{23} = 1.47 \times 10^{24}$ atoms

Additional info: These notes cover the foundational concepts of atomic theory, laws of chemical combination, subatomic particles, isotopes, atomic mass, periodic table structure, ion formation, and the mole concept, all essential for general chemistry students.