Atoms, Molecules, and Ions: Foundations of Atomic Theory and Chemical Formulas

Chapter 2: Atoms, Molecules, and Ions

Chapter Outline

• Early Ideas in Atomic Theory

• Evolution of Atomic Theory

• Atomic Structure and Symbolism

• Chemical Formulas

Early Ideas in Atomic Theory

Historical Development of Atomic Theory

The concept of atoms originated with the Greek philosophers Leucippus and Democritus in the fifth century BC. They proposed that matter is composed of indivisible units called atomos, meaning "indivisible." Later, Aristotle and others suggested that matter consisted of combinations of four elements: fire, earth, air, and water. The modern atomic theory was introduced by John Dalton in 1807.

Dalton’s Atomic Theory

Dalton's atomic theory laid the foundation for our understanding of matter. It consists of five key postulates:

• Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.

• An element consists of only one type of atom, which has a mass characteristic of the element and is the same for all atoms of that element.

• Atoms of one element differ in properties from atoms of all other elements.

• A compound consists of atoms of two or more elements combined in a small, whole-number ratio. In a given compound, the number of atoms of each element is always present in the same ratio.

• Atoms are neither created nor destroyed during a chemical change, but instead rearrange to yield different types of matter.

Key Laws Explained by Dalton’s Theory

• Law of Conservation of Matter: Atoms are neither created nor destroyed in chemical reactions, so the total mass remains constant.

• Law of Definite Proportions: All samples of a pure compound contain the same elements in the same proportion by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

Evolution of Atomic Theory

Discovery of Subatomic Particles

Modern atomic theory evolved through key experiments that revealed the existence of subatomic particles:

• Electron: Discovered by J.J. Thomson using cathode ray tubes. He determined that electrons are negatively charged particles much lighter than atoms.

• Charge of Electron: Measured by Robert Millikan in the oil drop experiment, finding the fundamental charge to be $1.602 imes 10^{-19}$ C.

• Nucleus: Discovered by Ernest Rutherford in the gold foil experiment, showing that atoms have a small, dense, positively charged nucleus surrounded by electrons.

• Proton: Positively charged particle located in the nucleus.

• Neutron: Discovered by James Chadwick, neutrons are uncharged particles with a mass similar to protons, also found in the nucleus.

• Isotopes: Atoms of the same element with different masses due to varying numbers of neutrons.

Atomic Structure and Symbolism

Subatomic Particles and Atomic Dimensions

• Proton: Mass = 1.0073 amu, Charge = +1

• Neutron: Mass = 1.0087 amu, Charge = 0

• Electron: Mass = 0.00055 amu, Charge = -1

The nucleus contains most of the atom’s mass, while electrons occupy most of its volume. The diameter of an atom is about $10^{-10}$ m, and the nucleus is about $10^{-15}$ m.

Atomic Number, Mass Number, and Isotopes

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in an atom.

• Isotopes: Atoms of the same element with different numbers of neutrons.

For a neutral atom, the number of protons equals the number of electrons. If not, the atom is an ion:

• Cation: Atom loses electrons, becomes positively charged.

• Anion: Atom gains electrons, becomes negatively charged.

Chemical Symbols and Isotope Notation

• Chemical Symbol: Abbreviation for an element (e.g., Hg for mercury).

• Isotope Notation: Mass number as superscript, atomic number as subscript to the left of the element symbol (e.g., $^{24}_{12} ext{Mg}$).

Chemical Formulas

Molecular and Empirical Formulas

• Molecular Formula: Shows the actual number of atoms of each element in a molecule (e.g., $C_6H_6$ for benzene).

• Empirical Formula: Shows the simplest whole-number ratio of atoms (e.g., $CH$ for benzene).

• Structural Formula: Shows how atoms are connected in a molecule.

Elements That Exist as Molecules

• Some elements exist as discrete molecules: $H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $Br_2$, $I_2$.

• Sulfur commonly exists as $S_8$.

Isomers

• Isomers: Compounds with the same chemical formula but different molecular structures and properties (e.g., acetic acid and methyl formate both have $C_2H_4O_2$).

The Mole and Avogadro’s Number

Definition and Use of the Mole

• Mole: The amount of substance containing $6.022 imes 10^{23}$ entities (atoms, molecules, or ions).

• Avogadro’s Number ($N_A$): $6.022 imes 10^{23}$

Molar Mass

• Molar Mass: Mass in grams of one mole of a substance (g/mol). Numerically equivalent to atomic or formula mass in amu.

Calculations Involving Moles, Mass, and Number of Particles

• To convert between mass, moles, and number of particles, use the following relationships:

Key Equations:

• $ ext{Number of moles} = rac{ ext{mass (g)}}{ ext{molar mass (g/mol)}}$

• $ ext{Number of particles} = ext{number of moles} imes N_A$

• $ ext{Mass} = ext{number of moles} imes ext{molar mass}$

Example Calculation

Given 47 g of potassium ($K$), molar mass = 39.10 g/mol:

• $ ext{Number of moles} = rac{47 ext{ g}}{39.10 ext{ g/mol}} = 1.20 ext{ mol}$

Summary Table: Properties of Subatomic Particles

Particle	Mass (amu)	Charge
Proton	1.0073	+1
Neutron	1.0087	0
Electron	0.00055	-1

Summary Table: Common Elements and Symbols

Element	Symbol
Aluminum	Al
Iron	Fe
Carbon	C
Oxygen	O
Hydrogen	H
Nitrogen	N

Key Concepts for Exam Preparation

• Understand Dalton’s atomic theory and its implications for chemical laws.

• Be able to identify and describe subatomic particles and their properties.

• Use atomic number, mass number, and isotope notation correctly.

• Distinguish between molecular, empirical, and structural formulas.

• Perform calculations involving moles, mass, and number of particles using Avogadro’s number and molar mass.

Additional info: The study notes expand on brief slide points to provide full academic context, definitions, and examples suitable for college-level General Chemistry students.