CHM113, ch2. Atoms, Molecules, and Ions: Foundations of Chemical Structure and Nomenclature

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Atoms, Molecules, and Ions

Introduction

This chapter introduces the fundamental concepts of atomic theory, the structure of atoms, the periodic table, and the basics of chemical nomenclature. Understanding these principles is essential for studying chemical reactions and properties of matter.

Laws of Chemical Combination

Law of Conservation of Mass

The Law of Conservation of Mass states that the total mass of substances present before a chemical reaction is equal to the total mass after the reaction. This principle, discovered by Antoine Lavoisier, is foundational to chemical equations and stoichiometry.

Key Point: Mass is neither created nor destroyed in a chemical reaction.
Example: The combustion of butane: 58 g butane + 208 g oxygen → 176 g carbon dioxide + 90 g water (total mass = 266 g before and after).

Law of Conservation of Mass: Butane combustion mass balance

Law of Constant Composition

The Law of Constant Composition (Joseph Proust) states that a compound always contains the same proportion of elements by mass. For example, every molecule of carbon dioxide (CO2) contains two oxygen atoms for every carbon atom, with a mass ratio of 2.66:1 (O:C).

Key Point: The composition of a compound is always fixed.
Example: CO2 always has a mass ratio of O to C of 32:12.01 = 2.66.

Law of Multiple Proportions

The Law of Multiple Proportions (John Dalton) states that if two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are small whole numbers.

Key Point: Different compounds formed from the same elements have mass ratios that are simple multiples of each other.
Example: O:C in CO2 is 2.66, in CO is 1.33; 2.66/1.33 = 2/1.

Dalton’s Atomic Theory

Postulates of Dalton’s Atomic Theory

Dalton’s atomic theory provides the basis for understanding chemical reactions and the nature of matter:

Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical in mass and properties, but differ from atoms of other elements.
Atoms are not created or destroyed in chemical reactions.
Atoms of different elements combine in simple whole-number ratios to form compounds.

Dalton's Atomic Theory illustrated

Structure of the Atom

Subatomic Particles

Atoms are composed of three main subatomic particles:

Protons: Positively charged, mass ≈ 1 amu, located in the nucleus.
Neutrons: Neutral, mass ≈ 1 amu, located in the nucleus.
Electrons: Negatively charged, mass ≈ 1/1836 amu, move around the nucleus.

The Nuclear Atom

Ernest Rutherford’s experiments led to the nuclear model of the atom, where a dense, positively charged nucleus is surrounded by electrons.

Atoms are 1–5 Å (100–500 pm) in diameter; the nucleus is much smaller (~10–4 Å).

Nuclear atom structure

Atomic Number, Mass Number, and Isotopes

Atomic Number and Mass Number

The atomic number (Z) is the number of protons in the nucleus and defines the element. The mass number (A) is the sum of protons and neutrons.

Atomic number is written as a subscript, mass number as a superscript before the element symbol.

Atomic number and mass number notation

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, and thus different mass numbers.

Isotopes have identical chemical properties but different physical properties (e.g., mass).
Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon.

Atomic Mass and Atomic Weight

Atomic Mass Unit (amu)

The atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 atom. 1 amu ≈ 1.66054 × 10–24 g.

Atomic Weight

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

Calculated as:
Example: Chlorine atomic weight calculation using 35Cl and 37Cl isotopes.

Atomic weight calculation for chlorine

The Periodic Table

Organization of the Periodic Table

The periodic table arranges elements in order of increasing atomic number. Rows are called periods, columns are groups or families. Elements in the same group often have similar chemical properties.

Periodic table of the elements

Reading the Periodic Table

Each box shows the atomic number (top), symbol (center), and atomic weight (bottom).

Periodic table element box

Classification of Elements

Metals: Left side, shiny, conduct heat/electricity, mostly solids.
Nonmetals: Right side, can be solid, liquid, or gas.
Metalloids: Along the staircase line, properties intermediate between metals and nonmetals.

Metals and nonmetals examples

Chemical Formulas and Models

Chemical Formulas

Chemical formulas indicate the types and numbers of atoms in a compound. Subscripts show the number of each atom.

Molecular formulas: Exact number of atoms (e.g., H2O).
Empirical formulas: Simplest whole-number ratio (e.g., CH2O for glucose).

Diatomic Molecules

Seven elements exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

Diatomic molecules and other molecular models

Structural and Molecular Models

Structural formulas show how atoms are connected; perspective, ball-and-stick, and space-filling models show three-dimensional arrangements.

Structural and molecular models of methane

Ions and Ionic Compounds

Formation of Ions

Atoms gain or lose electrons to form ions:

Cations: Positively charged, formed by loss of electrons (usually metals).
Anions: Negatively charged, formed by gain of electrons (usually nonmetals).

Writing Chemical Symbols for Ions

Symbol includes element, mass number (superscript), atomic number (subscript), and charge (superscript).
Example: for titanium ion with 22 protons, 26 neutrons, and 19 electrons.

Predicting Ionic Charges

Stable ions often have the same electron configuration as the nearest noble gas. Group 1 metals form +1 ions, group 2 metals form +2 ions, group 17 nonmetals form –1 ions, etc.

Periodic table with common ion charges

Common Ions

Some ions are so common they should be memorized, including polyatomic ions like NO3– (nitrate), SO42– (sulfate), and NH4+ (ammonium).

Ionic Compounds

Ionic compounds are formed from metals and nonmetals (or contain NH4+). The formula is written to ensure electrical neutrality.

Ionic compound lattice structure

Chemical Nomenclature

Naming Ionic Compounds

Name the cation first, then the anion.
For metals with variable charge, indicate the charge with Roman numerals (e.g., FeCl3 is iron(III) chloride).
Monatomic anions end in –ide (e.g., Cl– is chloride).
Polyatomic ions have specific names (e.g., NO3– is nitrate).

Naming Acids

If the anion ends in –ide, the acid name is hydro–[root]–ic acid (e.g., HCl is hydrochloric acid).
If the anion ends in –ite, the acid name is [root]–ous acid (e.g., HClO2 is chlorous acid).
If the anion ends in –ate, the acid name is [root]–ic acid (e.g., HNO3 is nitric acid).

Acid nomenclature chart

Naming Binary Molecular Compounds

The element farther left or lower in the periodic table is named first.
Prefixes indicate the number of atoms (mono-, di-, tri-, etc.).
The second element ends in –ide (e.g., CO2 is carbon dioxide).

Summary Table: Common Ions

Cation	Formula	Anion	Formula
Sodium	Na+	Chloride	Cl–
Calcium	Ca2+	Sulfate	SO42–
Ammonium	NH4+	Nitrate	NO3–
Iron(III)	Fe3+	Oxide	O2–

Conclusion

This chapter provides the essential foundation for understanding the structure of matter, the organization of elements, and the rules for naming chemical substances. Mastery of these concepts is critical for success in all areas of general chemistry.