Atoms, Molecules, and Ions: Foundations of Chemical Structure and Nomenclature

• Atoms, Molecules, and Ions

The Atomic Theory of Matter

The atomic theory of matter provides the foundation for understanding chemical structure and reactions. Early philosophers speculated about the existence of atoms, but scientific evidence emerged in the 18th and 19th centuries.

• Law of Constant Composition: Compounds always have the same proportion of elements by mass. Discovered by Joseph Proust.

• Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. Discovered by Antoine Lavoisier.

• 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. Discovered by John Dalton.

Dalton's Atomic Theory

John Dalton proposed a theory in the early 1800s to explain the laws of chemical combination.

1. All matter is composed of extremely small particles called atoms.

2. All atoms of a given element are identical in mass and properties, but different from atoms of other elements.

3. Atoms of different elements combine in simple whole-number ratios to form compounds.

4. Atoms are not created or destroyed in chemical reactions; they are rearranged.

Discovery of Subatomic Particles

Atoms are not indivisible; they are composed of smaller subatomic particles:

• Electrons (discovered by J.J. Thomson via cathode ray experiments)

• Protons (discovered by Ernest Rutherford)

• Neutrons (discovered by James Chadwick)

Radioactivity, discovered by Henri Becquerel and further studied by Marie Curie, revealed that atoms can emit high-energy radiation.

Cathode Rays and Electrons

J.J. Thomson's experiments with cathode ray tubes led to the discovery of the electron:

• Cathode rays are streams of negatively charged particles (electrons).

• Electrons have a negative charge and a very small mass.

• Thomson measured the charge-to-mass ratio of the electron.

Millikan's Oil Drop Experiment

Robert Millikan measured the charge of the electron using the oil drop experiment:

• Charge of electron: $1.602 \times 10^{-19}$ C

• Mass of electron: $9.10 \times 10^{-28}$ g

Radioactivity

Radioactivity is the spontaneous emission of radiation by atoms. Three types were identified by Ernest Rutherford:

• Alpha (α) particles: Positively charged, heavy

• Beta (β) particles: Negatively charged, light

• Gamma (γ) rays: Uncharged, high-energy radiation

The Nuclear Model of the Atom

Rutherford's gold foil experiment led to the nuclear model:

• Atoms have a small, dense, positively charged nucleus containing protons and neutrons.

• Electrons occupy the space around the nucleus.

• Most of the atom's volume is empty space.

Subatomic Particles

Particle	Charge	Mass (amu)	Relative Mass	Location
Proton	+1	1.0073	1	Nucleus
Neutron	0	1.0087	1	Nucleus
Electron	-1	0.00055	~0	Outside nucleus

Atomic Number and Mass Number

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

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Isotopes are atoms of the same element with different numbers of neutrons.

Isotopes

Isotopes have the same atomic number but different mass numbers.

Symbol	Number of Protons	Number of Electrons	Number of Neutrons
12C	6	6	6
13C	6	6	7
14C	6	6	8

Atomic Mass Unit (amu)

• 1 amu = $1.66054 \times 10^{-24}$ g

• Defined as 1/12 the mass of a $^{12}C$ atom

Atomic Weight

The atomic weight is the weighted average mass of all isotopes of an element.

• Calculated as: $\text{atomic weight} = \sum (\text{isotope mass}) \times (\text{fractional isotope abundance})$

The Mass Spectrometer

Mass spectrometry is used to determine atomic and molecular weights by measuring the mass-to-charge ratio of ions.

The Periodic Table

Structure and Organization

• Elements are arranged in order of increasing atomic number.

• Rows are called periods; columns are called groups or families.

• Elements in the same group have similar chemical properties.

Periodicity

Periodic trends are observed in the properties of elements, such as atomic radius, ionization energy, and electronegativity.

Groups

Group	Old System	New System	Examples
1A	IA	1	Li, Na, K
2A	IIA	2	Be, Mg, Ca
7A	VIIA	17	F, Cl, Br
8A	VIIIA	18	He, Ne, Ar

Metals, Nonmetals, and Metalloids

• Metals: Shiny, good conductors, mostly solids.

• Nonmetals: Poor conductors, can be solid, liquid, or gas.

• Metalloids: Properties intermediate between metals and nonmetals; found along the staircase line.

Chemical Formulas and Molecules

Chemical Formulas

• Show the types and numbers of atoms in a compound.

• Empirical formula: Simplest whole-number ratio of atoms.

• Molecular formula: Actual number of atoms of each element in a molecule.

• Structural formula: Shows how atoms are connected.

Diatomic Molecules

Seven elements exist naturally as diatomic molecules:

• Hydrogen (H2)

• Nitrogen (N2)

• Oxygen (O2)

• Fluorine (F2)

• Chlorine (Cl2)

• Bromine (Br2)

• Iodine (I2)

Ions and Ionic Compounds

• Cations: Positively charged ions (loss of electrons), usually formed by metals.

• Anions: Negatively charged ions (gain of electrons), usually formed by nonmetals.

• Polyatomic ions: Groups of atoms with a net charge (e.g., NH4+, SO42−).

Common Ions

Formula	Name
Na+	Sodium ion
Ca2+	Calcium ion
Fe3+	Iron(III) ion
Cl−	Chloride ion
SO42−	Sulfate ion

Writing Formulas for Ionic Compounds

• Combine cations and anions in ratios that yield a neutral compound.

• Empirical formulas are used (lowest whole-number ratio).

• Example: $\text{Mg}^{2+}$ and $\text{N}^{3-}$ combine to form $\text{Mg}_3\text{N}_2$.

Chemical Nomenclature

Inorganic Nomenclature – Ionic Compounds

• Name the cation first, then the anion.

• If the cation can have more than one charge, indicate with Roman numerals (e.g., iron(II) chloride).

• Anions ending in -ide are monatomic; polyatomic ions retain their names.

Oxyanion Nomenclature

• Oxyanions are anions containing oxygen.

• Fewer oxygens: -ite; more oxygens: -ate.

• Prefixes hypo- (least) and per- (most) are used when there are four oxyanions in a series.

• Examples: NO2− (nitrite), NO3− (nitrate), SO32− (sulfite), SO42− (sulfate)

Acids and Acid Nomenclature

• Acids yield H+ ions in water.

• If the anion ends in -ide, the acid name begins with hydro- and ends with -ic (e.g., HCl: hydrochloric acid).

• If the anion ends in -ate, the acid name ends with -ic (e.g., HNO3: nitric acid).

• If the anion ends in -ite, the acid name ends with -ous (e.g., HNO2: nitrous acid).

Binary Molecular Compound Nomenclature

• Composed of two nonmetals.

• Prefixes indicate the number of each atom (mono-, di-, tri-, etc.).

• The more electronegative element is named second and ends in -ide.

• Example: CO2 is carbon dioxide.

Prefix	Number
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6

Organic Compound Nomenclature

Alkanes

• Hydrocarbons with only single bonds.

• General formula: CnH2n+2

• Names end in -ane (e.g., methane, ethane, propane).

Alcohols

• Derived from alkanes by replacing a hydrogen with an -OH group.

• Names end in -ol (e.g., methanol, ethanol).

Additional info: These notes provide a comprehensive overview of atomic structure, periodicity, chemical formulas, and nomenclature, suitable for foundational study in General Chemistry.