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

Atoms, Molecules, and Ions

Dalton's Atomic Theory

Dalton's Atomic Theory laid the groundwork for our understanding of matter at the atomic level. It proposed that all matter is composed of tiny, indivisible particles called atoms, and described how atoms combine to form compounds.

• Atoms are the fundamental building blocks of matter, considered indivisible and indestructible in Dalton's time.

• All atoms of a given element are identical in mass and properties.

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

• Atoms can combine in more than one ratio to form different compounds (law of multiple proportions).

Note: Later discoveries showed that atoms are divisible (contain subatomic particles), so Dalton's first two postulates are no longer accepted. However, the remaining statements are still foundational in chemistry.

Thomson Model of the Atom

J.J. Thomson's experiments led to the discovery of subatomic particles, specifically electrons and protons, and the proposal of the "plum pudding" model of the atom.

• Electrons (e-): Negatively charged particles with a relative charge of -1.

• Protons (p+): Positively charged particles with a relative charge of +1.

• Thomson's plum pudding model depicted the atom as a sphere of positive charge with electrons embedded throughout, like plums in a pudding.

Rutherford Model of the Atom

Ernest Rutherford's gold foil experiment revolutionized atomic theory by demonstrating the existence of a dense, positively charged nucleus.

• Alpha particles were fired at thin gold foil; most passed through, but some were deflected, indicating a small, dense nucleus.

• The atom is mostly empty space, with positive charge concentrated in the nucleus.

• Neutrons (n0) were later discovered by James Chadwick; they are neutral particles with a mass similar to protons.

Subatomic Particles

Atoms are composed of three main subatomic particles: electrons, protons, and neutrons. Their properties are summarized below.

Particle	Symbol	Location	Relative Charge	Relative Mass
Electron	e-	Outside nucleus	-1	1/1836
Proton	p+	Inside nucleus	+1	1
Neutron	n0	Inside nucleus	0	1

Protons and Atomic Number

The number of protons in the nucleus of an atom determines its identity as an element and is called the atomic number (Z).

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

• For example, gold (Au) has 79 protons, so its atomic number is 79.

Atomic Number, Mass Number, and Isotopes

Atoms of the same element can have different numbers of neutrons, resulting in isotopes. The sum of protons and neutrons gives the mass number (A).

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

• Atomic number (Z): Number of protons.

• Mass number (A): Number of protons plus neutrons.

Isotope notation: $^A_ZX$ where X is the element symbol, A is the mass number, and Z is the atomic number.

Example: Isotopes Practice

• An isotope X contains 28 protons and 32 neutrons. Mass number = 28 + 32 = 60. The element with atomic number 28 is Nickel (Ni).

Ions

An ion is an atom or group of atoms with a net positive or negative charge, formed by gaining or losing electrons.

• Cation: Positive ion (lost electrons). Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$

• Anion: Negative ion (gained electrons). Example: $\text{O} + 2e^- \rightarrow \text{O}^{2-}$

For sodium (Na): p = 11, e = 11 (neutral); Na+: p = 11, e = 10. For oxygen (O): p = 8, e = 8 (neutral); O2-: p = 8, e = 10.

Periodic Table: Groups and Periods

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Group: Vertical column (also called family).

• Period: Horizontal row.

• There are 18 groups and 7 periods.

Valence Electrons

Valence electrons are the outermost electrons involved in chemical reactions. Their behavior determines the chemical properties of elements.

• Group 1 (IA) metals lose one valence electron to form 1+ ions: $\text{Na} \rightarrow \text{Na}^+ + e^-$

• Nonmetals gain electrons to form anions.

IUPAC Nomenclature and Classification of Compounds

The International Union of Pure and Applied Chemistry (IUPAC) sets rules for naming chemical compounds, ensuring consistency and clarity.

• Most inorganic compounds do not contain carbon (with few exceptions).

• Five common classes of inorganic compounds:

  1. Binary ionic

  2. Ternary ionic

  3. Binary molecular (covalent)

  4. Binary acid

  5. Ternary oxyacid

Binary Ionic Compounds

Binary ionic compounds consist of a metal and a nonmetal. The cation is named first, followed by the anion with the suffix "-ide".

• Example: $\text{KCl}$ is potassium chloride.

• Example: $\text{CaBr}_2$ is calcium bromide.

Binary Ionic Compounds (Transition Metals)

Transition metals can form cations with different charges, indicated by Roman numerals in the name.

• Example: $\text{CuF}$ is copper(I) fluoride.

• Example: $\text{Fe}_2\text{O}_3$ is iron(III) oxide.

Binary Covalent (Molecular) Compounds

Binary covalent compounds are formed between two nonmetals. Prefixes indicate the number of atoms of each element.

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

• Example: $\text{CO}$ is carbon monoxide.

• Example: $\text{CO}_2$ is carbon dioxide.

• Example: $\text{CCl}_4$ is carbon tetrachloride.

• Example: $\text{N}_2\text{O}_5$ is dinitrogen pentoxide.

Ternary Ionic Compounds

Ternary ionic compounds contain polyatomic ions. The cation is named first, followed by the polyatomic anion.

• Example: $\text{NaClO}_2$ is sodium chlorite.

• Example: $\text{CoCO}_3$ is cobalt(II) carbonate.

Common Polyatomic Ions

Name	Formula
Ammonium	NH4+
Hydroxide	OH-
Nitrate	NO3-
Sulfate	SO42-
Carbonate	CO32-
Permanganate	MnO4-
Phosphate	PO43-

Binary Acids

Binary acids are aqueous solutions containing hydrogen and a nonmetal. The name begins with "hydro-" and ends with "-ic acid".

• Example: $\text{HCl}$ (aq) is hydrochloric acid.

• Example: $\text{HF}$ (aq) is hydrofluoric acid.

• Example: $\text{HI}$ (aq) is hydroiodic acid.

Ternary Oxyacids

Ternary oxyacids are acids containing hydrogen, oxygen, and another element (usually a nonmetal). The name depends on the polyatomic ion present.

• If the anion ends in "-ate", the acid ends in "-ic acid". Example: $\text{HNO}_3$ (aq) is nitric acid (from nitrate ion).

• If the anion ends in "-ite", the acid ends in "-ous acid". Example: $\text{HNO}_2$ (aq) is nitrous acid (from nitrite ion).

• Example: $\text{H}_2\text{SO}_4$ (aq) is sulfuric acid (from sulfate ion).

Practice: Writing Formulas for Compounds

• Sodium nitride: $\text{Na}_3\text{N}$

• Nitrogen trihydride (ammonia): $\text{NH}_3$

• Zinc(IV) oxide: $\text{ZnO}_2$

• Ammonium sulfate: $(\text{NH}_4)_2\text{SO}_4$

• Calcium nitrate: $\text{Ca}(\text{NO}_3)_2$

• Aluminum carbonate: $\text{Al}_2(\text{CO}_3)_3$

• Hydrochloric acid: $\text{HCl}$ (aq)

Additional info: The notes cover foundational concepts in atomic theory, structure, and chemical nomenclature, suitable for introductory college-level General Chemistry.