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

Chapter 2 – Atoms, Molecules, and Ions

Dalton Atomic Theory

The Dalton Atomic Theory, proposed in the early 19th century, laid the groundwork for our understanding of matter at the atomic level. It consists of several key postulates about the nature and behavior of atoms.

• Atoms as Fundamental Particles: All matter is composed of tiny, indivisible, indestructible particles called atoms.

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

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

• Multiple Proportions: Atoms can combine in more than one ratio to form different compounds.

Modern Revisions: The first two postulates were later found to be incorrect due to the discovery of subatomic particles and isotopes. However, the remaining statements are still accepted today.

Thomson Model of the Atom

Following Dalton, J.J. Thomson discovered the existence of subatomic particles, leading to a new model of the atom.

• Discovery of Electrons: Negatively charged electrons ($e^-$) were identified.

• Discovery of Protons: Positively charged protons ($p^+$) were identified.

• Relative Charges: Electron: $-1$; Proton: $+1$.

• Plum Pudding Model: Thomson proposed that electrons were embedded in a positively charged 'pudding,' like plums in a pudding.

Rutherford Model of the Atom

Ernest Rutherford's gold foil experiment led to a new understanding of atomic structure.

• Gold Foil Experiment: Alpha particles were fired at thin gold foil; most passed through, but some were deflected, indicating a dense, positively charged nucleus.

• Nucleus: The atom's positive charge is concentrated in the nucleus.

• Discovery of Neutrons: James Chadwick discovered neutrons ($n^0$), which are similar in mass to protons but have no charge.

Subatomic Particles

Atoms are composed of three fundamental subatomic particles, each with distinct properties.

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

Protons ($p^+$) are positively charged particles located in the nucleus. The number of protons in an atom determines its atomic number ($Z$), which uniquely identifies each element.

• 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.

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

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

• Isotopes: Atoms with the same atomic number but different mass numbers.

Notation:

• $^{A}_{Z}X$ where $A$ = mass number, $Z$ = atomic number, $X$ = element symbol.

Isotopes Practice Example

• Problem: An isotope X contains 28 protons and 32 neutrons. What is its mass number and element?

• Solution: $A = 28 + 32 = 60$; 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 due to loss or gain of electrons.

• Cation: Positively charged ion (lost electrons). Example: Na ($p=11$, $e=11$) → Na+ ($p=11$, $e=10$)

• Anion: Negatively charged ion (gained electrons). Example: O ($p=8$, $e=8$) → O2- ($p=8$, $e=10$)

Periodic Table: Groups and Periods

The Periodic Table organizes elements by increasing atomic number and recurring chemical properties.

• Group (Family): Vertical column; elements in a group share similar chemical properties.

• Period: Horizontal row; elements in a period have the same number of electron shells.

• There are 18 groups and 7 periods.

Valence Electrons

Valence electrons are the outermost electrons involved in chemical reactions. Their number determines an element's chemical reactivity.

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

Ionic Charges

Metals and nonmetals form ions by losing or gaining valence electrons.

• Metals: Lose electrons to form cations.

• Nonmetals: Gain electrons to form anions.

IUPAC Nomenclature

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).

Classification of Compounds

Inorganic compounds are classified into five common types:

• Binary ionic

• Ternary ionic

• Binary molecular (covalent)

• Binary acid

• Ternary oxyacid

Binary Ionic Compounds

Binary ionic compounds consist of a metal and a nonmetal.

• Formula Writing: Combine cation and anion in the simplest whole-number ratio. Example: $\text{Ca}^{2+} + 2\text{Br}^- \rightarrow \text{CaBr}_2$

• Naming: Name the cation first, then the anion with the suffix "-ide." Example: $\text{KCl}$ – potassium chloride

Binary Ionic Compounds (Transition Metals)

• Transition metals may have multiple charges; use Roman numerals to indicate charge. Example: $\text{CuF}$ – copper(I) fluoride; $\text{Fe}_2\text{O}_3$ – iron(III) oxide

Binary Covalent (Molecular) Compounds

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

• Prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10)

• Naming: First element is named first; second element uses the root plus "-ide" and a prefix. The prefix "mono-" is not used for the first element. Examples: carbon monoxide ($\text{CO}$), carbon dioxide ($\text{CO}_2$), dinitrogen pentoxide ($\text{N}_2\text{O}_5$)

Ternary Ionic Compounds

Ternary ionic compounds contain polyatomic ions.

• Formula Writing: Combine cation and polyatomic anion. Example: $\text{Ca}^{2+} + \text{PO}_4^{3-} \rightarrow \text{Ca}_3(\text{PO}_4)_2$

• Naming: Name the cation, then the polyatomic anion. Example: $\text{NaClO}_2$ – sodium chlorite; $\text{CoCO}_3$ – cobalt(II) carbonate

Common Polyatomic Ions

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

Binary Acids

Binary acids are aqueous solutions containing hydrogen and a nonmetal.

• Naming: Prefix "hydro-" + root of nonmetal + "-ic acid". Examples: $\text{HCl(aq)}$ – hydrochloric acid; $\text{HF(aq)}$ – hydrofluoric acid

Ternary Oxyacids

Ternary oxyacids are aqueous solutions containing hydrogen, oxygen, and another element (usually a polyatomic ion).

• Naming: If the polyatomic ion ends in "-ate," change to "-ic acid"; if "-ite," change to "-ous acid." Examples: $\text{HNO}_2$ – nitrous acid ($\text{NO}_2^-$ = nitrite); $\text{HNO}_3$ – nitric acid ($\text{NO}_3^-$ = nitrate); $\text{H}_2\text{SO}_4$ – sulfuric acid ($\text{SO}_4^{2-}$ = sulfate)

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: Some content was expanded for clarity and completeness, including standard chemical notation and additional examples for nomenclature.