Atoms, Molecules, and Ions: Structured Study Notes

Chapter 2: Atoms, Molecules, and Ions

Introduction

This chapter introduces the foundational concepts of atomic structure, isotopes, the periodic table, molecules, ions, and chemical nomenclature. Mastery of these topics is essential for understanding chemical reactions and properties of matter.

Symbols of Elements

Atomic Number and Mass Number

• Symbol: Elements are represented by one or two letters (e.g., C for carbon).

• Atomic Number (Z): The number of protons in the nucleus; unique for each element.

• Mass Number (A): The sum of protons and neutrons in the nucleus.

Example: $^{12}_{6} ext{C}$ has 6 protons and 6 neutrons.

Isotopes

Definition and Properties

• Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different masses.

• All isotopes of an element have the same number of protons and electrons.

Isotope	Protons	Neutrons	Electrons
$^{11} ext{C}$	6	5	6
$^{12} ext{C}$	6	6	6
$^{13} ext{C}$	6	7	6
$^{14} ext{C}$	6	8	6

Atomic Weights

Calculation and Significance

• Atomic Weight: The weighted average mass of an element's isotopes based on their natural abundance.

• Calculated using: $ext{Atomic weight} = extstyleigg( ext{isotope mass} imes ext{fractional abundance}igg)$ (summed over all isotopes)

Isotope	Abundance	amu
$^{12} ext{C}$	98.93%	12.00
$^{13} ext{C}$	1.07%	13.00335

Example: Carbon's atomic weight: $(12.00 imes 0.9893) + (13.00335 imes 0.0107) = 12.01$ amu

The Periodic Table

Organization and Group Names

• Elements are arranged by increasing atomic number.

• Rows are called periods; columns are groups.

• Elements in the same group have similar chemical properties.

Group	Name
1A	Alkali metals
2A	Alkaline earth metals
6A	Chalcogens
7A	Halogens
8A	Noble gases

Diatomic Molecules

Definition and Examples

• Seven elements exist naturally as diatomic molecules (two atoms per molecule): H2, N2, O2, F2, Cl2, Br2, I2.

• These molecules are important for understanding elemental forms and reactions.

Molecules & Molecular Compounds

Formulas and Types

• Molecular compounds: Composed of molecules, typically only nonmetals (NM+NM).

• Molecular formula: Shows the exact number of atoms of each element (e.g., $ext{C}_2 ext{H}_4$ for ethylene).

• Empirical formula: Shows the simplest whole-number ratio of atoms (e.g., $ext{CH}_2$ for ethylene).

Structural and Perspective Formulas

• Structural formulas: Indicate the order in which atoms are bonded.

• Perspective drawings: Show the three-dimensional arrangement of atoms.

Ions & Ionic Compounds

Formation and Types

• Ions: Atoms that have gained or lost electrons.

• Cations: Positively charged ions (formed by metals).

• Anions: Negatively charged ions (formed by nonmetals).

• Ionic compounds: Formed from cations and anions, typically between metals and nonmetals (M+NM).

Particle Count of Ions

Ion	Protons	Neutrons	Electrons
$^7 ext{Li}$ (neutral)	3	4	3
$^7 ext{Li}^+$ (cation)	3	4	2
$^{16} ext{O}$ (neutral)	8	8	8
$^{16} ext{O}^{2-}$ (anion)	8	8	10

Writing Formulas for Ionic Compounds

Rules and Examples

• Compounds are electrically neutral; the sum of charges must equal zero.

• Use the criss-cross method to balance charges and determine subscripts.

• If subscripts are not in the lowest whole-number ratio, divide by the greatest common factor.

Example: $ext{Mg}^{2+}$ and $ext{O}^{2-}$ combine to form $ext{MgO}$.

Monatomic Ions

Cations and Anions

• Cations: Named after the metal (e.g., lithium ion).

• Transition metals: Use Roman numerals to indicate charge (e.g., iron(III)).

• Anions: Named by changing the ending to -ide (e.g., oxide, sulfide).

Polyatomic Ions

Common Polyatomic Ions

• Polyatomic ions consist of two or more atoms bonded together, carrying a charge.

• Examples include ammonium ($ext{NH}_4^+$), sulfate ($ext{SO}_4^{2-}$), nitrate ($ext{NO}_3^-$), hydroxide ($ext{OH}^-$).

Ion	Formula	Charge
Ammonium	NH4+	+1
Sulfate	SO42-	-2
Nitrate	NO3-	-1
Hydroxide	OH-	-1

Naming Ionic Substances

Practice and Rules

• Combine cation and anion names (e.g., magnesium chlorate for $ext{Mg(ClO}_3)_2$).

• For transition metals, indicate charge with Roman numerals (e.g., iron(III) nitrate).

Acid Nomenclature

Rules for Naming Acids

• If the anion ends in -ide: add prefix hydro- and change ending to -ic acid (e.g., hydrochloric acid).

• If the anion ends in -ite: change ending to -ous acid (e.g., nitrous acid).

• If the anion ends in -ate: change ending to -ic acid (e.g., chloric acid).

Nomenclature of Binary Molecular Compounds

Prefixes and Rules

• Use prefixes to indicate the number of atoms: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.

• The less electronegative atom is listed first.

• The ending of the more electronegative element is changed to -ide.

Prefix	Number
Mono-	1
Di-	2
Tri-	3
Tetra-	4
Penta-	5
Hexa-	6
Hepta-	7
Octa-	8
Nona-	9
Deca-	10

Summary of Chemical Nomenclature

• Naming Ionic Compounds: Name cation, then anion (ending in -ide for elements, or use polyatomic ion name). Use Roman numerals for variable charge cations.

• Naming Molecular Compounds: Use prefixes for number of atoms, except mono- for the first element.

• Naming Acids:

  • -ide: hydro-...ic acid

  • -ite: ...ous acid

  • -ate: ...ic acid

Graphic Organizer

Classification of compounds based on the first element:

• Hydrogen: Acid (H + NM or H + NM + O)

• Metal: Ionic (cation + anion)

• Non-metal: Molecular (NM + NM)

Example: Calcium chloride (ionic), carbon dioxide (molecular), hydrochloric acid (acid)

Additional info: These notes expand on the original slides by providing definitions, examples, and structured tables for clarity and completeness.