Atoms, Molecules, Ions, and Naming: Foundations of General Chemistry

Study Guide - Smart Notes

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Atoms and Models

Dalton's Atomic Theory

The concept of the atom as the fundamental building block of matter has evolved from ancient philosophy to modern science. Dalton's atomic theory, proposed in the early 1800s, laid the groundwork for our understanding of chemical reactions and the nature of elements.

All matter is composed of tiny, indivisible particles called atoms.
All atoms of a given element are identical in properties; atoms of different elements differ in properties.
Atoms cannot be created or destroyed in chemical reactions; they are simply rearranged.
Compounds are formed by the combination of atoms of different elements in fixed, whole-number ratios.

Dalton's theory explained the conservation of mass and the constant composition of compounds, though we now know atoms are divisible into subatomic particles.

Dalton's atomic theory applied to water formation

Development of Atomic Models

Key experiments and scientists contributed to the modern atomic model:

J.J. Thomson: Discovered the electron using cathode ray tubes, showing atoms contain negatively charged particles.
Robert Millikan: Measured the charge of the electron with the oil drop experiment, finding it to be C.
Ernest Rutherford: Discovered the nucleus by observing the deflection of alpha particles through gold foil, concluding atoms have a dense, positively charged center.

Law of Multiple Proportions analogy with nuts and bolts

Subatomic Particles

Protons, Neutrons, and Electrons

Atoms are composed of three main subatomic particles:

Proton: Positively charged, located in the nucleus, mass ≈ 1 amu.
Neutron: Neutral, located in the nucleus, mass ≈ 1 amu.
Electron: Negatively charged, found in the electron cloud, mass ≈ 0.00055 amu.

Analogy of proton/neutron and electron masses Table of subatomic particle properties

Atomic Number, Mass Number, and Isotopes

Atomic Number and Mass Number

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

Number of neutrons = Mass number - Atomic number

Calculating number of neutrons from mass and atomic number Atomic number, mass number, and atomic mass for carbon

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, and thus different mass numbers. They have identical chemical properties but different physical properties (such as mass).

Isotope cartoon

Atomic Mass and Average Atomic Mass

Calculating Average Atomic Mass

The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes, based on their relative abundances.

The formula is:

Chlorine isotopes and average atomic mass calculation Magnesium isotopes and average atomic mass calculation

The Periodic Table

Organization and Classification

The periodic table arranges elements by increasing atomic number in rows (periods) and columns (groups or families). Elements in the same group have similar chemical and physical properties.

Periodic table with labeled groups and families

Metals, Nonmetals, and Metalloids

Metals: Left side, shiny, good conductors, form cations.
Nonmetals: Right side, poor conductors, form anions.
Metalloids: Along the stair-step line, properties intermediate between metals and nonmetals.

Periodic table worksheet for classification

Chemical Formulas and Types of Compounds

Chemical, Molecular, Empirical, and Structural Formulas

Chemical formula: Shows the types and numbers of atoms in a compound (e.g., H2O).
Molecular formula: Actual number of atoms in a molecule (e.g., C2H4O2).
Empirical formula: Simplest whole-number ratio of atoms (e.g., CH2O for C2H4O2).
Structural formula: Shows how atoms are bonded (e.g., Lewis structures).

Examples of chemical and molecular formulas Lewis structure example

Diatomic Elements

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

Mnemonic for diatomic elements on the periodic table

Ionic vs. Molecular Compounds

Ionic compounds: Formed from metals and nonmetals, consist of cations and anions, represented by formula units (e.g., NaCl).
Molecular compounds: Formed from nonmetals, consist of molecules with covalent bonds (e.g., CO2).

Ionic lattice structure of NaCl

Cations and Anions

Formation and Prediction of Charges

Atoms become ions by gaining or losing electrons:

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

Chlorine atom gaining an electron to become chloride ion Sodium atom losing an electron to become sodium ion Magnesium atom losing two electrons to become magnesium ion

Periodic Table Patterns for Ion Charges

Main group elements form predictable charges based on their group number:

Group 1A: 1+ (e.g., Na+)
Group 2A: 2+ (e.g., Mg2+)
Group 7A: 1- (e.g., Cl-)
Group 6A: 2- (e.g., O2-)

Periodic table with common ion charges Periodic table with diagonal exceptions for ion charges

Naming Ions and Compounds

Monatomic and Polyatomic Ions

Monatomic ions: Single atom with a charge (e.g., Na+, Cl-).
Polyatomic ions: Group of atoms with a charge (e.g., NO3-, SO42-).

Table of common polyatomic ions

Naming Rules

Cations with one possible charge: Name of element + 'ion' (e.g., sodium ion).
Cations with multiple charges: Name (Roman numeral) ion (e.g., iron(III) ion) or Latin name with -ic/-ous endings (e.g., ferric/ferrous).
Anions: Root of element + '-ide' (e.g., chloride, oxide).
Polyatomic ions: Memorize names and formulas (e.g., nitrate, sulfate).

Formulas of Ionic Compounds

Writing Empirical Formulas

Ionic compounds are written so that the total positive and negative charges balance to zero. The simplest ratio of ions is used (empirical formula).

Example: Mg2+ and Cl- combine to form MgCl2.

Formation of MgCl2 from Mg and Cl ions

Swap and Drop Method

To determine the correct subscripts in an ionic compound, swap the magnitude of each ion's charge to become the subscript of the other ion, then reduce to the simplest ratio if necessary.

Naming Compounds

Ionic Compounds

Name the cation first, then the anion.
Use Roman numerals for transition metals with multiple charges.
For polyatomic ions, use the ion's name as is.

Acids

Binary acids: 'Hydro-' prefix + root + '-ic acid' (e.g., HCl is hydrochloric acid).
Oxyacids: Based on polyatomic ion name: '-ate' becomes '-ic acid', '-ite' becomes '-ous acid' (e.g., HNO3 is nitric acid, HNO2 is nitrous acid).

Naming acids: HCl(g) vs HCl(aq)

Molecular Compounds

Use prefixes to indicate the number of each atom (e.g., CO2 is carbon dioxide, N2O4 is dinitrogen tetroxide).
The second element ends in '-ide'.

Hydrates

Ionic compound name + prefix for number of water molecules + 'hydrate' (e.g., CuSO4·5H2O is copper(II) sulfate pentahydrate).

Hydrate vs. anhydrous compound

Summary Table: Subatomic Particles

Particle	Mass (g)	Mass (amu)	Charge (relative)	Charge (C)
Proton	1.67262 × 10-24	1.00727	+1	+1.60218 × 10-19
Neutron	1.67493 × 10-24	1.00866	0	0
Electron	0.00091 × 10-24	0.00055	-1	-1.60218 × 10-19

Additional info: Mastery of these foundational concepts is essential for understanding chemical reactions, stoichiometry, and the behavior of matter at the atomic and molecular level. Practice problems and memorization of ion names and charges are highly recommended for proficiency.