Fundamental Concepts in General Chemistry: Atomic Theory, Atomic Structure, and the Periodic Table

Early Atomic Theory

John Dalton's Atomic Theory

John Dalton proposed foundational postulates that led to the development of modern atomic theory. These principles describe the nature and behavior of atoms in matter.

• Matter consists of tiny particles called atoms.

• In any sample of pure elements, all atoms are identical in mass and other properties.

• Atoms of different elements have different masses and properties.

• Atoms combine in fixed ratios to form compounds.

• Atoms are indivisible in chemical processes.

How Did Atomic Theory Evolve?

The Cathode Ray Experiment (J.J. Thomson)

J.J. Thomson used a cathode ray tube to measure the charge-to-mass ratio of electrons, leading to the discovery of the electron.

• Measured the most charge-to-mass ratio: Value: $-1.76 \times 10^8$ C/g

• Key Insight: The nature of the gas did not affect the result, indicating the electron is a universal component of all atoms.

Rutherford's Gold Foil Experiment

Rutherford revealed the atomic nucleus by observing the scattering of alpha particles.

• Key Insight: Most of the atom is empty space; the nucleus is small and dense.

Millikan's Oil Drop Experiment

Millikan determined the charge of the electron by measuring the force on oil droplets in an electric field.

• Value: $-1.60 \times 10^{-19}$ C per electron

• Key Insight: The charge of the electron is quantized and fundamental.

Atomic and Mass Number

Definitions

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

• Mass Number (A): Sum of protons and neutrons in the nucleus. $A = Z + \text{number of neutrons}$

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

• Mass of Carbon-12: Defined as exactly 12 amu.

Symbols for Elements

Element Symbols and Atomic Number

• Each element is represented by a unique symbol (first letter capitalized, second letter lowercase if present).

• Atomic number is often omitted in chemical formulas.

Atoms and Ions

• Atoms can lose or gain electrons to form ions.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Example: Sodium atom (Na) loses one electron to become Na+ (cation); chlorine atom (Cl) gains one electron to become Cl- (anion).

Atomic Numbers and Mass Numbers

Problems and Calculations

• Given an ion or atom, you can determine the number of protons, neutrons, and electrons using atomic and mass numbers.

• Example: For Cr-52, find the number of protons, neutrons, and electrons.

Chemical Formulas

Empirical and Molecular Formulas

Chemical formulas represent the composition of compounds.

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

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

• Example:

  • Benzene: Molecular Formula = C6H6, Empirical Formula = CH

  • Acetic Acid: Molecular Formula = C2H4O2, Empirical Formula = CH2O

Isotopes

Definition and Notation

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

• Mass number is written as a superscript to the left of the element symbol.

• Atomic number is written as a subscript to the left of the element symbol.

• Most elements exist as mixtures of isotopes.

Example: Magnesium Isotopes

• Magnesium exists as several isotopes, each with a different mass number.

Isotopic Abundance and "Weight Mass"

Calculating Average Atomic Mass

Isotopic abundance tells us the fraction of the total number of atoms identified with a particular isotope. The average atomic mass is calculated using the masses and abundances of each isotope.

• Formula: $\text{Average Atomic Mass} = \sum (\text{fraction of isotope} \times \text{isotopic mass})$

• Example: Boron $\text{Boron Average Mass} = (1.199 \times 10.0129\, \text{amu}) + (0.801 \times 11.0093\, \text{amu}) = 10.81\, \text{amu}$

Example: Calculating Average Atomic Mass from Isotopic Abundances

• Chlorine has two main isotopes: $^{35}$Cl and $^{37}$Cl.

• Given: $^{35}$Cl (mass = 34.969 amu, abundance = 75.77%), $^{37}$Cl (mass = 36.966 amu, abundance = 24.23%)

• Calculation: $\text{Average Atomic Mass} = (34.969 \times 0.7577) + (36.966 \times 0.2423) = 35.45\, \text{amu}$

The Periodic Table

History and Organization

• The periodic table was discovered by Dmitry Mendeleev and Lothar Meyer.

• It organizes elements based on properties and atomic number.

• Elements are arranged in rows called periods and columns called groups.

Distribution of Metals, Nonmetals, and Metalloids

Classification

Elements are classified based on their physical and chemical properties.

• Metals: Shiny, malleable, good conductors of heat and electricity.

• Nonmetals: Dull appearance, poor conductors of heat and electricity.

• Metalloids: Conduct heat and electricity moderately well; possess properties intermediate between metals and nonmetals.

Examples

• Silicon is a metalloid, used in electronics (e.g., silicon chips).

• Iron is a metal, used in construction and manufacturing.

• Sulfur is a nonmetal, used in chemical production.

Common Polyatomic Ions

Table: Common Polyatomic Ions (Partial)

Polyatomic ions are ions composed of two or more atoms covalently bonded that act as a single charged entity in chemical reactions.

Name	Formula	Charge
Ammonium	NH4+	+1
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Hydroxide	OH-	-1
Carbonate	CO32-	-2
Phosphate	PO43-	-3
Chlorate	ClO3-	-1
Acetate	C2H3O2-	-1

Additional info: Table entries inferred from standard general chemistry curriculum.

Summary

• Atomic theory explains the structure and behavior of matter at the atomic level.

• Atoms are composed of protons, neutrons, and electrons; their arrangement determines the properties of elements and compounds.

• The periodic table organizes elements by atomic number and properties, facilitating the study of chemical behavior.

• Isotopes and ions are important for understanding chemical formulas and reactions.

• Polyatomic ions play a key role in the composition of many compounds.