Chapter 2 Atoms and Molecules Libre

Introduction to Atoms, Molecules, and Ions

• General Chemistry I by Professor McMahon

Question to Consider

• Dimethyl ether and ethanol both have the formula C2H6O.

• Dimethyl ether is used as a refrigerant and in wart removers.

• Ethanol is the main component in alcoholic beverages.

• The chemical structure of these two compounds explains their different properties and uses.

Section 2.1 Early Ideas in Atomic Theory

Objectives

• State postulates of Dalton’s atomic theory.

• Explain the laws of definite and multiple proportions using Dalton’s theory.

Dalton’s Atomic Theory

• Proposed in 1807 by John Dalton, building on earlier Greek ideas.

1. Matter is made of atoms—tiny particles that are the smallest unit of an element.

2. An element consists of only one type of atom with a characteristic mass.

3. Compounds are formed from two or more elements in constant ratios.

4. Atoms are neither created nor destroyed in chemical reactions but are rearranged.

Laws Explained by Dalton’s Theory

• Law of Conservation of Mass: Matter cannot be created or destroyed.

• Law of Definite Proportions: All samples of a pure compound contain the same elements in the same mass proportions.

  • Example: Isooctane (constant composition across samples).

• Law of Multiple Proportions: When elements form more than one compound, they do so in fixed, small whole-number ratios.

  • Example: Compounds of copper and chlorine.

Section 2.2 Evolution of Atomic Theory

Objectives

• Milestones in atomic theory development.

• Summarize Thomson, Millikan, and Rutherford's experiments.

• Describe subatomic particles and define isotopes.

J.J. Thomson + Electrons

• Discovered electrons in 1891 using cathode rays; established they are negatively charged.

• Electrons are much less massive than atoms.

Robert Millikan's Experiment

• In 1909, determined the charge of the electron (1.6 x 10^-19 C) through oil droplet experiments,

• This allowed calculation of electron mass: (9.107 x 10^-31 kg).

Ernest Rutherford + the Nucleus

• Conducted gold foil experiment, demonstrating that atoms are mostly empty space with a small, dense nucleus.

• Proposed nuclear theory of the atom.

Neutrons and Isotopes

• Proposed existence of neutrons in 1932 to explain atomic mass discrepancies.

• Isotopes: Variants of elements with the same number of protons but different numbers of neutrons.

Section 2.3 Atomic Structure and Symbolism

Objectives

• Write and interpret atomic symbols.

• Define and calculate average atomic mass.

Atomic Structure

• Nucleus contains protons and neutrons, while electrons orbit in the empty space.

• Atomic mass unit (amu) defined relative to carbon-12.

Electrons and Ions

• Atomic charge is defined as #protons - #electrons.

• Anions = gain of electrons (negative charge), Cations = loss of electrons (positive charge).

Chemical Symbols

• Notations such as Mg-24 or (24/12)Mg indicate atomic composition.

• Isotopes reflect variations in atomic mass.

Average Atomic Mass

• Calculated by weighted average of naturally occurring isotopes.

Section 2.4 The Periodic Table

Objectives

• Explain the periodic law and element organization.

• Classify metals, nonmetals, and metalloids.

Mendeleev and the Periodic Law

• Mendeleev arranged elements by increasing atomic mass and noted periodic properties.

• Modern periodic law organizes elements by atomic number.

Classification of Elements

• Metals: Shiny, malleable, ductile, conductive (right side of table).

• Nonmetals: Gas/liquid, nonconductive (left side).

• Metalloids: Exhibit properties of both.

Summary of Key Periodic Table Properties

• Main group elements exhibit predictable behaviors; transition metals do not.

• Common names for groups include: alkali metals, alkaline earth metals, halogens, noble gases.