Atomic Theory Study Guide (Public)

Honors Chemistry Packet: Atomic Theory

Lesson #1: Early Atomic Theory

Objectives

  • Compare and contrast a “scientific law” and a “theory”.

  • State Dalton’s Atomic Theory.

  • Use Dalton’s Atomic Theory to explain early chemistry laws.

  • Summarize Cathode-Ray Tube Experiments.

  • Relate how Cathode-Ray Tube Experiments led to a change in Atomic Theory.

  • Summarize Rutherford’s-Gold Foil Experiment.

  • Relate how Rutherford’s-Gold Foil Experiment led to a change in Atomic Theory.

Lesson #1-Part A

  • Objective: Compare and contrast a "scientific law" and a "theory".

Scientific Inquiry

Scientific Law

  • A generalization of scientific observations.

  • Describes what happens but does not explain why it happens.

Theory (Model)

  • A set of assumptions used to explain observations and predict new observations.

  • Theories can never be proven 100% correct.

  • Theories may change and evolve with new information.

  • Considered successful when they explain and predict observations.

Lesson #1-Part B

  • Objective: State Dalton's Atomic Theory.

  • Use Dalton's Atomic Theory to explain early chemistry laws.

Philosophical Precursors to Atomic Theory

Democritus (460-370 BC)

  • Greek philosopher who viewed matter as consisting of tiny particles called atoms.

  • His view lacked experimental evidence.

Aristotle (384-322 BC)

  • Greek philosopher who viewed matter as continuous.

  • His views dominated until the 18th century and were also not based on experimental evidence.

Dalton's Atomic Theory

John Dalton (1766-1844)

  • English schoolteacher who proposed the first Atomic Theory in 1803.

Assumptions of Dalton’s Atomic Theory

  • All matter is composed of extremely small particles (atoms).

  • Atoms of the same element are identical in size and mass; atoms of different elements differ.

  • Atoms cannot be subdivided, created, or destroyed.

  • Atoms combine in simple whole-number ratios to form chemical compounds.

  • In chemical reactions, atoms are combined, separated, or rearranged.

Importance of Dalton’s Atomic Theory

  • Helped explain chemical laws and predict new observations, stimulating further research.

Law of Conservation of Mass

  • Matter is not created or destroyed in a chemical reaction.

Law of Definite Proportions

  • Compounds contain the same elements in the same proportions by mass.

Law of Multiple Proportions

  • Masses of one element that combine with a constant mass of another element, forming more than one compound, are in a ratio of small whole numbers.

Improvements of Dalton's Theory

  • Experimental evidence showed:

    • Atoms can have different masses (isotopes).

    • Atoms can be subdivided into smaller particles.

    • Atoms of one element can be transformed into atoms of another element in nuclear reactions.

Lesson #1-Part C

  • Objectives: Summarize Cathode-Ray Tube Experiments and their impact on Atomic Theory.

Cathode-Ray Tube Experiments (Late 1800's)

Setup

  • A glass tube containing a gas at low pressure with a cathode (negative) and anode (positive).

  • High voltage applied causes a cathode ray to glow.

Observations and Implications

  • Experiment 1: Deflection of a cathode ray towards a positive plate indicates a negative charge.

  • Experiment 2: Cathode ray moves a paddle-wheel showing it has mass.

J. J. Thomson (1856-1940)

  • Conducted cathode-ray experiments in 1897.

  • Discovered electrons, identifying them as the first subatomic particles.

Thomson Model of the Atom (Plum Pudding Model)

  • Proposed an atom as a positive cloud with negatively charged electrons embedded within.

Lesson #1-Part D

  • Objectives: Summarize Rutherford’s Gold Foil Experiment

  • Understand its impact on Atomic Theory.

Rutherford's Gold Foil Experiment (1908-1909)

Setup

  • Bombarded thin gold foil with alpha particles.

Observations

  • Most particles passed through, some deflected at large angles, and a few redirected.

Implications

  • Proposed the existence of a nucleus containing dense, positively charged matter.

Rutherford Model of the Atom

  • Proposed a nuclear model where electrons orbit a dense nucleus composed of protons and neutrons.

  • Electrons are located outside the nucleus.

Lesson #2: Modern Atomic Theory

Essential Question
Objectives

  • Outline the Modern Model of the Atom, represent isotopes, define “atomic mass,” and calculate average atomic mass.

Modern Model of the Atom

  • An atom is composed of subatomic particles:

    • Protons: Positive charge, found in the nucleus.

    • Neutrons: Neutral charge, found in the nucleus.

    • Electrons: Negative charge, found in the space around the nucleus.

Characteristics of Atoms

  • Atoms are mostly empty space.

  • Atomic number = number of protons, determines element identity.

  • Isotopes have the same number of protons but different numbers of neutrons.

Atomic Mass

  • Defined as the average mass of isotopes of an element, given in atomic mass units (amu). Average atomic mass reflects naturally occurring isotopes.

Lesson #3: The Mole

Objectives

  • Define “mole”, compare it to other counting units, and learn about “molar mass.”

Definition of Mole

  • A counting unit quantifying amount of substance, similar to a dozen but represents 6.022 × 10²³ units.

  • Avogadro’s Number: 6.022 × 10²³ molecules or atoms.

Molar Mass

  • Mass of one mole of atoms (g/mol).

  • Numerically equal to atomic mass in amu but with different units.

Conversions Between Mass, Moles, and Atoms

  • Use molar mass and Avogadro’s number for conversions between grams, moles, and number of particles.