Week3

Earliest Atomic Theories

• Philosophical Origins:

  • Concept of atoms first suggested by Greek philosophers Leucippus and Democritus (5th century BC).

  • The term "atomos" means "indivisible" in Greek.

  • Aristotle proposed that matter was composed of four elements: fire, earth, air, and water.

• John Dalton's Contribution:

  • Dalton introduced his atomic theory in 1807, marking a significant advancement in chemical understanding.

Dalton’s Atomic Theory

• Five Postulates of Dalton’s Atomic Theory:

  1. Matter Composition: Matter is made of exceedingly small particles called atoms; these atoms are the smallest units participating in chemical change.

  2. Element Composition: Each element consists of one type of atom, which has a specific mass unique to that element.

• Example: A pre-1982 copper penny is composed of approximately 3 × 10²² copper atoms.

Continued Postulates

3. Diversity of Atoms: Atoms of different elements differ in properties from one another.

4. Compound Formation: A compound is made up of atoms of two or more elements combined in a fixed, whole-number ratio.

• Visual Example: Copper(II) oxide exhibits a 1:1 ratio of copper and oxygen atoms.

Final Postulate of Dalton

5. Chemical Changes: Atoms are neither created nor destroyed in chemical reactions; they are rearranged to form new types of matter.

• Visual Example: The reaction of copper and oxygen forms a new compound.

Law of Conservation of Matter

• Dalton's theory explains the macroscopic properties of matter:

  • Atoms remain constant in mass throughout chemical changes, a principle known as the law of conservation of matter.

Complete vs. Incomplete Reactions

• Assumption: In this section, reactions are treated as completely converting reactants into products.

• In Practice: Many reactions are incomplete, leaving some reactants unreacted.

Law of Definite Proportions

• Definition: All samples of a pure compound contain the same elements in the same proportion by mass.

• Illustration: Supported by French chemist Joseph Proust's experiments.

Concept of Constant Composition

• For any compound, its elemental composition remains consistent regardless of sample size.

• Table Example: Mass ratios of carbon to hydrogen across different samples showing consistent ratios.

Law of Multiple Proportions

• Definition: When two elements form multiple compounds, a fixed mass of one element combines with varying masses of the other in ratios of small whole numbers.

• Example: Two compounds with copper and chlorine exhibit different mass ratios.

Visualization of the Law of Multiple Proportions

• Visual Figures: Illustrate different compounds of copper and chlorine, showing their varied atomic ratios.

Discovery of the Electron

• J.J. Thomson’s Experiments: Utilized cathode ray tubes to discover cathode rays, which were deflected by positive charges, indicating the presence of a negatively charged particle (the electron).

Thomson’s Findings

• Cathode ray particles were much lighter than atoms and identical regardless of their source.

• Defined as electrons, negatively charged subatomic particles significantly lighter than atoms.

Thomson’s Cathode Ray Tube

• Visual description of the cathode ray tube experiment showing beam deflection and calculation of mass-to-charge ratios.

Millikan’s Oil Drop Experiment

• Overview: Evaluated the charge of individual oil droplets, allowing for the determination of the electron's charge.

Page 19: Millikan’s Apparatus

• Visual Data: Illustrates apparatus used in Millikan’s experiment, highlighting different charge measurements of oil drops.

Page 20: Millikan’s Results

• Established that the charge of an oil drop is always a multiple of 1.6 × 10⁻¹⁹ C, defining the charge of a single electron.

• Reiterates Thomson’s mass-to-charge ratio of the electron.

Page 21: Discovering the Nucleus

• Rutherford’s Gold Foil Experiment: Analyzing alpha particle scattering provided insights into atomic structure.

• Focus on the detection of alpha particles scattering through gold foil.

Page 22: Rutherford’s Experimental Setup

• Visual Schematic: Diagram showing how alpha particles interact with gold foil, illustrating scattering patterns.

Page 23: Rutherford’s Conclusions

• Atoms contain vast amounts of empty space with a small, dense, positively charged nucleus at their center, which contains most of an atom's mass.

• Electrons are dispersed around the nucleus.

Page 24: Visualization of Rutherford’s Findings

• Figure Description: Shows α particle behavior in relation to gold nucleus, delineating space occupied by atoms.

Page 25: Other Important Discoveries

• Isotopes: Discovered by Frederick Soddy, differing mass atoms of the same element.

• Neutrons: Identified by James Chadwick as uncharged particles located in the nucleus, similar in mass to protons.

Page 26: Section 2.3 Learning Objectives

• Write and interpret atomic symbols, including atomic number and mass number.

• Define atomic mass unit; calculate average atomic mass and isotopic abundance.

Page 27: Atomic Structure Overview

• Composition: Nucleus: contains most mass (protons + neutrons).

• Structure: Electrons occupy nearly all atomic volume.

• Size Ratios: Diameter of a typical atom is ~ 10⁻¹⁰ m; the nucleus is ~ 10⁻¹⁵ m (100,000 times smaller).

Page 28: Visualizing Atomic Scale

• Metaphor: If an atom were the size of a football stadium, the nucleus would compare to a blueberry.

Page 29: Atomic Units and Measurements

• Scale: Atoms and particles are extremely lightweight (e.g., carbon weighs less than 2 × 10⁻²³ g).

• Defined Atomic Mass Unit (amu): 1 amu = 1.6605 x 10⁻²⁴ g; for example, a carbon-12 atom weighs 12 amu.

Page 30: Properties of Subatomic Particles

• Protons: 1.0073 amu, Charge = +1

• Neutrons: 1.0087 amu, Charge = 0

• Electrons: 0.00055 amu, Charge = -1

Page 31: Atomic Number Definition

• Atomic number (Z) indicates the number of protons in an atom's nucleus, defining the element.

• Example: An atom with six protons is carbon (Z = 6).

Page 32: Neutral Atoms

• Neutral atoms must have equal numbers of protons and electrons; thus, atomic number also indicates electron count.

Page 33: Mass Number Explained

• Mass number (A) = Total protons + neutrons.

• Relationship: A = Z + Neutrons, where Z is atomic number.

Page 34: Ions Overview

• Ions form when the number of protons and electrons differ, altering electrical charge.

• Charge Calculated: Charge of an atom = Protons - Electrons.

Page 35: Cations and Anions Defined

• Anions: Mass gain causes negative charge from electron addition.

  • Example: Oxygen atom gains electrons (2− charge).

• Cations: Mass loss results in positive charge from electron removal.

  • Example: Sodium atom loses an electron (1+ charge).

Page 36: Chemical Symbols

• Symbols denote elements, example: Hg (mercury).

• Chemical symbols derived from element names, often with one or two letters; three-letter symbols exist as well.

Page 37: Importance of Chemical Symbols

• The chemical symbol represents the element regardless of the quantity (single atom or larger mass).

Page 38: Common Elements and Symbols

• Table of common elements:

  • Aluminum (Al), Iron (Fe), Bromine (Br), Lead (Pb), etc.

Page 39: Symbolizing Isotopes

• Isotope notation includes mass number as a left superscript and atomic number as a subscript.

• Example: Magnesium isotopes (24Mg, 25Mg, 26Mg) share proton count with differences in neutron count.

Page 40: Atomic Symbols Explained

• Atomic symbol format:

  • Element symbol (1 or 2 letters), mass number (superscript), atomic number (subscript), and charge

Page 41: Isotopic Abundance of Hydrogen

• Hydrogen Isotopes:

  • Protium (99.989%), Deuterium (0.0115%), Tritium (trace).

Page 42: Atomic Mass Fundamentals

• Atomic mass reflects protons and neutrons’ approximate weight of ~1 amu; varies due to isotopic mixtures.

Page 43: Calculating Atomic Mass

• Weighted averages of isotopes based on natural abundance.

• Example: Boron with isotopes 10B (19.9%) and 11B (80.1%).

Page 44: Mass Spectrometry Method

• Mass spectrometry determines isotopes’ occurrences and natural abundances by ionizing atoms and separating them by mass and charge.

Page 45: Mass Spectrometer Visual

• Mass Spectrum Sample: Shows different zirconium isotopes detected in a mass spectrometer.