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:
Matter Composition: Matter is made of exceedingly small particles called atoms; these atoms are the smallest units participating in chemical change.
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
Diversity of Atoms: Atoms of different elements differ in properties from one another.
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
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.