Atomics-Quiz One
Chemistry Study Notes
Page 1: Definition of Chemistry
Chemistry: The study of the composition of matter and the changes that matter undergoes.
Page 2: Pure vs. Applied Chemistry
Pure Chemistry:
Pursuit of chemical knowledge.
Research can lead to new applications.
Applied Chemistry:
Research directed toward practical goals.
Applications may exist before understanding the underlying chemistry.
Page 3: Alchemy
Precursor to Modern Chemistry:
Objective: Creation of the philosopher's stone.
Capable of turning base metals into gold or silver.
Acts as an elixir of life for youth and immortality.
Contributions:
Development of tools and techniques still in use.
Introduction of experimentation and result analysis.
Page 4: Experimentation Process
Steps:
Make observations (collect factual data).
Form a hypothesis (educated guess).
Test the hypothesis through experimentation.
Form a conclusion (true or false).
Repeat experiments for consistency.
Share results and reformulate hypothesis if necessary.
Page 5: Analyzing Data
Graphical Representation:
Include title, x/y axis labels (with units), and uniform scale.
Ensure points fill the graph to reveal trends.
Page 6: Variables in Experiments
Example Study: Effect of sleep on test scores.
Independent Variable: Length of time spent sleeping.
Dependent Variable: Students' test scores.
Control: The test itself.
Hypothesis: More sleep leads to better exam performance.
Page 8: Problem Solving in Chemistry
Key Points:
Chemistry involves math and problem-solving.
Problems often require multiple steps.
Planning is essential for effective problem-solving.
Page 9: Dimensional Analysis
Method:
Use conversion factors (ratios of numbers and units that equal one).
Example: 12 inches = 1 foot.
Page 10: Unit Conversions
Examples:
2.00 gallons = 7.57 liters.
200 g = 0.441 lbs.
Page 12: Atomic Composition
Matter: Anything with mass and volume.
Chemistry Focus: Study of matter's composition and changes.
Page 13: Early Studies of Matter
Ancient Greeks:
Democritus: Proposed matter is made of indivisible atoms ("atomos").
Aristotle: Believed matter was continuous and composed of four elements.
Page 14: Change of Thought
Robert Boyle: Identified elements, moving away from Aristotle's four-element theory.
Element Definition: Substance that cannot be decomposed into simpler substances.
Page 15: Dalton’s Atomic Theory (1808)
Key Points:
Atoms cannot be subdivided, created, or destroyed.
Atoms combine in whole-number ratios to form compounds.
In reactions, atoms are rearranged.
All matter consists of atoms.
Page 16: Modern Atomic Theory
Changes to Dalton’s Theory:
Atoms of an element have a characteristic average mass (isotopes).
Atoms cannot be created or destroyed in ordinary reactions, but can in nuclear reactions.
Page 19: Thomson’s Atomic Model
Plum Pudding Model: Electrons embedded in a positively charged "pudding."
Page 20: Conclusions from Electron Studies
Findings:
Cathode rays have identical properties across elements.
Atoms must contain positive particles to balance electrons.
Most of the atom's mass comes from particles other than electrons.
Page 21-22: Radioactivity
Discoveries:
X-rays by W. C. Roentgen (1895).
Spontaneous radiation from uranium by Henri Becquerel (1896).
Isolation of radioactive components by Marie Curie and Pierre Curie.
Page 24: Rutherford’s Findings
Gold Foil Experiment:
Nucleus is small, dense, and positively charged.
Atoms are mostly empty space.
Page 25: Trouble with the Nuclear Atom
Mass Discrepancy:
Difficulty explaining atomic mass with only protons.
James Chadwick discovered the neutron in 1932.
Page 28: Atomic Particles
Particles:
Electron: Charge -1, mass 9.109 x 10^-31 kg.
Proton: Charge +1, mass 1.673 x 10^-27 kg.
Neutron: Charge 0, mass 1.675 x 10^-27 kg.
Page 29: Metric System
Origin: Developed in France during the French Revolution.
Purpose: Standardized measurements for the public.
Page 30: SI Prefixes
Common Prefixes:
Kilo (k): 10^3
Centi (c): 10^-2
Milli (m): 10^-3
Page 35: Scientific Notation
Usage: For extremely small or large numbers.
Format: One non-zero digit to the left of the decimal.
Example: 3,050,000 = 3.05 x 10^6.
Page 38: The Atomic Scale
Mass Distribution: Most mass in the nucleus; electrons in the electron cloud.
Page 39: Atomic Numbers and Mass Numbers
Definitions:
Mass Number = Protons + Neutrons.
Atomic Number = Number of Protons.
Page 40: Isotopes
Definition: Atoms of the same element with different mass numbers (different neutrons).
Page 43: Ions
Stability: Atoms gain or lose electrons to achieve stability (8 valence electrons).
Metals: Typically lose electrons; Nonmetals gain electrons.
Page 46: Atomic Mass Measurement
Method: Measured accurately using a mass spectrometer.
Page 50: Average Atomic Mass
Calculation: Weighted average of isotopes based on natural abundance.
Page 56: Isotope Abundance Calculation
Example: Given average atomic mass, calculate the abundance of