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CHEM209+Lecture+2.5+Review+2024-2025+Spring+Eberts.pdf
Page 1: Course Information
Student Instructors in SLC
Monday: 9 am – 1 pm (Dina)
Wednesday: 9 am – 12 pm (Joann)
Assessments
Elements & Symbols Quiz
Exam 1: next Monday
Study Materials
Study Guide: posted
Lecture Review: today
Lab Schedule
Focus: Measurement & Density
Practice: Good practice for exam
Report: Must be uploaded as a single pdf file
Next Two Weeks: Lab will occur without prelab or synopsis
Page 2: Elements & Symbols Quiz
Quiz on Elements & Symbols (Details forthcoming)
Page 3: OpenStax Readings
Page 4: Chemistry in Context
Chemistry examines the composition, properties, and interactions of matter.
Historical Context:
2500-year history, beginning with the Greeks who claimed matter consists of:
Earth
Air
Fire
Water
Alchemists attempted to convert base metals into noble metals.
Page 5: Scientific Method
Chemistry is based on observation and experimentation.
Hypothesis: A tentative explanation of observations.
Laws of Science: Summarize vast experimental observations and predict natural phenomena.
Theory: Comprehensive and testable explanation of nature; hypotheses are less substantiated than theories.
Page 6: Scientific Method - Concept Identification
Identify if statements represent a hypothesis, law, or theory:
(a) A truck’s gas mileage has dropped significantly, probably due for a tune-up. (Hypothesis)
(b) The pressure of a gas is directly proportional to its temperature. (Law)
(c) Matter consists of tiny particles combining in specific ratios. (Theory)
Page 7: Scientific Method Process
Observations
Hypothesis
Experiment
Analysis
Theory
Laws
Page 8: Chemistry Domains
Chemists study matter and energy in three domains:
Macroscopic: Large enough to be sensed (detection by sight/touch).
Microscopic: Often imagined; observable through a microscope.
Symbolic: Specialized chemical language representing components (e.g., chemical symbols).
Page 9: Examples of Chemistry Domains
Macroscopic: Moisture in the air, icebergs, ocean represent water.
Microscopic: Gas (molecules disorganized and far apart), solid (molecules close together and organized), liquid (molecules close but disorganized).
Symbolic: Formula H2O (water); phases denoted by (g), (s), (l).
Page 10: OpenStax Readings
Page 11: Phases and Classification of Matter
Matter: Anything occupying space and having mass.
States of Matter:
Solid: Rigid with a definite shape.
Liquid: Flows, taking the shape of its container.
Gas: Takes both shape and volume of its container.
Page 12: Law of Conservation of Matter
Definition: No detectable change in total quantity of matter during conversions (chemical and physical changes).
Mass remains constant in processes like electricity production involving lead and sulfuric acid.
Page 13: Elements
Element: Pure substance that cannot be broken down by chemical changes, with 118 known elements in total.
90 elements occur naturally; others created in laboratories.
Page 14: Pure Substances and Mixtures
Characteristics of Pure Substances: Constant composition; includes elements (Au, P, O) and compounds (H2O, C6H12O6).
Differences: Properties of compounds differ from uncombined elements.
Page 15: Mixtures
Mixture: Composed of two or more types of matter; can be separated by physical changes.
Types of Mixtures:
Homogeneous: Uniform composition, also known as solutions.
Heterogeneous: Composition varies (e.g., oil and vinegar).
Page 16: Mixture Examples
(a) Heterogeneous: Oil and vinegar dressing.
(b) Homogeneous: Commercial sports drink.
Page 17: Classification of Matter
Matter can be classified based on its properties as either homogeneous mixture, heterogeneous mixture, compound, or element.
Page 18: Physical and Chemical Properties
Definition: Properties that distinguish substances.
Physical Property: Not associated with changes in chemical composition (e.g., density, color).
Physical Change: State change with no change in chemical composition.
Page 19: Chemical Properties
Chemical Property: Indicates a change in chemical composition (e.g., flammability, reactivity).
(a) Iron: Rusts (chemical property).
(b) Chromium: Does not rust (chemical property).
Page 20: Physical Change
Definition: Change in one or more physical properties without a chemical composition change (e.g., boiling water).
Page 21: Chemical Change
Definition: Substance transforms into new substances with different properties (e.g., electrolysis of water).
Page 22: Types of Properties
Extensive Properties: Depends on amount (e.g., mass, volume).
Intensive Properties: Independent of amount (e.g., density, temperature).
Page 23: OpenStax Readings
Page 24: Measurement
Measurements underpin hypotheses, theories, and laws in chemistry, providing:
Size or magnitude (a number).
Standard of comparison (a unit).
Indication of uncertainty.
Page 25: Base Units of the SI System
Length: meter (m)
Mass: kilogram (kg)
Time: second (s)
Temperature: kelvin (K)
Amount of Substance: mole (mol)
Importance of units in making sense of measurements highlighted.
Page 26: Common Unit Prefixes
Prefix and Factors:
femto (f): 10–15
pico (p): 10–12
nano (n): 10–9
micro (µ): 10–6
milli (m): 10–3
centi (c): 10–2
deci (d): 10–1
kilo (k): 10^3
mega (M): 10^6
giga (G): 10^9
tera (T): 10^{12}
Page 27: Derived SI Units: Density
Definition: Ratio of mass to volume.
Standard SI Unit: kilogram per cubic meter (kg/m3).
Common Units: g/cm3 (solids, liquids) and g/L (gases).
Page 28: Measurement Uncertainty, Accuracy, Precision
Counting: Free from uncertainty; results are exact.
Defined quantities also exact (e.g., 1 ft = 12 in).
Page 29: Measurement Uncertainty, Accuracy, Precision
Other measurements contain uncertainty requiring reporting of significant figures.
Record all certain digits and first uncertain digit.
Page 30: Reading Graduated Cylinder
Example: Volume of liquid in graduated cylinder requires estimating the meniscus reading.
Precision determined by place of uncertain digit (21.6 ml).
Page 31: Significant Figures
Record every digit in a measurement, including uncertain last digit.
Page 32: Significant Figures Rules
Always Significant:
Nonzero digits
Captive zeros
Trailing zeros (with decimal)
Never Significant:
Leading zeros
Trailing zeros (without decimal).
Page 33: Adding and Subtracting Significant Figures
Round result based on least number of decimal places in the least precise value.
Page 34: Multiplying and Dividing Significant Figures
Round based on the least number of significant figures in any number used in the calculation.
Page 35: Rounding Numbers
Rounding rules based on the value of the digit to be dropped (5 or more rounds up).
Page 36: Density Calculation Example
Use water displacement to calculate density.
Density calculation based on mass and volume results.
Page 37: Accuracy and Precision
Precision: Similar results when measured repeatedly.
Accuracy: Result close to true value.
Page 38: OpenStax Readings
Page 39: Mathematical Treatment of Measurement Results
Using dimensional analysis for indirect measurements: align units with mathematical operations.
Page 40: Conversion Factors and Dimensional Analysis
Definition: Ratio of two equivalent quantities with different measurement units.
Example: 2.54 cm = 1 inch.
Page 41: Using Conversion Factors
Use appropriate conversion factors for unit changes to ensure original units cancel correctly.
Page 42: OpenStax Readings on Atomic Theory
Page 43: Early Ideas in Atomic Theory
Concept: Atoms proposed by Leucippus and Democritus in the 5th century BC.
Term 'atomos' means indivisible.
Later views by Aristotle suggested combinations of four elements.
John Dalton’s Theory (1807):
Proposed atomic nature of matter.
Page 44: Dalton’s Atomic Theory - Postulates
Matter is made of small particles called atoms, the smallest unit that can engage in changes.
Each element consists of one atom type defined by characteristic mass.
Atoms of different elements have different properties.
Compounds consist of distinct atom ratios.
Atoms are rearranged but neither created nor destroyed in chemical reactions.
Page 45: Dalton’s Atomic Theory - Continued
Emphasizes the ratio of elements within a compound is consistent.
Page 46: Dalton’s Atomic Theory - Continued
Atoms rearranging confirm the conservation of matter in changes.
Page 47: Law of Conservation of Matter
Discusses how Dalton’s theory aligns with observed properties and conservation in chemical changes.
Page 48: Laws of Proportions
Law of Definite Proportions: Pure compounds have consistent elemental ratios.
Law of Multiple Proportions: When two elements form multiple compounds, fixed ratios occur for mass interactions.
Page 49: Law of Multiple Proportions - Example
Confirms ratios observed between different compounds formed from chlorine and copper.
Page 50: Discovery of the Electron: J.J. Thomson
Experiment using Cathode Ray Tube demonstrated properties of electrons and their resemblance across materials.
Page 51: Thomson's Conclusions
Cathode rays are negatively charged and significantly lighter than atoms; named electrons.
Page 52: Thomson's Atomic Model
Proposed 'Plum Pudding' model where electrons inhabit a positively charged sphere.
Page 53: Discovery of the Nucleus: Rutherford
Conducted the Gold Foil Scattering Experiment, revealing atomic structure and nucleus existence.
Page 54: Rutherford's Experiment Results
Most alpha particles passes through gold foil while a few deflected dramatically, indicating the nucleus.
Page 55: Rutherford’s Conclusions
Atoms possess empty space with dense positively charged nuclei at their centers, surrounding electrons.
Page 56: Advances in Atomic Theory
Isotopes: Atoms with varying masses discovered by Frederick Soddy (Nobel Prize).
Neutrons: Identified by James Chadwick (1932); uncharged particles in the nucleus.
Page 57: Atomic Structure and Symbolism
Nucleus: Contains most mass; protons and neutrons heavier than electrons.
Atomic Size: Diameter and nucleus size on widely different scales.
Page 58: Properties of Subatomic Particles
Proton: Mass = 1.0073 amu, Charge = +1
Neutron: Mass = 1.0087 amu, Charge = 0
Electron: Mass = 0.00055 amu, Charge = -1
Page 59: Atomic Number (Z)
Number of protons defines the element's identity; e.g., Carbon = 6 protons.
Page 60: Neutral Atoms
Equal numbers of protons and electrons in neutral atoms, indicating identity.
Page 61: Mass Number (A) & Isotopes
Mass Number (A): Total protons + neutrons.
Atomic Number (Z): Defines protons; Neutron count calculated from mass number and atomic number.
Page 62: Ions
Charged atoms result from electron loss or gain affecting balance between protons and electrons.
Page 63: Cations and Anions
Anion: Gains electrons, exhibits negative charge.
Cation: Loses electrons, exhibits positive charge.
Page 64: Protons & Electrons Assignment
Determining the number of protons and electrons for given ions (Na+, Cl-, O2-, Al3+, P3-).
Page 65: Chemical Symbols
Abbreviations representing elements; examples include Hg (mercury).
Generally one to two letters; capitalized appropriately.
Page 66: Element Symbols for Memorization
Common elements and their symbols (e.g., H for hydrogen, O for oxygen, Au for gold).
Page 67: Atomic Mass
Atoms consist of protons/neutrons weighing ~1 amu; electrons negligible.
Periodic table reflects weighted average based on isotopes.
Page 68: Determining Atomic Mass - Chlorine Example
Calculation based on isotope prevalence and contributions toward average atomic mass.
Page 69: Chemical Formulas
Molecular formula indicates atom types and counts; structural formula shows how atoms connect.
Page 70: Elements Existing as Molecules
Elements can exist as single atoms or molecules (e.g., diatomic molecules).
Page 71: Empirical vs. Molecular Formula
Empirical formula: simplest atomic ratios; molecular formula: actual atom counts per molecule.
Page 72: The Mole
The mole allows connection between mass and number of atoms/molecules, defined by carbon-12 standards.
Page 73: Avogadro's Number
Defined as 6.022 x 10^23, linking count of molecules to moles.
Page 74: Molar Mass
Molar mass of a substance represents mass of one mole in g/mol, linked to atomic mass in amu.
Page 75: Calculations - Dimensional Analysis
Methods to compute substance quantities using mole relationships.
Page 76: Calculations Example - Potassium Requirement
Inquiry into dietary potassium conversion from grams to moles.
Page 77: Calculations Example - Argon Mass
Compute mass from molar quantities in air calculations.
Page 78: Copper Atoms in Copper Wire
Mass leads to calculation of total copper atoms using Avogadro's.
Page 79: Molar Mass Contributions Example
Breakdown of molecular weight determinations through dimensional analysis.
Page 80: Vitamin C Mass Calculation
Related daily dietary allowances and molar mass conversions for vitamin C.
Page 81: Mass to Atoms Process
Methodology translating atomic counts to mass.
Page 82: Mass to Molecules Calculation
Conversion techniques yielding mass for given molecule counts in chemical analysis.
NoteStudied by 11 peopleNoteStudied by 52 peopleNoteStudied by 48 peopleNoteStudied by 7 peopleNoteStudied by 10 peopleNoteStudied by 153 people