CHM 111 General Chemistry - Introduction to Chemistry, Scientific Method, and Measurement

Course Overview and Academic Resources

• Course Designation: CHM 111
• Academic Support Contacts and Campus Resources:
  • STEM Center: ST318, Phone: 985-5654
  • Tutoring Services: Academic Link located in Brown Library
  • Counseling Services: Kimberly Stafford, Phone: 857-6711
  • Writing Center
• Core Course Platforms and References:
  • Canvas Learning Management System
  • Official Course Syllabus
  • Instructor Office Hours
  • MindTap Online Homework System
  • Khan Academy Online Resources
  • LibreTexts Chemistry Textbook Platform

Strategies for Success in Science Courses

• Dedicated Study Time: Allocate approximately $3\,\text{hours}$ of study time outside of class for every $1\,\text{credit}$ hour of enrollment.
• Textbook Reading: Consistently read assigned textbook sections to solidify concepts.
• Time Management: Avoid last-minute cramming; maintain a steady study schedule.
• Deep Understanding: Focus on understanding core concepts rather than relying exclusively on rote memorization.
• Critical Analysis: Actively sharpen tools for critical analysis and logical reasoning.
• Visual Data Literacy: Carefully analyze graphs, tables, and figures from lectures and readings.
• Laboratory Preparation: Prepare thoroughly for laboratory experiments prior to attending lab sessions.
• Practice Execution: Solve extensive practice problems repeatedly to master calculations and concepts.
• Active Inquiry: Ask questions immediately when encountering confusing or unclear material.

Chapter 1 Overview: Chemistry and Matter on the Atomic Scale

• Chapter Scope:
  • Section 1.1: What Is Chemistry?
  • Section 1.2: Classification of Matter
  • Section 1.3: Units and Measurement
  • Section 1.4: Unit Conversions
• Chemistry Scale Comparison:
  • A human is approximately $9$ orders of magnitude larger than an aspirin molecule.
  • This scale difference corresponds to a factor of $10^9$, or $1\,\text{billion}$ times larger.

Scientific Method and Research Methodology

• Sequential Steps of the Scientific Method:
  • Observation: Identifying a phenomenon in the physical world.
  • Question: Formulating a specific inquiry based on physical observations.
  • Research: Investigating existing literature surrounding the observed trend.
  • Hypothesis: Establishing a testable explanation guided by inductive reasoning.
  • Experiment: Testing predictions under controlled experimental conditions.
  • Data Collection: Gathering objective quantitative and qualitative measurements.
  • Conclusion: Analyzing collected data to corroborate or refute the hypothesis.
• Categorization of Observations:
  • Subjective Observations: Statements based on personal feelings, judgments, or opinions rather than verifiable facts.
  • Example: Starbucks Coffee is amazing.
  • Example: Nirvana is an awesome band.
  • Example: Tennessee is an awesome football team.
  • Example: It is about $2\,\text{miles}$ to I-581 from campus.
  • Objective Observations: Statements grounded strictly in measurable or verifiable facts (required for scientific inquiry).
  • Example: Fishburn Hall is beside Brown Library.
  • Example: 81 is an interstate highway.
  • Example: There is a Starbucks Coffee store on the VWCC campus.
  • Example: There are $100\,\text{centimeters}$ in $1\,\text{meter}$.
• Literature Research Rules:
  • Review peer-reviewed journal articles related to the initial observation.
  • Evaluate whether other researchers have observed similar trends and proposed valid explanations.
  • Utilize gathered information to structure a testable hypothesis and design controlled experiments.
  • Avoid non-peer-reviewed sources or open-wiki websites.
• Scientific Hypotheses:
  • Definition: A rational accounting for a set of observations, guided by inductive reasoning ("an explanation on trial").
  • Must yield specific predictions that can be empirically tested via further observations or controlled experimentation.
• Experimental Variables and Controls:
  • Controlled Experiment: Compares an experimental group against an untreated control group.
  • Experimental Variables: Factors that are altered or measured within the experimental system.
  • Independent Variable: The specific factor directly manipulated by researchers.
  • Dependent Variable: The observed effect or response resulting from manipulation of the independent variable.
• Data Analysis and Reporting:
  • Analyze data using mathematical equations, graphical representations, data tables, and error analysis.
  • Draw logical conclusions that either corroborate (support) or refute the starting hypothesis.
  • Publish and report findings to the scientific community.

Classification and States of Matter

• Atomic Scale Classification of Matter:
  • Pure Substance: Matter composed of only a single type of atom or molecule.
  • Element: A pure substance that cannot be broken down into simpler substances by chemical reactions (e.g., copper atoms on a silica surface visualized via Scanning Tunneling Microscope).
  • Compound: A pure substance composed of two or more distinct elements chemically combined in fixed proportions.
• Categorization Practice (Elements vs. Compounds):
  • $\text{Fe}$: Element
  • $\text{CH}_4$: Compound
  • $\text{O}_2$: Element
  • $\text{FeCl}_3$: Compound
• Macroscopic States of Matter:
  • Solid State:
  • Characterized by a rigid, defined shape and fixed volume.
  • Examples: Ice, iron ($\text{Fe}$), gold ($\text{Au}$).
  • Crystalline Solid: Atoms or molecules are arranged in long-range, repeating geometric patterns (e.g., diamond, graphite).
  • Amorphous Solid: Atoms or molecules lack any long-range positional order (e.g., glass, plastic).
  • Liquid State:
  • Characterized by a definite volume but no definite shape; takes the shape of its container while maintaining a free surface.
  • Particles are packed closely together but possess freedom of movement past one another.
  • Gas State:
  • Characterized by having neither a definite shape nor a definite volume.
  • Expands completely to occupy the shape and volume of its container; easily compressible due to large interparticle spacing.
• Scale Categorization Practice:
  • An RNA molecule: Atomic scale
  • A mercury atom: Atomic scale
  • A sample of liquid mercury: Macroscopic scale

Physical and Chemical Changes and Properties

• Physical vs. Chemical Changes:
  • Physical Change: A process that alters state or appearance without altering chemical composition.
  • Example: Liquid water boiling on a stove into gaseous steam:

$\text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{O}(g)$

• Chemical Change: A process that alters the chemical composition of matter, converting it into entirely different substances.
  • Example: An iron nail rusting, where iron combines with oxygen to yield iron(II) oxide:

$2\,\text{Fe}(s) + \text{O}_2(g) \rightarrow 2\,\text{FeO}(s)$

• Physical vs. Chemical Properties:
  • Physical Property: A property displayed by a substance without changing its chemical composition.
  • Examples: Melting point, boiling point, density, odor.
  • Chemical Property: A property displayed by a substance only through changing its chemical composition via a chemical reaction.
  • Examples: Flammability, corrosiveness, acidity, toxicity, reaction of alkali metals with water.
• Property Classification Practice:
  • Evaporation of rubbing alcohol: Physical change (Physical property)
  • Burning of lamp oil: Chemical change (Chemical property)
  • Bleaching of hair with hydrogen peroxide: Chemical change (Chemical property)
  • Forming of frost on a cold night: Physical change (Physical property)

Classification and Separation of Mixtures

• Mixture Definition: Matter composed of two or more different types of atoms or molecules combined in continuously variable proportions.
• Types of Mixtures:
  • Heterogeneous Mixture: A mixture in which the physical composition varies from one region to another.
  • Homogeneous Mixture (Solution): A mixture in which the physical composition is uniform throughout.
• Physical Separation Methods:
  • Separation by Density: Decantation and centrifugation
  • Separation by Boiling Point: Distillation
  • Separation by State of Matter: Filtration
  • Separation by Intermolecular Forces: Chromatography
  • Separation by Vapor Pressure: Evaporation
  • Separation by Magnetism: Magnets
  • Separation by Solubility: Filtration

Measurement, Units, and SI Prefixes

• Units Overview: Standard quantities used to express physical measurements across the English system, Metric system, and International System of Units (SI).
• Mandatory SI Prefixes (Atto- to Exa-):
  • atto- ($\text{a}$): $10^{-18}$
  • femto- ($\text{f}$): $10^{-15}$
  • pico- ($\text{p}$): $10^{-12}$
  • nano- ($\text{n}$): $10^{-9}$
  • micro- ($\mu$): $10^{-6}$
  • milli- ($\text{m}$): $10^{-3}$
  • centi- ($\text{c}$): $10^{-2}$
  • deci- ($\text{d}$): $10^{-1}$
  • deka- ($\text{da}$): $10^{1}$
  • hecto- ($\text{h}$): $10^{2}$
  • kilo- ($\text{k}$): $10^{3}$
  • mega- ($\text{M}$): $10^{6}$
  • giga- ($\text{G}$): $10^{9}$
  • tera- ($\text{T}$): $10^{12}$
  • peta- ($\text{P}$): $10^{15}$
  • exa- ($\text{E}$): $10^{18}$
  • Note: SI prefix factors must be fully memorized. English-to-English and metric-to-English conversions are provided on exams.
• Standard Physical Units:
  • Length: The SI base unit is the meter ($\text{m}$).
  • Defined by the exact speed of light in a vacuum:

$c = 299\,\!792\,\!458\,\text{m/s}$

* Length equalities:

$100\,\text{cm} = 1\,\text{m}$

$1\,\text{cm} = 10^{-2}\,\text{m}$

$1\,\text{nm} = 10^{-9}\,\text{m}$

• Volume: Measure of space occupied by matter.
  • Expressed in cubic centimeters ($\text{cm}^3$), liters ($\text{L}$), and milliliters ($\text{mL}$).
  • Essential identity:

$1\,\text{cm}^3 = 1\,\text{mL}$

• Mass: Measure of matter contained within an object (distinct from weight).
  • Units: Milligrams ($\text{mg}$), grams ($\text{g}$), kilograms ($\text{kg}$).

Temperature Scales and Energy

• Temperature and Heat Flow:
  • Temperature determines the direction of heat flow; heat flows spontaneously from high temperature to low temperature.
  • Chemistry temperature scales: Celsius ($^{\circ}\text{C}$) and Kelvin ($\text{K}$).
  • Temperature conversion formulas:

$t_{\text{F}} = 1.8 \times t_{\text{C}} + 32^{\circ}$

$T_{\text{K}} = t_{\text{C}} + 273.15$

• Temperature Scale Comparison Points:
  • Water Boiling Point: $212^{\circ}\text{F} = 100^{\circ}\text{C} = 373.15\,\text{K}$
  • Water Freezing Point: $32^{\circ}\text{F} = 0^{\circ}\text{C} = 273.15\,\text{K}$
  • Coincidence Point: $-40^{\circ}\text{F} = -40^{\circ}\text{C} = 233.15\,\text{K}$
  • Temperature Intervals: An interval of $180\,\text{Fahrenheit degrees}$ equals $100\,\text{Celsius degrees}$ and $100\,\text{kelvins}$.
• Energy Principles:
  • Energy: The capacity to do work.
  • Total Energy: Sum of kinetic energy and potential energy.
  • Work: Action of a force through a distance.
  • Kinetic Energy: Energy associated with object motion.
  • Potential Energy: Energy associated with an object's position or chemical composition.
  • Thermal Energy: Energy associated with temperature (a form of kinetic energy).
  • Law of Conservation of Energy: Energy is neither created nor destroyed.
  • SI Unit of Energy: Joule ($\text{J}$), defined as the energy needed to accelerate a $1\,\text{kg}$ mass using a $1\,\text{N}$ force across $1\,\text{m}$:

$1\,\text{J} = 1\,\text{kg} \times \frac{\text{m}^2}{\text{s}^2} = 1\,\text{N} \cdot \text{m}$

Density and Physical Quantities

• Density Definition and Formula:
  • Density is an intensive physical property of matter:

$\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Standard units: $\text{g/mL}$ or $\text{g/cm}^3$ for solids and liquids; $\text{g/L}$ for gases.
  • Intensive vs. Extensive Properties:
• Intensive Property: Independent of sample size (e.g., density, temperature).
• Extensive Property: Dependent on sample size (e.g., mass, volume).
  • Density Comparison Example:
• Material A: $\text{Mass} = 10.0\,\text{g}$, $\text{Volume} = 2.00\,\text{L} = 2000\,\text{mL}$

$d_{\text{A}} = \frac{10.0\,\text{g}}{2000\,\text{mL}} = 0.00500\,\text{g/mL}$

• Material B: $\text{Mass} = 5.00\,\text{g}$, $\text{Volume} = 10.0\,\text{cm}^3$

$d_{\text{B}} = \frac{5.00\,\text{g}}{10.0\,\text{cm}^3} = 0.500\,\text{g/cm}^3$

• Material C: Sinks in ethanol ($d = 0.789\,\text{g/cm}^3$) and floats on water ($d = 1.00\,\text{g/cm}^3$)

$0.789\,\text{g/cm}^3 < d_{\text{C}} < 1.00\,\text{g/cm}^3$

• Conclusion: Material C possesses the largest density.

Significant Figures, Precision, and Accuracy

• Significant Figures Principles:
  • Reflect experimental uncertainty in measured values.
  • Examples: $8.00\,\text{mL}$ ($3$ significant figures), $760 \pm 10\,\text{mg}$, $5.47 \times 10^{-2}\,\text{mL}$ ($3$ significant figures).
• Mathematical Rules for Calculations:
  • Multiplication and Division: Answer retains the same number of significant figures as the factor with the fewest significant figures.
  • Practice 1:

$5.03\,\text{mL} \times 2.1\,\text{mL} \times 10.04\,\text{mL} = 106.05252\,\text{mL}^3$

* Smallest significant figures in input is $2$ (from $2.1\,\text{mL}$). Reported answer: $1.1 \times 10^2\,\text{mL}^3$ (or $110\,\text{mL}^3$).
* Practice 2:

$\frac{2.543\,\text{cm} \times 5.81\,\text{cm}}{0.02\,\text{cm} \times 1.1\,\text{cm}} = 671.5831$

* Smallest significant figures in input is $1$ (from $0.02\,\text{cm}$). Reported answer: 7 \times 10^2$.\n  * Addition and Subtraction: Answer retains the same number of decimal places as the term with the fewest decimal places.\n    * Practice:\n\n1.12 + 0.1 + 2.567 = 3.787\n\n    * Smallest decimal places in input is 1$(from$0.1$). Reported answer:$3.8$.

• Intermediate Rounding Rule:
  • Keep all digits during multi-step calculations; underline the least significant digit at each step, and round only at the final step.
• Exact Numbers:
  • Defined values or counts do not limit the significant figures of a calculation.
  • Example: Force on a dropped $10.1\,\text{kg}$ mass:

$F = m \times g = 10.1\,\text{kg} \times 9.8\,\text{m/s}^2 = 98.98\,\text{kg} \cdot \text{m/s}^2$

• Reported to $2$ significant figures: $99\,\text{N}$.
  • Precision vs. Accuracy:
• Precision: Closeness of agreement among multiple experimental measurements under identical conditions.
• Accuracy: Closeness of a measured value to the true or accepted value.

General Problem-Solving Strategy

• Step 1: Write down given values with explicit units and labels.
• Step 2: Clearly identify the target variable to solve for.
• Step 3: Determine required formulas, conversion factors, and constants.
• Step 4: Perform calculations using dimensional analysis and verify that the final answer is physically reasonable.

Comprehensive Practice Problems and Solutions

• Practice Problem 1: Converting Units Raised to Powers
  • Task: Convert $4.85\,\text{cm}^3$ to $\text{in}^3$.
  • Calculation:

$4.85\,\text{cm}^3 \times \left(\frac{1\,\text{in}}{2.54\,\text{cm}}\right)^3 = 4.85\,\text{cm}^3 \times \frac{1\,\text{in}^3}{16.387\,\text{cm}^3} = 0.29596\,\text{in}^3$

• Final Answer (3 sig figs): $0.296\,\text{in}^3$

  • Practice Problem 2: Metric Unit Conversions

• Convert $7.5\,\text{cm}$ to $\text{m}$:

$7.5\,\text{cm} \times \frac{10^{-2}\,\text{m}}{1\,\text{cm}} = 0.075\,\text{m}$

• Convert $1.75\,\text{mg}$ to $\text{g}$:

$1.75\,\text{mg} \times \frac{10^{-3}\,\text{g}}{1\,\text{mg}} = 0.00175\,\text{g}$

• Convert $2.95 \times 10^6\,\text{L}$ to $\text{ML}$:

$2.95 \times 10^6\,\text{L} \times \frac{1\,\text{ML}}{10^6\,\text{L}} = 2.95\,\text{ML}$

• Practice Problem 3: Density Unit Conversion (Example 1.4.3)
  • Task: Use the density of diamond ($d = 3.52\,\text{g/cm}^3$) to find the volume of a $1.2\text{-carat}$ diamond ($1\,\text{carat} = 0.200\,\text{g}$).
  • Mass calculation:

$1.2\,\text{carats} \times \frac{0.200\,\text{g}}{1\,\text{carat}} = 0.24\,\text{g}$

• Volume calculation:

$V = \frac{m}{d} = \frac{0.24\,\text{g}}{3.52\,\text{g/cm}^3} = 0.06818\,\text{cm}^3$

• Final Answer (2 sig figs): $0.068\,\text{cm}^3$

  • Practice Problem 4: Distance Unit Conversion

• Task: Convert $125\,\text{miles}$ to meters.

• Calculation:

$125\,\text{mi} \times \frac{1.60934\,\text{km}}{1\,\text{mi}} \times \frac{1000\,\text{m}}{1\,\text{km}} = 201\,\!167.5\,\text{m}$

• Final Answer (3 sig figs): $2.01 \times 10^5\,\text{m}$

  • Practice Problem 5: Nickel Atom Calculations

• Part A: Radius of a nickel atom is $125\,\text{pm}$. Convert to centimeters.

$125\,\text{pm} \times \frac{10^{-12}\,\text{m}}{1\,\text{pm}} \times \frac{100\,\text{cm}}{1\,\text{m}} = 1.25 \times 10^{-8}\,\text{cm}$

• Part B: Single nickel atom mass is $9.75 \times 10^{-23}\,\text{g}$. Determine how many nickel atoms lined up side-by-side are required to span $48.5\,\text{cm}$ ($19\,\text{inches}$).
  • Atomic Diameter:

$\text{Diameter} = 2 \times r = 2 \times (1.25 \times 10^{-8}\,\text{cm}) = 2.50 \times 10^{-8}\,\text{cm/atom}$

* Number of Atoms:

$\text{Atoms} = \frac{48.5\,\text{cm}}{2.50 \times 10^{-8}\,\text{cm/atom}} = 1.94 \times 10^9\,\text{atoms}$

• Part C: Mass of nickel required for this span in grams:

$\text{Mass} = (1.94 \times 10^9\,\text{atoms}) \times (9.75 \times 10^{-23}\,\text{g/atom}) = 1.8915 \times 10^{-13}\,\text{g}$

• Final Answer (3 sig figs): $1.89 \times 10^{-13}\,\text{g}$