General Chemistry 1 — Comprehensive Notes

Introduction to General Chemistry 1

• Presented by Dennis M. Canonizado, MST Chem; Adviser - STEM fields.

General Chemistry 1 vs General Chemistry 2

• General Chemistry 1 topics: Properties of Matter, Classification, Chemical Reactions, Chemical Kinetics, Thermochemistry, Thermodynamics, Chemical Equilibria.

• General Chemistry 2 topics are not enumerated here but are contrasted with the list above.

What is Chemistry?

• Chemistry is the study of matter – its properties, composition, structure, changes it undergoes, energy involved, and the laws and principles that govern such changes.

Chemistry is the Central Science; Chemistry is a Servant Science

• Emphasizes the foundational role of chemistry in connecting and enabling other sciences.

Interdisciplinary Connections (Broader Fields Related to Chemistry)

• ENVIRONMENTAL SCIENCE: Ecology, Pollution Studies, Greenhouse Effect.

• BIOLOGY: Cell Biology, Genetics, Botany, Zoology, Microbiology; Biochemistry overlaps with Molecular Biology, Immunology, Endocrinology.

• GEOLOGY: Mineralogy, Petrology, Petrography, Paleontology, Structural Geology.

• MEDICINE AND ALLIED HEALTH SCIENCES: Pharmacology, Nutrition, Radiology.

• ASTRONOMY: Solar Astronomy, Galactic Astronomy, Astrophysics, Astrometry.

• PHYSICS: Atomic Physics, Nuclear Physics, Quantum Mechanics, Biomechanics, Material Science, Spectroscopy.

• NUCLEAR CHEMISTRY: Radio Chemistry, Body Imaging, Nuclear Medicine.

• Note: These items illustrate the broad application and integration of chemical principles across disciplines.

The Multidisciplinary Nature of Chemistry (as seen in the slides)

• CHEMISTRY sits at the crossroads with environmental science, biology, geology, medicine, astronomy, physics, and nuclear chemistry.

Branches of Chemistry

• INORGANIC: Everything except carbon-containing compounds.

• ORGANIC: Carbon-containing compounds.

• BIOCHEMISTRY: Chemistry of living things.

• PHYSICAL: Measuring physical properties of substances.

• ANALYTICAL: Qualitative and quantitative analysis.

Visual/Fragmentary Versions of “Branches” (repeated slides)

• The content repeats branches and cross-links to emphasize scope; core definitions above apply.

Father of Chemistry (Historical Perspectives)

• Page explains there is no universal agreement on a single father of chemistry.

• Popular choices:

  • Jābir ibn Ḥayyān (Jabir ibn Hayyan): Systematically classified chemicals; prepared inorganic chemicals from organic compounds.

  • Antoine Lavoisier: Discovered the role of oxygen in combustion; listed chemical elements; observed the conservation of mass in chemical reactions.

• “Father of Modern Chemistry” (as per slides):

  • Demise of phlogiston theory of combustion; modern concepts of element, compound & mixture; production of a table of elements; naming of oxygen and hydrogen; nomenclature; modern methods of chemical analysis (e.g., stoichiometry); conservation of mass.

The Beginning of Chemistry: Early Thinkers and Alchemy

• The Greeks believed in four elements; early practical chemistry involved household goods, weapons, soap, wine, basic medicine: earth, wind, air, fire.

• ALCHEMY: Medieval pursuit to transmute base metals into gold; sought the Philosopher’s Stone (~500–1300 C.E.); associated symbolisms for substances (Gold, Silver, Copper, Iron, Sand).

• Transmutation: the idea that one substance could become another; deconstruction of alchemical goals as impossible for elements to transmute into different elements in ordinary chemical reactions (illustrated via examples Copper → Gold and Philoopher’s Stone).

• Alchemy spread across regions (China, the Middle East, Western Europe ~500 C.E.); Modern chemistry evolved from alchemy.

• Contributions of Alchemists:

  • Experimental techniques; development of new glassware; information about elements; development of several alloys.

Uses of Chemistry

• Food Production; Clothing; Shelter; Medicine and Pharmaceuticals; Travel and Exploration; Communication; Population Control; Biotechnology; Leisure and Sports; Cosmetics; Warfare.

The Scientific Method (Mneumonic) – 7 Steps

• The mnemonic: "People Really Hate Elephants On Compact Cars" (memory aid).

• Steps:

  1. Problem – The thing you want to know; the question to answer.

  2. Research – Information about the problem; what is already known?

  3. Hypothesis – An educated guess or prediction of the outcome of the experiment.

  4. Experiment – Tests the hypothesis.

  5. Observations – Data collected during the experiment.

  6. Conclusions – Determine whether the hypothesis is correct.

  7. Communicate – Present results in a talk or paper.

Significant Figures (SF)

• SF definition: comprise all digits known with certainty plus the first digit that is uncertain.

• All nonzero digits are significant: e.g., $879.22$ has five SF; $24.46$ has four SF.

• Leading zeros are not significant: $0.0056$ has two SF.

• Zeros between nonzero digits are significant: e.g., $30{,}002$ has five SF; $0.004005$ has four SF.

• Trailing zeros after nonzero digits and to the right of the decimal point are significant: e.g., $192.0$ has four SF; $0.530$ has three SF.

• Trailing zeros without a decimal point are ambiguous and may not be significant: $45000$ has two SF; $4030700$ has five SF.

Significant Figures in Multiplication and Division

• The result should not have more SF than the measurement with the fewest SF among the factors.

• Example (as in slides):

  • $76.02 imes 4.35 imes 0.190 = 62.83053 o 62.8 ext{ (4 SF)}$; the slide shows 4 SF, 3 SF, 3 SF for the factors.

  • Division example: $125 <br>\, ÷ \, 30 = 4.16666667 o 4 ext{ (3 SF)} \, (1 ext{ SF})$

• Note: The slide’s numeric labels may reflect an instructional example; standard rule is the limiting SF among operands.

Significant Figures in Addition and Subtraction

• The result should not have more decimal places than the least precise measurement involved.

• Example combinations (as shown):

  • Addition: 13.6 (1 digit past decimal) + 4.082 (3 digits past decimal) = 17.682 → 17.7

  • Subtraction: 65.41 (2 digits past decimal) − 52.8 (1 digit past decimal) = 12.61 → 12.6

Scientific Notation

• Used to express very large or very small numbers and to indicate precision (significant figures).

• Form: $( ext{number from 1 to } 9.999) imes 10^{ ext{exponent}}$

• Examples from slides:

  • 800 = $8 imes 10^{2}$

  • 2531 = $2.531 imes 10^{3}$

  • 0.0014 = $1.4 imes 10^{-3}$

• Converting to standard form examples:

  • $1.87 imes 10^{-5} = 0.0000187$

  • $3.7 imes 10^{8} = 370{,}000{,}000$

  • $7.88 imes 10^{?} = 78.8$ (note: slide shows “7.88 x 101 = 78.8” which implies a typographical simplification; actual correct is $7.88 imes 10^{1} = 78.8$)

  • $2.164 imes 10^{-2} = 0.02164$

• Change to scientific notation:

  • $12{,}340 = 1.234 imes 10^{4}$

  • $0.369 = 3.69 imes 10^{-1}$

  • $0.008 = 8 imes 10^{-3}$

  • $1{,}000{,}000{,}000 = 1 imes 10^{9}$

Exponent Key and Notation Tools

• Slides reference an “Exponent Key” with key labels (EE, EXP, etc.) for entering exponents.

• Practical note: Use scientific notation to manage significant figures when dealing with very large or very small numbers.

Measurement, Conversion Rules, and Unit Cancellations

• Core rule: Use equivalent measures and multiply or divide to convert units.

• Example: Inches to centimeters

  • $12 ext{ in} imes 2.54 rac{ ext{cm}}{ ext{in}} = 30.48 ext{ cm}$

• Example: Centimeters to inches

  • $5.12 ext{ cm} imes rac{1 ext{ in}}{2.54 ext{ cm}} o 2.013… ext{ in}$

• Common units and abbreviations (from the table):

  • US Customary: in, ft, yd, mi; fl oz, pt, qt, gal; oz, lb.

  • Metric: mm, cm, m, km; L, mL; g, kg.

• Heads-up on approximations in slides: 1 fl oz = 29.574 mL; 1 pt = 0.473 L; 1 qt = 0.946 L; 1 gal = 3.785 L; 1 in = 2.54 cm; 1 ft = 30.48 cm; 1 yd = 0.914 m; 1 mi = 1.609 km; 1 L = 1 dm³; 1 mL = 1 cm³; 1 kg = 2.205 lb.

• Note: Some values in slides contain minor typographical inconsistencies; the standard conversions above are conventional.

SI Prefixes and Units

• SI prefixes (examples):

  • kilo- (k) = 10^3

  • hecto- (h) = 10^2

  • deca- (da) = 10^1

  • deci- (d) = 10^{-1}

  • centi- (c) = 10^{-2}

  • milli- (m) = 10^{-3}

  • micro- (µ) = 10^{-6}

  • nano- (n) = 10^{-9}

  • pico- (p) = 10^{-12}

  • femto- (f) = 10^{-15}

  • atto- (a) = 10^{-18}

  • zepto- (z) = 10^{-21}

  • yocto- (y) = 10^{-24}

• Note: 1 mL = 1 cm³ and 1 L = 1 dm³ are standard equivalencies.

Conversion Factors (Common Equivalences)

• Length and Distance:

  • 1 in = 2.54 cm

  • 1 ft = 30.48 cm

  • 1 yd = 0.914 m

  • 1 mi = 1.609 km

• Volume and Capacity:

  • 1 L = 1 dm³

  • 1 L ≈ 1.057 qt; 1 qt ≈ 0.946 L; 1 gal ≈ 3.785 L

• Mass and Weight:

  • 1 oz = 28.350 g

  • 1 lb = 0.454 kg

  • 1 kg = 2.205 lb

  • 1 metric ton = 1.102 tons; 1 ton ≈ 0.907 metric tons

• Additional notes:

  • 1 cm³ = 1 mL; 1 L = 1000 mL

  • 1 m = 1.0936 yd; 1 km = 1000 m; 1 cm = 0.394 in; 1 m = 39.37 in (from standard conversions, with slide values present as references)

Conversion Strategy: Unit Cancellation

• Example problems (from slides):

  • How many cm are in 1.32 meters?

  • 1.32 m = 132 cm; thus 1.32 m × (100 cm / 1 m) = 132 cm.

  • How many m is 8.72 cm?

  • 8.72 cm = 0.0872 m; thus 8.72 cm × (1 m / 100 cm) = 0.0872 m.

  • How many kilometers is 15{,}000 decimeters?

  • 15{,}000 \, ext{dm} = 1.5 \, ext{km} (since 1 dm = 0.1 m and 1 km = 1000 m).

  • How many seconds is 4.38 days?

  • 4.38 d × (24 h/d) × (3600 s/h) = 378{,}432 s.

• Note on significant figures in these conversions: if applying significant figures, adjust accordingly.

Simple Math with Conversions

• Area and Volume:

  • Area of a rectangle: A = L × W

  • Example: L = 4.6 cm, W = 9.1 cm → A = 4.6 × 9.1 = 41.86 cm² → rounded to 42 cm² (as shown in slide)

• Unit conversions for area: 42 cm² × (1 m² / 10{,}000 cm²) = 0.0042 m²

• For volume: V = L × W × H

  • Example rectangular solid: L = 14.2 cm, W = 8.6 cm, H = 21.5 cm → V = 14.2 × 8.6 × 21.5 = 2600 cm³

• Converting V to mm³: 2600 cm³ × (1000 mm³ / 1 cm³) = 2{,}600{,}000 mm³ = 2.6 × 10^6 mm³

• Note: 1 cm² = 100 mm²; 1 cm³ = 1000 mm³; 1 cm = 10 mm.

Precision and Accuracy

• Definitions:

  • Accuracy: how closely measurements agree with the true value.

  • Precision: how close individual measurements are to the average value.

• Diagrams show combinations of high/low accuracy and precision:

  • Accurate and precise; accurate but not precise; precise but not accurate; neither.

• Visual diagrams demonstrate the concepts with labeled examples (descriptions provided in slides).

Properties of Matter

• Two broad property categories:

  • Extensive Properties: depend on the amount of material (mass, volume, weight).

  • Intrinsic/Intrinsic Properties: inherent to the substance and do not change with sample amount (odor, taste, color).

• Two more property classifications:

  • Intensive Properties: do not depend on the amount of material (e.g., density, melting point, boiling point).

  • Physical Properties: relate to physical appearance and transformation (mass, melting point, boiling point, solubility, hardness).

  • Chemical Properties: relate to a substance’s tendency to undergo chemical change (e.g., rusting of iron, ability of water to undergo electrolysis).

• Extrinsic Properties: vary for different samples of the same material (mass, volume, shape, size).

• Intrinsic Properties: do not change for different samples of the same substance (odor, taste, color).

Kinds of Matter by Phase

• Phases: Solid, Liquid, Gas, Plasma, Bose-Einstein Condensate.

• Phase Parameters (qualitative features):

  • Solid: Shape definite, Volume fixed, Compressibility low, Heat expansion low, Density high.

  • Liquid: Shape indefinite, Volume fixed, Compressibility medium, Heat expansion medium, Density medium.

  • Gas: Shape indefinite, Volume indefinite, Compressibility large, Heat expansion large, Density low.

• Note: A concise phase table is provided in slides; the above bullets summarize the intended content.

Kinds of Matter by Components

• PURE SUBSTANCE vs MIXTURE

  • Pure Substance: uniform composition; cannot be separated into simpler substances by physical means; can be decomposed chemically.

  • Mixture: can be physically separated; composition varies; includes Homogeneous Mixtures (solutions), Colloids, and Suspensions; Element is included as a fundamental kind.

• Additional nodes:

  • Elements: basic substances that cannot be decomposed into simpler substances by chemical means.

  • Compounds: substances composed of two or more elements in fixed ratios.

Changes in Matter

• Physical Change: changes in form or state without changing composition.

• Chemical Change: transforms into one or more different substances; may involve energy exchange.

• Nuclear Change: changes in the nucleus of atoms (not detailed in slides but listed).

Energetics Accompanying Changes of Matter

• Endothermic: reaction or change that absorbs heat (examples: photosynthesis, melting, vaporization).

• Exothermic: reaction or change that releases heat (examples: decomposition, neutralization, burning, combustion, corrosion).

Energy Involved in Chemical Reactions

• Three main energy forms:

  • Radiant Energy (Electromagnetic Radiation): includes radio waves, microwaves, X-rays, UV, visible light, infrared.

  • Heat Energy (Thermal Energy).

  • Chemical Energy (stored in chemical bonds).

Inspirational Quote

• Learning perspective: "Learning is not the product of teaching. Learning is the product of the activity of learners." — John Caldwell Holt.

Summary of Key Formulas and Concepts (LaTeX-friendly recap)

• Area of a rectangle: $A = L \times W$

• Volume of a rectangular solid: $V = L \times W \times H$

• Conversion principle: use equivalent measures to cancel units (unit analysis).

• Standard conversions (typical references):

  • Length: $1\text{ in} = 2.54\text{ cm}$; $1\text{ ft} = 30.48\text{ cm}$; $1\text{ yd} = 0.914\text{ m}$; $1\text{ mi} = 1.609\text{ km}$.

  • Volume: $1\text{ L} = 1\text{ dm}^3$; $1\text{ mL} = 1\text{ cm}^3$.

  • Mass: $1\text{ oz} = 28.350\text{ g}$; $1\text{ lb} = 0.454\text{ kg}$; $1\text{ kg} = 2.205\text{ lb}$.

  • Temperature and energy conversions are context-dependent and are not explicitly listed here.

Endnotes

• The slides contain some typographical inconsistencies (e.g., digits in exponent notation, minor unit conversions). When studying, rely on standard chemistry conventions for conversions and SF rules, but note what is presented in the transcript as the source material for this exam review.

• The content above is organized to reflect the transcript’s structure while preserving all major and minor points raised for comprehensive study.