Chemistry Vocabulary Flashcards - CHM 101 Notes (Chapter 1)

Chapter 1: Matter, Energy & Measurement — Objectives (Pages 7–8, 22, 63)

• 1) Classify matter by composition and state

• 2) Explain how chemical and physical properties describe/classify matter

• 3) Define kinetic and potential energy in chemical systems

• 4) Describe the metric system and the use of significant figures in reporting measurements

• 5) Convert measurements correctly using dimensional analysis

Matter: Substances and Mixtures (Pages 9–15)

• Definitions

  • Matter: anything that has mass and takes up space

  • Pure substance: distinct properties; composition does not vary sample to sample

• Subtypes of pure substances

  • Element: cannot be decomposed into simpler substances; Atoms are the building blocks; Atoms of an element are identical

  • Molecule: two or more atoms bound in a specific shape

  • Some molecules contain only atoms of one element

• Compound: composed of two or more elements; Molecules contain atoms of more than one element

• Law of Constant Composition (Definite Proportions)

  • The elemental composition of a compound is always the same

  • Example: one molecule of methane gas has exactly one carbon atom and four hydrogen atoms

• Mixture: two or more substances combined; Components retain chemical properties

  • Homogeneous mixtures (solutions): uniform throughout

  • Heterogeneous mixtures: not uniform

Pure Substances, Substances Summary (Page 15–16)

• Visual schematic: atoms, molecules, compounds, mixtures

• Homogeneous vs heterogeneous classification in a decision flow

Classification of Matter by Composition (Page 17)

• Decision tree outline

  • Matter > Is it uniform throughout?

  • YES: Homogeneous

    • If contains more than one kind of atom? YES → Compound; NO → Element or Homogeneous substance

  • NO: Heterogeneous mixture

• Labels: Pure substance, Element, Compound, Mixture, Homogeneous, Heterogeneous

States of Matter (Pages 18–21)

• Three states: Gas, Liquid, Solid

• Gas

  • No fixed shape or volume

  • Fills container; can be expanded or compressed

  • Atoms/molecules moved rapidly and are far apart

• Liquid

  • Definite volume, no fixed shape

  • Fills portion of container; not easily compressed

  • Molecules packed loosely but move quickly

• Solid

  • Definite shape and volume

  • Not compressible; molecules locked in place

Properties of Matter (Page 23)

• Physical properties: observed without changing composition

  • Examples: boiling point, density, mass, volume, odor, hardness

• Chemical properties: observed when a substance changes into another substance

  • Examples: flammability, corrosiveness, reactivity with acid

Changes to Matter (Pages 24–25)

• Physical changes: do not change the composition

  • Examples: phase changes, temperature changes, volume changes

• Chemical changes: form new substances

  • Examples: combustion (with oxygen), oxidation (rusting), decomposition

Physical Change Example (Page 25)

• Ice melting or water evaporating are physical changes; composition (H2O) remains the same: two H atoms and one O atom per molecule

Chemical Change Example (Page 26)

• In a chemical reaction, reactants are converted to new substances; hydrogen and oxygen combine to form water

Determining Chemical Changes (Page 27)

• Which of the following are chemical changes?

  • I. Burning wood – Yes (chemical change)

  • II. Pulverizing rock salt – No (physical change)

  • III. Dissolving sugar in tea – No (physical change, unless reaction occurs)

  • IV. Melting a popsicle – No (physical change)

• Answer choices provided: A B C D (correct option would be the one corresponding to I only, if presented)

Intensive vs Extensive Properties (Page 28)

• Intensive properties: independent of amount of substance

  • Examples: color, density, boiling point

• Extensive properties: depend on amount of substance

  • Examples: mass, volume, energy

Energy in Chemical Processes (Page 30)

• Energy: the ability to transfer heat or to do work

• Kinetic energy: energy of motion; Ek = rac{1}{2} m v^2

  • where m = mass, v = velocity (speed)

• Potential energy: stored energy due to position

Scientific Notation and Number Representation (Pages 32–38)

• Scientific notation: number written as a product where the first factor is between 1 and 10 and the second factor is a power of 10

  • Example: 65,307.2 = 6.53072 imes 10^4

• Standard form to scientific notation and vice versa (examples shown):

  • 6.53072 imes 10^4

  • 3.72 imes 10^{-2}

  • If exponent is negative, move decimal left; if positive, move right

• Operations in scientific notation:

  • Multiplication:

  • 
    (a imes 10^{m}) (b imes 10^{n}) = (ab) imes 10^{m+n}

  • Division:

  • 
    rac{a imes 10^{m}}{b imes 10^{n}} = rac{a}{b} imes 10^{m-n}

  • Addition/Subtraction require same exponent (power of ten) alignment before combining numbers

• Practical examples (from slides):

  • 
    (3.8 imes 10^{4}) (2.0 imes 10^{3}) = 7.6 imes 10^{7}

  • 
    (4.6 imes 10^{6}) (2 imes 10^{-2}) = ?

• Practice problems and multiple-choice prompts are shown throughout, including:

  • 2.3 imes 10^{8} vs 9.2 imes 10^{8} etc.

• Important reminder: when converting, keep track of exponent rules and significant figures during the calculation

SI Units and Metric Prefixes (Pages 39–41, 69–71)

• SI base units and derived units

• Prefixes and their meaning (powers of ten)

  • Peta (P) = 10^{15}

  • Tera (T) = 10^{12}

  • Giga (G) = 10^{9}

  • Mega (M) = 10^{6}

  • Kilo (k) = 10^{3}

  • Deci (d) = 10^{-1}

  • Centi (c) = 10^{-2}

  • Milli (m) = 10^{-3}

  • Micro (μ or mu) = 10^{-6}

  • Nano (n) = 10^{-9}

  • Pico (p) = 10^{-12}

  • Femo (f) = 10^{-15}

  • Atto (a) = 10^{-18}

  • Zepto (z) = 10^{-21}

• Watt and joule relations

  • The watt (W) is the SI unit of power; the rate at which energy is generated or consumed

  • The joule (J) is the SI unit of energy; 1 J = 1 kg·m^2/s^2; 1 W = 1 J/s

  • Note: mu (μ) is the micro prefix; “mu” is used for the micro symbol

Mass, Length, Temperature, Volume, and Density (Pages 42–46)

• Mass

  • Measure of the amount of material in an object

  • SI base unit: kilogram (kg); metric base unit: gram (g)

• Length

  • Measure of distance; base unit: meter (m)

• Temperature

  • Scientific measures use Celsius and Kelvin scales

  • Celsius: 0 °C = freezing point of water; 100 °C = boiling point of water

  • Kelvin: SI unit of temperature; no negative Kelvin values; absolute zero is 0 K

  • Conversion: K = °C + 273.15

• Volume

  • Not a base SI unit; derived from length: ext{volume} = m imes m imes m = m^3

  • Common units: liter (L) and milliliter (mL)

  • 1 cm^3 = 1 mL

• Density

  • Physical property with units derived from mass/volume

  • Common units: g/mL or g/cm^3

• Energy units

  • Joule (J): 1 J = 1 kg·m^2/s^2

  • Calorie (cal): 1 cal = 4.184 J

Measurement Uncertainty, Accuracy, and Precision (Pages 47–49)

• Exact numbers and inexact (measured) numbers

  • Exact numbers are counted or defined (e.g., 12 eggs in 1 dozen)

  • Inexact numbers depend on measurement and instrument precision

• Uncertainty in measurements

  • All measurements have some degree of inaccuracy due to instrument limitations

• Accuracy vs. Precision

  • Accuracy: closeness to true value

  • Precision: closeness among repeated measurements

Significant Figures (Pages 50–57)

• Definition: digits that were measured, used to express measurement precision

• Final digit is always estimated and uncertain

• Counting sig figs

  • Non-zero integers always count as sig figs

  • Zeros rules:

  • Leading zeros do not count

  • Captive zeros (in the middle) count

  • Trailing zeros count if a decimal point is present

  • Exact numbers (e.g., conversion factors) have infinite sig figs

• Examples

  • 3456 has 4 sig figs

  • 0.0486 has 3 sig figs

  • 16.07 has 4 sig figs

  • 9.300 has 4 sig figs; 9,300 has 2 sig figs (decimal point presence matters)

• Practice questions shown in slides

  • Question: How many sig figs in 3270? Answer: 3 (C)

• Using sig figs in operations

  • Multiplication/Division: number of sig figs in result = least number of sig figs in any operand

  • Example: 6.38 imes 2.0 = 12.76
    ightarrow 13 (2 sig figs)

  • Addition/Subtraction: number of decimal places in result = least precise decimal place among operands

  • Example: 6.8 + 11.934 = 18.734
    ightarrow 18.7 (3 sig figs)

• Example problem: 6.578 − 4.5 = ? → 2 (sig figs) or 2.1 depending on decimal placement; see slide for options

Dimensional Analysis (Pages 64–77)

• Purpose: convert one quantity to another using conversion factors

• Core idea: multiply by conversion factors so units cancel, ending with desired unit

• Steps (Dimensional Analysis: How-to)

  • Identify target unit

  • Find a conversion factor relating given unit to target unit

  • Set up expression so units cancel, leaving the desired unit

  • Consider significant figures when counting digits

• Example 1: Convert 3.000 m to cm

  • Conversion factors used: 1 ext{ cm} = 10^{-2} ext{ m}

  • Calculation:

  • Start with 3.000 m

  • Multiply by 1 cm / 10^{-2} m

  • Result: 3.000 ext{ m} imes rac{1 ext{ cm}}{10^{-2} ext{ m}} = 300.0 ext{ cm}

  • Note: Do not count conversion factors when determining sig figs

• Useful conversion factors (Table of equivalencies)

  • Length:

  • 1 ext{ km} = 0.62137 ext{ mi}

  • 1 ext{ mi} = 5280 ext{ ft} = 1.6093 ext{ km}

  • 1 ext{ m} = 1.0936 ext{ yd}

  • 1 ext{ in} = 2.54 ext{ cm}

  • 1 ext{ cm}^3 = 1 ext{ mL}

  • Mass:

  • 1 ext{ kg} = 2.2046 ext{ lb}

  • 1 ext{ lb} = 16 ext{ oz}

  • Volume:

  • 1 ext{ L} = 1000 ext{ cm}^3 = 1.0567 ext{ qt}

  • 1 ext{ gal} = 4 ext{ qt} = 3.7854 ext{ L}

• Example 2: Convert 6.78 cm to inches

  • Use conversions: 1 in = 2.54 cm; 1 in ≈ 2.54 cm

  • Setup to compute: 6.78 cm × (1 in / 2.54 cm) = 2.67 in (approx)

• Example 3: Dimensional analysis for volume

  • Teacup contains 12.0 in³ of tea; convert to mL

  • Step 1: Use 1 in³ = 16.3871 cm³ and 1 cm³ = 1 mL

  • Calculation: 12.0 in³ × 16.3871 cm³/in³ × 1 mL/cm³ = 197 mL

Dimensional Analysis: Two or More Conversions (Pages 75–77)

• When no direct conversion exists, use multiple factors

• Example worked: 3.000 m to in

  • Steps include: 1 cm = 10^{-2} m; 1 in = 2.54 cm

  • Result: 118.1 in

Example: Dimensional Analysis – Volume conversions (Teacup Case) – Revisited (Page 76)

• Start with 12.0 in³

• Use: 1 in³ = 16.3871 cm³ and 1 cm³ = 1 mL

• Final units: mL

• Calculation structure demonstrates cancellation to reach desired unit

Additional Notes: Practice Problems and Quick References

• Practice prompts throughout the slides emphasize:

  • Decision-making on sign figs in operations

  • Correct unit cancellations in dimensional analysis

  • Reading and interpreting the table of metric prefixes

  • Basic conversions between SI units and common units (lb, oz, qt, gal, etc.)

Key Formulas and Equations (Collected)

• Kinetic energy: E_k = frac{1}{2} m v^2

• Temperature conversion: K = °C + 273.15

• Density:
  ho = rac{m}{V}

• Joule definition: 1 ext{ J} = 1 ext{ kg}\, ext{m}^2/ ext{s}^2

• Power and energy relation: 1 ext{ W} = rac{1 ext{ J}}{ ext{s}}

• Conversion factor framework: Use factors like 1 ext{ cm} = 10^{-2} ext{ m}, 1 ext{ in} = 2.54 ext{ cm}, 1 ext{ L} = 1000 ext{ cm}^3 = 1.0567 ext{ qt}

• Sig figs operation rules:

  • Multiplication/Division: number of sig figs in result = least number of sig figs among the factors

  • Addition/Subtraction: decimal places in result = fewest in any term

Connections to Foundational Principles and Real-World Relevance

• Matter states and changes underpin material science, chemical engineering, and materials selection for technology (e.g., solar cells, LEDs, medical plastics)

• Dimensional analysis is a foundational tool for ensuring correct units and magnitudes in calculations across chemistry and physics

• Significant figures reflect measurement precision, crucial in lab contexts and in communicating experimental results

• SI units, prefixes, and conversions enable global standardization in science and engineering

Ethical, Philosophical, and Practical Implications

• Proper reporting of measurements and uncertainty is essential for reproducibility and trust in scientific results

• Clear distinction between physical and chemical properties supports safe material handling and hazard assessment

• Understanding energy, work, and efficiency informs sustainability and technology development (e.g., energy conversion, storage, and ecologically responsible design)

Quick Reference: Common Conversions (Selected)

• Length/Distance:

  • 1\text{ m} = 100\text{ cm}

  • 1\text{ cm} = 10^{-2}\text{ m}

  • 1\text{ in} = 2.54\text{ cm}

  • 1\text{ km} = 1000\text{ m}

  • 1\text{ mi} = 5280\text{ ft} \approx 1.6093\text{ km}

• Mass:

  • 1\text{ kg} = 2.2046\text{ lb}

  • 1\text{ lb} = 16\text{ oz}

• Volume:

  • 1\text{ L} = 1000\text{ cm}^3 = 1.0567\text{ qt}

• Energy and Power:

  • 1\text{ J} = 1\text{ kg m}^2/\text{s}^2

  • 1\text{ W} = 1\text{ J/s}

• Density and Volume examples:

  • Common density units: g/mL or g/cm³

  • 1 cm³ = 1 mL