Example 1.4 Density Calculation" (Answer: ≈ 0.869 g/mL, matching density of toluene) - Interactive Example 1.5: Using density to relate mass and volume: - Problem: isopropyl alcohol, d = 0.785 g/mL; mass needed = 43.7 g; volume to measure: $$V = rac{m}{ ho} = rac{43.7}{0.785} ext{ mL} \

1.1 Modern Chemistry: A Brief Glimpse

• Chemistry defined: science of the composition and structure of materials and the changes they undergo.
• Modern chemistry roots:
  • Ancient technology: fire, ceramics, glass, metals, dyes, medicines.
  • Tyrian purple: ancient dye from sea snail; one ounce required >200,000 snails; later synthesized from simpler molecule (aniline).
• Central principle of modern chemistry:
  • Matter is composed of exceedingly small particles called atoms.
  • The arrangement of atoms into molecules or more complex structures accounts for materials’ characteristics.
  • Ability to synthesize molecules and correlate molecular structure with material properties.
• Key technologies illustrating chemical principles:
  • Liquid-crystal displays (LCDs): rely on alignment of rodlike liquid-crystal molecules to control light; layered alignment via grooves; electrical control changes light transmission; high contrast and millions of colors.
  • Liquid crystals: form a state intermediate between liquids and solid crystals.
• Applications showcase breadth and relevance:
  • Chemistry as a practical tool for lifesaving drugs (e.g., cisplatin family) and broader technology.
  • Chemistry as an intellectual pursuit: seeking chemical explanations for phenomena.
  • Interdisciplinary impact: medicine, biology, physics, environmental science, and technology.
• Biological connection:
  • DNA stores hereditary information; consists of two intertwined chains with four bases; order of bases encodes information similar to characters on a page.
• Everyday context:
  • All matter is made of materials (papers, plastics, rocks, water, biological substances).
• Preview of what’s coming:
  • Atomic theory of matter (to be discussed in Chapter 2).
  • Foundational vocabulary, measurement, and units for quantitative work.

1.2 Experiment and Explanation

• Heart of chemical research: experiment and explanation.
  • Experiment: observation under controlled conditions where variables (e.g., temperature, amounts of substances) can be controlled; results are duplicable.
  • Example: Rosenberg et al. studied electricity’s effect on bacterial growth; platinum electrodes with current stopped cell division, due to a platinum-containing substance produced by the electrode.
• Outcomes of experiments:
  • A regularity or relationship observed → if simple and fundamental, can be stated as a law.
  • Law example: Law of Conservation of Mass: mass remains constant during a chemical change.
• From observation to explanation:
  • A hypothesis: tentative explanation of an observed regularity.
  • If hypothesis passes many tests, it becomes a theory (e.g., molecular theory of gases).
  • Theories are not proven absolutely; new data can refine or replace them (Newtonian mechanics gave way to relativity and quantum mechanics in certain regimes).
• The scientific method flow (general steps):
  • Observation → Hypothesis → Experimental tests → Results → Explanation/Model → New hypotheses and further experiments.
  • Rosenberg’s work illustrates iterative testing: identify platinum compounds, test anticancer activity, refine hypothesis, expand experiments.
• Everyday example of scientific creativity:
  • The Birth of the Post-it Note: accidental adhesive discovery led to a reusable, removable bookmark and later a new office product via iterative experimentation and collaboration.
• Relationship between experiment and explanation:
  • Experiments generate data; explanations organize knowledge and guide new experiments; theory emerges from accumulated, tested explanations.
• Important caveat:
  • The scientific method is not a rigid protocol; creativity and individual approaches drive experimental design and interpretation.

1.3 Law of Conservation of Mass

• Historical development:
  • Balances became a standard tool in eighteenth-century chemistry.
  • Antoine Lavoisier demonstrated that total mass is conserved in chemical changes.
• Law statement:
  • The total mass of substances reacting (reactants) equals the total mass of substances formed (products): m_ ext{reactants} = m_ ext{products}.
• Practical example: combustion of mercury oxide:
  • Mercury + oxygen → mercury(II) oxide (HgO).
  • When HgO is heated, it decomposes back into Hg and O2.
  • Example problem illustration: heating 2.53 g of mercury in air yields 2.73 g of red-orange residue; mass of oxygen that reacted is 0.20 ext{ g} = 2.73 ext{ g} - 2.53 ext{ g}.
• Verification technique:
  • Arithmetic check: 2.53 g Hg + 0.20 g O2 should equal 2.73 g HgO when combined.
• Related concept: mass vs weight
  • Weight is the force of gravity on a mass; weight varies with location due to gravity; mass is invariant.
• Conceptual note:
  • Law applies to chemical reactions; measurement precision is essential for confirming mass balance.

1.4 Matter: Physical State and Chemical Composition

• Two main classifications of matter:
  • Physical state: solid, liquid, gas.
  • chemical composition: element, compound, mixture.
• States of matter definitions:
  • Solid: rigid, relatively incompressible, fixed shape and volume.
  • Liquid: relatively incompressible; fixed volume but no fixed shape.
  • Gas: easily compressible; fills container; highly fluid.
• Vapor: for the gaseous state of matter that normally exists as a liquid or solid.
• Physical vs chemical changes and properties:
  • Physical change: change in form, not chemical identity (e.g., phase changes, dissolving).
  • Chemical change: one or more kinds of matter transformed into new kinds of matter (e.g., rusting).
  • Physical properties: can be observed without changing chemical identity (e.g., melting point, color).
  • Chemical properties: involve chemical change (e.g., reactivity with oxygen to form rust).
• Substances vs mixtures:
  • Substance: a kind of matter that cannot be separated into other kinds by physical processes (e.g., pure water, sodium chloride).
  • Mixture: a material that can be separated by physical means into two or more substances; composition can vary.
• Types of mixtures:
  • Heterogeneous: physically distinct parts (e.g., potassium dichromate crystals with iron filings; salt and sugar mixed visibly).
  • Homogeneous (solution): uniform composition (e.g., sodium chloride in water; air as a gaseous solution).
• Phase concept:
  • A phase is a homogeneous part of a system; a sample can contain multiple phases (e.g., ice in water; ice in saline solution).
• Nonstoichiometric compounds: some compounds do not follow definite proportions; discussed briefly ( Chapter 11 references).
• Separation and analysis techniques:
  • Distillation: physical separation of components based on different volatilities (e.g., separating water from sodium chloride).
  • Chromatography: separation based on differential movement through a stationary phase; various forms include paper chromatography and gas chromatography.
  • Gas chromatography (GC): rapid separation using gas as mobile phase and a solid/liquid stationary phase; retention time helps identify substances; chromatograms show peaks corresponding to components (e.g., chocolate contains >800 flavor compounds).
• DNA and biology context:
  • DNA represents a universal molecular information carrier across life forms.
• Conceptual map (Figure 1.16): relationships among elements, compounds, and mixtures; processes connect these categories via physical or chemical changes.
• Concept Check 1.1 (visual model exercise): represent matter as units (elements, compounds, mixtures) using a simple schematic.
• Additional notes:
  • Distinguish between substances and mixtures when considering separation methods and properties.

1.5 Measurement and Significant Figures

• Why measurements matter: chemists characterize substances via physical measurements (mass, volume, temperature, etc.).
• Precision vs accuracy:
  • Precision: closeness of multiple measurements to each other.
  • Accuracy: closeness of a measurement to the true value.
• Example of precision:
  • A simple rod measured with a centimeter ruler might be read as 9.12 cm, with repeated readings 9.11 cm and 9.13 cm; the measurement is between 9.11 and 9.13 cm, indicating precision.
• Significant figures concept:
  • Significant figures are digits in a measured or calculated value that include all certain digits plus a final digit with some uncertainty.
  • Example: 9.12 cm has three significant figures; 9.120 cm has four, but only if the trailing zero is meaningful (decimal point is present).
• Rules for counting significant figures:
  • All digits are significant except leading zeros and possibly trailing zeros not indicated by a decimal point.
  • Trailing zeros to the right of the decimal point are significant (e.g., 9.00 cm has three sig figs).
  • Trailing zeros without a decimal point may or may not be significant (e.g., 900 cm is ambiguous); scientific notation clarifies precision (e.g., 9.0 × 10^2 cm vs. 9.00 × 10^2 cm).
• Scientific notation:
  • Representation: A imes 10^n where A has one nonzero digit to the left of the decimal point.
  • Examples: 9.0 × 10^2 cm (two significant figures in the mantissa imply two sig figs), 3.00 × 10^8 m/s (three sig figs).
• Significant figures in calculations:
  • Multiplication and division: the result should have as many significant figures as the operand with the least number of sig figs.
  • Addition and subtraction: the result should have the same number of decimal places as the value with the least decimal places.
• Example: solubility calculation for cisplatin:
  • Given: solubility = 0.0634 g in 25.31 g water; extrapolate to 100.0 g water.
  • Calculation: rac{0.0634 ext{ g}}{25.31 ext{ g}} imes 100.0 ext{ g} = 0.250493875 ext{ g}
  • Final report: 0.250 g (three significant figures, because 0.0634 g has three sig figs).
• Addition example: 184.2 g + 2.324 g = 186.524 g; least decimal places is one (184.2 has one decimal place), so the result is 186.5 g.
• Exact numbers:
  • Exact numbers arise from counting (e.g., defined constants, or counted items) and have infinite significant figures for practical purposes.
• Exercise/Problems: problems 1.61, 1.62, etc., reinforce significant figures rules and rounding conventions.

1.6 SI Units

• Historical context: early measurement used body-based units; lack of standard units.
• SI and the metric system:
  • In 1791, the metric system established; in 1960, the General Conference on Weights and Measures adopted the SI system with seven base units.
• SI base units (four base quantities discussed here):
  • Length: meter (m)
  • Mass: kilogram (kg)
  • Time: second (s)
  • Temperature: kelvin (K)
  • Also relevant base quantities in broader context: amount of substance (mol), electric current (ampere, A), luminous intensity (candela, cd)
• SI prefixes (Table 1.2):
  • Mega (M) = 10^6, kilo (k) = 10^3, deci (d) = 10^-1, centi (c) = 10^-2, milli (m) = 10^-3, micro (μ) = 10^-6, nano (n) = 10^-9, pico (p) = 10^-12, etc.
• Length and time units:
  • The meter (m) is the SI base unit of length; small-length units include nanometer (nm = 10^-9 m) and picometer (pm = 10^-12 m).
  • The angstrom (Å) is a non-SI unit equal to 10^-10 m; often used for atomic-scale distances (e.g., oxygen atom diameter ~ 1.3 Å).
  • The second (s) is the SI base unit of time; with prefixes, allows measurement of very fast events (e.g., picoseconds).
• Kilogram and related mass units:
  • The kilogram is the SI base unit of mass; although it is a unit with a prefix in its name, it is a base unit.
  • Other mass units formed by applying prefixes to gram (e.g., mg = 10^-3 g).
• Temperature scales and conversions:
  • Celsius (°C) and Kelvin (K) are tied; conversion: K = °C + 273.15
  • Fahrenheit (°F) conversion: F = rac{9}{5}°C + 32 and °C = rac{5}{9}(F - 32)
• Temperature examples and practical notes:
  • Room temperature ~ 20°C ≈ 293 K.
  • 0°C = 273.15 K; 100°C = 373.15 K; 0°C = 32°F; 100°C = 212°F.
• Exercise 1.4: converting numbers to SI prefixes to express in appropriate units.
• Temperature in context: practical problem-solving often requires converting between Kelvin, Celsius, and Fahrenheit.

1.7 Derived Units

• Concept: once base units exist, other units are derived from them via definitions/relations.
• Common derived units (Table 1.3):
  • Area: A = ext{length}^2
    ightarrow ext{Unit: } ext{m}^2
  • Volume: V = ext{length}^3
    ightarrow ext{Unit: } ext{m}^3
  • Density:
    ho = rac{m}{V}
    ightarrow ext{Unit: } rac{ ext{kg}}{ ext{m}^3} (also expressed as g/cm^3 in many contexts)
  • Speed: v = rac{ ext{distance}}{ ext{time}}
    ightarrow ext{Unit: } rac{ ext{m}}{ ext{s}}
  • Acceleration: a = rac{ ext{velocity}}{ ext{time}}
    ightarrow rac{ ext{m}}{ ext{s}^2}
  • Force: F = m a
    ightarrow ext{Unit: } ext{N} = ext{kg m s}^{-2}
  • Pressure: P = rac{F}{A}
    ightarrow ext{Unit: } ext{Pa} = rac{ ext{N}}{ ext{m}^2} = rac{ ext{kg}}{ ext{m s}^2}
  • Energy: E = F d
    ightarrow ext{Unit: } ext{J} = ext{kg m}^2 ext{s}^{-2}
• Volume and density in practice:
  • 1 L = 1 dm^3 = 1000 cm^3; 1 mL = 1 cm^3.
  • Density examples:
  • Water: ≈ 1.000 g/cm^3 at 4°C; ≈ 0.998 g/cm^3 at 20°C.
  • Lead: ≈ 11.3 g/cm^3 at 20°C.
  • Oxygen gas density: ≈ 1.33 × 10^-3 g/cm^3 at normal pressure and 20°C (≈ 1.33 g/L).
• Density usage and examples:
  • Density as a physical property helps identify substances and assess purity.
  • Density-based reasoning: E.g., a gold bar’s density helps assess purity vs. impurities.
• Example 1.4: Calculating density from mass and volume:
  • Given m = 30.5 g, V = 35.1 mL; density $$
    ho = rac{m}{V} = rac{30.5}{35.1} ext{ g/mL} \