Week 1 Notes: Chemistry in Context and Fundamentals

1.1 Chemistry in Context

• Chemistry is relevant to daily life and everyday activities (e.g., making coffee, cooking eggs, toasting bread).
  • Products we use (soap, shampoo, fabrics, electronics, gasoline) involve chemical substances and processes.
  • Chemistry is part of art and culture: cleaning, conserving, restoring art; authenticating artifacts; polymers for 3D printing; dyes and pigments.
  • Occupational health and safety is a common career area for chemists due to hazardous materials in art supplies, industry, etc.
• Chemistry and art are interconnected: artworks and conservation rely on chemical principles; chemistry underpins authentication and dating techniques.
• Week 1 objectives (end of module):
  • Outline the historical development of chemistry.
  • Provide examples of the importance of chemistry in everyday life.
  • Describe the scientific method.
  • Differentiate among hypotheses, theories, and laws.
• Historical development overview:
  • Early humans manipulated matter by changing shape (e.g., flint tools, carved wood).
  • With knowledge growth, humans began changing composition: clay to pottery; hides to garments; ores to metals; grain to bread.
  • Chemistry emerged with fire control (cooking, pottery, smelting metals).
  • Isolation of drugs from natural sources; dyes from plant/animal matter; formation of alloys (e.g., Cu–Sn to bronze); smelting to iron.
  • Alkalis from ashes; soaps from fats and alkalis; fermentation and distillation to produce alcohol.
  • As early as the 6th century BCE, Greeks proposed water as a basis of matter and elements concept (earth, air, fire, water).
  • Alchemists attempted to transform base metals into noble metals and to create elixirs, contributing to later chemical technologies; not scientific by modern standards.
• From alchemy to modern chemistry:
  • Isolation of drugs, metallurgy, and the dye industry are historical progressions.
  • Modern chemistry is central to many STEM fields and interacts with biology, medicine, materials science, forensics, environmental science, astronomy, and more.
  • Chemistry is referred to as a central science because of its connections to physics, mathematics, computer science, biology, etc.
  • Interdisciplinary connections include: chemical physics, nuclear chemistry, biochemistry, chemical engineering, materials science, nanotechnology, agriculture, food science, veterinary science, brewing and wine making, medicine, pharmacology, biotechnology, botany, environmental science, geology, oceanography, atmospheric science, and astronomy.
• The central idea: chemistry helps explain and control matter and its interactions, underpinning both theoretical frameworks and practical applications in society.

1.1 The Scientific Method and Key Concepts

• Chemistry is based on observation and experimentation; questions are answered in terms of chemical laws and theories via accepted procedures.
• There is no single route to answering a question; all approaches rely on reproducible experiments.
• Core components of the scientific method include:
  • Hypothesis: a tentative explanation guiding data gathering and testing.
  • Experimentation, calculation, and comparison with others’ experiments to verify results.
  • Iterative refinement of hypotheses.
  • Relationship between hypotheses, laws, and theories: laws summarize many observations; theories are well-substantiated, comprehensive explanations that can be modified with new data.
• The path from observation to law or hypothesis to theory, with experimental verification, constitutes the scientific method.
• Scientific progress is non-linear and involves open inquiry and reworking questions/ideas in response to findings.
• Figure references: the central science diagram emphasizes the interconnectedness of chemistry with many fields; the scientific method diagram outlines components and order.

1.2 Phases and Classification of Matter

• By the end of this section you should be able to:
  • Describe the basic properties of solid, liquid, and gas.
  • Classify matter as an element, compound, homogeneous mixture, or heterogeneous mixture with respect to its state and composition.
  • Define and give examples of atoms and molecules.
• Matter is anything that occupies space and has mass; solids, liquids, and gases are the three common states; plasma is a fourth state.
• States of matter:
  • Solid: rigid, definite shape.
  • Liquid: flows, takes shape of container, flat or curved surface due to gravity.
  • Gas: takes both shape and volume of its container; expands to fill space.
  • Plasma: a high-temperature state with many charged particles; found in stars and other high-energy environments (e.g., lightning, some TV screens, analytical instruments).
• Special notes:
  • In zero gravity, liquids tend to form spheres due to lack of gravity flattening; sand can behave like a liquid when poured; clouds are mixtures (gas + tiny liquid/solid water particles).
• Mass and weight:
  • Mass measures the amount of matter; weight is the force of gravity on an object: $W = m g$ where g is local gravitational acceleration.
  • Mass is constant; weight varies with gravity (e.g., on the Moon, weight is ~1/6 of Earth’s, while mass remains unchanged).
• Classification of matter:
  • Pure substances have fixed composition; two main classes: elements and compounds.
  • Elements: cannot be broken down by chemical changes (e.g., Fe, O, C, S).
  • Compounds: can be decomposed into elements or other compounds via chemical changes (e.g., Mercury(II) oxide decomposes to Hg and O2; sucrose decomposes under certain processes to C and H2O).
  • Mixtures: combinations of two or more types of matter in varying proportions; can be separated by physical changes.
  • Heterogeneous mixtures: composition varies throughout (Italian dressing, granite, chocolate chip cookies).
  • Homogeneous mixtures (solutions): uniform composition (sports drinks, air, gasoline, salt in water).
• Key examples and implications:
  • Sucrose composition by mass: 42.1% C, 6.5% H, 51.4% O (by mass).
  • Sucrose and water can be formed by chemical combination of elements; sugar is a compound; white sodium chloride is formed from Na and Cl.
  • Silver chloride (AgCl) is a compound whose decomposition can be induced by light, useful in photography and photochromic lenses.
  • A sample of matter can show properties of more than one phase (e.g., sand acts like a liquid when poured; clouds are gas with dispersed particles).
• Earth’s elemental composition (approximate):
  • Oxygen ~ 49.20%; Silicon ~ 25.67%; Aluminum ~ 7.50%; Iron ~ 4.71%; Calcium ~ 3.39%; Sodium ~ 2.63%; Potassium ~ 2.40%; Magnesium ~ 1.93%; Phosphorus ~ 0.11%; Manganese ~ 0.09%; Others ~ 0.47% (Table 1.1 summary).
• Atoms and molecules:
  • An atom is the smallest unit of an element that retains its properties; e.g., a gold atom is the fundamental unit of gold.
  • Dalton’s atomic theory provided quantitative support and formed the basis for modern chemistry; although some details have evolved, the concept of atoms remains central.
  • A molecule is two or more atoms bonded together; molecules can be of the same element (O2, N2) or different elements (H2O, C6H12O6).
  • Some elements exist as diatomic molecules in the gaseous state (H2, N2, O2, Cl2, etc.), while others form more complex units (P4, S8).
• Size and mass scale highlights:
  • An ordinary spider web strand is about 0.0001 cm in diameter; a carbon atom is ~0.000000015 cm in diameter; to span the spider web with carbon atoms would require about 7,000 atoms across.
  • A dime-sized carbon atom would imply a cross-section vastly larger than familiar macroscopic scales; this illustrates the enormous size difference between macroscopic and atomic scales.
• Useful population of units and prefixes (Table 1.2 and Table 1.3):
  • SI base units: length (meter, m); mass (kilogram, kg); time (second, s); temperature (kelvin, K); electric current (ampere, A); amount of substance (mole, mol); luminous intensity (candela, cd).
  • Common prefixes (Table 1.3):
  • femto 10^{-15}, pico 10^{-12}, nano 10^{-9}, micro 10^{-6}, milli 10^{-3}, centi 10^{-2}, deci 10^{-1}
  • kilo 10^3, mega 10^6, giga 10^9, tera 10^{12}
• SI base unit definitions and examples:
  • Length: meter (m) – defined by light travel in vacuum; 1 m ≈ 39.37 in; 1 km = 10^3 m; 1 cm = 10^{-2} m; 1 mm = 10^{-3} m.
  • Mass: kilogram (kg) – originally mass of a liter of water; now defined by a platinum-iridium cylinder held in France; 1 kg ≈ 2.2046 lb; 1 g = 10^{-3} kg.
  • Temperature: kelvin (K); Celsius is allowed; 0 °C = 273.15 K; 100 °C = 373.15 K.
  • Time: second (s); units scale with prefixes (µs, Ms, etc.).
• Volume and density (Derived SI units):
  • Volume: derived from length; 1 m^3 is a cube of edge 1 m; 1 L = 1 dm^3; 1 L ≈ 1.06 quarts; 1 cm^3 = 1 mL.
  • Density: mass per volume; SI unit: kg/m^3; common practical units: g/cm^3 for solids/liquids; g/L for gases.
  • Typical densities (examples):
  • Water: 1.0 g/cm^3
  • Ice (at 0 °C): 0.92 g/cm^3
  • Air (dry): 1.20 g/L
  • Iron: 7.9 g/cm^3
  • Gold: 19.3 g/cm^3
  • Helium: 0.000178 g/L (approx)
  • Density calculation example: density = mass/volume; example calculation:
  • volume of lead cube with edge 2.00 cm: $V = 2.00\text{ cm} \times 2.00\text{ cm} \times 2.00\text{ cm} = 8.00\text{ cm}^3$
  • density: $\rho = \frac{m}{V} = \frac{90.7\text{ g}}{8.00\text{ cm}^3} = 11.3\text{ g cm}^{-3}$
• Table 1.4 highlights include the densities listed above for common substances (solids, liquids, gases) to show the wide range in densities.

1.3 Physical and Chemical Properties

• Distinguishing properties is the core idea:
  • Physical properties: characteristics not involving a change in chemical composition (e.g., density, color, hardness, melting/boiling points, electrical conductivity).
  • Some physical properties can be observed without changing state (e.g., density, color); others are observable only during a physical change (e.g., melting temperature, boiling temperature).
  • Physical changes: changes in state or properties without changing the chemical composition (e.g., wax melting, sugar dissolving, steam condensing).
• Chemical properties and changes:
  • Chemical property: a substance’s ability to undergo a chemical change (e.g., flammability, toxicity, acidity, reactivity).
  • Examples: iron rusts in the presence of oxygen and water; chromium does not oxidize under the same conditions (as a comparison of chemical reactivity).
  • Chemical changes produce new kinds of matter (e.g., rust forms a different substance than iron, oxygen, and water initially present).
  • Examples of chemical changes include reactions in the lab (copper with nitric acid), all combustion, cooking, digestion, and rotting (Fig. 1.20).
• Elements, compounds, and mixtures (revisited):
  • Elements: pure substances that cannot be broken down into simpler substances by chemical changes.
  • Compounds: pure substances that can be broken down into simpler substances (elements or other compounds) by chemical changes.
  • Mixtures: combinations of two or more substances that retain their own properties; can be separated by physical means.
• Examples and illustrations:
  • Rust is a chemical change forming different matter; nitroglycerin is a chemical property/behavior related to explosive potential.
  • Examples of chemical changes: copper + nitric acid -> copper nitrate + nitrogen dioxide gas; combustion of a match (cellulose + oxygen -> CO2 + H2O); meat cooking changes color due to oxidation of myoglobin; banana browning due to chemical changes.
• Periodic table and element types:
  • Metals: good conductors of heat and electricity.
  • Nonmetals: poor conductors.
  • Metalloids: intermediate conductivities.
  • The periodic table groups elements by similar properties; color coding often used to indicate metal/metalloid/nonmetal and state at room temperature (solid/liquid/gas).

1.4 Measurements

• The role of measurements in chemistry:
  • Measurements provide quantity (magnitude), units (standard of comparison), and uncertainty (precision/accuracy context).
  • Numbers can be represented in decimal or scientific notation:
  • Example: maximum takeoff weight of a Boeing 777-200ER: $2.98\times 10^{5}\text{ kg}$
  • Mosquito mass: $2.5\times 10^{-6}\text{ kg}$
• SI base units (Table 1.2):
  • Length: meter (m)
  • Mass: kilogram (kg)
  • Time: second (s)
  • Temperature: kelvin (K)
  • Electric current: ampere (A)
  • Amount of substance: mole (mol)
  • Luminous intensity: candela (cd)
• Derived and common units:
  • Volume: derived from length; 1 m^3; 1 dm^3 = 1 L; 1 L ≈ 1.06 quarts; 1 cm^3 = 1 mL.
  • Density: mass per volume; SI unit kg/m^3; common units g/cm^3 for solids/liquids and g/L for gases.
• Unit prefixes (Table 1.3) and scaling:
  • 10^{-15} to 10^{12} scale examples: 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}).
• Base unit definitions and examples:
  • Meter: defined by light in vacuum; 1 m ≈ 39.37 inches; 1 km = 10^3 m; 1 cm = 10^{-2} m; 1 mm = 10^{-3} m.
  • Kilogram: defined by a platinum-iridium cylinder (IPK); 1 kg ≈ 2.2046 lb; 1 g = 10^{-3} kg.
  • Temperature: Kelvin (K); 0 °C = 273.15 K; 100 °C = 373.15 K.
  • Time: second (s).
• Volume and density (Derived SI Units, continued):
  • Density: $\rho = \frac{m}{V}$ with SI unit in kg/m^3; common practical units: $\text{g cm}^{-3}$ for solids and liquids; $\text{g L}^{-1}$ for gases.
• Densities of common substances (examples from Table 1.4):
  • Solids: ice 0.92 g/cm^3; oak wood 0.60–0.90 g/cm^3; iron 7.9 g/cm^3; copper 8.96 g/cm^3.
  • Liquids: water 1.0 g/cm^3; ethanol 0.79 g/cm^3; glycerin 1.26 g/cm^3; acetone 0.79 g/cm^3.
  • Gases: dry air 1.20 g/L; nitrogen 1.14 g/L; helium 0.18 g/L approx; neon 0.83 g/L; radon 9.1 g/L (approx).
• Density calculation example (density demonstration):
  • Problem: density of lead cube with edge 2.00 cm and mass 90.7 g.
  • Solution:
  • Volume: $V = (2.00\text{ cm})^3 = 8.00\text{ cm}^3$
  • Density: $\rho = \frac{m}{V} = \frac{90.7\text{ g}}{8.00\text{ cm}^3} = 11.3\ \text{g cm}^{-3}$

1.5 Accuracy and Precision

• Definitions:
  • Precision: closeness of repeated measurements to each other.
  • Accuracy: closeness of a measurement to the true or accepted value.
  • These concepts can be illustrated with archery examples:
  • (a) Accurate and precise: arrows near bull's eye and near each other.
  • (b) Precise but not accurate: arrows cluster together but far from target.
  • (c) Neither precise nor accurate: scattered results far from target.
• Practical example (Table 1.5): volume delivery by three dispensers (each tested five times):
  • Dispenser #1: all values around 283–284 mL; precise but not accurate (target 296 mL).
  • Dispenser #2: values within ~3 mL of 296 mL but spread more widely; accurate but less precise.
  • Dispenser #3: all values within ~0.1–0.2 mL of 296 mL; both accurate and precise.
• Takeaway:
  • Accuracy relates to closeness to the true value; precision relates to reproducibility.
  • In quality control, both high accuracy and high precision are desirable for reliable measurements.