Chapter 1 Notes — Matter, Energy and Measurement

Chapter 1: Matter, Energy and Measurement (Notes)

• Focus areas of Chapter 1:
  • Introduction to science and chemistry
  • Classifications of matter
  • Properties of matter
  • Units of measurement
  • Uncertainty in measurement
  • Dimensional analysis (unit conversions)

What is Chemistry?

• Chemistry is the study of matter.
• Origins debated across cultures:
  • Egyptian: chemia – from alchemia meaning “land of rich black soil”
  • Greek: chemia – from Oxford English Dictionary meaning “pouring or infusion”
• History of chemistry:
  • Ancient times: use of fire and stones (Stone Age ~2.6 million years ago), metallurgy (Bronze Age ~3300 B.C.), Iron Age (~1200 B.C.)
  • Chemistry and alchemy intermingled for centuries; focus on philosopher’s stone and transmutation
  • Modern chemistry: Robert Boyle in the 17th century; The Sceptical Chymist (1661) redefined study of matter

Chemistry – The Study of Matter

• Chemistry involves the properties and behavior of matter.
• If it exists, chemistry is involved; if you can buy it, a chemist was involved somewhere.
• Submicroscopic world vs macroscopic world:
  • Atoms & molecules are usually too small to observe directly (submicroscopic)
  • Macroscopic world: objects we can see and measure
• Water molecule: example of a common molecule composed of atoms

Matter: Definitions and Building Blocks

• Matter: the physical material of the universe; anything with mass and that takes up space
• Matter is composed of atoms and molecules
• Atoms: infinitesimally small building blocks of matter
• Molecules: two or more atoms joined together (same or different)
• Elements: substances made of a unique kind of atom
• Key takeaway: different elements have unique atoms; compounds are formed from atoms of two or more elements

Elements and the Periodic Table

• The periodic table organizes elements by properties and atomic structure
• Categories and sample elements (illustrative, not exhaustive):
  • Alkali metals (e.g., Li, Na, K)
  • Alkaline earth metals (e.g., Be, Mg, Ca)
  • Transition metals (e.g., Fe, Cu, Zn)
  • Nonmetals (e.g., H, C, N, O, Cl)
  • Lanthanides & Actinides (series in separate blocks)
• Atomic data commonly shown: atomic number, symbol, relative atomic mass (numerical values shown on charts in class materials)

States and Classifications of Matter

• Classifications of matter (three basic states):
  • Gas: no definite shape or volume; easy to compress; particles far apart
  • Liquid: definite volume, takes shape of container; hard to compress; particles medium distance apart
  • Solid: definite shape and volume; hard to compress; particles close together
• Pure Substances vs Mixtures:
  • Pure Substance: same properties and composition throughout; all particles are the same; two types:
  • Elements: cannot be decomposed into simpler substances; consist of one type of atom; found on the periodic table
  • Compounds: composed of atoms from more than one element; can be broken down chemically (e.g., H2O, CH4, NaCl)
  • Mixtures: composed of more than one type of particle (more than one element or compound); two types:
  • Heterogeneous: not uniform throughout (e.g., cement, wood, salad)
  • Homogeneous: uniform throughout (e.g., sugar water, air, bronze)

Law of Constant Composition (Definite Proportions)

• Proclaimed by Joseph Proust (1754–1826): Law of Definite Proportions
• A pure substance has a fixed elemental composition
• Examples:
  • Water is always H2O (2 H, 1 O)
  • Methane is always CH4 (1 C, 4 H)
  • Sulfuric acid is always H2SO4 (2 H, 1 S, 4 O)
• Notation examples: H2O, CH4, H2SO4

Mixtures in Detail

• Mixtures consist of more than one type of particle; components retain their own properties
• Separation of mixtures is possible because components have different properties
• Separation methods discussed: filtration, distillation, chromatography

Visualizing the Differences in Matter

• Visual distinctions between:
  • Atoms of an element
  • Molecules of an element
  • Molecules of a compound
  • Mixtures: mixture of elements, mixtures of atoms, elements and compounds
  • Pure substances vs mixtures

Review: Classifications (Practice Questions)

• Classify items as pure substance or mixture; element or compound; homogeneous or heterogeneous:
  • Water: pure substance; compound; homogeneous
  • Nitrogen: pure substance; element; homogeneous
  • Soft drink: mixture; heterogeneous or homogeneous depending on composition; usually homogeneous
  • Diamond: pure substance; element (if elemental diamond) or compound in some cases; homogeneous depending on form
  • Gold (14k): mixture (alloy); homogeneous

Properties of Matter

• Properties distinguish substances; two main types:
  • Physical Properties: observed without changing the substance into another substance
  • Examples: boiling point, density, mass, volume, color, hardness
  • Chemical Properties: observed only when the substance is changed into another substance
  • Examples: flammability, corrosiveness, reactivity with acids
• Physical properties can be extensive or intensive

Intensive vs Extensive Physical Properties

• Intensive properties: do not depend on the amount of substance present
  • Examples: density, boiling point, temperature, color, brittleness
• Extensive properties: depend on the amount of substance present
  • Examples: mass, volume, energy, length

Changes of Matter

• Physical Changes: do not change the composition of a substance
  • Examples: changes of state (melting, boiling), temperature change, volume change
• Chemical Changes (Reactions): create new substances
  • Examples: combustion, oxidation, decomposition
• Everyday example: cooking eggs (chemical change) in heat

Physical vs Chemical Changes: Quick Comparison

• Physical: changes in state; no new substance formed
• Chemical: new substances formed; bonds broken/formed

Sequences of States (Illustrative)

• Physical changes vs chemical changes chart (examples):
  • Water (H2O) changes state: physical
  • Hydrogen gas (H2) reacting with oxygen (O2) to form water: chemical

Chemical Reactions (Basics)

• In chemical reactions, reactants are converted to products
• Example: hydrogen gas + oxygen gas → water
  • Reactants: H2 + O2
  • Products: H2O
• Stoichiometry concepts are introduced in practice problems later in the course

Separation of Mixtures (Techniques)

• Filtration: separates solids from liquids
• Distillation: separates homogeneous mixtures based on differences in boiling points
  • Example: salt water (NaCl in water); water boils at 100°C; common salt boils at a much higher temperature (~1465°C)
• Chromatography: separation based on differential adherence to a solid (stationary phase) by components in a mixture
  • Thin Layer Chromatography (TLC) is a common form

Chromatography: How it Works (Conceptual)

• A mixture is carried by a solvent (mobile phase) through an adsorbent (stationary phase)
• Different compounds interact with the stationary phase to different extents
• Outcome: separation of components as they travel at different rates

SI Units and Base Quantities

• SI (Système International) Base Units:
  • Mass: kilogram, with symbol kg
  • Length: meter, symbol m
  • Time: second, symbol s or sec
  • Temperature: Kelvin, symbol K
  • Amount of substance: mole, symbol mol
  • Electric current: ampere, symbol A or amp
  • Luminous intensity: candela, symbol cd
• SI prefixes (used to scale base units):
  • Prefixes for larger units: kilo (10^3), hecto (10^2), deka (10^1)
  • Prefixes for smaller units: deci (10^-1), centi (10^-2), milli (10^-3)
• The chart also includes thousands of other prefixes (micro, nano, etc.) used in science texts

Metric System Prefixes (Practical Use)

• Converting units by moving the decimal point:
  • From base to a larger unit: move decimal left by the number of prefix steps
  • From base to a smaller unit: move decimal right by the number of prefix steps
• Example prefixes (core set):
  • 1 km = 1000 m; 1 m = 100 cm; 1 cm = 10 mm
  • 1 L = 1000 mL; 1 mL = 1 cm^3

Practice Unit Conversions and Problems

• Practice problems illustrate simple conversions between units and applying prefix rules
• Examples include converting g to kg, cm to m, mL to L, etc.

Temperature Scales and Conversions

• Kelvin (K) as SI unit for temperature; 0 K is absolute zero (no negative temperatures on Kelvin scale)
• Absolute zero: 0 K = -273.15 °C = -459.67 °F
• Temperature conversions:
  • Celsius to Kelvin: $K = °C + 273.15$
  • Fahrenheit to Celsius: $°C = \frac{5}{9}(°F - 32)$
  • Celsius to Fahrenheit: $°F = \frac{9}{5}°C + 32$
• Common reference points (for intuition):
  • 32 °F ≈ 0 °C
  • 65 °F ≈ 18 °C
  • 77 °F ≈ 25 °C
  • 100 °F ≈ 38 °C
  • 212 °F = 100 °C

Mass, Volume, and Density

• Mass: amount of matter; common SI unit is the kilogram (kg); grams (g) and milligrams (mg) also used
• Volume: amount of space occupied; common units are liter (L) and milliliter (mL)
  • 1 L is a cube 1 dm on each side
  • 1 mL is a cube 1 cm on each side
  • SI unit for volume is cubic meter (m^3)
• Density: mass per unit volume; an intensive property; formula:
  • $\rho = \frac{m}{V}$
  • Typical density units: kg/m^3, g/mL
• Density example problem: If a sample weighs 45.4 g and displaces 40.9 mL of water, density is calculated as $\rho = \frac{m}{V} = \frac{45.4\,\text{g}}{40.9\,\text{mL}}$

Scientific Notation

• Purpose: express very large or very small numbers concisely
• Form: $a \times 10^{n}$ where a is the mantissa and n is the exponent
• Examples:
  • 6.02 × 10^23
  • 0.000000000023 = 2.3 × 10^{−11}
  • 475,000,000 = 4.75 × 10^{8}
• Converting between standard and scientific notation:
  • Positive numbers < 1: negative exponent
  • Positive numbers > 1: positive exponent

Significant Figures

• Definition: the number of digits that carry meaning contributing to precision
• The last digit in a measurement is uncertain
• Counting numbers have no uncertainty
• Rules of thumb:
  • All nonzero digits are significant
  • Zeros between significant digits are significant
  • Leading zeros are not significant
  • Zeros at the end of a number are significant if a decimal point is written
• Examples (significant figures):
  • 102 has 3 sig figs
  • 600.007 has 6 sig figs
  • 1.705 × 10^−3 has 4 sig figs
  • 0.124 has 3 sig figs
  • 0.003807 has 4 sig figs (3.807 × 10^−3)
• Calculations with sig figs:
  • Add/subtract: round to the fewest number of decimal places among the terms
  • Multiply/divide: round to the fewest number of significant figures among the terms
  • For multi-step calculations, keep extra sig figs during calculation and round only at the end
• Important note: conversion factors are exact and do not limit sig figs

Uncertainty in Measurements

• All measurements have some uncertainty; the last digit is estimated
• Differences between exact and inexact numbers:
  • Counted numbers (e.g., number of people) are exact
  • Defined numbers (e.g., 360 degrees in a circle) are exact
• The level of uncertainty is tied to the number of significant figures
• Example concept: “the 7 in 27 °C is estimated, so it has uncertainty”
• Different measuring devices have different accuracies; device selection affects uncertainty
• Examples of instruments (from the image): pipette, volumetric flask, graduated cylinder, stopcock, burette, syringe
  • These have varying degrees of precision for delivering/measuring volumes

Uncertainty and Significant Figures: Practical Implications

• The more digits shown, the less uncertainty in the reported value
• When reporting calculated results, reflect uncertainty through appropriate sig figs
• Dimensional analysis and unit consistency help manage uncertainty by ensuring correct unit cancellations

Dimensional Analysis (Conversions)

• Also called factor-label method or unit analysis
• Core idea: multiply by unit conversion factors so that units cancel to give the desired units
• Steps:
  • Identify what is given and what is desired
  • Choose conversion factors so that units cancel appropriately
  • Multiply through and simplify
• Key rule: conversion factors are chosen so that your undesired units cancel
• Example pattern (as taught): Set up the math factors with units on top and bottom such that the desired units remain at the end

Practice Problems and Key Calculations

• Practice problems emphasize converting units with correct significant figures and proper rounding rules
• Example conversions to practice: g to kg, cm to m, mL to L, s to hours or years
• Dimensional analysis practice: converting 8.18 × 10^−3 L to nL (nanoliters)
• Additional practice: volume of a cube with side length in inches; using 1 inch = 2.54 cm
• Time conversions: seconds in 2.5 years; using 1 year = 365 days, 1 day = 24 h, 1 h = 3600 s

Key Concepts for Chapter 1 (Summary)

• Definition of matter and energy
• Types and states of matter (solid, liquid, gas)
• Pure substances vs mixtures (elements vs compounds; homogeneous vs heterogeneous)
• Properties of matter (physical vs chemical; intensive vs extensive)
• Base SI units and common metric prefixes
• Temperature, volume, density, and energy concepts
• Metric system and prefixes for scaling units
• Scientific notation and significant figures concepts and rules
• Energy and units (general mention; see course materials for specific units)
• Dimensional analysis (conversions) as a fundamental tool in chemistry

Quick Reference Formulas and Facts

• Density: $\rho = \frac{m}{V}$
• Mass units: kg, g, mg
• Volume units: L, mL; 1 L = 1000 mL; 1 L = (10 cm)^3; 1 mL = 1 cm^3
• Temperature conversions:
  • $K = °C + 273.15$
  • $°C = \frac{5}{9}(°F - 32)$
  • $°F = \frac{9}{5}°C + 32$
• Absolute zero: $0\,\text{K} = -273.15\,°\text{C} = -459.67\,°\text{F}$
• Average: $\bar{x} = \frac{1}{n} \sum<em>{i=1}^n x</em>i$
• Standard deviation: $\sigma = \sqrt{\frac{1}{n-1} \sum<em>{i=1}^n (x</em>i - \bar{x})^2}$
• Percent error: $\%\text{ error} = \left|\frac{\text{measured} - \text{true}}{\text{true}}\right| \times 100\%$
• Scientific notation: mantissa × 10^exponent, e.g., $6.02 \times 10^{23}$
• Significant figures rules (summary): nonzero digits, zeros between significant digits, decimal point presence, leading/trailing zeros depending on decimals
• Dimensional analysis key rule: units cancel out; keep desired units until the end

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