Chapter 1 Notes: Introduction to Matter and Measurement

Introduction to Matter and Measurement

• Chapter 1 overview: Introduction to Matter and Measurement

• Core idea: Chemistry is the study of matter and how it changes; energy is central to chemical processes.

• Matter basics: made up of atoms; atoms form molecules, ions, and salts; energy governs chemical reactivity.

Why is Chemistry Important to You?

• Chemistry is integral to many careers and everyday scenarios:

  • Military — Antiterrorism

  • Biology — Bacteria in ice cream

  • Engineering — Hydration (water) effects on concrete

  • Medical doctor — Role of beta-amyloid in Alzheimer's

  • Nurse — Ibuprofen and quality of patient care

  • Dentist — Nickel leaching from orthodontia

  • Optician — Sodium fluorescein for eye pressure monitoring

  • Pharmacist — Drug/supplement interactions in patients

  • Farmer — Pesticide levels on fruit

  • Athlete — Hyponatremia (low sodium)

  • Food — Reference to Alton Brown and Good Eats

Chemistry the Central Science

• Chemistry is foundational for understanding many fields:

  • Biology, Physics (macroscopic properties), Medicine, Agriculture, Astronomy, Forensics

  • Homeland security

  • Business management? (considered, though less central)

How Many Chemical Compounds?

• Major databases and scale of chemical space:

  • ChemSpider ~ 1.29×10^8 compounds

  • PubChem (NIH) ~ 1.19×10^8 compounds

  • Chemical Abstracts registered compounds ~ 2.79×10^8

  • Chemical Space ~ 1.06×10^60 (contextual figure for scale)

What is Chemistry?

• Chemistry is the study of matter and its properties.

• Key questions:

  • What is matter made of?

  • What are the properties of matter?

• Matter components:

  • Individual atoms

  • Combinations of atoms → molecules, ions, salts

• Energy is very important to chemistry and helps dictate why chemicals react the way they do.

We Rely on Chemicals

• Top chemicals produced in the U.S. (annual production in billions of pounds) and end uses:

  • Sulfuric acid, H₂SO₄ → ~70; Fertilizers, chemical manufacturing

  • Ethylene, C₂H₄ → ~50; Plastics, antifreeze

  • Lime, CaO → ~45; Paper, cement, steel

  • Propylene, C₃H₆ → ~35; Plastics

  • Ammonia, NH₃ → ~18;

  • Chlorine, Cl₂ → ~21;

  • Phosphoric acid, H₃PO₄ → ~20

  • Sodium hydroxide, NaOH → ~16; Fertilizers; Bleaches, plastics, water purification

  • Aluminum (Al) production, soap → ~15

Atoms and Molecules

• The Chemical Elements:

  • Atoms are the fundamental building blocks of ordinary matter.

  • Molecules are combinations of atoms connected by covalent bonds.

  • Example: Liquid water is H₂O, containing two hydrogen atoms and one oxygen atom held by covalent bonds.

  • Online periodic table: https://pubchem.ncbi.nlm.nih.gov/periodic-table/

The Chemical Compounds

• Names of compounds can be extremely long and unwieldy; systematic naming is complex:

  • Example: Tetrahydrocannabinol (THC)

  • Very long IUPAC-like name: (6aR,10aR)-6,6,9-trimethyl-3-pentyl-6a,7,8,10a-tetrahydrobenzo[c]chromen-1

  • Alternative representations: CYQFCXCEBYINGO-IAGOWNOFSA-N, CCCCCC1=CC(=C2C3C=C(CCC3C(OC2=C1)(C)C)C)O

  • PubChem entry: https://pubchem.ncbi.nlm.nih.gov/compound/16078

The Scientific Approach to Knowledge

• How science advances:

  • Existing knowledge

  • Develop a hypothesis

  • Design experiment

  • Run experiment

  • Interpret results

  • Use results to form a new theory or adapt an existing one

• Science is not exact; there is always some error to account for.

Classifications of Matter

• Matter can be classified by composition and by uniformity/separability.

• Important concepts to know:

  • Pure Substance vs Mixture

  • Element vs Compound

  • Homogeneous vs Heterogeneous

  • Separable into simpler substances? (Yes/No)

  • Uniform throughout? (Yes/No)

What is Matter?

• Matter is anything that has mass and occupies space:

  • Mass measured on a balance

  • Occupies three dimensions

States of Matter

• Solid: definite shape and volume; high density; not easily compressed

• Liquid: definite volume; shape adapts to container; medium density

• Gas: no fixed shape or volume; low density; highly compressible

Composition of Matter

• Atom: smallest representative particle of an element

• Molecule: two or more atoms joined by covalent bonds

• Ions: electrically charged atoms or groups of atoms

• Examples:

  • Neon (Ne) as a free atom

  • Oxygen (O₂) as a molecule

  • Sodium ion Na⁺ and Chloride ion Cl⁻ as ions

Classification of Matter

• Pure Substance: fixed composition and distinct properties

• Mixture: combination of two or more substances

• Homogeneous mixture: uniform throughout (a solution)

• Heterogeneous mixture: non-uniform composition, visible phases

Heterogeneous/Homogeneous

• How to decide: look for multiple phases, solids, etc.

• Examples:

  • Seawater: heterogeneous (solution + particles)

  • Filtered seawater: homogeneous (solution only)

  • Tap water: homogeneous (solution)

  • Atmosphere: heterogeneous (gas + water)

  • Air (N₂, O₂, CO₂, etc.): homogeneous (gas solution)

  • Brass: homogeneous (alloy)

  • Concrete: heterogeneous (sand + gravel + mortar)

  • Jewelry gold: homogeneous (Au/Cu solution)

Classification Scheme (Visual)

• Matter

  • Variable composition? No → Pure Substance

  • Yes → Mixture

• Separable into simpler substances? No → Element

• Yes → Compound

• Homogeneous vs Heterogeneous determined by uniformity and phases

Separation of Mixtures

• Distillation: separate liquids with different boiling points

• Filtration: solid/liquid separation

• Chromatography: separation of organic mixtures

• General idea: separate components and collect pure fractions

Properties of Matter

• Physical properties and transformations

Intensive and Extensive Properties

• Intensive properties: independent of amount of sample (e.g., temperature, density)

• Extensive properties: depend on amount of sample (e.g., mass, volume)

Physical Transformations

• Physical transformations keep composition unchanged but change state or form, often with energy exchange:

  • Solid → Liquid: Melting

  • Liquid → Gas: Boiling

  • Gas → Liquid: Condensation

  • Liquid → Solid: Freezing

  • Gas → Solid: Sublimation (direct)

Physical Changes Diagram (conceptual)

• Water: liquid to gas phase change demonstrates energy exchange without changing chemical composition.

A Chemical Reaction

• Definition: Initial substances are converted into different substances via rearrangement of atoms/electrons

• We say a reaction occurs; involves breaking and forming bonds

Energy

• Kinetic Energy (KE): motion

  • Formula: KE = \frac{1}{2} m v^{2}

  • m = mass (g), v = velocity (m/s)

• Potential Energy (PE): stored energy (gravitational, electrical, chemical)

• Thermal Energy: related to temperature and microscopic motion

• Radiant Energy: light energy

• Note: Chemical energy relates to bonds within molecules

Units of Measurement

• SI system used for scientific measurements

• Key base units: length (m), mass (kg), time (s), temperature (K), amount of substance (mol), electric current (A), luminous intensity (cd)

Scientific Notation and Units

• Exponential notation to write large/small numbers:

  • Example: N_A = 6.02214076 \times 10^{23} \ \text{mol}^{-1}

  • “6.02214076 × 10^23” particles per mole; exact constant as of 2019

• Atomic mass example: m(^{12}\text{C per atom}) = 1.992 \times 10^{-23} \ \text{g}

  • Note: per-atom mass used in specific contexts

Systeme International (SI) Units

• See standard table:

  • Length: meter (m)

  • Mass: kilogram (kg)

  • Time: second (s)

  • Temperature: Kelvin (K)

  • Amount of substance: mole (mol)

  • Electric current: ampere (A)

  • Luminous intensity: candela (cd)

SI Prefixes

• Prefixes from Exa (E) to Atto (a):

  • exa (E) = 10^{18}

  • 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) = 10^{-6}

  • nano (n) = 10^{-9}

  • pico (p) = 10^{-12}

  • femto (f) = 10^{-15}

  • atto (a) = 10^{-18}

Converting Temperature

• Temperature scales:

  • Celsius to Fahrenheit: TC = \dfrac{TF - 32}{1.8}

  • Celsius to Kelvin: TK = TC + 273.16

• Note: The conversion factor 1.8 is the ratio between Fahrenheit degrees and Celsius degrees

Dimensional Analysis

• Important problem-solving tool: check units to ensure consistency

• Use conversion factors to cancel units

• Examples of conversion factors:

  • 1\ \text{km} = 1 \times 10^{3}\ \text{m}

  • 1\ \text{nm} = 1 \times 10^{-9}\ \text{m}

  • 1\ \text{inch} = 2.54\ \text{cm}

  • 1\text{ m} = 3.281\ \text{ft}

Dimensional Analysis (Unit Conversions)

• Writing the sequence of unit changes helps plan calculations

• Example data: density = 0.99 g/cm³; speed = 55 miles/hr

• Use these values to set up unit cancellations and conversions

Example: Calculation of Density

• Problem example: density of chocolate pudding; 1 cup of pudding weighs 0.2884 kg

• Key relationship: density, ρ, is mass per volume

• Formula: \rho = \frac{m}{V}

• Typical conversions: many units can express the same density (g/cm³, kg/m³, etc.)

• Note: density values can be tabulated in many units; choose a consistent set

Example: Density Conversions (Summary from table in notes)

• Common density units include g/cm³, g/mL, kg/L, kg/m³

• Example conversion: 1 g/mL = 1 g/cm³

• Specific pudding density examples shown in a lengthy table; core idea is to practice unit conversions to obtain consistent density values

Reliability of Measurements

• Measurements carry uncertainty; no measurement is perfectly exact

• Uncertainty affects significant digits and reported values

• Concept: Garbage in → garbage out

Uncertainty in Measurements

• Every data point has some uncertainty

• Report values with a correct number of significant digits to reflect precision

• Significant digits rule: number of significant digits equals number of certain digits plus one uncertain digit

• Example: burette reading 20.46 mL; balance reading 1.0785 g

Accuracy and Precision

• Accuracy: closeness to true value (often unknown)

• Precision: reproducibility of repeated measurements under the same conditions

Accuracy/Precision and Apparatus (Measurement Tools)

• Tools and typical uses include:

  • Stopcock-based devices to control fluid flow

  • Volumetric tools to deliver a specific volume

  • Graduated cylinders, syringes, burettes, pipettes, volumetric flasks

Counting Significant Digits

• Guidelines for identifying significant digits:

  • Nonzero digits are always significant

  • Leading zeros are not significant

  • Captive zeros (between significant digits) are significant

  • Trailing zeros are significant only if the number contains a decimal point

  • In scientific notation, all digits in the coefficient are significant

• Exact numbers have infinite significant digits (context determines exactness)

Rounding and Significant Digits

• Rounding rules:

  • If the immediate right digit is 0–4, round down

  • If it is 6–9, round up

  • If it is 5 with trailing zeros, round to the nearest even digit (banker’s rounding)

  • For 5 followed by non-zero digits, round up

• Examples:

  • 3.73176 → 3.7 (one decimal place)

  • 3.73176 → 3.73 (two decimal places)

  • 3.73176 → 3.732 (three decimal places)

  • 3.7500 → 3.8 (round to one decimal place, even rule)

  • 8.652 → 8.7

• Exact numbers concept recap: exact values have infinite decimals; many everyday constants are not exact except when defined by definition (e.g., counting numbers like 1 elephant is exact by definition)