Chapter 1 Matter and Measurements

Chapter 1 - Matter and Measurements

Section 1.1-1.2: Chemistry + The Scientific Method

• Objectives:

  • Explain the scientific method.

  • Differentiate between hypotheses, theories, and laws.

  • Examples of macroscopic, microscopic, and symbolic domains.

Chemistry: The Central Science

• Definition: Study of composition, properties, and interactions of matter.

• Focus: Understanding atomic/microscopic behavior to explain macroscopic phenomena.

The Scientific Method

• Foundation: Based on observation and experimentation; reproducibly verifies results.

1. Start with a hypothesis: Tentative explanation of observations.

   • Must be testable and falsifiable.

2. Experiments lead to laws: Summarize consistent observations.

3. Eventually leads to theories: Comprehensive explanations of natural behaviors.

The Domains of Chemistry

• Macroscopic Domain: Everyday life; observable size (e.g., food, raw materials).

• Microscopic Domain: Not directly visualized; requires microscopes (e.g., bacteria, viruses).

• Submicroscopic/Atomic Domain: Understood through experimentation (e.g., atoms, molecules).

• Symbolic Domain: Language of chemistry representing atomic substances.

Phases and Classification of Matter

Section Objectives

• Describe properties of solids, liquids, gases.

• Define atoms and molecules; classify matter as elements, compounds, or mixtures.

• Distinction between mass and weight; apply the Law of Conservation of Matter.

Describing Matter by its Phase

• Matter: Occupies space and has mass.

• States:

  • Solids: Fixed volume/shape; incompressible.

  • Liquids: Fixed volume; indefinite shape; incompressible.

  • Gases: Indefinite shape/volume; compressible.

Weight vs Mass

• Mass: Constant; how much matter is present.

• Weight: Force of gravity; variable.

• Law of Conservation of Matter: Matter cannot be created or destroyed, only converted.

Atoms and Molecules

• Atoms: Simplest form of matter; cannot be broken down.

• Molecules: Composed of 2+ chemically united atoms.

  • Atomic Elements: Single atoms.

  • Molecular Elements: Naturally found in pairs or larger groups.

Classifying Matter: Composition

• Pure Substances: One type of atom/molecule; constant composition; separable via chemical changes.

• Mixtures: Combination of 2+ types; variable composition; separable via physical changes.

  • Examples of Pure Substances: Hydrogen, water.

  • Examples of Mixtures: Salt water, air.

Closer Looks: Elements and Compounds

• Elements: Cannot be further simplified; examples include Iron,

• Compounds: Chemically bonded molecules; properties differ from free/uncombined state.

Closer Looks: Mixtures

• Heterogeneous Mixtures: Visibly distinguishable components; non-uniform.

• Homogeneous Mixtures (Solutions): Uniform composition; indistinguishable components.

Chemical and Physical Properties

Section Objectives

• Identify chemical/physical properties and changes.

• Differentiate between extensive and intensive properties.

Physical Properties and Changes

• Definition: Characteristics not associated with chemical composition.

• Examples: Density, color, melting/boiling points.

• Most Changes: Reversible (e.g., ice melting).

Chemical Properties and Changes

• Definition: Characteristics associated with a change in chemical composition.

• Examples: Flammability, reactivity.

• Most Changes: Not reversible (e.g., combustion).

Extensive vs Intensive Properties

• Extensive Properties: Depend on amount (e.g., mass, volume).

• Intensive Properties: Do not depend on amount (e.g., density).

Measurements

Section Objectives

• Explain measurement process and identify basic quantity components.

• Describe properties and units of measurements (length, mass, volume, etc.).

Measurements

• Scientific Notation: Used for large/small measurements.

  • Example: 298,000 kg = 2.98 x 10^5 kg.

• Units: SI Units for standard measurements (m, kg, s, K, mol).

SI Units: Prefixes

• Used to modify base units for convenience (e.g., kilo, centi, milli).

Derived SI Units

• Volume: cm³ (1 cm³ = 1 mL).

• Density: Intensive property, crucial for identifying substances.

Measurement, Uncertainty, Accuracy and Precision

Section Objectives

• Define accuracy and precision; distinguish between exact and uncertain numbers.

Accuracy and Precision

• Precision: Repeatability of measurements.

• Accuracy: Closeness to the true value.

Uncertainty in Measurements

• Measurements include uncertainty; values reported with last digit estimated.

• Counted/defined values have no uncertainty (e.g., 12 inches = 1 foot).

Significant Figures

• Rules:

  • Non-zero numbers are significant.

  • Interior zeroes are significant.

  • Trailing zeroes significant if decimal present.

  • Leading zeroes insignificant.

Dimensional Analysis

Section Objectives

• Explain dimensional analysis for calculations and unit conversions.

Dimensional Analysis/Factor Label Method

• Set up calculations to ensure units cancel.

Common Conversion Factors

• Length: 1 m = 1.0936 yd.

• Volume: 1 L = 1.0567 qt.

• Mass: 1 kg = 2.2046 lb.

Multi-Step Problems

• May require multiple calculations for unit conversion; do not round until fully completed.

Temperature Scales

• Common Scales: Fahrenheit, Celsius, Kelvin.

• Conversions: From one scale to another; formulate equations.