Chemistry Chapter 1 Notes - Review

Chapter 1: Introduction to Matter, Energy, and Measurement

Chapter Objectives

1. Describe and classify matter and changes.

2. Communicate uncertainty in measurements.

3. Utilize dimensional analysis to convert units.

1.1 What is Chemistry

• Definition: Chemistry is the study of matter, its properties, and the changes it undergoes.

• Importance: Central to understanding many science-related fields.

• Examples: burning wood, iron rust.

1.2 Classifications of Matter

• Matter: Anything that has mass and takes up space (e.g., computer, coffee, oxygen gas).

  • Pure Substances: Fixed composition and distinct properties.

  • Cannot vary from sample to sample.

  • Examples of pure substances: elements and compounds.

  • Mixtures: Combinations of two or more different substances that are not chemically combined.

  • Examples: air, salad, and alloys.

States of Matter

• Three States: Solid, Liquid, and Gas.

  • Example: H2O can exist as ice (solid), liquid water (liquid), and water vapor (gas).

Classification of Matter: Pure Substances

• Elements: Substances that cannot be decomposed into simpler substances.

  • Characteristics: Made of unique kinds of atoms.

  • Examples: Copper (Cu), Oxygen (O2).

• Compounds: Substances that can be decomposed into simpler substances (made of atoms from two or more different elements).

  • Examples: Water (H2O), Carbon Dioxide (CO2).

Representing Elements

• Symbols: Chemists represent elements by symbols (one or two letters, first letter capitalized).

• Some Common Elements:

  • Carbon (C), Hydrogen (H), Nitrogen (N), Oxygen (O), Iron (Fe), etc.

1.3 Properties of Matter

• Physical Properties: Can be observed without changing the substance.

  • Examples: color, odor, density, and melting/boiling points.

• Chemical Properties: Observed when a substance changes into another.

  • Example: flammability.

Physical and Chemical Changes

• Physical Changes: Changes that do not alter the composition (e.g., state changes, temperature).

• Chemical Changes: Produces new substances (e.g., combustion, oxidation).

1.4 Energy --- Key Definitions

• Energy: Capacity to do work or transfer heat.

  • Kinetic Energy (KE): Energy of motion, dependent on mass (m) and velocity (v).

  • KE = rac{1}{2} mv^2

  • Potential Energy (PE): Stored energy, depends on position or chemical composition.

  • PE = mgh

1.5 Units of Measurement

• Importance: Chemistry involves quantitative measurements.

• Major Concepts:

  • Units of measurement, quantities, uncertainty, significant figures, dimensional analysis.

SI Units

• Base Units for fundamental physical quantities:

  • Length: Meter (m)

  • Mass: Kilogram (kg)

  • Temperature: Kelvin (K)

  • Time: Second (s)

  • Amount of substance: Mole (mol)

Derived Units: Volume and Density

• Volume Units:

  • 1 m^3 = 1000 L

  • Common metric units: Liter (L) and milliliter (mL).

  • 1 mL = 1 cm^3

• Density:

  • Density = \frac{mass}{volume} (commonly in g/mL or g/cm³).

  • Example substances: Water - 1.00 g/cm³, Gold - 19.32 g/cm³.

Measurement Uncertainty

• Precision vs. Accuracy:

  • Precision: Agreement of measurements with each other.

  • Accuracy: Agreement of measurements with the true value.

Significant Figures

• Definition: Digits that contribute to the precision of a measurement.

• Rules:

  1. All non-zero digits are significant.

  2. Zeros between significant digits are significant.

  3. Leading zeros are not significant.

  4. Trailing zeros are significant if there's a decimal.

Dimensional Analysis

• Description: Used to convert units based on known relationships.

• Key Steps:

  1. Identify the given and needed units.

  2. Use conversion factors to cancel units.

  3. Set up the calculation to solve for the needed unit.