Foundations of General Chemistry: Matter, Measurement, and Atomic Structure

The Scientific Method

Overview

The scientific method is a systematic approach used in scientific investigation to acquire new knowledge and validate existing knowledge. It involves making observations, forming hypotheses, conducting experiments, analyzing data, and drawing conclusions.

• Hypothesis: A tentative explanation or prediction that can be tested by experiments.

  • Falsifiable: A hypothesis must be structured so that it can be proven false by evidence.

• Experiments: Controlled procedures carried out to test the validity of a hypothesis.

• Data Analysis: The process of interpreting experimental results to determine if they support or refute the hypothesis.

• Drawing Conclusions: Summarizing findings and determining the next steps, such as revising the hypothesis or developing a theory.

Scientific Laws and Theories

Definitions

• Scientific Law: A concise statement that describes a fundamental relationship or regularity of nature, often expressed mathematically. Example: The Law of Conservation of Mass.

• Scientific Theory: A well-substantiated explanation of some aspect of the natural world that is based on a body of evidence and has stood up to repeated testing. Example: Atomic Theory.

States of Matter

Classification

Matter exists in three primary physical states, each with distinct properties:

• Solid: Has a definite shape and volume. Particles are closely packed and vibrate in place.

  • Crystalline: Particles are arranged in an orderly, repeating pattern (e.g., salt, diamond).

  • Amorphous: Particles lack a long-range order (e.g., glass, plastic).

• Liquid: Has a definite volume but takes the shape of its container. Particles are close but can move past one another.

• Gas: Has neither definite shape nor volume. Particles are far apart and move freely.

Types of Matter

Classification

• Mixture: A physical combination of two or more substances.

  • Homogeneous: Uniform composition throughout (e.g., saltwater).

  • heterogeneous: Non-uniform composition (e.g., salad, sand and iron filings).

• Pure Substance: Matter with a fixed composition and distinct properties. Includes elements and compounds.

Properties and Changes of Matter

Properties

• Physical Properties: Characteristics observed without changing the substance's identity. Examples: Color, melting point, density.

• Chemical Properties: Characteristics observed when a substance undergoes a chemical change. Examples: Flammability, reactivity with acid.

Changes

• Physical Change: Alters the state or appearance but not the chemical composition. Example: Melting ice, dissolving sugar in water.

• Chemical Change: Alters the chemical structure, resulting in new substances. Example: Rusting of iron, burning of wood.

Significant Figures

Rules and Applications

• Zeroes that matter:

  • Interior (captive) zeroes: Always significant (e.g., 205 has three significant figures).

  • Leading zeroes: Never significant (e.g., 0.0025 has two significant figures).

  • Trailing zeroes: Significant only if there is a decimal point (e.g., 2.300 has four significant figures).

• Scientific Notation: Used for very large or small numbers to clearly indicate significant figures (e.g., $3.00 \times 10^2$).

• Rounding: Round to the correct number of significant figures based on the operation performed.

Unit Conversions

Dimensional Analysis

• Cross-method (factor-label method): A systematic approach to converting units using conversion factors. Example: To convert 10 inches to centimeters, use $10 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 25.4 \text{ cm}$

Law of Definite and Multiple Proportions

Fundamental Laws of Chemistry

• Law of Definite Proportions: A chemical compound always contains the same elements in the same proportion by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

The Structure of the Atom

Basic Atomic Structure

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles orbiting the nucleus.

Isotopes

Understanding Isotopes

• Number of Protons: Equal to the atomic number (Z) of the element.

• Number of Neutrons: Calculated as $\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$

• Natural Abundance: The relative proportion of each isotope found in nature.

Reading the Periodic Table

Groups and Charges

• Groups: Vertical columns in the periodic table; elements in the same group have similar chemical properties.

• Determining Ionic Charges: Main group elements often form ions with predictable charges based on their group number (e.g., Group 1 forms +1 ions, Group 17 forms -1 ions).

Ions

Types and Naming

• Cations: Positively charged ions (e.g., Na+).

• Anions: Negatively charged ions (e.g., Cl-).

• Polyatomic Ions: Ions composed of more than one atom (e.g., SO42-).

• Determining Charge: Based on the loss or gain of electrons relative to the neutral atom.

• Naming: Cations use the element name (e.g., sodium ion), anions often end in "-ide" (e.g., chloride), and polyatomic ions have specific names (e.g., sulfate).

Formula Writing and Naming Compounds

Rules and Examples

• Formula Writing: Combine cations and anions in ratios that yield a neutral compound.

• Naming Compounds: Use systematic rules for ionic and molecular compounds.

  • KI: Potassium iodide

  • Sr(NO3)2: Strontium nitrate

  • Carbon tetrachloride: CCl4

  • Hydrogen sulfate: HSO4-

  • Copper(II) oxide: CuO

  • Co(NO3)2: Cobalt(II) nitrate

  • CaCO3: Calcium carbonate

In-Class Math Review

Sample Calculations

• 3.5 - 2.396 = 1.104

• 2.341 × 376 × 0.007 = 6.16

• (7826 + 23 - 5.2) × 7.5 = 58,627.5

• (5.95 × 3.7628) - 4.25 = 17.12

Additional info: Calculations above are rounded to the appropriate number of significant figures where possible.