Introduction to Chemistry: Study Guide for Chapters 1–4

Scientific Method and Foundations of Chemistry

Basic Steps in the Scientific Method

The scientific method is a systematic approach used by scientists to explore observations, answer questions, and solve problems.

• Observation: Gathering information through the senses or instruments.

• Hypothesis: A testable explanation or prediction based on observations.

• Experiment: Testing the hypothesis through controlled procedures.

• Analysis: Interpreting data and drawing conclusions.

• Conclusion: Determining whether the hypothesis is supported or refuted.

• Communication: Sharing results with the scientific community.

Example: Testing whether plants grow faster under blue light compared to red light.

Hypothesis, Theory, and Law

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

• Theory: A well-substantiated explanation of some aspect of the natural world, based on a body of evidence.

• Law: A statement that describes a consistently observed phenomenon, often expressed mathematically.

Example: The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

Matter and Its Classification

Definition of Matter

Matter is anything that has mass and occupies space.

Atoms, Molecules, Elements, and Compounds

• Atom: The smallest unit of an element that retains its chemical properties.

• Molecule: Two or more atoms chemically bonded together.

• Element: A pure substance made of only one kind of atom.

• Compound: A substance composed of two or more different elements chemically combined in fixed proportions.

Example: Water (H2O) is a compound made of hydrogen and oxygen atoms.

Classification of Matter

• Pure Substances: Elements and compounds with a fixed composition.

• Mixtures: Physical combinations of two or more substances; can be homogeneous (solutions) or heterogeneous.

Example: Air is a homogeneous mixture; sand and iron filings form a heterogeneous mixture.

States of Matter

• Solid: Definite shape and volume; particles vibrate in place; lowest kinetic energy.

• Liquid: Definite volume, no definite shape; particles move more freely.

• Gas: No definite shape or volume; particles move rapidly; highest kinetic energy.

Chemical vs. Physical Properties and Changes

• Physical Property: Can be observed without changing the substance (e.g., melting point, density).

• Chemical Property: Describes a substance's ability to undergo chemical changes (e.g., flammability).

• Physical Change: Does not alter the chemical composition (e.g., melting ice).

• Chemical Change: Produces new substances (e.g., rusting iron).

Chemistry and Measurements

Scientific Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• Format: $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Example: 0.00056 = $5.6 \times 10^{-4}$

Properties of Measurements and Significant Figures

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Uncertainty: The degree of doubt in a measurement, reflected in the last significant figure.

• Expressing to 3 Significant Figures: Round or write measurements so that only three digits are significant.

Example: 0.00456 has three significant figures.

Metric Units and Conversions

• Mass: Measured in grams (g), kilograms (kg), milligrams (mg).

• Volume: Measured in liters (L), milliliters (mL).

• Length: Measured in meters (m), centimeters (cm), millimeters (mm).

To convert between units, use conversion factors based on powers of ten.

Measurement of Mass, Volume, and Temperature

• Mass: Measured with a balance.

• Volume: Measured with a graduated cylinder, pipet, or buret.

• Temperature: Measured with a thermometer; units include Celsius (°C) and Kelvin (K).

Density

Density is the ratio of mass to volume.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Finding Volume or Mass: Rearranged formulas:

  • $\text{Volume} = \frac{\text{Mass}}{\text{Density}}$

  • $\text{Mass} = \text{Density} \times \text{Volume}$

Example: If a sample has a mass of 10 g and a volume of 2 mL, its density is $5 \text{ g/mL}$.

Temperature Conversions

• Celsius to Kelvin: $K = ^\circ C + 273.15$

• Kelvin to Celsius: $^\circ C = K - 273.15$

Energy and Chemical Changes

Definition of Energy and Law of Conservation of Energy

• Energy: The capacity to do work or produce heat.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed from one form to another.

Energy Changes in Chemical Reactions

• Chemical reactions often involve absorption or release of energy (as heat, light, etc.).

Calculating Energy Change (q)

• Formula: $q = m c \Delta t$

• Where:

  • $q$ = heat energy (Joules, J)

  • $m$ = mass (grams, g)

  • $c$ = specific heat capacity (J/g·°C)

  • $\Delta t$ = change in temperature (°C)

• Can solve for any variable if the others are known.

Example: How much energy is needed to raise 50 g of water by 10°C? (Assume $c = 4.18 \text{ J/g·°C}$)

$q = 50 \times 4.18 \times 10 = 2090 \text{ J}$

Atoms, Elements, and the Periodic Table

Names and Symbols of Elements

• Each element has a unique symbol (e.g., H for hydrogen, O for oxygen).

• Some elements exist as diatomic molecules (e.g., H2, N2, O2, F2, Cl2, Br2, I2).

Periodic Table: Families and Sections

• Groups (Families): Vertical columns; elements in a group have similar properties.

• Periods: Horizontal rows.

• Common Families:

  • Alkali metals (Group 1)

  • Alkaline earth metals (Group 2)

  • Halogens (Group 17)

  • Noble gases (Group 18)

Metals, Nonmetals, and Metalloids

Type	General Properties	Examples
Metals	Shiny, good conductors, malleable, ductile	Iron (Fe), Copper (Cu)
Nonmetals	Dull, poor conductors, brittle	Oxygen (O), Sulfur (S)
Metalloids	Properties intermediate between metals and nonmetals	Silicon (Si), Boron (B)

Compounds: Formulas and Atomic Composition

• A chemical formula shows the types and numbers of atoms in a compound (e.g., CO2).

• Subscripts indicate the number of each atom present.

Example: In H2SO4, there are 2 hydrogen, 1 sulfur, and 4 oxygen atoms.

Additional info: Some content (such as the full list of element names and symbols) is referenced as a handout and not included here. Students should refer to their course materials for the complete list.