Chapter 1: Chemistry in Our Lives

The Scientific Method

The scientific method is a systematic approach used in scientific investigation. It involves making observations, forming hypotheses, conducting experiments, and drawing conclusions.

• Observation: Gathering information using senses or instruments.

• Hypothesis: A proposed explanation based on observations.

• Experiment: Testing the hypothesis through controlled procedures.

• Conclusion: Analyzing results to accept or reject the hypothesis.

• Example: Observing that ice melts at room temperature, hypothesizing that heat causes melting, testing by changing temperature, and concluding based on results.

Chapter 2: Chemistry and Measurements

Scientific Notation and Expanded Notation

Scientific notation is used to express very large or very small numbers in a compact form. Expanded notation shows the full value.

• Converting to Scientific Notation: Move the decimal point to create a number between 1 and 10, then multiply by a power of ten.

• Converting to Expanded Notation: Multiply the coefficient by the power of ten.

• Example:

Comparing Numbers

Numbers in scientific or expanded notation can be compared by converting both to the same format.

• Example: is larger than

Metric System and Units

The metric system uses base units for measurement. Each measurement consists of a number and a unit.

• Base Units: Meter (length), gram (mass), liter (volume), second (time).

• Prefixes: Used to modify base units by powers of ten (e.g., milli-, centi-, kilo-).

• Example: 1 kilometer = 1000 meters

Measured vs. Exact Numbers

Measured numbers are obtained by measurement and have uncertainty. Exact numbers are counted or defined and have no uncertainty.

• Example: 12 eggs (exact), 2.5 cm (measured)

Significant Figures

Significant figures reflect the precision of a measured value.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros in a decimal are significant.

• Rounding: Round to the specified number of significant figures based on calculation rules.

• Example: 0.00450 has three significant figures.

Dimensional Analysis and Conversion Factors

Dimensional analysis uses conversion factors to solve measurement problems.

• Equality: 1 m = 100 cm

• Conversion Factors: or

• Example: Convert 250 cm to meters:

Density and Specific Gravity

Density is the mass per unit volume. Specific gravity compares the density of a substance to the density of water.

• Density Formula:

• Specific Gravity:

• Example: If a substance has a mass of 10 g and a volume of 2 mL, its density is .

Chapter 3: Matter and Energy

Classification of Matter

Matter can be classified as pure substances or mixtures.

• Pure Substance: Has a fixed composition (element or compound).

• Mixture: Contains two or more substances (homogeneous or heterogeneous).

• Example: Water (compound), air (homogeneous mixture), salad (heterogeneous mixture)

States of Matter

Solids, liquids, and gases have distinct properties.

• Solids: Definite shape and volume.

• Liquids: Definite volume, no definite shape.

• Gases: No definite shape or volume.

Chemical vs. Physical Properties and Changes

Physical properties can be observed without changing the substance. Chemical properties involve a change in composition.

• Physical Change: Melting, boiling, dissolving.

• Chemical Change: Burning, rusting.

Energy: Potential and Kinetic

Energy exists as potential (stored) or kinetic (motion).

• Potential Energy: Stored energy (e.g., chemical bonds).

• Kinetic Energy: Energy of motion (e.g., moving particles).

Changes of State

Substances change state through melting, freezing, condensation, vaporization, sublimation, and deposition.

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Condensation: Gas to liquid

• Vaporization: Liquid to gas

• Sublimation: Solid to gas

• Deposition: Gas to solid

Chapter 4: Atoms and Elements

Chemical Symbols and the Periodic Table

Elements are represented by chemical symbols. The periodic table organizes elements by periods (rows) and groups (columns).

• Symbol: One or two letters; first capitalized, second lowercase.

• Groups: Similar chemical properties.

• Example: Na (sodium), Cl (chlorine)

Classification of Elements

Elements are classified as metals, nonmetals, or metalloids. Main-group elements and transition metals are distinguished by their position.

• Metals: Left and center of periodic table.

• Nonmetals: Right side.

• Metalloids: Border between metals and nonmetals.

Atomic Structure

Atoms consist of protons, neutrons, and electrons.

• Protons: Charge +1, mass 1 amu.

• Neutrons: Charge 0, mass 1 amu.

• Electrons: Charge -1, negligible mass.

• Atomic Number: Number of protons.

• Mass Number: Protons + neutrons.

• Isotope: Atoms of same element with different numbers of neutrons.

Electron Arrangement and Periodic Trends

Electron arrangement determines chemical properties. Group number equals number of valence electrons for main-group elements. Atomic size and ionization energy show periodic trends.

• Atomic Size: Increases down a group, decreases across a period.

• Ionization Energy: Decreases down a group, increases across a period.

Chapter 5: Nuclear Chemistry

Types of Radiation

Common particles in nuclear chemistry include alpha, beta, and positron.

• Alpha Particle: or

• Beta Particle: or

• Positron: or

Nuclear Equations and Half-Life

Nuclear equations must be balanced. Half-life is the time for half of a radioactive sample to decay.

• Example:

• Half-Life Calculation:

Nuclear Fission and Fusion

Fission splits heavy nuclei; fusion combines light nuclei. Fission produces more radioactive waste than fusion.

• Fission: Used in nuclear reactors.

• Fusion: Powers the sun.

Chapter 6: Ionic and Molecular Compounds

Ionic and Covalent Bonds

Ionic bonds form between metals and nonmetals; covalent bonds form between nonmetals or nonmetals and metalloids.

• Ionic Bond: Transfer of electrons.

• Covalent Bond: Sharing of electrons.

Formation of Ions

Metals lose electrons to form cations; nonmetals gain electrons to form anions. Only electrons are gained or lost.

• Ion Charge: Number of protons minus number of electrons.

• Example: Sodium (Na) forms Na+ by losing one electron.

Naming and Writing Formulas

Ionic compounds are named based on the ions present. Roman numerals indicate cation charge when necessary. Polyatomic ions are recognized by their specific names and formulas.

• Example: FeCl2 is iron(II) chloride.

• Polyatomic Ion: SO42- (sulfate)

Molecular Compounds and Electronegativity

Molecular compounds contain covalent bonds. Electronegativity trends help classify bonds as polar or nonpolar.

• Polar Covalent: Unequal sharing of electrons.

• Nonpolar Covalent: Equal sharing of electrons.

Intermolecular Forces

Intermolecular forces include hydrogen bonding, dipole-dipole attractions, and dispersion forces.

• Hydrogen Bonding: Strongest

• Dispersion Forces: Weakest